Download Chapter 4 - Atomic Structure - A

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
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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
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Early Models of the Atom
4.1
Dalton’s Atomic Theory
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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
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Calculating Atomic mass 4.3
Must know 3 values to determine
atomic mass based on amount
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3.
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# 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