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Chapter 1000A Spring 2014 Chapter 2 & 25 Atomic structure and Nuclear Chemistry History of the Atom – Greek philosophers – Thales: All matter was water – Empedocles: All matter was a fusion of opposites. – Aristotles: philosophical elements (air, earth, fire & water and principles (Hot, cold, moist & dry) ~2000 yrs – Leucippus and Democritus: Atomos: The point at which matter can no longer be subdivided. a = not & tomos = cut • Phlogistons (Bechner & Stahl) • Combustion • Metal burns • Wood burns • Joseph Priestly burned Hg with air in a sealed container making HgO. Priestly discovered O2 . Antoine Lavoisier (1743-1794): observed no change in mass. • 1785, Law of Conservation of Mass: . • 1794, Joseph Proust (1754 -1826): – Law of definite proportions: • In a given chemical compound the proportions of the elements that compose it are fixed integers (independent of its origin or preparation). – lead sulfide – Law of multiple proportions: • The same elements can combine in different proportions to form different chemical compounds (independent of its origin or preparation). – water & peroxide – carbon oxides I & II – (do calc Wt% - chalk talk) • 1808, John Dalton (1766-1844) adopts ideas of Democritus, Lavoisier & Proust to formulate Dalton’s Atomic theory: 1. All matter consist of solid and indivisible atoms 2. Atoms are indestructible and retain their identity in all chemical reactions. 3. All atoms of a given element are identical in mass and other properties. – *Different elements have different kinds of atoms and differ in mass from others. 4. Compounds consist of elements combined in small whole number ratios. • 1808, John Dalton (1766-1844) adopts ideas of Democritus, Lavoisier & Proust to formulate Dalton’s Atomic theory: 1. All matter consist of solid and indivisible atoms 2. Atoms are indestructible and retain their identity in all chemical reactions. 3. All atoms of a given element are identical in mass and other properties. 4. Compounds consist of elements combined in small whole number ratios. • Any theory is only as good as the current data/evidence. Table 2.2 Properties of the 3 Key Subatomic Particles Charge Name (Symbol) Relative Absolute (C)* Proton (p+) 1+ Neutron (n0) Electron (e-) Mass Relative (u)† Absolute (g) Location in Atom +1.60218x10-19 1.00727 1.67262x10-24 Nucleus 0 0 1.00866 1.67493x10-24 Nucleus 1- -1.60218x10-19 0.00054858 9.10939x10-24 Outside nucleus Charge = JJ Thomson – Eugen Goldstein – James Chadwick – * The † coulomb (C) is the SI unit of charge. The atomic mass unit (u) equals 1.66054x10-24 g. Rutherford’s a-scattering experiment and discovery of the atomic nucleus. Rutherford proposed that atoms consisted of tiny, massive, positive nucleus surrounded by electrons Most of an atom is empty space: (volume of 5 x 10-46 m3) • A tiny nucleus (about 1/1,000,000,000,000,000th of the atom’s volume) • contains most of the atom’s mass: protons and neutrons • bound together in a region of positive charge • electrons travel around the nucleus, balancing the overall charge of the atom. • Give analogy: Defining an element • The relative mass and charge of a particle is described by the notation: • An element . A = mass number Z = charge of the particle Example: electron proton neutron 0 e -1 1 p 1 1 n 0 A Z X • Every atom has an atomic number and a mass number: – Mass number (A) = # nucleons* C – Atomic number (Z) = # protons 12 6 *nucleon is a general term for a proton or neutron (i.e. a particle in the nucleus) • Atoms of the same element must have the same atomic number but different mass numbers are possible. • These are . Most elements have more than one naturally occurring isotope: H H H 1 2 3 1 1 1 12 13 14 6 C 6 C 6 C 12 Isotopes & applications • Isotope ratios • Climatology • A forensic application1 developed in Sweden – Atmospheric 14C levels increased sharply during testing of nuclear bombs C C C (1955-1963). – 14C levels decreased since banning of tests 12 13 14 6 6 6 K.L. Spalding, B.A. Buchholz, L.-E. Bergman, H. Druid and J. Frisén “Age Written in Teeth by Nuclear Tests” Nature 437, 333 - 334 (14 Sep 2005). 1 13 The amount of 14C incorporated in tooth enamel can be used to determine the year in which the enamel was formed and, therefore, the year of birth (to within 1.6 years). How Can We Measure Isotopes? • The average atomic mass for a chlorine atom is 35.4527 amu. How is the isotopes present in a sample determined? – Revisit the main difference – Take advantage of that difference 15 The end result is a spectrum showing the proportion of atoms in the sample belonging to each isotope: This is mass spectrometry! (the “MS” in GC/MS) 16 • Most elements occur as mixtures of several isotopes. These are treated as consisting of “averaged” atoms with “averaged” masses. • Atomic mass – Mav % abundance of isotope Misotope 100% Where % abundance # atoms of isotope 100% total # atoms of element Exceptions: 19F, 23Na, 31P 17 “Average” Samples • Silicon has three naturally occurring isotopes: – 92.23% 28Si (27.9769 u) – 4.67% 29Si (28.9765 u) – 3.10% 30Si (29.9738 u). First estimate then calculate the average atomic mass of silicon. 18 • First, the isotope present > % is • Ave. atomic weight Si = • Ave. atomic weight Si = • Ave. atomic weight Si = . Quick revisiting of the “mole” Concept. • How many atoms of Galium (Ga) are in 2.85 x 10-3 mol Ga? • 2.85 x 10-3 mol Ga x • = . atoms • Gallium has two naturally occurring isotopes and an average atomic mass of 69.723 u. 69Ga has an atomic mass of 68.926 u, and 71Ga has an atomic mass of 70.925 u. • First predict which isotope is more abundant • then calculate the natural abundance of each isotope of gallium. 21 69Ga 68.926 u the average amu Ga is 69.723 u 71Ga 70.925 u. • First calculate the differences from the average reported: 0.797u and 1,202 therefore it would be 69Ga. There are 2 variables , 2 equations are needed. MWGa = Or using mole concept & mole fraction MWGa = Also have the second equation x+y=1