Inorganic Chemistry
THE DISCOVERY OF ATOMIC PARTICLES
The 3 fundamental atomic particles
are protons, neutrons and electrons.
nucleons (particles
residing in the nucleus)
-
+
proton
(positive)
neutron
(neutral)
Democritus (an ancient Greek philosopher) taught that if any
object was repeatedly cut into smaller and smaller pieces,
eventually a smallest particle would be obtained that could not
be further divided. He called this smallest particle of matter an
‘atom’. (Gr. ‘a’ = not, ‘tom’ = cuttable).
electron
(negative)
Democritus (BC): 'The atom is the
smallest indivisible particle.'
John Dalton (1807), a British schoolteacher, pictured atoms as solid billiard ball-like spheres. He
measured the masses of elements that reacted to form various compounds and proposed his
‘atomic hypothesis’:
1. an element is composed of only 1 kind of
atom, e.g., the element carbon contains only
carbon atoms.
2. Atoms of different elements have unique
(different) masses, e.g., a carbon atom has a
mass of 12 atomic mass units (amu), a
hydrogen atom has a mass of 1 amu.
John Dalton (1807): atoms are spherical
1. elements contain only 1 kind of atom
2. atoms of different elements have
different masses
3. compounds have fixed ratios of
atoms, e.g., 2:1ratio of H:O in H 2O
4. mass is conserved in reactions
O
4. Atoms are exchanged (not created or
destroyed) in chemical reactions, e.g.,
H2O +
Na 
H
H
3. Chemical compounds are formed from
specific ratios of different elements, e.g., H2O
always forms in the ratio of 2 hydrogen
atoms per oxygen atom.
Elemental C contains only C, each
C atom with a mass of 12 amu.
Water (H2O)
H = 1 amu
O = 16 amu
NaOH + H2
We know this as the law of conservation of mass. Thus, chemists always balance equations.

