The Electron

The Electron

Electron, is lightest stable subatomic particle known. It carries a negative charge of 1.602176634 × 10−19 coulomb, which is considered the basic unit of electric charge. The rest mass of the electron is 9.1093837015 × 10−31 kg, which is only 1/1,836the mass of a proton. An electron is therefore considered nearly mass-less in comparison with a proton or a neutron, and the electron mass is not included in calculating the mass number of an atom.

During the 1870s, the English chemist and physicist Sir William Crookes developed the first cathode ray tube to have a high vacuum inside. He then showed in 1874 that the cathode rays can turn a small paddle wheel when placed in their path. Therefore, he concluded that the rays carried momentum. Furthermore, by applying a magnetic field, he was able to deflect the rays, thereby demonstrating that the beam behaved as though it were negatively charged.

The German-born British physicist Arthur Schuster expanded upon Crookes' experiments by placing metal plates parallel to the cathode rays and applying an electric potential between the plates. The field deflected the rays toward the positively charged plate, providing further evidence that the rays carried negative charge.

In 1897, the British physicist J. J. Thomson, with his colleagues John S. Townsend and H. A. Wilson, performed experiments indicating that cathode rays really were unique particles, rather than waves, atoms or molecules as was believed earlier. Thomson made good estimates of both the charge e and the mass m, finding that cathode ray particles, which he called "corpuscles," had perhaps one thousandth of the mass of the least massive ion known: hydrogen. He showed that their charge-to-mass ratio, e/m, was independent of cathode material.

The electron's charge was more carefully measured by the American physicists Robert Millikan and Harvey Fletcher in their oil-drop experiment of 1909, the results of which were published in 1911.

By 1914, experiments by physicists Ernest Rutherford, Henry Moseley, James Franck and Gustav Hertz had largely established the structure of an atom as a dense nucleus of positive charge surrounded by lower-mass electrons. In 1913, Danish physicist Niels Bohr postulated that electrons resided in quantized energy states, with their energies determined by the angular momentum of the electron's orbit about the nucleus. The electrons could move between those states, or orbits, by the emission or absorption of photons of specific frequencies. By means of these quantized orbits, he accurately explained the spectral lines of the hydrogen atom.

Chemical bonds between atoms were explained by Gilbert Newton Lewis, who in 1916 proposed that a covalent bond between two atoms is maintained by a pair of electrons shared between them

In 1927 George Paget Thomson, discovered the interference effect was produced when a beam of electrons was passed through thin metal foils and by American physicists Clinton Davisson and Lester Germer by the reflection of electrons from a crystal of nickel.

Louis de Broglie hypothesized that under the appropriate conditions, electrons and other matter would show properties of either particles or waves.

De Broglie's prediction of a wave nature for electrons led Erwin Schrödinger to postulate a wave equation for electrons moving under the influence of the nucleus in the atom. In 1926, this equation, the Schrödinger equation, successfully described how electron waves propagated.

Solutions of Schrödinger's equation, like Heisenberg's, provided derivations of the energy states of an electron in a hydrogen atom that were equivalent to those that had been derived first by Bohr in 1913, and that were known to reproduce the hydrogen spectrum.

The wave-like behavior of a bound electron is described by a function called an atomic orbital. Each orbital has its own set of quantum numbers such as energy, angular momentum and projection of angular momentum, and only a discrete set of these orbitals exist around the nucleus. According to the Pauli exclusion principle each orbital can be occupied by up to two electrons, which must differ in their spin quantum number.

Electrons can transfer between different orbitals by the emission or absorption of photons with an energy that matches the difference in potential. Other methods of orbital transfer include collisions with particles, such as electrons, and the Auger effect. To escape the atom, the energy of the electron must be increased above its binding energy to the atom. This occurs, for example, with the photoelectric effect, where an incident photon exceeding the atom's ionization energy is absorbed by the electron.

The chemical bond between atoms occurs as a result of electromagnetic interactions, as described by the laws of quantum mechanics. The strongest bonds are formed by the sharing or transfer of electrons between atoms, allowing the formation of molecules. Within a molecule, electrons move under the influence of several nuclei, and occupy molecular orbitals; much as they can occupy atomic orbitals in isolated atoms. A fundamental factor in these molecular structures is the existence of electron pairs. These are electrons with opposed spins, allowing them to occupy the same molecular orbital without violating the Pauli exclusion principle (much like in atoms). Different molecular orbitals have different spatial distribution of the electron density. For instance, in bonded pairs (i.e. in the pairs that actually bind atoms together) electrons can be found with the maximal probability in a relatively small volume between the nuclei. By contrast, in non-bonded pairs electrons are distributed in a large volume around nuclei.

Since an electron has charge, it has a surrounding electric field, and if that electron is moving relative to an observer, said observer will observe it to generate a magnetic field. Electromagnetic fields produced from other sources will affect the motion of an electron according to the Lorentz force law. Electrons radiate or absorb energy in the form of photons when they are accelerated.

Evolution of early chemical science

Evolution of early chemical science

During the 19th and 20th century, this transformation was credited to the work of the French chemist Antoine Lavoisier (the "father of modern chemistry"). However, recent work on the history of early modern chemistry considers the chemical revolution to consist of gradual changes in chemical theory and practice that emerged over a period of two centuries. The so-called scientific revolution took place during the sixteenth and seventeenth centuries whereas the chemical revolution took place during the seventeenth and eighteenth centuries.

Several factors led to the first chemical revolution. First, there were the forms of gravimetric analysis that emerged from alchemy and new kinds of instruments that were developed in medical and industrial contexts. In these settings, chemists increasingly challenged hypotheses that had already been presented by the ancient Greeks. For example, chemists began to assert that all structures were composed of more than the four elements of the Greeks or the eight elements of the medieval alchemists. The Irish alchemist, Robert Boyle, laid the foundations for the Chemical Revolution, with his mechanical corpuscular philosophy, which in turn relied heavily on the alchemical corpuscular theory and experimental method dating back to pseudo-Geber.

