2.1: Historical Development of Atomic Theory

• Leucippus and Democritus argued that all matter was composed of small, finite particles that they called atomos, a term derived from the Greek word for invisible.
• Aristotle and others came to the conclusion that matter consisted of various combination of the four elements - fire, earth, air, and water - and could be infinitely divided.
• The Aristotelian view of the composition of matter held sway until English schoolteacher John Dalton helped revolutionize chemistry with his hypothesis that the behavior of matter could be explained using atomic theory.
• Dalton's hypothesis about the microscopic features of matter are still valid in modern atomic theory.
• Matter is composed of small particles called atoms, which are the smallest unit of an element that can participate in a chemical change.
• An element consists of only one type of atom, which has a mass that is characteristic of the element and is the same for all atoms of that element.
• Atoms of one element differ in properties from atoms of all other elements.
• A compound consists of atoms of two or more elements combined in a small, whole-number ratio.
• Atoms are neither created nor destroyed during a chemical change, but are instead rearranged to yield substances that are different from those present before the change.
• The Law of Conservation of Mass states that mass cannot be created nor destroyed, rather transformed.
• The Law of Definite Proportions or the Law of Constant Composition states that all samples of a pure compound contain the same elements in the same proportion by mass.
• Samples that have the same mass ratio are not necessarily the same substance.
• The Law of Multiple Proportions states that when two elements react to form more than one compound, a fixed mass of one element will react with masses of the other element in a ratio of small, whole numbers.
• In the late 1800s, English physicist J. J. Thomson investigated electrical discharges produced in low-pressure gases using a cathode ray tube.
• When high voltage was applied across the electrodes, a visible cathode ray was deflected toward the positive charge and away from the negative charge.
• Measurements of the extent of deflection of the cathode ray and the magnetic field strength allowed the calculation of the charge-to-mass ratio of the particles.
• Particles are attracted by positive charges and repelled by negative charges, so they must be negatively charged.
• Particles are less massive than atoms and indistinguishable, regardless of the source material, so they must be fundamental, subatomic constituents of all atoms.
• Thomson's cathode ray particle is what we now call an electron.
• Millikan created microscopic oil droplets, which could be electrically charged by friction or x-rays.
• By adjusting the electric field strength, Millikan was able to determine the charge on individual oil drops.
• Millikan concluded that the charge of an oil droplet was a multiple of 1.6 x 10^-19 C, the charge of a single electron.
• Since the charge of an electron was known due to Millikan's research, and the charge-to-mass ratio was known due to Thomson's research, the mass of an electron was able to be calculated, 9.107 x 10^-31 kg.
• Thomson proposed the "plum pudding" model of atoms, which described a positively charged mass with an equal amount of negative charge in the form of electrons embedded in it.
• Nagaoka postulated a Saturn-like atom, consisting of a positively charged sphere surrounded by a halo of electrons.
• Ernest Rutherford performed a series of experiments using a beam of high-speed, positively charged alpha particles that were produced by the radioactive decay of radium.
• Rutherford, Geiger, and Marsden aimed a beam of alpha particles at a thin gold foil and examined the scattering of the particles using a luminescent screen.
• Most particles passed through the foil, but some were diverted, and a small number were deflected back towards the source.
• Rutherford concluded that the volume occupied by an atom must consist of a large amount of empty space.
• Rutherford concluded that a small, relatively heavy, positively charged body, the nucleus, must be at the center of each atom.
• Rutherford proposed a model which an atom consists of a very small, positively charged nucleus, in which most of the mass of the atom is concentrated, surrounded by a cloud of negatively charged electrons.
• Rutherford discovered that the nuclei of other elements contain the hydrogen nucleus as a "building block," named the proton.
• A proton is a positively charged subatomic particle found in the nucleus.
• Frederick Soddy realized that an element could have types of atoms with different masses that were chemically indistinguishable, otherwise known as isotopes.
• James Chadwick found evidence of neutrons, uncharged, subatomic particles with a mass approximately the same as that of protons.
• Isotopes differ in mass because they have different numbers of neutrons, but they are chemically identical because they have the same number of protons.