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Chapter 1
Basic concepts: atoms
TOPICS
 Fundamental particles.
 Atomic number, mass number and isotopes.
 An overview of quantum theory.
 Orbitals of the hydrogen atom and quantum
numbers.
 The multi-electron atom, the aufbau principle
and electronic configurations.
 The periodic table.
 Ionization energies and electron affinities.
1.1 Introduction
 Inorganic chemistry is the chemistry of all elements
except carbon.
 Inorganic chemistry is not simply the study of elements
and compounds; it is also the study of physical
principles.
Examples
1- In order to understand solubility, we apply laws of
thermodynamics.
2- To propose details of a reaction mechanism, then
a knowledge of reaction kinetics is needed.
3- Overlap between physical and inorganic chemistry
is also significant in the study of molecular structure.
Inorganic chemistry: it is not an isolated branch of
chemistry
1.2 Fundamental particles of an atom
An atom is the smallest unit quantity of an element that
is capable of existence, either alone or in chemical
combination with other atoms of the same or another
element. The fundamental particles of which atoms are
composed are the proton , electron and neutron.
 The nucleus of an atom consists of protons and (with the
exception of protium) neutrons, and is positively charged;
the nucleus of protium consists of a single proton.
 The electrons occupy a region of space around the
nucleus.
 Nearly all the mass of an atom is concentrated in the
nucleus.
The volume of the nucleus is only a tiny fraction of that of
the atom.
The radius of the nucleus is
about 10-15 m while the atom
itself is about 105 times larger
than this.
1.3 Atomic number, mass number and isotopes
Relative atomic mass
Atomic mass unit (amu) is 1/12 the mass of 12C (1.660 × 10-27 kg)
Relative atomic masses ( Ar) are thus all stated relative to
12C
= 12.0000. The masses of the proton and neutron can be
considered to be 1 u where u is the atomic mass unit (1 u
1.660 x10-27 kg).
Isotopes
Nuclides of a particular element that differ in the number of
neutrons and, therefore, their mass number, are called isotopes .
Isotopes of some elements occur naturally while others may be
produced artificially.
Elements that occur naturally with only one nuclide are monotopic
Elements that exist as mixtures of isotopes
Since the atomic number is constant for a given element, isotopes
are often distinguished only by stating the atomic masses,
e.g.12C and 13C.
Isotopes and allotropes
Do not confuse isotope and allotrope ! Sulfur exhibits both
isotopes and allotropes.
Allotropes of an element are different structural modifications of
that element.
Allotropes of sulfur include cyclic structures, e.g. S6 and S8, and
Sx-chains of various lengths (poly catenasulfur).
Allotropes: element's atoms are bonded together in a different
manner
Isotopes can be separated by mass spectrometry
Figure 1.1b shows a mass spectrometric trace for molecular S8,
the structure of which is shown in Figure 1.1c; five peaks are
observed due to combinations of the isotopes of sulfur.
1.4 Successes in early quantum theory
The development of quantum theory took place in two stages. In
the older theories (1900–1925), the electron was treated as a
particle, and the achievements of greatest significance to
inorganic chemistry were the interpretation of atomic spectra and
assignment of electronic configurations.
In more recent models, the electron is treated as a wave (hence
the name wave mechanics) and the main successes in chemistry
are the elucidation of the basis of stereochemistry and methods
for calculating the properties of molecules
Some important successes of classical quantum theory
Planck suggested that energy could be absorbed or emitted only
in quanta of Magnitude E related to the frequency of the
radiation,  ,
The Planck constant (h = 6.626 x10-34 J s)
The hertz, Hz, is the SI unit of frequency.
 is the wavelength
c is the speed of light in a vacuum (c = 2.998 x108ms-1)
Rutherford–Bohr model of the atom. When an electric discharge is
passed through a sample of dihydrogen, the H2 molecules
dissociate into atoms, and the electron in a particular excited H
atom may be promoted to one of many high energy levels.
These states are transient and the electron falls back to a lower
energy state, emitting energy as it does so. The consequence is
the observation of spectral lines in the
emission spectrum of hydrogen
Balmer pointed out that the wavelengths of the spectral lines
observed in the visible region of the atomic spectrum of
hydrogen obeyed the following equation.
n is an integer 3, 4, 5 .
Other series of spectral lines occur in the ultraviolet (Lyman
series) and infrared (Paschen, Brackett and Pfund series)
All lines in all the series obey the general expression given in the
following equation:
For the Lyman series, n =1,
For the Balmer series, n = 2, and
For the Paschen, Brackett and Pfund series, n = 3, 4 and 5
respectively
Energy is absorbed or emitted only w he n an electron moves
from one stationary state to another and the energy change is
given by the following equation:
If we apply the Bohr model to the H atom, the radius of each
allowed circular orbit can be determined
An increase in the principal quantum number from n = 1 to n =
has a special significance.
Ionization energy, IE, can be determined
Although the SI unit of energy is the joule, ionization energies are
often expressed in electron volts (eV) (1 eV = 96:4853 = 96:5 kJ
mol-1). Therefore, the ionization energy of hydrogen can also be
given as 13.60 eV.
