General Chemistry

1. ATOMIC STRUCTURE
AND PERIODICITY
AP Chem: Unit 7

2. ELECTROMAGNETIC
RADIATION

3. ELECTROMAGNETIC RADIATION
One of the ways that energy travels through
space is by electromagnetic radiation.
 light from the sun
 X-rays
 microwaves

4. WAVE CHARACTERISTICS
 Wavelength (λ) – is the distance between
two consecutive peaks or troughs in a wave.
 Frequency (ν) – is the number of waves
(cycles) per second that pass a given point
in space.
 units – hertz or waves/sec (s-1)
 Speed (c) – all types of electromagnetic
radiation travel at the speed of light.
 2.9979 x 108 m/s
 c = λν

5. ELECTROMAGNETIC RADIATION

6. ELECTROMAGNETIC RADIATION

7. THE NATURE OF
MATTER

8. WAVE AND PARTICLE DUALITY
Planck found that matter could only absorb or
emit energy in whole number multiples of
the quantity hν.
 h is Planck’s constant = 6.626 x10-34 Js
 ΔE = hν
 Transfer of energy is not continuous but is
quantized and can occur only in discrete
amounts called quantum. Thus energy has
particle properties as well as wave
properties.

9. EINSTEIN

10. WAVE AND PARTICLE DUALITY
Einstein proposed that electromagnetic
radiation was also quantized and could
be viewed as a stream of “particles”
called photons.
Ephoton = hv = hc/λ

11. THE PHOTOELECTRIC EFFECT
The photoelectric effect refers to the
phenomenon in which electrons are emitted
from the surface of a metal when light
strikes it.
1. No electrons are emitted by a metal below
a specific threshold frequency (vo)
2. For light with frequency lower than the
threshold frequency, no electrons are
emitted regardless of intensity of the light.

12. THE PHOTOELECTRIC EFFECT
3. For light with frequency greater than the
threshold frequency, the number of
electrons emitted increases with the
intensity of the light.
4. For light with frequency greater than the
threshold frequency, the kinetic energy of
the emitted electrons increases directly
with frequency of the light.

13. THE PHOTOELECTRIC EFFECT
These observations can be explained by
assuming that electromagnetic radiation is
quantized (consists of photons), and that
the threshold frequency represents the
minimum energy required to remove the
electron from the metal’s surface.
 Minimum energy required to remove an
electron = Eo = hvo
 KEelectron = ½ mv2 = hv – hvo

14. PLANCK AND EINSTEIN
CONCLUSIONS
Energy is quantized. It can occur only
in discrete units called quanta.
Electromagnetic radiation, which was
previously thought to exhibit only wave
properties, seems to show certain
characteristics of particulate matter as
well. This phenomenon is sometimes
referred to as the dual nature of light.

15. WAVE PARTICLE DUALITY
The main significance of the equation E
= mc2 is that energy has mass.
m = E/c2

16. LOUIS DE BROGLIE (1892-1987)
Since light which previously was thought
to be purely wavelike, was found to
have certain characteristics of
particulate matter. But is the opposite
also true? Does matter have that is
normally assumed to be particulate
exhibit wave properties?

17. LOUIS DE BROGLIE (1892-1987)
de Broglie’s equation allows us to
calculate the wavelength for a particle:

18. DE BROGLIE’S PROOF

19. LOUIS DE BROGLIE (1892-1987)
Conclusion: Energy is really a form of
matter, and all matter shows the same
types of properties. All matter exhibits
both particulate and wave properties.

20. THE ATOMIC SPECTRUM OF
HYDROGEN

21. SPECTRUM
A continuous spectrum results when
white light passes through a prism and
all wavelengths (colors) are shown.
An emission spectrum produces only
a few lines of color that is limited to
discrete wavelengths produced by an
atom. This is called a line spectrum
and is specific to each atom.

22. HYDROGEN LINE SPECTRUM
The significance of the line spectrum is
that it indicates that only certain
energies are allowed for the electron in
the hydrogen atom. In other words the
energy of the electron in the hydrogen
atom is quantized

23. HYDROGEN LINE SPECTRUM

24. THE BOHR MODEL

25. NIELS BOHR
Bohr developed a
quantum model for
the hydrogen atom
that allowed for only
specific energy
levels around the
atom that
corresponded with
specific radii.

26. NIELS BOHR (1885-1962)
The most important equation to come
from Bohr’s model is the expression for
the energy levels available to the electron
in the hydrogen atom.

Z is the nuclear charge, n is the energy
level.

27. NIELS BOHR (1885-1962)
The most important equation to come from
Bohr’s model is the expression for the
energy levels available to the electron in
the hydrogen atom.

the negative sign calculates a lower energy
closer to the atom, not the radiation of
negative energy.

