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Chapter 4 Atomic Structure Chemistry 2 Defining the Atom 4.1 Early Models of the Atom 4.1 • The smallest particle of an element that retains its identity in a chemical reaction • All matter composed of • Atomos = indivisible • Democritus’s Atomic Philosophy – – – – – Greek philosopher (460 B.C. – 370 B.C) 1st to suggest atom Indivisible and indestructible Did not explain chemical behavior Did not use scientific method • Early Models of the Atom 4.1 Dalton’s Atomic Theory – – – 1. 2. 3. 4. 2000 yrs. Later John Dalton (1766 -1844) English chemist & schoolteacher Used experimental methods to develop a theory All elements composed of tiny indivisible particles = atoms Atoms in the same element are identical; atoms from 1 element are different form atoms of another element Atoms of different elements can physically mix or chemically combine in whole-number ratios Chemical reactions = atoms joined, separated, or rearranged. Atoms of 1 element NEVER change into atoms of another element Sizing up the Atom 4.1 • Pure copper coin about a penny size = 2.4 x 1022 atoms – …earth’s population = 6 x 109… – There are about 4 x 1012 times as many atoms in a penny as there are people on earth • Radii of most atoms = 5 x 10-11 – 2 x 10-10m http://broadeducation.com/htmlDemos/ AbsorbChem/HistoryAtom/page.htm Structure of the Nuclear Atom 4.2 Subatomic Particles 4.2 • Change in Dalton’s atomic theory = atoms are divisible = electrons, protons, and neutrons • Electrons – negatively charged particle – J.J. Thomson 1856-1940 – English physicist – Experiment involved passing electric current thru gases at low pressure sealed glass tube with electrodes (metal disks) anode, + electrode, and cathode, - electrode = glowing beam, cathode ray (beam traveled from cathode to anode) • Beam attract positive charge = ray is a stream of tiny negatively charged particle moving at a high speed (Figure 4.5) – Robert A. Millikan 1868-1953 • Calculated mass of electron 1/1840 of a H atom Subatomic Particle 4.2 • Eugene Goldstein (1850 – 1930) – Found positively charged particle in 1886 – Observed cathode-ray tube found rays traveling in opposite direction = PROTON • Mass = 1840 xs electron • James Chadwick (1891 - 1974) English Physicist – Found the neutron = no charge, mass nearly equal to proton The Atomic Nucleus 4.2 • Many theories of structure of atom – JJ Thomson = plum pudding model • Electrons stuck into a lump of + charge (similar to raisins stuck in dough!) • Ernest Rutherford (1871-1937) - Student of Thomson • Rutherford’s Gold-Foil Experiment – Used He atoms that lost 2 electrons double + charge (massive alpha particles) – Directed narrow beam of alpha a particles at a thin sheet of gold foil Rutherford’s Atomic Model 4.2 • According to theory: particles should pass thru w/ slight deflection • HOWEVER, majority passed thru w/o deflection and some bounced back at large angle • Rutherford Concluded: 1)atoms mostly empty space (lack of deflection 2)Most mass concentrated in small region (deflection) – Called region the NUCLEUS, composed of p & n • Model = NUCELAR ATOM – p & n located in nucleus, small – E distributed around nucleus = occupy most of volume – Ex of atom: football stadium, nucleus = marble Technology and Society • Ernest Ruska & Max Knoll (1931) built 1st e microscope – Uses e beam and lenses of magnetic or electric field – Can magnify up to 100,000 times • Light microscope can magnify 1000 times • Biochemistry = uses to look at DNA molecules • Microelectronics = use to measure and analyze characteristics of microcircuits Distinguishing Among Atoms 4.3 Atomic Number 4.3 • Number of p in the nucleus – Ex: H – 1 p = atomic number of 1 – # of p = # of e when neutral Mass Number 4.3 • Total # of p + n – Hint: most of the mass of an atom is where? – Nucleus – What's in the nucleus? – p + n = most of the mass of an atom = MASS # • # of n = mass # - atomic # • # of n = (p +n) – p Isotopes 4.3 • Same # of p and e, different # of n • Different mass # (mass # = p + n) • Chemically alike (for the most part) b/c p & e are responsible for chemical behavior Atomic Mass 4.3 • Weighted average mass of the atoms in a naturally occurring sample of the element • Use a mass spectrometer to find mass of atom – Mass of p or n = 1.67x 10-24 g – Mass of e = 9.11 x 10-28 g – TOO SMALL TO WORK WITH • Compare relative masses of atoms w/ reference isotope, carbon-12 (6p + 6n) – Atomic Mass Unit (amu) – 1/12 of the mass of carbon 12atom • Mass of proton or neutron is 1/12 of 12 amu, or about 1 amu – Atomic Mass is normally not a whole number b/c different # of isotopes • Calculating Atomic mass 4.3 Must know 3 values to determine atomic mass based on amount 1. 2. 3. – # of stable isotopes Mass of isotopes Natural % abundance of each isotope Multiply mass of each isotope by natural abundance add products The Periodic Table – A Preview 4.3 • An arrangement of elements separated into groups based on properties • Elements listed in order of increasing atomic number • 7 Periods = horizontal line – Properties of elements vary • Group (family) = vertical line – Similar chemical and physical properties Archaeologist • Detectives of the past –sift for clues of past civilization • Excavate ancient cities, artifacts, tools • Sometimes draw conclusions from indirect clues • Use radiometric dating (sample dated by measuring concentration of certain isotopes • Chemical tests to determine artifacts composition • Need background in history and science