REDEEMER'S UNIVERSITY

REDEEMER’S UNIVERSITY

COLLEGE OF NATURAL SCIENCES

DEPARTMENT OF CHEMICAL SCIENCES

INORGANIC CHEMISTRY I

COURSE: CHE104

STUDY GUIDE

2016/17

SEMESTER 2

Lecturers

Professor G A Kolawole

Dr. T T OYETUNDE

Compiled by Prof. GA Kolawole

Professor G A Kolawole Contact: 07063178352

e-mail: [email protected] Office:Room 18, CNS

Dr. T Oyetunde Contact: 08057647836

e-mail: [email protected] Office:Room 3, CNS

1. Course Title: Inorganic Chemistry 1

2. Module code: CHE104

4. Number of Unit: 2 (30 lectures)

7. Type of Course: Compulsory

9. Purpose of the Course:

This course is designed to introduce students to the structure of the Periodic Table, periodicity of the properties

of elements in the Periodic Table based of the electronic configuration, arrangement and position of elements in

the table and periodicity as a predictive tool for the properties of elements. The course would also cover basic

principles of bonding and properties of covalent bonding, and structures of solids. At the end of the course

students would have been introduced to periodicity along Periods 2 and 3, Groups 1 and 2 and the first row

transition metals within the Periodic Table that will form a conceptual foundation to Inorganic Chemistry at

higher levels of their studies, be able to predict the nature of bonding between two elements based on their

positions in the Periodic Table and describe the nature of structures of different types of covalent and ionic

solids.

10. Course Outcomes:

By the end of the course students should be able to:

Describe the structure of the Periodic Table and relate the position of elements in the Periodic Table to

the electronic configuration of the elements.

Classify elements in the PT according to their electronic configuration

Describe variations of selected physical properties, like radius, ionisation energy, electronegativity,

etc, along a Period and down a Group, and account for any discontinuity.

Relate reactivity and properties of elements in Periods 2 and 3, Groups 1and 2, and the first row

transition metal to the electronic configurations of the elements. Account for the differences and similarities in the properties of the elements in the respective selected

families of elements to their positions in the Periodic Table.

Describe the formation of ionic compounds and simple covalent compounds and relate the preference

of either to the position of the elements forming the bonds in the Periodic Table and their respective

electron configuration.

Describe the structures of simple ionic and covalent compounds and conceptualise bonding as a

continuum of polarity.

Undertake basic experiments in qualitative analysis, with appropriate demonstrable skills, and present

their results following standard formats.

11. Prescribed text books

(A) S. S. Zumdahl & S. A. Zumdahl: Chemistry, An Atoms First Approach, Brooks/Cole, 2012.

(B) K. W. Whitten, R. E. Davis, M. L. Peck & G. G. Stanley: General Chemistry, 7th

Edition, Thomson,

Brooks/Cole, 2004.

Students are urged to consult other helpful books available in the library. The lecture notes are available to be

photocopied on request from Professor Kolawole. Interested students can request for the soft copy by sending

an e-mail to Prof. GA Kolawole. The lecture notes only compliment the textbooks. Each student should

endeavour to buy a copy of one of the prescribed textbooks. Come to class with the lecture notes and one of

the textbooks. You will also find a Periodic Table handy. Attendance will be taken at every lecture and

assessment.

12. Work Schedule

UNIT Coverage Background reading Completion

Date (Week

ending)

UNIT1: The structure

of the Periodic Table

1. The electronic configuration

of elements and the Periodic

Table

2. Classification of elements in

the Periodic Table

(A) Chap. 2: 2.6-2.11, pg. 76-90

(B) Chap.5:5.15-5.17, pg 205-218

(A) Chap. 2: 2.11

Weeks 1

UNIT 2: Periodicity 1. 1. Physical properties of elements:

atomic/ionic sizes, ionization

energy, electron affinity;

2. Metals, metalloids, non-metals

(A) Chapter 2: 2.12, pg. 91-96;

4.4, pg 154-157

(B) Chap 6 pg 230 - 264

(A) Chapter 2: 2.13, pg. 96-97

See lecture note

Week 2

UNIT 3: Chemical

periodicity:

Periods 2& 3

Graduation of chemical and

physical properties of elements of

Periods 2&3.

(A) Chapter 2: 2.13, pg. 92

See lecture note

Weeks 3& 4

Assessment 1 (45 minutes)

Unit 4:Chemical

periodicity: s-Block,

Groups 1 & 2

Gradation in the chemical and

physical properties of s-block

elements: Groups IA& IIA

(A) Chapter 2: 2.13, pg. 96-99

(B) Chap. 23: 23.1-23.6, pg915-

925

See lecture note

Weeks5 &6

UNIT 5. Chemical

periodicity:

Introduction to first

row transition metals

Gradation of physical properties

and review of their unique chemical

properties. Introduction to

Coordination Chemistry

(A) Chap. 21, 21.1-21.6, pg 928-

956

Weeks 7 &8

MID-SEMESTER ASSESSMENT (1½ hours)

UNIT 6: Chemical

bonding

1. Types of chemical

bonding

2. Electronegativity

3. Polarity of covalent bonds

4. Covalent bonding

5. Lewis structures

6. Valence Bond theory:

molecular shapes and

bonding

7. VSEPR theory

8. Hybridization

9. Ionic bonding & ionic

compounds

(A) Chap. 4:4.1, pg 145-149

(A) Chap. 2:4.2, pg 149- 150

(A) Chap. 2:4.3, pg 151-153; 4.6,

pg 161-162

(A) Chap 2:4.7, pg 162-167

(B) Chap 7: 7.3, pg 275-276

(A) Chap 2: 4.9-4.12, pg 168-182

(B) Chap 7: 7.4-7.11, pg 276-296

(A) Chap 5: 5.1, pg 193-205

(B) Chap 8, pg 302-347

(A) Chap 5:5.1, pg 193-205

(B) Chap 8, pg 302-347

(A) Chap 5:5.1, pg 205-217

(B) Chap 7: 7.2, pg 268-274

Weeks

9&10

UNIT 7: Solid state 1. Structure & bonding in

metals

2. Carbon and silicon

network

3. Ionic solids

(A) Chap 8: 8.4, pg 339-345

(B) Chap 13: 13.16, pg 509

(A) Chap 8: 8.5, pg 345-349

(B) Chap 13: 13.16, pg 517-518

(A) Chap 4, 4.5, pg 157-161;

Chap 8: 8.3, 8.7, pg 333-338;

356-359

(B) Chap 13: 13.14-13.16, pg

503-517

Weeks 11&

12

Assessment 3 (1 hour)

REVISION Week 13

13. Methods of Assessment to be used: (% weighting)

Activity %

Attendance at lectures 05

Other Assessments 20

Mid-semester assessment 15

Semester Examination 60

Total 100

NOTE

80% attendance is mandatory to qualify to write the final semester examinations,

for which a score of five is awarded. Any attendance less than 80% attracts zero

score.

The other assessments will be in form of Tutorial Quizzes such that a set of

questions will be provided to guide your reading and periodically students are

given quizzes on the questions during one of the lectures to ascertain that each

student has read the note and has worked through the problems.

You are also given reading assignments that would warrant the use of the

library, for which a short typed report is submitted every two weeks.

14. Plagiarism

Plagiarism involves copying printed work or idea developed by other people without acknowledging the source

of your information. This amounts to stealing the intellectual property of other people and is punishable.

It is easily detected when you copy other peoples‘ work verbatim (word for word) or even when you reframe it

without acknowledging the origin of the information you use in your work. You need to keep this in mind when

you are given an assignment to do that involves consulting books, scientific journals or even newspapers and

use of information contained in the printed lecture notes given to you. When detected, you can easily lose

critical marks due to you in an assignment or even face disciplinary action.

Like in all human endeavours, intellectual honesty is a virtue and it pays to cultivate the habit right now and

keep it up for the rest of your life.

To acknowledge the source of information, you give a reference number after the statement and at the end of the

work you provide a list of references corresponding to the numbered references (There are different ways of

doing this depending on your discipline, which you will learn as you progress in your studies).

References normally should include: names of authors (Initials and surname); title of article; name of journal,

book, newspaper, etc; volume of the journal, book, newspaper, etc; pages of the article in the source and, if a

book, and the publisher, the edition and year of publication.

If the information is from a website, quote the website and the date you download the information in addition to

the above.

The University has recently acquired a software, ‗Turn-it-in‘, for detecting plagiarism and was first deployed to

assess final year projects last year. It will be applied in all assignments that warrant literature work at all levels

this year. The Committee managing the software runs training courses for staff and students every Thursday at

the Virtual Library. Dr. Osho of Biological Sciences is the Chairman of the Committee and can be contacted by

staff and students to book a place in the weekly training sessions.

UNIT 1: The Structure of the Periodic Table

1.1. Electronic Configuration

Three fundamental particles are identified as constituents of the atom: electron, proton, and neutron. While

neutron and protons reside in the nucleus of the atom electrons are located in the orbital surrounding the

nucleus. You are advised to undertake a reading on quantum theory as it relates to atomic structure. This

concept is critical to the understanding of the behaviour of atoms of elements we are discussing in this course.

Electron configuration refers to the arrangement of electrons in the shell and orbitals surrounding the

nucleus.

The Periodic Table is constructed based on these configurations and the chemical and physical properties of

atoms are dictated by its electron configuration. Indeed if the electron configuration of an atom is known, one

could easily predict the position of the element in the Periodic Table and hence its chemical and physical

properties. It is therefore very important that any Chemistry student should understand this concept and always

carry a Periodic Table when studying Chemistry.

Let us now discuss some basic terms required in the description of the electron configuration.

1.1.1. Orbitals and quantum numbers

The description of the probable location of an electron around the nucleus of an atom

requires four quantum numbers. Each electron in an atom has its own four unique quantum

numbers that can be used to identify which electron we are talking about and to which atomic

orbital it belongs.

The principalquantum number identifies the main energy level known as the

shell(denoted asn).

The sub-shell (angular or azimuthal) quantum number identifies the sub-levels of

energy within the main energy level, known as sub-shell or orbitals (denoted as s, p,

d, f, etc).

Theorbital (magnetic) quantum number pins down the location of individual electrons

in orbitals (denoted as ml).

The spin quantum number provides the possible orientations of an electron in the

orbital (denoted as s).

All orbitals that have the same value of n are said to be in the same shell. Thus the

shell with n = 1 is called the first shell, the shell with n = 2 is the second shell and so

forth.

The various shells are often identified by letters beginning with K for the first shell:

N 1 2 3 4

Shell K L M N

The quantum numbers

Principal quantum numbers (shell), n, (p. q. n) roughly correspond to the n of the energy

levels in the Bohr atom. It has whole number values, n = 1, 2, 3, …. They correspond tothe

main energy levelsin an atom. As n increases electrons are generally farther from the

nucleus and have higher energy; n ranges from 1 to 7 in the ground state of the atom. At the

excited state it can range from 2 to (infinity).

Subsidiary orangular or azimuthalquantum number (sub-shell), l,defines the different

energy sub-shells/sub-levels/orbitals, within the main, or n, level and indicates the shapes of

different types of orbital. Only certain values of l are allowed, which depend on n: l = 0, 1, 2,

3, ∙∙∙∙ (n-1). The total number of possible sub-shells or orbitals in each level is equal to n.

The sub-shells or orbitals are given by the values of l, which can assume values 0, 1, 2, 3, 4,

corresponding to s, p, d, and f orbitals respectively.

The energies of the orbitals increase from s to f, i.e.,s<p<d<f.

Assignment 1: Draw the shapes of s, p, d, and f orbitals. What is the degeneracy of each of

these orbitals?

Magnetic or orbital quantum number, ml defines the regions in space that can be occupied

by an electron, governed by ml.The number of allowed ml values depends on l, ml = -l to +l.