Balance the preceding equation.
The Atom
1
Inorganic Chemistry
Electrons: In 1897, the first subatomic particle (the electron) was discovered by a British
physicist, J. J. Thomson using cathode ray tubes. Two electrodes are sealed in a glass tube
containing gas at a low pressure. When a high voltage is applied across the electrodes, current
flows as a visible stream of electrons are emitted from the negative electrode (cathode) to the
positive electrode (anode). Thomson found that the rays were the same regardless of the metal
used for the cathode and he correctly concluded that the particles were part of the makeup of all
atoms. Thomson found that cathode rays were deflected by nearby electric and magnetic fields.
Cathode Ray Tube:
containing inert gas at low pressure (JJ Thompson, 1897)
-
+
S
N
+
-
+
anode
cathode
-
The cathode ray (beam of electrons) is deflected by both electric and magnetic fields.
Thus electrons have electric charge. In fact,
the charge on an electron is the smallest unit
of electric charge that can exist (1.602  10-19
Coulombs).
Although electrons have a
negative charge, atoms have an overall
charge of zero. Therefore scientists around
1900 knew that each atom must contain
enough positive charge to cancel out the
negative charge.
Thomson proposed a
‘plum pudding’ model of the atom, in which
the positive charge was distributed evenly
throughout the atom and the negative
charges were pictured as being imbedded in
the atom like plums in a pudding.
The Atom
Thomson's Plum Pudding Model (1900)
electrons imbedded
in the atom
+
-
+
+
+
-
-
-
positive charges
evenly distributed
+
+
-
2
Inorganic Chemistry
Atomic Nucleus: By 1909, Ernest Rutherford had determined that alpha () particles (helium
nuclei, He+2) are positively charged particles and are emitted by some radioactive atoms – atoms
that spontaneously disintegrate. Rutherford bombarded a thin gold foil with  particles from a
radioactive source. A fluorescent, ZnS screen was placed around the foil to observe the scattering
of the  particles by the gold atoms. Scintillations (flashes) on the screen caused by the impact of
individual  particles were counted to determine the relative number of  particles deflected at
various angles of deflection.
Rutherford's Discovery of the Nucleus
As expected, most 
particles passed through the
foil with little or no
deflection, however, to his
amazement, a few were
deflected at large angles
and a few  particles
bounced straight back at the
source.
ZnS scintillation screen
thin Au film
-particle beam
Pb shield
radioactive
-emitter (Ra)
Rutherford proposed that
the positive charge in atoms
is not evenly distributed but
exists as dense, point-like
centers surrounded by a
large volume of empty
space.
scintillations
(sparking)
Pb shield
Rutherford Model of the Atom
atom is mostly
empty space
dense positive
nucleus
tiny electrons
-
-
-
 particles
(helium nuclei)
Rutherford named these centers of positive
charge – ‘atomic nuclei’. He was able to
calculate the magnitude of the positive
charges and estimate the diameter of the
nucleus at ca. 1/100,000 of an atom.
Since the atom is mostly empty space, the
atomic nucleus, containing virtually all the
mass, is extremely dense. In fact, an
atomic nucleus the size of a grain of sand
would weigh ca. 50  106 tons!!
Alpha () particles are Helium nuclei, He+2.
M ost alpha particles pass through the gold foil undeflected
because the atom is m ostly em pty space, how ever, a few
alpha particles are deflected at sharp angles w hen a nucleus
is approached by the He nuclei (like charges repel).
The Atom
3
Inorganic Chemistry
A Lithium atom
ee-
n n +
+ p
np n
+
p
e-
3
Li
We now know that every nucleus contains
an integral (whole) number of protons
equal to the number of electrons in the
atom (atoms are electrically neutral). The
number of protons in an atom (called the
atomic number, symbol (Z) determines an
atoms identity, e.g., all atoms with 3
protons are lithium atoms.
Atomic number, Z, (number of protons)
determines an atom's identity.
All atoms with 3 protons are lithium atoms.
Neutrons: The 3rd kind of fundamental particle
was discovered by James Chadwick in 1932. He
bombarded beryllium atoms with high-energy 
particles and dislodged uncharged particles
(neutrons) from the nucleus. It was soon after
understood that the nuclei of all atoms (except the
common form of hydrogen) contain 1 or more
neutrons. Neutrons are almost identical in size and
mass with protons. Both neutrons and protons are
collectively termed ‘nucleons’ since they both
reside in the nucleus of the atom.
Identifying the Elements:
H.G.J. Moseley
directed high-energy electrons at samples of pure
elements. Electrons decelerate rapidly on impact
and in so doing emit x-rays. The x-rays emitted are
recorded photographically as a series of lines –
their patterns varying with the atomic mass of the
element. On the basis of mathematical analysis of
these x-ray data, it was concluded that each
element (H, He, Li, Be, etc.) differs from the
preceding element by having one more positive
charge in its nucleus. For the first time it was
possible to arrange all know elements in order of
increasing nuclear charge.
The Atom
neutrons
alpha particles
Be metal
x-rays
high energy
electrons
any metal
4
Inorganic Chemistry
Properties of Subatomic Particles
Particle
Symbol
Charge*
Mass (g)
Mass (amu)**
electron
e
-1
9.109  10-28
0.00055
proton
p
+1
1.673  10-24
1.0078
neutron
n
0
1.675  10-24
1.0090
* charges are given as multiples of the charge on an electron (1.602  1019 Coulombs)
** amu (atomic mass unit) = 1/12 of the mass of carbon 12 (12C), the most common form of C.
The Mass Spectrometer and Isotopes: The mass spectrometer (or mass spec) is one of the
most powerful analytical instruments available. It permits chemists to identify and quantify
(measure the concentration of) all known elements and almost all known compounds.
A portion of the element to be analyzed is injected into a heated sample chamber. Its
vapors are drawn into the evacuated instrument and bombarded with high-energy electrons from a
cathode ray. The colliding electrons dislodge electrons from the sample atoms producing positive
ions of the element. These electrically charged atoms (positive ions) are accelerated through the
instrument by a strong electric field applied between 2 metal grids. The ions’ speeds vary with
their masses, lighter ions reaching higher speeds. The path of the ions is bent as they travel
between the poles of a variable electromagnet. As the magnetic field is varied, each type of
positive ion is, in turn, directed to a detector which produces electric signals integrated into a
‘mass spectrum’ – a plot of concentration (signal intensity) versus atomic (positive ion) mass.
Mass Spectrometer
N
varying magnetic field
deflects cations as per
their mass/charge ratio
slit
detector
cations accelerated
through spectrometer
by electric field
+
S
Mass Spectrum
electron gun dislodges electrons
from sample creating cations
mass
fraction
sample injected into heated chamber
mass
The Atom
5
Inorganic Chemistry
Using mass specs, the atomic mass of all 112 known elements have been measured with great
accuracy. The use of early mass specs led to the discovery of isotopes. Researchers found that
not all atoms of a single element have the same mass. For example, all atoms of boron have 5
protons, however, in a sample of pure boron, 20.0% of the atoms have 5 neutrons while the
remaining 80.0% have 6 neutrons. Atoms of the same atomic number (same number of protons,
hence same element) but different number of neutrons (hence different mass) are called isotopes.
Naturally Occurring Isotopic Abundance of Some Elements
Element
Boron
Carbon
Silicon
Nuclide
Symbol of
Isotope
Number of
Neutrons
% Natural
Abundance
Atomic
Mass
(amu)
10
5
B
5
20.0
10.01294
11
5
B
6
80.0
11.00931
12
6
C
98.89
12.0000
13
6
C
1.11
13.0034
28
14
Si
92.21
27.9769
29
14
Si
4.70
28.9765
30
14
Si
3.09
29.9738
Protium
1
1
H
99.98
1.0078
Deuterium
2
1
H
0.02
2.0141
Tritium
3
1
H
trace
3.0
Notation of Nuclide Symbols:
A
Z
E
Weighted
Average
Mass (amu)
10.81
E = symbol of the Element
Z = atomic number (number of p)
A = mass number (number of p + n)*
* On the periodic table, the symbol A refers to the atomic mass, i.e., the weight average atomic
weight of the element in its naturally occurring form – as a mixture of isotopes. Atomic mass is not
an integral (whole) number whereas mass number is a count of the number of protons + neutrons
and is always an integral number
Problem: Complete the empty cells in the table of Isotopic Abundance.
Problem: Write the nuclide symbol and state the number of protons, neutrons, and electrons:
a. an atom of nitrogen 14
b. an atom of iron 56
c. an atom of uranium 236
Problem: The weight average mass of gallium is 69.72 amu. The masses of the naturally
occurring isotopes are 68.9257 for 69Ga and 70.9249 for 71Ga. Calculate the % abundance of each
isotope. (Answer: 69 Ga = 60.0%, 71Ga = 40.0%)
The Atom
6
Inorganic Chemistry
Atomic Weight Scale and Atomic Weights: Even before the masses of various kinds of atoms
could be measured (as with the mass spec) scientists determined a relative scale of atomic
masses for many of the elements. For example, experiments showed that carbon and hydrogen
have relative atomic masses (atomic weights) of 12 to 1, respectively.
The atomic weight scale approved in 1962 by the International Union of Pure and Applied
Chemists (IUPAC) is based on the carbon-12 isotope.
One amu is exactly 1/12 of the mass of a C-12 atom.
One 12C atom weighs 12 amu.
This is approximately the mass of one atom of protium (1H), the lightest isotope of the element with
lowest mass.
Problem: Calculate the mass of 1 amu in grams. Recall Avogadro’s Number, N = 6.022  1023
atoms per mole. (Answer = 1.66110-24g)
Quantum Mechanics: The Rutherford model of the atom, while basically correct, did not answer
important questions such as the following.
 Why do different elements have different physical and chemical properties?
 Why and how does chemical bonding occur?
 Why do atoms of different elements give off or absorb light of characteristic colors?
Early scientists found that classical mechanics (Newton’s laws) which successfully describe the
motion of visible objects like balls and planets, fail when applied to electrons in atoms. New laws,
which came to be known as quantum mechanics were developed in the early 1900’s.
Electromagnetic Radiation: All of the previous unanswered questions (above) can be explained
with an understanding of electron arrangement (configuration) within atoms. What we currently
know about electron configuration has largely been determined from the analysis of
electromagnetic radiation emitted or absorbed by substances (spectroscopy)
Electromagnetic radiations are forms of radiant energy, some natural and some synthetic, that
possess no mass or weight and are electrically neutral. They also share 4 other common
characteristics:
1) all pass through a vacuum in wavelike motion;
2) all travel at the speed of light (3.00  108 m/s), denoted ‘c’
3) all give off electric and magnetic fields
4) all have different energies, wavelengths, and frequencies
Problem: Calculate the speed of light in miles per hour. 1 mile = 1.609 km. (Ans. = 6.71108mi/h)
Problem: How many minutes will it take light from the sun to reach earth assuming an average
distance of 93  106 miles?(Ans. = 8.31min.)
The Atom
7
Inorganic Chemistry
The spectrum of electromagnetic radiation, in order of increasing energy (decreasing wavelength),
includes radio and TV waves, microwaves, infrared radiation, visible light, ultraviolet radiation, Xrays, and gamma rays. No clear-cut separation exists between the bands so overlap of
wavelengths, as shown below, is reported in various literature sources.
The Electromagnetic Spectrum
ENERGY
Wavelength (m)
-14
-12
10
-10
10
10
Gamma rays
-8
10-6
10
10-4
10-2
Visible light
1
102
Microwaves
104
AM radio
Ultraviolet
X-rays
1022
Infrared (heat)
1020
1018
1016
1014
1012
TV and FM radio
1010
108
106
104
Frequency (Hz)
As far as we know, there is neither and upper nor a lower limit to the wavelength of EMR.
Since all types of electromagnetic radiation (EMR) travel as waves, they can be described in terms
of their frequency (, Gr. ‘nu’) and wavelength (, Gr. ‘lambda’).
Wavelength is the distance between any 2 identical points of a wave, for instance, 2 adjacent
crests. The frequency is the number of wave crests passing a given point per unit time, usually
expressed in cycles/second (cps or s-1 or Hertz, Hz).

e.g., radio
wave
a)
amplitude (A) = maximum
displacement from the rest position
A
wavelength () = distance between
any 2 identical points on a wave,
e.g., crest to crest, trough to trough
e.g. visible light
frequency (v) = cycles per second
b)
a) has longer  but lower v than b).
a) and b) travel at the same velocity
(3.00 X 10 8 m/s).

