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Atoms, Molecules, and Ions
Chemistry Timeline #1
B.C.
400 B.C. Demokritos and Leucippos use the term "atomos”

2000 years of Alchemy
1500's
 Georg Bauer: systematic metallurgy
 Paracelsus: medicinal application of minerals
1600's
Robert Boyle:The Skeptical Chemist. Quantitative experimentation, identification of
elements
1700s'
 Georg Stahl: Phlogiston Theory
 Joseph Priestly: Discovery of oxygen
 Antoine Lavoisier: The role of oxygen in combustion, law of conservation of
mass, first modern chemistry textbook
Chemistry Timeline #2
1800's
Joseph Proust: The law of definite proportion (composition)
 John Dalton: The Atomic Theory, The law of multiple proportions
Joseph Gay-Lussac: Combining volumes of gases, existence of diatomic molecules
Amadeo Avogadro: Molar volumes of gases
Jons Jakob Berzelius: Relative atomic masses, modern symbols for the elements
 Dmitri Mendeleyev: The periodic table
 J.J. Thomson: discovery of the electron
 Henri Becquerel: Discovery of radioactivity
1900's
 Robert Millikan: Charge and mass of the electron
 Ernest Rutherford: Existence of the nucleus, and its relative size
 Meitner & Fermi: Sustained nuclear fission
 Ernest Lawrence: The cyclotron and trans-uranium elements
Laws
• Conservation of Mass
• Law of Definite Proportion –
– compounds have a constant composition.
– They react in specific ratios by mass.
• Multiple Proportions-
– When two elements form more than one
compound, the ratios of the masses of the
second element that combine with one
gram of the first can be reduced to small
whole numbers.
Proof
• Mercury has two oxides.
– One is 96.2 % mercury by mass, the
other is 92.6 % mercury by mass.
• Show that these compounds follow
the law of multiple proportion.
• Speculate on the formula of the two
oxides.
Dalton’s Atomic Theory (1808)
 All matter is composed of extremely
small particles called atoms
 Atoms of a given element are
identical in size, mass, and other
properties; atoms of different
John Dalton
elements differ in size, mass, and
other properties
 Atoms cannot be subdivided, created, or destroyed
 Atoms of different elements combine in simple
whole-number ratios to form chemical compounds
 In chemical reactions, atoms are combined,
separated, or rearranged
Modern Atomic Theory
Several changes have been made to Dalton’s theory.
Dalton said:
Atoms of a given element are identical in
size, mass, and other properties; atoms of
different elements differ in size, mass, and
other properties
Modern theory states:
Atoms of an element have a characteristic
average mass which is unique to that
element.
Modern Atomic Theory #2
Dalton said:
Atoms cannot be subdivided, created, or destroyed
Modern theory states:
Atoms cannot be subdivided, created, or destroyed
in ordinary chemical reactions. However, these
changes CAN occur in nuclear reactions!
Atomic Particles
Particle
Charge
Mass (kg)
Location
Electron
-1
9.109 x 10-31
Electron
cloud
Proton
+1
1.673 x 10-27
Nucleus
0
1.675 x 10-27
Nucleus
Neutron
The Atomic
Scale
 Most of the mass of the
atom is in the nucleus
(protons and neutrons)
 Electrons are found
outside of the nucleus (the
electron cloud)
 Most of the volume of
the atom is empty space
“q” is a particle called a “quark”
About Quarks…
Protons and neutrons are
NOT fundamental particles.
Protons are made of
two “up” quarks and
one “down” quark.
Neutrons are made of
one “up” quark and
two “down” quarks.
Quarks are held together
by “gluons”
Isotopes
Isotopes are atoms of the same element having
different masses due to varying numbers of neutrons.
Isotope
Protons
Electrons
Neutrons
Hydrogen–1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
1
1
Hydrogen-3
(tritium)
1
1
2
Nucleus
Atomic Masses
Atomic mass is the average of all the naturally
isotopes of that element.
Carbon = 12.011
Symbol
Composition of
the nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
6 protons
7 neutrons
1.11%
Carbon-14
14C
6 protons
8 neutrons
<0.01%
Isotope
Molecules
Two or more atoms of the same or different
elements, covalently bonded together.
Molecules are discrete structures, and their
formulas represent each atom present in the
molecule.
Benzene, C6H6
Covalent Network Substances
Covalent network substances have covalently
bonded atoms, but do not have discrete
formulas.
Why Not??
Graphite
Diamond
Ions
 Cation: A positive ion
• Mg2+, NH4+
 Anion: A negative ion
 Cl-, SO42-
 Ionic Bonding: Force of attraction between
oppositely charged ions.
 Ionic compounds form crystals, so their
formulas are written empirically (lowest whole
number ratio of ions).
Periodic Table with Group Names
This slide contains
classified material and
cannot be shown to high
school students. Please
continue as if everything is
normal.
Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube
to deduce the presence of a negatively charged
particle.
Cathode ray tubes pass electricity through a gas
that is contained at a very low pressure.
Thomson’s Atomic Model
Thomson believed that the electrons were like plums
embedded in a positively charged “pudding,” thus it was
called the “plum pudding” model.
Rutherford’s Gold Foil Experiment
 Alpha particles are helium nuclei
 Particles were fired at a thin sheet of gold foil
 Particle hits on the detecting screen (film) are
recorded
The Puzzle of the Atom
 Protons and electrons are attracted to each
other because of opposite charges
 Electrically charged particles moving in a
curved path give off energy
 Despite these facts, atoms don’t collapse
Electromagnetic radiation propagates
through space as a wave moving at the
speed of light.
c = 
C = speed of light, a constant (3.00 x 108 m/s)
 = frequency, in units of hertz (hz, sec-1)
 = wavelength, in meters
Types of electromagnetic radiation:
Long
Wavelength
=
Low Frequency
=
Low ENERGY
Short
Wavelength
=
High Frequency
=
High ENERGY
Wavelength Table
Toupee?
The Wave-like Electron
The electron propagates
through space on an energy
wave. To understand the
atom, one must understand
the behavior of
electromagnetic waves.
Louis deBroglie
The Great Niels Bohr
(1885 - 1962)
Spectroscopic analysis of the visible spectrum…
…produces all of the colors in a continuous spectrum
Spectroscopic analysis of the hydrogen
spectrum…
…produces a “bright line” spectrum
Electron transitions
involve jumps of
definite amounts of
energy.
This produces bands
of light with definite
wavelengths.
Bohr Model Energy Levels
Schrodinger Wave Equation

