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2.3 Atomic Theories
Greeks (5th Century B.C.) – coined the term
“atoms” to describe invisible particles of which
substances were composed
Aristotle (3rd Century B.C.) – believed the
universe was made of only 4 substances:
Earth, Air, Water and Fire
John Dalton (1803) – Atomic Theory of Matter




Matter composed of indivisible particles called atoms
Elements contain identical atoms
Different elements contain different atoms
Atoms can combine from two or more elements to
form new substances
J.J. Thompson (1897) – atoms contained
negatively charged particles called electrons;
envisioned a positive sphere with embedded
electrons; sphere had a net charge of “zero”;
termed the “Raisin Bun” Model
H.Nagaoka (1904) – envisioned a positive
sphere with a ring of electrons orbiting it (similar
to the rings of Saturn)
Ernest Rutherford (1914) – envisioned a very
small positively charged nucleus surrounded by
electrons; nucleus consisted of 1/1000th of the
total space of the atom
Niels Bohr (1921) – used Rutherford’s nuclear
model with electrons ‘quantized’ in specific
energy levels; became known as the BohrRutherford Model (looked similar to planets
orbiting the Sun)
Erwin Schrodinger (1926) - Quantum
Mechanics Theory – electrons were not in
definite places, rather in “probability clouds”;
similar to rotating fan blades
James Chadwick (1932) – nucleus of the atom
contained neutral particles called “neutrons”;
had the same mass as protons and shared the
nucleus with them
Isotopes
Frederick Soddy (1913) – discovered the
existence of isotopes
Isotopes are a form of the same element in
which the number of protons and electrons is the
same, but the number of neutrons is different
(example: carbon-12 and carbon-13)
In other words, isotopes have the same atomic
number but different atomic mass
Carbon-12
Carbon-13

#p
#e
atomic natural
#n mass abundance
6
6
6
6
6
7
Average atomic mass:
Bromine-79 35 35 44
Bromine-81 35 35 46

Average atomic mass:
12
13
98.89%
1.11%
12.011 a.m.u.
79
81
50.69%
49.31%
79.904 a.m.u.
Bohr’s Theory of Atomic Structure
Each electron in an atom have a fixed amount of
energy related to the circular orbit in which it is
found
Electrons cannot exist between orbits, but they
can move into unfilled orbits if a “quantum” of
energy is absorbed or released
The higher the energy level, the further it is from
the nucleus
The maximum number of electrons in the first
three levels is: 2, 8, 8
Example: aluminum

Atomic number: 13 (13 protons p+, 13 electrons e-)

Electrons must be distributed amoungst 3 orbits around the
nucleus using the 2,8,8 rule

Diagram:
Al

3e8e2e13p+
(3rd level – “valence” level)
(2nd level)
(1st level)
(nucleus)
Formation of Monatomic Ions
Ions – atoms which have either gained or lost electrons
to become stable; unlike atoms, ions always have a net
charge
The reason atoms gain or lose electrons to become ions
is to attain a filled outermost (valence) shell
Metals typically lose electrons to become positively
charged (+); while non-metals typically gain electrons to
become negatively charged (-)
We will limit our discussion to the first 20 elements for
simplicity reasons
Metal Ions
Group 1 metals (e.g. Li, Na, K) donate one valence
electron to become +1 ions
donates to a non-metal to become…
E.g.

sodium atom
1e8e2e11p+
Net
Charge:
0
Symbol: Na
sodium ion
(3rd level)
(2nd level)
(1st level)
(nucleus)
8e2e11p+
+1
Na+
Metal Ions
Group 2 metals (e.g. Be, Mg, Ca) donate two valence
electrons to become +2 ions
donates to a non-metal to become…
E.g. magnesium atom