For a wave traveling at some speed, the wavelength and frequency are related to each other by…
   = speed of propagation
or
   = c = 3.00  108 m/s
Thus  and  are inversely proportional to each other. A shorter the  equals a higher . For
water waves it is the surface of the water that oscillates. For a vibrating guitar string it is the string
that moves repetitively. EMR consists of regular, repetitive variation in electrical and magnetic
fields.
The Atom
8
Inorganic Chemistry
The EMR most familiar to us is visible light.
400 m (violet) to 800 m (red).
It has wavelengths varying from about
Problem: Calculate the frequency of:
a. violet light (Ans. = 7.51014s-1)
b. red light (Ans. = 3.81014Hz)
Quanta and Photons: In addition to behaving as waves, EMR can be described as particles
called photons. Max Plank, 1900, discovered that each photon has a fixed amount (a quantum)
of energy. The amount of energy possessed by a photon depends on its frequency and
wavelength.
The energy of a photon of light (e) is given by Plank’s Equation.
e=h
or
c
eh

h = Plank’s constant = 6.62  10-34Js
 = frequency of radiation in Hz
 = wavelength of radiation in m
Problem:
a. Calculate the energy of 1 photon of violet light ( = 7.31  1014 s-1) [Ans. = 4.8410-19 J]
b. Calculate the energy of 1 photon of x-rays with  = 2.5 m. [Ans. = 7.9410-17 J]
The Photoelectric Effect: Evidence for the particle
nature of light came from the photoelectric effect.
When light of sufficiently high energy strikes the
negative electrode (cathode) in an evacuated tube,
electrons are knocked off the electrode surface and
travel to the positive electrode (anode) creating an
electric current in the circuit. However, the following
behaviors were noted …
a. no electrons were ejected unless the
incident radiation has a frequency above a
certain value characteristic of the metal
cathode, no matter how long or how brightly
the light shines.
Photoelectric Effect
anode
electrons
emitted from
cathode
+
A
+
EMR (h v)
DC
voltage
-
cathode
b. the electric current (number of electrons
emitted per second) increases with
increasing brightness (intensity) of the light.
Classical physics said that even low energy light should cause a current to flow if the metal
is irradiated long enough. Electrons should accumulate energy and be released when they have
enough energy to escape from the metal atoms. This is not observed. In addition, old theory
suggests that if light is more energetic, then the current should increase even though the light
intensity remains the same. This also is not observed.
The Atom
9
Inorganic Chemistry
The answer to the puzzle was provided by Albert Einstein. In 1905 he extended Plank’s idea that
light behaves as though it were composed of photons, each with a particular amount (quantum) of
energy. According to Einstein, each photon can transfer its energy to a single electron in a
collision. If the energy of the photon is equal to or greater than the amount needed to liberate the
electron, then the electron can escape the metal surface. Increased intensity means that the
number of photons striking a given area per second is increased.
Atomic Spectra and the Bohr Atom:
Incandescent (red hot or white hot) solids, liquids, and high-pressure gases emit continuous
spectra. For example, a white hot (nearly 1000 C) tungsten light bulb filament emits a continuous
band of visible radiation (white).
However, when an electric current is passed through a gas in a vacuum tube at very low pressure,
the light that the gas emits is dispersed by a prism into distinct lines. Such emission spectra are
described as bright line spectra. The lines can be recorded photographically and the wavelength
of light that produced each line can be calculated.
Similarly, we can shine a beam of white light (containing a continuous spectrum) through a gas
and analyze the beam that emerges. We find that only certain wavelengths have been absorbed.
The wavelengths that are absorbed in this absorption spectrum are the same as those given off
in the emission experiment.
Each element displays its own characteristic set of lines in its emission or absorption spectrum.
These spectra can serve as 'fingerprints' to allow us to identify different elements in a sample,
even in trace amounts.
Emission spectra of various elements were intensely studied by scientists. J.J. Rydberg, a British
school teacher discovered that the wavelengths of the hydrogen spectrum can be related by a
mathematical equation:
The Rydberg equation:
410
434
410
486
 1
1
 R   2  2  where R = 1.097  107 m-1 (the Rydberg constant)

 n1 n2 
1
656 nm
The bright line emission spectrum of
hydrogen
shows
emission
at
 = 410, 434, 486 and 656 nm. Each
element has a unique spectrum.
In 1913, Neils Bohr, a Danish physicist, provided an explanation for Rydberg's observations. He
wrote equations that described the electron of a hydrogen atom as revolving around the nucleus of
the atom in circular orbits (planetary model of the atom). He included assumptions that the
electronic energy is quantized; that is, only certain values of electron energy are possible. This led
him to the suggestion that electrons can only be in certain discrete orbits, and that they absorb or
emit energy in discrete amounts as they move from one orbit to another. Each orbit thus
corresponds to a definite energy level for the electron. When an electron is promoted from a lower
energy level to a higher one, it absorbs a definite (or quantized) amount of energy. When the
electron falls back to the original energy level, it emits exactly the same amount of energy it
absorbed in moving from the lower to the higher energy level.
The Atom
10
Inorganic Chemistry
Bohr's equation for the energy of each orbit was:
E
2   m e4
n2h2
h = Plank's constant
m = mass of an electron
n = (1, 2, 3 … +n), i.e., the various allowed orbits
where
The larger the value of n, the farther from the nucleus is the orbit being described. For orbits
farther from the nucleus, the electronic potential energy is higher (less negative - the electron is in
a higher energy level or less stable state). As n approaches infinity, the electron is completely
removed from the nucleus.
With this equation, Bohr was able to predict the wavelengths observed in the hydrogen emission
spectrum. Although the Bohr theory explained the spectra of hydrogen and other species
containing only one electron (He+, Li+2) it could not calculate the wavelengths observed in spectra
of more complex species. Bohr's approach was doomed to failure because it modified classical
mechanics. It was a contrived solution. There was a need to literally invent a new physics,
quantum mechanics, to deal with subatomic particles.
However, Bohr's theory did support the ideas that only certain energy levels are possible and that
energy levels could be described by quantum numbers.
Wave Particle Duality of Matter: Once it was learned that EMR can exhibit both wave properties
and particle properties, French scientist, Louie de Broglie (1925) suggested that all particles have
wavelength properties. de Broglie predicted that a particle with a mass m and velocity v should
have a wavelength given by …
where h = Plank’s constant (6.62  10-34 Js)

