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NOTES 2013 Atomic Theory 2 Atomic Theory Development of Theory Historical Atomic Models Democritus Greek philosopher who postulated that matter is comprised of atoms as the smallest part (ca 400 BC) John Dalton atoms are the smallest, indivisible part of an element, are alike for one kind of element, combine chemically in whole ratios, and compounds are composed of elements (1803) Max Planck first postulated that energy is quantized, originator of quantum theory (1900) J.J. Thompson “plum-pudding” model – negative electrons (plums) are located in a positively charged pudding (1903) Hantaro Nagaoka “Saturnian” model – large nucleus with electrons orbiting in rings (1904) Robert Millikan measured the charge and mass of an electron (1908) Ernest Rutherford atom has a small, positive, central nucleus, which contains the mass and is surrounded by a cloud of negative electrons [correct model] (1911) Discovered the proton. (1917) Neils Bohr “planetary” model – the nucleus is surrounded by electrons orbiting in rings (1913) Edwin Schroedinger mathematical wave equation led to prediction of the possible states for an electron [correct, part of orbital theory] (1926) Werner Heisenberg “uncertainty” principle correctly states that it is impossible to predict the exact position and momentum of an electron (1927) James Chadwick discovered the neutron which served as a catalyst for the growth of nuclear physics (1932) Rutherford Experiment method: involved shooting alpha particles (He2+) at a sheet of gold foil postulate: the deflection of the alpha particles would determine the location of the mass within the atom results: most particles went straight through, while some deflected back conclusions: atom is mostly empty space {why many particles went straight}, with almost all the mass in a small positively charged nucleus {why some particles were deflected backwards} Nucleus Central portion of the atom, positively charged and contains the mass . Electron cloud The bulk of the space of an atom, mostly empty, negatively charged. Basic Subatomic Particles electron negative charge (–) located in electron cloud 0 u proton positive charge (+) located in nucleus 1 u neutron neutral ( ) located in nucleus 1 u Note that for an individual atom, the number of protons and neutrons never changes in ordinary reactions. The number of electrons can change, which effects the charge of the atom, but the nucleus does not change. Atomic Mass, Y the sum of the protons and neutrons. p + n e.g. an atom with 6 protons and 7 neutrons has an atomic mass of 13 Atomic Number, Z The number of protons. This defines the element. For example carbon always has 6 protons, but is known to have 6 neutrons (Y=12) or 7 neutrons (Y=13) or 8 neutrons (Y=14) See isotopes below. Isotope Atoms with the same number of protons (same element), but with different number of neutrons. Percent Abundance the percentage of one isotope for an element Average Atomic Mass, Yavg a weighted average of all known isotopic masses for an element Yavg Y1 Y2 where X = percent abundance as a decimal Y1 and Y2 are isotopic masses Charge Some particles emit an electric force creating a field of force around the particle. This field attracts opposite fields while repelling similar fields. These fields are positive (+) or negative (-). The absence of the field is neutral ( ). Ion a charged atom or molecule Cation positive ion, lost electrons Anion negative ion, gained electrons To calculate charge atom – number of excess protons or electrons. e.g. if an atom contains 6 protons and 7 electrons, thus 6 + charges and 7- charges with a net of 1- charge molecule – the sum of the oxidation numbers for each atom is the charge, e.g. sulfate is comprised of one S at a +6 oxidation number and four O’s each at a -2 oxidation number, thus sulfate has a 2- charge because 1S + 4O’s = (+6) + 4(-2) = -2 Oxidation Number the apparent charge of an atom in the molecule. Some oxidation numbers can often be found from the atom’s location on the periodic table, Group 1 is +1, Group 2 is +2, H is +1 (or -1 in hydrides), O is -2 (-1 is peroxides), in binary ionic compounds the halogens are -1. Otherwise the oxidation number is calculated . example: given, NaClO4, where Na = +1 (Group 1), O = -2, since Na + Cl + 4 O = 0, then Cl = +7. Note that for a single atom the charge is the oxidation number. Oxidation the loss of electrons in a chemical change; increase in oxidation number Reduction the gain of electrons in a chemical change; decrease in oxidation number Note that oxidation cannot happen without reduction and vice versa. Electrons Electron Spin from probability, electrons are said to spin up (↿) or spin down (⇂). Electron Pair (↿⇂) - combination of a spin up (↿) with a spin down (⇂). Pairing requires energy. Valence electrons electrons in outermost energy level. These are the electrons involved in bonding and reactions. Aufbau Principle lowest energy orbitals fill first with electrons (fill diagram from the bottom – up, lower energy state is preferred) Hund’s Rule of Multiplicity if two or more orbitals of equal energy are available, electrons will occupy them singly before filling them by pairing. Electron pairing requires energy, the lower energy state is always preferred so the electrons stay single. Following the axiom that the lower energy state is preferred the electrons will pair when the choice is pairing or moving to an orbital that is higher in energy Paule Exclusion Principle no two identical electrons can occupy the same orbital – means that only electrons of opposite spin may be in the same orbital Energy Level, n A discrete distance from the nucleus. The further the distance, the greater the energy needed for an electron to be located there. Electrons must gain a set amount of energy to move to a higher energy level further from the nucleus (the set amount of energy is said to be quantized). The energy is released as a set amount of energy if the electron moves closer to the nucleus to a lower level. Ground State All electrons in the lowest energy level possible. Excited State One or more electrons not at ground state. Orbitals A defined region (shape) of space, where it is most probable to find an electron. Each orbital contains 0, 1, or 2 e’s. There are four classes of orbitals: s, p d, f. Each class of orbital can have certain types. For instance the p-orbital has 3 types: px, py, pz. Each orbital is hourglass shaped and is aligned along an axis of space. s: 1 type, total of 2 e’s, 1 pr p: 3 types, total of 6 e’s, 3 pr’s d: 5 types, total of 10 e’s, 5 pr’s f: 7 types, total of 14 e’s, 7 prs Sublevel, l Indicates the class of orbital (s, p, d or f) present in the energy level. l = 0, …(n-1) Orbital type, m l Indicates the specific orbital (e.g. px, py, pz). m l = - l, …, + l 1 1 Electron spin, ms Electrons spin up (↑), ms = 2 , or down (↓ ) ms = - 2 (n)s (n)p (n-1)d (n-2)f Electron configuration states the arrangement of electrons within the electron cloud; includes the energy level, orbital type and number of electrons. examples: H = 1s1 N = 1s2 2s2 2p3 Notes - All families have the same valence electron configuration noble gas configuration ns2np6 halogen configuration ns2np5 chalcogen (O-family) configuration ns2np4 3 Atom Stability Nuclear Radioactivity the release of energy and/or particles resulting from an unstable nucleus There is no set rule for stability, but from experiment stability is based on n the neutron to proton ration, the further the value of p is from 1, the more likely the isotope is radioactive. 12 13 14 6 C 6 C 6 C 6 7 8 = 6 6 6 furthest from 1, so most likely radioacitve Nuclear Transformations a change in the number of protons and /or neutrons in the nucleus as a result of radioactive decay Half-life The time it takes for half of a sample of a radioactive isotope to decay. For example, the half-life of 32P is 14 days. So after 14 days a 50 g sample of 32P is now 25 g of 32P and 25 g of 32S. (see beta decay below) Types of radioactive decay 4 223 4 219 o , Alpha Particle, 2 He positive He nucleus ejected from the nucleus , 87 Fr 2 He + 85 At 0 - - 32 32 o Beta Decay, 1 e high energy e is ejected from the nucleus (n p + e ), 15 P + 16 S o , Gamma Rays high energy photon emitted as nucleus moves from excited to lower energy 232 132 state 90Th* 90Th + (*=excited state) 0 1 30 30 o Positron Emission, 1 e positive particle ejected from nucleus (p 0 n + ), 15 P + 14 Si o EC-electron capture e- falls into nucleus combining with a proton and forming a neutron, 202 0 202 81Tl + 1e 80 Hg 4 Periodic Table History Dmitri Mendeleev Wrote the 1st periodic table based on increasing atomic mass and similar properties. Left gaps where necessary in order to line-up families with similar properties. Predicted products of missing elements that, when discovered, would fill-in the gaps Henry Mosely Created the modern periodic table based on increasing atomic number Periodic Law The physical and chemical properties of the elements are periodic functions of their atomic number. Layout Period Horizontal rows A period is likened to an energy level when completing energy level diagrams. Moving left to right, the effective nuclear charge (the attraction between the valence electrons and the nucleus) increases, this causes the atomic radius to decrease, and electronegativity and ionization energy to increase.
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