Earlier works by chemists such as John Baptist van Helmont helped to shift the belief in theory that air existed as a single element to that of one in which air existed as a composition of a mixture of distinct kinds of gasses. John Baptist van Helmont’s data analysis also suggests that he had a general understanding of the law of conservation of mass in the 17 century. Furthermore, work by Jean Rey in the early 17 century with metals like tin and lead and their oxidation in the presence of air and water helped pinpoint the contribution and existence of oxygen in the oxidation process.

Other factors included new experimental techniques and the discovery of 'fixed air' (carbon dioxide) by Joseph Black in the middle of the 18th century. This discovery was particularly important because it empirically proved that 'air' did not consist of only one substance and because it established 'gas' as an important experimental substance. Nearer the end of the 18th century, the experiments by Henry Cavendish. Henry Cavendish was a British scientist noted for his discovery of hydrogen or what he called "inflammable air". He described the density of inflammable air, which formed water on combustion, in a 1766 paper "On Factitious Airs". Antoine Lavoisier later reproduced Cavendish's experiment and gave the element its name.

The latter stages of the revolution was fuelled by the 1789 publication of Lavoisier's Traité Élémentaire de Chimie (Elements of Chemistry). Beginning with this publication and others to follow, Lavoisier synthesised the work of others and coined the term "oxygen". Antoine Lavoisier represented the chemical revolution not only in his publications, but also in the way he practiced chemistry. Lavoisier's work was characterized by his systematic determination of weights and his strong emphasis on precision and accuracy. While it has been postulated that the law of conservation of mass was discovered by Lavoisier, this claim has been refuted by scientist Marcellin Berthelot. Earlier use of the law of conservation of mass has been suggested by Henry Guerlac, noting that scientist Jan Baptist van Helmont had implicitly applied the methodology to his work in the 16th and 17th centuries. Earlier references of the law of conservation of mass and its use were made by Jean Rey in 1630. Although the law of conservation of mass was not explicitly discovered by Lavoisier, his work with a wider array of materials than what most scientists had available at the time allowed his work to greatly expand the boundaries of the principal and its fundamentals.

Lavoisier also contributed to chemistry a method of understanding combustion and respiration and proof of the composition of water by decomposition into its constituent parts. He explained the theory of combustion, and challenged the phlogiston theory with his views on caloric. The Traité incorporates notions of a "new chemistry" and describes the experiments and reasoning that led to his conclusions. Like Newton's Principia, which was the high point of the Scientific Revolution, Lavoisier's Traité can be seen as the culmination of the Chemical Revolution.

Humphry Davy was an English chemist and a professor of chemistry at the London's Royal Institution in the early 1800's. There he performed experiments that cast doubt upon some of Lavoisier's key ideas such as the acidity of oxygen and the idea of a caloric element. Davy was able to show that acidity was not due to the presence of oxygen using muriatic acid (hydrochloric acid) as proof. He also proved that the compound oxymuriatic acid contained no oxygen and was instead an element, which he named chlorine. Through his use of electric batteries at the Royal Institution Davy first isolated chlorine, followed by the isolation of elemental iodine in 1813. Using the batteries Davy was also able to isolate the elements sodium and potassium. From these experiments Davy concluded that the forces that join chemical elements together must be electrical in nature. Davy was also a proponent against the idea that caloric was an immaterial fluid, arguing instead that heat was a type of motion

John Dalton was an English chemist that developed the idea of atomic theory of chemical elements. Dalton's atomic theory of chemical elements assumed that each element had unique atoms associated with and specific to that atom. This was in opposition to Lavoisier's definition of elements which was that elements are substances that chemists could not break down further into simpler parts. Dalton's idea also differed from the idea of corpuscular theory of matter, which believed that all atoms were the same, and had been a supported theory since the 17th century. To help support his idea, Dalton worked on defining the relative weights of atoms in chemicals in his work New System of Chemical Philosophy, published in 1808. His text showed calculations to determine the relative atomic weights of Lavoisier's different elements based on experimental data pertaining to the relative amounts of different elements in chemical combinations. Dalton argued that elements would combine in the simplest form possible. Water was known to be a combination of hydrogen and oxygen, thus Dalton believed water to be a binary compound containing one hydrogen and one oxygen.

Dalton was able to accurately compute the relative quantity of gases in atmospheric air. He used the specific gravity of azotic (nitrogen), oxygenous, carbonic acid (carbon dioxide), and hydrogenous gases as well as aqueous vapor determined by Lavoisier and Davy to determine the proportional weights of each as a percent of a whole volume of atmospheric air. Dalton determined that atmospheric air contains 75.55% azotic gas, 23.32% oxygenous gas, 1.03% aqueous vapor, and 0.10% carbonic acid gas.

Jöns Jacob Berzelius was a Swedish chemist who studied medicine at the Univseristy of Uppsala and was a professor of chemistry in Stockholm. He drew on the ideas of both Davy and Dalton to create an electrochemical view of how elements combined together. Berzelius classified elements into two groups, electronegative and electropositive depending which pole of a galvanic battery they were released from when decomposed. He created a scale of charge with oxygen being the most electronegative element and potassium the most electropositive. This scale signified that some elements had positive and negative charges associated with them and the position of an element on this scale and the element's charge determined how that element combined with others. Berzelius's work on electrochemical atomic theory was published in 1818 as Essai sur la théorie des proportions chimiques et sur l'influence chimique de l'électricité. He also introduced a new chemical nomenclature into chemistry by representing elements with letters and abbreviations, such as O for oxygen and Fe for iron. Combinations of elements were represented as sequences of these symbols and the number of atoms were represented at first by superscripts and then later subscripts.

The Chemical revolution

The Chemical revolution

The chemical revolution, also called the first chemical revolution, was the early modern reformulation of chemistry that culminated in the law of conservation of mass and the oxygen theory of combustion. During the 19th and 20th century, this transformation was credited to the work of the French chemist Antoine Lavoisier (the "father of modern chemistry").

However, recent work on the history of early modern chemistry considers the chemical revolution to consist of gradual changes in chemical theory and practice that emerged over a period of two centuries.