h
mv
the product mv = linear momentum
Two years later, C. Davisson and L.H.
Germer at Bell Telephone laboratory
demonstrated the wavelike character
of electrons. They directed a beam of
electrons at a crystal of nickel. The
regular array of Ni atoms in the crystal
with centers separated by 250 m acts
as a grid that diffracts waves. A
diffraction pattern was observed.
Electron diffraction is the principle that
an instrument called an electron
microscope uses to determine the
structure of molecules and solid
surfaces.
The Atom
diffraction
pattern
beam of
electrons
Ni crystal
11
Inorganic Chemistry
Diffraction: is the bending of waves around the corner or edge of a solid object.
 Water waves are not interrupted by swimmers. The wave front reforms after passing
a swimmer (or other small objects). Small objects create no permanent break
(shadow) in a wave front.
 Sound bends (diffracts) around corners. Even in the absence of reflection, we can
hear sounds around the corner of large buildings. Sound has long wavelengths.
 Diffraction of light, although less obvious, is also common. Light bends (slightly)
around corners and through small openings. This produces light and dark fringes
(interference patterns) as wave fronts rejoin. As a crest meets another crest a bright
fringe is formed due to constructive interference. As a crest meets a trough, a dark
fringe appears due to destructive interference.
thick ne s s
Diffraction
(bending of waves)

shadow
When  is much larger than the thickness
of an object, the wavefronts bend around
a small object and reform.
No 'shadow' is cast.
When  is smaller than the thickness
of an object, the wavefronts do not reform.
A 'shadow' is cast.
Points where crests meet crests (or troughs meet troughs)
are points of constructive interference producing bright regions.
Points where crests meet troughs are points of destructive
interference producing dark regions.
 is larger than the opening so wavefronts bend and rejoin.
Problem:
a. Calculate the  of an electron traveling at 1/100 the speed of light. The mass of an electron is
9.11  10-28 g. [Ans. = 200pm-about the same size as many atoms-interference patterns are seen]
b. Calculate the  of a baseball of mass 5.25 oz. traveling at 92.5 mi/h.
16 oz = 1 lb, 1 kg = 2.205 lb. [Ans. = 1.0810-34m]
Wavelength of a large object is imperceptibly small in the macroscopic world.
The Atom
12
Inorganic Chemistry
Heizenburg Uncertainty Principle & Quantum Mechanics:
Through the work of de Broglie and others, we know that electron movement in atoms can
be better understood as wave motion rather than small particles traveling in circular orbits around
the nucleus (Bohr’s model). Classical laws of mechanics (Newton’s Laws) do not hold. A different
kind of mechanics, called quantum mechanics, has been developed to describe the wave
properties of small particles.
In classical mechanics, a particle has a definite position in space and has a definite
trajectory (path) when moving. This is not true for a wave. Think of the wave of a guitar string.
The wave is spread out all along the string, not localized at a precise point.
The Heizenburg Uncertainty Principle (1927, by Werner Heizenburg)
states that the location and momentum (trajectory & velocity) of an
electron cannot be known simultaneously.
If we know a particle is here at one instant, we can say nothing about where it will be an instant later.
A standing wave is a wave that does not travel and therefore has at least one point with
zero amplitude, called a node. As an example, consider the various ways a guitar string
can vibrate when plucked.
a) 1 half-wavelength
b) 2 half-wavelengths
(1 full wave)
nodes
c) 3 half-wavelengths
d) 1 1/4 wavelengths
(not possible)
Only integral (whole) numbers of half-wavelength vibrations are possible because the
ends remain fixed. Similarly, in the space around a nucleus, only certain wave forms
can exist. Each allowed waveform corresponds to an energy state, i.e., an orbital.
The Atom
13
Inorganic Chemistry
Some of the basic ideas of quantum mechanics include the following…

Atoms and molecules can exist only in certain energy states (at fixed energy levels). When an
atom or molecule changes its energy state, it must emit or absorb just enough energy to bring it
to the new energy level.

Atoms or molecules emit or absorb radiation (light) as they change their energies. The
frequency of the light emitted or absorbed is related to the energy change by the equation:
E = h
Plank’s Equation:
or
E  h 
c

Energy is gained (or lost) by an atom when its electrons move to higher (or lower) energy states
(orbitals).
Because of the Heizenburg uncertainty, scientists use a statistical approach to describing electron
position and motion in an atom. Erwin Schrodinger (1926) modified existing mathematical
equations of a 3-dimensional standing wave (called wave functions) to apply them to electron
motion.
E  V 
h 2   2  2  2