28. EXAMPLE
What is the change in energy if an
electron in level 6 (excited state)
returns to level 1 (ground state) in
a hydrogen atom?
 ni=6; nf=1; Z=1 (hydrogen
nucleus contains a single proton)

29. EXAMPLE
What is the change in energy if an
electron in level 6 (excited state)
returns to level 1 (ground state) in
a hydrogen atom?

30. EXAMPLE
ΔE=Ef – Ei = E1 – E6=-2.117 x 10-18J
The negative sign for the change in energy
indicates that the atom has lost energy and
is now more stable. This loss of energy
produces a photon.

31. EXAMPLE
What is the corresponding wavelength
for the energy produced from the
electron jump?
E = -2.117 x 10-18J

9.383x10-8 m

32. BOHR MODEL CONCLUSIONS
 The model correctly fits the quantized energy
levels of the hydrogen atom and postulates
only certain allowed circular orbits for the
electrons.
 As the electron becomes more tightly bound,
its energy becomes more negative relative to
the zero-energy reference state. As the
electron is brought closer to the nucleus,
energy is released from the system.

33. BOHR MODEL CONCLUSIONS

34. BOHR MODEL CONCLUSIONS
The energy levels calculated by Bohr
closely agreed with the values
obtained from the hydrogen emission
spectrum but does not apply well to
other atoms. The Bohr’s model is
fundamentally incorrect but is very
important historically because it paved
the way for our current theory of
atomic structure.

35. THE QUANTUM
MECHANICAL MODEL OF
THE ATOM

36. QUANTUM MECHANICS
Quantum Mechanics or Wave
Mechanics were developed by three
physicists: Heisenberg, de Broglie,
and Schrodinger.
Emphasis was given to the wave
properties of the electron.
The electron bound to the nucleus
behaves similar to a standing wave.

37. QUANTUM MECHANICS
 Like a standing wave,
electrons can travel in
patterns that allow for
a common node. In
other words, wave
patterns around the
nucleus must be in
whole number wave
patterns. But their
exact movement is
not known.

38. HEISENBERG UNCERTAINTY
PRINCIPLE
There is a fundamental limitation to
just how precisely we can know both
the position and momentum of a
particle at a given time. This
limitation is small for large particles
but substantial for electrons.

39. PROBABILITY DISTRIBUTION
A probability
distribution is used to
indicate the
probability of finding
an electron in a
specific position.
Electron density map
Radial probability
distribution

40. PROBABILITY DISTRIBUTION
For the hydrogen 1s orbital, the maximum
radial probability occurs at a distance of
5.29x10-2nm or .529Å from the nucleus.
This is the exact radius of the innermost
orbit calculated in the Bohr Model.
The definition most often used by chemists
to describe the size of the hydrogen 1s
orbital is the radius of the sphere that
encloses 90% of the total electron
probability

41. QUANTUM NUMBERS

42. QUANTUM NUMBERS
Each orbital is characterized by a series
of numbers called quantum numbers,
which describe various properties of
an orbital:
1. Principal quantum number (n)- has
integral values : 1,2,3,4. It describes
the size and energy of the orbital.
Energy Level

43. QUANTUM NUMBERS
2. Angular momentum quantum number
(l) – has integral values from 0 to n-1.
This is related to shape of the atomic
orbitals. Sublevel
o l =0 is s
o l =1 is p
o l =2 is d
o l =3 is f
o l =4 is g

44. QUANTUM NUMBERS
3. Magnetic quantum number (ml)
has values between l and – l ,
including 0. The value of ml is related
to the orientation of the orbital in
space. Axis designation

45. QUANTUM NUMBERS
4. Electron spin quantum number
(ms)- can only have one of two
values, +½, -½. Electrons can
spin in one of two opposite
directions.

46. QUANTUM NUMBERS
In a given atom no two electrons can
have the same set of four quantum
numbers (n, l, ml , ms). This is called
the Pauli exclusion principle; an orbital
can only hold two electrons, and they
must have opposite spins.

47. QUANTUM NUMBERS

49. ORBITAL SHAPES AND
ENERGIES

50. S ORBITALS
The s orbitals have
a characteristic
spherical shape
and contain areas
of high probability
separated by areas
of zero probability.
These areas are
called nodal
surfaces, or nodes.

51. S ORBITALS
The number of
nodes increases as
n increases. The
number of nodes =
n - 1.

52. P ORBITALS
 P orbitals each have two lobes separated by a
node at the nucleus. The p orbitals are labeled
according to the axis of the xyz coordinate
system along which the lobes lie.

53. P ORBITALS
Cross section of electron probability of a p orbital

54. D ORBITALS
 The five d orbitals first occur in energy level
3. They have two fundamental shapes.
Four of the orbitals (dxz, dyz, dxy, and dx2-y2)
have four lobes centered in the plane
indicated in the orbital label. dx2-y2 lie along
the x and y axes and dxy lie between the
axes. The fifth orbital dz2 has a unique
shape with two lobes along the z axis and a
belt centered in the xy plane.

