science(biology,physics,chemistry,agriculture) ,USA education: chemistry ,What Is Chemistry?A SIMPLE VIEW OF ATOMIC STRUCTURE, Atom and Electromagnetic Radiation,MOLECULAR STRUCTURES,Shape of Molecules,matter(gases,liquids,solids),Solutions,Chemical Reactivity,

define chemistry?
Chemistry is a branch of physical science, is the study of the composition, structure, properties and change of matter.
What Is Chemistry?
Chemistry is the study of matter and energy and the interactions between them. This is also the definition for physics, by the way. Chemistry and physics are specializations of physical science. Chemistry tends to focus on the properties of substances and the interactions between different types of matter, particularly reactions that involve electrons. Physics tends to focus more on the nuclear part of the atom, as well as the subatomic realm. Really, they are two sides of the same coin.
Chemistry as science
Under the influence of the new empirical methods propounded by Sir Francis Bacon and others, a group of chemists at Oxford, Robert Boyle,Robert Hooke and John Mayow began to reshape the old alchemical traditions into a scientific discipline. Boyle in particular is regarded as the founding father of chemistry due to his most important work, the classic chemistry text The Sceptical Chymist where the differentiation is made between the claims of alchemy and the empirical scientific discoveries of the new chemistry

Prior to his work, though, many important discoveries had been made, specifically relating to the nature of 'air' which was discovered to be composed of many different gases. The Scottish chemist Joseph Black (the first experimental chemist) and the Dutchman J. B. van Helmontdiscovered carbon dioxide, or what Black called 'fixed air' in 1754; Henry Cavendish discovered hydrogen and elucidated its properties andJoseph Priestley and, independently, Carl Wilhelm Scheele isolated pure oxygen.
English scientist John Dalton proposed the modern theory of atoms; that all substances are composed of indivisible 'atoms' of matter and that different atoms have varying atomic weights.
The development of the electrochemical theory of chemical combinations occurred in the early 19th century as the result of the work of two scientists in particular, J. J. Berzelius and Humphry Davy, made possible by the prior invention of the voltaic pile by Alessandro Volta. Davy discovered nine new elements including the alkali metals by extracting them from their oxides with electric current.
British William Prout first proposed ordering all the elements by their atomic weight as all atoms had a weight that was an exact multiple of the atomic weight of hydrogen. J. A. R. Newlands devised an early table of elements, which was then developed into the modern periodic table of elements by the German Julius Lothar Meyer and the Russian Dmitri Mendeleev in the 1860s. The inert gases, later called the noble gases were discovered by William Ramsay in collaboration with Lord Rayleigh at the end of the century, thereby filling in the basic structure of the table.
chemical structure:

At the turn of the twentieth century the theoretical underpinnings of chemistry were finally understood due to a series of remarkable discoveries that succeeded in probing and discovering the very nature of the internal structure of atoms. In 1897, J. J. Thomson of Cambridge Universitydiscovered the electron and soon after the French scientist Becquerel as well as the couple Pierre and Marie Curie investigated the phenomenon of radioactivity. In a series of pioneering scattering experiments Ernest Rutherford at the University of Manchester discovered the internal structure of the atom and the existence of the proton, classified and explained the different types of radioactivity and successfully transmuted the first element by bombarding nitrogen with alpha particles.
His work on atomic structure was improved on by his students, the Danish physicist Niels Bohr and Henry Moseley. The electronic theory ofchemical bonds and molecular orbitals was developed by the American scientists Linus Pauling and Gilbert N. Lewis.
The year 2011 was declared by the United Nations as the International Year of Chemistry. It was an initiative of the International Union of Pure and Applied Chemistry, and of the United Nations Educational, Scientific, and Cultural Organization and involves chemical societies, academics, and institutions worldwide and relied on individual initiatives to organize local and regional activities.
Element

Standard form of the periodic table of chemical elements. The colors represent different categories of elements
A chemical element is a pure substance which is composed of a single type of atom, characterized by its particular number of protons in the nuclei of its atoms, known as the atomic number and represented by the symbol Z. The mass number is the sum of the number of protons and neutrons in a nucleus. Although all the nuclei of all atoms belonging to one element will have the same atomic number, they may not necessarily have the same mass number; atoms of an element which have different mass numbers are known as isotopes. For example, all atoms with 6 protons in their nuclei are atoms of the chemical element carbon, but atoms of carbon may have mass numbers of 12 or 13
Atom

A diagram of an atom based on the Rutherford modelThe atom is the basic unit of chemistry. It consists of a dense core called the atomic nucleus surrounded by a space called the electron cloud. The nucleus is made up of positively charged protons and uncharged neutrons (together called nucleons), while the electron cloud consists of negatively-charged electrons which orbit the nucleus. In a neutral atom, the negatively-charged electrons balance out the positive charge of the protons. The nucleus is dense; the mass of a nucleon is 1,836 times that of an electron, yet the radius of an atom is about 10,000 times that of its nucleus.
The atom is also the smallest entity that can be envisaged to retain the chemical properties of the element, such as electronegativity, ionization potential, preferred oxidation state(s), coordination number, and preferred types of bonds to form (e.g., metallic, ionic, covalent)..
Compound

Carbon dioxide(CO₂), an example of a chemical compound
Main article: Chemical compoundA compound is a pure chemical substance composed of more than one element. The properties of a compound bear little similarity to those of its elements.^[44] The standard nomenclature of compounds is set by the International Union of Pure and Applied Chemistry (IUPAC). Organic compoundsare named according to the organic nomenclature system.^[45] Inorganic compounds are named according to the inorganic nomenclature system.^[46] In addition the Chemical Abstracts Service has devised a method to index chemical substances. In this scheme each chemical substance is identifiable by a number known as its CAS registry number.
Molecule

A ball-and-stick representation of the caffeine molecule (C₈H₁₀N₄O₂).
A molecule is the smallest indivisible portion of a pure chemical substance that has its unique set of chemical properties, that is, its potential to undergo a certain set of chemical reactions with other substances. However, this definition only works well for substances that are composed of molecules, which is not true of many substances (see below). Molecules are typically a set of atoms bound together by covalent bonds, such that the structure is electrically neutral and all valence electrons are paired with other electrons either in bonds or in lone pairs.
Thus, molecules exist as electrically neutral units, unlike ions. When this rule is broken, giving the "molecule" a charge, the result is sometimes named a molecular ion or a polyatomic ion. However, the discrete and separate nature of the molecular concept usually requires that molecular ions be present only in well-separated form, such as a directed beam in a vacuum in a mass spectrometer. Charged polyatomic collections residing in solids (for example, common sulfate or nitrate ions) are generally not considered "molecules" in chemistry.
Main Branches of Chemistry
Although many would say that there are FIVE main branches of chemistry, namely: Physical, Analytical, Biochemistry, Organic and Inorganic chemistry many would argue that the science of chemistry actually links out to other branches or sub-branches that include Materials Chemistry, Theoretical Chemistry, Macromolecular (Polymer) Chemistry, Nuclear Chemistry, Metallurgy, Forensic Chemistry, Medicinal Chemistry and more.
It is important to note that often sub-branches fall under one or more of the main branches of chemistry.
Let’s start by taking a look at the 5 main branches of chemistry and then delve deeper into chemistry’s many sub-branches:
Analytical chemistry is the study involving how we analyze the chemical components of samples. How much caffeine is really in a cup of coffee? Are there drugs found in athlete’s urine samples? What is the pH level of my swimming pool? Examples of areas using analytical chemistry include forensic science, environmental science, and drug testing.
Analytical chemistry is divided into two main branches: qualitative and quantitative analysis.
Qualitative analysis employs methods/measurements to help determine the components of substances. Quantitative analysis on the other hand, helps to identify how much of each component is present in a substance.
Both types of analysis can be used to provide important information about an unidentified sample and help to identify what the sample is.
Biochemistry
The study of life or more aptly put, of chemical processes in living organisms. Biochemists research includes cancer and stem cell biology, infectious disease as well as membrane and structural biology and spans molecular biology, genetics, mechanistic biochemistry, genomics, evolution and systems biology.
Biochemistry, according to many scientists can also be explained as a discipline in which biological phenomena are examined in chemical terms. Examples are digestion and cellular respiration.
For this reason biochemistry is also known as Chemical Biology or Biological Chemistry.
Under the main umbrella of biochemistry many new sub-branches have emerged that modern chemists may specialize in solely. Some of these disciplines include:

Enzymology (study of enzymes)

Endocrinology (study of hormones)

Clinical Biochemistry (study of diseases)

Molecular Biochemistry (Study of Biomolecules and their functions).

There are also others like Pharmacological Biochemistry, Agricultural Biochemistry and more.
Click the informative links below to learn more about biochemistry:
Chemists in this field focus on elements and compounds other than carbon or hydrocarbons. Simply put, inorganic chemistry covers all materials that are not organic and are termed as non-living substances – those compounds that do not contain a carbon hydrogen (C-H) bond.
Compounds studied by inorganic chemists include crystal structures, minerals, metals, catalysts, and most elements on the periodic table. An example is the strength of a power beam used to carry a specific weight or investigating how gold is formed in the earth.
Branches of inorganic chemistry include:

Bioinorganic chemistry (study of role of metals in biology)

Coordination chemistry (study of coordination compounds and interactions of ligands)

Geochemistry (study of the earth’s chemical composition, rocks, minerals & atmosphere)

Inorganic technology (synthesizing new inorganic compounds)

Nuclear chemistry (study of radioactive substances)

Organometallic chemistry (study of chemicals that contain bonds between a metal and carbon – overlaps into organic chemistry)

Solid-state chemistry/materials chemistry (study of the forming, structure, and characteristics of solid phase materials)

Synthetic inorganic chemistry (study of synthesizing chemicals)

Industrial inorganic chemistry (study of materials used in manufacturing. E.g.: fertilizers)

Organic chemistry
The study of carbon compounds such as fuels, plastics, food additives, and drugs. An opposite of inorganic chemistry that focuses on non-living matter and non-carbon based substances, organic chemistry deals with the study of carbon and the chemicals in living organisms. An example is the process of photosynthesis in a leaf because there is a change in the chemical composition of the living plant.
Organic chemists are often the ones who devise experimental methods to isolate or synthesize new materials, or to study their properties, and usually work and research in a lab. Some examples on the work they do include formulating a conditioner that keeps hair softer, developing a better drug for headaches and creating a non-toxic home cleaning product.
The branches of organic chemistry involve many different disciplines including the study of ketones, aldehydes, hydrocarbons (alkenes, alkanes, alkynes) and alcohols.

Stereochemistry (study of the 3-dimensional structure of molecules)

Medicinal chemistry (deals with designing, developing and synthesizing pharmaceutical drugs)

Organometallic chemistry (study of chemicals that contain bonds between a carbon and a metal)

Physical organic chemistry (study of structure and reactivity in organic molecules)

Polymer chemistry (study of the composition and creation of polymer molecules)

The study of the physical properties of molecules, and their relation to the ways in which molecules and atoms are put together. Physical chemistry deals with the principles and methodologies of both chemistry and physics and is the study of how chemical structure impacts physical properties of a substance. An example is baking brownies, as you’re mixing materials and using heat and energy to get the final product.
Physical chemists would typically study the rate of a chemical reaction, the interaction of molecules with radiation, and the calculation of structures and properties.
Sub-branches of physical chemistry include:

Electrochemistry (study of the interaction of atoms, molecules, ions and electric current)

Photochemistry (study of the chemical effects of light; photochemical reactions)

Surface chemistry (study of chemical reactions at interfaces)

Chemical Kinetics (study of rates of chemical reactions)

Thermodynamics/Thermochemistry (study of how heat relates to chemical change)

Quantum Mechanics/Quantum Chemistry (study of quantum mechanics and how it relates to chemical phenomena.
The periodic table of elements is a chart that outlines all the basic elements of chemistry that make up our world according to their atomic numbers, the number of electrons each element has, and their predominant chemical properties.
Each element is lined up from low to high atomic number, which simply refers to the number of protons it has. Most periodic table of elements charts are laid out in this fashion: A tabular grid of 18 by 7 that houses all of the major elements over another two rows of elements below it.
The table can also be broken down into 4 distinct parts or blocks: the s-block on the left, the p-block on the right, the d-block towards the middle and the f-block at the bottom.
The table rows are referred to as "periods" and the columns (s, p and d blacks) are called "groups." Some groups also have specific names such as the noble gases, or the halogens.
The name "periodic" table suggests that the table itself is open to being updated on a periodic basis, so it's not only used to uncover how each of the elements relate to one another but also to discover the characteristics of new elements or yet to be found or synthesized elements.
Therefore, the periodic table proves to be an important guide and resource when it comes to showcasing all the basic elements and studying chemical tendencies, and is commonly used not only in the science of chemistry but other fields of science as well.
Although other forms of the periodic table have been known to exist, Dmitri Mendeleev is typically recognized as the pioneer for publishing the first periodic table of elements in 1869. He designed the table to show similarities in the properties of the elements that were known back in the day. He also forecasted the properties of undiscovered elements back then and marked their place on the table, and in fact most of his claims and estimations proved true when the elements were discovered as time passed. Since the 1800's the periodic table has grown and improved with new elements being found and new theories explaining the way chemicals behave.
Elements from atomic number 1 to 118, hydrogen to ununoctium, have either been discovered or created. From all of these, all the elements you see up to californium occur naturally. Others have been created in labs. Chemists continue their pursuit to synthesize new elements way beyond ununoctium, but the presence of these synthesized chemicals having their place on the periodic table is still a question of continued disagreement and debate. Synthetic versions of elements that naturally occur in the earth have also been produced in chemical laboratories.
Periodic Table of Elements

The above image of an eighteen column period table of elements structure is now the most commonly and broadly used format because it’s been so popular and widely accepted.
Also called the “long form” periodic table layout, it differs from Mendeleev’s short form design by removing the groups three to twelve and inserting them instead into the other major groups. The long form layout actually includes the actinides and the lanthanides in its structure instead of placing them below and away from the main table body. This wider layout table also adds two more periods, periods 8 and 9, and also incorporates the superactinides.
Periodic Table of Elements

One thing to keep in mind is that the periodic table only documents chemical elements. It does not account for subatomic particles, or elements combined together such as mixtures and compounds. Each element's atomic number depends on how many protons it has in its nucleus. Isotopes are two or more variants of the same chemical element. They contain the same number of protons but carry a different number of neutrons in their nucleus. As the number of protons stays equal, the atomic number does not change. For example, carbon has three isotopes that occur naturally. Most have 6 protons but the number of neutrons can vary between 6 and 8. However, isotopes are always shown together under the element they belong to in the periodic table. Elements with no stable isotopes carry the atomic mass of the most stable isotope of that element, where the mass is displayed in parentheses.
Let’s take a closer look at the arrangement of elements in the table. As mentioned before, all the elements are arranged according to their rising atomic number, or their increasing number of protons. A new period or row begins when an electron shell gets its first electron. Groups or columns are arranged according to the number of electrons the atom carries. Also, elements that display comparable chemical characteristics usually also fall within the same column, and in the f and d blocks elements that lie in the same row to some extent also display similar characteristics. Therefore if you know the properties of a particular element it is fairly simple to figure out the properties of other elements that surround it in the table.
Here are some more facts about the periodic table:
According to the most updated version in 2012, the periodic table is said to have 118 elements. Of these 114 are official and have been named and documented by the International Union of Pure and Applied Chemistry (IUPAC).
Ninety-eight elements are naturally occurring out of which eighty-four are known as primordial elements. The remaining fourteen occur in decay chains of these primordial elements.
Even though elements like livermorium, flerovium and those listed from einsteinium to copernicium do not naturally occur in the earth and have been synthesized, they are still recognized by the IUPAC.
Other elements like 113, 115, 117 and 118 are known to be supposedly formulated in labs but the information has not yet been validated. Therefore these elements are only recognized by their element name, depending on their atomic number. As of the year 2012, there have been no reports of any element being synthesized keeping the count as of today at 118.

A SIMPLE VIEW OF ATOMIC STRUCTURE

This page revises the simple ideas about atomic structure that you will have come across in an introductory chemistry course (for example, GCSE). You need to be confident about this before you go on to the more difficult ideas about the atom which under-pin A'level chemistry.