8m  x 2 y 2 z 2



Schrodinger’s wave equations were used calculate the probability of finding a particle at a
particular location, i.e., they define a 3-dimensional region of space in which an electron will reside
90 – 95% of the time. These regions of space are what we commonly call orbitals. Each solution
of the wave equations generates a set of 3 values (called quantum numbers) which together
describe each energy level (orbital) of an atom.
In 1928, Paul Dirac reformulated quantum equations to take into account relativity and this gave
rise to a 4th quantum number for each solution.
The Atom
14
Inorganic Chemistry
Quantum Numbers:
The Schrodinger and Dirac equations can only be exactly solved for 1 electron systems (H,
He+, Li+2). Simplifying assumptions were necessary to solve these equations for more complex
atoms and molecules. The most common and useful approximation of the wave equations is the
orbital approximation. This approximation is well supported by experimental evidence from
spectroscopy and chemical bonding behavior of the elements. The approximated wave equation
solutions yield 4 quantum numbers:
1. The principal quantum number, n, describes the principal (main) energy level an electron
occupies, basically its distance from the nucleus. It can only be a positive integer. These are
called shells.
n
1
2
3
4
shell
K
L
M
N
2. The 2nd quantum number, l, (also called azimuthal or angular momentum quantum number)
describes the divisions found in each of the main shells, i.e., ‘subshells’. The 2nd quantum
number, l, may take positive integral values from 0 up to and including (n-1). Each value of l
corresponds to a different type of orbital as shown in the following table.
l
0
1
2
3
(n-1)
subshell
s
p
d
f
orbital type
The s, p, d, and f designations arise from the characteristics of spectral emission lines
produced by electrons occupying the orbitals, s (sharp), p (principal), d (diffuse), and
f (fundamental).
In the 1st shell (n = 1), l = 0. There is only 1 possible subshell – the 1s subshell.
In the 2nd shell (n = 2), l = 0, 1. There are 2 possible subshells – the 2s and 2p subshells
In the 3rd shell (n = 3), l = 0, 1, 2. There are 3 possible subshells – the 3s, 3p and 3d subshells
In the 4th shell (n = 4), l = 0, 1, 2, 3. There are 4 possible subshells – 4s, 4p, 4d & 4f subshells
3. The 3rd quantum number, called the magnetic quantum number, ml, gives the spatial
orientation of an atomic orbital. ml is an integral number from – l through zero up to and
including + l, ml = (-l) … 0 … (+l)
For l = 0, there is only 1 value for ml (ml = 0). This indicates that there is only one orientation
for an s orbital, which is spherical around the nucleus. There is only one s orbital in each shell.
For l = 1, there are 3 values for ml (ml = -1, 0, +1). These correspond to three distinct regions
of space, the px, py and pz orbitals. There are three p orbitals in the 2nd and all higher shells.
For l = 2, there are 5 values for ml (ml = -2, -1, 0, +1, +2). These correspond to five distinct
d orbitals. There are five d orbitals in the 3rd and all higher shells.
For l = 3, there are 7 values for ml (ml = -3, -2, -1, 0, +1, +2, +3). These correspond to seven
distinct f orbitals. There are seven f orbitals in the fourth and all higher shells.
The Atom
15
Inorganic Chemistry
4.
The 4th quantum number, called the spin quantum number, ms, refers to the direction of spin
of an electron. For every set of n, l, and ml quantum values, ms can be either +½ or -½.
(clockwise spin or counter clockwise spin)
The values of n, l, ml describe a particular atomic orbital. Each atomic orbital can accommodate
no more than 2 electrons, one with ms = +½ and another with ms = -½.
Summary of Quantum Numbers
n values
(shell)
l values
[0 - (n-1)]
subshell
(orbital type)
ml values
[(-l) … 0 … (+l)]
# of
orbitals
max. # e-'s
in subshell
max. # e-'s
in shell
1 (K)
0
1s
0
1
2
2
2 (L)
0
2s
0
1
2
8
1
2p
-1, 0, +1
3
6
0
3s
0
1
2
1
3p
-1, 0, +1
3
6
2
3d
-2, -1, 0, +1, +2
5
10
0
4s
0
1
2
1
4p
-1, 0, +1
3
6
2
4d
-2, -1, 0, +1, +2
5
10
3
4f
-3,-2,-1, 0,+1,+2,+3
7
14
0
5s
0
1
2
1
5p
-1, 0, +1
3
6
2
5d
-2, -1, 0, +1, +2
5
10
3
5f
-3,-2,-1, 0,+1,+2,+3
7
14
4
5g*
-4…0…+4
9
18
0
6s
0
1
2
1
6p
-1, 0, +1
3
6
2
6d
-2, -1, 0, +1, +2
5
10
3
6f
-3,-2,-1, 0,+1,+2,+3
7
14
4
6g*
-4…0…+4
9
18
5
6h*
-5…0…+5
11
22
2l + 1
2[2l + 1]
3 (M)
4 (N)
5 (O)
6 (P)
18
32
50
72
2n2
*These orbitals are not used in the ground state of any known elements.
Problem: Determine the element that has the following set of quantum numbers:
1.
n = 4, l = 3, ml = 2, ms = +½
2.
n = 3, l = 1, ml = 0, ms = -½
The Atom
16
Inorganic Chemistry
3.
Shapes and Orientation of Orbitals:
's' orbitals are spherical and centered around the nucleus
'p' orbitals are propeller shaped, i.e., a twin-bladed propeller. They have a region of zero electron
density (a node) between the two blades. There are 3 different p orbitals, px, py, & pz, oriented
along the x, y and z axes of the 3-dimensional molecule, respectively.
'd' and 'f' orbitals have more complex shapes.
Ground State Electron Configuration of Elements:
In writing ground state (lowest energy) electron configuration, 3 principles are followed: the Aufbau
principle, the Pauli exclusion principle, and Hund's Rule.
Aufbau Principle (German = 'building up'): For each atom, the correct number of electrons are
added to fill atomic orbitals in order of lowest to highest energy, i.e., lowest energy orbitals are
filled first.
Energy Level Diagram of Atomic Orbitals (Showing Overlap of Energy of Shells)
8s
7p
6d
5f
7s
6p
5d
6s
4f
5p
4d
5s
4p
3d
4s
3p
E
N
E
R
G
Y
3s
2p
2s
1s
l=0
The Atom
l=1
l=2
l=3
17
Inorganic Chemistry
Two Memory Aids for the Aufbau Filling Order of Atomic Orbitals
1s
ns
2s
2p
3s
3p
3d
4s
4p
4d
5s
5p
5d
6s
6p
6d
7s
7p
Write all orbitals of the same shell on the same
horizontal line.
Write all orbitals of the same type in the same
vertical column.
Draw parallel arrows diagonally from upper right
to lower left.
Arrows are read from bottom to top, from tail to
head.
(n-1)d
np
H
1s
Li
2s
2p
Ne
Na
3s
3p
Ar
K
4s
3d
4p
Kr
Rb
5s
4d
5p
Xe
Cs
6s
4f
5d
6p
Rn
Fr
7s
5f
6d
7p
Uuo
4f
5f
(n-2)f
He
Look at a periodic table. Shown above is the layout
of the s-, p-, d- and f-blocks on the periodic table,
i.e., this is the filling order. Note the patterns.
The filling order is ns, (n-2)f, (n-1)d, np.
The p-orbitals begin filling after the 2s orbital.
The d-orbitals begin filling after the 4s orbital.
The f-orbitals begin filling after the 6s orbital.
Note that the 4s orbital is slightly lower in energy than (and filled before) the 3d orbital. In general,
the ns orbital is filled before the (n-1)d orbital. This is referred to as the (n-1) rule.
Note that the ns orbital is filled immediately before the (n-2)f orbital, i.e., [6s immediately before 4f]
and [7s immediately before 5f]
Problem: Write out the filling order of atomic orbitals from 1s to 8s.
Pauli Exclusion Principle: A maximum of two electrons can reside in an orbital. When two
electrons occupy the same orbital, their spins are paired (opposite). It is sometimes stated as: No
two electrons can have the same set of 4 quantum numbers.
Moving electrons produce a magnetic field ('magnetic induction'). Electrons are negatively
charged. Like charges repel. Electrons occupying the same orbital repel each other. Two
electrons in the same orbital have the least repulsion when their spins are paired (opposite)
producing opposite magnetic fields. Opposite magnetic fields attract each other - like the north
and south poles of a permanent magnet.
The electron configuration of atoms is shown using a notation in which the number of electrons in
each orbital is written as a superscript. The orbital is shown as a line, _ or as a circle, O. Each
electron in the orbital is written as an arrow, . The direction of the arrow is either up,  ,
(indicating clockwise rotation) or down,  , (indicating counterclockwise rotation). See below.
The Atom
18
Inorganic Chemistry
Full Orbital Notation
1s
2s
2px
2py
Simplified Orbital Notation
2pz
1H
__
1s1
2He
__
1s2
3Li
__
__
2
4Be
__
__
5B
__
__
__
__
__
or
[He]
1s
1
2s
or
[He]
2s1
1s2
2s2
or
[He]
2s2
1s2
2s2
or
[He]
2s2
2p1
2p1
In Simplified Orbital Notation, only the electrons of the outermost shell (the valence electrons) are
listed. The inner (core) electrons are represented by the symbol of the noble gas with the same
electron configuration, e.g., [He], [Ne], [Ar], etc.
Hund's Rule of Maximum Multiplicity: All orbitals of the same energy (degenerate orbitals) are
singly filled before any are doubly filled. This is because electrons repel each other and thus
naturally spread out and do not pair up until forced to.
Full Orbital Notation
Simplified Orbital Notation
1s
2s
2px
2py
2pz
7N
__
__
__
__
__
__
__
__
__
__
1s2
1s2
2s2
2s2
2p2
2p3
or
or
[He]
[He]
2s2
2s2
2p2
2p3
8O
__
__
__
__
__
1s2
2s2
2p4
or
[He]
2s2
2p4
9F
__
__
__
__
__
__
__
__
1s2
1s2
2s2
2s2
2p5
2p6
or
or
[He]
[Ne]
2s2
2p5
10Ne
__
__
11Na
__
__
__
__
__
1s2
2s2
2p6
3s1
[Ne]
3s1
6C
3s
__
Elements with unpaired electrons (e.g., Li, B etc.) are termed paramagnetic. They are attracted
by magnetic fields as the spinning of the unpaired electrons becomes aligned with the applied
magnetic field. In many cases the effect is weak and hardly noticeable.
Elements without unpaired electrons (e.g., He, Be, Ne, etc.) are termed diamagnetic. They are
not attracted (and slightly repelled by) external magnetic fields.
Problem: Write out the electron configuration for Mg through Ar in both 'Full Orbital Notation' and
'Simplified Orbital Notation'.
Full Orbital Notation
3s
12Mg
[Ne]
3px
3py
Simplified Orbital Notation
3pz
[Ne]
3s2
13Al
14Si
15P
16S
17Cl
18Ar
The Atom
19
Inorganic Chemistry
Problem: Using both 'Full Orbital Notation' and 'Simplified Orbital Notation', write out the ground
state electron configuration of K through Kr. Recall the (n-1) rule. Chemical and spectrographic
evidence shows that the configurations of Cr & Cu have only 1 electron in their 4s orbital. Halffilled and filled sets of orbitals have special stability
Full Orbital Notation
4s
19K
3dxy
3dxz
[Ar]
3dyz
3dx2-y2
Simplified
3dz2
4px
4py
4pz
[Ar]
4s1
20Ca
21Sc
22Ti
23V
24Cr
25Mn
26Fe
27Co
36Kr
Layout of the Periodic Table:
Elements are listed in order of increasing atomic number (1 through 114) from left to right and top
to bottom.
7 horizontal rows, called periods, correspond to 7 principal energy levels (shells), n = 1 to 7.
The are 18 vertical columns, 8 of which are labeled as 'A' group elements and 10 of which are
labeled as 'B' group elements.
Elements in the same group (vertical column) are called families because they have similar
chemical and physical properties (because they have similar outer electron configurations).

