Early History of Chemistry Courtesy: www.lab-initio.com
Greeks • 400 B.C. – Matter was composed of four fundamental substances: fire, earth, water, and air. • Greeks considered whether matter was continuous (infinitely divisible into smaller pieces) or composed of small indivisible particles. • Democritus (460 – 370 B.C.) used the term atomos (which later became atoms) to describe the small indivisible particles. • No experiments to test theories so no conclusion reached about divisibility of matter.
Alchemy • Next 2000 years were dominated by a pseudoscience called alchemy. • Some alchemists were fakes who were obsessed with turning base metals into gold. • Others were serious scientists who made several advances such as the discovery of several elements.
Sixteenth Century • The first “chemist” to perform true quantitative experiments was RobertBoyle (1627-1691). • He measured the relationship between pressure and volume of air. • Published The Skeptical Chymist in 1661.
Seventeenth and Eighteenth Centuries • German chemist Georg Stahl (1660-1734) suggested a substance burning in a closed container stopped when the air became saturated with “phlogiston” ( a substance that flowed out of a burning material). • Joseph Priestley (1733-1804) discovered oxygen gas. • Oxygen gas, he found, supported combustion and was supposed to be low in phlogiston. Oxygen was originally called “dephlogisticated air”.
Fundamental Chemical Laws • Antoine Lavoisier (1743-1794), a French chemist, explained the true nature of combustion. • He also verified the law of conservation of mass. • Law of conservation of mass– mass is neither created nor destroyed. • Lavoisier was also a tax collector and was executed on the guillotine in 1794.
After 1800, chemists used experimentation to study chemical reactions and determine composition of chemical compounds. • Joseph Proust (1754-1826), a French chemist, showed that a given compound always contains exactly the same proportion of elements by mass. • This principle is known as the law of definite proportions. • For example, oxygen makes up 8/9 of the mass of any sample of water, while hydrogen makes up the remaining 1/9 of the mass.
John Dalton (1766-1844), an English schoolteacher, proposed that atoms are the particles that compose elements. • Dalton discovered a principle that became known as the law of multiple proportions. • Law of multiple proportions: when two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers.
Example: The Law of Multiple Proportions The following data was collected for two compounds of hydrogen and oxygen: water (H2O) and hydrogen peroxide, H2O2. Mass of Oxygen That Combines with 1 gram of Hydrogen __________________________________________________ Water, H2O 7.92 g Hydrogen Peroxide, H2O2 15.84 g For the law of multiple proportions to be true, the ratios of the masses of oxygen combining with 1 gram of hydrogen should be a small whole number: This is a “small whole number” that supports the law of multiple proportions.
Dalton’s Atomic Theory • In 1808 Dalton published A New System of Chemical Philosophy, in which he presented his theory of atoms: How does this relate to chemistry? There are 6.02 x 1023 atoms in 55.85 g of iron. Although graphite and diamond have different properties, they are both composed of carbon. The carbon atoms are identical. Each element is made up of tiny particles called atoms. The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways.
Chemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms. Chemical reactions involve reorganization of the atoms – changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction. C + O2→ CO2; CO2 is not CO, CO3 or Fe2O3. 2H2 + O2 → 2H2O, not CS2 or NaCl.
Dalton also prepared the first table of atomic masses. • Many of the masses were proved to be wrong but the construction of the table was an important step. • Joseph Gay-Lussac (1778-1850) and Amadeo Avogadro (1776-1856) provided the keys to the determination of absolute formulas for compounds. Gay-Lussac Avogadro
Early Experiments to Characterize the Atom • Based on the work of Dalton, Gay-Lussac, and Avogadro, the concept of the atom was becoming generally accepted. • The next questions to be addressed were: • What is an atom made of? • How do the atoms of the various elements differ? • The first important experiments on the composition of the atom were conducted by J. J. Thomson (1856-1940) from 1898 to 1903.
J. J. Thomson (1898-1903) • Postulated the existence of electrons using cathode-ray tubes. • Determined the charge-to-mass ratio of an electron. • The atom must also contain positive particles that balance the exact negative charge carried by particles we now know as electrons.
The cathode ray was deflected by the negative pole of an applied electrical field, Thomson postulated the ray was a stream of negatively charged particles, now called electrons.
Thomson knew atoms were electrically neutral so he assumed atoms contain some positive charge to balance the negative charge. • Developed the “plum pudding” model.
Robert Millikan (1909) • Performed experiments involving charged oil drops. • Determined the magnitude of the charge on a single electron. • Calculated the mass of the electron. Millikan put a charge on a drop of oil and measured how strong an applied electrical field had to be in order to stop the drop from falling. Determined the mass and the force of gravity on the drop and from this calculated the charge.
Ernest Rutherford (1911) • Explained the nuclear atom. • Atom has a dense center of the positive charge called the nucleus. • Performed “gold foil experiment”.
If Thomson’s “plum pudding” model were correct all alpha particles would pass through. • Rutherford found some particles were deflected. • Postulated must have encountered some positive charge in center of atom. • Called this center of positive charge the “nucleus”.