The sub-atomic particlesProtons, neutrons and electrons.

	relative mass	relative charge
proton	1	+1
neutron	1	0
electron	1/1836	-1

Beyond A'level: Protons and neutrons don't in fact have exactly the same mass - neither of them has a mass of exactly 1 on the carbon-12 scale (the scale on which the relative masses of atoms are measured). On the carbon-12 scale, a proton has a mass of 1.0073, and a neutron a mass of 1.0087.

The behaviour of protons, neutrons and electrons in electric fields
What happens if a beam of each of these particles is passed between two electrically charged plates - one positive and one negative? Opposites will attract.
Protons are positively charged and so would be deflected on a curving path towards the negative plate.
Electrons are negatively charged and so would be deflected on a curving path towards the positive plate.
Neutrons don't have a charge, and so would continue on in a straight line.

Exactly what happens depends on whether the beams of particles enter the electric field with the various particles having the same speeds or the same energies
If the particles have the same energy
If beams of the three sorts of particles, all with the same energy, are passed between two electrically charged plates:

Protons are deflected on a curved path towards the negative plate.

Electrons are deflected on a curved path towards the positive plate.
The amount of deflection is exactly the same in the electron beam as the proton beam if the energies are the same - but, of course, it is in the opposite direction.

Neutrons continue in a straight line.

If the electric field was strong enough, then the electron and proton beams might curve enough to hit their respective plates.

If the particles have the same speeds
If beams of the three sorts of particles, all with the same speed, are passed between two electrically charged plates:

Protons are deflected on a curved path towards the negative plate.
Electrons are deflected on a curved path towards the positive plate.
If the electrons and protons are travelling with the same speed, then the lighter electrons are deflected far more strongly than the heavier protons.
Neutrons continue in a straight line.

Note: This is potentially very confusing! Most chemistry sources that talk about this give either one or the other of these two diagrams without any comment at all - they don't specifically say that they are using constant energy or constant speed beams. But it matters!
If this is on your syllabus, it is important that you should know which version your examiners are going to expect, and they probably won't tell you in the syllabus. You should look in detail at past questions, mark schemes and examiner's reports which you can get from your examiners if you are doing a UK-based syllabus. Information about how to do this is on the syllabuses page.If in doubt, I suggest you use the second (constant speed) version. This actually produces more useful information about both masses and charges than the constant energy version.

The nucleusThe nucleus is at the centre of the atom and contains the protons and neutrons. Protons and neutrons are collectively known as nucleons.
Virtually all the mass of the atom is concentrated in the nucleus, because the electrons weigh so little.

Working out the numbers of protons and neutrons

No of protons = ATOMIC NUMBER of the atom
The atomic number is also given the more descriptive name of proton number.

No of protons + no of neutrons = MASS NUMBER of the atom
The mass number is also called the nucleon number.

This information can be given simply in the form:

How many protons and neutrons has this atom got?
The atomic number counts the number of protons (9); the mass number counts protons + neutrons (19). If there are 9 protons, there must be 10 neutrons for the total to add up to 19.

The atomic number is tied to the position of the element in the Periodic Table and therefore the number of protons defines what sort of element you are talking about. So if an atom has 8 protons (atomic number = 8), it must be oxygen. If an atom has 12 protons (atomic number = 12), it must be magnesium.
Similarly, every chlorine atom (atomic number = 17) has 17 protons; every uranium atom (atomic number = 92) has 92 protons.

Isotopes
The number of neutrons in an atom can vary within small limits. For example, there are three kinds of carbon atom ¹²C, ¹³C and ¹⁴C. They all have the same number of protons, but the number of neutrons varies.

	protons	neutrons	mass number
carbon-12	6	6	12
carbon-13	6	7	13
carbon-14	6	8	14

These different atoms of carbon are calledisotopes. The fact that they have varying numbers of neutrons makes no difference whatsoever to the chemical reactions of the carbon.
Isotopes are atoms which have the same atomic number but different mass numbers. They have the same number of protons but different numbers of neutrons.

The electronsWorking out the number of electrons
Atoms are electrically neutral, and the positiveness of the protons is balanced by the negativeness of the electrons. It follows that in a neutral atom:

no of electrons = no of protons
So, if an oxygen atom (atomic number = 8) has 8 protons, it must also have 8 electrons; if a chlorine atom (atomic number = 17) has 17 protons, it must also have 17 electrons.
The arrangement of the electrons
The electrons are found at considerable distances from the nucleus in a series of levels called energy levels. Each energy level can only hold a certain number of electrons. The first level (nearest the nucleus) will only hold 2 electrons, the second holds 8, and the third also seems to be full when it has 8 electrons. At GCSE you stop there because the pattern gets more complicated after that.
These levels can be thought of as getting progressively further from the nucleus. Electrons will always go into the lowest possible energy level (nearest the nucleus) - provided there is space.
To work out the electronic arrangement of an atom

Look up the atomic number in the Periodic Table - making sure that you choose the right number if two numbers are given. The atomic number will always be the smaller one.

This tells you the number of protons, and hence the number of electrons.

Arrange the electrons in levels, always filling up an inner level before you go to an outer one.

e.g. to find the electronic arrangement in chlorine

The Periodic Table gives you the atomic number of 17.

Therefore there are 17 protons and 17 electrons.

The arrangement of the electrons will be 2, 8, 7 (i.e. 2 in the first level, 8 in the second, and 7 in the third).

The electronic arrangements of the first 20 elements

After this the pattern alters as you enter the transition series in the Periodic Table.
Two important generalisations
If you look at the patterns in this table:

The number of electrons in the outer level is the same as the group number. (Except with helium which has only 2 electrons. The noble gases are also usually called group 0 - not group 8.) This pattern extends throughout the Periodic Table for the main groups (i.e. not including the transition elements).
So if you know that barium is in group 2, it has 2 electrons in its outer level; iodine (group 7) has 7 electrons in its outer level; lead (group 4) has 4 electrons in its outer level.
Noble gases have full outer levels. This generalisation will need modifying for A'level purposes.

Dots-and-crosses diagrams
In any introductory chemistry course you will have come across the electronic structures of hydrogen and carbon, for example, drawn as:

Note: There are many places where you could still make use of this model of the atom at A'level. It is, however, a simplification and can be misleading. It gives the impression that the electrons are circling the nucleus in orbits like planets around the sun. As you will find when you look at the A'level view of the atom, it is impossible to know exactly how they are actually moving.

The circles show energy levels - representing increasing distances from the nucleus. You could straighten the circles out and draw the electronic structure as a simple energy diagram.
Carbon, for example, would look like this:

Thinking of the arrangement of the electrons in this way makes a useful bridge to the A'level view.

The Atom and Electromagnetic Radiation:
Fundamental Subatomic Particles

Particle	Symbol	Charge	Mass
electron	e^-	-1	0.0005486 amu
proton	p⁺	+1	1.007276 amu
neutron	n^o	0	1.008665 amu

The number of protons, neutrons, and electrons in an atom can be determined from a set of simple rules.

The number of protons in the nucleus of the atom is equal to the atomic number (Z).

The number of electrons in a neutral atom is equal to the number of protons.

The mass number of the atom (M) is equal to the sum of the number of protons and neutrons in the nucleus.

The number of neutrons is equal to the difference between the mass number of the atom (M) and the atomic number (Z).

Examples: Let's determine the number of protons, neutrons, and electrons in the following isotopes.

¹²C

¹³C

¹⁴C

¹⁴N

The different isotopes of an element are identified by writing the mass number of the atom in the upper left corner of the symbol for the element. ¹²C, ¹³C, and ¹⁴C are isotopes of carbon (Z = 6) and therefore contain six protons. If the atoms are neutral, they also must contain six electrons. The only difference between these isotopes is the number of neutrons in the nucleus.
¹²C: 6 electrons, 6 protons, and 6 neutrons
¹³C: 6 electrons, 6 protons, and 7 neutrons
¹⁴C: 6 electrons, 6 protons, and 8 neutrons
Electromagnetic Radiation
Much of what is known about the structure of the electrons in an atom has been obtained by studying the interaction between matter and different forms of electromagnetic radiation. Electromagnetic radiation has some of the properties of both a particle and a wave.
Particles have a definite mass and they occupy space. Waves have no mass and yet they carry energy as they travel through space. In addition to their ability to carry energy, waves have four other characteristic properties: speed, frequency, wavelength, and amplitude. The frequency (v) is the number of waves (or cycles) per unit of time. The frequency of a wave is reported in units of cycles per second (s^-1) or hertz (Hz).
The idealized drawing of a wave in the figure below illustrates the definitions of amplitude and wavelength. The wavelength (l) is the smallest distance between repeating points on the wave. The amplitude of the wave is the distance between the highest (or lowest) point on the wave and the center of gravity of the wave.
Diagram

If we measure the frequency (v) of a wave in cycles per second and the wavelength (l) in meters, the product of these two numbers has the units of meters per second. The product of the frequency (v) times the wavelength (l) of a wave is therefore the speed (s) at which the wave travels through space.
vl = s
Electromagnetic Radiation
Much of what is known about the structure of the electrons in an atom has been obtained by studying the interaction between matter and different forms of electromagnetic radiation. Electromagnetic radiation has some of the properties of both a particle and a wave.
Particles have a definite mass and they occupy space. Waves have no mass and yet they carry energy as they travel through space. In addition to their ability to carry energy, waves have four other characteristic properties: speed, frequency, wavelength, and amplitude. The frequency (v) is the number of waves (or cycles) per unit of time. The frequency of a wave is reported in units of cycles per second (s^-1) or hertz (Hz).
The idealized drawing of a wave in the figure below illustrates the definitions of amplitude and wavelength. The wavelength (l) is the smallest distance between repeating points on the wave. The amplitude of the wave is the distance between the highest (or lowest) point on the wave and the center of gravity of the wave.
Diagram

If we measure the frequency (v) of a wave in cycles per second and the wavelength (l) in meters, the product of these two numbers has the units of meters per second. The product of the frequency (v) times the wavelength (l) of a wave is therefore the speed (s) at which the wave travels through space.
vl = s.
Light and Other Forms of Electromagnetic Radiation
Light is a wave with both electric and magnetic components. It is therefore a form of electromagnetic radiation.
Visible light contains the narrow band of frequencies and wavelengths in the portion of the electro-magnetic spectrum that our eyes can detect. It includes radiation with wavelengths between about 400 nm (violet) and 700 nm (red). Because it is a wave, light is bent when it enters a glass prism. When white light is focused on a prism, the light rays of different wavelengths are bent by differing amounts and the light is transformed into a spectrum of colors. Starting from the side of the spectrum where the light is bent by the smallest angle, the colors are red, orange, yellow, green, blue, and violet.
As we can see from the following diagram, the energy carried by light increases as we go from red to blue across the visible spectrum.
Diagram

Because the wavelength of electromagnetic radiation can be as long as 40 m or as short as 10^-5 nm, the visible spectrum is only a small portion of the total range of electromagnetic radiation.
Diagram

The electromagnetic spectrum includes radio and TV waves, microwaves, infrared, visible light, ultraviolet, x-rays, g-rays, and cosmic rays, as shown in the figure above. These different forms of radiation all travel at the speed of light (c). They differ, however, in their frequencies and wavelengths. The product of the frequency times the wavelength of electromagnetic radiation is always equal to the speed of light.
vl = c
As a result, electromagnetic radiation that has a long wavelength has a low frequency, and radiation with a high frequency has a short wavelength.
ATOM
Matter has mass and takes up space. Atoms are basic building blocks of matter, and cannot be chemically subdivided by ordinary means.
The word atom is derived from the Greek word atom which means indivisible. The Greeks concluded that matter could be broken down into particles to small to be seen. These particles were called atoms
Atoms are composed of three type of particles: protons, neutrons, and electron. Protons and neutrons are responsible for most of the atomic mass e.g in a 150 person 149 lbs, 15 oz are protons and neutrons while only 1 oz. is electrons. The mass of an electron is very small (9.108 X 10^-28grams).
Both the protons and neutrons reside in the nucleus. Protons have a postive (+) charge, neutrons have no charge --they are neutral. Electrons reside in orbitals around the nucleus. They have a negative charge (-).
It is the number of protons that determines the atomic number, e.g., H = 1. The number of protons in an element is constant (e.g., H=1, Ur=92) but neutron number may vary, so mass number (protons + neutrons) may vary.
The same element may contain varying numbers of neutrons; these forms of an element are called isotopes. The chemical properties of isotopes are the same, although the physical properties of some isotopes may be different. Some isotopes are radioactive-meaning they "radiate" energy as they decay to a more stable form, perhaps another element half-life: time required for half of the atoms of an element to decay into stable form. Another example is oxygen, with atomic number of 8 can have 8, 9, or 10 neutrons.

What are elements?
All matter is made up of elements which are fundamental substances which cannot be broken down by chemical means. There are 92 elements that occur naturally. The elements hydrogen, carbon, nitrogen and oxygen are the elements that make up most living organisms. Some other elements found in living organisms are: magnesium, calcium, phosphorus, sodium, potassium.
By the late 1800's many elements had already been discovered. The scientist Dmitri Mendeleev, a Russian chemist, proposed an arrangement of know elements based on their atomic mass. The modern arrangement of the elements is known as the Periodic Table of Elements and is arranged according to the atomic number of elements. Here is an Interactive Table of Elements where you can learn more about each of the elements.

What makes each element unique?
Every atom would like to have an electron configuration like a noble gases. In noble gases the outer electron shell is complete. This makes the element chemically inert. Helium is an example of a noble (inert) gas. It is not present in organisms because it is not chemically reactive.

Historical Models of the atom BOHR MODEL
Bohr model shows electrons circling the nucleus at different levels or orbitals much like planets circle the sun. Electrons move from one energy state to another but can only exist aft defineite energy levels. The energy absorbed or released when electrons change states is in the form of electromagnetic radiation.
THE WAVE MODEL AND QUANTUM THEORY
The Bohr model was only able to explain the very simplest atoms, like hydrogen. Today's modern day theory is based on mathematics and the properties of waves. The wave model forms the basis for the Quantum Theory
. This theory gives the probability of locating electrons in a particular location, unlike assuming electrons orbit the nucleus as in the Bohr model.

How are electrons organized around the nucleus? All atoms would like to attain electron configurations like noble gases. That is, have completed outer shells. Atoms can form stable electron configurations like noble gases
by:

losing electrons

sharing electrons

gaining electrons.

For a stable configuration each atom must fill its outer energy level. In the case of noble gases that means eight electrons in the last shell (with the exception of He which has two electrons).
Atoms that have 1, 2 or 3 electrons in their outer levels will tend to lose them in interactions with atoms that have 5, 6 or 7 electrons in their outer levels. Atoms that have 5, 6 or 7 electrons in their outer levels will tend to gain electrons from atoms with 1, 2 or 3 electrons in their outer levels. Atoms that have 4 electrons in the outer most energy level will tend neither to totally lose nor totally gain electrons during interactions.
This Periodic Table of Elements will show you the electron configuration for any element you click on.

Visualizing Atomic Orbitals The atomic orbitals of the hydrogen atom can be visualized as a cloud around the nucleus. The orbital represents a probability of finding the electron at a particular location. Darker regions signify a greater probability. Shown below are the 1s (lowest orbital and the 2s orbital.

2s Atomic orbitals do not always have the shape of a sphere. Higher orbitals have very unusual shapes.

2px

3px

These orbitals were prepared by Dr. Yue-Ling Wong from the University of Florida for more images click here.

Remember molecular orbitals are 3-Dimensional 3D models of atomic orbitals

MOLECULAR STRUCTURES

This page describes how the physical properties of substances having molecular structures varies with the type of intermolecular attractions - hydrogen bonding or van der Waals forces.

Important! There's not much point in reading this page unless you are reasonably happy about the origin of hydrogen bondingand van der Waals forces. Follow these links first if you aren't sure about these.