The Irish alchemist, Robert Boyle, laid the foundations for the Chemical Revolution, with his mechanical corpuscular philosophy. 

In 1756,Joseph Black heated lime stone and produced an air which he called fixed air. It did not support burning. Joseph Black demonstrated experimentally that the air fixed in certain reactions is chemically different from common air. Black wanted to know why slaked quicklime (hydrated calcium oxide) was neutralized when exposed to the atmosphere. He found that it absorbed only one component of the atmosphere, carbon dioxide, which he called “fixed air.” Black’s work marked the beginning of investigative efforts devoted to identifying chemically distinct airs, an area of research that grew rapidly during the latter half of the century.

This discovery was particularly important because it empirically proved that 'air' did not consist of only one substance and because it established 'gas' as an important experimental substance. Nearer the end of the 18th century, the experiments by Henry Cavendish and Joseph Priestley further proved that air is not an element and is instead composed of several different gases.

Antoine-Laurent Lavoisier, a meticulous experimenter, revolutionized chemistry. He established the law of conservation of mass, determined that combustion and respiration are caused by chemical reactions with what he named “oxygen,” and helped systematize chemical nomenclature, among many other accomplishments.

Lavoisier’s research in the early 1770s focused upon weight gains and losses in calcination. It was known that when metals slowly changed into powders (calxes), as was observed in the rusting of iron, the calx actually weighed more than the original metal, whereas when the calx was “reduced” to a metal, a loss of weight occurred.

Lavoisier also contributed to chemistry a method of understanding combustion and respiration and proof of the composition of water by decomposition into its constituent parts. He explained the theory of combustion, and challenged the phlogiston theory with his views on caloric.

In attempting to dismantle phlogiston theory and implement his own theory of combustion, Lavoisier utilized multiple apparatuses. These included a red-hot iron gun barrel which was designed to have water run through it and decompose, and an alteration of the apparatus which implemented a pneumatic trough at one end, a thermometer, and a barometer. The precision of his measurements was a requirement in convincing opposition of his theories about water as a compound, with instrumentation designed by himself implemented in his research.

Lavoisier is most noted for his discovery of the role oxygen plays in combustion. He recognized and named oxygen (1778) and hydrogen (1783), and opposed the phlogiston theory.

Lavoisier also translated the names of chemical substance into a new nomenclatural language more appealing to scientists of the nineteenth century. Such changes took place in an atmosphere in which the industrial revolution increased public interest in learning and practicing chemistry.

Much of the reasoning behind Antoine Lavoisier being named the "father of modern chemistry" and the start of the chemical revolution lay in his ability to mathematize the field, pushing chemistry to use the experimental methods utilized in other "more exact sciences." Lavoisier changed the field of chemistry by keeping meticulous balance sheets in his research, attempting to show that through the transformation of chemical species the total amount of substance was conserved. Lavoisier used instrumentation for thermometric and barometric measurements in his experiments, and collaborated with Pierre Simon de Laplace in the invention of the calorimeter, an instrument for measuring heat changes in a reaction.

Lavoisier's work was not immediately accepted and it took several decades for it gain momentum.This transition was aided by the work of Jöns Jakob Berzelius, who came up with a simplified shorthand to describe chemical compounds based on John Dalton's theory of atomic weights. Many people credit Lavoisier and his overthrow of phlogiston theory as the traditional chemical revolution, with Lavoisier marking the beginning of the revolution and John Dalton marking its culmination.

Antoine Lavoisier, in a collaborative effort with Louis Bernard Guyton de Morveau, Claude Louis Berthollet, and Antoine François de Fourcroy, published Méthode de nomenclature chimique in 1787. This work established a terminology for the "new chemistry" which Lavoisier was creating, which focused on a standardized set of terms, establishment of new elements, and experimental work. Méthode established 55 elements which were substances that could not be broken down into simpler composite parts at the time of publishing. By introducing new terminology into the field, Lavoisier encouraged other chemists to adopt his theories and practices in order to use his terms and stay current in chemistry.

To propagate his ideas, in 1789 he published a textbook, Traité élémentaire de chimie, and began a journal, Annales de Chimie, which carried research reports about the new chemistry almost exclusively.

In1803, Dalton gave the concept of atom. All matter is made up of atoms, he said. He made a list of relative atomic weights. Taking the weight of hydrogen as unity, he calculated the weights of oxygen, nitrogen, carbon and so on.

Proust gave the law of definite proportion. “Elements combine in definite proportion by mass, while forming compounds”, he said.

One of science’s great odd couples — British minister Joseph Priestley and French tax administrator Antoine Lavoisier — together discover a fantastic new gas called oxygen, overturning the reigning theory of chemistry and triggering a worldwide search for new elements. Soon caught up in the hunt is science’s first great showman, a precocious British chemist named Humphry Davy, who dazzles London audiences with his lectures, introduces them to laughing gas and turns the battery into a powerful tool in the search for new elements.

Humphry Davy established that Electricity is the fundamental property of matter.

Humphry Davy discovered new elements, using the Voltaic pile as electric energy to isolate elements of their natural compounds. The list of elements grew over the time. He wrote the book, the elements of Chemistry. The elements were symbolically represented by Jakob Berzelius and used symbols for writing chemical formulae.

Many doubts relating to chemical science were discussed in the first world chemical conference held in 1860 . A list of atomic weights was presented in the meet by Cannizzaro. The young Russian scientist Mendeleev attended this meet and later in1869, he constructed the periodic table of [known] elements. He arranged elements with their increasing atomic weights in rows and columns in his table of elements.

Thus modern chemistry flourished between 1756 and 1869; A century of chemical revolution, Chemistry as a fundamental Science of Matter was established.

