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Pedagogical Comment For Learners
As a learner, the pre-assessment test evaluates your present level of chemistry
knowledge as a link to that knowledge you are to acquire in this Inorganic chemistry
1. Your test score should help in identifying your competence and indicate areas where
you need special emphasis on. The basics of understanding Inorganic chemistry 1 lies
in appreciating the effects of electronic configurations and their concomitant interac-
tions in atoms, ions, molecules, compounds etc as they direct the periodicity of the
elemental properties. A learner who scores 40 percent or less in the pre-assessment
test is likely to encounter difficulties comprehending the contents of this module and
is, therefore, advised to review Introductory Chemistry 1, which is a prerequisite to
this course. However, your performance index is not in anyway intended to make
you be discouraged or be complacent; it is for you to appreciate how much effort you
need to put in this work, be ready to make that extra mile.
Key Concepts
Atomic number. Is the number of protons in the nucleus or the number of electrons
in an atom.
An ion. A charged atom or molecule. An ion is positive (cation) if it has lost electrons or negative (anion) if it has gained electrons
Isotopes. One member of a (chemical-element) family of atomic species which has
two or more nuclides with the same number of protons ( Z) but a different number
of neutrons ( N). Because the atomic mass is determined by the sum of the number
of protons and neutrons contained in the nucleus, isotopes differ in mass. Since they
contain the same number of protons (and hence electrons), isotopes have the same
chemical properties.
Allotropes. One or more forms of an elementary substance. Examples are Graphite and diamond are both allotropes of carbon. O and ozone, O , are allotropes of oxygen.
2
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Electronic ground state. This is electronic configuration of an atom with the lowest
energy orbitals all accupied according to Hund’s rule.
Isoelectronic series. A series for atoms or ions that have the same electronic arran-
gements/configuration.
Electroneutrality. The principle expresses the fact that all pure substances carry a
net charge of zero. That is the overall charge in a molecule like [Na+Cl-]0 is zero.
Chemical change. A change that results in the formation of a new substance, such
as the burning of wood.
Catalyst. Anything/substance which creates a situation in which change can occur
at a faster rate.
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Atomic mass unit. An atomic mass unit (symbolized AMU or amu) is defined as
precisely 1/12 the mass of an atom of carbon-12. The carbon-12 (C-12) atom has six
protons and six neutrons.
Chemical bond. An attractive force that holds atoms together to form molecules
or Electrical interaction between electrons of one atom and the positive nucleus of
another atom that result in the binding of atoms together in a stable unit.
Alloy. Is a homogeneous mixture of two or more elements, at least one of which
is a metal, and where the resulting material has metallic properties. The resulting
substance usually has different properties (sometimes substantially different) from
those of its components.
Base. A substance that ionises in water to form hydroxide ions and a cation (there
are more fundamental definitions of the term).
Bond polarity. The extent to which the bonding electron pairs between the two atoms
is displaced towards one of the atoms.
.
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X. learning activities
Learning activity # 1
Title of Learning Activity : PERIOdIC TABlE Of ElEmEnTS
At the end of this Unit, the learner should be able to;
1. Describe and predict the position of an element in the periodic table by use
of the atomic numbers.
2. Classify the elements into s-, p-, d-, and f-block elements also as metals, non-
metals and metalloids according to location in the periodic table.
4. Use the periodic table to classify the elements according to IUPAC system.
5. Use different labelling sytems for the periodic table.
Summary of the learning activity
Being the first topic covered in this module, activity 1 includes the historical deve-
lopment of the Periodic table of elements by arranging the elements in horizontal
rows according to their atomic weights. Identification of columns (groups) and rows
(periods) of the periodic table and mark the metallic, non-metallic, and metalloids re-
gions of the table and also as s-, p-, d-, and f-block elements portions are subsequently
disccused. In addition, the elements will then be classified according to IUPAC system
and lastly, different numbering for the modern periodic table will be discussed. At
the end of each topic, relevant worked examples and excersises will follow to to aid
you in development of conceptual and quantitative problem solving skills.