The physical properties of molecular substancesMolecules are made of fixed numbers of atoms joined together by covalent bonds, and can range from the very small (even down to single atoms, as in the noble gases) to the very large (as in polymers, proteins or even DNA).
The covalent bonds holding the molecules together are very strong, but these are largely irrelevant to the physical properties of the substance. Physical properties are governed by the intermolecular forces - forces attracting one molecule to its neighbours - van der Waals attractions or hydrogen bonds.
Melting and boiling points
Molecular substances tend to be gases, liquids or low melting point solids, because the intermolecular forces of attraction are comparatively weak. You don't have to break any covalent bonds in order to melt or boil a molecular substance.

Note: This is really important! You can make yourself look extremely stupid if you imply in an exam that boiling water, for example, splits it into hydrogen and oxygen by breaking covalent bonds. Exactly the same water molecules are present in ice, water and steam.

The size of the melting or boiling point will depend on the strength of the intermolecular forces. The presence of hydrogen bonding will lift the melting and boiling points. The larger the molecule the more van der Waals attractions are possible - and those will also need more energy to break.

Solubility in water
Most molecular substances are insoluble (or only very sparingly soluble) in water. Those which do dissolve often react with the water, or else are capable of forming hydrogen bonds with the water.
Why doesn't methane, CH₄, dissolve in water?
The methane itself isn't the problem. Methane is a gas, and so its molecules are already separate - the water doesn't need to pull them apart from one another.
The problem is the hydrogen bonds between the water molecules. If methane were to dissolve, it would have to force its way between water molecules and so break hydrogen bonds. That costs a reasonable amount of energy.
The only attractions possible between methane and water molecules are the much weaker van der Waals forces - and not much energy is released when these are set up. It simply isn't energetically profitable for the methane and water to mix.
Why does ammonia, NH₃, dissolve in water?
Ammonia has the ability to form hydrogen bonds. When the hydrogen bonds between water molecules are broken, they can be replaced by equivalent bonds between water and ammonia molecules.
Some of the ammonia also reacts with the water to produce ammonium ions and hydroxide ions.

The reversible arrows show that the reaction doesn't go to completion. At any one time only about 1% of the ammonia has actually reacted to form ammonium ions. The solubility of ammonia is mainly due to the hydrogen bonding and not the reaction.
Other common substances which are freely soluble in water because they can hydrogen bond with water molecules include ethanol (alcohol) and sucrose (sugar).

Solubility in organic solvents
Molecular substances are often soluble in organic solvents - which are themselves molecular. Both the solute (the substance which is dissolving) and the solvent are likely to have molecules attracted to each other by van der Waals forces. Although these attractions will be disrupted when they mix, they are replaced by similar ones between the two different sorts of molecules.

Electrical conductivity
Molecular substances won't conduct electricity. Even in cases where electrons may be delocalised within a particular molecule, there isn't sufficient contact between the molecules to allow the electrons to move through the whole solid or liquid.

Some individual examplesIodine, I₂
Iodine is a dark grey crystalline solid with a purple vapour. M.Pt: 114°C. B.Pt: 184°C. It is very, very slightly soluble in water, but dissolves freely in organic solvents.

Iodine is therefore a low melting point solid. The crystallinity suggests a regular packing of the molecules.

The structure is described as face centred cubic - it is a cube of iodine molecules with another molecule at the centre of each face.
The orientation of the iodine molecules within this structure is quite difficult to draw (let alone remember!). If your syllabus and past exam papers suggests that you need to remember it, look carefully at the next sequence of diagrams showing the layers.

Note: If you are studying a UK-based syllabus and haven't got a copy of yoursyllabus or copies of recent past papers, follow this link to find out how to get them.

Notice that as you look down on the cube, all the molecules on the left and right hand sides are aligned the same way. The ones in the middle are aligned in the opposite way.
All these diagrams show an "exploded" view of the crystal. The iodine molecules are, of course, touching each other. Measurements of the distances between the centres of the atoms in the crystal show two different values:

The iodine atoms within each molecule are pulled closely together by the covalent bond. The van der Waals attraction between the molecules is much weaker, and you can think of the atoms in two separate molecules as just loosely touching each other.

Ice
Ice is a good example of a hydrogen bonded solid.
There are lots of different ways that the water molecules can be arranged in ice. This is one of them, but NOT the common one - I can't draw that in any way that makes sense! The one below is known as "cubic ice", or "ice Ic". It is based on the water molecules arranged in a diamond structure.

This is just a small part of a structure which extends over huge numbers of molecules in three dimensions. In the diagram, the lines represent hydrogen bonds. The lone pairs that the hydrogen atoms are attracted to are left out for clarity.
Cubic ice is only stable at temperatures below -80°C. The ice you are familiar with has a different, hexagonal structure. It is called "ice Ih".

Note: Don't worry about this problem. If asked to draw ice in an exam at this level (16 - 18 year olds), don't try to be too clever. It is probably best not to go beyond the top five molecules in the above diagram. This will show the essential features of the bonding in the structure without getting bogged down in stuff which is far beyond this level.If you are interested in following this up, try a Google search using the search term ice structure hexagonal cubic (or something similar). This will throw up lots of information together with an assortment of fairly dreadful diagrams which I for one don't have the visual imagination to unscramble!

The unusual density behaviour of water
The hydrogen bonding forces a rather open structure on the ice - if you made a model of it, you would find a significant amount of wasted space. When ice melts, the structure breaks down and the molecules tend to fill up this wasted space.
This means that the water formed takes up less space than the original ice. Ice is a very unusual solid in this respect - most solids show an increase in volume on melting.
When water freezes, the opposite happens - there is an expansion as the hydrogen bonded structure establishes. Most liquids contract on freezing.
Remnants of the rigid hydrogen bonded structure are still present in very cold liquid water, and don't finally disappear until 4°C. From 0°C to 4°C, the density of water increases as the molecules free themselves from the open structure and take up less space. After 4°C, the thermal motion of the molecules causes them to move apart and the density falls. That's the normal behaviour with liquids on heating.

Note: You can find more about water (particularly its abnormally high boiling point) in the page on hydrogen bonding.

Polymers
Bonding in polymers
Polymers like poly(ethene) - commonly called polythene - consist of very long molecules. Poly(ethene) molecules are made by joining up lots of ethene molecules into chains of covalently bound carbon atoms with hydrogens attached. There may be short branches along the main chain, also consisting of carbon chains with attached hydrogens. The molecules are attracted to each other in the solid by van der Waals dispersion forces.
By controlling the conditions under which ethene is polymerised, it is possible to control the amount of branching to give two distinct types of polythene.
High density polythene
High density polythene has virtually unbranched chains. The lack of branching allows molecules to lie close together in a regular way which is almost crystalline.
Because the molecules lie close together, dispersion forces are more effective, and so the plastic is relatively strong and has a somewhat higher melting point than low density polythene.
High density polythene is used for containers for household chemicals like washing-up liquid, for example, or for bowls or buckets.
Low density polythene
Low density polythene has lots of short branches along the chain. These branches prevent the chains from lying close together in a tidy arrangement. As a result dispersion forces are less and the plastic is weaker and has a lower melting point. Its density is lower, of course, because of the wasted space within the unevenly packed structure.
Low density polythene is used for things like plastic bags.

The Shape of Molecules

The three dimensional shape or configuration of a molecule is an important characteristic. This shape is dependent on the preferred spatial orientation of covalent bonds to atoms having two or more bonding partners. Three dimensional configurations are best viewed with the aid of models.

In order to represent such configurations on a two-dimensional surface (paper, blackboard or screen), we often use perspective drawings in which the direction of a bond is specified by the line connecting the bonded atoms.

In most cases the focus of configuration is a carbon atom so the lines specifying bond directions will originate there. As defined in the diagram on the right, a simple straight line represents a bond lying approximately in the surface plane. The two bonds to substituents A in the structure on the left are of this kind. A wedge shaped bond is directed in front of this plane (thick end toward the viewer), as shown by the bond to substituent B; and a

hatched bond is directed in back of the plane (away from the viewer), as shown by the bond to substituent D. Some texts and other sources may use a dashed bond in the same manner as we have defined the hatched bond, but this can be confusing because the dashed bond is often used to represent a partial bond (i.e. a covalent bond that is partially formed or partially broken). The following examples make use of this notation, and also illustrate the importance of including non-bonding valence shell electron pairs (colored blue) when viewing such configurations .

Methane

Ammonia

Water

Bonding configurations are readily predicted by valence-shell electron-pair repulsion theory, commonly referred to as VSEPR in most introductory chemistry texts. This simple model is based on the fact that electrons repel each other, and that it is reasonable to expect that the bonds and non-bonding valence electron pairs associated with a given atom will prefer to be as far apart as possible. The bonding configurations of carbon are easy to remember, since there are only three categories.

Configuration

Bonding Partners

Bond Angles

Example

Tetrahedral

4

109.5º

Trigonal

3

120º

Linear

2

180º

In the three examples shown above, the central atom (carbon) does not have any non-bonding valence electrons; consequently the configuration may be estimated from the number of bonding partners alone. For molecules of water and ammonia, however, the non-bonding electrons must be included in the calculation. In each case there are four regions of electron density associated with the valence shell so that a tetrahedral bond angle is expected. The measured bond angles of these compounds (H2O 104.5º & NH3 107.3º) show that they are closer to being tetrahedral than trigonal or linear. Of course, it is the configuration of atoms (not electrons) that defines the the shape of a molecule, and in this sense ammonia is said to be pyramidal (not tetrahedral). The compound boron trifluoride, BF3, does not have non-bonding valence electrons and the configuration of its atoms is trigonal. Nice treatments of VSEPR theory have been provided by Oxford and Purdue . Click on the university name to visit their site.

The best way to study the three-dimensional shapes of molecules is by using molecular models. Many kinds of model kits are available to students and professional chemists. Some of the useful features of physical models can be approximated by the model viewing applet Jmol. This powerful visualization tool allows the user to move a molecular stucture in any way desired. Atom distances and angles are easily determined. To measure a distance, double-click on two atoms. To measure a bond angle, do a double-click, single-click, double-click on three atoms
Isomers

Structural Formulas

It is necessary to draw structural formulas for organic compounds because in most cases a molecular formula does not uniquely represent a single compound. Different compounds having the same molecular formula are called isomers, and the prevalence of organic isomers reflects the extraordinary versatility of carbon in forming strong bonds to itself and to other elements.

When the group of atoms that make up the molecules of different isomers are bonded together in fundamentally different ways, we refer to such compounds as constitutional isomers. There are seven constitutional isomers of C4H10O, and structural formulas for these are drawn in the following table. These formulas represent all known and possible C4H10O compounds, and display a common structural feature. There are no double or triple bonds and no rings in any of these structures.. Note that each of the carbon atoms is bonded to four other atoms, and is saturated with bonding partners.

Structural Formulas for C4H10O Isomers

Kekulé Formula

Condensed Formula

Shorthand Formula

Simplification of structural formulas may be achieved without any loss of the information they convey. In condensed structural formulas the bonds to each carbon are omitted, but each distinct structural unit (group) is written with subscript numbers designating multiple substituents, including the hydrogens. Shorthand (line) formulas omit the symbols for carbon and hydrogen entirely. Each straight line segment represents a bond, the ends and intersections of the lines are carbon atoms, and the correct number of hydrogens is calculated from the tetravalency of carbon. Non-bonding valence shell electrons are omitted in these formulas.

Developing the ability to visualize a three-dimensional structure from two-dimensional formulas requires practice, and in most cases the aid of molecular models. As noted earlier, many kinds of model kits are available to students and professional chemists, and the beginning student is encouraged to obtain one.
Analysis of Molecular Formulas

Although structural formulas are essential to the unique description of organic compounds, it is interesting and instructive to evaluate the information that may be obtained from a molecular formula alone. Three useful rules may be listed:
The number of hydrogen atoms that can be bonded to a given number of carbon atoms is limited by the valence of carbon. For compounds of carbon and hydrogen (hydrocarbons) the maximum number of hydrogen atoms that can be bonded to n carbons is 2n + 2 (n is an integer). In the case of methane, CH4, n=1 & 2n + 2 = 4. The origin of this formula is evident by considering a hydrocarbon made up of a chain of carbon atoms. Here the middle carbons will each have two hydrogens and the two end carbons have three hydrogens each. Thus, a six-carbon chain (n = 6) may be written H-(CH2)6-H, and the total hydrogen count is (2 x 6) + 2 = 14. The presence of oxygen (valence = 2) does not change this relationship, so the previously described C4H10O isomers follow the rule, n=4 &2n + 2 = 10. Halogen atoms (valence = 1) should be counted equivalent to hydrogen, as illustrated by C3H5Cl3, n = 3 & 2n + 2 = 8 = (5 + 3). If nitrogen is present, each nitrogen atom (valence = 3) increases the maximum number of hydrogens by one.

Some Plausible
Molecular Formulas

C7H16O3, C9H18, C15H28O3, C6H16N2

Some Impossible
Molecular Formulas

C8H20O6, C23H50, C5H10Cl4, C4H12NO

For stable organic compounds the total number of odd-valenced atoms is even. Thus, when even-valenced atoms such as carbon and oxygen are bonded together in any number and in any manner, the number of remaining unoccupied bonding sites must be even. If these sites are occupied by univalent atoms such as H, F, Cl, etc. their total number will necessarily be even. Nitrogen is also an odd-valenced atom (3), and if it occupies a bonding site on carbon it adds two additional bonding sites, thus maintaining the even/odd parity.

Some Plausible
Molecular Formulas

C4H4Cl2, C5H9OBr, C5H11NO2, C12H18N2FCl

Some Impossible
Molecular Formulas

C5H9O2, C4H5ClBr, C6H11N2O, C10H18NCl2
The number of hydrogen atoms in stable compounds of carbon, hydrogen & oxygen reflects the number of double bonds and rings in their structural formulas. Consider a hydrocarbon with a molecular structure consisting of a simple chain of four carbon atoms, CH3CH2CH2CH3. The molecular formula is C4H10 (the maximum number of bonded hydrogens by the 2n + 2 rule). If the four carbon atoms form a ring, two hydrogens must be lost. Similarly, the introduction of a double bond entails the loss of two hydrogens, and a triple bond the loss of four hydrogens.

From the above discussion and examples it should be clear that the molecular formula of a hydrocarbon (CnHm) provides information about the number of rings and/or double bonds that must be present in its structural formula. By rule #2 m must be an even number, so if m < (2n + 2) the difference is also an even number that reflects any rings and double bonds. A triple bond is counted as two double bonds.

The presence of one or more nitrogen atoms or halogen substituents requires a modified analysis.
Resonance

Kekulé structural formulas are

essential tools for understanding organic chemistry. However, the structures of some compounds and ions cannot be represented by a single formula. For example, sulfur dioxide (SO2) and nitric acid (HNO3) may each be described by two equivalent formulas (equations 1 & 2). For clarity the two ambiguous bonds to oxygen are given different colors in these formulas.

1) sulfur dioxide

2) nitric acid

If only one formula for sulfur dioxide was correct and accurate, then the double bond to oxygen would be shorter and stronger than the single bond. Since experimental evidence indicates that this molecule is bent (bond angle 120º) and has equal length sulfur : oxygen bonds (1.432 Å), a single formula is inadequate, and the actual structure resembles an average of the two formulas. This averaging of electron distribution over two or more hypothetical contributing structures (canonical forms) to produce a hybrid electronic structure is called resonance. Likewise, the structure of nitric acid is best described as aresonance hybrid of two structures, the double headed arrow being the unique symbol for resonance.

The above examples represent one extreme in the application of resonance. Here, two structurally and energetically equivalent electronic structures for a stable compound can be written, but no single structure provides an accurate or even an adequate representation of the true molecule. In cases such as these, the electron delocalization described by resonance enhances the stability of the molecules, and compounds or ions composed of such molecules often show exceptional stability.