For example, the Group 1A elements (Alkali Metals) are all soft, low melting, very reactive
metals with similar outer electron configuration (ns1):
Li = [He] 2s1, Na = [Ne] 3s1, K = [Ar] 4s1, Rb = [Kr] 5s1, Cs = [Xe] 6s1 Fr = [Rn] 7s1

For example, the group 8A elements (Noble Gases) are all unreactive (inert), monatomic gases
with an ns2np6 outer electron configuration (a stable octet).
He = 1s2, Ne = 2s2 2p6, Ar = 3s2 3p6, Kr = 4s2 4p6, Xe = 5s2 5p6, Ra = 6s2 6p6.
The Atom
20
Inorganic Chemistry
Learn the names of the eight A-group families:
1A = alkali metals
2IA = alkaline earth(s) metals
3A = aluminum group
4A = carbon group
5A = pnicogens
6A = chalcogens
7A = halogens
8A = noble gases
Elements in the eight A-group columns collectively are referred to as the representative elements.
Their group numbers equal the number of electrons in the outermost shell (highest
n-value). These outer electrons are least strongly held by the nucleus are thus are the electrons used
in bonding. Bonding electrons are called valence electrons. For example, all halogens (Group 7A)
have 7 valence electrons, i.e., ns2 np5.
For B-group elements, the group number equals the number of [ns + (n-1)d] electrons for the first six
groups only (Sc to Fe groups). For example, Mn (in Group 7B) has 7 outer electrons (4s2 + 3d5). This
does not apply to the last four B groups (Co through Zn groups).
The periodic table is also divided into blocks:
s block: electrons are filling the ns orbital (includes group 1A & 2A)
p block: electrons are filling the np orbital (includes group 3A to 8A)
d block: electrons are filling the (n-1)d orbital. d-block elements are also called transition
metals since their properties are transitional between reactive metals on the left side and less
reactive metals and nonmetals on the right. There are 4 series of 10 transition metals:
1. (4s 3d)
1st transition series: 21Sc through 30Zn
2. (5s 4d)
2nd transition series: 39Y through 48Cd
3. (6s 4f 5d)
3rd transition series: 57La and 72Hf through 80Hg
4. (7s 5f 6d)
4th transition series: 89Ac and 104Rf through 112Uub
In general, transition metals are high melting (mp > 1000 C), less reactive metals forming
brightly colored aqueous solutions. Some are noble metals, e.g., Au, Ag, Pt, Pd.
f block: electrons are filling the (n-2)f orbital. f-block elements are also called inner transition
elements. There are 2 series of 14 inner transition metals:
1. (6s 5d1 4f)
1st inner transition series - the lanthanides or rare earths
58Ce through 71Lu
2. (7s 6d1 5f)
2nd inner transition series - the actinides - 90Th through 103Lr
f-block elements are placed below the periodic table. If they were placed within the table in
their filling order, the periodic table would be too wide to fit on a single page. Separating them
also emphasizes the unique properties of these groups.
Adding inner f electrons appears to have little effect on chemical properties. All lanthanides
are fairly reactive metals. All actinides are radioactive and only radium and thorium are found
in appreciable amounts in nature (others are produced in controlled nuclear reactions).
Uranium and plutonium are used as fuels in nuclear reactors and nuclear weapons.
Problem: Write out the electron configuration in simplified notation of
51
Sb, 55Cs, 79Au, 81Tl, 82Pb, 83Bi, 84Po & 87Fr.
The Atom
42
Mo,
48
Cd,
47
Ag,
40
Zr,
49
21
In,
Inorganic Chemistry
Periodic Trends in the Periodic Table (Periodicity):
All physical and chemical behavior of the elements is based ultimately on the electron configuration of
the atoms. Properties such as mp, bp, volume, acidity, and reactivity generally increase or decrease
in a recurring manner through the Periodic Table. These consistent trends within groups and periods
is referred to as periodicity.
Atomic radii: range from 31 pm in He to 262 pm in Cs.
All are smaller than the wavelengths of visible light (400 to 800 nm) and hence all atoms
are invisible to visible light - even the most powerful optical microscope could not resolve
the atom. X-rays (with  < 1 pm) can be now be used to 'see' atoms.
Atomic radii increase down each group as more layers (shells) of electrons are added to the atom.
Atomic radii decrease from left to right across the periods as more electrons are added to the same
shell. Increasing atomic number means an increasing positive nuclear charge (more protons) which
pulls the outer shell of electrons in closer to the nucleus. The shielding effect of the inner core of
electrons remains the same across each period, but the nuclear charge is increasing (with more
protons). Hence the net core charge increases across the periods and decreases the atomic radii.
3 Li
4 Be
5B
3P
6C
7N
8O
9F
10 Ne
4P
5P
6P
7P
8P
9P
10P
134 pm
125 pm
90 pm
77 pm
75 pm
73 pm
71 pm
69 pm
520 kJ/m ol
899
801
1086
1402
1314
1681
2081
+1
+2
+3
+4
+5
+6
+7
+8
EN = 1.0
1.5
2.0
2.5
3.0
4.0
-----
11 Na
11P
1st shell (1s)
2nd shell (2s & 2p)
3rd shell (3s, 3p & 3d)
atomic radius
1st ionization energy
223 pm
net core charge
3.5
The core is all of the atom excluding its valence electrons.
The net charge of the core = (# protons - # inner electrons),
where inner electrons are all except the outer, valence shell
electrons.
As the net core charge increases left to right across each
row of the periodic table, atomic radii decrease and the
first ionization energies increase.
Net core charge is constant down each group but atomic
radii increase (as more shells of electrons are added) and
ionization energies decrease (since the valence electrons are
progressively farther from the attractive force of the nucleus.
496 kJ/mol
+1
electronegativity
EN = 1.0
Problem: Calculate the net core charge for Na to Ar.
The Atom
22
Inorganic Chemistry
250
227
Atomic Radii (pm)
197
200
186
160
143
Atomic 150
Radius
(pm) 100
132
118
112
110 103
85
50
77
75
73 72
N
O
100 98
71
37 31
0
H
He
Li
Be
B
C
F
Ne Na Mg
Al
Si
P
S
Cl
Ar
K
Ca
Atomic Radius vs. Atomic Number
Atomic Radius (pm)
300
1st
Transitio n
Series
250
2nd
Transitio n
Series
Rb
3rd
Transitio n
Series
Cs
K
200
Na
Li
150
Rn
Xe
Kr
100
Ar
Lanthanides
Ne
50
He
0
0
10
20
30
40
50
60
70
80
90
100
Atomic Number
The Atom
23
Inorganic Chemistry
Ionization Energy (IE): (called the 1st ionization potential) is the amount of energy required to
remove the most loosely bound electron from an isolated atom to form a cation of +1 charge.
e.g.,
Ca (g) + 590 kJ  Ca+ + 1 e-
The 2nd ionization energy is the amount of energy required to remove a second electron and is always
higher than the 1st (because it is more difficult to remove an electron from a cation than a neutral
atom.
e.g.,
Ca+(g) + 1145 kJ  Ca+2 + 1 e-
A low IE indicates that electrons are easily removed. Low IE is characteristic of metals. Alkali metals
have the lowest IE. In general metals react by donating (losing) electrons.
IE increases left to right across each period and decreases down each group. Thus the lower left
corner of the periodic table contains the most reactive metals.
First Ionization Energy vs. Atomic Number
2500
First Ionization Energy (kJ/mol)
He
Ne
2000
Ar
1500
Kr
H
Xe
O
1000
Zn
Cd
S
Se
Te
B
Al
500
LI
Ga
Na
In
Rb
K
Cs
0
0
10
20
30
40
50
60
Atomic Number
Group 3A elements (B, Al, Ga, In, Tl) are exceptions to the general horizontal trends. Their IE's are
lower than those of group 2A because the 3A elements have only a single electron in their outermost
p orbitals and less energy is required to remove a single p-orbital electron than the second s-orbital
electron from the same shell because the ns orbital is lower in energy than the np orbital.
Group 6A elements (O, S, Se, Te, Po), like the 3A elements are exceptions to the horizontal trend.
They have slightly lower IE than the 5A elements in the same periods. This tells us that less energy
is required to remove a paired electron from a 6A element than to remove an unpaired p electron
from a 5A element. Removal of one electron from the 6A elements gives a half-filled set of p orbitals.
Half-filled and completely-filled orbitals have special stability.
One factor that favors an atom of a representative element forming an ion in a compound is the
formation of a stable noble gas electron configuration. Atoms generally gain, lose or share electrons
to become isoelectronic with the nearest noble gas.
The Atom
24
Inorganic Chemistry
Knowledge of the relative values of IE assists us in predicting whether an element is likely to form
ionic or covalent molecular compounds.