List of Required Readings
Text books
1. Alan G. Sharpe; Inorganic Chemistry, 3rd Edition. Longman Singapore Pu-
blisher, 1992.
2. Catherine E. Housecroft and Alan G. Sharpe; Inorganic Chemistry. Prentice-
Hall International, USA. 2000.
3. J. D. Lee, Concise Inorganic Chemistry, 4th edition. Chapman & Hall, New
York. USA. 1993.
4. Thomas R. Gilbert, Rein V. Kirss, and Geoffrey Davies; Chemistry, The
science in context. W.W. Norton and company NY, USA. 2004.
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List of relevant useful links:
1. http://www.dayah.com/periodic/
2. http://en.wikipedia.org/wiki/Periodic_table
3. http://www.chem1.com/acad/webtext/atoms/atpt-6.html#ORG
4. http://chemistry.about.com/od/elementgroups/a/metals.htm
5. http://chem.lapeer.org/Chem1Docs/Mendeleev.html
6. http://en.wikipedia.org/wiki/Electron_configuration
List of relevant MULTIMEDIA resources:
- Computer with internet connecting facility to access relevant links and free
source resourses.
- Multi-media resourses such as CD players, VCD etc.
- CD-ROM for this module for compulsory reading and demonstrations.
Learning activities
Introduction and historical aspects of Periodic table:
The periodic table of the chemical elements is a tabular method of displaying the
chemical elements. Although precursors to this table exist, its invention is generally
credited to a Russian chemist Dmitri Mendeleev in 1869. Mendeleev invented the
table to illustrate recurring («periodic») trends in the properties of the elements. The
layout of the table has been refined and extended over time, as new elements have
been discovered, and new theoretical models have been developed to explain che-
mical behavior. The periodic table is now ubiquitous within the academic discipline
of chemistry, providing an extremely useful framework to classify, systematize and
compare all the many different forms of chemical behavior. The current, as of October
2006, standard table contains 117 elements (while element 118 has been synthesized,
element 117 has not). Ninety-two elements are found naturally on Earth, and the rest
are synthetic elements that have been produced artificially in particle accelerators.
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Structure of the periodic table
Some definitions:
1. A group is a vertical column in the periodic table of the elements.
2. A period is a horizontal row in the periodic table of the elements.
See the Figure 1.1 below. Aslo available at: http://en.wikipedia.org/wiki/Image:800px-
PTable.png.
figure 1.1: Periodic table of Elements showing the outermost shells and s, p, d, and
f-block regions.
There are more than one way of designation for the groups in the periodic table.
Table 1.1 below compares labelling for the rest of available Periodic tables to that
of figure 1.1.
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Table 1.1; The different ways of labellings used in modern periodic table.
Arabic +
1A
2A
3B
4B
5B
6B
7B
8B
8B
8B
1B
2B
3A
4A
5A
6A
7A
8A
Alphabets
Arabic No.
1
2
1
2
3
4
5
6
7
8
9
10
3
4
5
6
7
8
Roman
IA
IIA
IIIB
IVB
VB
VIB
VIIB
VIII
VIII
VIII
IB
IB
I IA
IVA
VA
VIA
VIIA
VIIIA
No
IUPAC
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
Key. In Arabic numbers (second row), in Blue are for main block while in black for
transition block elements.
n.B. Since the last category is the IUPAC recommended, it will be used for all
subsequent discussion.
Recommended name for Groups of the periodic table
Group number
Recommended name
1
Alkali metals
2
Alkali earth metals
15
Pnicogens
16
Chalcogens
17
Halogens
18
Noble gases / inert gases
Electronic Configuration
Definition. Electronic configuration refers to the way in which electrons are arranged
in the atomic orbitals.
Currently it is more fashinable to look at a Periodic table as determined by the num-
ber and the arrangement of the electrons of the elements. The primary determinant
of an element’s chemical properties is its electron configuration, particularly the
valence shell (outer most) electrons. In addition, the type of orbital in which the
atom’s outermost electrons reside determines the «block» to which it belongs. The
number of valence shell electrons determines the family, or group, to which the
element belongs.