The number of orbitals in each sub-shell is equal to the number of values of ml, which is

equal to 2l + 1. Thus for s orbital,ml= 1; p orbitals, ml = 3; d orbitals, ml = 5 and for f

orbitals, ml = 7.

This results in one s orbital, three p orbitals, five d orbitals and seven f orbitals.

Spin quantum numbers, ms(or s)

An electron within an orbit rotates along an orbit but also, like a magnet, spins about its

axis.The rotation and spinning of electrons around the nucleus can be likened to the motion of

the sun around the earth. The spinning causes each electron to behave like a tiny magnet.

Electron spin therefore has two possible orientations corresponding to two possible values of

ms, +½, -½.

For each main level, only a specific number of atomic orbitals are allowed; thus for n = 1, l

has only one possible value, l = 0 and ml also has only one possible value, ml = 0. Therefore

at the n = 1 level, only one single atomic orbital is present, an s orbital.

Let us summarize:

Table 1a

For n

n = l = orbital =

1 0 s

2 0, 1 s, p

3 0, 1, 2 s, p, d

4 0, 1, 2, 3 s, p, d, f

Table 1b

For l

l = l ml = -l, (-l +1),…0, …, l-2, l-1, +l

0 (s) 0 one s orbital

1 (p) -1, 0, +1 three p orbitals

2 (d) -2, -1, 0, +1, +2 five d orbitals

3 (f) -3, -2, -1, 0, +1, +2, +3 seven f orbitals

An atomic orbital is designated as nl, where n is the principal quantum number and l is sub-

shell quantum number, expressed as s, p, d, or f.

The second table shows how the sub-shells determine the number of available orbitals.

Let us summarize what we have learnt on quantum numbers in another way:

Table 1c

1. n = 1 2 3 4….

K L M N….

2. l = 0 1 2 3 4 5…

Letter symbols/sub-shell/orbitals s p d f g h…

3. Sub-shells: s p d f g…

No. of mlvalues (sub-orbitals)1 3 5 7 9…

Can you recognize the arithmetic progression?

Table 1d: Summary of the relationship among the three quantum numbers

Value of

n

Value of l Values of ml Sub-shell/orbital Number of

orbitals

1 0 0 1s 1

2 0

1

0

-1, 0, +1

2s

2p

1

3

3 0

1

2

0

-1, 0, +1

-2, -1, 0, +1, +2

3s

3p

3d

1

3

5

4 0

1

2

3

0

-1, 0, +1

-2, -1, 0, +1, +2

-3, -2, -1, 0, +1, +2, +3

4s

4p

4d

4f

1

3

5

7

The spin quantum number will be useful when we consider the arrangement of electron in the

orbitals. For each ml value, there are two values of s, ½.

How many electrons can occupy a shell?

An orbital cannot accommodate more than two electrons. Where two electrons are in the

same orbital their spin quantum number must be different, i.e. they would have opposite

spins.

From the four quantum numbers it is possible to predict the number of electrons in an orbital

and therefore the number of electrons in a shell.

Number of electrons in an orbital = 4l +2, where l is the subsidiary quantum number.

Number of electrons in a shell = 2n2, where n is the principal quantum number.

These are summarized in the Table 1e below.

From the table note that two electrons can have the same n, l and mlbut the msmust be

different, i.e. no two electrons in the same atom can have all the four quantum numbers

equal (Pauli Exclusion Principle).

Table 1 e

n l ml ms Electrons in

an orbital = 4l +2

Electron in a

shell = 2n2

1 0 (1s) 0 2 2

2 0 (2s)

1 (2p)

0

-1, 0, +1

+½, -½

+½, -½

2

6

8

3 0 (3s)

1 (3p)

2 (3d)

0

-1, 0, +1

-2,-1, 0, +1, +2

+½, -½

½ for

each ml

2

6

10

18

4 0 (4s)

1 (4p)

2 (4d)

3 (4f)

0

-1, 0, +1

-2,-1,0,+1,+2

-3,2,1,0,+1,+2,+3

+½, -½

½ for

each ml

2

6

10

14

32

1.2: Electron configuration

Some important rules in writing electron configuration

The Aufbau principle: Electronsoccupy the lowest-energy orbitals available to them; they

enter higher energy orbitals only when the lowest energy orbitals are filled. This is also

known as the lowest energy principle. The sequence of energies for atomic orbitals is given

in Figure 2.

Hunds rule: When electrons are placed in a set of orbitals of equal energy (i. e., degenerate

orbitals), all the orbitals must be occupied singly first before any pairing can occur.

Pauli exclusion theory: No two electrons in the same atom can have the same set of four

quantum numbers.

In instances where two electrons occupy the same orbital (n, l and ml same) their spin

quantum number must be different.

For example: Helium, He, has two electrons.

n = 1, l = 0, ml = 0, but s = +½ or -½.

Thus two electrons are in the s orbital but with opposite spins.

i.e. 1s2 or

No orbital (as specified by n, l, ml) can ever contain more than two electrons.

For n = 2, l = 0, 1

For l = 0, ml= 0, ms = ½

2s orbital; configuration is 2s2 2 electrons.

For l = 1, ml = -1, 0, +1

p orbitals

For ml = -1, s = ½, 2 electrons

ml = 0, s = ½, 2 electrons

ml =+1, s = ½, 2 electrons

6 electrons

2s 2p

Figure 1: Relative energies of the orbitals

n = 7

n = 6

n = 5

n = 4

n = 3

n = 2

n = 1

1s<2s<2p<3s<3p<4s3d<4p<5s4d<5p<6s<4f<5d<6p<7s<5f<6d<7p

The configuration is 2p6 (Recall that there are three p orbitals, px, py, pz, so 2p

6 can be written

as 2px2 2py

2 2pz

2).

Thus for n = 2 there are 8 electrons: two in 2s, six in 2p. The three p orbitalsare degenerate.

The p orbital istherefore said to be triply degenerate.

The d orbital has five degenerate orbitals.

These orbitals are degenerate only at the ground state.

The maximum number of electrons for each principal quantum level and each orbital are

summarized in Table 2. The last column indicates that the maximum number of electrons in

each n level is 2n2.

Table 2: Summary of relationship among n, l, and ml

Value

of n

Value

of l

Values of ml Orbital Number

of sub-

orbitals

Electron

occupation

Maximum

electron in

a shell

1

2

3

4

0

0

1

0

1

2

0

1

2

3

0

0

+1, 0, -1

0

+1, 0, -1

+2, +1, 0, -1, -2

0

+1, 0, -1

+2, +1, 0, -1, -2

+3, +2, +1, 0, -1, -2, -3

1s

2s

2p

3s

3p

3d

4s

4p

4d

4f

1

1

3

1

3

5

1

3

5

7

2

2

6

2

6

10

2

6

10

14

2

8

18

32

The electron configuration of an atom is the distribution among the orbitals of all the

electrons in the atom. The following orbital notation is used: nlx, where n represents the

principal quantum level (as a number 1, 2, 3, ··); l represents the orbital as letters (s, p, d, f,

··) and x is the number of electrons.

The following orbital notations are used when writing the electron configuration of an atom:

nlx

e.g. 4d6

Example: 3d3 means ‗in the n = 3 level, the d orbitals are occupied by 3 electrons‘.

Examples:

H 1s1 (n = 1, l = 0, ml = 0, s = +½)

He 1s2 (n = 1, l = 0, ml = 0, s = +½ for 1st electron and n = 1, l = 0, ml = 0, s = -½ for 2nd

electron)

When p orbital is involved you must recall that the p orbital has three sub-orbitals (mlvalues

also known as degenerate levels). This can also be illustrated with box diagrams:

Orbital quantum number

No. of electrons in an

orbital

Principal

quantum number

are occupied by six electrons

the d orbitals In the n = 4

level/shell

B 1s22s

22p

1

C 1s22s

22p

2

N 1s22s

22p

3

1s 2s 2px 2py 2pz

The electron configurations of the 18 elements are easy to write but for elements of higher

atomic numbers, Figure 1(page 8) would be a useful guide in ordering the energy levels of

the orbitals.

Summary of rules for writing electron configuration

No two electrons can have the same four quantum numbers (Pauli Exclusion Principle)

Electrons occupy the lowest-energy orbitals available to them; they enter higher energy

orbtals only when the lower energy orbitals are filled.

Orbitals of equal energy (degenerate) are each occupied by single electron first before a

second electron, which will have the opposite spin quantum number, enters any of them.

Exceptions to the ideal configurations predicted by the energy sequence above are in Cr

(Z = 24), Cu (Z = 29) and Ag (Z = 47):

Cr (Z = 24), 1s2 2s

2 2p

6 3s

2 3p

63d

5 4s

1 instead of 1s

2 2s

2 2p

6 3s

2 3p

6 4s

24s

23d

4.

Cu (Z =29), 1s2 2s

2 2p

6 3s

2 3p

63d

10 4s

1 instead of …4s

23d

9.

Ag (Z = 47), 1s2 2s

2 2p

6 3s

2 3p

64s

23d

10 4p

6 4d

105s

1instead of …5s

2 4d

9.[Kr] 4d

10 5s

1 instead

of [Kr] 4d95s

2.

These configurations can be written as [Ar]3d5 4s

1, [Ar]3d

10 4s

1and [Kr] 4d

10

5s1respectively.

Note 1s2 2s

2 2p

6 3s

2 3p

6 is the configuration of element 18, Ar and 1s

2 2s

2 2p

6 3s

2 3p

64s

2 3d

10

4p6 is that of krypton, Kr.

This is attributed to the possible greater stability of half-filled or full-filled d obitals (e.g. 3d5

instead of 3d4 in chromium, because the half-filled 3d

5 is more stable than the 3d

4; and

3d10

/4d10

instead of 3d9/4d

10 in copper/silver because the fully-filled 3d

10/4d

10 is more stable

than 3d9/4d

9.

Generally half filled and fully filled configurations in any set of degenerate orbitals are

more stable than any other partially filled configuration.

1.3: Electron configurations and the Periodic Table

A modern Periodic Table is given in Figure 3 below.

The Periodic Table is built up by arrangement of elements in order of increasing

atomic number such that all elements that have similar chemical properties fall into

the same column and each row corresponds to the filling of electrons into an atomic

shell.

The columns are called Groups and the rows Periods. Elements within a group are

said to belong to the same Family.

The Periodic Table can be explained by the electron configuration of atoms of the

elements.

Each Period contains the number of elements that correspond to the number of

electrons contained in a shell.

The number of electrons accommodated by the sub-shells that become filled as the

atomic number increases determines the length of each row.

All the orbitals of the last element of each Period are fully filled by 2 electrons each.

The Group number corresponds to the number of electrons in the outermost shell of

an atom.

All elements in the same group have similar electron configurations in their outermost

shell, i.e. their outer electron configurations are similar.

Each Period corresponds to a principal quantum level: Period 1 (n = 1); Period 2 (n =

2); Period 3 (n = 3); Period 4 (n = 4), etc.

The outermost shell is the shell corresponding to the energy level of highest principal

quantum number.

Example: Consider Li (Z = 3) and Na (Z = 11)

Li: 1s22s

1, the higher principal quantum number in Li = 2

Na: 1s2 2s

2 2p

63s

1, the highest principal quantum number in Na = 3.

Both have 1 electron in the outermost s orbital, hence they are in group 1; however, Li is in

Period 2 whereas Na is in Period 3. Study the Table below:

n Period No. of electrons No. of elements

1 1 2 2

2 2 8 8

3 3 18 8 (18*)

4 4 32 18 (32**)

5 5 50 18 (50**)

6 6 72 32 (72**)

7 7 98 32 (98***)

*The apparent fewer elements than the periods should accommodate follow from the relative

energies of the orbitals given earlier (See Fig. 1).

