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You too can be as smart as Einstein (almost) The elements Earth – dry, heavy Water – wet, heavy Air – cool, light Fire – warm, light The composition of a substance could be estimated from its properties. These ideas were based on observation, logic and reason, but not experimentation. Democritus (460 B.C. - 370 B.C.) Matter is made of small, hard indivisible particles called atoms, which exist in the void. These atoms differ in size and shape, but not in any other way. Quantitative differences (how much) vs. Qualitative differences (what kind) Antoine Lavoisier (1743-1794) Discoverer of Oxygen (disputed) His work refuted the phlogiston theory Responsible for the law of conservation of matter. Claude Louis Berthollet Joseph Louis Proust Berthollet – “compounds do not have a fixed composition”. Cu + S CuxSy Every time he tried the experiment he got a different result. Proust - compounds have a fixed composition. 2H2 + O2 2H2O He always got the same result. Proust’s argument is called The Law of Definite Proportions. He was proved to be right. John Dalton (ca. 1804) The father of modern atomic theory Schoolteacher Colorblind – studied colorblindness The points of Dalton’s theory All matter is made of atoms Atoms are indivisible and indestructible All atoms of one element are exactly alike, and atoms of different elements are different. Atoms combine in small whole number ratios to form compounds. The Law of Multiple Proportions: If two elements combine to make two different compounds, the ratios of the elements involved are small whole numbers. Examples: CO and CO2 CuS and Cu2S H2O and H2O2 J. J. Thomson and the Electron (1897) Thomson discovered the electron - he called it a “corpuscle”. He used an instrument called a Crookes tube. Cathode (-) Evacuated tube Anode (+) He noticed a stream of charged particles coming from the cathode, called cathode rays. Thomson proposed the "plum pudding" atomic model - negatively charged corpuscles swarm inside a cloud of massless positive charge. The gold foil experiment (1909) Most of the alpha particles went straight through, and a few were bounced straight back. Rutherford’s interpretation: The atom has a small, hard, dense and positively charged nucleus. The electrons are outside the nucleus. Discovery of the proton: Henry Moseley (1913) Moseley bombarded metals with x-rays Each successive element had one more positive charge – called “atomic number” Rutherford proved that the nucleus of nitrogen contains hydrogen nuclei – a “proton” (1918-19) Discovery of the neutron – James Chadwick (1932) Name Charge Mass (amu) Location Discoverer Electron -1 1/2000 outside nucleus Thomson Proton +1 1 nucleus Moseley/ Rutherford Neutron 0 1 nucleus Chadwick Atomic number = number of protons in the nucleus Atomic number determines the identity of the element Mass number = protons + neutrons Number of electrons = number of protons Isotopes: two atoms of the same element with different numbers of neutrons C-12 and C-13 are isotopes of carbon Nuclear symbols 13C 6 Write the nuclear symbol for lead-206. 206Pb 82 Atomic number 20 Symbol Name Ca Calcium 40.078 Average atomic mass Average mass of all the isotopes of an element Average is weighted Example: Boron has two isotopes, B-10 and B-11 B-10: 19.9% B-11: 80.1% Average atomic mass of boron: 10x0.199 = 1.99 11x0.801 = 8.811 Average atomic mass = 1.99 + 8.811 = 10.8amu Niels Bohr and the stepwise atom (ca. 1918) Rutherford suggested that electrons orbit around the nucleus like planets around the sun. This did not explain emission spectra, which gave sharp lines. He theorized that electrons could only travel in certain sized orbits, and not anywhere in between. The orbits were called energy levels. Each orbit has a specific energy. Electrons can jump from one level to another; as they do, they absorb or emit energy. Erwin Schrödinger and probable cause (ca. 1935) Schrödinger’s work showed that electrons do not move in actual “orbits”. Electrons move randomly and form “probability clouds”. The shape of these clouds is similar to the shape of Bohr’s orbits. The position and momentum of an electron cannot be determined simultaneously (Heisenberg Uncertainty Principle) Schrödinger’s “electron cloud” Bohr suggested that electrons inhabit energy levels around the nucleus. Each level has a specific energy associated with it. The outermost (highest energy) level is called the “valence shell”. The electrons in the valence shell are called the “valence electrons”. The valence electrons are the most important electrons in the chemistry of the atom. The number of levels depends on the number of electrons. The first level (K) holds two electrons. The second level holds eight electrons. The third level holds 18, and the fourth 32. No atom can have more than eight electrons in its valence shell. When the valence shell reaches eight electrons, the next two electrons are put in a higher level. Then the lower level can be filled. Lewis dot structures show how many electrons are in the valence shell of an atom. Lewis dot structure for sodium The first electron always goes to the right of the symbol. The second is paired with the first. Lewis dot structure of magnesium The third goes on top. Lewis dot structure of aluminum The fourth goes on the left, and is not paired. The fifth goes on the bottom, and successive electrons are paired until a total of eight is reached. Lewis dot structure of silicon Lewis dot structure of oxygen Bohr’s model based on atomic spectra Obtaining emission atomic spectra Energy is applied to a gas or liquid sample. Flame test (for samples in solution) Gas discharge tube The energy makes an electron or two jump to a higher energy level. The electrons fall back down to a lower level, and give off energy in the form of light – bright lines against a dark background. Absorption spectra – light is passed through a sample and analyzed – looks like a rainbow with dark lines Interpreting atomic spectra The light given off is viewed through a spectroscope. The spectroscope has either a prism or a grating, which splits the light into its component colors. Only a few sharp lines appear in the spectrum. Each line corresponds to a specific electron transition. Transition = jump from one energy level to another Light energy travels in the form of waves. Color depends on frequency. High frequency = violet end of spectrum Low frequency = red end of spectrum Energy also depends on frequency, so each color has its own energy. Blue or violet is higher energy than red or green. When a specific color line is seen in a spectrum, the energy of the electron transition responsible can be calculated. Bohr reasoned that since only certain lines are seen in atomic spectra, only certain energies must be allowed in electron orbits.