2 Atomic Theory Development of Theory

Transcription

1 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 Max Planck J.J. Thompson Hantaro Nagaoka 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 (803) first postulated that energy is quantized, originator of quantum theory (900) plum-pudding model negative electrons (plums) are located in a positively charged pudding (903) Saturnian model large nucleus with electrons orbiting in rings (904) Robert Millikan measured the charge and mass of an electron (908) Ernest Rutherford Neils Bohr atom has a small, positive, central nucleus, which contains the mass and is surrounded by a cloud of negative electrons [correct model] (9) Discovered the proton. (97) planetary model the nucleus is surrounded by electrons orbiting in rings (93) Edwin Schroedinger mathematical wave equation led to prediction of the possible states for an electron [correct, part of orbital theory] (96) Werner Heisenberg James Chadwick uncertainty principle correctly states that it is impossible to predict the exact position and momentum of an electron (97) discovered the neutron which served as a catalyst for the growth of nuclear physics (93) Rutherford Experiment method: involved shooting alpha particles (He + ) at a sheet of gold foil postulate: results: the deflection of the alpha particles would determine the location of the mass within the atom 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 u neutron neutral ( ) located in nucleus u

2 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 3 Atomic Number, Z Isotope Percent Abundance Average Atomic Mass, Y avg Charge Ion Cation Anion The number of protons. This defines the element. For example carbon always has 6 protons, but is known to have 6 neutrons (Y=) or 7 neutrons (Y=3) or 8 neutrons (Y=4) See isotopes below. Atoms with the same number of protons (same element), but with different number of neutrons. the percentage of one isotope for an element a weighted average of all known isotopic masses for an element Y Y Y where X = percent abundance as a decimal avg Y and Y are isotopic masses 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 ( ). a charged atom or molecule positive ion, lost electrons 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 - 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 - oxidation number, thus sulfate has a - charge because S + 4O s = (+6) + 4(-) = - Oxidation Number Oxidation Reduction 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 is +, Group is +, H is + (or - in hydrides), O is - (- is peroxides), in binary ionic compounds the halogens are -. Otherwise the oxidation number is calculated. example: given, NaClO 4, where Na = + (Group ), O = -, since Na + Cl + 4 O = 0, then Cl = +7. Note that for a single atom the charge is the oxidation number. the loss of electrons in a chemical change; increase in oxidation number 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 ( ).

3 Electron Pair ( ) - combination of a spin up ( ) with a spin down ( ). Pairing requires energy. Valence electrons Aufbau Principle Hund s Rule of Multiplicity Paule Exclusion Principle Energy Level, n electrons in outermost energy level. These are the electrons involved in bonding and reactions. lowest energy orbitals fill first with electrons (fill diagram from the bottom up, lower energy state is preferred) 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 no two identical electrons can occupy the same orbital means that only electrons of opposite spin may be in the same orbital 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 Excited State Orbitals All electrons in the lowest energy level possible. One or more electrons not at ground state. A defined region (shape) of space, where it is most probable to find an electron. Each orbital contains 0,, or 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: p x, p y, p z. Each orbital is hourglass shaped and is aligned along an axis of space. s: type, total of e s, pr p: 3 types, total of 6 e s, 3 pr s d: 5 types, total of 0 e s, 5 pr s f: 7 types, total of 4 e s, 7 prs Sublevel, l Orbital type, m l Indicates the class of orbital (s, p, d or f) present in the energy level. l = 0, (n-) Indicates the specific orbital (e.g. p x, p y, p z ). m l = - l,, + l Electron spin, m s Electrons spin up ( ), m s =, or down ( ) m s = - (n)s (n)p (n-)d (n-)f Electron configuration states the arrangement of electrons within the electron cloud; includes the energy level, orbital type and number of electrons. examples: H = s N = s s p 3

4 Notes - All families have the same valence electron configuration noble gas configuration ns np 6 halogen configuration ns np 5 chalcogen (O-family) configuration ns np 4 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 the neutron to proton ration, the further the value of p n is from, the more likely the isotope is radioactive. 6 C 3 6 C 4 6 C n = p furthest from, so most likely radioacitve Nuclear Transformations Half-life a change in the number of protons and /or neutrons in the nucleus as a result of radioactive decay The time it takes for half of a sample of a radioactive isotope to decay. For example, the half-life of 3 P is 4 days. So after 4 days a 50 g sample of 3 P is now 5 g of 3 P and 5 g of 3 S. (see beta decay below) Types of radioactive decay o, Alpha Particle, He positive He nucleus ejected from the nucleus, 87 Fr He + 85 At o Beta Decay, 0 e high energy e - is ejected from the nucleus (n p + e - 3 ), 5 P 0 3 e + 6 S o, Gamma Rays o Positron Emission, o EC-electron capture high energy photon emitted as nucleus moves from excited to lower energy 3 3 state 90Th * 90 Th + (*=excited state) 0 e positive particle ejected from nucleus (p 0 n e ), 5 P 0 30 e + Si e - falls into nucleus combining with a proton and forming a neutron, 0 8 Tl e 80 Hg 4 4 Periodic Table History Dmitri Mendeleev Henry Mosely Periodic Law Wrote the st 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 Created the modern periodic table based on increasing atomic number The physical and chemical properties of the elements are periodic functions of their atomic number.

5 Layout Period Group/Family Trends Electron shielding Effective nuclear charge 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. A vertical column Elements in the same family have the same valence e-config, and thus similar properties When moving down a group the distance (# of energy levels) between the nucleus and the valence electrons increases causing the attraction between them to decrease, so atomic radius increases down a group while the electronegativity and ionization energy decrease. the masking of the nucleus by the kernel electrons. Shielding is constant within a period, but grows down a group the charge felt by each valence electron. Calculated by protons kernel electrons Increases left to right across a period, but is constant in a group Electronegativity the ability to attract electrons in a covalent bond trend = First Ionization Energy the energy needed to remove one electron trend = Atomic Radius distance from the nucleus to the valence energy level trend = examples: Which is more electronegative, K or Cl? ans = Cl Which has the larger atomic radius, S or As? ans = As