Elements with low IE (metals) readily form ionic compounds when reacting with elements
which gain electrons (non metals).

Elements with intermediate IE generally form covalent molecular compounds by sharing
electrons with other elements.

Elements with high IE (non metals of Group 6A & 7A) often gain electrons from metals
forming ionic compounds or share electrons with other non metals forming covalent
molecular compounds.
Electron Affinity (EA): is the amount of energy required for an isolated gaseous atom to accept an
electron and form an anion with a -1 charge. For most elements this process is exothermic,
particularly for non metals, which need electrons to complete their octet.
Electronegativity: is a measure of the force of an atom’s attraction for electrons that it shares in a
chemical bond with other atoms. Electronegativity is much more useful than electron affinity.
In the 1930’s, Linus Pauling assigned electronegativity values to all elements relative to F (the most
electronegative element), which he gave a value of 4.0 .
Linus Pauling's Table of Electronegativities
H
2.1
Li
1.0
Na
1.0
K
0.9
Rb
0.9
Cs
0.8
Fr
0.8
He
Be
1.5
Mg
1.2
Ca
1.0
Sr
1.0
Ba
1.0
Ra
1.0
Sc
1.3
Y
1.2
La
1.1
Ac
1.1
Ti
1.4
Zr
1.3
Hf
1.3
V
1.5
Nb
1.5
Ta
1.4
Cr
1.6
Mo
1.6
W
1.5
Mn
1.6
Tc
1.7
Re
1.7
Fe
1.7
Ru
1.8
Os
1.9
Co
1.7
Rh
1.8
Ir
1.9
Ni
1.8
Pd
1.8
Pt
1.8
Cu
1.8
Ag
1.6
Au
1.9
Zn
1.6
Cd
1.6
Hg
1.7
B
2.0
Al
1.5
Ga
1.7
In
1.6
Tl
1.6
C
2.5
Si
1.8
Ge
1.9
Sn
1.8
Pb
1.7
N
3.0
P
2.1
As
2.1
Sb
1.9
Bi
1.8
O
3.5
S
2.5
Se
2.4
Te
2.1
Po
1.9
F
4.0
Cl
3.0
Br
2.8
I
2.5
At
2.1
Ne
Ar
Kr
Xe
Rn
In general, both ionization energies and electronegativities are low for elements at the lower left of the
periodic table and high for those at the upper right.
One way to estimate the degree of ionic or covalent character in a chemical bond is to compare
electronegativities of atoms involved. The less electronegative element gives up its electrons to the
more electronegative element. Two non metals with similar electronegativities share electrons to
form covalent bonds.
The Atom
25
Inorganic Chemistry
Generalizations Regarding the Elements:
Metals:
Of the 114 known elements, 89 are metals, i.e., those to the left of the staircase including Al,
excluding metalloids and H.
Metals include:
 Groups 1A and 2A
 Heavier Group 3A elements (Al, Ga, In, Tl), Group 4A (Sn, Pb) and Group 5A (Bi)
 The transition elements (d block) and the Lanthanides & Actinides - f block (inner transition
elements).
All metals possess to varying degrees, the following physical properties:
1. High Electrical Conductivity: Silver has the highest electrical conductivity (lowest electrical
resistance). Mercury is one of the poorest metallic conductors but finds many applications as
a liquid electrical switch.
2. High Thermal Conductivity: Among solids, metals are by far the best conductors of heat.
3. Luster: Most metals have a silvery white appearance (when polished) indicating that light of
all wavelengths is reflected. Gold and copper absorb some light in the blue region of the
spectrum and hence appear yellow and orange, respectively.
4. Ductility, Malleability: Most metals are ductile (capable of being drawn out into a wire) and
malleable (capable of being hammered into thin sheets).
Non Metals:
Clustered toward the upper right hand corner of the Periodic Table are 17 nonmetals. They have few
metallic properties.