55. D ORBITALS

56. F ORBITALS
The f orbitals first occur in level 4 and have
shapes more complex than those of the d
orbitals. These orbitals are not involved in
the bonding in any of the compounds that
we will consider.

57. F ORBITALS

58. ORBITAL ENERGIES
For the hydrogen atom, the energy of
a particular orbital is determined by its
value of n. Thus all orbitals with the
same value of n have the same energy
– they are said to be degenerate.

59. POLYELECTRONIC
ATOMS

60. POLYELECTRONIC ATOMS
Polyelectronic atoms are atoms with more
than one electron. To look at these atoms,
three energy contributions must be
considered:
Kinetic energy of the electrons as they move
around the nucleus.
The potential energy of attraction between
the nucleus and the electrons.
The potential energy of repulsion between
the two electrons.

61. POLYELECTRONIC ATOMS
Since electron pathways are unknown,
dealing with the repulsions between
electrons cannot be calculated exactly.
This is called the electron correlation
problem.

62. POLYELECTRONIC ATOMS
The electron correlation problem occurs
with all polyelectronic atoms. To deal
with this, we assume each electron is
moving in a field of charge that is the
net result of the nuclear attraction and
the average repulsions of all the other
electrons.
In other words,…..

63. POLYELECTRONIC ATOMS
A valence electron is attracted to the
highly charged nucleus and still
repelled by the other ‘inner’ electrons.
The net effect is that the electron is not
bound nearly as tightly to the nucleus
as it would be if it were alone.
 This is a screened or shielded affect.

64. POLYELECTRONIC ATOMS
Because of this shielded affect. orbitals
within a principal energy level do not
have the same energy (degenerate).
Sublevels vary in energy within a
principal quantum level.
s<p<d<f

65. POLYELECTRONIC ATOMS
Hydrogen vs. Polyelectronic

66. HISTORY OF THE PERIODIC
TABLE

67. EARLY GREEKS
Earth
Air
Fire
Water

68. DOBEREINER (1780-1849)
Johann Dobereiner was the first chemist
to recognize patterns and found
several groups of three elements that
have similar properties.
chlorine, bromine and iodine
called triads.

69. NEWLANDS
John Newlands suggested that
elements should be arranged in
octaves, based on the idea that certain
properties seemed to repeat for every
eighth element in a way similar to the
musical scale.

70. MEYER AND MENDELEEV
The present form of the periodic table
was conceived independently by two
chemists: Meyer and Mendeleev.
Usually Mendeleev is given most of
the credit, because it was he who
emphasized how useful the table could
be in predicting the existence and
properties of still unknown elements.

71. MEYER AND MENDELEEV
In 1872 when Mendeleev first
published his table, the elements
gallium, scandium, and germanium
were unknown. Mendeleev correctly
predicted the existence and properties
of these elements from the gaps in his
periodic table. Mendeleev also
corrected the atomic masses of
several elements.

72. MENDELEEV’S TABLE

73. THE AUFBAU PRINCIPLE
AND THE PERIODIC
TABLE

74. THREE RULES FOR ORBITAL
CONFIGURATION
Aufbau principle – As protons are added,
so are electrons, and fill in orbitals in order
of energy levels.
Pauli Exclusion – Two electrons with
opposite spins can occupy an orbital.
Hund’s rule – The lowest energy
configuration for an atom is the one with
one unpaired electrons in each degenerate
orbital. (Electrons don’t like roommates)

75. VALENCE ELECTRONS
Valence electrons are the electrons in
the outermost principal quantum level
of an atom. These are the most
important electrons because they are
involved in bonding.
The inner electrons are known as core
electrons.

76. VALENCE ELECTRONS
The elements in the same group have
the same valence electron
configuration. Elements with the same
valence electron configuration show
similar chemical behavior.

77. TRANSITION METALS
Transition metals have electron
configurations that fill in the order of 4s
before 3d. Copper and Chromium have a
configuration that is observed different
than what is expected.
Expected: Cr: 1s22s22p63s23p64s23d4
 Observed: 1s22s22p63s23p64s13d5
Expected: Cu:1s22s22p63s23p64s23d9
 Observed: 1s22s22p63s23p64s13d10

78. TRANSITION METALS

79. ADDITIONAL ORBITAL RULES
The (n+1)s orbital always fills before the nd
orbitals. The s orbitals fill prior to the d orbitals
due to the vicinity of the nucleus.
After lanthanum, which has the configuration of
[Xe] 6s25d1, a group of 14 elements called the
lanthanide series, or the lanthanides occurs.
This seris of elements corresponds to the filling
of the seven 4f orbitals.