Fundamentals of Chemistry

Fundamentals of Chemistry

Chemistry as a body of knowledge distinct from alchemy began to emerge when a clear differentiation was made between them by Robert Boyle in his work The Sceptical Chymist (1661). While both alchemy and chemistry are concerned with matter and its transformations, the crucial difference was given by the scientific method that chemists employed in their work. Chemistry is considered to have become an established science with the work of Antoine Lavoisier, who developed a law of conservation of mass that demanded careful measurement and quantitative observations of chemical phenomena. The history of chemistry is intertwined with the history of thermodynamics, especially through the work of Willard Gibbs.
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 in the 1860s by Dmitri Mendeleev and independently by several other scientists including Julius Lothar Meyer. 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.
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 University discovered 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 of chemical bonds and molecular orbitals was developed by the American scientists Linus Pauling and Gilbert N. Lewis.
Organic chemistry was developed by Justus von Liebig and others, following Friedrich Wöhler's synthesis of urea which proved that living organisms were, in theory, reducible to chemistry. Other crucial 19th century advances were; an understanding of valence bonding (Edward Frankland in 1852) and the application of thermodynamics to chemistry (J. W. Gibbs and Svante Arrhenius in the 1870s).

Chemical transformation
A chemical reaction is a transformation of some substances into one or more different substances. The basis of such a chemical transformation is the rearrangement of electrons in the chemical bonds between atoms. It can be symbolically depicted through a chemical equation, which usually involves atoms as subjects. The number of atoms on the left and the right in the equation for a chemical transformation is equal. (When the number of atoms on either side is unequal, the transformation is referred to as a nuclear reaction or radioactive decay.) The type of chemical reactions a substance may undergo and the energy changes that may accompany it are constrained by certain basic rules, known as chemical laws.
Energy and entropy considerations are invariably important in almost all chemical studies. Chemical substances are classified in terms of their structure, phase, as well as their chemical compositions. They can be analyzed using the tools of chemical analysis, e.g. spectroscopy and chromatography. Scientists engaged in chemical research are known as chemists.[13] Most chemists specialize in one or more sub-disciplines. Several concepts are essential for the study of chemistry; some of them are:
Bonding
Chemical bond
Atoms sticking together in molecules or crystals are said to be bonded with one another. A chemical bond may be visualized as the multipole balance between the positive charges in the nuclei and the negative charges oscillating about them. More than simple attraction and repulsion, the energies and distributions characterize the availability of an electron to bond to another atom.
A chemical bond can be a covalent bond, an ionic bond, a hydrogen bond or just because of Van der Waals force. Each of these kinds of bonds is ascribed to some potential. These potentials create the interactions which hold atoms together in molecules or crystals. In many simple compounds, valence bond theory, the Valence Shell Electron Pair Repulsion model (VSEPR), and the concept of oxidation number can be used to explain molecular structure and composition.
Ionic bond
An ionic bond is formed when a metal loses one or more of its electrons, becoming a positively charged cation, and the electrons are then gained by the non-metal atom, becoming a negatively charged anion. The two oppositely charged ions attract one another, and the ionic bond is the electrostatic force of attraction between them. For example, sodium (Na), a metal, loses one electron to become an Na+ cation while chlorine (Cl), a non-metal, gains this electron to become Cl−. The ions are held together due to electrostatic attraction, and that compound sodium chloride (NaCl), or common table salt, is formed.
Covalent bond:
In the methane molecule (CH4), the carbon atom shares a pair of valence electrons with each of the four hydrogen atoms. Thus, the octet rule is satisfied for C-atom (it has eight electrons in its valence shell) and the duet rule is satisfied for the H-atoms (they have two electrons in their valence shells).
In a covalent bond, one or more pairs of valence electrons are shared by two atoms: the resulting electrically neutral group of bonded atoms is termed a molecule. Atoms will share valence electrons in such a way as to create a noble gas electron configuration (eight electrons in their outermost shell) for each atom. Atoms that tend to combine in such a way that they each have eight electrons in their valence shell are said to follow the octet rule. However, some elements like hydrogen and lithium need only two electrons in their outermost shell to attain this stable configuration; these atoms are said to follow the duet rule, and in this way they are reaching the electron configuration of the noble gas helium, which has two electrons in its outer shell.
Similarly, theories from classical physics can be used to predict many ionic structures. With more complicated compounds, such as metal complexes, valence bond theory is less applicable and alternative approaches, such as the molecular orbital theory, are generally used.
 Energy
In the context of chemistry, energy is an attribute of a substance as a consequence of its atomic, molecular or aggregate structure. Since a chemical transformation is accompanied by a change in one or more of these kinds of structures, it is invariably accompanied by an increase or decrease of energy of the substances involved. Some energy is transferred between the surroundings and the reactants of the reaction in the form of heat or light; thus the products of a reaction may have more or less energy than the reactants.
A reaction is said to be exergonic if the final state is lower on the energy scale than the initial state; in the case of endergonic reactions the situation is the reverse. A reaction is said to be exothermic if the reaction releases heat to the surroundings; in the case of endothermic reactions, the reaction absorbs heat from the surroundings.
Chemical reactions are invariably not possible unless the reactants surmount an energy barrier known as the activation energy. The speed of a chemical reaction (at given temperature T) is related to the activation energy E, by the Boltzmann's population factor  e^{-E/kT}  – that is the probability of a molecule to have energy greater than or equal to E at the given temperature T. This exponential dependence of a reaction rate on temperature is known as the Arrhenius equation. The activation energy necessary for a chemical reaction to occur can be in the form of heat, light, electricity or mechanical force in the form of ultrasound.
A related concept free energy, which also incorporates entropy considerations, is a very useful means for predicting the feasibility of a reaction and determining the state of equilibrium of a chemical reaction, in chemical thermodynamics. A reaction is feasible only if the total change in the Gibbs free energy is negative; if it is equal to zero the chemical reaction is said to be at equilibrium.
There exist only limited possible states of energy for electrons, atoms and molecules. These are determined by the rules of quantum mechanics, which require quantization of energy of a bound system. The atoms/molecules in a higher energy state are said to be excited. The molecules/atoms of substance in an excited energy state are often much more reactive; that is, more amenable to chemical reactions.