3) formaldehyde

The electronic structures of most covalent compounds do not suffer the inadequacy noted above. Thus, completely satisfactory Kekulé formulas may be drawn for water (H2O), methane (CH4) and acetylene C2H2). Nevertheless, the principles of resonance are very useful in rationalizing the chemical behavior of many such compounds. For example, the carbonyl group of formaldehyde (the carbon-oxygen double bond) reacts readily to give addition products. The course of these reactions can be explained by a small contribution of a dipolar resonance contributor, as shown in equation 3. Here, the first contributor (on the left) is clearly the best representation of this molecular unit, since there is no charge separation and both the carbon and oxygen atoms have achieved valence shell neon-like configurations by covalent electron sharing. If the double bond is broken heterolytically, formal charge pairs result, as shown in the other two structures. The preferred charge distribution will have the positive charge on the less electronegative atom (carbon) and the negative charge on the more electronegative atom (oxygen). Therefore the middle formula represents a more reasonable and stable structure than the one on the right. The application of resonance to this case requires a weighted averaging of these canonical structures. The double bonded structure is regarded as the major contributor, the middle structure a minor contributor and the right hand structure a non-contributor. Since the middle, charge-separated contributor has an electron deficient carbon atom, this explains the tendency of electron donors (nucleophiles) to bond at this site.

The basic principles of the resonance method may now be summarized.

For a given compound, a set of Lewis / Kekulé structures are written, keeping the relative positions of all the component atoms the same. These are the canonical forms to be considered, and all must have the same number of paired and unpaired electrons.

The following factors are important in evaluating the contribution each of these canonical structures makes to the actual molecule.
The number of covalent bonds in a structure. (The greater the bonding, the more important and stable the contributing structure.)
Formal charge separation. (Other factors aside, charge separation decreases the stability and importance of the contributing structure.)
Electronegativity of charge bearing atoms and charge density. (High charge density is destabilizing. Positive charge is best accommodated on atoms of low electronegativity, and negative charge on high electronegative atoms.)

The stability of a resonance hybrid is always greater than the stability of any canonical contributor. Consequently, if one canonical form has a much greater stability than all others, the hybrid will closely resemble it electronically and energetically. This is the case for the carbonyl group (eq.3). The left hand C=O structure has much greater total bonding than either charge-separated structure, so it describes this functional group rather well. On the other hand, if two or more canonical forms have identical low energy structures, the resonance hybrid will have exceptional stabilization and unique properties. This is the case for sulfur dioxide (eq.1) and nitric acid (eq.2).

4) carbon monoxide

5) azide anion

To illustrate these principles we shall consider carbon monoxide (eq.4) and azide anion (eq.5). In each case the most stable canonical form is on the left. For carbon monoxide, the additional bonding is more important than charge separation. Furthermore, the double bonded structure has an electron deficient carbon atom (valence shell sextet). A similar destabilizing factor is present in the two azide canonical forms on the top row of the bracket (three bonds vs. four bonds in the left most structure). The bottom row pair of structures have four bonds, but are destabilized by the high charge density on a single nitrogen atom.
Atomic and Molecular Orbitals

A more detailed model of covalent bonding requires a consideration of valence shell atomic orbitals. For second period elements such as carbon, nitrogen and oxygen, these orbitals have been designated 2s, 2px, 2py & 2pz. The spatial distribution of electrons occupying each of these orbitals is shown in the diagram below.

Very nice displays of orbitals may be found at the following sites:

J. Gutow, Univ. Wisconsin Oshkosh

R. Spinney, Ohio State

M. Winter, Sheffield University

The valence shell electron configuration of carbon is 2s2, 2px1, 2py1 & 2pz0. If this were the configuration used in covalent bonding, carbon would only be able to form two bonds. In this case, the valence shell would have six electrons- two shy of an octet. However, the tetrahedral structures of methane and carbon tetrachloride demonstrate that carbon can form four equivalent bonds, leading to the desired octet. In order to explain this covalent bonding, Linus Pauling proposed an orbital hybridization model in which all the valence shell electrons of carbon are reorganized.

Hybrid Orbitals

In order to explain the structure of methane (CH4), the 2s and three 2p orbitals are converted to four equivalent hybrid atomic orbitals, each having 25% s and 75% p character, and designated sp3. These hybrid orbitals have a specific orientation, and the four are naturally oriented in a tetrahedral fashion. Thus, the four covalent bonds of methane consist of shared electron pairs with four hydrogen atoms in a tetrahedral configuration, as predicted by VSEPR theory.

Molecular Orbitals

Just as the valence electrons of atoms occupy atomic orbitals (AO), the shared electron pairs of covalently bonded atoms may be thought of as occupying molecular orbitals (MO). It is convenient to approximate molecular orbitals by combining or mixing two or more atomic orbitals. In general, this mixing of n atomic orbitals always generates n molecular orbitals. The hydrogen molecule provides a simple example of MO formation. In the following diagram, two 1s atomic orbitals combine to give a sigma (σ) bonding (low energy) molecular orbital and a second higher energy MO referred to as an antibonding orbital. The bonding MO is occupied by two electrons of opposite spin, the result being a covalent bond.

The notation used for molecular orbitals parallels that used for atomic orbitals. Thus, s-orbitals have a spherical symmetry surrounding a single nucleus, whereas σ-orbitals have a cylindrical symmetry and encompass two (or more) nuclei. In the case of bonds between second period elements, p-orbitals or hybrid atomic orbitals having p-orbital character are used to form molecular orbitals. For example, the sigma molecular orbital that serves to bond two fluorine atoms together is generated by the overlap of p-orbitals (part A below), and two sp3 hybrid orbitals of carbon may combine to give a similar sigma orbital. When these bonding orbitals are occupied by a pair of electrons, a covalent bond, the sigma bond results. Although we have ignored the remaining p-orbitals, their inclusion in a molecular orbital treatment does not lead to any additional bonding, as may be shown by activating the fluorine correlation diagram below.

Another type of MO (the π orbital) may be formed from two p-orbitals by a lateral overlap, as shown in part A of the following diagram. Since bonds consisting of occupied π-orbitals (pi-bonds) are weaker than sigma bonds, pi-bonding between two atoms occurs only when a sigma bond has already been established. Thus, pi-bonding is generally found only as a component of double and triple covalent bonds. Since carbon atoms involved in double bonds have only three bonding partners, they require only three hybrid orbitals to contribute to three sigma bonds. A mixing of the 2s-orbital with two of the 2p orbitals gives three sp2 hybrid orbitals, leaving one of the p-orbitals unused. Two sp2 hybridized carbon atoms are then joined together by sigma and pi-bonds (a double bond), as shown in part B.

The manner in which atomic orbitals overlap to form molecular orbitals is actually more complex than the localized examples given above. These are useful models for explaining the structure and reactivity of many organic compounds, but modern molecular orbital theory involves the creation of an orbital correlation diagram. Two examples of such diagrams for the simple diatomic elements F2 and N2 will be drawn above when the appropriate button is clicked. The 1s and 2s atomic orbitals do not provide any overall bonding, since orbital overlap is minimal, and the resulting sigma bonding and antibonding components would cancel. In both these cases three 2p atomic orbitals combine to form a sigma and two pi-molecular orbitals, each as a bonding and antibonding pair. The overall bonding order depends on the number of antibonding orbitals that are occupied. The subtle change in the energy of the σ2p bonding orbital, relative to the two degenerate π-bonding orbitals, is due to s-p hybridization that is unimportant to the present discussion.

The p-orbitals in these model are represented by red and blue colored spheres or ellipses, which represent different phases, defined by the mathematical wave equations for such orbitals.

Finally, in the case of carbon atoms with only two bonding partners only two hybrid orbitals are needed for the sigma bonds, and these sp hybrid orbitals are directed 180º from each other. Two p-orbitals remain unused on each sp hybridized atom, and these overlap to give two pi-bonds following the formation of a sigma bond (a triple bond), as shown below.

States of Matter(Gases, liquids and solids)

Gases, liquids and solids are all made up of microscopic particles, but the behaviors of these particles differ in the three phases. The following figure illustrates the microscopic differences.


Microscopic view of a gas.	Microscopic view of a liquid.	Microscopic view of a solid.

Note that:

Particles in a:

gas are well separated with no regular arrangement.
liquid are close together with no regular arrangement.
solid are tightly packed, usually in a regular pattern.

Particles in a:

gas vibrate and move freely at high speeds.
liquid vibrate, move about, and slide past each other.
solid vibrate (jiggle) but generally do not move from place to place.

Liquids and solids are often referred to as condensed phases because the particles are very close together.

The following table summarizes properties of gases, liquids, and solids and identifies the microscopic behavior responsible for each property.

Some Characteristics of Gases, Liquids and Solids and the Microscopic Explanation for the Behavior
gas	liquid	solid
assumes the shape and volume of its container particles can move past one another	assumes the shape of the part of the container which it occupies particles can move/slide past one another	retains a fixed volume and shape rigid - particles locked into place
compressible lots of free space between particles	not easily compressible little free space between particles	not easily compressible little free space between particles
flows easily particles can move past one another	flows easily particles can move/slide past one another	does not flow easily rigid - particles cannot move/slide past one another

Solutions and Dissolving

What is a solution?

A solution is a specific type of mixture where one substance is dissolved into another. A solution is the same, or uniform, throughout which makes it a homogeneous mixture . Go here to learn more about mixtures.

A solution has certain characteristics:

It is uniform, or homogeneous, throughout the mixture

It is stable and doesn't change over time or settle

The solute particles are so small they cannot be separated by filtering

The solute and solvent molecules cannot be distinguished by the naked eye

It does not scatter a beam of light

Example of a Solution

One example of a solution is salt water which is a mixture of water and salt. You cannot see the salt and the salt and water will stay a solution if left alone.

Parts of a Solution

Solute - The solute is the substance that is being dissolved by another substance. In the example above, the salt is the solute.

Solvent - The solvent is the substance that dissolves the other substance. In the example above, the water is the solvent.

A solution is a type of homogeneous mixture

Dissolving

A solution is made when one substance called the solute "dissolves" into another substance called the solvent. Dissolving is when the solute breaks up from a larger crystal of molecules into much smaller groups or individual molecules. This break up is caused by coming into contact with the solvent.

In the case of salt water, the water molecules break off salt molecules from the larger crystal lattice. They do this by pulling away the ions and then surrounding the salt molecules. Each salt molecule still exists. It is just now surrounded by water molecules instead of fixed to a crystal of salt.

Solubility

Solubility is a measure of how much solute can be dissolved into a liter of solvent. Think of the example of water and salt. If you keep pouring salt into water, at some point the water isn't going to be able to dissolve the salt.

Saturated

When a solution reaches the point where it cannot dissolve any more solute it is considered "saturated." If a saturated solution loses some solvent, then solid crystals of the solute will start to form. This is what happens when water evaporates and salt crystals begin to form.

Concentration

The concentration of a solution is the proportion of the solute to solvent. If there is a lot of solute in a solution, then it is "concentrated". If there is a low amount of solute, then the solution is said to be "diluted."

Miscible and immiscible

When two liquids can be mixed to form a solution they are called "miscible." If two liquids cannot be mixed to form a solution they are called "immiscible." An example of miscible liquids is alcohol and water. An example of immiscible liquids is oil and water. Have you ever heard the saying "oil and water don't mix"? This is because they are immiscible.

Interesting Facts about Solutions

There is a solvent called aqua regia which can dissolve the noble metals including gold and platinum.
You can't see a beam of light when shining it through a true solution. This means fog is not a solution. It is a colloid.
Solutions can be liquid, solid, or gas. An example of a solid solution is steel.
Solids are generally more soluble at higher temperatures.

Carbonated beverages are made by dissolving carbon dioxide gas into liquid at high pressure.

solution:

DEFINITION OF SOLUTION

A homogeneous mixture, which may be liquid, gas or solid, formed by dissolving one or more substances.

Properties can be considered colligative only if their properties are dependent on the amount of solute present in the solution, disregarding the identity of the solute itself.
To better visualize the effect of solute on the vapor pressure of a solution, consider a pure solvent.
Subjected to temperatures below the solvent's boiling point, the molecules changing to the gaseous phase are mostly situated on the top layer of the solution.
Now consider a solution composed of both solvent and solute.
In an ideal solution, the vapor pressure is dependent on the vapor pressure of each chemical component and the mole fraction of the component present in thesolution.
The vapor pressure of an electrolytic solution is dependent on the ratio of solute to solvent molecules in a solution.
Molar solubility, which is directly related to the solubility product, is the number of moles of the solute that can be dissolved per liter of solution before the solutionbecomes saturated.
Once a solution is saturated, any additional solute precipitates out of the solution.
Solution: The balanced equation for the reaction is:$AgI (s) \leftrightarrow Ag^+ (aq) + I^- (aq)$The formula for Ksp is:Ksp = [Ag2+][I-]Ksp = s2 = 8.5 x 10-17 where s is the concentration of each ion at equilibrium.
Molar solubility is the number of moles of a solute that can be dissolved per liter ofsolution before the solution becomes saturated.
Dilution refers to the reduction in concentration of a chemical, often in the gaseous, vapor or solution form.
In doing so, the amount of solute remains constant, but the amount of solutionincreases effectively diluting the final concentration of the solution.
Dilution of a solution of a given chemical entity can also be realized by mixingsolution of higher concentration with a solution of lower concentration.
The relationship of the concentrated and new solution are given by the equation: , where M1 denotes the concentration of the original solution, M2 represents the concentration of the desired solution, V1 represents the volume of the originalsolution (usually how much of the original solution needed to make the newsolution) and V2 represents the volume of the new solution.
Another type of dilution is a serial dilution which is the stepwise dilution of a substance in solution.
In the process of dilution, the concentration of the solute decreases through the addition of additional solvent to the solution.
Commonly used units are listed in the table hereafter: Applications of Molarity To determine the molarity of a solution, the number of moles of solute added must be divided by the number of liters of total solution produced.
To determine the number of moles in a given solution of known molarity, simply multiply the molarity times the volume used, where V is the volume in liters.For example, how many moles of potassium chloride (KCl) are in 4L of a 0.65Msolution?
Dilution (equation) Dilution is a reduction in the concentration of a chemical (gas, vapor, solution).
To dilute a solution means to add more solvent without the addition of more solute.
The resulting solution is thoroughly mixed so as to ensure that all parts of thesolution are identical.
Molarity is defined as the moles of a solute per volume of total solution.