**The present Periodic Table cannot accommodate more than 32 elements in a period.

***There are prospects of discovering more elements!

The modern Periodic Table is given in Figure 2

1.4: Classification of Elements in the Periodic Table

The chemical and physical properties of an element are governed by the number and

arrangement of the orbital electrons, i.e. by their atomic number. The atomic number

determines the position in the Periodic Table, where elements are arranged in order of

increasing atomic number. Such arrangement leads to certain elements falling in a particular

column, the Group, and some horizontal arrangement or rows, the Period.

Each Group shows similarities in chemical properties and gradation in physical properties.

Thus it is only necessary to learn the general properties of each group, and the trends as the

group is descended, rather than the properties of each individual element in the group. The

group to which an element belongs is determined by the number of electrons in its highest

energy electron shell, i.e. its outermost electron shell.

Figure 2: The Modern Periodic Table

Groups

s-block

Group I – one s-electron in the outermost sub-shell.

Group II – two s-electrons in the outermost sub-shell.

p-block

Group III – two s-electrons + one p-electron in the outermost shell = 3 electrons.

Group IV – two s-electrons + two p-electron in the outermost shell = 4 electrons.

Group V – two s-electrons + three p-electron in the outermost shell = 5 electrons

Group VI – two s-electrons + four p-electron in the outermost shell = 6 electrons

Group VII – two s-electrons + five p-electron in the outermost shell = 7 electrons

Group VIII – two s-electrons + six p-electron in the outermost shell = 8 electrons

Group VIII or group 0 has a full outer electron shell so that the next shell is empty, hence the

name Group 0.

Groups III to VIII all have their p orbitalsbeing filled with electrons and because their

properties depend on the presence of p-electrons, they are jointly called the p-block elements.

The Periodic Table Contains 16 groups (Group VIIIB contains three subgroups), making it 18

groups effectively. Of these, 8 are main groups (long) and 10 are subgroups (short).

Periods

The elements in the Periodic Table are arranged in order of increasing atomic number, i.e. in

order of increasing nuclear charge, so that each element contains one more orbital electron

than the preceding element.

Each row/period begins with an alkali metal and ends with a noble gas.

The sequence in which the various levels are filled determines the number of elements in

each period, and the Periodic Table can be divided into four main regions according to

whether s, p, d, or f levels are being filled.

There are seven periods corresponding to the filling of:

Period 1: 1s2 2 elements

Period 2: 2s22p

6 8 elements

Period 3: 3s23p

6 8 elements

Period 4: 4s23d

104p

6 18 elements

Period 5: 5s24d

105p

6 18 elements

Period 6: 6s24f

145d

106p

6 32 elements

Period 7: 7s25f

146d

107p

6 32 elements

The periods of the Periodic Table can therefore be broken down into:

3 short periods – corresponding to principal quantum numbers 1, 2, 3.

2 long periods – corresponding to principal quantum numbers 4, 5.

2 very long periods – corresponding to principal quantum numbers 6, 7.

The most recent Periodic Table reveals that 117 elements are now discovered, leaving one

more possible elements to fill the present Periodic Table (i.e. to completely fill Period 7).

Thus there is the prospect of discovering more new elements in the future.

Classifications of elements

Elements are classified according to:

(a) Electron configuration

s-block – Groups IA (ns1) and IIA (ns

2); maximum 2 electrons in s-sub-shell.

p-block – Groups IIIB – VIIIB (ns2np

x, x = 1-6); maximum 6 electrons in psub-shell.

d-block – Groups IIIA – VIIIA, IB – IIB (ns2(n-1)d

x, x = 1-10); maximum 10

electrons in the dsub-shell.

f-block – Lanthanides and actinides [ns2(n-1)d

0-1(n-2)f

x, x = 1-14]; maximum 14

electrons in the f-sub-shell.

(b) Metals – s-block, d-block, f-block and bottom of p-block.

Non-metals – Top of p-block.

Metalloids – located between metals and non-metals.

(c) Trivial class names

s-block – alkali metals (Group IA) and alkaline earth metals (Group IIA).

p-block – pnicogens, rarely used (Group V), chalcogens (Group VI), halogens (Group

VII), noble gases (Group VIII or 0).

d-block – transition metals.

f-block – Inner transition metals (lanthanides and actinides).

The Periodic Table puts hydrogen in two groups, Group IA and Group VIIB. Some put it in

no group! Hydrogen resembles Group IA in the sense that it forms a positive ion, H+,

carrying a single charge like Group IA metals. It also forms a negative ion, H-, like Group

VIIB elements.

The Periodic Table can also be classified into four blocks (Figure 3):

s-block: Groups where the outermost orbital is s. They are to the left of the Periodic Table.

Since the s sub-shell can accommodate a maximum of 2 electrons, there are only two groups

for s block, Groups I and II.

p-block: Here outermost electrons are located in the p orbitals. They are located to the right

of the Periodic Table. The p sub-shell can accommodate a maximum of 6 electrons (2 in

each of px, py, and pz). There are six groups in the p-block, Groups III to VIII.

d-block: This block is located between the s and the p blocks, and corresponds to electrons

being fed into the d sub-shell, which can accommodate 10 electrons; hence there are 10

elements in each row of the d block. There are three rows altogether hence there are 30 d

block elements. The d block elements are called Transition Elements, from their location

between s and p orbitals.

f-block: They correspond to the filling of the f sub-shell, which can accommodate 14

electrons, hence 14 elements of each of the two rows in f-block. There are 28 elements in the

f-block

The classification of elements into blocks is presented in Figure 3.

Figure 3: Classification of elements of the Periodic Table

s-block

d-block

f-block

p-block

Representative elements areelements for which all inner sub-shells are fully filled and the

outer s and p sub-shells are filling are called representative elements or main group

elements. They correspond to the first 16 elements, 18 inclusive of the Noble gases (Group

VIII), in the Periodic Table.

Periodicity and Oxidation states

The maximum oxidation number attainable for each Group corresponds to the Group number.

However, while for groups 1-5 these oxidation states are positive and stable at the top of each

group, for Groups 6-7, the oxidation state corresponding to the group number is oxidising and

therefore unstable. Thus the most stable oxidation state corresponds to Group Number-8, i.e.,

for Groups 6 and 7, the -2 and -1 oxidation states are the most stable.

UNIT 2: Chemical Periodicity

The chemical properties of elements are related to their electronic configuration and therefore

to their position in the Periodic Table. The variation of physical and chemical properties

within the Periodic Table follows some regular patterns. This is covered in the periodic law:

The properties of the elements vary periodically with their atomic numbers.

2.1: Metals, non-metals and metalloids

As mentioned earlier the elements in the Periodic Table can be classified into metals

(over 80 elements), metalloids (8 elements) and non-metals (17 elements). Each class

has its unique properties that will be discussed when we treat chemistry of the

elements.

Within a Period, the metals are over 80% of all elements located on the left side of the

Periodic Table; the metalloids are between metals and non-metals and non-metals are

on the right side of the Periodic Table.

One characteristic of all metals is that they have electrons that can be easily lost,

usually from s, p, d (for transition metals) and f (for lanthanides). This makes metals

to be conductors. Metalloids are semiconductors and non-metals are insulators.

As a group is descended, the elements become more and more metallic; thus at the top

of a group we find non-metal/less metallic elements followed by metalloids and

finally by metals. Even in the groups containing predominantly non-metals they do

contain elements with metallic properties at the bottom of the groups.

Figure 4: Some of the elements to which we will make references in this course

III IV V VI VII

I

IIH H

Li

Na

K

Rb

Cs

Fr

Be

Mg

Ca

Sr

Ba

Ra

B C N O F

Al Si P S Cl Ar

Sc Ti V Cr Mn Fe Co Ni

VIII

Ne

Cu Zn

He

Ga Ge As Se Br Kr

Xe

Rn

I

At

Period 2

Period 3

Period 4

Period 5

Period 6

Period 7

Period 1

Te

Metalloids are written in red

2.2: Atomic and ionic radii

The radius of the atom is larger than that of the cation whereas it is smaller than that

of the anion. Cations are formed by metals by loss of electrons whereas anions are

formed from non-metals by gain of electrons.

Generally cations are smaller than anions.

Metals have larger sizes than metalloids and metalloids larger than non-metals. Thus:

Atomic sizes decrease from left to right of a Period and increases from top to bottom

of a Group.

o In the p-block elements, the second and third members are similar in size and

properties because of the intervention of the d-transition metals.

o In the d-transition metals the two heaviest members are similar in size and properties

because of the intervention of the lanthanides (lanthanide contraction).

o The first element in each representative group differs from the other elements in the

family because of the smaller size of its atoms.

On the other hand ionic sizes increase from left to right of a Period and also increase

down a Group subject to the exceptions indicated above and provided the charges are the

same.

Factors that influence radii

The radius of an atom or an ion is mainly influenced by effective nuclear charge and the

number of energy levels (shells) occupied by electrons in the atom. The bonding

environment of the atom also has some effect.

The effective nuclear charge is the portion of the nuclear charge that acts on a given

electron. As nuclear charge increases (i.e. number of protons increases), electrons are

attracted more strongly to the nucleus.

The screening effect is the decrease in the nuclear charge acting on an electron due to the

effects of other electrons. Screening effect varies depending on the location of the

electrons that screen an outermost electron. Electrons closest to the nucleus screen

strongest and decreases as the closed shells approach the outermost shell. Within the

outermost shell, the type of sub-shell also affects screening: The order is s>p>d>f

because of the shapes of these orbitals.

Ions that have the same electronic configuration are said to be iso-electronic. Examples

are N3-

, O2-

, F-, Na

+, Mg

2+ and Al

3+. N

3-, O

2-and F

-gain electrons to attain the neon

configuration whereas Na+, Mg

2+ and Al

3+ lose their valence electrons to attain neon

configuration.

Isoelectronic ions N3-

O2-

F-Na

+ Mg

2+ Al

3+

Z (Nuclear charge) 7 8 9 11 12 13

Radii (nm) 0.171 0.1260.119 0.116 0.085 0.068

No. of electrons 10 10 10 10 10 10

The decreases in ionic radii as the ions become heavier are solely due to greater effective

nuclear attraction by the highly charged nucleus.

2.3: Ionization energy (IE)

The ionization energy (also called ionization potential) is the enthalpy change for the

removal of the most loosely bound electron from an atom or an ion in the gaseous state.

Ionization energies are given per mole of atoms or ions of a given type.

The first ionization energy is the energy required to remove one electron from a

neutral atom. E.g. Na(g) Na+

(g) + e- H

o = 495.8 kJ mol

-1 = the first IE

Such reactions are always endothermic and IEs are always positive. Energy is always

required to pull an electron from the attraction of the nucleus.Energies for the removal of

additional electrons are the second, third IEs, and so on.

Removal of subsequent electrons requires larger energies:

Al(g) Al+

(g) + e- H

o = 578 kJ mol

-1 (1

st IE)

Al+

(g) Al2+

(g) + e Ho = 1817 kJ mol

-1(2

nd IE)

Al2+

(g) Al3+

(g) + e Ho = 2743 kJ mol

-1 (3

rd IE)

The effective nuclear charge increases as successive electrons are removed hence increase

in IEs from first to second to third.

IEs are, like radii, influenced by the effective nuclear charge and the electron

configuration. Increase in effective nuclear charge leads to increase in IE.

Generally IE is smaller for larger atoms/ions and thus IEs are smaller in metals than

metalloids and non-metals; i.e. decreases from left to right of a Period and from top to

bottom of a Group.

Stable electron configurations like are found in the Noble gases require larger

energies to ionize; hence the IEs of the Noble gases are generally the largest in a

Period. Similarly all ions that attain the noble gas configurations will have large IEs.