The phase of a substance is invariably determined by its energy and the energy of its surroundings. When the intermolecular forces of a substance are such that the energy of the surroundings is not sufficient to overcome them, it occurs in a more ordered phase like liquid or solid as is the case with water (H2O); a liquid at room temperature because its molecules are bound by hydrogen bonds. Whereas hydrogen sulfide (H2S) is a gas at room temperature and standard pressure, as its molecules are bound by weaker dipole-dipole interactions.
The transfer of energy from one chemical substance to another depends on the size of energy quanta emitted from one substance. However, heat energy is often transferred more easily from almost any substance to another because the phonons responsible for vibrational and rotational energy levels in a substance have much less energy than photons invoked for the electronic energy transfer. Thus, because vibrational and rotational energy levels are more closely spaced than electronic energy levels, heat is more easily transferred between substances relative to light or other forms of electronic energy. For example, ultraviolet electromagnetic radiation is not transferred with as much efficacy from one substance to another as thermal or electrical energy.
The existence of characteristic energy levels for different chemical substances is useful for their identification by the analysis of spectral lines. Different kinds of spectra are often used in chemical spectroscopy, e.g. IR, microwave, NMR, ESR, etc. Spectroscopy is also used to identify the composition of remote objects – like stars and distant galaxies – by analyzing their radiation spectra.
The term chemical energy is often used to indicate the potential of a chemical substance to undergo a transformation through a chemical reaction or to transform other chemical substances.
Chemical reaction
During chemical reactions, bonds between atoms break and form, resulting in different substances with different properties. In a blast furnace, iron oxide, a compound, reacts with carbon monoxide to form iron, one of the chemical elements, and carbon dioxide.
When a chemical substance is transformed as a result of its interaction with another substance or with energy, a chemical reaction is said to have occurred. A chemical reaction is therefore a concept related to the "reaction" of a substance when it comes in close contact with another, whether as a mixture or a solution; exposure to some form of energy, or both. It results in some energy exchange between the constituents of the reaction as well as with the system environment, which may be designed vessels—often laboratory glassware.
Chemical reactions can result in the formation or dissociation of molecules, that is, molecules breaking apart to form two or more molecules or rearrangement of atoms within or across molecules. Chemical reactions usually involve the making or breaking of chemical bonds. Oxidation, reduction, dissociation, acid-base neutralization and molecular rearrangement are some of the commonly used kinds of chemical reactions.
A chemical reaction can be symbolically depicted through a chemical equation. While in a non-nuclear chemical reaction the number and kind of atoms on both sides of the equation are equal, for a nuclear reaction this holds true only for the nuclear particles viz. protons and neutrons.
The sequence of steps in which the reorganization of chemical bonds may be taking place in the course of a chemical reaction is called its mechanism. A chemical reaction can be envisioned to take place in a number of steps, each of which may have a different speed. Many reaction intermediates with variable stability can thus be envisaged during the course of a reaction. Reaction mechanisms are proposed to explain the kinetics and the relative product mix of a reaction. Many physical chemists specialize in exploring and proposing the mechanisms of various chemical reactions. Several empirical rules, like the Woodward–Hoffmann rules often come in handy while proposing a mechanism for a chemical reaction.
Ions and salts
The crystal lattice structure of potassium chloride (KCl), a salt which is formed due to the attraction of K+ cations and Cl− anions. Note how the overall charge of the ionic compound is zero.
An ion is a charged species, an atom or a molecule, that has lost or gained one or more electrons. When an atom loses an electron and thus has more protons than electrons, the atom is a positively charged ion or cation. When an atom gains an electron and thus has more electrons than protons, the atom is a negatively charged ion or anion. Cations and anions can form a crystalline lattice of neutral salts, such as the Na+ and Cl− ions forming sodium chloride, or NaCl. Examples of polyatomic ions that do not split up during acid-base reactions are hydroxide (OH−) and phosphate (PO43−).
Plasma is composed of gaseous matter that has been completely ionized, usually through high temperature.
Acid–base reaction
A substance can often be classified as an acid or a base. There are several different theories which explain acid-base behavior. The simplest is Arrhenius theory, which states that acid is a substance that produces hydronium ions when it is dissolved in water, and a base is one that produces hydroxide ions when dissolved in water. According to Brønsted–Lowry acid-base theory, acids are substances that donate a positive hydrogen ion to another substance in a chemical reaction; by extension, a base is the substance which receives that hydrogen ion.
A third common theory is Lewis acid-base theory, which is based on the formation of new chemical bonds. Lewis theory explains that an acid is a substance which is capable of accepting a pair of electrons from another substance during the process of bond formation, while a base is a substance which can provide a pair of electrons to form a new bond. According to this theory, the crucial things being exchanged are charges. There are several other ways in which a substance may be classified as an acid or a base, as is evident in the history of this concept.
Acid strength is commonly measured by two methods. One measurement, based on the Arrhenius definition of acidity, is pH, which is a measurement of the hydronium ion concentration in a solution, as expressed on a negative logarithmic scale. Thus, solutions that have a low pH have a high hydronium ion concentration and can be said to be more acidic. The other measurement, based on the Brønsted–Lowry definition, is the acid dissociation constant (Ka), which measures the relative ability of a substance to act as an acid under the Brønsted–Lowry definition of an acid. That is, substances with a higher Ka are more likely to donate hydrogen ions in chemical reactions than those with lower Ka values.
Redox
Redox (reduction-oxidation) reactions include all chemical reactions in which atoms have their oxidation state changed by either gaining electrons (reduction) or losing electrons (oxidation). Substances that have the ability to oxidize other substances are said to be oxidative and are known as oxidizing agents, oxidants or oxidizers. An oxidant removes electrons from another substance. Similarly, substances that have the ability to reduce other substances are said to be reductive and are known as reducing agents, reductants, or reducers.
A reductant transfers electrons to another substance and is thus oxidized itself. And because it "donates" electrons it is also called an electron donor. Oxidation and reduction properly refer to a change in oxidation number—the actual transfer of electrons may never occur. Thus, oxidation is better defined as an increase in oxidation number, and reduction as a decrease in oxidation number.