The enthalpy of solution, enthalpy of dissolution, or heat of solution is the enthalpy change associated with the dissolution of a substance in a solvent at constant pressure resulting in infinite dilution.
The enthalpy of solution is most often expressed in kJ/mol at constant temperature.
The temperature of the solution then decreases to that of the surroundings.
When a saturated solution of a gas is heated, gas comes out of solution.
The value of the enthalpy of solution is the sum of these individual steps.
Heat of solution is the energy released or absorbed during dissolution of a substance in a solvent.
Acid Dissociation Constant (Ka)An acid dissociation constant, Ka, is a quantitative measure of the strength of an acid in solution.
Acid dissociation constants are associated with weak acids, or acids that do not completely dissociate in solution; strong acids that completely dissociate insolution exist in their dissociated form when in equilibrium.In the chemical species HA (a generic acid), A- (the conjugate base of the acid) and H+ (the hydrogen ion or proton) are said to be in equilibrium when their concentrations do not change with the passing of time.
A strong acid is almost completely dissociated in aqueous solution; it is dissociated to the extent that the concentration of the undissociated acid becomes undetectable. pKa values for strong acids can be estimated by theoretical means or by extrapolating from measurements in non-aqueous solvents with a smaller dissociation constant, such as acetonitrile and dimethylsulfoxide.Acetic acid is a weak acid.
When it is introduced into water, it partially ionizes into acetate ions and hydrogen ions .At equilibrium, the Ka of acetic acid in solution is 1.8X10-5.
The acid dissociation constant, Ka, is the measure of the strength of an acid insolution.
The Effect of pH on Solubility The pH of an aqueous solution can affect the solubility of the solute.
By changing the pH of the solution, you can change the charge start of the solute.
If the pH of the solution is such that a particular molecule carries no net electric charge, the solute often has minimal solubility and precipitates out of the solution.
When the proteins are added to the solution and current is applied, they migrate toward the electrode with the opposite charge.
By changing the pH of the solution, you can change the charge state of the solute.
This allows for quantitative analysis of the concentration of an unknown acid or basesolution.
It is often wrongly assumed that neutralization should result in a solution with pH 7.0; this is only the case when the acid and based used have similar strengths.The table in shows the pK values of typical acid-base systems.
Because H2O serves as a proton sink to any acid in which the proton free energy level is greater than zero, the strong acids, such as HCl and H2SO4, cannot exist (as acids) in aqueous solution.
This is the basis of the leveling effect, which states that the strongest acid that can exist in aqueous solution is H3O+.
A strong acid will react with a strong base to form a neutral (pH = 7) solution.
Electrolytic PropertiesWhen electrodes are placed in an electrolyte solution and a voltage is applied, the electrolyte will conduct electricity.
The ions in the electrolyte neutralize these charges, enabling the electrons to keep flowing and the reactions to continue.For example, in a solution of ordinary table salt (sodium chloride, NaCl) in water, the cathode reaction will be:$2H_{2}O + 2e^{-} \rightarrow 2OH^{-} + H_{2}$and hydrogen gas will bubble up.
For example: p-Benzoquinone can be reduced to hydroquinone at the cathode: $+ 2 e^{-} + 2 H^{+} \rightarrow$ In the last example, H+ ions (hydrogen ions) also take part in the reaction, and are provided by an acid in the solution or by the solvent itself (water, methanol, etc.).
It is possible to have electrolysis involving gases.In order to determine which species in solution will be oxidized and which will be reduced, the standard electrode potential of each species can often be obtained from a table of standard electrode potentials, a small sampling of which is shown here:
When electrodes are placed in an electrolyte solution and a voltage is applied, the electrolyte will conduct electricity.
I
n chemistry, the molality, b (or m), of a solution is defined as the amount of substance of solute (in moles), nsolute, divided by the mass in kg of the solvent, msolvent:$bM_{solute}=\frac{m_{solute}}{m_{solvent}}=\frac{w_{solute}}{w_{solvent}}$Molality is an intensive property of solutions, such as sodium chloride (table salt) in water .OriginThe earliest definition of molality was most likely coined by G.
The two words "molality" and "molarity" are apt to be confused with one another, and in fact the molality and molarity of a weak aqueous solution happen to be nearly the same, as one kilogram of water (the solvent) occupies one liter of volume at room temperature, and the small amount of solute would have little effect on the volume of the solvent.UnitThe SI unit for molality is mol/kg.
A solution with a molality of 3 mol/kg is often described as "3 molal" or "3 m."
Despite this, their recommendation has not been universally implemented in scientific literature, and many references to molality still exist.Usage ConsiderationsCompared to molar concentration or mass concentration, the preparation of a solution of a given molality requires only a good scale: both solvent and solute need to be weighed, as opposed to measured volumetrically.
Molality is a property of a solution that indicates the moles of solute per kilogram of solvent.electrochemistry
The branch of chemistry that deals with the relations between electrical and chemical phenomena.
or
Electrochemistry is a branch of chemistry that studies chemical reactions which take place in a solution at the interface of an electronconductor (the electrode: a metal or a semiconductor) and an ionic conductor (the electrolyte). These reactions involve electron transfer between the electrode and the electrolyte or species in solution.Thus electrochemistry deals with interactions between electrical energy and chemical change and vice versa.
If a chemical reaction is driven by an externally applied voltage, as in electrolysis, or if a voltage is created by a chemical reaction as in abattery, it is an electrochemical reaction. In contrast, chemical reactions where electrons are transferred between molecules are calledoxidation-reduction (redox) reactions. In general, electrochemistry deals with situations where redox reactions are separated in space or time, connected by an external electric circuit.
The term "redox" stands for reduction-oxidation. It refers to electrochemical processes involving electron transfer to or from a molecule or ion changing its oxidation state. This reaction can occur through the application of an external voltage or through the release of chemical energy. Oxidation and reduction describe the change of oxidation state that takes place in the atoms, ions or molecules involved in an electrochemical reaction. Formally, oxidation state is the hypothetical charge that an atom would have if all bonds to atoms of different elements were 100% ionic. An atom or ion that gives up an electron to another atom or ion has its oxidation state increase, and the recipient of the negatively charged electron has its oxidation state decrease.
For example, when atomic sodium reacts with atomic chlorine, sodium donates one electron and attains an oxidation state of +1. Chlorine accepts the electron and its oxidation state is reduced to −1. The sign of the oxidation state (positive/negative) actually corresponds to the value of each ion's electronic charge. The attraction of the differently charged sodium and chlorine ions is the reason they then form an ionic bond.
The loss of electrons from an atom or molecule is called oxidation, and the gain of electrons is reduction. This can be easily remembered through the use of mnemonicdevices. Two of the most popular are "OIL RIG" (Oxidation Is Loss, Reduction Is Gain) and "LEO" says "GER" (Lose Electrons: Oxidation, Gain Electrons: Reduction). Oxidation and reduction always occur in a paired fashion such that one species is oxidized when another is reduced. For cases where electrons are shared (covalent bonds) between atoms with large differences in electronegativity, the electron is assigned to the atom with the largest electronegativity in determining the oxidation state.
The atom or molecule which loses electrons is known as the reducing agent, or reductant, and the substance which accepts the electrons is called the oxidizing agent, oroxidant. Thus, the oxidizing agent is always being reduced in a reaction; the reducing agent is always being oxidized. Oxygen is a common oxidizing agent, but not the only one. Despite the name, an oxidation reaction does not necessarily need to involve oxygen. In fact, a fire can be fed by an oxidant other than oxygen; fluorine fires are often unquenchable, as fluorine is an even stronger oxidant (it has a higher electronegativity and thus accepts electrons even better) than oxygen.
For reactions involving oxygen, the gain of oxygen implies the oxidation of the atom or molecule to which the oxygen is added (and the oxygen is reduced). In organic compounds, such as butane or ethanol, the loss of hydrogen implies oxidation of the molecule from which it is lost (and the hydrogen is reduced). This follows because the hydrogen donates its electron in covalent bonds with non-metals but it takes the electron along when it is lost. Conversely, loss of oxygen or gain of hydrogen implies reduction.
Electrochemical reactions in water are better understood by balancing redox reactions using the ion-electron method where H+, OH– ion, H2O and electrons (to compensate the oxidation changes) are added to cell's half-reactions for oxidation and reduction.
Acidic medium:
In acid medium H+ ions and water are added to half-reactions to balance the overall reaction. For example, when manganese reacts with sodium bismuthate.
Unbalanced reaction: Mn2+(aq) + NaBiO3(s) → Bi3+(aq) + MnO4–(aq)
Oxidation: 4 H2O(l) + Mn2+(aq) → MnO4–(aq) + 8 H+(aq) + 5 e–
Reduction: 2 e– + 6 H+(aq) + BiO3–(s) → Bi3+(aq) + 3 H2O(l)
Finally, the reaction is balanced by multiplying the number of electrons from the reduction half reaction to oxidation half reaction and vice versa and adding both half reactions, thus solving the equation.
8 H2O(l) + 2 Mn2+(aq) → 2 MnO4–(aq) + 16 H+(aq) + 10 e–
10 e– + 30 H+(aq) + 5 BiO3–(s) → 5 Bi3+(aq) + 15 H2O(l)
Reaction balanced:
14 H+(aq) + 2 Mn2+(aq) + 5 NaBiO3(s) → 7 H2O(l) + 2 MnO4–(aq) + 5 Bi3+(aq) + 5 Na+(aq)
Basic medium:
In basic medium OH– ions and water are added to half reactions to balance the overall reaction. For example, on reaction between potassium permanganate and sodium sulfite.
Unbalanced reaction: KMnO4 + Na2SO3 + H2O → MnO2 + Na2SO4 + KOH
Reduction: 3 e– + 2 H2O + MnO4– → MnO2 + 4 OH–
Oxidation: 2 OH– + SO32– → SO42– + H2O + 2 e–
The same procedure as followed on acid medium by multiplying electrons to opposite half reactions solve the equation thus balancing the overall reaction.
6 e– + 4 H2O + 2 MnO4– → 2 MnO2 + 8 OH–
6 OH– + 3 SO32– → 3 SO42– + 3 H2O + 6e–
Equation balanced:
2 KMnO4 + 3 Na2SO3 + H2O → 2 MnO2 + 3 Na2SO4 + 2 KOH
Neutral medium:
The same procedure as used on acid medium is applied, for example on balancing using electron ion method to complete combustion of propane.
Unbalanced reaction: C3H8 + O2 → CO2 + H2O
Reduction: 4 H+ + O2 + 4 e– → 2 H2O
Oxidation: 6 H2O + C3H8 → 3 CO2 + 20 e– + 20 H+
As in acid and basic medium, electrons which were used to compensate oxidation changes are multiplied to opposite half reactions, thus solving the equation.
20 H+ + 5 O2 + 20 e– → 10 H2O
6 H2O + C3H8 → 3 CO2 + 20 e– + 20 H+
Equation balanced:
C3H8 + 5 O2 → 3 CO2 + 4 H2O
Electrochemical cells
An electrochemical cell is a device that produces an electric current from energy released by a spontaneous redox reaction. This kind of cell includes the Galvanic cell or Voltaic cell, named after Luigi Galvani and Alessandro Volta, both scientists who conducted several experiments on chemical reactions and electric current during the late 18th century.
Electrochemical cells have two conductive electrodes (the anode and the cathode). The anode is defined as the electrode where oxidation occurs and the cathode is the electrode where the reduction takes place. Electrodes can be made from any sufficiently conductive materials, such as metals, semiconductors, graphite, and even conductive polymers. In between these electrodes is the electrolyte, which contains ions that can freely move.
The galvanic cell uses two different metal electrodes, each in an electrolyte where the positively charged ions are the oxidized form of the electrode metal. One electrode will undergo oxidation (the anode) and the other will undergo reduction (the cathode). The metal of the anode will oxidize, going from an oxidation state of 0 (in the solid form) to a positive oxidation state and become an ion. At the cathode, the metal ion in solution will accept one or more electrons from the cathode and the ion's oxidation state is reduced to 0. This forms a solid metal that electrodeposits on the cathode. The two electrodes must be electrically connected to each other, allowing for a flow of electrons that leave the metal of the anode and flow through this connection to the ions at the surface of the cathode. This flow of electrons is an electrical current that can be used to do work, such as turn a motor or power a light.
A galvanic cell whose electrodes are zinc and copper submerged in zinc sulfate and copper sulfate, respectively, is known as a Daniell cell.
Half reactions for a Daniell cell are these:
Zinc electrode (anode): Zn(s) → Zn2+(aq) + 2 e–
Copper electrode (cathode):

Cu2+(aq) + 2 e– → Cu(s)

A modern cell stand for electrochemical research. The electrodes attach to high-quality metallic wires, and the stand is attached to apotentiostat/galvanostat (not pictured). A shot glass-shaped container is aerated with a noble gas and sealed with the Teflonblock.
In this example, the anode is zinc metal which oxidizes (loses electrons) to form zinc ions in solution, and copper ions accept electrons from the copper metal electrode and the ions deposit at the copper cathode as an electrodeposit. This cell forms a simple battery as it will spontaneously generate a flow of electrical current from the anode to the cathode through the external connection. This reaction can be driven in reverse by applying a voltage, resulting in the deposition of zinc metal at the anode and formation of copper ions at the cathode.
To provide a complete electric circuit, there must also be an ionic conduction path between the anode and cathode electrolytes in addition to the electron conduction path. The simplest ionic conduction path is to provide a liquid junction. To avoid mixing between the two electrolytes, the liquid junction can be provided through a porous plug that allows ion flow while reducing electrolyte mixing. To further minimize mixing of the electrolytes, a salt bridge can be used which consists of an electrolyte saturated gel in an inverted U-tube. As the negatively charged electrons flow in one direction around this circuit, the positively charged metal ions flow in the opposite direction in the electrolyte.
A voltmeter is capable of measuring the change of electrical potential between the anode and the cathode.
Electrochemical cell voltage is also referred to as electromotive force or emf.
A cell diagram can be used to trace the path of the electrons in the electrochemical cell. For example, here is a cell diagram of a Daniell cell:
Zn(s) | Zn2+ (1M) || Cu2+ (1M) | Cu(s)
First, the reduced form of the metal to be oxidized at the anode (Zn) is written. This is separated from its oxidized form by a vertical line, which represents the limit between the phases (oxidation changes). The double vertical lines represent the saline bridge on the cell. Finally, the oxidized form of the metal to be reduced at the cathode, is written, separated from its reduced form by the vertical line. The electrolyte concentration is given as it is an important variable in determining the cell potential.
Standard electrode potential:
To allow prediction of the cell potential, tabulations of standard electrode potential are available. Such tabulations are referenced to the standard hydrogen electrode (SHE). The standard hydrogen electrode undergoes the reaction
2 H+(aq) + 2 e– → H2
which is shown as reduction but, in fact, the SHE can act as either the anode or the cathode, depending on the relative oxidation/reduction potential of the other electrode/electrolyte combination. The term standard in SHE requires a supply of hydrogen gas bubbled through the electrolyte at a pressure of 1 atm and an acidic electrolyte with H+ activity equal to 1 (usually assumed to be [H+] = 1 mol/liter).
The SHE electrode can be connected to any other electrode by a salt bridge to form a cell. If the second electrode is also at standard conditions, then the measured cell potential is called the standard electrode potential for the electrode. The standard electrode potential for the SHE is zero, by definition. The polarity of the standard electrode potential provides information about the relative reduction potential of the electrode compared to the SHE. If the electrode has a positive potential with respect to the SHE, then that means it is a strongly reducing electrode which forces the SHE to be the anode (an example is Cu in aqueous CuSO4 with a standard electrode potential of 0.337 V). Conversely, if the measured potential is negative, the electrode is more oxidizing than the SHE (such as Zn in ZnSO4 where the standard electrode potential is −0.76 V).
Standard electrode potentials are usually tabulated as reduction potentials. However, the reactions are reversible and the role of a particular electrode in a cell depends on the relative oxidation/reduction potential of both electrodes. The oxidation potential for a particular electrode is just the negative of the reduction potential. A standard cell potential can be determined by looking up the standard electrode potentials for both electrodes (sometimes called half cell potentials). The one that is smaller will be the anode and will undergo oxidation. The cell potential is then calculated as the sum of the reduction potential for the cathode and the oxidation potential for the anode.
E°cell = E°red(cathode) – E°red(anode) = E°red(cathode) + E°oxi(anode)
For example, the standard electrode potential for a copper electrode is:
Cell diagram
Pt(s) | H2(1 atm) | H+(1 M) || Cu2+ (1 M) | Cu(s)
E°cell = E°red(cathode) – E°red(anode)
At standard temperature, pressure and concentration conditions, the cell's emf (measured by a multimeter) is 0.34 V. By definition, the electrode potential for the SHE is zero. Thus, the Cu is the cathode and the SHE is the anode giving
Ecell = E°(Cu2+/Cu) – E°(H+/H2)
Or,
E°(Cu2+/Cu) = 0.34 V
Changes in the stoichiometric coefficients of a balanced cell equation will not change E°red value because the standard electrode potential is an intensive property.
Spontaneity of redox reaction:
During operation of electrochemical cells, chemical energy is transformed into electrical energy and is expressed mathematically as the product of the cell's emf and theelectric charge transferred through the external circuit.
Electrical energy = EcellCtrans
where Ecell is the cell potential measured in volts (V) and Ctrans is the cell current integrated over time and measured in coulombs (C); Ctrans can also be determined by multiplying the total number of electrons transferred (measured in moles) times Faraday's constant (F).
The emf of the cell at zero current is the maximum possible emf. It is used to calculate the maximum possible electrical energy that could be obtained from a chemical reaction. This energy is referred to as electrical work and is expressed by the following equation:
Wmax = Welectrical = –nF·Ecell,
where work is defined as positive into the system.
Since the free energy is the maximum amount of work that can be extracted from a system, one can write:
ΔG = –nF·Ecell
A positive cell potential gives a negative change in Gibbs free energy. This is consistent with the cell production of an electric current from the cathode to the anode through the external circuit. If the current is driven in the opposite direction by imposing an external potential, then work is done on the cell to drive electrolysis.
A spontaneous electrochemical reaction (change in Gibbs free energy less than zero) can be used to generate an electric current in electrochemical cells. This is the basis of all batteries and fuel cells. For example, gaseous oxygen (O2) and hydrogen (H2) can be combined in a fuel cell to form water and energy, typically a combination of heat and electrical energy.
Conversely, non-spontaneous electrochemical reactions can be driven forward by the application of a current at sufficient voltage. The electrolysis of water into gaseous oxygen and hydrogen is a typical example.
The relation between the equilibrium constant, K, and the Gibbs free energy for an electrochemical cell is expressed as follows:
ΔG° = –RT ln(K) = –nF·E°cell
Rearranging to express the relation between standard potential and equilibrium constant yields

