Atoms and Elements
Electrostatics had been
discovered by the time of
Benjamin Franklin (late 1700’s)
Radioactivity
1896- Henri Becquerel:
Discovered uranium ore emitted
rays that exposed photographic
plates protected by black paper
Electrical Charge is fundamental to our understanding of the
atom
Suggested particles emitted through
atomic disintegration
1898- Marie/Pierre Curie: Isolated
polonium and radium which
emitted the same kind of rays
α, β and γ radiation
Contradicted Dalton’s atomic theory
which states that atoms are indivisible
Radiation Movie 1
Radiation Movie 2
Components of the Atom
Electrons
Primarily discovered by J.J. Thomson in
1897 by discovery of “negatively
charged” cathode rays using Cathode
Ray Tube (Crookes tube)
Determined charge to mass ratio of the
electron.
Thomson’s method for determining the
charge to mass ratio (q/m) for the electron
Thomson developed the “plum pudding”
model of the atom
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Robert Millikan (1909): Millikan’s Oil-drop
Experiment
Discovered the charge on the electron
(-1.602x10-19C)
Millikan’s
Experiment Movie
Eugen Goldstein (1886) – Discovered
positively charged “canal” rays
(positively charged nuclei of atoms, with
hydrogen having the smallest mass).
Protons and the Nucleus
1910- Ernest Rutherford: Gold foil
experiment
-Also experimented with alpha
irradiated gaseous elements: Alpha
deflection proportional to atomic mass
-Bombardment of nitrogen gas
produced particles consistent with that
of hydrogen (deemed a fundamental
particle)
-Developed Nuclear Model of the Atom
1919- Proton officially proclaimed
The Neutron
1932- James Chadwick: Discovered the
neutron by bombarding beryllium with
alpha particles
Gold Foil Movie 1
Gold Foil Movie 2
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Atomic Number and Atomic Mass
Isotopes
Atoms of the same element (i.e. the same
number of protons), that differ in their
number of neutrons.
(Compare to isobars (atoms with the
same mass number but different atomic
number) and isotones (atoms with the
same number of neutrons but different
number of protons)
Atomic Mass Unit (u or amu) = 1/12 the
mass of a C-12 isotope. (aka 1dalton (Da))
About the mass of a proton or neutron (1u =
1.661x10-24g)
The mass number of an isotope is not the sum
of the masses of the individual particles due to
the mass defect
(E =mc2), which is the binding energy.
Equations:
Percent abundance = (number of atoms of a
given isotope / total number of atoms of all
isotopes of that element)x100%
Atomic weight (weighted average)= (fractional
abundance of isotope 1)(mass of isotope 1) +
…
Isotopic Abundance is
Determined by mass spectroscopy
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Mass spectroscopy of a molecule
Sample Problems
a) Argon has three isotopes with 18, 20
and 22 neutrons, respectively. What are
the mass number and symbols of these
three isotopes?
b) Gallium has two isotopes:Ga-69 and
Ga-71. How many protons and neutrons
are in the nuclei of each of these
isotopes? If the abundance of Ga-69 is
60.1%, what is the abundance of Ga-71?
Answers:
36
18
Sample Problems
40
Ar 1838 Ar 18
Ar
69
31
Ga
71
31
Ga
p
31
31
n
38
40
e
31
31
1. There are three naturally occurring isotopes of
neon. Their percent abundances and atomic masses
are: neon-20, 90.51%, 19.99244u; neon-21, 0.27%,
20.99395u; neon-22, 9.22%, 21.99138u. Calculate
the weighted average atomic mass of neon.
2. The two naturally occurring isotopes of copper are
copper-63, mass 62.9298u, and copper-65, mass
64.9278u. What is the percent abundances of each
of the two isotopes? (The atomic mass of copper
listed on the periodic table is 63.546u)
Ga-71 39.9%
Answers:
(.9051)(19.99244u)+ (.0027)(20.99395u) + (.0922)(21.99138u)
= 20.1794 = 20.18u
1 mole ≡ The quantity of things as there
are atoms in exactly 12.0g of the C-12
isotope.
This must be determined experimentally.
(62.9298u)(X) + (64.9278u)(1-X) = 63.546u
62.9298X + 64.9278u – 64.9278X = 63.546u
-1.998X = -1.3818u
X = .6916
Cu-63 = 69.16%
Cu-65 = 30.84%
NA (Avogadro’s number) = 6.022x1023 =
mole
1
4
Mass number of an element
1) Taken in amu’s (u) = atomic mass (mass
of 1 atom)
2) Taken in grams (g) = molar mass (mass
of 1 mole of atoms)
24.30Mg
If you have 24.30g of Mg you have
6.022x1023 atoms of Mg or 1mole
Determining the molar mass of a
compound
Add the molar masses of the individual
elements that the compound contains.
Grams
(multiply by MM)
(divide by NA)
Moles
Basic Units (atoms, molecules, formula units, etc.)
(divide by MM)
(multiply by NA)
Sample Problem
a) What is the mass, in grams, of 1.5 mol
of silicon?
b) What amount (moles) of sulfur is
represented by 454g? How many atoms?
c) What is the average mass (in grams) of
one sulfur atom?
Answers:
a) g Si = 1.5mol Si (28.09g/mol) = 42.135g =
42g
b) Mol S = 454g (1mol/32.07g) =14.1565 =
14.2mol
# S atoms = 14.1565mol (6.022x1023) =
8.53x1024atoms
c) Mass of 1 sulfur atom = 32.07g/mol (1 mol /
6.022x1023atom)
= 5.325x10-23g
Sample problem
The density of gold, Au, is 19.32g/cm3.
What is the volume (in cm3) of a piece of
gold that contains 2.6x1024 atoms? If the
piece of metal is a square with a
thickness of 0.10cm, what is the length (in
cm) of one side of the piece?
Answers:
2.6x1024atoms (1mol/6.022x1023atoms)
(197.0g/1mol)(1cm3 / 19.32g) = 44.02 =
44cm3
X2(0.10cm) = 44cm3
X = 20.982 = 21cm
5
The Periodic Table
Dmitri Mendeleev: Russian Schoolteacher
Father of the modern periodic table (1869)
Table based on mass instead of atomic number (prior to
understanding of atomic particles)
Trends allowed for prediction of elements that had not yet
been discovered.
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Diatomic elements
Allotropes
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