Chemical laws
Chemical reactions are governed by certain laws, which have become fundamental concepts in chemistry. Some of them are:
Avogadro's law
Beer–Lambert law
Boyle's law (1662, relating pressure and volume)
Charles's law (1787, relating volume and temperature)
Fick's laws of diffusion
Gay-Lussac's law (1809, relating pressure and temperature)
Le Chatelier's principle
Henry's law
Hess's law
Law of conservation of energy leads to the important concepts of equilibrium, thermodynamics, and kinetics.
Law of conservation of mass continues to be conserved in isolated systems, even in modern physics. However, special relativity shows that due to mass–energy equivalence, whenever non-material "energy" (heat, light, kinetic energy) is removed from a non-isolated system, some mass will be lost with it. High energy losses result in loss of weighable amounts of mass, an important topic in nuclear chemistry.
Law of definite composition, although in many systems (notably biomacromolecules and minerals) the ratios tend to require large numbers, and are frequently represented as a fraction.
Law of multiple proportions
Raoult's law.

Solubility of substances

Solubility of substances

The solubility of most substances depends strongly on the temperature and, in the case of gases, on the pressure. The solubility of most solid or liquid solutes increases with increasing temperature. The components of a mixture can often be separated using fractional crystallization, which separates compounds according to their solubilities. The solubility of a gas decreases with increasing temperature. Henry’s law describes the relationship between the pressure and the solubility of a gas.

The solubility of gases in liquids decreases with increasing temperature. Attractive inter-molecular interactions in the gas phase are essentially zero for most substances. When a gas dissolves, it does so because its molecules interact with solvent molecules. Because heat is released when these new attractive interactions form, dissolving most gases in liquids is an exothermic process ( ΔH soln<0 ). Conversely, adding heat to the solution provides thermal energy that overcomes the attractive forces between the gas and the solvent molecules, thereby decreasing the solubility of the gas.

The relationship between pressure and the solubility of a gas is described quantitatively by Henry’s law, which is named for its discoverer, the English physician and chemist, William Henry (1775–1836):                C=kP

where,
C  is the concentration of dissolved gas at equilibrium,
P  is the partial pressure of the gas, and
k  is the Henry’s law constant, which must be determined experimentally for each combination of gas, solvent, and temperature.

Gases that react chemically with water, such as  HCl  and the other hydrogen halides,  H2S , and  NH3 , do not obey Henry’s law.

Human red blood cells contain a protein called hemoglobin that specifically binds  O2  and facilitates its transport from the lungs to the tissues, where it is used to oxidize food molecules to provide energy. The concentration of hemoglobin in normal blood is about 2.2 mM, and each hemoglobin molecule can bind four  O2  molecules. Although the concentration of dissolved  O2  in blood serum at 37°C (normal body temperature) is only 0.010 mM, the total dissolved  O2  concentration is 8.8 mM, almost a thousand times greater than would be possible without hemoglobin.

Classical chemistry

Classical chemistry

In Europe, the study of chemistry was conducted by alchemists with the goals of transforming common metals into gold or silver and inventing a chemical elixir that would prolong life. Although these goals were never achieved, there were some important discoveries made in the attempt.

Robert Boyle(1627-1691) studied the behavior of gases and discovered the inverse relationship between volume and pressure of a gas. He also stated that “all reality and change can be described in terms of elementary particles and their motion,” an early understanding of atomic theory.
In 1661, he wrote the first chemistry textbook, “The Sceptical Cymist,” which moved the study of substances away from mystical associations with alchemy and toward scientific investigation. the first steps in a new science: chemistry.
By the eighteenth century the alchemists who had once tried in vain to find the philosopher's stone and the elixir of life were only a memory; the search for knowledge was now conducted along genuinely scientific lines. Many new elements (such as nickel, platinum, manganese, nitrogen, and chlorine) and a few compounds (such as carbon dioxide and hydrochloric acid) had been identified. In England science was so far advanced that plans had been made for the industrial production of sulfuric acid.
The mysterious substance: phlogiston
Despite this increasing knowledge, scientists were still hampered in their researches by an unfortunate theory which today seems to be a last relic of the age of the alchemists. It was thought that all combustible materials contained a mysterious substance called phlogiston, the principle of fire. When materials burned, it was the phlogiston that was seen to escape in the form of flames. Even when it was shown that the products of combustion weighed more than the original material, this was accounted for by the belief that the phlogiston that had escaped must have a negative weight – that is, it weighed less than nothing.
The first strong evidence against phlogiston theory came from pneumatic chemists in Britain during the later half of the 18th century. Joseph Black, Joseph Priestley and Henry Cavendish all identified different gases that composed air; however, it was not until Antoine Lavoisier discovered in the fall of 1772 that, when burned, sulphur and phosphorus “gain[ed] in weight” that the phlogiston theory began to unravel.