.
The previous equation can use Briggsian logarithm as shown below:

Cell emf dependency on changes in concentration:
The standard potential of an electrochemical cell requires standard conditions (ΔG°) for all of the reactants. When reactant concentrations differ from standard conditions, the cell potential will deviate from the standard potential. In the 20th century German chemist Walther Nernst proposed a mathematical model to determine the effect of reactant concentration on electrochemical cell potential.
In the late 19th century, Josiah Willard Gibbs had formulated a theory to predict whether a chemical reaction is spontaneous based on the free energy
ΔG = ΔG° + RT·ln(Q)
Here ΔG is change in Gibbs free energy, ΔG° is the cell potential when Q is equal to 1, T is absolute temperature(Kelvin), R is the gas constant and Q is reaction quotientwhich can be found by dividing products by reactants using only those products and reactants that are aqueous or gaseous.
Gibbs' key contribution was to formalize the understanding of the effect of reactant concentration on spontaneity.
Based on Gibbs' work, Nernst extended the theory to include the contribution from electric potential on charged species. As shown in the previous section, the change in Gibbs free energy for an electrochemical cell can be related to the cell potential. Thus, Gibbs' theory becomes
nFΔE = nFΔE° – RT ln(Q)
Here n is the number of electrons/mole product, F is the Faraday constant (coulombs/mole), and ΔE is cell potential.
Finally, Nernst divided through by the amount of charge transferred to arrive at a new equation which now bears his name:
ΔE = ΔE° – (RT/nF)ln(Q)
Assuming standard conditions (T = 25 °C) and R = 8.3145 J/(K·mol), the equation above can be expressed on base—10 logarithm as shown below:[24]

Main article: Concentration cell
A concentration cell is an electrochemical cell where the two electrodes are the same material, the electrolytes on the two half-cells involve the same ions, but the electrolyte concentration differs between the two half-cells.
An example is an electrochemical cell, where two copper electrodes are submerged in two copper(II) sulfate solutions, whose concentrations are 0.05 M and 2.0 M, connected through a salt bridge. This type of cell will generate a potential that can be predicted by the Nernst equation. Both can undergo the same chemistry (although the reaction proceeds in reverse at the anode)
Cu2+(aq) + 2 e– → Cu(s)
Le Chatelier's principle indicates that the reaction is more favorable to reduction as the concentration of Cu2+ ions increases. Reduction will take place in the cell's compartment where concentration is higher and oxidation will occur on the more dilute side.
The following cell diagram describes the cell mentioned above:
Cu(s) | Cu2+ (0.05 M) || Cu2+ (2.0 M) | Cu(s)
Where the half cell reactions for oxidation and reduction are:
Oxidation: Cu(s) → Cu2+ (0.05 M) + 2 e–
Reduction: Cu2+ (2.0 M) + 2 e– → Cu(s)
Overall reaction: Cu2+ (2.0 M) → Cu2+ (0.05 M)
The cell's emf is calculated through Nernst equation as follows:

The value of E° in this kind of cell is zero, as electrodes and ions are the same in both half-cells.
After replacing values from the case mentioned, it is possible to calculate cell's potential:

or by:

However, this value is only approximate, as reaction quotient is defined in terms of ion activities which can be approximated with the concentrations as calculated here.
The Nernst equation plays an important role in understanding electrical effects in cells and organelles. Such effects include nerve synapses and cardiac beat as well as the resting potential of a somatic cell.
Battery
Many types of battery have been commercialized and represent an important practical application of electrochemistry. Early wet cells powered the first telegraph andtelephone systems, and were the source of current for electroplating. The zinc-manganese dioxide dry cell was the first portable, non-spillable battery type that madeflashlights and other portable devices practical. The mercury battery using zinc and mercuric oxide provided higher levels of power and capacity than the original dry cell for early electronic devices, but has been phased out of common use due to the danger of mercury pollution from discarded cells.
The lead acid battery was the first practical secondary (rechargeable) battery that could have its capacity replenished from an external source. The electrochemical reaction that produced current was (to a useful degree) reversible, allowing electrical energy and chemical energy to be interchanged as needed. Common lead acid batteries contain a mixture of acid and water, as well as lead plates. The most common mixture used today is 30% acid. One problem however is if left uncharged acid will crystallize within the lead plates of the battery rendering it useless. These batteries last an average of 3 years with daily use however it is not unheard of for a lead acid battery to still be functional after 7-10 years. Lead-acid cells continue to be widely used in automobiles.
All the preceding types have water-based electrolytes, which limits the maximum voltage per cell. The freezing of water limits low temperature performance. The lithium battery, which does not (and cannot) use water in the electrolyte, provides improved performance over other types; a rechargeable lithium ion battery is an essential part of many mobile devices.
The flow battery, an experimental type, offers the option of vastly larger energy capacity because its reactants can be replenished from external reservoirs. The fuel cell can turn the chemical energy bound in hydrocarbon gases or hydrogen directly into electrical energy with much higher efficiency than any combustion process; such devices have powered many spacecraft and are being applied to grid energy storage for the public power system.
Corrosion
Main article: Corrosion
Corrosion is the term applied to steel rust caused by an electrochemical process. Most people are likely familiar with the corrosion of iron, in the form of reddish rust. Other examples include the black tarnish on silver, and red or green corrosion that may appear on copper and its alloys, such as brass. The cost of replacing metals lost to corrosion is in the multi-billions of dollars per year.
For iron rust to occur the metal has to be in contact with oxygen and water, although chemical reactions for this process are relatively complex and not all of them are completely understood, it is believed the causes are the following: Electron transferring (reduction-oxidation)
One area on the surface of the metal acts as the anode, which is where the oxidation (corrosion) occurs. At the anode, the metal gives up electrons.
Fe(s) → Fe2+(aq) + 2 e–
Electrons are transferred from iron reducing oxygen in the atmosphere into water on the cathode, which is placed in another region of the metal.
O2(g) + 4 H+(aq) + 4 e– → 2 H2O(l)
Global reaction for the process:
2 Fe(s) + O2(g) + 4 H+(aq) → 2 Fe2+(aq) + 2 H2O(l)
Standard emf for iron rusting:
E° = E°cathode – E°anode
E° = 1.23V – (−0.44 V) = 1.67 V
Iron corrosion takes place on acid medium; H+ ions come from reaction between carbon dioxide in the atmosphere and water, forming carbonic acid. Fe2+ ions oxides, following this equation:
4 Fe2+(aq) + O2(g) + (4+2x)H2O(l) → 2 Fe2O3·xH2O + 8 H+(aq)
Iron(III) oxide hydrated is known as rust. The concentration of water associated with iron oxide varies, thus chemical representation is presented as Fe2O3·xH2O. The electric circuit works as passage of electrons and ions occurs, thus if an electrolyte is present it will facilitate oxidation, this explains why rusting is quicker on salt water.
Coinage metals, such as copper and silver, slowly corrode through use. A patina of green-blue copper carbonate forms on the surface of copper with exposure to the water and carbon dioxide in the air. Silver coins or cutlery that are exposed to high sulfur foods such as eggs or the low levels of sulfur species in the air develop a layer of blackSilver sulfide.
Gold and platinum are extremely difficult to oxidize under normal circumstances, and require exposure to a powerful chemical oxidizing agent such as aqua regia.
Some common metals oxidize extremely rapidly in air. Titanium and aluminium oxidize instantaneously in contact with the oxygen in the air. These metals form an extremely thin layer of oxidized metal on the surface. This thin layer of oxide protects the underlying layers of the metal from the air preventing the entire metal from oxidizing. These metals are used in applications where corrosion resistance is important. Iron, in contrast, has an oxide that forms in air and water, called rust, that does not stop the further oxidation of the iron. Thus iron left exposed to air and water will continue to rust until all of the iron is oxided.
Attempts to save a metal from becoming anodic are of two general types. Anodic regions dissolve and destroy the structural integrity of the metal.
While it is almost impossible to prevent anode/cathode formation, if a non-conducting material covers the metal, contact with the electrolyte is not possible and corrosion will not occur.
Coating
Metals can be coated with paint or other less conductive metals (passivation). This prevents the metal surface from being exposed to electrolytes. Scratches exposing the metal substrate will result in corrosion. The region under the coating adjacent to the scratch acts as the anode of the reaction.
Sacrificial anodes
A method commonly used to protect a structural metal is to attach a metal which is more anodic than the metal to be protected. This forces the structural metal to be cathodic, thus spared corrosion. It is called "sacrificial" because the anode dissolves and has to be replaced periodically.
Zinc bars are attached to various locations on steel ship hulls to render the ship hull cathodic. The zinc bars are replaced periodically. Other metals, such as magnesium, would work very well but zinc is the least expensive useful metal.
To protect pipelines, an ingot of buried or exposed magnesium (or zinc) is buried beside the pipeline and is connected electrically to the pipe above ground. The pipeline is forced to be a cathode and is protected from being oxidized and rusting. The magnesium anode is sacrificed. At intervals new ingots are buried to replace those lost.
Electrolysis
The spontaneous redox reactions of a conventional battery produce electricity through the different chemical potentials of the cathode and anode in the electrolyte. However, electrolysis requires an external source of electrical energy to induce a chemical reaction, and this process takes place in a compartment called an electrolytic cell.
When molten, the salt sodium chloride can be electrolyzed to yield metallic sodium and gaseous chlorine. Industrially this process takes place in a special cell named Down's cell. The cell is connected to an electrical power supply, allowing electrons to migrate from the power supply to the electrolytic cell.
Reactions that take place at Down's cell are the following:
Anode (oxidation): 2 Cl– → Cl2(g) + 2 e–
Cathode (reduction): 2 Na+(l) + 2 e– → 2 Na(l)
Overall reaction: 2 Na+ + 2 Cl–(l) → 2 Na(l) + Cl2(g)
This process can yield large amounts of metallic sodium and gaseous chlorine, and is widely used on mineral dressing and metallurgy industries.
The emf for this process is approximately −4 V indicating a (very) non-spontaneous process. In order for this reaction to occur the power supply should provide at least a potential of 4 V. However, larger voltages must be used for this reaction to occur at a high rate.
Water can be converted to its component elemental gasses, H2 and O2 through the application of an external voltage. Water doesn't decompose into hydrogen and oxygen spontaneously as the Gibbs free energy for the process at standard conditions is about 474.4 kJ. The decomposition of water into hydrogen and oxygen can be performed in an electrolytic cell. In it, a pair of inert electrodes usually made of platinum immersed in water act as anode and cathode in the electrolytic process. The electrolysis starts with the application of an external voltage between the electrodes. This process will not occur except at extremely high voltages without an electrolyte such as sodium chloride orsulfuric acid (most used 0.1 M).
Bubbles from the gases will be seen near both electrodes. The following half reactions describe the process mentioned above:
Anode (oxidation): 2 H2O(l) → O2(g) + 4 H+(aq) + 4 e–
Cathode (reduction): 2 H2O(g) + 2 e– → H2(g) + 2 OH–(aq)
Overall reaction: 2 H2O(l) → 2 H2(g) + O2(g)
Although strong acids may be used in the apparatus, the reaction will not net consume the acid. While this reaction will work at any conductive electrode at a sufficiently large potential, platinum catalyzes both hydrogen and oxygen formation, allowing for relatively mild voltages (~2 V depending on the pH).
Electrolysis in an aqueous is a similar process as mentioned in electrolysis of water. However, it is considered to be a complex process because the contents in solution have to be analyzed in half reactions, whether reduced or oxidized.
Electrolysis of a solution of sodium chloride
The presence of water in a solution of sodium chloride must be examined in respect to its reduction and oxidation in both electrodes. Usually, water is electrolysed as mentioned in electrolysis of water yielding gaseous oxygen in the anode and gaseous hydrogen in the cathode. On the other hand, sodium chloride in water dissociates in Na+and Cl– ions, cation, which is the positive ion, will be attracted to the cathode (-), thus reducing the sodium ion. The anion will then be attracted to the anode (+) oxidizingchloride ion.
The following half reactions describes the process mentioned:
1. Cathode: Na+(aq) + e– → Na(s) E°red = –2.71 V
2. Anode: 2 Cl–(aq) → Cl2(g) + 2 e– E°red = +1.36 V
3. Cathode: 2 H2O(l) + 2 e– → H2(g) + 2 OH–(aq) E°red = –0.83 V
4. Anode: 2 H2O(l) → O2(g) + 4 H+(aq) + 4 e– E°red = +1.23 V
Reaction 1 is discarded as it has the most negative value on standard reduction potential thus making it less thermodynamically favorable in the process.
When comparing the reduction potentials in reactions 2 and 4, the reduction of chloride ion is favored. Thus, if the Cl– ion is favored for reduction, then the water reaction is favored for oxidation producing gaseous oxygen, however experiments show gaseous chlorine is produced and not oxygen.
Although the initial analysis is correct, there is another effect that can happen, known as the overvoltage effect. Additional voltage is sometimes required, beyond the voltage predicted by the E°cell. This may be due to kinetic rather than thermodynamic considerations. In fact, it has been proven that the activation energy for the chloride ion is very low, hence favorable in kinetic terms. In other words, although the voltage applied is thermodynamically sufficient to drive electrolysis, the rate is so slow that to make the process proceed in a reasonable time frame, the voltage of the external source has to be increased (hence, overvoltage).
Finally, reaction 3 is favorable because it describes the proliferation of OH– ions thus letting a probable reduction of H+ ions less favorable an option.
The overall reaction for the process according to the analysis would be the following:
Anode (oxidation): 2 Cl–(aq) → Cl2(g) + 2 e–
Cathode (reduction): 2 H2O(l) + 2 e– → H2(g) + 2 OH–(aq)
Overall reaction: 2 H2O + 2 Cl–(aq) → H2(g) + Cl2(g) + 2 OH–(aq)
As the overall reaction indicates, the concentration of chloride ions is reduced in comparison to OH– ions (whose concentration increases). The reaction also shows the production of gaseous hydrogen, chlorine and aqueous sodium hydroxide.
Quantitative aspects of electrolysis were originally developed by Michael Faraday in 1834. Faraday is also credited to have coined the terms electrolyte, electrolysis, among many others while he studied quantitative analysis of electrochemical reactions. Also he was an advocate of the law of conservation of energy.
First law
Faraday concluded after several experiments on electrical current in non-spontaneous process, the mass of the products yielded on the electrodes was proportional to the value of current supplied to the cell, the length of time the current existed, and the molar mass of the substance analyzed. In other words, the amount of a substance deposited on each electrode of an electrolytic cell is directly proportional to the quantity of electricity passed through the cell.[28]
Below is a simplified equation of Faraday's first law:

Where
m is the mass of the substance produced at the electrode (in grams),
Q is the total electric charge that passed through the solution (in coulombs),
n is the valence number of the substance as an ion in solution (electrons per ion),
M is the molar mass of the substance (in grams per mole).
Second law
Faraday devised the laws of chemical electrodeposition of metals from solutions in 1857. He formulated the second law of electrolysis stating "the amounts of bodies which are equivalent to each other in their ordinary chemical action have equal quantities of electricity naturally associated with them." In other words, the quantities of different elements deposited by a given amount of electricity are in the ratio of their chemical equivalent weights.
An important aspect of the second law of electrolysis is electroplating which together with the first law of electrolysis, has a significant number of applications in the industry, as when used to protect metals to avoid corrosion.
Applications
There are various extremely important electrochemical processes in both nature and industry, like the coating of objects with metals or metal oxides through electrodeposition and the detection of alcohol in drunken drivers through the redox reaction of ethanol. The generation of chemical energy through photosynthesis is inherently an electrochemical process, as is production of metals like aluminum and titanium from their ores. Certain diabetes blood sugar meters measure the amount of glucose in the blood through its redox potential.
The action potentials that travel down neurons are based on electric current generated by the movement of sodium and potassium ions into and out of cells. Specialized cells in certain animals like the electric eel can generate electric currents powerful enough to disable much larger animals.
Chemical Reactivity Worksheet:
The Chemical Reactivity Worksheet (CRW) is a free software program you can use to find out about the chemical reactivity of thousands of common hazardous chemicals.
Download the Chemical Reactivity Worksheet for Windows, Mac, or iPad.
Reactivity is the tendency of substances to undergo chemical change, which can result in hazards—such as heat generation or toxic gas byproducts. The CRW predicts possible hazards from mixing chemicals and is designed to be used by emergency responders and planners, as well as the chemical industry, to help prevent dangerous chemical incidents.
The chemical datasheets in the CRW database contain information about the intrinsic hazards of each chemical and about whether a chemical reacts with air, water, or other materials. It also includes case histories on specific chemical incidents, with references.
You can also create your own custom chemical datasheets, as you might do, for instance, if your facility produces a proprietary chemical that is not in the CRW database.
The CRW also includes a reactivity prediction worksheet that you use to virtually "mix" chemicals to simulate accidental chemical mixtures, such as in the case of a train derailment, to learn what dangers could arise from the accidental mixing. For example, if the reaction is predicted to generate gases, the CRW will list the potential gaseous products, along with literature citations related to the prediction.