Generally too, half-filled and fully filled orbital configurations are relatively stable

and will require higher IEs if electrons are to be removed from such configurations.

Figure 5: First ionization energies

III IV V VI VII

I

II

H

H

Li

Na

K

Rb

Cs

Fr

Be

Mg

Ca

Sr

Ba

Ra

B C N O F

Al Si P S Cl

Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br

I

At

He

Ne

Ar

Period 2

Period 3

Period 4

Period 5

Period 6

Period 7

1312

520

496

419

403

376

899 801 1086 1402 1681

2372

1314 2081

738 578 786 1012 1000 1251 1521

590

550

503

509

2.4: Electron affinity (EA)

Electron affinity is the enthalpy change for the addition of one electron to an atom or ion in

the gaseous state. EAs are also given per mole of atoms or ions.

Example:Cl(g) + e- Cl

-(g) H

o = -349 kJ mol

-1

The addition of an electron to a chlorine atom to attain a noble gas configuration occurs

readily and is exothermic. Unlike IE, electron affinities can be positive or negative. Energy

is required to add electron to a stable configuration. EAs for Group II elements (ns2

configuration) and Group VIII elements (ns2np

6) are positive (endothermic). All second EAs

are positive because electrons are being added to an already negative species because energy

is required to overcome the repulsion between electron and an already negatively charged

ion.

Energy is usually released when an electron is added to an isolated neutral atom. The energy

which is associated with the change

X(g) + e X-(g)

is called the electron affinity (EA). In IE, energy is added to remove an electron. Energy

will be released if an electron is added to a neutral atom. Addition of electron to an atom is

therefore an exothermic process. Some EA values are given below:

Element EA (kJmol-1

)* Process

Fluorine -344 F(g) + e F-(g)

Chlorine -349 Cl(g) + e Cl-(g)

Bromine -325 Br(g) + e Br-(g)

Oxygen -142

+844 O(g) + e O

-(g)

O(g) + e O2-

(g)

Hydrogen -72 H(g) + e H-(g)

Sodium -50 Na(g) + e Na-(g)

* A negative sign indicates that energy is evolved when an electron is added.

2.5: Electronegativity

Electronegativity is the ability of an atom, in a covalent bond, to attract electrons to it. The

electronegativity values can be used to predict whether a covalent bond is polarized or not

(See Unit 5). Within a Periodic Table,

Electronegativity increases from left to right of a Period.

Electronegativity decreases from top to bottom of a Group.

The most electronegative element is F; at the topmost right hand corner of the

Periodic Table (excluding the noble gases).

The least electronegative element is at the bottom of the left hand corner of the

Periodic Table, Fr.

An electronegative atom tends to acquire a partial negative charge in a covalent bond

or to form negative ion. Non-metals are generally electronegative.

An electropositive atom tends to acquire a partial positive charge in a covalent bond

or form positive ions. Metals are generally electropositive.

In compounds formed by two non-metals of unequal electronegativity, the less

electronegative atom carries the partial positive charge and the more electronegative

carries the partial negative charge; e.g. in ICl, I is less electronegative and therefore

carries a partial positive charge, Cl carries partial negative charge.

The polarity of a bond increases with increase in the difference between the

electronegativities of the two atoms that bond.

III IV V VI VII

I

II

HH

Li

Na

K

Rb

Cs

Fr

Be

Mg

Ca

Sr

Ba

Ra

B C N O F

Al Si P S Cl

Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br

I

At

2.1

1.0

0.9

0.8

0.8

0.7

0.7

1.5

1.2

1.0

0.9

0.9

2.0 2.5 3.0 3.5 4.0

1.5 1.8 2.1 2.5 3.0

Least electronegative

Most electronegative

1.0

UNIT 3: Chemical Periodicity

3.1 Period 1

3.1.1 Physical Properties

Element Li Be B C N O F Ne

Atomic Number 3 4 5 6 7 8 9 10

Electron configuration [He] 2s1 2s

2 2s

22p

1 2s

22p

2 2s

22p

3 2s

22p

4 2s

22p

5 2s

22p

6

Atomic radii/pm 157 112 79corr 71* 73

** 74

*** 71 -

First I.E./kJmol-1

526 906 807 1093 1407 1320 1687 -

Electron Affinity/kJmol-1

-57 66 -15 -121 31 -142 -333 +99

* Radius for graphite (77 pm for diamond).

** Based on covalent bond length of N-N in N2H4.

*** Calculated from O-O covalent bond length in H2O2.

General comments

As we move from Li Ne the second shell is being filled with electrons.

Each addition of electrons increases the number of protons in the nucleus by one; hence

the atoms are neutral.

Because the electrons go into the same shell (n=2) their addition does not have much

influence on screening effect, i.e. screening effect is almost constant.

Effective nuclear charge on the outer electrons therefore increases from Li F.

Atomic sizes therefore decreases.

Although the first ionisation energy is expected to increase as the effective nuclear

charge on the electrons increases, there are two discontinuities ("saw-tooth") at boron

and oxygen.

The values of B, C and N are lower than expected (Plot the data in the table above and

extrapolate the lines). This is due to the fact that the p electrons being removed in these

elements are less penetrating than the s electrons that are removed from Li and Be.

The p electrons are thus more effectively shielded from the nucleus and are hence easily

ionised.

The discontinuity in O, F and Ne is due to the removal of the first electron from the

paired p electrons. There is some repulsion before pairing takes place (a "spin-pairing

energy required) and thus the paired electron is less tightly bound. The I.E for O, F, and Ne

are again lower than the values expected by extrapolation from B, C and N values. The N

value is particularly high because the electron being removed is from a relatively stable half-

filled p-orbital.

The electron affinity increases steadily indicating the increasing tendency of the

elements to accept electrons. It also conforms with the change from metal to non-metal.

The positive values in Be and N illustrate the extra stability enjoyed by filled s

orbital and half-filled p orbitals respectively.

The electronegativities are generally high, increasing across the period. Compounds of

the elements therefore tend to be covalent. The covalent character of bonding is most

pronounced at the middle of the period.

3.1.2 Chemical periodicity of Period II elements

The elements can be divided into three groups:

Metal - Li; Metalloid - Be, B; Non-metals - C, N, O, F.

Classification can also be made using other criteria:

(a) electrical conductivity

(b) behaviour of oxides.

Using (a) Li and Be are metals; B a metalloid (because its conductivity increases with

increase in temperature) and the rest are non-metals.

Using (b) Li is a metal (Li2O is basic), Be and B are metalloids (BeO and B2O3 are

amphoteric) and the rest are non-metals (their oxides are acidic e.g. CO2, NO2) except

that the oxide of O (O2) is neutral. F does not have an oxide but a fluoride of oxygen

OF2! (Why?)

2.1.3. Compounds

The Oxides

Generally oxides can be basic, acidic, amphoteric or neutral, depending on their

behaviour with water, aqueous acids and alkalis.

Within a period oxides change from strongly basic (for pure metal) and strongly acidic

(for non-metals). In passing through a period the transition is through weakly basic,

amphoteric, and weekly acidic and this is gradual.

Oxide Normal state Type Product on reacting with

water (hydrolysis)

Li2O solid basic LiOH

BeO solid amphoteric-basic Be(OH)2

B2O3 solid amphoteric B(OH)3

CO2 gas weakly acid H2CO3

CO gas neutral -

N2O5 gas strongly acidic HNO3

NO2(N2O4) gas acidic HNO3 + HNO2

NO gas neutral -

O2 gas - -

OF2* gas acidic HF

* not an oxide but a fluoride of oxygen.

Reactions of oxides with water

Li2O + H2O LiOH

BeO + H2O Be(OH)2

B2O3 + 3H2O 2B(OH)3 (2H3BO3)

CO2 + H2O H2CO3

N2O5 + H2O 2HNO3 , etc.

Reactions with alkalis

Only the amphoteric and acidic oxides react.

BeO + OH- + H2O [Be(OH)4]

2- beryllate

B2O3 + 2OH- + 3H2O 2[B(OH)4]

- borate

N2O5 + 2OH- → 2NO3

- + H2O

3.1.4 The Hydroxides

The metal forms basic hydroxide, which can be easily decomposed by heat,

2LiOH Li2O + H2O

Be(OH)2 is amphoteric and can be decomposed by heat too,

Be(OH)2 BeO + H2O

The hydroxides of B, C and N are acidic.

These non-metals form oxy-acids as their 'hydroxides'. The acids are often wrongly

written as H3BO3, H2CO3, HNO3 instead of (OH)3B, (OH)2CO, (OH)NO2 respectively.

In all these the hydrogen is attached to the O not the elements.

Chlorides

LiCl BeCl2 BCl3 CCl4 NCl3 Cl2O ClF

solid solid solid liquid liquid gas gas

ionic ionic polar cov cov cov cov cov

+water soluble soluble sparingly

soluble

insol. insol. reacts reacts

pH of soln. acidic acidic weakly

acidic

- - acidic acidic

Changes from solidliquidgas correspond to changes in bonding (ionic covalent).

Ionic chlorides are soluble in water giving acidic solutions. This is due to the hydrolysis

of the small metal ions.

Cl2O and ClF react to give acidic solutions

Cl2O + H2O 2HOCl

ClF + H2O HF + HOCl

2.2. Period 3

3.2.1 General properties

Element Na Mg Al Si P S Cl Ar

Atomic Number 11 12 13 14 15 16 17 18

Electron configuration [Ne] 3s1 3s

2 3s

23p

1 3s

23p

2 3s

23p

3 3s

23p

4 3s

23p

5 3s

23p

6

Atomic radii/pm 186 160 143 117 110 104 99 191

First I.E./kJmol-1

496 737 577 786 1012 999 1256 1521

Trends in physical properties and explanations for them are similar to Period II.

3.2.2 Classification

Metals: Na Mg Al (Al is, however, amphoteric)

Metalloid: Si

Non-metal: P S Cl

The reactivity of the elements falls under the basic reactions expected for metals, metalloids

and non-metals. The following reactions are, however, important:

3.2.3 Reactions with alkali

Al + OH- + 3H2O [Al(OH)4]

- + 1½H2

Aluminate(III) ion

Si + 2OH- + H2O SiO3

2- + 2H2

Silicate(IV) ion

P4 + 3OH- + 3H2O PH3 + 3H2PO2

-

dihydrogenphosphate(I) ion

3S + 6OH- 2S

2- + SO3

2- + 3H2O

sulphate(IV) ion

Cl2 + 2OH- Cl

- + ClO

- + H2O

chlorate(I) ion Note that in all the cases, one of the products is an oxo-anion, or a hydroxyl-anion, containing

the element. The last two reactions are examples of disproportionation.

3.2.4 Compounds

3.2.4.1 Hydride

NaH MgH2 (AlH3)x SiH4 PH3 H2S HCl

ionic ionic cov cov cov cov cov

basic weakly

basic

polymeric

(amph)

? neutral weakly

acidic

acidic

solid solid gas gas gas gas gas

+water reacts*

(soluble)

reacts reacts reacts insol. slightly

soluble

very

soluble * NaH + H2O NaOH + H2

3.2.4.2 Oxides

Na2O MgO Al2O3 SiO2 P4O10 SO3 Cl2O7

strongly

basic

basic amph. weakly

acidic

acidic strongly

acidic

strongly

acidic

+water soluble slightly

soluble

insoluble insoluble soluble soluble soluble

+acid soluble

(reacts)

soluble

(reacts)

soluble

(reacts)

slightly

soluble

no

reaction

no

reaction

no

reaction

+alkali - - soluble

(reacts)

slightly

soluble

soluble

(reacts)

soluble

(reacts)

soluble

(reacts)

ionic ionic cov cov cov cov cov

solid solid solid solid solid gas gas

Note:

Basic oxides react with acid but not with alkali; acidic oxides react with alkali but not

with acid; amphoteric oxides react with both acid and alkali.