2e(3rd level)
8e(2nd level)
2e(1st level)
12p+
(nucleus)
Net
Charge: 0
Symbol: Mg
magnesium ion
8e2e12p+
+2
Mg2+
Transition Metals
Transition metals (groups 3-12) are very
different from other metals in that their
charges are much less predictable and
often can have more than one ion charge
(e.g. copper ions - Cu+, Cu2+)
Metal Ions
Group 13 metals (e.g. Al) donate three valence electrons
to become +3 ions
donates to a non-metal to become…
E.g. aluminum atom

3e(3rd level)
8e(2nd level)
2e(1st level)
13p+
(nucleus)
Net
Charge: 0
Symbol: Al
aluminum ion
8e2e13p+
+3
Al3+
A note about Group 14
Since there are no metals in group 14
within the first 20 elements, we will move
our discussion to non-metals
Note: There are 3 metals in group 14
beyond the first 20 elements (Ge, Pb, Sn);
however, their ion charges are somewhat
unpredictable. We will treat them similar to
the transition metals and look up their
charges instead of trying to predict them
Non-metal Ions
Group 15 non-metals (e.g. N, P) accept three valence
electrons to become -3 ions
accepts electrons from a metal to become…
E.g. phosphorus atom

5e(3rd level)
8e(2nd level)
2e(1st level)
15p+
(nucleus)
Net
Charge: 0
Symbol:
P
phosphide ion
8e8e2e15p+
-3
P3-
Non-metal Ions
Group 16 non-metals (e.g. O, S, Se) accept two valence
electrons to become -2 ions
accepts electrons from a metal to become…
E.g. sulfur atom

6e8e2e16p+
Net
Charge: 0
Symbol:
S
(3rd level)
(2nd level)
(1st level)
(nucleus)
sulfide ion
8e8e2e16p+
-2
S2-
Non-metal Ions
Group 17 non-metals (e.g. F, Cl, Br, I) accept one
valence electron to become -1 ions
accepts an electron from a metal to become…
E.g. chlorine atom

7e8e2e17p+
Net
Charge: 0
Symbol:
Cl
(3rd level)
(2nd level)
(1st level)
(nucleus)
chloride ion
8e8e2e17p+
-1
Cl-
Noble Gases
Group 18 elements (e.g. He, Ne, Ar, Kr, Xe, Rn) were
“born happy” will a filled outermost shell and therefore do
not react with anyone
E.g. argon atom

8e8e2e18p+
(3rd level)
(2nd level)
(1st level)
(nucleus)
A note about hydrogen…
Hydrogen is unique in that it can either GAIN or LOSE an electron to
become stable
donates to a non-metal to become…
E.g.

hydrogen atom
ep+
hydrogen ion
(1st level)
(nucleus)
Charge: 0
Symbol: H
p+
+1
H+
accepts an electron from a non-metal to become…
E.g.

hydrogen atom
ep+
Charge: 0
Symbol: H
(1st level)
(nucleus)
hydride ion
2ep+
-1
H-
Homework
Worksheet #4.
Ionic Compounds
Ionic compounds are formed when
metals donate electrons to non-metals
Metals are left with a positive charge and
are called cations (e.g. Na+, Mg2+ )
Non-metals are left with a negative charge
and are called anions (e.g. Cl-, N3-)
Ionic Compounds
Group 1 elements (Li, Na, K) react very readily with Group 17
elements (F, Cl, Br, I) because an exchange of one electron results
in both ions having a filled outermost shell; e.g.
sodium (atom) + chlorine (atom)
Na
Cl
e8e2e11p+
7e8e2e17p+
sodium (ion) + chloride (ion)
Na +
Cl -
8e2e11p+
8e8e2e17p+
Ionic Compounds
Group 2 elements (Be, Mg, Ca) react very readily with Group 16
elements (F, Cl, Br, I) because an exchange of two electrons results
in both ions having a filled outermost shell; e.g.
calcium (atom) + oxygen (atom)
Ca
O
2e8e8e2e20p+
6e2e8p+
calcium (ion) +
Ca 2+
8e8e2e20p+
oxide (ion)
O 2-
8e2e8p+