Except for Se or the graphite form of C, they are nonconductors of electricity and heat (electrical
and thermal insulators).

With few exceptions, notably diamond, crystals of nonmetals have a dull rather than shiny
appearance.

All solid nonmetals shatter if drawn out or hammered.

Of the nonmetals, Group 8A (the noble gases) are unique. They exist as monatomic gases and
show no tendency to combine with one another or other elements. Their ns2np6 electron
configuration is unusually stable.

In contrast to the noble gases, most other nonmetals form polyatomic molecules in the gaseous
state (N2, P4, O2, S8, F2, Cl2, Br2, and I2)
Metalloids (Semimetals):
With the exception of Al, all elements touching 2 sides of the staircase (on the periodic table) are
metalloids, i.e., B, Si, Ge, As, Sb, Te, Po and At (8 in all). Po and At do not occur naturally.

Metalloids have properties between those of metals and nonmetals. All show metallic luster.

Metalloids typically are semiconductors, although As and Sb actually have electrical conductivities
which approach those of metals. Si and Ge are most important as semiconductors. In contrast
to metals, their electrical conductivity increases when temperature is raised.
The Atom
26
Inorganic Chemistry
The Inert-Pair Effect: Although both Al and In are in Group 3A, Al forms Al+3 ions only, whereas In
forms both In+3 and In+ ions. The tendency to form ions two units lower in charge than expected from
the group number is called the inert-pair effect.
Other examples of the inert-pair effect are found in Group 4A: tin forms both SnO and the more stable
oxide, SnO2. Likewise lead forms both PbO and PbO2.
The inert-pair effect is due in part to the different energies of the valence p- and s-electrons. In the
later periods, valence p-electrons are relatively high in energy because of the shielding (reduced
nuclear attraction) provided by the (n-1)d-electrons. Thus p-electrons are more readily removed than
s-electrons of the same shell.
Study a periodic table and note this tendency in the lower periods of Groups 3A, 4A and 5A.
Diagonal Relationships:
H
Li and Mg are chemically similar. Both react
directly with nitrogen to form nitrides.
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
Be and Al are chemically similar. Both are
amphoteric (react with acids and bases).
K
Ca
Ga
Ge
As
Se
Br
Kr
B through At are all metalloids.
Rb
Sr
In
Sn
Sb
Te
I
Xe
Cs
Ba
Tl
Pb
Bi
Po
At
Rn
Fr
Ra
These diagonal similarities make some
sense if we recall that metallic properties
increase down each group but decrease left
to right across each period. By moving
diagonally down and to the right we
encounter chemical similarities. What are
the EN values along the diagonal lines?
Acid-base Behavior of the Element Oxides:
Acidity of the element oxides increases up each group and left to right across each period.
Basicity of the element oxides increases down each group and right to left across each period.
Metal oxides are basic.
When dissolved in water they produce OH- ion.
Na2O + H2O  2 NaOH
CaO + H2O  Ca(OH)2
Nonmetal oxides are acidic.
When dissolved in water they produce acids.
CO2 + H2O  H2CO3
SO3 + H2O  H2SO4
Oxides of intermediate elements are amphoteric.
Al2O3 + HCl  AlCl3 + H2O
Al2O3 +
NaOH 
H2O
Li2O
BeO
B2O3
CO2
N2O5
F2O
Na2O
MgO
Al2O3
SiO2
P4O10
SO3
Cl2O7
K2O
CaO
Ga2O3
GeO2
As2O5
SeO3
Br2O7
Rb2O
SrO
In2O3
SnO2
Sb2O5
TeO3
I2O7
Cs2O
BaO
Tl2O3
PbO2
Bi2O5
PoO3
At2O7
Fr2O
RaO
Shaded oxides are amphoteric.
NaAlO2 + H2O
Problem: Write balanced chemical equations for the reaction of the water with the following metal
oxides: Li2O, K2O, MgO, BaO, Al2O3 [producing Al(OH)3], CO2, SiO2, N2O5, P4O6, P4O10, As2O3,
As2O5, SO2, SO3, Cl2O7, Br2O7, I2O5 and I2O7.
The Atom
27
Inorganic Chemistry
Using the Periodic Table we can often correctly predict the formulas of compounds based on known
formulas of analogous compounds containing elements in the same groups as the known
compounds. For example, given that the normal chloride salt of magnesium = MgCl2, we are not
surprised to find the other Alkaline Earth salts to be BeCl2, CaCl2, SrCl2, BaCl2, and RaCl2. Caution
must be exercised. For example, 2nd period oxyacids often have unique formulas. The lower periods
are more consistent.
Problem: Compounds containing elements from the same groups of the Periodic Table often have
similar formulas..
Given the formulas for sodium chlorate (NaClO3), barium chromate (BaCrO4) and sodium phosphate
(Na3PO4), predict the formulas of
a. potassium arsenate
b. strontium tungstate
c. rubidium bromate
d. sodium molybdate
Problem: Predict the formula of selenic acid.
Trends in Acidity of Oxy Acids:
Oxyacids, like H2SO4, contain, in addition to a nonmetal such as S, both hydrogen and oxygen. Since
they contain 3 different types of atoms, oxyacids are also called ternary acids. The trends in acidity
of oxyacids follow the same pattern as seen in the acidity of nonmetal oxides.
Lower pKa means more acidic. See the chart below. For polyprotic acids (with more than one acidic
H’s), like H2SO4, H3PO4, etc., the pKa listed is for the dissociation of the first H, i.e., pKa1
Group
nd
2
Period
3A
H3BO3
Group
4A
Group
5A
Group
6A
Group
7A
H2CO3
carbonic acid
pKa = 6.4
HNO3
nitric acid
pKa = -1.4
H2SiO3
H2SO4
sulfuric acid
HClO4
perchloric acid
pKa = 9.8
H3PO4
phosphoric
acid (ortho)
pKa = 2.1
pKa = -5
pKa = -7
4
Period
H4GeO4
germanic acid
pKa = 8.6
H3AsO4
arsenic acid
pKa = 2.2
H2SeO4
selenic acid
pKa = 1.7
HBrO4
perbromic acid
pKa = ca. 2
5th
Period
H2SnO3 or
SnO2H2O
stannic acid or
stannic oxide
hyrate
amphoteric
H7Sb(OH)6
antimonic acid
H6TeO6
telluric acid
pKa = 7.7
HIO4
periodic acid
pKa = 1.6
boric acid (ortho)
pKa = 9.2
3rd
Period
th
The Atom
Al(OH)3 or
Al2O33H2O
amphoteric
silicic acid (meta)
H2O
(not a ternary acid)
28