There are other oxides for P, S and Cl in addition to the typical ones given above.

3.2.4.3 Hydroxides

(a) Metals

NaOH Mg(OH)2 Al(OH)3

alkaline basic amphoteric

+heat not decomposed decomposed

decomposed

Mg(OH)2 MgO + H2O

2Al(HO)3 Al2O3 + 3H2O

Note:

Basic strength decreases across the period.

Stability of hydroxides to heat decreases.

(b) Non-metals

The 'hydroxides' of non-metals are oxy-acids: the hydrogen is attached to the oxygen not

to the central element, i.e.

(HO)4Si not H4SiO4

(HO)3P(O) not H3PO4

(HO)2S(O)2 not H2SO4

(HO)Cl(O3) not HClO4

(HO)4Si (HO)3PO (HO)2SO2 HOClO3

silicic acid phosphoric acid sulphuric acid perchloric acid

It is impossible to rationalise from classical analytical composition the formulae H4SiO4,

H3PO4, H2SO4 and HClO4. The structures are

Si Cl, , ,

HO OH

OHHO

P O

HO

HO

OH

S

O

O

HO

HO

O

O

O

HO

On reacting with an alkali they form:

SiO44-

PO43-

SO42-

ClO4-

silicate(IV) phosphate(V) sulphate(VI) chlorate(VII)

ion ion ion ion

Note:

Basicity of the 'hydroxides' decreases from 4 to 1 from Group IV to Group VII.

There are other oxy-acids for P, S and Cl.

The acid strength of the acids increases from left to right.

The above oxo-anions can be represented as XO4n-

where n = 8- group number.

The order of acid strength of the oxy-acids is related to the number of double bonded O

to the central atom. Oxygen, being more electronegative withdraws electrons from the

central atom which is then relayed to the oxygen in O-H and consequently to the H atom.

For example:

H O

S

H O

O

+ 2H2O 2H3O+ +

O

S

O

OO

O

3.2.4.4 Chlorides

NaCl MgCl2 Al2Cl6 SiCl4 PCl5 S2Cl2 Cl2

M.p. 808 714 192(p) -68 148(p) -80 -101

B.p. 1465 1418 180(s) 57 164 136 -35

Bond

type

ionic solid ionic solid polar cov

volatile

solid

polar cov

liquid

cov

liquid

cov

liquid

cov

gas

Structure crystalline layer dimeric molecular

+water soluble soluble

(hydrolyse)

soluble

(hydrolyse)

hydrolyse hydrolyse hydrolyse reacts

Trends in properties of the Chlorides

Bond: Changes from ionic to covalent.

M.p/b.p.:

Both depend on the bonding.

The crystal lattice in the involatile ionic chlorides cannot be easily broken.

The intermolecular forces in the covalent chlorides become increasingly weak as the

chlorides become more covalent (i.e. less polar).

The weak van der Waals forces become important at the extreme right where the

chlorides are most covalent.

The sudden change in m.p. between AlCl3 and SiCl4 corresponds to a dramatic change

in structure.

Interaction with water:

Ionic chlorides dissolve in water to give hydrated metal ions and chloride ions.

The metal ions, depending on size, can undergo hydrolysis, giving acidic solutions

(cations hydrolyse to give acidic solutions).

NaCl + H2O Na+(aq) + Cl

-(aq)

Na+does not hydrolyse.

MgCl2 + H2O Mg2+

(aq) + 2Cl-(aq)

AlCl3 + H2O Al3+

(aq) + 3Cl-(aq)

Mg2+

and Al3+

ions undergo hydrolysis in water.

[Mg(H2O)x]2+

+ 2H2O [Mg(H2O)x-2(OH)2] + 2H3O+

[Mg (H2O)x-2(OH)2] Mg(OH)2 + (x-2) H2O sparingly soluble

[Al(H2O)x]3+

+ 3H2O [Al(H2O)x-3(OH)3] + 3H3O+

[Al(H2O)x-3(OH)3] Al(OH)3 + (x-3)H2O insoluble

Covalent chlorides dissolve in organic solvents, but they hydrolyse in water to give

hydrogen chloride and the oxide, oxyacid or oxychloride of the element involved.

PCl5 + H2O POCl3 + 2HCl

POCl3 + 2H2O (HO)3PO + 3HCl

(H3PO4)

SiCl4 + 4H2O (HO)4Si

(HO)4Si SiO2.2H2O

2S2Cl2 + 2H2O 3S + SO2 + 4HCl

Cl2 + H2O HCl + HOCl

UNIT 4: Chemical Periodicity: Groups I & II

s-block elements

These are elements, which have the outermost electrons in the s-orbital. The elements are in

Groups IA and IIA. Group IA elements are known as alkali metals while Group IIA

elements are known as alkaline earth metals.

4.1 GROUP IA: ALKALI METALS

Physical properties

Element Li Na K Rb Cs Fr*

Atomic Number 3 11 19 37 55 87

Electron configuration [He]2s1 [Ne]3s

1 [Ar]4s

1 [Kr]5s

1 [Xe]6s

1 [Rn]7s

1

1st I.E./kJmol

-1- 520 496 419 403 376 ~375

2nd

I.E./kJmol-1

7298 4563 3051 2632 2420 -

Electronegativity 1.0 0.9 0.8 0.8 0.7 -

Metal radii/pm 157 186 227 248 265 -

Ionic radii/pm 76 102 138 152 167 (180)

4.1.1 General comments

The metals are all very soft and can be cut very easily with a knife. Li is the hardest.

The alkali metals form a homogeneous group of extremely reactive elements which

illustrate the similarities and trends expected within a group.

The physical and chemical properties are dependent on their simple electronic

configuration, ns1 (i.e. all of them have one electron in the outermost shell).

The single electrons can be readily released; they are therefore

strong reducing agents,

extremely reactive, and

most of their compounds are ionic.

The loss of one electron leads to all the elements having an oxidation state of +1.

The metals have a high lustre when freshly cut but tarnish rapidly in air due to reaction

with O2 and moisture.

4.1.2 Trends in their properties

From the above table:

4.1.2a Ionization energy

They have the lowest first ionization energy within any period because of their large

sizes, which make the outermost electron only weekly held by the nucleus.

First ionization energy decreases as the group is descended.

This is due to the shielding effect of the closed shell preceding the outermost shell.

With addition of filled shell as the group is descended the outer electron gets farther

and farther away from the influence of the nuclear charge.

The effective nuclear charge therefore decreases.

The second ionization energy is exceedingly large because the electron being removed is

from a stable noble gas configuration

The formation of M2+

ions is therefore ruled out

4.1.2b Electronegativity

The electronegativity is generally low.

They have the lowest electronegativity within a period

They are therefore electropositive metals.

As a result of their electropositivity they react with non-metals( electronegative

elements) readily to form ionic compounds.

(For the formation of ionic compounds the electronegativity difference between the two

elements should be large, usually greater than 1.7 in most cases (See Unit 5).

4.1.2c Metal/ionic radius

They have the largest atomic sizes within a period.

The ionic sizes are smaller than the atomic sizes.

The loss of a single electron results in more protons in the nucleus than elctrons in the

orbitals.

The outermost shell has been removed.

The number of protons (i.e. positive charge) in the nucleus is greater than the number

of electrons. The effective nuclear charge on the remaining electrons is therefore

greater than in the atom.

The atomic and the ionic sizes increase down the group because of the addition of

extra shell as the group is descended.

4.1.3 Trends in the chemical properties of the metals

4.1.3a Reactions with air/O2

They all tarnish in air and therefore are stored under oil.

A series of reactions are possible in air: For example, Na

It reacts with O2 in air to form the oxide.

The oxide can absorb moisture to form NaOH.

Both the Na2O and the NaOH can absorb CO2 and get converted to NaHCO3.

The hydrogen carbonate can be decomposed by heat from the sun and get converted

to Na2CO3, which depending on humidity can be hydrated to Na2CO3. 10 H2O.

As the weather becomes dry the hydrated sodium carbonate loses water until it finally

becomes a white powder, Na2CO3. H2O:

Na →Na2O → NaOH → NaHCO3 → Na2CO3 → Na2CO3.10 H2O → Na2CO3.H2O

The process by which Na2CO3 .10 H2O is converted to Na2CO3 .H2O is known as

efflorescence (Not all the water is lost, hence it is not dehydration)

The product obtained when the metals react with pure oxygen depends on the metal:

Li forms normal oxide, 4Li + O2 2Li2O

Na forms both the normal oxide and peroxide,

4Na + O2 2Na2O normal oxide

2Na + O2 Na2O2peroxide

K, Rb and Cs form superoxide, in addition to the normal oxide and the peroxide,

4M + O2 2M2O

2M + O2 2M2O2

M + O2 MO2

All the oxides are ionic and therefore dissolve and ionize in water:

Reactions of the oxides with water,

M2O + H2O 2MOH

M2O2 + 2H2O MOH +H2O2

2MO2 + 2H2O 2MOH + H2O2 + O2

Ionisation of the oxides in water,

M2O 2M+ + O

2-, O has an oxidation number of -2

M2O2 2M+ +O2

2-, i.e. each O has an oxidation number of –1

MO2 M+ + O2

-, i.e. O has an oxidation number of -½

As the metals get larger the tendency to form the higher oxides increases; i.e, as the group

is descended, the tendency to form higher oxides increases.

4.1.3b Reaction with water

They all react readily with water to form their respective hydroxides and hydrogen.

Li is the least reactive and Cs is the most reactive, i.e. the reaction with water becomes

progressively more violent as the group is descended.

Li reacts smoothly.

Na darts round violently on the surface of water and forms a silvery ball.

K reacts very violently and the large heat of reaction causes the hydrogen gas

liberated to catch fire and explode.

Rb and Cs react even more violently.

One could predict that the radioactive Fr would be the most reactive.

The general equation for the reaction with water is

2M + H2O 2MOH + H2

All the hydroxides are very soluble in water.

4.1.4 Compounds

The oxides have been discussed above.

4.1.4a Hydroxides

NaOH and KOH are called caustic soda and caustic potash respectively, because of their

corrosive properties.

They are the strongest hydroxides known.

Because they are soluble in water they are also known as alkalis.

It is from this property that the group earns its name as alkali metals.

NaOH, KOH, RbOH and CsOH are very soluble in water but LiOH is the least soluble.

The solubility (in g/100g of H2O) of the hydroxides increases down the group, LiOH,

13.0(25o); NaOH, 108.3(25

o); KOH, 112.8(25

o); RbOH, 197.6(30

o); CsOH, 385.6(15

o).

Although LiOH is a strong alkali it is the weakest of the lot.

Some of the reactions of the hydroxides

They react with acids to form salts,

MOH + HCl → MCl + H2O

They liberate ammonia from ammonium salts,

NaOH + NH4Cl → NaCl +NH3 + H2O

The hydroxides are stable to heat but LiOH decomposes on heating,

LiOH + heat → Li2O + H2O

There are many other reactions, which you will come across in the future lectures.

4.1.4.b Carbonates, bicarbonates, and nitrates

They form water-soluble carbonates, M2CO3 (M = Li, Na, Rb, Rb, Cs), solubility

increases down the group.

The carbonates are thermally stable but Li2CO3 decomposes on strong heating, and

stability increases as the group is descended.

Li2CO3 + heat → Li2O + CO2

With the exception of LiHCO3, which exists in solution, they form solid hydrogen

carbonates.

The hydrogen carbonates are readily decomposed by heat,

2MHCO3 + heat→ Na2CO3 + CO2 + H2O

They form nitrates

LiNO3 decomposes to the oxide and NO2.