d
V 
8  m dx
h
2
2
2
2
 E
Equation for probability of a
single electron being found
along a single axis (x-axis)
Erwin Schrodinger
Heisenberg Uncertainty Principle
“One cannot simultaneously
determine both the position
and momentum of an electron.”
You can find out where the
electron is, but not where it
is going.
Werner
Heisenberg
OR…
You can find out where the
electron is going, but not
where it is!
Quantum Numbers
Each electron in an atom has a unique
set of 4 quantum numbers which describe
it.
 Principal quantum number
(n)
 Angular momentum quantum number (l)
 Magnetic quantum number (m)
 Spin quantum number
(s)
Principal Quantum Number
Generally symbolized by n, it denotes the shell
(energy level) in which the electron is located.
Number of electrons
that can fit in a shell:
2n2
Angular Momentum Quantum Number
The angular momentum quantum number, generally
symbolized by l, denotes the orbital (subshell) in
which the electron is located.
l =3
f
Magnetic Quantum Number
The magnetic quantum number, generally
symbolized by m, denotes the orientation of the
electron’s orbital with respect to the three axes in
space.
Assigning the Numbers
 The three quantum numbers (n, l, and m) are
integers.
 The principal quantum number (n) cannot be
zero.
 n must be 1, 2, 3, etc.
 The angular momentum quantum number (l )
can be any integer between 0 and n - 1.
 For n = 3, l can be either 0, 1, or 2.
 The magnetic quantum number (ml) can be any
integer between -l and +l.
 For l = 2, m can be either -2, -1, 0, +1, +2.
Principle, angular momentum, and magnetic
quantum numbers: n, l, and ml
Pauli Exclusion Principle
No two electrons in an atom
can have the same four
quantum numbers.
Wolfgang
Pauli
Spin Quantum Number
Spin quantum number denotes the behavior
(direction of spin) of an electron within a
magnetic field.
Possibilities for electron spin:
1