In a series of experiments on the combustion of mercury, Lavoisier showed that in fact when mercury "burned" it combined with a part of the air – oxygen – and increased in weight. By the use of the balance he showed that the weight of the oxide formed was equal to the weight of the mercury and the oxygen that had combined with it. His experiments finally disproved the phlogiston theory, and the way now lay open to the proper understanding of other chemical reactions.
In later experiments Lavoisier discovered the nature of air and water. He has justly been called the father of modern chemistry.
By the 1700s, the Age of Enlightenment had taken root all over Europe. Joseph Priestley (1733-1804) disproved the idea that air was an indivisible element. He showed that it was, instead, a combination of gases when he isolated oxygen and went on to discover seven other discreet gases. Jacques Charles continued Boyles’ work and is known for stating the direct relationship between temperature and pressure of gases. In 1794, Joseph Proust studied pure chemical compounds and stated the Law of Definite Proportions — a chemical compound will always have its own characteristic ratio of elemental components. Water, for instance, always has a two-to-one ratio of hydrogen to oxygen.
Antoine Lavoisier (1743-1794) was a French chemist who made important contributions to the science. While working as a tax collector, Lavoisier helped to develop the metric system in order to ensure uniform weights and measures. He was admitted to the French Academy of Sciences in 1768. Two years later, at age 28, he married the 13-year-old daughter of a colleague. Marie-Anne Lavoisier is known to have assisted her husband in his scientific studies by translating English papers and doing numerous drawings to illustrate his experiments.
Lavoisier’s insistence on meticulous measurement led to his discovery of the Law of Conservation of Mass. In 1787, Lavoisier published "Methods of Chemical Nomenclature," which included the rules for naming chemical compounds that are still in use today. His "Elementary Treatise of Chemistry" (1789) was the first modern chemistry textbook. It clearly defined a chemical element as a substance that cannot be reduced in weight by a chemical reaction and listed oxygen, iron, carbon, sulfur and nearly 30 other elements then known to exist. The book did have a few errors though; it listed light and heat as elements.
Amedeo Avogadro (1776-1856) was an Italian lawyer who began to study science and mathematics in 1800. Expanding on the work of Boyle and Charles, he clarified the difference between atoms and molecules. He went on to state that equal volumes of gas at the same temperature and pressure have the same number of molecules. The number of molecules in a 1-gram molecular weight (1 mole) sample of a pure substance is called Avogadro’s Constant in his honor. It has been experimentally determined to be 6.023 x 1023 molecules and is an important conversion factor used to determine the mass of reactants and products in chemical reactions.
In 1803, an English meteorologist began to speculate on the phenomenon of water vapor. John Dalton (1766-1844) was aware that water vapor is part of the atmosphere, but experiments showed that water vapor would not form in certain other gases. He speculated that this had something to do with the number of particles present in those gases. Perhaps there was no room in those gases for particles of water vapor to penetrate. There were either more particles in the “heavier” gases or those particles were larger. Using his own data and the Law of Definite Proportions, he determined the relative masses of particles for six of the known elements: hydrogen (the lightest and assigned a mass of 1), oxygen, nitrogen, carbon, sulfur and phosphorous. Dalton explained his findings by stating the principles of the first atomic theory of matter.
⦁ Elements are composed of extremely small particles called atoms.
⦁ Atoms of the same element are identical in size, mass and other properties. Atoms of different elements have different properties.
⦁ Atoms cannot be created, subdivided or destroyed.
⦁ Atoms of different elements combine in simple whole number ratios to form chemical compounds.
⦁ In chemical reactions atoms are combined, separated or rearranged to form new compounds.

On 6 March 1869, Mendeleev  made a formal presentation to the Russian Chemical Society, titled The Dependence between the Properties of the Atomic Weights of the Elements, which described elements according to both atomic weight (now called relative atomic mass) and valence.This presentation stated that:
⦁ The elements, if arranged according to their atomic weight, exhibit an apparent periodicity of properties.
⦁ Elements which are similar regarding their chemical properties either have similar atomic weights (e.g., Pt, Ir, Os) or have their atomic weights increasing regularly (e.g., K, Rb, Cs).
⦁ The arrangement of the elements in groups of elements in the order of their atomic weights corresponds to their so-called valencies, as well as, to some extent, to their distinctive chemical properties; as is apparent among other series in that of Li, Be, B, C, N, O, and F.
⦁ The elements which are the most widely diffused have small atomic weights.
⦁ The magnitude of the atomic weight determines the character of the element, just as the magnitude of the molecule determines the character of a compound body.
⦁ We must expect the discovery of many yet unknown elements – for example, two elements, analogous to aluminum and ⦁ silicon, whose atomic weights would be between 65 and 75.
⦁ The atomic weight of an element may sometimes be amended by a knowledge of those of its contiguous elements. Thus the atomic weight of ⦁ tellurium must lie between 123 and 126, and cannot be 128. (Tellurium's atomic weight is 127.6, and Mendeleev was incorrect in his assumption that atomic weight must increase with position within a period.)
⦁ Certain characteristic properties of elements can be foretold from their atomic weights.

Mendeleev published his periodic table of all known elements and predicted several new elements to complete the table in a Russian-language journal. Only a few months after, Meyer published a virtually identical table in a German-language journal. Mendeleev has the distinction of accurately predicting the qualities of what he called ekasilicon, ekaaluminium and ekaboron (germanium, galliumand scandium, respectively).

History Of The Internet

History Of The Internet

The first TCP/IP-wide area network was created on January 1, 1983 when all hosts on the ARPANET were switched over from the older protocols to TCP/IP.
In 1984, the National Science Foundation (NSF) commissioned the construction of a 1.5 megabit/second network which became known as NSFNET. In 1989 the US Federal Networking Council approved the interconnection of the NSFNET to the commercial MCI Mail system.
Soon after, other commercial e-mail services were connected such as OnTyme, Telemail, and CompuServe. Three Internet Service Providers (ISPs) were also created: UUNET, PSINET, and CERFNET. More and more seperate networks were created that eventually interconnected with this large, growing network of networks.
The ability of TCP/IP to work over virtually any pre-existing communication networks allowed for a great ease of growth, although the rapid growth of the Internet was due primarily to the availability of commercial routers from companies such as Cisco Systems, Proteon and Juniper, the availability of commercial Ethernet equipment for local-area networking and the widespread implementation of TCP/IP on the UNIX operating system.
Although the basic applications and guidelines that make the Internet possible had existed for almost a decade, the network did not gain public face until the 1990s. On August 6, 1991, the European Organization for Nuclear Research, (CERN), which straddles the border between France and Switzerland, publicized the new World Wide Web project. The web was invented by English scientist Tim Berners-Lee in 1989.
 By 1996 usage of the word “Internet” had become commonplace, and consequently, so had its use as a reference to the World Wide Web. Over the course of the decade, the Internet successfully accommodated the majority of previously existing public computer networks (although some networks have remained seperate).