The others decompose to the nitrites,

2MNO3 + heat → 2MNO2 + O2

On stronger heating NaNO3 decomposes to Na2O, O2 and N2

4NaNO3 + strong heat → 2Na2O + 5O2 + N2

The thermal stability of the nitrates increases down the group.

Nitrates are generally oxidizing agents and their reduction results in the nitrites,

NaNO3 + C + heat → 2NaNO2 + CO2

NaNO3 + Zn + heat → KNO2 +ZnO

4.1.4c Halides and polyhalides

All of them form halides of the type MX, where X = F, Cl, Br, I.

LiF is anhydrous, The other halides of lithium are trihydrate, LiX.3H2O

The halides of the other metals are anhydrous.

The tendency for Li to form hydrated halides is due to the small ionic size.

The heavier metals form poly-halides. Recall that they also form super oxides.

KI + I2 → KI3

The reason I2, which is only sparingly soluble in water, dissolves very readily in KI is

because I2 reacts with KI as in the equation above.

4.1.4d Some differences between lithium and other Group IA metals

Li has the highest melting and boiling points.

Li is harder than the other metals.

Li is the least reactive.

It forms only the normal mono-oxide, whereas the others form peroxides and super-

oxides.

LiOH is the least soluble and the least basic, and therefore forms less stable salts (Recall

that most compounds of Li decompose on heating).

Li is the only metal in the group, which reacts with N2 to form Li3N.

Although we are not discussing complexes, Li forms more stable complexes than the

other metals.

While all the salts of the other metals are soluble, LiF, Li2CO3, and Li3PO4 are insoluble

and LiOH is only sparingly soluble.

The halides of Li are more covalent, while the halides of the others are totally ionic.

Li+ ion is more heavily hydrated because of its small size.

4.2 GROUP IIA: ALKALINE EARTH METALS

Element Be Mg Ca Sr Ba Ra*

Atomic Number 4 12 20 38 56 88

Electron configuration [He]2s2 [Ne]3s

2 [Ar]4s

2 [Kr]5s

2 [Xe]6s

2 [Rn]7s

2

1st I.E./kJmol

-1- 899 737 590 549 503 509

2nd

I.E./kJmol-1

1757 1450 1145 1064 965 979

3rd

I.E./kJmol-1

14847 1731 4910 - - (3281)

E’vity 1.5 1.2 1.0 1.0 0.9 -

Metal radii/pm 112 160 197 215 222 -

Ionic radii/pm 31* 72 100 118 135 148

4.2.1 General comments

They are silvery white metals and harder than Group IA metals.

They have higher density.

They have higher melting/boiling points because two electrons are involved in

metallic bonding. This gives a higher binding energy than Group IA metals.

Similar trends in physical and chemical properties are observed as are observed in the

alkali metals, except for the differences associated with their smaller sizes.

They are generally less reactive than the Group IA metals because more energy is

required to remove two electrons.

The reactivity, however, increases down the group.

Their physical and chemical properties are dependent on the electronic configuration

of ns2 (i.e. all the elements have two electrons in the outermost shell).

The loss of the two outer electrons results in all the elements having a fixed oxidation

state of +2.

M M2+

+ 2e

Their compounds are therefore generally ionic.

Be, the first element in the group, is also unique and differs much from other metals

as Li does from other Group IA metals.

The reasons for the uniqueness of Be are:

Be atom and ion are extremely small.

Its electronegativity is relatively high, and

The maximum number of electrons which Be (in Period II) can

accommodate is 8. Other members of the group can have more than 8

electrons due to availability of empty orbitals.

Its charge density is high.

4.2.2 Trends in other properties

Electrons are more tightly held; therefore the first ionisation energy is greater than for

Group IA metals.

Loss of the first electron results in greater effective nuclear charge, which makes the

second ionization energy almost twice the first one.

The third ionisation energy is very high, since it is being removed from a closed shell;

hence M3+

is never formed.

The second ionisation energy of Be is very high, hence Be2+

forms covalent

compounds.

4.2.3 Electronegativity

The electronegativities are low, but much higher than for Group IA.

They still form ionic compounds when they react with non-metals, because the

electronegativity difference is still large.

Only BeF2 would have been ionic because of the large electronegativity difference

between Be and F. It, however, has very low conductivity when fused and is

therefore covalent.

4.2.4 Metallic/ionic radii

The radii are much smaller than the radii of Group IA metals but increases down the

group, as a result of additional shells.

The ionic radii of Li+ and Mg

2+, just as those of Be

2+ and Al

3+, are comparable.

o The chemical properties of Li and Mg are therefore similar but similarity is not as

pronounced as within a group.

o This phenomenon is known as a diagonal relationship.

If you recall that size decreases across the period and increases down the group, moving

diagonally these effects cancel each other and there is therefore no marked difference in

properties. Consequently the type of bonds formed and the properties of their compounds are

often similar, even though the oxidation states differ.

4.2.5 Chemical properties

Reactions with water

Be is the least electropositive and therefore the least metallic. Be, therefore, does not

react with water.

Mg does not react with cold water but reacts with steam to give Mg(OH)2 and H2.

The Eo values for the other members are within the range recorded for the Group IA

metals. They therefore react readily with cold water, liberating hydrogen and the

hydroxides.

M + 2H2O M(OH)2 + H2 (M = Ca, Sr, and Ba)

They are, however, less reactive than Group IA; this is because of the lower solubility

of the hydroxides which tends to cover the surface of the metal.

When Mg, for example, is amalgamated with mercury the surface MgO is removed

and the metal reacts smoothly with cold water.

4.2.6 Compounds

4.2.6a Hydroxides

Be(OH)2 is amphoteric.

The hydroxides of the others are basic.

Basic properties increase down the group.

Mg(OH)2 is a mild base and is used to treat acid indigestion.

Ca(OH)2 is called lime water, and is used to test for CO2 (CO2 turns lime water turbid).Ca(OH)2 + CO2 CaCO3 + H2O

white ppt The white precipitate dissolves on passing excess CO2 through the mixture,

CaCO3 + CO2 + H2O Ca(HCO3)2

soluble Ba(OH)2 is called baryta water, and is turned turbid by CO2 as well.

The decomposition temperatures of the hydroxides decrease as the group is

descended.

The hydroxides are less basic than the hydroxides of Group IA.

4.2.6b Oxides and peroxides

All the elements react with O2 to form oxides. They react faster when in the powder

form.

They also burn in air to form both oxides and nitrides. Recall that only Li (in Group

IA) reacts with N2 to form the nitride. (Write the formula of magnesium nitride)

Mg burns in air/oxygen with dazzling flame with evolution of large amount of heat.

The dazzling flame is used in flash photography.

The Group IA metals are more reactive with air and oxygen.

Ba forms barium peroxide, BaO2, as well as the normal oxide. Others do not form

peroxides.

Unlike Group IA no superoxide is known.

All the (solid) oxides

- have high melting points; m.p. of BeO is ~2500, MgO is ~2800 oC.

- have very low vapour pressure.

-are good heat conductors.

-are electrical insulators.

They are all therefore used for lining furnaces.

CaO, SrO, and BaO react with water with evolution of heat, e.g.,

CaO + H2O Ca(HCO3)2

soluble This is usually accompanied with swelling, and the calcium hydroxide obtained from this

reaction is known as slaked lime.

4.2.6c Carbonates

The decomposition temperatures of the carbonates are given below.

Carbonate BeCO3 MgCO3 CaCO3 SrCO3 BaCO3

Dec.

Temp./oC

<100 540 900 1290 1360

The carbonates are ionic.

They are insoluble in water, except BeCO3.

All the carbonates can be decomposed by heat.

As can be seen above, the decomposition temperatures increases as the group is

descended. This is because the size of the ions increases down the group. The smaller

cations polarize the large CO32-

anion and therefore introduce some degree of covalency

into the bonding, which makes it easy to decompose. Thus the carbonates of the smallest

Be2+

is the least thermally stable.

4.2.6d The sulphates

BeSO4 and MgSO4 are soluble in water.

CaSO4 is sparingly soluble.

The sulphates of the other metals are virtually insoluble.

Thus solubility decreases down the group.

The higher solubility of BeSO4 and MgSO4 is due to the high hydration energies of smaller

Be2+

and Mg2+

ions.

MgSO4.7H2O is Epsom salt and it is used as a laxative.

CaSO4.2H2O is gypsum and can be readily converted to CaSO4.½H2O, which is

plaster of Paris.

Their decomposition temperatures increase down the group, indicating that the more

basic the metal the more stable the sulphate.

BaSO4 is used in medicine in taking the X-ray of the intestine of patients that suffer

from ulcer.

Ba2+

ions (from BaCl2 or Ba(NO3)2, are used to test for SO42-

. White precipitate is given.

CO32-

also gives a white precipitate but the BaCO3 formed dissolves on addition of dilute

HCl.

Ba2+

+ SO42-

→ BaSO4(s) (white precipitate)

Ba2+

+ CO32-

→ BaCO3(s) (white precipitate)

BaSO4 + HCl (there is no visible reaction)

BaCO3 + 2HCl → BaCl2 + CO2 + H2O (dissolves in dilute HCl)

UNIT 5: TRANSITION METAL CHEMISTRY

5.1: General comments

There are two transition metal series:

- the d-block, generally referred to as transition metal series

- the f-block, generally referred to as inner transition metal series.

The d-block fall in the centre of the Periodic table in the 4th, 5th and 6th periods (corresponding to

filling the 3d, 4d and 5d orbitals of electrons), between the alkaline earth metals (Group II) and the

boron family element (Group III). There are 10 elements per period, giving 30 transition elements.

The 4th, 5th and 6th period transition metal series are called first, second, and third series

respectively.

There are two f-block series - the lanthanides and the actinides. The lanthanides follow the third d-

series after lanthanum, while the actinides follow actinum (period 7). There are 14 elements per

series, giving 28 f-block elements. There are therefore 58 transition metals altogether.

Our interest in this course is in the first row d-transition series, scandium to zinc.

5.2 Definition: A d-transition element may be defined as an element whose atom or at least one

of its ions has a partially filled d-orbital.

Recall that the d-orbital has a degeneracy of 5 i.e. has 5 sub-orbitals defined by the values of the

magnetic quantum numbers. Each sub-orbital accommodates two electrons each, hence 10 electrons

corresponding to the 10 elements in the first d-series.

Physical Properties

Element Atomic

number

Atomic

radii/pm

*Ionic

radii/pm

Density

gcm-3

M.P.

C

B.P.

C

Config. Abundance

ppm

Ca 20 197 100(II) 1.54 839 1484 [Ar]3d04s

2 46600

Sc 21 162 74.5 3.0 1539 2748 [Ar]3d14s

2 25

Ti 22 147 67 4.5 1667 3285 [Ar]3d24s

2 6320

V 23 134 64 6.11 1915 3350 [Ar]3d34s

2 136

Cr 24 128 61.5 7.14 1900 2690 [Ar]3d54s

1 122

Mn 25 127 64.5 7.43 1244 2060 [Ar]3d54s

2 1060

Fe 26 126 64.5 7.87 1535 2750 [Ar]3d64s

2 62000

Co 27 125 61 8.86 1495 3100 [Ar]3d74s

2 29

Ni 28 124 60(II) 8.91 1455 2920 [Ar]3d84s

2 99

Cu 28 128 73(II) 8.95 1083 2570 [Ar]3d10

4s1 68

Zn 30 134 74 7.13 419 907 [Ar]3d10

4s1 76

* For M3+

Chemical Properties

Element E'vity Ionisation energies* Oxn E

E

3

1st 2nd 3rd 4th state M2+/M M

3+/M

Ca 1.05 590 1146 4941 6464 +2 -2.87 -

Sc 1.3 633 1235 2388 7130 +3 - -2.077

Ti 1.5 659 1309 2648 4171 +3,+4 -1.63 -1.21

V 1.6 650 1414 2866 4631 +2,+3 -1.18 -0.85

Cr 1.6 653 1519 2992 4861 +2,+3 -0.91 -0.74

Mn 1.5 717 1509 3259 5021 +2,+3 -1.18 -0.28

Fe 1.8 762 1561 2958 5502 +2,+3 -0.44 -0.04

Co 1.8 759 1644 3230 5104 +2,+3 -0.28 +0.40

Ni 1.8 736 1751 3391 5400 +2,+3 -0.25 -

Cu 1.9 745 1958 3556 5681 +2,+1 +0.34 -

Zn 1.6 906 1732 3828 5983 +2- -0.76 -

* in kJmol-1

5.3 General Comments They are all metals and hence display all the characteristics of metals, e.g. malleable, ductile, etc.

