The first unit we will undertake is Structures & Properties of Matter. Although some of the material will be familiar (ionic, polar covalent and pure covalent bonding), most of this is honest-to-goodness grade 12 stuff. Woo!
Learning Goals:
So, let's get cracking:
Atomic Theory
- Democritus (460-370 BC)
Greek philosophers wondered whether matter could be divided endlessly. Most philosophers, including Aristotle and Plato, believed that matter was infinitely divisible. However, Democritus disagreed with this viewpoint. He argued that that matter is composed of very small, indivisible particles which he called atomos, meaning indivisible. Unfortunately, Democritus’ proposition was not widely accepted until the early nineteenth century.
- John Dalton (1803-1807)
During the period of 1803-1807, John Dalton published an atomic theory that was designed to explain experimental data. His theory, of which most is still accepted today, consists of the following postulates:
1. Each element is composed of extremely small particles called atoms.
2. All atoms of a given element are identical; the atoms of different elements are different and have different properties.
3. Atoms of an element are not changed into different types of atoms by chemical reactions; atoms are neither created nor destroyed in chemical reactions. (Law of Conservation of Mass)
4. Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms. (Law of Definite Proportions)
- J. J. Thomson (1897)
In the mid-1800s, scientists began to suspect that the atom was made of small particles. They began to study electrical discharge through partially evacuated glass tubes. This radiation was known as cathode rays, because it originated from the negative electrode or cathode. Although the rays were invisible, they caused certain substances painted on the inside of the tubes to fluoresce and thus, their presence could be detected. Cathode rays travel in straight lines, unless induced to bend by an outside magnetic or electric field. Regardless of the source of radiation, the cathode rays always bent in a manner expected for negatively charged particles. This led to the conclusion that the streams of negative particles (electrons) are a basic component of all matter. Using a cathode ray tube, J. J. Thomson measured the ratio of the electrical charge to mass of the electron. Thomson proposed the ‘raisin bun’ or ‘plum pudding’ model of the atom, in which raisins (negative electrons) are embedded in the dough (region of positive charge).
- Ernest Rutherford (1910)
One of Thomson’s students, Ernest Rutherford, performed an experiment in 1910 that led to the downfall of Thomson’s atomic model. Rutherford was studying the angles at which alpha particles were scattered as they passed through thin gold foil. If Thomson’s model was true, the alpha particles should pass directly through the foil without deflection. However, a few particles were deflected and a few even bounced straight back toward the source. Rutherford postulated that most of the mass of the atom, and all of its positive charge, reside in a very small, very dense region which he called the nucleus. Most of the atom is empty space in which electrons move around the nucleus. Thus, most of the alpha particles passed through the foil unaffected because they were not close to the tiny nucleus. A few of the positively charged alpha particles, however, came close to the positive nucleus and were deflected by the force of repulsion.
Later study by Rutherford and Thomson, using positive rays in the cathode ray tube led to the discovery of the proton in 1914. They were able to determine the proton had a charge equal in magnitude but opposite in sign to that of the electron. It was also determined that the proton is 1836 times as massive as the electron.
- James Chadwick (1932)
James Chadwick, working with Rutherford, was attempting to calculate the masses of nuclei by bombarding elements with alpha particles. When a comparison of the masses of the nuclei was compared to the sum of the masses of the protons in the element, they were not the same. He hypothesized that the nucleus of the atom consisted of protons, as well as neutrally charged particles, called neutrons.
During 1919-1925, Francis Aston, using a mass spectrometer, found that some elements had atoms of different masses. Although he could not explain his findings, he named the atoms of different mass isotopes. With the discovery of the neutron, the existence of the isotope could now be explained. Isotopes have different numbers of neutron in the nucleus, which resulted in different nuclear stability and mass.
- Quantum Effects and Photons
When solids are heated, they emit radiation, as seen in the light emitted by a tungsten filament in an incandescent light bulb. The wavelength distribution depends on temperature, with a red-hot object being cooler than a white-hot one. In the late 1800s, many physicists were trying to understand the relationship between the temperature and intensity (brightness) and wavelengths of the emitted radiation. Max Planck solved this problem by making the assumption that energy can be released or absorbed by atoms only in packets of some minimum size, which he gave the name quantum. According to Planck’s theory, energy is always released in whole number multiples of the quantum, rather than as a continuous stream of energy. Allowed energies are therefore quantized, which is analogous to climbing a ladder, you can only stop on rungs, not between them. The quantum for any given object is quite small, and therefore the gain or loss of quantum energy is all but unnoticeable for macroscopic objects. However, the impact of quantized energies is far more significant for matter at the atomic level.
- The Photoelectric Effect
In 1887, Heinrich Hertz noticed that light shining on a metal surface liberates electrons from the surface, but he had no explanation for this observation. In 1905, Einstein, using Planck’s quantum theory, reasoned that light consisted of a stream of energy packets, or photons (with each colour of light having a certain amount of energy, red less than yellow less than blue, etc). Albert Einstein deduced that the size of a quantum of electromagnetic energy depends directly on its frequency. When a photon strikes the metal, its energy is transferred to a metal electron. The electron, to break free of the metal atom, uses a portion of the energy and the remainder appears as the kinetic energy of the ejected electron.
- Atomic Spectra
Most radiation sources produce radiation containing many different wavelengths. When radiation from such sources is separated into its different wavelength components, a spectrum is produced. For instance, when white light passes through a prism, a continuous spectrum (rainbow) is produced. In 1859, Bunsen and Kirchhoff studied heated samples of elements and discovered that when energy was released by the sample, they did not produce a continuous spectrum, but rather a series of coloured lines separated by black regions, called a bright line spectrum. Conversely, when energy is absorbed, an absorption or dark line spectrum is produced and the dark lines on this spectrum are in the same position as the coloured lines from the bright line spectrum.
- Bohr's Model of the Atom
In 1913, Neils Bohr suggested that only certain quanta of light could be emitted or absorbed by an atom. He said that this meant that the energy of the electrons in an atom must also be quantized. He proposed the planetary atomic model with a nucleus at the centre with each electron orbit at a fixed distance with a fixed energy. To explain atomic spectra, he stated that electrons could change their energy only by undergoing a transition from one stationary state to another. Thus, the energy released or absorbed by the electron was equivalent to the size of the ‘jump’ between stationary states, resulting in unique line spectra for all elements.
- Problems with the Bohr Model
1. He was wrong in postulating that an electron moves in an orbit of fixed radius, which changes only when the electron jumps to another orbit having a different fixed radius.
2. He was wrong when he supposed that the electron is a particle whose position and motion can be specified exactly at any given time.
3. The model applied to the simplest possible atom – hydrogen, which has only one electron. His model did not apply to multi-electron atoms.
Homework: Typically, each day's homework is divided by the dotted line. There will also be additional pages occasionally, but they will be included in that day's lesson.
Answer Keys:
#1-3
#4, 5