The Mixture Manager screen allows you to search for chemicals in the CRW's database, see a preview of the information on the chemical datasheets, and create a virtual mixture of chemicals. You can also access all the other features of the program from this screen, including the Compatibility Chart and hazards report for any mixture you create, reference information about the reactive groups used in the CRW, and information about absorbent incompatibilities with certain chemicals.

The Compatibility Chart shows the predicted hazards of mixing the chemicals in a mixture in an easy-to-use graphical interface. The reactivity predictions are color coded, and the cells on the chart can be clicked to find more information about specific predicted reactions. General hazard statements, predicted gas products, and literature documentation for the selected pair of chemicals are shown at the bottom of the chart.
More Information about the Chemical Reactivity Worksheet
NOAA has partnered with chemical industry experts to release a major update to the CRW (CRW version 3.0). This newest version of the program is the result of a two-year-long collaboration between NOAA chemical response specialists and expert chemists from the Dow Chemical Company.
All versions of the Chemical Reactivity Worksheet (CRW) were developed by NOAA's Office of Response and Restoration, in collaboration with the Environmental Protection Agency and the Center for Chemical Process Safety.
Chemical Reactivity Worksheet Fact Sheet [PDF, 967 KB]: Learn more about basic features of the CRW in this short fact sheet.
CAMEO Software Suite: Find out more about the several other programs in the CAMEO software suite related to the Chemical Reactivity Worksheet.

What Is Chemical Reactivity

Chemical reactivity refers to the ease with which a chemical compound tends to react when it comes into contact with other compounds. An element or a compound will be said to have high reactivity if it reacts fast when it interact with another substance. Such a compound may also be viewed as having high reactivity in terms of how many substances it can react with. Some of the common substances used to determine the chemical reactivity of elements include water, hydrochloric acid, sulphuric acid, chlorine gas, pure oxygen, air and many others.
Reactivity can be influenced by several internal and external factors. Some of these factors include temperature, pressure, surface area, amount of the substance and catalysts. Most substances react faster in high temperatures. However, there are substances whose reactivity will increase with decrease in temperatures. The same case applies to pressure. Reducing the size of particles in a substance increases its surface area with which it comes into contact with a reagent increasing its activity. Catalysts also enhance the reaction process. The process may release or absorb heat depending on the elements involved.
This process is very important in the study of science. In construction of the periodic table, reactivity is one of the main guiding factors. Scientists also came up with the reactivity series which place elements in a certain order depending on the stability of the elements. Some of the most reactive elements include potassium, sodium and fluorine. Noble gases are considered to be stable and hence most uncreative.
Reactivity of chemicals has very many applications in the world today. In medicine, elements are combined under controlled environment to form compounds that are used for treatment of diseases. A simple example is the milk of magnesia (magnesium hydroxide) used to treat stomach upset caused by excess acidity in the stomach. Making of weapons is another area that relies heavily on chemical reactivity. Most of the technology used in making weapons depends on the ability of some compounds to react vigorously and explode when a spark or some heat is introduced. Biological weapons also involve chemical compounds reacting to form poisonous substances. These may be in powder, liquid or gaseous state. An example of a compound used in making of ammunition is potassium cyanide.
The process is also very useful in the manufacturing industry where many consumer goods are made using the same technology. Understanding chemical reactivity is also very important in preventing and managing disasters. For instance, nuclear reactors are potentially very dangerous with the ability to cause mass destruction and death if not handled carefully. Such a disaster occurred in Chernobyl in Ukraine and is considered the worst chemical disaster in the world. It emitted radio active particles in the atmosphere which spread in a very large region of Europe and USSR.A more recent catastrophe involving nuclear reactors happened in Fukushima which caused death and its effects are still being felt today. However, nuclear reactors are very useful source of electric power in some countries. They produce large amount of power which cannot be produced through other means. They are therefore important but have to be handled carefully.
OR
see detail:
Organic chemistry encompasses a very large number of compounds ( many millions ), and our previous discussion and illustrations have focused on their structural characteristics. Now that we can recognize these actors ( compounds ), we turn to the roles they are inclined to play in the scientific drama staged by the multitude of chemical reactions that define organic chemistry.
We begin by defining some basic terms that will be used frequently as this subject is elaborated.
Chemical Reaction: A transformation resulting in a change of composition, constitution and/or configuration of a compound ( referred to as the reactant or substrate ).
Reactant or Substrate: The organic compound undergoing change in a chemical reaction. Other compounds may also be involved, and common reactive partners ( reagents ) may be identified. The reactant is often ( but not always ) the larger and more complex molecule in the reacting system. Most ( or all ) of the reactant molecule is normally incorporated as part of the product molecule.
Reagent: A common partner of the reactant in many chemical reactions. It may be organic or inorganic; small or large; gas, liquid or solid. The portion of a reagent that ends up being incorporated in the product may range from all to very little or none.
Product(s) The final form taken by the major reactant(s) of a reaction.
Reaction Conditions The environmental conditions, such as temperature, pressure, catalysts & solvent, under which a reaction progresses optimally. Catalysts are substances that accelerate the rate ( velocity ) of a chemical reaction without themselves being consumed or appearing as part of the reaction product. Catalysts do not change equilibria positions.

Chemical reactions are commonly written as equations:

Classifying Organic Chemical Reactions

If you scan any organic textbook you will encounter what appears to be a very large, often intimidating, number of reactions. These are the "tools" of a chemist, and to use these tools effectively, we must organize them in a sensible manner and look for patterns of reactivity that permit us make plausible predictions. Most of these reactions occur at special sites of reactivity known as functional groups, and these constitute one organizational scheme that helps us catalog and remember reactions.
Ultimately, the best way to achieve proficiency in organic chemistry is to understand how reactions take place, and to recognize the various factors that influence their course.
This is best accomplished by perceiving the reaction pathway or mechanism of a reaction.

1. Classification by Structural Change

First, we identify four broad classes of reactions based solely on the structural change occurring in the reactant molecules. This classification does not require knowledge or speculation concerning reaction paths or mechanisms.
The letter R in the following illustrations is widely used as a symbol for a generic group. It may stand for simple substituents such as H– or CH3–, or for complex groups composed of many atoms of carbon and other elements.

Four Reaction Classes

Addition
Elimination

Substitution
Rearrangement

In an addition reaction the number of σ-bonds in the substrate molecule increases, usually at the expense of one or more π-bonds. The reverse is true of elimination reactions, i.e.the number of σ-bonds in the substrate decreases, and new π-bonds are often formed. Substitution reactions, as the name implies, are characterized by replacement of an atom or group (Y) by another atom or group (Z). Aside from these groups, the number of bonds does not change. A rearrangement reaction generates an isomer, and again the number of bonds normally does not change.
The examples illustrated above involve simple alkyl and alkene systems, but these reaction types are general for most functional groups, including those incorporating carbon-oxygen double bonds and carbon-nitrogen double and triple bonds. Some common reactions may actually be a combination of reaction types. The reaction of an ester with ammonia to give an amide, as shown below, appears to be a substitution reaction ( Y = CH3O & Z = NH2 ); however, it is actually two reactions, an addition followed by an elimination.

The addition of water to a nitrile does not seem to fit any of the above reaction types, but it is simply a slow addition reaction followed by a rapid rearrangement, as shown in the following equation. Rapid rearrangements of this kind are called tautomerizations.

2. Classification by Reaction Type
At the beginning, it is helpful to identify some common reaction types that will surface repeatedly as the chemical behavior of different compounds is examined. This is not intended to be a complete and comprehensive list, but should set the stage for future elaborations.
It is useful to begin a discussion of organic chemical reactions with a review of acid-base chemistry and terminology for several reasons. First, acid-base reactions are among the simplest to recognize and understand. Second, some classes of organic compounds have distinctly acidic properties, and some other classes behave as bases, so we need to identify these aspects of their chemistry. Finally, many organic reactions are catalyzed by acids and/or bases, and although such transformations may seem complex, our understanding of how they occur often begins with the functioning of the catalyst.
Organic chemists use two acid-base theories for interpreting and planning their work: the Brønsted theory and the Lewis theory.
Brønsted Theory
According to the Brønsted theory, an acid is a proton donor, and a base is a proton acceptor. In an acid-base reaction, each side of the equilibrium has an acid and a base reactant or product, and these may be neutral species or ions.
H-A + B:(–)

A:(–) + B-H
(acid1) (base1) (base2) (acid2)
Structurally related acid-base pairs, such as {H-A and A:(–)} or {B:(–) and B-H} are called conjugate pairs. Substances that can serve as both acids and bases, such as water, are termed amphoteric.
H-Cl + H2O

Cl:(–) + H3O(+)
(acid) (base) (base) (acid)

H3N: + H2O

NH4(+) + HO(–)
(base) (acid) (acid) (base)
The relative strength of a group of acids (or bases) may be evaluated by measuring the extent of reaction that each group member undergoes with a common base (or acid). Water serves nicely as the common base or acid for such determinations. Thus, for an acid H-A, its strength is proportional to the extent of its reaction with the base water, which is given by the equilibrium constant Keq.

H-A + H2O

H3O(+) + A:(–)

Since these studies are generally extrapolated to high dilution, the molar concentration of water (55.5) is constant and may be eliminated from the denominator. The resulting K value is called the acidity constant, Ka. Clearly, strong acids have larger Ka's than do weaker acids. Because of the very large range of acid strengths (greater than 1040), a logarithmic scale of acidity (pKa) is normally employed. Stronger acids have smaller or more negative pKa values than do weaker acids.

Examples of Brønsted Acid-Base Equilibria
Acid-Base Reaction Conjugate
AcidsConjugate
BasesKapKa
HBr + H2O

H3O(+) + Br(–) HBr
H3O(+) Br(–)
H2O 105 -5
CH3CO2H + H2O

H3O(+) + CH3CO2(–) CH3CO2H
H3O(+) CH3CO2(–)
H2O 1.77*10-5 4.75
C2H5OH + H2O

H3O(+) + C2H5O(–) C2H5OH
H3O(+) C2H5O(–)
H2O 10-16 16
NH3 + H2O

H3O(+) + NH2(–) NH3
H3O(+) NH2(–)
H2O 10-34 34
In all the above examples water acts as a common base. The last example ( NH3 ) cannot be measured directly in water, since the strongest base that can exist in this solvent is hydroxide ion. Consequently, the value reported here is extrapolated from measurements in much less acidic solvents, such as acetonitrile.
Since many organic reactions either take place in aqueous environments ( living cells ), or are quenched or worked-up in water, it is important to consider how a conjugate acid-base equilibrium mixture changes with pH. A simple relationship known as the Henderson-Hasselbalch equation provides this information.

When the pH of an aqueous solution or mixture is equal to the pKa of an acidic component, the concentrations of the acid and base conjugate forms must be equal ( the log of 1 is 0 ). If the pH is lowered by two or more units relative to the pKa, the acid concentration will be greater than 99%. On the other hand, if the pH ( relative to pKa ) is raised by two or more units the conjugate base concentration will be over 99%. Consequently, mixtures of acidic and non-acidic compounds are easily separated by adjusting the pH of the water component in a two phase solvent extraction.
For example, if a solution of benzoic acid ( pKa = 4.2 ) in benzyl alcohol ( pKa = 15 ) is dissolved in ether and shaken with an excess of 0.1 N sodium hydroxide ( pH = 13 ), the acid is completely converted to its water soluble ( ether insoluble ) sodium salt, while the alcohol is unaffected. The ether solution of the alcohol may then be separated from the water layer, and pure alcohol recovered by distillation of the volatile ether solvent. The pH of the water solution of sodium benzoate may then be lowered to 1.0 by addition of hydrochloric acid, at which point pure benzoic acid crystallizes, and may be isolated by filtration.
Basicity
The basicity of oxygen, nitrogen, sulfur and phosphorus compounds or ions may be treated in an analogous fashion. Thus, we may write base-acid equilibria, which define a Kb and a corresponding pKb. However, a more common procedure is to report the acidities of the conjugate acids of the bases ( these conjugate acids are often "onium" cations ). The pKa's reported for bases in this system are proportional to the base strength of the base. A useful rule here is: pKa + pKb = 14.
We see this relationship in the following two equilibria:
Acid-Base Reaction Conjugate
AcidsConjugate
BasesKpK
NH3 + H2O

NH4(+) + OH(–) NH4(+)
H2O NH3
OH(–) Kb = 1.8*10-5 pKb = 4.74
NH4(+) + H2O

H3O(+) + NH3 NH4(+)
H3O(+) NH3
H2O Ka = 5.5*10-10 pKa = 9.25
Tables of pKa values for inorganic and organic acids ( and bases) are available in many reference books, and may be examined here by clicking on the appropriate link:
Inorganic Acidity Constants
Organic Acidity Constants
Basicity Constants
Although it is convenient and informative to express pKa values for a common solvent system (usually water), there are serious limitations for very strong and very weak acids. Thus acids that are stronger than the hydronium cation, H3O(+), and weak acids having conjugate bases stronger than hydroxide anion, OH(–), cannot be measured directly in water solution. Solvents such as acetic acid, acetonitrile and nitromethane are often used for studying very strong acids. Relative acidity measurements in these solvents may be extrapolated to water. Likewise, very weakly acidic solvents such as DMSO, acetonitrile, toluene, amines and ammonia may be used to study the acidities of very weak acids. For both these groups, the reported pKa values extrapolated to water are approximate, and many have large uncertainties.
Lewis Theory
According to the Lewis theory, an acid is an electron pair acceptor, and a base is an electron pair donor. Lewis bases are also Brønsted bases; however, many Lewis acids, such as BF3, AlCl3and Mg2+, are not Brønsted acids. The product of a Lewis acid-base reaction, is a neutral, dipolar or charged complex, which may be a stable covalent molecule. As shown at the top of the following drawing, coordinate covalent bonding of a phosphorous Lewis base to a boron Lewis acid creates a complex in which the formal charge of boron is negative and that of phosphorous is positive. In this complex, boron acquires a neon valence shell configuration and phosphorous an argon configuration. If the substituents (R) on these atoms are not large, the complex will be favored at equilibrium. However, steric hindrance of bulky substituents may prohibit complex formation. The resulting mixture of non-bonded Lewis acid/base pairs has been termed "frustrated", and exhibits unusual chemical behavior.
Two examples of Lewis acid-base equilibria that play a role in chemical reactions are shown in equations 1 & 2 below.

In the first example, an electron deficient aluminum atom bonds to a covalent chlorine atom by sharing one of its non-bonding valence electron pairs, and thus achieves an argon-like valence shell octet. Because this sharing is unilateral .