2
1
2
An orbital is a region within an atom where there
is a probability of finding an electron. This is a
probability diagram for the s orbital in the first
energy level…
Orbital shapes are defined as the surface that
contains 90% of the total electron probability.
Sizes of s orbitals
Orbitals of the same shape (s, for instance) grow
larger as n increases…
Nodes are regions of low probability within an
orbital.
Orbitals in outer energy levels DO penetrate into
lower energy levels. Penetration #1
This is a probability
Distribution for a
3s orbital.
What parts of the
diagram correspond
to “nodes” – regions
of zero probability?
The s orbital has a spherical shape centered around
the origin of the three axes in space.
s orbital shape
P orbital shape
There are three peanut-shaped p orbitals in
each energy level above n = 1, each assigned to
its own axis (x, y and z) in space.
d orbital shapes
Things get a bit more
complicated with the five d
orbitals that are found in
the d sublevels beginning
with n = 3. To remember
the shapes, think of:
“double peanut”
…and a “peanut
with a donut”!
Shape of f orbitals
Things get even more
complicated with the seven f
orbitals that are found in the
f sublevels beginning with n
= 4. To remember the
shapes, think of:
Flower
Orbital filling table
Element
Lithium
Configuration
notation
1s22s1
[He]2s1
____
1s
Beryllium
____
____
2p
____
____
2s
____
____
2p
____
[He]2s2p2
____
2s
____
____
2p
____
1s22s2p3
[He]2s2p3
____
2s
____
____
2p
____
1s22s2p4
[He]2s2p4
____
2s
____
____
2p
____
1s22s2p5
[He]2s2p5
____
1s
Neon
____
2s
1s22s2p2
____
1s
Fluorine
____
[He]2s2p1
____
1s
Oxygen
____
2p
1s22s2p1
____
1s
Nitrogen
____
[He]2s2
____
1s
Carbon
____
2s
1s22s2
____
1s
Boron
Noble gas
notation
Orbital notation
____
2s
____
____
2p
____
1s22s2p6
[He]2s2p6
____
1s
____
2s
____
____
2p
____
Electron configuration of the
elements of the first three series
Irregular confirmations of Cr and Cu
Chromium steals a 4s electron to half
fill its 3d sublevel
Copper steals a 4s electron to FILL
its 3d sublevel
energy is emitted
2. energy is absorbed
3. no change in energy occurs
4. light is emitted
5. none of these
1.
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5.
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4.
ow
3.
s
2.
gamma rays
microwaves
radio waves
infrared radiation
x-rays
ra
y
1.
18
19
20
1. 2
2. 5
3. 10
4. 18
1
2
3
4
5
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13
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6
0%
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0%
5
2
0%
10
5. 6
20
1
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0%
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0%
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42
0½
½
–½
½
–½
½
43
–2
–½
0
0
–1
–2
0
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–1
½
1
0
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2
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0–
½
s
11
0½
m
s
1
2. 3
3. 2
4. 4
5. 4
1.
l
nlm
n
20
1s22s22p63s23p64s23d104p65s24d105p15d10
2. 1s22s22p63s23p64s23d104d104p1
3. 1s23s22p63s23p64s24d104p65s25d105p1
4. 1s22s22p63s23p64s23d104p65s24d105p1
5. none of these
1.
1
2
3
4
5
6
7
8
9
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21
22
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13
14
15
0%
0%
0%
0%
1s
22
s2
2p
63
s2
3p
1s
.. .
22
s2
2p
63
s2
3p
1s
.. .
23
s2
2p
63
s2
3p
1s
.. .
22
s2
2p
63
s2
3p
.. .
no
ne
of
th
es
e
0%
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20
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Periodicity
}
Atomic
Radius
Radius = half the distance between
two nuclei of a diatomic molecule.
 Influenced
by three factors.
 Energy Level
› Higher energy level is further away.
 Charge
on nucleus
› More charge pulls electrons in closer.
 Shielding
› Layers of electrons shield from nuclear
pull.

The electron on the
outside energy level
has to look through all
the other energy levels
to see the nucleus
The electron on the
outside energy level
has to look through
all the other energy
levels to see the
nucleus.
 A second electron
has the same
shielding.

As we go down a
group
 Each atom has
another energy
level,
 So the atoms get
bigger.