Atomic numbers 21-30 (calcium is included for comparison only; calcium is not a transition metal).

Electron configuration: [Ar]3dn4s

2 (n=1-10) except [Ar]3d

54s

1 (for Cr) and [Ar]3d

104s

1 (for Cu)

where the half-filled and fully-filled d configurations respectively are more stable.

Atomic radii: A gradual decrease in atomic radii as one moves from left to right of the series. Here

electrons go into penultimate shell. The d electron shield the outer electrons pretty well from the

increasing nuclear charge, so the effective nuclear charge felt by the outer electrons (4s electrons)

increases very slowly, hence low decrease in size. A minimum seems to be reached because as the d

orbital becomes more than half-filled, electron repulsion force the d-orbital to gradually expand in

size which causes the sizes of the last few atoms to swell.

Ionic radii: They are less than the atomic radii and generally M2+

>M3+

>M4+

.

Oxidation State: The 3d transition metals are characterised by availability of incompletely filled 3d

energy level that is close to the 4s energy level. Electrons in both the 3d and 4s levels are therefore

close to the outer regions of the electron clouds and are available to interact with their surroundings

(i.e. are involved in bonding).

They all therefore exhibit variable oxidation states. The energy required to unpair and promote the

inner d electrons for use in bonding is low (see the ionisation energies). All show oxidation state of

+2 (except Sc) when both 4s electrons are involved in bonding. For oxidation states greater than 2, 3d

electrons are used in addition to both 4s electrons.

IIIB IVB VB VIB VIIB VIIIB IB IIB

Sc Ti V Cr Mn Fe Co Ni Cu Zn

7

6 6 6

5 5 5 5 5

4* 4* 4 4 4 4 4

3* 3 3 3* 3 3* 3 3 3

2 2 2 2* 2 2* 2* 2* 2*

1 1 1 1 1 1 1 1

(Important oxidation states in bold; *The most important oxidation state).

The lower oxidation states are generally reducing and the higher oxidation states are oxidising,

particularly where such redox leads to attainment of more stable configuration. Sc3+

attains noble gas

configuration and hence resembles group 3 metals. Zn has +2 state and resembles s-block. The

number of known oxidation states for the remaining elements increases with the number of unpaired

electrons, reaching maximum in manganese and iron. Manganese has the largest number of known

oxidation states (+1 to +7) of all the transition elements. From scandium to manganese, the maximum

oxidation state equals the sum of the number of 3d and 4s electrons, e.g.

Sc Ti V Cr Mn

3d14s

2 3d

24s

2 3d

34s

2 3d

54s

1 3d

54s

2

+3 +4 +5 +6 +7

No correlation between maximum oxidation state and configuration is found beyond manganese.

Generally the highest oxidation states occur in oxo-ions. For Cr, Mn and Fe the highest known

oxidation states are the least stable. Thus the dichromate ion (Cr2O72-

), the permanganate ion MnO4-

and the ferrate ion (FeO42-

) are all strong oxidising agents. The low oxidation states are found in

simple salts and are generally reducing unless in cases where their d configuration is stable (d5 and

d10

).

Melting point/boiling point: They are generally high. The 3d and 4s electrons are used in metallic

bonding, hence the increase up to d3. As pairing occurs there is decrease, i.e. less d electrons are

available for bonding. Cr is relatively low because the stable 3d5 configuration makes it more

difficult for the d electrons to be made available for metallic bonding. Mn is abnormally low partly

due to its complex crystalline structure which involves less efficient packing of the metallic atoms. It

is possible that the d-electrons are involved, to some extent , in covalent bonding within the Mn

metallic lattice.

Density: Generally high, increasing across the series. This is typical of most metals.

Abundance: Odd atomic number metals are less abundant than the even atomic number metals.

Ionisation energy: The first ionisation energy gradually increases across the series. The second

ionisation energy is larger than the first and also increases across the series. Both are higher than for

Ca. Abnormal increases are observed for the second ionisation energy in Cr and Cu corresponding to

the unusual stability of half-filled and fully filled orbital,

i.e. 3d54s

1(Cr) and 3d

104s

1(Cu)

e 3d

5(Cr

I) and 3d

10(Cu

I).

The 3rd ionisation energy for Ca is greater than the 3rd for all the transition series because the 3rd

electron from Ca is removed from a noble gas configuration. An increase of the 3rd ionisation energy

is observed across the series but note a bigger increase on Mn, a depression on iron, and a large

increase on Zn. Explain these observations.

The 4th is generally much larger than the 3rd in all but unusually high on Fe. Why? Thus generally

successive ionisation energies of the transition metal atom increase gradually. The big increase from

3rd to 4th and the gradual increase between 2nd and 3rd explain why +2 and +3 oxidation states are

common.

Electrode potential: Standard reduction potential increases (becomes less negative) across the series.

These E(M

2+/M) indicate a decreasing tendency to form individual cations across the series. In this

regard Mn is an exception which is the strongest reducing agent, i.e. Mn Mn2+

+ 2e is most

favoured. The reason is the attainment of stable d5 configuration.

E values indicate that all metals should be reducing agents (except Cu) and should react with

dilute non-oxidising acids (they are above H in the electrochemical series). Ti and V are passive

to dilute non-oxidising acids at room temperature. Cu is the least reducing, i.e. Cu2+

is the

strongest oxidising agent in aqueous solution.

For M3+

/M system, the Mn3+and Co

3+ ions are the strongest oxidising agent in aqueous solution.

Ti2+

and Cr2+

will liberate hydrogen from a dilute acid.

2Cr + 2H 2Cr + H (g).(aq)

2+

(aq) (aq)

3+

2

Electronegativity: Their electronegativities are intermediate between s-block and p-block element

and therefore can form both ionic and covalent bonds. Electronegativity increases from Sc to Zn.

The elements thus become less metallic in character from left to right. (This is also reflected in the

tendency to positive redox potentials on crossing the series).

5.4 Other general properties

Formations of interstitial compounds: The transition metals form interstitial compounds with non-

metals with small radii, e.g. H (37 pm), B (79 pm), C (77 pm) and N (73 pm). The structure of an

interstitial compound is different from that of the original metal, hence there are strong bonding forces

between the metals and the non-metals. Interstitial compounds often have non-stoichiometric

compositions and do not correspond to the normal oxidation states of the metal.

Examples: TiH1.7, PdH0.6 , VH0.56 . They are therefore not strictly compounds. However, there are

examples of some stoichiometric interstitial compounds, e.g. TiC, TiN, VN, Mn4N, Fe8N, TiH2. A

number of oxides also fall into this category.

Generally interstitial compounds are

inert chemically except towards oxidising agents

very hard

high melting e.g. TiC, 3410C; TiN, 3200C

good electrical conductors (although slight reduction due to less mobility of electrons)

less malleable because the gliding planes are "pegged" in position by the non-metallic atoms.

Carbon steels are interstitial compounds. The interstitial carbon atoms prevent the iron atoms in the

lattice from readily sliding over one another. This makes the iron harder and stronger but more brittle.

Paramagnetism: This is the attraction of magnetic lines of force by a compound. Where the

electrons are all paired there is no attraction and such a compound is said to be diamagnetic. There is

some permanent magnetic effect in some metals and compounds, e.g. Fe, Co, Ni. These metals can be

magnetised. They are said to be ferromagnetic. Ferromagnetic metals can be magnetised.

A paramagnetic compound is attracted by a magnet. This is due to the spin of one or more unpaired

electrons in an atom. The spinning of electrons about their axis generates magnetic moment. The

magnetic moment of an ion increases with the number of unpaired electrons. For the first row

transition metals, magnetic moments increase from Sc to Mn and then drop off.

V3+

Cr3+

Mn3

+

Fe2+

Ion Sc3+

Ti3+

Ti2+

V2+

Cr2+

Mn2

+

Fe2+

Co2+

Ni2+

Cu2+

Zn2+

Configuration d0 d

1 d

2 d

3 d

4 d

5 d

6 d

7 d

8 d

9 d

10

No of unpaired

electrons

0 1 2 3 4 5 4 3 2 1 0

Magnetic moments 0 1.73 2.83 3.87 4.90 5.92

The number of unpaired electrons can be derived by using box configuration. Recall that the d orbital

has a degeneracy of 5 (i.e. 5 sub-orbitals). Electrons are arranged following the Hund's rule.

For example, d6 ______ 4 unpaired electrons

The magnetic moment can be calculated by using the formula

= 2 s(s + 1) B.M. ...Eqn. 2

where s is the sum of the spins of all the unpaired electrons, i.e. n x ½ (n = number of unpaired

electron).

Example: Calculate the magnetic moment of Mn3+

. Configuration of Mn3+

is [Ar]3d4 or simply 3d

4

In box

form:

There are therefore 4 unpaired electrons,

= 4(4 + 2) (see equation 1)

= 24 = 4.90 B.M.

Alternatively using equation 2

s = n x ½ = 4 x ½ = 2

= 2 1s s( ) = 2 2 3( )

= 2 6 = 4.90 B.M.

5.5 Formation of complex ions Transition metal ions form complexes with polar molecules or ions known as ligands. Ligands are

either negatively charged or they have an atom with lone pair of electrons. Such atoms are known as

donor atoms. Examples of donor atoms include O, N, S, P, etc. Some of the ligands of interest are

CN-, H2O, NH3, Cl

-. When a set of ligands bond to a central metal ion the compound formed is

known as a complex. A complex can be neutral, anionic or cationic. Neutral ligands are normally

attached to the central metal ion by means of coordinated bond. Complexes are therefore also

referred to as coordination compounds. The number of ligands attached to the central atom is called

the coordination number/stoichiometry. Coordination compounds are not restricted to transition

metals only; other metal ions do form. However, for a metal ion to form a coordination compound,

the following conditions have to be fulfilled:

the size of the cation must be small;

the cation must carry a comparatively high charge;

there must be empty orbitals of the right energy.

These conditions favour acceptance of lone pairs of electrons. Where a particular metal forms more

than one ion, the one with higher charge forms more stable complexes, i.e. it has greater attraction for

electrons, e.g. Co2+

does not form a stable complex with ammonia, Co3+

does.

N and O donor atoms form more stable complexes. At oxidation state +2 complexes of the latter half

of the first transition series form complexes with a particular ligand with progressively increasing

stability from Mn2+

to Cu2+

. Zn2+

is, however, less stable than Cu2+

.

Complexes are usually written within a square bracket, which represent its coordination sphere.

Any species outside the coordination sphere is not part of the complex.

Consider a complex [Co(NH3)6]Cl3. Here the complex species in this salt is [Co(NH3)6]3+

, a cation.

Hence [Co(NH3)6]Cl3 is a cationic complex, even though it is a neutral salt.

What information can be derived from this complex?

The coordination compound is hexaamminecobalt(III) chloride.

The formula is [Co(NH3)6]Cl3 even though it was originally formulated as

CoCl36NH3.