80. ADDITIONAL ORBITAL RULES

81. ADDITIONAL ORBITAL RULES
After actinium, a group of 14 elements
called the actinide series or actinides
occurs.
The groups 1A, 2A, 3A…, the group
numbers indicate the total number of
valence electrons for the atoms in these
groups.

82. ADDITIONAL ORBITAL RULES
After actinium, a group of 14 elements
called the actinide series or actinides
occurs.
The groups 1A, 2A, 3A…, the group
numbers indicate the total number of
valence electrons for the atoms in these
groups.

83. ADDITIONAL ORBITAL RULES

84. PERIODIC TRENDS IN
ATOMIC PROPERTIES

85. IONIZATION ENERGY
Ionization energy is the energy required
to remove an electron from a gaseous
atom or ion when the atom or ion is
assumed to be in its ground state:
 X(g) X+
(g) + e-

86. IONIZATION ENERGY
It is always the highest-energy
electron (the one bound least tightly)
that is removed first. The first
ionization energy (I1) is the energy
required to remove that first electron.
The second ionization energy (I2) is
considerably larger.

87. IONIZATION ENERGY
The first electron is removed from a
neutral atom, the second from a +1 cation.
The increase in positive charge binds the
electrons more firmly and the ionization
energy increases. The trend continues for
consecutive electrons removed.
Core electrons are always held tighter
than valence.

88. IONIZATION ENERGY
First ionization energy increases from left
to right across a period.
First ionization energy decreases in going
down a group.

89. IONIZATION ENERGY

90. ELECTRON AFFINITY
Electron Affinity is the change in
energy change associated with
the addition of an electron to a
gaseous atom:
 X(g) + e-  X-
(g)

91. ELECTRON AFFINITY
If the addition of the electron is
exothermic the corresponding value for
electron affinity will carry a negative sign.
The more negative the energy, the
greater the quantity of energy
released.

92. ELECTRON AFFINITY
Electron affinities generally become
more negative from left to right across a
period and becomes more positive down
a group.
 As with Ionization energy. Some
exceptions occur due to repulsions
and electron configuration.

93. ELECTRON AFFINITY

94. ELECTRON AFFINITY

95. ATOMIC RADIUS
Atomic radii are measured by the distances
between atoms in chemical compounds.
Covalent atomic radii are assumed to be
half the distance between atoms in covalent
bonds.
For metallic atoms, the metallic radii are
obtained from half the distance between
metal atoms in a solid metal crystal

96. ATOMIC RADIUS
Atomic radii decrease in going from left to
right across a period because of increasing
nuclear charge and decreasing shielding.
Atomic radius increases down a group,
because of the increases in the orbitals
sizes associated with principal quantum
numbers.

97. ATOMIC RADIUS

98. ATOMIC RADIUS

99. THE PROPERTIES OF A
GROUP: THE ALKALI
METALS

100. INFORMATION AND THE PERIODIC
TABLE
It is the number and type of valence
electrons that primarily determine an
atom’s chemistry
The organization of the period table
allows the prediction of electron
configuration without memorization.

101. INFORMATION AND THE PERIODIC
TABLE
Groups on the periodic table have
specialized names: Alkali metals, Alkaline
earth metals, Halogens, …etc.
The most basic division of elements in the
periodic table is into metals and non-
metals. This division affects chemical
properties.
 Metals tend to give up electrons and have low
ionization energies. The opposite is true for
non-metals.

102. INFORMATION AND THE PERIODIC
TABLE

103. INFORMATION AND THE PERIODIC
TABLE
Metalloids are elements along the division
line and exhibit both metallic and
nonmetallic properties under certain
circumstances. These elements are
sometimes called semimetals.

104. THE ALKALI METALS
Lithium, sodium, potassium, rubidium,
cesium, and francium are the most
chemically reactive of the metals.
Hydrogen is found in group 1 but behaves
as a nonmetal because its very small and
the electron is bound tightly to the
nucleus.

105. THE ALKALI METALS
Going down the group the first ionization
energy decreases and the atomic radius
increases. The overall density increases
due to the increase of atomic mass
relative to atomic size (therefore more
mass per unit volume).

106. THE ALKALI METALS
There is a smooth decrease in melting
point and boiling points in Group 1 that is
not typical for other groups.
The most important chemical property of
Group 1 is its ability to lose its valence
electrons. Group 1 are very reactive.

107. THE ALKALI METALS
 Hydration energy of an ion represents the change
in energy that occurs when water molecules attach
to the metal cation.
 The hydration energy is greatest with Li+
because it has the most charge density (charge
per unit volume). This means that polar water
molecules are more strongly attracted to the
small Li+ ions
 The order of reducing abilities in an aqueous
reaction is Li > K > Na

108. THE ALKALI METALS

109. THE ALKALI METALS

110. THE END