A terminology related to the Lewis acid-base nomenclature is often used by organic chemists. Here the term electrophile corresponds to a Lewis acid, and nucleophile corresponds to a Lewis base.
Electrophile: An electron deficient atom, ion or molecule that has an affinity for an electron pair, and will bond to a base or nucleophile.
Nucleophile: An atom, ion or molecule that has an electron pair that may be donated in bonding to an electrophile (or Lewis acid).
(chlorine contributes both electrons), both the aluminum and the chlorine have formal charges, as shown. If the carbon chlorine bond in this complex breaks with both the bonding electrons remaining with the more electronegative atom (chlorine), the carbon assumes a positive charge. We refer to such carbon species as carbocations. Carbocations are also Lewis acids, as the reverse reaction demonstrates.
Many carbocations (but not all) may also function as Brønsted acids. Equation 3 illustrates this dual behavior; the Lewis acidic site is colored red and three of the nine acidic hydrogen atoms are colored orange. In its Brønsted acid role the carbocation donates a proton to the base (hydroxide anion), and is converted to a stable neutral molecule having a carbon-carbon double bond.
A parallel and independent method of characterizing organic reactions is by oxidation-reduction terminology. Carbon atoms may have any oxidation state from –4 (e.g. CH4 ) to +4 (e.g. CO2 ), depending upon their substituents. Fortunately, we need not determine the absolute oxidation state of each carbon atom in a molecule, but only the change in oxidation state of those carbons involved in a chemical transformation. To determine whether a carbon atom has undergone a redox change during a reaction we simply note any changes in the number of bonds to hydrogen and the number of bonds to more electronegative atoms such as O, N, F, Cl, Br, I, & S that has occurred. Bonds to other carbon atoms are ignored. This count should be conducted for each carbon atom undergoing any change during a reaction.
If the number of hydrogen atoms bonded to a carbon increases, and/or if the number of bonds to more electronegative atoms decreases, the carbon in question has been reduced (i.e. it is in a lower oxidation state).
If the number of hydrogen atoms bonded to a carbon decreases, and/or if the number of bonds to more electronegative atoms increases, the carbon in question has been oxidized (i.e. it is in a higher oxidation state).
If there has been no change in the number of such bonds, then the carbon in question has not changed its oxidation state. In the hydrolysis reaction of a nitrile shown above, the blue colored carbon has not changed its oxidation state.
These rules are illustrated by the following four addition reactions involving the same starting material, cyclohexene. Carbon atoms colored blue are reduced, and those colored red are oxidized. In the addition of hydrogen both carbon atoms are reduced, and the overall reaction is termed a reduction. Peracid epoxidation and addition of bromine oxidize both carbon atoms, so these are termed oxidation reactions. Addition of HBr reduces one of the double bond carbon atoms and oxidizes the other; consequently, there is no overall redox change in the substrate molecule.

Since metals such as lithium and magnesium are less electronegative than hydrogen, their covalent bonds to carbon are polarized so that the carbon is negative (reduced) and the metal is positive (oxidized). Thus, Grignard reagent formation from an alkyl halide reduces the substituted carbon atom. In the following equation and half-reactions the carbon atom (blue) is reduced and the magnesium (magenta) is oxidized.

3. Classification by Functional Group
Functional groups are atoms or small groups of atoms (usually two to four) that exhibit a characteristic reactivity when treated with certain reagents
. A particular functional group will almost always display its characteristic chemical behavior when it is present in a compound. Because of this, the discussion of organic reactions is often organized according to functional groups. The following table summarizes the general chemical behavior of the common functional groups. For reference, the alkanes provide a background of behavior in the absence of more localized functional groups.
Functional ClassFormulaCharacteristic Reactions
Alkanes C–C, C–H Substitution (of H, commonly by Cl or Br)
Combustion (conversion to CO2 & H2O)
Alkenes C=C–C–H Addition
Substitution (of H)
Alkynes C≡C–H Addition
Substitution (of H)
Alkyl Halides H–C–C–X Substitution (of X)
Elimination (of HX)
Alcohols H–C–C–O–H Substitution (of H); Substitution (of OH)
Elimination (of HOH); Oxidation (elimination of 2H)
Ethers (α)C–O–R Substitution (of OR); Substitution (of α–H)
Amines C–NRH Substitution (of H);
Addition (to N); Oxidation (of N)
Benzene Ring C6H6 Substitution (of H)
Aldehydes (α)C–CH=O Addition
Substitution (of H or α–H)
Ketones (α)C–CR=O Addition
Substitution (of α–H)
Carboxylic Acids (α)C–CO2H Substitution (of H); Substitution (of OH)
Substitution (of α–H); Addition (to C=O)
Carboxylic Derivatives (α)C–CZ=O
(Z = OR, Cl, NHR, etc.) Substitution (of Z); Substitution (of α–H)
Addition (to C=O)
This table does not include any reference to rearrangement, due to the fact that such reactions are found in all functional classes, and are highly dependent on the structure of the reactant. Furthermore, a review of the overall reaction patterns presented in this table discloses only a broad and rather non-specific set of reactivity trends. This is not surprising, since the three remaining categories provide only a coarse discrimination (comparable to identifying an object as animal, vegetable or mineral). Consequently, apparent similarities may fail to reflect important differences. For example, addition reactions to C=C are significantly different from additions to C=O, and substitution reactions of C-X proceed in very different ways, depending on the hybridization state of carbon.
The Variables of Organic Reactions
In an effort to understand how and why reactions of functional groups take place in the way they do, chemists try to discover just how different molecules and ions interact with each other as they come together. To this end, it is important to consider the various properties and characteristics of a reaction that may be observed and/or measured as the reaction proceeds . The most common and useful of these are listed below:
1. Reactants and Reagents

A. Reactant Structure: Variations in the structure of the reactant may have a marked influence on the course of a reaction, even though the functional group is unchanged. Thus, reaction of 1-bromopropane with sodium cyanide proceeds smoothly to yield butanenitrile, whereas 1-bromo-2,2-dimethylpropane fails to give any product and is recovered unchanged. In contrast, both alkyl bromides form Grignard reagents (RMgBr) on reaction with magnesium.

B. Reagent Characteristics: Apparently minor changes in a reagent may lead to a significant change in the course of a reaction. For example, 2-bromopropane gives a substitution reaction with sodium methylthiolate but undergoes predominant elimination on treatment with sodium methoxide.

2. Product Selectivity

A. Regioselectivity: It is often the case that addition and elimination reactions may, in principle, proceed to more than one product. Thus 1-butene might add HBr to give either 1-bromobutane or 2-bromobutane, depending on which carbon of the double bond receives the hydrogen and which the bromine. If one possible product out of two or more is formed preferentially, the reaction is said to be regioselective.

Simple substitution reactions are not normally considered regioselective, since by definition only one constitutional product is possible. However, rearrangements are known to occur during some reactions.

B. Stereoselectivity: If the reaction products are such that stereoisomers may be formed, a reaction that yields one stereoisomer preferentially is said to be stereoselective. In the addition of bromine to cyclohexene, for example, cis and trans-1,2-dibromocyclohexane are both possible products of the addition. Since the trans-isomer is the only isolated product, this reaction is stereoselective.

C. Stereospecificity: This term is applied to cases in which stereoisomeric reactants behave differently in a given reaction. Examples include:(i) Formation of different stereoisomeric products, as in the reaction of enantiomeric 2-bromobutane isomers with sodium methylthiolate, shown in the following diagram.

Here, the (R)-reactant gives the configurationally inverted (S)-product, and (S)-reactant produces (R)-product. The (R) and (S) notations for configuration are described in a later sectionof this text.

(ii) Different rates of reaction, as in the base-induced elimination of cis & trans-4-tert-butylcyclohexyl bromide (equation 1 below).

(iii) Different reaction paths leading to different products, as in the base-induced elimination of cis & trans-2-methylcyclohexyl bromide (equation 2 below).

3. Reaction Characteristics

A. Reaction Rates: Some reactions proceed very rapidly, and some so slowly that they are not normally observed. Among the variables that influence reaction rates are temperature (reactions are usually faster at a higher temperature), solvent, and reactant / reagent concentrations. Useful information about reaction mechanisms may be obtained by studying the manner in which the rate of a reaction changes as the concentrations of the reactant and reagents are varied. This field of study is called kinetics.

B. Intermediates: Many reactions proceed in a stepwise fashion. This can be convincingly demonstrated if an intermediate species can be isolated and shown to proceed to the same products under the reaction conditions. Some intermediates are stable compounds in their own right; however, some are so reactive that isolation is not possible. Nevertheless, evidence for their existence may be obtained by other means, including spectroscopic observation or inference from kinetic results.
4. Factors that Influence Reactions
It is helpful to identify some general features of a reaction that have a significant influence on its facility. Some of the most important of these are:

A. Energetics: The potential energy of a reacting system changes as the reaction progresses. The overall change may be exothermic ( energy is released ) or endothermic ( energy must be added ), and there is usually an activation energy requirement as well. Tables of Standard Bond Energies are widely used by chemists for estimating the energy change in a proposed reaction. As a rule, compounds constructed of strong covalent bonds are more stable than compounds incorporating one or more relatively weak bonds.

B. Electronic Effects: The distribution of electrons at sites of reaction (functional groups) is a particularly important factor. Electron deficient species or groups, which may or may not be positively charged, are attracted to electron rich species or groups, which may or may not be negatively charged. We refer to these species as electrophiles & nucleophiles respectively. In general, opposites attract and like repel.
The charge distribution in a molecule is usually discussed with respect to two interacting effects: An inductive effect, which is a function of the electronegativity differences that exist between atoms (and groups); and a resonance effect, in which electrons move in a discontinuous fashion between parts of a molecule.

C. Steric Effects: Atoms occupy space. When they are crowded together, van der Waals repulsions produce an unfavorable steric hindrance. Steric hindrance may influenceconformational equilibria, as well as destabilizing transition states of reactions.

D. Stereoelectronic Effects: In many reactions atomic or molecular orbitals interact in a manner that has an optimal configurational or geometrical alignment. Departure from this alignment inhibits the reaction.

E. Solvent Effects: Most reactions are conducted in solution, not in a gaseous state. The solvent selected for a given reaction may exert a strong influence on its course. Remember, solvents are chemicals, and most undergo chemical reaction under the right conditions.
Mechanisms of Organic Reactions
A detailed description of the changes in structure and bonding that take place in the course of a reaction, and the sequence of such events is called the reaction mechanism. A reaction mechanism should include a representation of plausible electron reorganization, as well as the identification of any intermediate species that may be formed as the reaction progresses. These features are elaborated in the following sections.
1. The Arrow Notation in Mechanisms
Since chemical reactions involve the breaking and making of bonds, a consideration of the movement of bonding ( and non-bonding ) valence shell electrons is essential to this understanding. It is now common practice to show the movement of electrons with curved arrows, and a sequence of equations depicting the consequences of such electron shifts is termed a mechanism. In general, two kinds of curved arrows are used in drawing mechanisms:
A full head on the arrow indicates the movement or shift of an electron pair:

A partial head (fishhook) on the arrow indicates the shift of a single electron:

The use of these symbols in bond-breaking and bond-making reactions is illustrated below. If a covalent single bond is broken so that one electron of the shared pair remains with each fragment, as in the first example, this bond-breaking is called homolysis. If the bond breaks with both electrons of the shared pair remaining with one fragment, as in the second and third examples, this is calledheterolysis.
Bond-Breaking Bond-Making

Other Arrow Symbols
Chemists also use arrow symbols for other purposes, and it is essential to use them correctly.

The Reaction Arrow

The Equilibrium Arrow

The Resonance Arrow

The following equations illustrate the proper use of these symbols:

2. Reactive Intermediates
The products of bond breaking, shown above, are not stable in the usual sense, and cannot be isolated for prolonged study. Such species are referred to as reactive intermediates, and are believed to be transient intermediates in many reactions. The general structures and names of four such intermediates are given below.

A pair of widely used terms, related to the Lewis acid-base notation, should also be introduced here.

Electrophile: An electron deficient atom, ion or molecule that has an affinity for an electron pair, and will bond to a base or nucleophile.
Nucleophile: An atom, ion or molecule that has an electron pair that may be donated in bonding to an electrophile (or Lewis acid).
Using these definitions, it is clear that carbocations ( called carbonium ions in the older literature ) are electrophiles and carbanions are nucleophiles. Carbenes have only a valence shell sextet of electrons and are therefore electron deficient. In this sense they are electrophiles, but the non-bonding electron pair also gives carbenes nucleophilic character. As a rule, the electrophilic character dominates carbene reactivity. Carbon radicals have only seven valence electrons, and may be considered electron deficient; however, they do not in general bond to nucleophilic electron pairs, so their chemistry exhibits unique differences from that of conventional electrophiles. Radical intermediates are often called free radicals.
The importance of electrophile / nucleophile terminology comes from the fact that many organic reactions involve at some stage the bonding of a nucleophile to an electrophile, a process that generally leads to a stable intermediate or product. Reactions of this kind are sometimes called ionic reactions, since ionic reactants or products are often involved. Some common examples of ionic reactions and their mechanisms may be examined by Clicking Here
The shapes ideally assumed by these intermediates becomes important when considering the stereochemistry of reactions in which they play a role. A simple tetravalent compound like methane, CH4, has a tetrahedral configuration. Carbocations have only three bonds to the charge bearing carbon, so it adopts a planar trigonal configuration. Carbanions are pyramidal in shape ( tetrahedral if the electron pair is viewed as a substituent ), but these species invert rapidly at room temperature, passing through a higher energy planar form in which the electron pair occupies a p-orbital. Radicals are intermediate in configuration, the energy difference between pyramidal and planar forms being very small. Since three points determine a plane, the shape of carbenes must be planar; however, the valence electron distribution varies.
Analytical Chemistry
Inorganic chemistry

Physical chemistry

Vapor Pressure of Electrolyte Solutions

Molar Solubility and Relative Solubility

Dilutions of Solutions

Molarity

Heat of Solution and Dilution

Acid Dissosociation Constant (Ka)

The Effect of pH on Solubility

Strong Acid-Strong Base Titrations

Electrolytic Properties

Molality

Oxidation and reduction

Balancing redox reactions:

Nernst equation
Concentration cells
Iron corrosion:
Corrosion of common metals
Prevention of corrosion
Electrolysis of molten sodium chloride
Electrolysis of water
Electrolysis of aqueous solutions
Quantitative electrolysis and Faraday's laws

Chemical Reactivity
Acidity and Basicity

Oxidation and Reduction Reactions

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chemistry ,What Is Chemistry?A SIMPLE VIEW OF ATOMIC STRUCTURE, Atom and Electromagnetic Radiation,MOLECULAR STRUCTURES,Shape of Molecules,matter(gases,liquids,solids),Solutions,Chemical Reactivity,

The Shape of Molecules

Solutions and Dissolving

A solution has certain characteristics:

It is uniform, or homogeneous, throughout the mixture

It is stable and doesn't change over time or settle

The solute particles are so small they cannot be separated by filtering

The solute and solvent molecules cannot be distinguished by the naked eye

It does not scatter a beam of light

Solute - The solute is the substance that is being dissolved by another substance. In the example above, the salt is the solute.

Solvent - The solvent is the substance that dissolves the other substance. In the example above, the water is the solvent.

solution:

1 comment:

Pages

chemistry ,What Is Chemistry?A SIMPLE VIEW OF ATOMIC STRUCTURE, Atom and Electromagnetic Radiation,MOLECULAR STRUCTURES,Shape of Molecules,matter(gases,liquids,solids),Solutions,Chemical Reactivity,

The Shape of Molecules

Solutions and Dissolving

A solution has certain characteristics:

It is uniform, or homogeneous, throughout the mixture It is stable and doesn't change over time or settle The solute particles are so small they cannot be separated by filtering The solute and solvent molecules cannot be distinguished by the naked eye It does not scatter a beam of light

Solute - The solute is the substance that is being dissolved by another substance. In the example above, the salt is the solute. Solvent - The solvent is the substance that dissolves the other substance. In the example above, the water is the solvent.

solution:

1 comment:

It is uniform, or homogeneous, throughout the mixture

It is stable and doesn't change over time or settle

The solute particles are so small they cannot be separated by filtering

The solute and solvent molecules cannot be distinguished by the naked eye

It does not scatter a beam of light

Solute - The solute is the substance that is being dissolved by another substance. In the example above, the salt is the solute.

Solvent - The solvent is the substance that dissolves the other substance. In the example above, the water is the solvent.