Exercise 1: Arrange the following orbitals; s, p, d, and f interms of their energy levels
(closeness to the nucleus) starting with the lowest (one nearest to the nucleus).
A). s, p, d, f;
B) p, s, f, d;
C) d, f, p, s;
D) f, d, p, s.
Solution is (A).
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Electronic configuration of any element is intimately related to the position of the ele-
ment in the periodic table. The chemical properties of an atom are largely determined
by the arrangement of the electrons in its outermost «valence» shell (although other
factors, such as atomic radius, atomic mass, and increased accessibility of additio-
nal electronic states also contribute to the chemistry of the elements as atomic size
increases) therefore elements in the same table group are chemically similar because
they contain the same number of «valence» electrons.
Summary of the quantum numbers
The state of an electron in an atom is given by four quantum numbers Three of these
are integers and are properties of the atomic orbital in which it sits.
number
denoted allowed values
represents
Partly the overall energy of the orbital,
principal
and by extension its general distance
quantum
n
integer, 1 or more
from the nucleus. In short, the energy
number
level it is in. (1+)
The orbital’s angular momentum, also
azimuthal
seen as the number of nodes in the den-
quantum
l
integer, 0 to n-1
sity plot. Otherwise known as its orbital.
number
(s=0, p=1...)
Determines energy shift of an atomic
magnetic
orbital due to external magnetic field
integer, - l to + l,
quantum
m
(Zeeman effect). Indicates spatial orien-
including zero.
number
tation.
Spin is an intrinsic property of the elec-
+½ or -½ (some-
tron and independent of the other num-
spin quan-
m
times called «up»
bers. s and l in part determine the elec-
tum number
s
and «down»)
tron’s magnetic dipole moment.
n.B. According to Pauli Exclusion Principle; No two electrons in one atom can have
the same set of these four quantum numbers.
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Shells and subshells
Shells and subshells (also called energy levels and sublevels) are defined by the
quantum numbers, not by the distance of its electrons from the nucleus, or even their
overall energy. In larger atoms, shells above the second shell overlap ie. the restriction
is nolonger valid (see Aufbau principle).
States with the same value of n are related, and said to lie within the same electron shell.
States with the same value of n and also l are said to lie within the same electron subshell, and those electrons having the same n and l are called equivalent electrons.
If the states also share the same value of m, they are said to lie in the same atomic orbital.
Because electrons have only two possible spin states, an atomic orbital cannot contain
more than two electrons (Pauli exclusion principle).
A subshell can contain up to 4 l + 2 electrons; a shell can contain up to 2 n 2 electrons; where n equals the shell number.
Worked example
Here is the electron configuration for a filled fifth shell:
Shell Subshell Orbitals
No. of orbitals
Max No. of
electrons
n = 5
l = 0
m = 0
1 type s orbital
2
l = 1
m = -1, 0, +1
3 type p orbitals
6
l = 2
m = -2, -1, 0, +1, +2
5 type d orbitals
10
m = -3, -2, -1, 0, +1,
l = 3
7 type f orbitals
14
+2, +3
m = -4, -3 -2, -1, 0, +1,
l = 4
9 type g orbitals
18
+2, +3, +4
Aufbau principle
Aufbau is a German word meaning to fill-up. It states, electrons enter into states in
order of the states’ increasing energy; i.e., the first electron goes into the lowest-energy
state, the second into the next lowest, and so on. The order in which the states are
filled is as follows:
The order is; 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 5d ≈ 4f < 6p < 7s < 6d ≈5f.
The orbital labels s, p, d, and f originate from a now-discredited system of categorizing spectral lines as sharp, principal, diffuse, and fundamental, based on their observed fine structure. When the first four types of orbitals were described, they were associated with these spectral line types, but there were no other names. The designation
g was derived by following alphabetical order. Shells with more than five subshells
are theoretically permissible, but this covers all discovered elements. Some call the
s and p orbitals spherical and peripheral.