Earth crust composition

Earth crust composition

More than 90% on the crust is composed of silicate minerals. Most abundant silicates are feldspars (plagioclase (39%) and alkali feldspar (12%)). Other common silicate minerals are quartz (12%) pyroxenes (11%), amphiboles (5%), micas (5%), and clay minerals (5%). The rest of the silicate family comprises 3% of the crust. Only 8% of the crust is composed of non-silicates — carbonates, oxides, sulfides, etc.
Clay minerals are too small to be shown individually. Even with a light microscope you will see only mud or dust depending on whether these minerals are wet or dry. Clay minerals are silicates that are the products of weathering of other silicate minerals, mostly feldspars.
The most abundant rock types in the crust
Rocks are divided into three broad groups: igneous, metamorphic, and sedimentary rocks. The oceanic crust is largely composed of basaltic igneous rocks which are covered by a thin veneer of sediments which are thickest near the margins of the continental landmasses. The continental crust is much thicker and older. The continental crust is also much more variable and structurally very complex. Virtually all the rock types known to man occur in the continental crust. Even meteorites, xenoliths from the mantle, and ophiolites (fragment of former oceanic crust) are constituents of the continental crust because that’s where we found them.
Roughly three fourths of the continental crust is covered by sedimentary rocks and almost all of it is covered by loose sediments (soil, sand, dirt, etc.). We are most likely to encounter these materials, but it is important to understand that despite being so ubiquitous on the surface, they make up only about 8% of the whole mass of the crust. Sediments consolidate to sedimentary rocks after burial. Sand turns to sandstone, limy mud to limestone, clay to claystone. Sedimentary rocks are stable only in the upper parts of the crust. High pressure and temperature in the deeper parts metamorphoses them (minerals recrystallize) to various metamorphic rocks. The bulk of the continental crust is made of metamorphic rocks. Igneous rocks are also common on the surface in volcanically active regions, but they also occur deeper in the crust as granitic (mostly) intrusions.
Important sediments are sand, clay, mud (wet mixture of clay and fine sand), and limy mud. Widespread sedimentary rocks are limestone (2% of the crust by volume), sandstone (1.7%), claystone (4.2%) which are lithified versions of the loose sediments mentioned before. Chemical sediments like halite and gypsum are important as well, but their overall volume is clearly less than 1% of the crust. Important igneous rocks are granite, granodiorite, gabbro, basalt, diorite, andesite, etc. It is very difficult to say what is the percentage of these rocks. Important metamorphic rocks are metamorphosed equivalents of widespread sedimentary and igneous rocks. Common metamorphic rocks are slate (metamorphosed claystone), schist (met. claystone, higher grade than slate) quartzite (met. sandstone), marble (met. limestone), gneiss (met. igneous rock or sedimentary rocks), amphibolite (met. basaltic rocks).
Gypsum is an evaporite mineral. Evaporites are water-soluble chemical sediments that crystallize out of concentrated (high salinity) seawater in lagoons. 

James Watt’s Steam Engine:

James Watt’s Steam Engine:

In 1763, James Watt was working as instrument maker at the University of Glasgow when he was assigned the job of repairing a model Newcomen engine and noted how inefficient it was.

Out walking one Sunday afternoon in 1765, Watt sees the solution. He envisions a separate chamber in which the steam can condense. This means no more need for cooling and reheating, making the engine faster and more fuel efficient. His insights will convert a machine of limited use into one that will power the industrial revolution.

In 1765, Watt conceived the idea of equipping the engine with a separate condensation chamber, which he called a "condenser". Because the condenser and the working cylinder were separate, condensation occurred without significant loss of heat from the cylinder. The condenser remained cold and below atmospheric pressure at all times, while the cylinder remained hot at all times.

Steam was drawn from the boiler to the cylinder under the piston. When the piston reached the top of the cylinder, the steam inlet valve closed and the valve controlling the passage to the condenser opened. The condenser being at a lower pressure, drew the steam from the cylinder into the condenser where it cooled and condensed from water vapor to liquid water, maintaining a partial vacuum in the condenser that was communicated to the space of the cylinder by the connecting passage. External atmospheric pressure then pushed the piston down the cylinder.

The separation of the cylinder and condenser eliminated the loss of heat that occurred when steam was condensed in the working cylinder of a Newcomen engine. This gave the Watt engine greater efficiency than the Newcomen engine, reducing the amount of coal consumed while doing the same amount of work as a Newcomen engine.

Watt's next improvement to the Newcomen design was to seal the top of the cylinder and surround the cylinder with a jacket. Steam was passed through the jacket before being admitted below the piston, keeping the piston and cylinder warm to prevent condensation within it. The second improvement was the utilisation of steam expansion against the vacuum on the other side of the piston. The steam supply was cut during the stroke, and the steam expanded against the vacuum on the other side. This increased the efficiency of the engine, but also created a variable torque on the shaft which was undesirable for many applications, in particular pumping. Watt therefore limited the expansion to a ratio of 1:2 (i.e. the steam supply was cut at half stroke). This increased the theoretical efficiency from 6.4% to 10.6%, with only a small variation in piston pressure. Watt did not use high pressure steam because of safety concerns.

These improvements led to the fully developed version of 1776 that actually went into production.

As fully developed, it used about 75% less fuel than a similar Newcomen one.

The Watt steam engine, alternatively known as the Boulton and Watt steam engine, was the first practical steam engine and was one of the driving forces of the industrial revolution. James Watt developed the design sporadically from 1763 to 1775 with support from Matthew Boulton. Watt's design saved significantly more fuel compared to earlier designs that they were licensed based on the amount of fuel they would save. Watt never ceased developing the steam engine, introducing double-acting design.

James Watt's improvements to the steam engine "converted it from a prime mover of marginal efficiency into the mechanical workhorse of the Industrial Revolution". The availability of efficient, reliable motive power made whole new classes of industry economically viable, and altered the economies of continents. In doing so it brought about immense social change, attracting millions of rural families to the towns and cities.

In 1785 he and Boulton were elected fellows of the Royal Society of London.Watt established in 1794 the new firm of Boulton & Watt, which built the Soho Foundry to manufacture steam engines more competitively. In 1784 Watt makes further improvements to the steam engine and patents a steam locomotive. By 1790, both Boulton & Watt are wealthy men.