1.5 An introduction to wave mechanics
The wave-nature of electrons
In 1924, Louis de Broglie argued that if light were composed of particles and yet
showed wave-like properties, the same should be true of electrons and other
particles. This phenomenon is referred to as wave–particle duality.
momentum mv ( m = mass and v = velocity of the particle)
The uncertainty principle
Heisenberg’s uncertainty principle:
it is impossible to know exactly both
the momentum and position of the
electron at the same instant in time
In order to get around this problem, rather than trying to define its
exact position and momentum, we use the probability of finding
the electron in a given volume of space.
1.6 Atomic orbitals
Degenerate orbitals possess the same energy.
The radial part of the wavefunction,R (r )
Fig. 1.5 Plots of the radial parts of the wavefunction, R ( r ), against distance, r ,
from the nucleus for (a) the 1s and (b) the 2 s atomic orbitals of the hydrogen
atom; the nucleus is at r = 0. The vertical scales for the two plots are different
but the horizontal scales are the same.
Fig. 1.6 Plots of radial parts of the wavefunction R (r ) against r for the 2 p,3p,4p
and 3d atomic orbitals; the nucleus is at r =0.
The radial distribution function, 4  r2R(r )2
A useful way of depicting the probability density is to plot aradial
distribution function and this allows us to envisage the region in
space in which the electron is found.
The probability of finding the electron somewhere in
space is taken to be 1.
Each function is zero at the nucleus, following from the r2 term
and the fact that at the nucleus r = 0.
Each plot of 4r2R(r)2 shows at least one maximum value for the
function, corresponding to a distance from the nucleus at
which the electron has the highest probability of being found.
Points at which 4r2R(r)2 = 0 (ignoring r = 0) correspond to radial
nodes where R(r) = 0.
The angular part of the wavefunction, A (, )
The amplitude of a wavefunction may be positive or negative;
this is shown using + and - signs, or by shading the lobes in
different colours.
chemists often indicate the amplitude by a sign or by shading
because of the importance of the signs of the wavefunctions with
respect to their overlap during bond formation
Orbital energies in a hydrogen-like species
For the hydrogen atom, Z = 1, but for the hydrogen-like He+ ion,
Z = 2. The dependence of E on Z2 therefore leads to a significant
lowering of the orbitals on going from H to He+.
The spin quantum number and the magnetic spin quantum number
Angular momentum, the inner quantum number, j , and spin--orbit coupling
1.7 Many-electron atoms
Ground state electronic configurations: experimental data
The following sequence is approximately true for the relative
energies (lowest energy first) of orbitals in neutral atoms :
Some irregularities: see Table 1.3 in textbook
Penetration and
shielding
Effective nuclear charge and Slater’s rules
Thus, an electron in the 4s (rather than the 3d) atomic
orbital is under the influence of a greater effective
nuclear charge and in the ground state of potassium,
it is the 4s atomic orbital that is occupied.
Collective names for some of the groups of elements in the
periodic table are given in Table 1.4.
The periodic table in Figure 1.13 shows elements up to Z =112.
Three additional elements with Z = 114, 116 and 118 have been
reported, but still await IUPAC authentication.
1.9 The aufbau principle (aufbau means ‘building up’ in German)
Ground state electronic configurations
. Orbitals are filled in order of energy, the lowest energy orbitals
being filled first.
. Hund’s first rule (often referred to simply as Hund’s rule): in a set
of degenerate orbitals, electrons may not be spin-paired in an
orbital until each orbital in the set contains one electron;
electrons singly occupying orbitals in a degenerate set have
parallel spins, i.e. they have the same values of ms.
. Pauli exclusion principle: no two electrons in the same atom may
have the same set of n, l, ml and ms quantum numbers; it follows
that each orbital can accommodate a maximum of two electrons
with different ms values (different spins = spin-paired).
Valence and core electrons
The configuration of the outer or valence electrons is of particular significance.
These electrons determine the chemical properties of an element
Electrons that occupy lower energy quantum levels are called core electrons
Diagrammatic representations of electronic configurations
In this case, the electrons
are represented by arrows
with the direction of the
arrow corresponding to
1.10 Ionization energies and electron affinities
The first ionization energy, IE1, of an atom is the internal energy change
at 0 K, U(0 K), associated with the removal of the first valence electron
(equation 1.17); the energy change is defined for a gas-phase process.
The units are kJ mol-1 or electron volts (eV)
IE2
IE3
Several repeating patterns are apparent and some features to note are:
. the high values of IE1 associated with the noble gases;
. the very low values of IE1 associated with the group 1 elements;
. the general increase in values of IE1 as a given period is crossed;
. the discontinuity in values of IE1 on going from an element in group 15
to its neighbour in group 16;
. the decrease in values of IE1 on going from an element in group 2 or
12 to its neighbour in group 13;
. the rather similar values of IE1 for a given row of d-block elements.
Electron affinities
The first electron affinity (EA1) is minus the internal energy change for
the gain of an electron by a gaseous atom. The second electron affinity
of atom Y is defined for process. Each reaction occurs in the gas phase.