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N.B. The energies of different orbitals are close together for higher values of n (1, 2,
3,..) and their relative energies change significantly upon ionization.
Hunds rule
In general the hund’s rule of maximum multiplicity states that every orbital in a
subshell is singly occupied with one electron before any one orbital is doubly occu-
pied, and all electrons in singly occupied orbitals have the same spin.
Pauli exclusion principle
Simply stated: No two electrons in the same atom can be in the same quantum state.
This means that no two electrons can have the same set of quantum states of: 1)
energy, 2) angular momentum magnitude, 3) angular momentum orientation, and 4)
orientation of intrinsic spin.
The order of increasing energy of the subshells can be constructed by going through
downward-leftward diagonals of the table above (also see the diagram at the top of the
page), going from the topmost diagonals to the bottom. The first (topmost) diagonal
goes through 1s; the second diagonal goes through 2s; the third goes through 2p and
3s; and so on. This explains the ordering of the blocks in the periodic table.
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Notation and the Filling of orbitals
You can think of an atom as a very bizarre house (like an inverted pyramid!) - with
the nucleus living on the ground floor, and then various rooms (orbitals) on the higher
floors occupied by the electrons. On the first floor there is only 1 room (the 1s orbital);
on the second floor there are 4 rooms (the 2s, 2p , 2p and 2p orbitals); on the third
x
y
z
floor there are 9 rooms (one 3s orbital, three 3p orbitals and five 3d orbitals); and so
on. But the rooms aren’t very big . . . Each orbital can only hold 2 electrons.
In the notation, a subshell is written in the form nxy, where n is the shell number (i.e
1, 2, 3..), x is the subshell label (i.e s, p, d, f..) and y is the number of electrons in
the subshell. An atom’s subshells are written in order of increasing energy – in other
words, the sequence in which they are filled. For instance, ground-state Lithium has
two electrons in the 1s subshell and one in the (higher-energy) 2 s subshell, so its
ground-state configuration is written 1 s 2 2 s 1. Phosphorus (atomic number 15), is as follows: 1 s 2 2 s 2 2 p 6 3 s 2 3 p 3.
For atoms with many electrons, this notation can become lengthy and so the noble
gas notation is used. It is often abbreviated by noting that the first few subshells are
identical to those of one or another noble gas. Phosphorus, for instance, differs from
neon (1 s 2 2 s 2 2 p 6) only by the presence of a third shell. Thus, the electron configuration of neon is pulled out, and phosphorus is written as follows: [Ne]3 s 2 3 p 3.
An even simpler version is simply to quote the number of electrons in each shell, e.g.
(again for phosphorus): 2-8-5.
Valence and core electrons
The electronic configuration of the outermost (valence) electrons is significant. These
electrons determine the chemical properties of the element. Electrons that occupy
lower energy quantum levels are called core electrons. That of oxygen is 1s22s22p4.
The core electrons of oxygen are those in 1s atomic orbital; the six electrons with n
= 2 are the valence electrons.
Exercise 1
Write the electronic configuration for the following elements. In brackets are their
Atomic numbers.
(a) Be (4), Na (11), Rb (37).
(b) B (5), N (7), P(15).
(c) Sc (21), Co (27).
(d) He (2), Ne(10), Ar(18).
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Solution
a) Be,4: 1s22s2; na, 11: 1s22s22p63s1; Rb, 37: [Kr]5s1
b) B, 5:1s22s22p1; n, 7: 1s22s22p3; P, 15: 1s22s22p63s23p3 ≡ [Ne]3s23p3
c) Sc, 21: [Ar]4s23d1, Co, 27: [Ar]4s23d7
d) He, 2: 1s2, ne, 10: 1s22s22p6, Ar, 18: 1s22s22p63s23p6
Relation to the structure of the periodic table
Electron configuration is intimately related to the structure of the periodic table. The
chemical properties of an atom are largely determined by the arrangement of the
electrons in its outermost «valence» shell (although other factors, such as atomic
radius, atomic mass, and increased accessibility of additional electronic states also