Download 1 - Atomic Theory - Crestwood Local Schools

An ATOM is the
smallest unit of an
element that maintains
the properties of that
element.
• Postulated existence of atoms and
void (makes change possible)
• Elements differ in shape, position,
and arrangement
• Believed that atoms are:
–
–
–
–
Microscopic
Indestructible
Entirely solid (no holes or gaps)
Homogeneous with no internal
structure
• Aristotle and other well-known
Greek philosophers didn’t believe
Democritus
• The late 1700’s –definitions and basic laws had been
discovered and accepted by chemists.
– Element – substance that cannot be broken down by
ordinary chemical means.
– Chemical Reaction – transformation of substance or
substances into one or more new substances.
Mass cannot be created or destroyed just changed from one
form to another. (Antoine Lavosier)
A chemical compound contains exactly the same elements in
the same proportion regardless of sample size. (Joseph Proust
from work of Gay-Lussac & Amadeo Avogadro – 1802/1804)
If two or more different compounds are composed of the
same two elements, then the ratio of the masses of those
elements will always exist as a ratio of small whole numbers.
(John Dalton – 1808)
1. All elements are composed of tiny
indivisible particles called atoms.
2. Atoms of the same element are identical.
3. The atoms of one element are different
from the atoms of another element.
4. Atoms combine in simple whole-number ratios.
5. Atoms are separated, joined or rearranged in
chemical reactions. Atoms of one element
are never changed into atoms of another
element as a result of a chemical reaction.
• Rutherford, Geiger & Marsden (1912) -showed that most of
the atom was empty space, but that atoms had a solid,
positive core.
• Most particles passed through the foil undisturbed (black arrows).
• A few were deflected (red arrows).
• Rutherford reasoned that each atom in the foil contained a small,
dense, positively charged nucleus surrounded by electrons.
• Rutherford, Geiger &
Marsden (1912) -showed
that most of the atom was
empty space, but that atoms
had a solid, positive core.
Rutherford’s
Nuclear Atom
Thomson’s “Plum Pudding”
Model of the Atom
Electrons
Uniform, positively
charged sphere
Nucleus
Electron Cloud
• Atomic Number
– Represents the number of protons
– Defines the identity of the element
• Mass Number
– Represents the total number
of protons and neutrons
• Atomic Mass
– average mass of the isotopes
(based on relative frequency in nature)
Proton
+
Proton
• Symbol: p+
• Charge: 1+
Neutron
Neutron
• Symbol: n0
• Charge: 0
Electron
e-
Electron
• Symbol:
• Charge: 1-
• Location: Nucleus
• Relative Mass: ~1 amu
• Mass: 1.67 x 10-24g
• Location: Nucleus
• Relative Mass: ~1 amu
• Mass: 1.67 x 10-24g
• Location: Electron cloud
• Relative Mass: 1/1840
amu
An ISOTOPE is an atom with a
different number of Neutrons and
therefore a different atomic mass.
Example: C-12 vs. C-14
An ION is an atom that has lost or
gained electrons, resulting in a
positively or negatively charged
particle.
Atomic Number = # of protons
8
O
Atomic Symbol
Mass Number = #
protons & neutrons
(round to 16)
15.999
Mass Number (unrounded) =
average mass of the isotopes
o A scale designed for atoms gives their small atomic
masses in atomic mass units (amu)
o An atom of 12C was assigned an exact mass of 12.00
amu
o Relative masses of all other atoms was determined by
comparing each to the mass of 12C
o An atom twice as heavy has a mass of 24.00 amu. An
atom half as heavy is 6.00 amu.
o Listed on the periodic table
o Gives the mass of “average” atom
of each element compared to 12C
o Average atom based on all the
isotopes and their abundance %.
o Atomic mass is not a whole
number due to isotopes.
Na
22.99
Oxygen-16
Oxygen-17
Oxygen-18
• 8 protons
• 8 electrons
• 8 neutrons
• Mass = 16
• 8 protons
• 8 electrons
• 9 neutrons
• Mass = 17
• 8 protons
• 8 electrons
• 10 neutrons
• Mass = 18
 Percent(%) abundance of isotopes
 Mass of each isotope of that element
 Weighted average =
mass isotope1(%) + mass isotope2(%) + …
100
100
Isotopes
Mass of Isotope
Abundance
24Mg
=
24.0 amu
78.70%
25Mg
=
25.0 amu
10.13%
26Mg
=
26.0 amu
11.17%
Atomic mass (average mass) Mg = 24.3 amu
Mg
24.3
The element copper has naturally occurring
isotopes with mass numbers of 63 and 65.
Isotope
Atomic Mass
Relative Abundance
Cu-63
63 amu
69.2%
Cu-65
65 amu
30.8%
Calculate the average atomic mass of copper.
Cu-63:
(63 amu)(0.692) = 43.596 amu
Cu-65:
(65 amu)(0.308) = 20.02 amu
Total:
63.616 amu
Naturally occurring boron is 80.20% B-11 and 19.80% of a
different isotope. What must the mass of this isotope be if
the average atomic mass of boron is 10.81 amu?
Isotope
Atomic Mass
Relative Abundance
B-11
11 amu
80.20%
B-??
? amu
19.80%
Total:
=
10.81 amu
B-11:
(11 amu)(0.8020)
=
8.822 amu
B-X:
(X amu)(0.1980)
=
1.988 amu
X = (1.988 amu)/(0.1980) = 10.04 amu
Other isotope = B-10!!!