H
Li
Na
K
Rb
As you go across a period the radius
gets smaller.
 Same energy level.
 More nuclear charge.
 Outermost electrons are closer.

Na
Mg
Al
Si
P
S Cl Ar
Table of
Atomic
Radii
Cations form by losing electrons.
 Cations are smaller that the atom they
come from.
 Metals form cations.
 Cations of representative elements have
noble gas configuration.

Anions form by gaining electrons.
 Anions are bigger that the atom they
come from.
 Nonmetals form anions.
 Anions of representative elements have
noble gas configuration.

Rb
K
Atomic Radius (nm)
Na
Li
Kr
Ar
H
Ne
10
Atomic Number
The amount of energy required to
completely remove an electron from a
gaseous atom.
 Removing one electron makes a +1 ion.
 The energy required is called the first
ionization energy.

The second ionization energy is the
energy required to remove the second
electron.
 Always greater than first IE.
 The third IE is the energy required to
remove a third electron.
 Greater than 1st of 2nd IE.

Symbol First
H
He
Li
Be
B
C
N
O
F
Ne
1312
2731
520
900
800
1086
1402
1314
1681
2080
Second
Third
5247
7297
1757
2430
2352
2857
3391
3375
3963
11810
14840
3569
4619
4577
5301
6045
6276
The greater the nuclear charge the
greater IE.
 Distance from nucleus increases IE
 Filled and half filled orbitals have lower
energy, so achieving them is easier,
lower IE.
 Shielding

 As
you go down a group first IE
decreases because
 The electron is further away.
 More shielding.
All the atoms in the same period have
the same energy level.
 Same shielding.
 Increasing nuclear charge
 So IE generally increases from left to right.
 Exceptions at full and 1/2 fill orbitals.

He
He has a greater IE
than H.
 same shielding
 greater nuclear
charge
First Ionization energy

H
Atomic number
He
l
l
l
First Ionization energy
l
Li has lower IE than H
more shielding
further away
outweighs greater nuclear charge
H
Li
Atomic number
He
l
l
First Ionization energy
l
H
Be has higher IE than Li
same shielding
greater nuclear charge
Be
Li
Atomic number
He
l
l
l
First Ionization energy
l
H
B has lower IE than Be
same shielding
greater nuclear charge
By removing an electron we make s
orbital half filled
Be
Li
B
Atomic number
First Ionization energy
He
H
Be
Li
C
B
Atomic number
He
First Ionization energy
N
H
C
Be
Li
B
Atomic number
He

First Ionization energy
N
H
C O
Be
Li
Breaks the pattern
because
removing an
electron gets to
1/2 filled p orbital
B
Atomic number
He
First Ionization energy
N F
H
C O
Be
Li
B
Atomic number
Ne
He
First Ionization energy
N F
H
C O
Be
Li
Ne has a lower IE
than He
 Both are full,
 Ne has more
shielding
 Greater distance

B
Atomic number
Ne
He
N F
l
First Ionization energy
l
H
C O
Be
Li
B
l
l
Na
Atomic number
Na has a lower
IE than Li
Both are s1
Na has more
shielding
Greater distance
Atomic number
First Ionization energy
Full Energy Levels are very low energy.
 Noble Gases have full orbitals.
 Atoms behave in ways to achieve noble
gas configuration.

Electron Affinity - the energy change
associated with the addition of an electron
 Affinity tends to increase across a period
 Affinity tends to decrease as you go down
in a period
Electrons farther from the nucleus
experience less nuclear attraction
Some irregularities due to repulsive
forces in the relatively small p orbitals
Table of Electron Affinities
The tendency for an atom to attract
electrons to itself when it is chemically
combined with another element.
 How fair it shares.
 Big electronegativity means it pulls the
electron toward it.
 Atoms with large negative electron
affinity have larger electronegativity.

The further down a group the farther the
electron is away and the more electrons
an atom has.
 More willing to share.
 Low electronegativity.

Metals are at the left end.
 They let their electrons go easily
 Low electronegativity
 At the right end are the nonmetals.
 They want more electrons.
 Try to take them away.
 High electronegativity.

Ionization energy, electronegativity
Electron affinity INCREASE
Atomic size increases, shielding constant
Ionic size increases
Another Way to Look at Ionization Energy
Yet Another Way to Look at Ionization Energ
Summary of Periodic Trends