The complex ion is [Co(NH3)6]3+

.

o The central atom is cobalt.

o The oxidation state of cobalt is +3.

o The ligands are six ammonia molecules.

o The coordination number is 6.

o The structure/geometry/stereochemistry is octahedron, i.e.

Co

NH3

NH3

NH3

NH3

H3N

H3N

Co

NH3

NH3

NH3

NH3

H3N

H3N

3+3+

The arrows represent coordinate bonds or dative bonds.

5.6 Types of Ligands

Simple ligands such as H2O, NH3, CN-, Cl

- are called monodentate ligands, since they can only form

one coordinate bond. Monodentate means "one tooth". Bidentate ligands have "two teeth" i.e. they

form two coordinate bonds.

An example of a bidentate ligand is ethylenediamine, H2N-CH2-CH2-NH2(en). The formula for

[Co(en)3]3+

is

H2C CH2

H2N NH2

Co

CH2

H2C

NH2

H2N

CH2

CH2

H2N

H2N

3+

There are tridentate (three teeth, 3 coordinate bonds); tetradentate (for 4 coordinate bonds);

pentadentate (for 5) and hexadentate (for 6), etc.

5.7 Charges on a complex The charges on an ion are delocalized over the whole complex ion. The charge is the algebraic sum

of the charge on the central ion and the charges on the ligands. E.g. [Fe(CN)6] : Here CN- carries -1

charge; the complex carries -4. Since there are 6 CN- ions, i.e. -6 charge, then the Fe carries +2

charge. Similarly in [Fe(H2O)6]3+

, Fe has +3 charge since H2O is a neutral (zero charge) molecule.

The overall charge on complex ions are calculated thus:

[Fe(CN)6]3-

: charge = (2+) from Fe + 6(-1) from CN- from CN

- = -4

[Fe(H2O)6]3+

: charge = (3+) + 6(0) = +3

The oxidation state of the central atom can also be calculated this way:

[Fe(CN)6]3-

: Let the oxidation state of Fe = x

x + 6(-1) = -3

x – 6= -3

x = +3.

5.8 Chelation

Some ligands are able to form ring structures with central atoms. This property is known as chelation

and the resulting compound is known as chelating compounds. Ethylenediamine is an example of a

chelating ligand.

5.9 Coordination number (C.N.) The structure/geometry/stereochemistry of a coordination compound depends on the coordination

number.

The most common coordination numbers are 6 and 4.

C.N. GEOMETRY EXAMPLE

2 Linear [CuCl2]-, [Cl-Cu-Cl]

-

[Ag(NH3)2]+, [H3N-Ag-NH3]

+

3 Trigonal Not common

4 Tetrahedral [Ni(CO)4]

4 Square Planar [Ni(CN)4]2-

5 Square pyramidal TlCl5

5 Trigonal bipyramidal [CuCl5]3-

6 Octahedral [Co(NH3)6]3+

5.10 Stability of a complex The relative stability of a complex is measured by its stability constant. When stability constants are

large, complex is stable; when small, complex is unstable.

Consider

[Fe(H2O)6]3+

+ 6CN-

[Fe(CN)6]3-

+ 6 H2O

kstab =[[Fe(CN)6]

3-]

[[Fe(H2O)6]2+

][CN-]6 = 10

37dm

18mol

-6 at 298 K

This stability constant is very large and indicates that the formation of the hexacyano complex goes

virtually to completion. It also shows that [Fe(CN)6]3-

is more stable than [Fe(H2O)6]3+

. Stability

constants are also referred to as formation constants.

The log of stability constants, log kstab., are often tabulated as indicated in the table below:

COMPLEX logkstab

[Mn(en)3]2+

5.7

[Fe(en)3]2+

9.6

[Co(en)3]2+

13.8

[Ni(en)3]2+

18.1

[Cu(en)3]2+

18.7

Recall that the stability of the complexes increases across the series from manganese to

copper.

5.11 Colour Compounds and ions of the d-block elements tend to be coloured. Why are they coloured?

When white light falls on a substance it may be totally reflected. Such a substance appears white. If

it is totally absorbed the substance appears black. In certain substances certain wavelengths are

absorbed and others are reflected. Such substances appear coloured. The colour is due to the

reflected wavelengths.

The absorption of light is due to the presence of unpaired d-electrons in the transition metal ions.

Aqueous solutions of Sc3+

and Zn2+

are colourless. These ions do not contain unpaired electrons.

Unpaired electrons absorb light energy by becoming promoted from the ground-state energy levels to

their excited state energy levels.

The wavelength of the light absorbed depends on the energy difference E between the ground

state and excited state.

E

Ground state

d-electron in ground state

h

Excited state

In a transition metal octahedral complex six ligands approach the metal ion. As they approach

the degeneracy of the d-orbitals split into two:

E = h

= frequencyh = Plunck's constant

dn

t2g

eg

The separation E depends on a number of factors:

the ligand

the metal ion

the charge on the metal ion (oxidation state)

the structure of the complex.

As each of the above changes, E changes, i.e. the frequency of the light absorbed, , changes and

hence colour changes. Examine the table below to illustrate the above factors.

Variation of ligands Variation of metal ions

Complex Colour Complex Colour

[Co(H2O)6]2+

pink [Co(H2O)6]2+

pink

[CoCl4]2-

blue [Fe(H2O)6]2+

pale green

As ligand changes from H2O to Cl-

colour changes

[Cu(H2O)6]2+

[Ni(H2O]6]2+

blue

green

As the metal ions change colour changes.

Note that the ligand is the same.

The charge on ion The structure of complex

Complex Colour Complex Structure Colour

[Fe(H2O)6]2+

pale green [Ni(H2O]6]2+

Octahedral green

[Fe(H2O)6]3+

pale yellow [Ni(HDMG)2] Square Planar red

Here the oxidation state of Fe

changes from +2 to +3

The colours of gemstones are due to the presence of traces of d-block ions:

Gemstone Colour Ion

Emerald green Cr3+

blue sapphire blue V3+

or Co3+

Ruby red Cr3+

Amethyst purple Mn3+

Topaz yellow Fe3+

Turquoise blue-green Cu2+

5.12 Catalytic Activity Catalytic activity is due to the ease with which electrons are lost and gained or moved from one

energy level to another. Many transition metals, or their compounds, are important catalysts.

For example:

nickel in hydrogenation;

vanadium(V) oxide, V2O5, in the contact process for the production ofsulphuric acid;

iron in the Haber process to synthesise ammonia;

catalysts containing cobalt, platinum, palladium, rhodium or titanium are employed in many

industrial processes for the production of organic compounds from unsaturated hydrocarbons.

The transition metals are therefore very important in the chemical industry.

5.13 Hydrolysis of transition metal ions Hydrolysis is reaction with water. Transition metal salts with strong acids, e.g. chlorides,

sulphates and nitrates, give acidic solutions when dissolved in water. For the same metal, the

acidity increases with increase in the charge carried by its ions. Example, for the same

concentration of Fe2+

and Fe3+

in water, the Fe3+

solution will be more acidic. For ions carrying

the same charge there is an increase in acidity in order of increasing atomic number of the

transition metal. Hydrolysis occurs in a number of other non-transition metal ions. Generally

hydrolysis process can be represented as

[M(H2O)x]n+

+ H2O [M(H2O)x-1OH](n-1)+

+ H3O+

It is the formation of H3O+ in water which makes the aqueous solution acidic.

5.14 Naming of transition metal complexes

Ligands

(a) Ions

Cl- chloro CN

- cyano

Br- bromo SCN

- thiocyanato

I-

iodo NO2

- nitro

(b) Neutral

H2O aqua/aquo NH3 ammine

(c) Number of ligands

Two, bi-(or di-); three, tri-; four, tetra-; five, penta- and six, hexa- as prefixes to the

ligands.

Complexes

(a) Cationic:

Metals bear their ordinary names, followed by their oxidation state (in Roman numerals) put between

brackets e.g. Fe2+

, iron(II), Cr3+

, chromium(III).

Anionic

Metals bear their latin names ending with -ate, followed by oxidation state in brackets, e.g. iron

becomes ferrate, chromium becomes chromate, copper becomes cuprate, etc.

Sequence of naming

Name the cation first, then the anion, where the complex is a salt.

(a) Cationic complexes

[Co(NH3)6]Cl3: Here the complex ion is [Co(NH3)6]3+

, a cation, and the anion is Cl-. The anion

is outside the coordination sphere therefore it is not a ligand in this complex. It is a chloride ion.

Note that there are six ammonia molecules as ligands, hence hexa-ammine (or hexaammine).

This complex is named hexaamminecobalt(III) chloride. Note that the complex cation is

written as one word and there is a space between the cation and anion.

[Cu(H2O)6]2+

hexaaquacopper(II) ion

[FeCl(H2O)5]+ pentaaquachloroiron(II) ion

(b) Anionic complexes

K3[Fe(CN)6]: The cation here is K+ and the complex ion carries a negative charge,

[Fe(CN)6]3-

. Its name is potassium hexacyanoferrate(III)

[CoCl4]2-

tetrachlorocobaltate(II) ion

[PtCl6]2-

hexachloroplatinate(IV) ion

(c) Neutral complexes

[Ni(CO)4] tetracarbonylnickel(0)

[Pt(NH3)2Cl4] diamminetetrachloroplatinum(IV)

You can also write the formula of a complex from its name:

hexaaquatitanium(IV) [Ti(H2O)6]4+

tetrabromocobaltate(II) [CoBr4]

2-

5.15 Trends in the oxides of transition metals Generally oxides at low oxidation states are basic, those of intermediate oxidation states are

amphoteric and those at high oxidation states are acidic. ZnO is, however, amphoteric, even

though the oxidation state of zinc is +2. For manganese, for example, known oxides are

MnO - basic, Mn2O3-basic, MnO2--amphoteric, Mn2O5 and MnO3 are rather unstable but would

be expected to be acidic; Mn2O7-strongly acidic. Mn3O4 is a mixed oxide, MnO.Mn2O3 i.e.

manganese(II)manganese(III) oxide.

UNIT 6: CHEMICAL BONDING

6.1: Trends in chemical bonding

A chemical bond is a force that acts strongly enough between atoms or groups of atoms to

hold them together in a different species that has measurable properties.

The types of chemical bonds formed by atoms of different elements can be related to the

positions of the elements in the periodic table.

6.2 Bond types

6.2.1 Metallic bond

Metals other than those in the p, d and f blocks have only one or two electrons in their highest

energy levels. In general all metals give up their outer electrons more easily than non-metals.

In metallic bonding each metal atom contributes its valence electrons to what has been called

―a sea of electrons‖. These valence electrons once contributed do not belong to any

particular metal atom; the electrons are delocalized.

Metallic bonding is therefore the attraction between positive metal ions and surrounding

freely mobile electrons. These electrons are from s (for Groups I and II), s and p (for p-block

metals), s and d (for d-block metals) and s and f (for f-block metals) orbitals.

Thus a metallic crystal can be pictured as an assembly of spherical atoms packed together and

bonded to each other equally in all directions.

e- e

-

e-

e-

e-

e-

e-

e-

+ ++

+ +

+ +

+

cations

sea of electrons

+

++

+

e-

e-

e-

e-

e-

The energy involved in metallic bond is called binding energy.

The higher the valence electrons involved in metallic bonding the larger the binding

energy and the higher the m.p. and b.p. of the metal.

The binding energy also dictates whether the metal is soft or hard. Na is softer than

Mg and Mg is softer than Al where 1, 2, 3 valence electrons respectively are involved

in metallic bonding.

6.2.2 The ionic bond

Ionic bonding is the attraction between positive and negative ions.

Ionic compounds are collections of ions held together by electrostatic attraction

between cations and anions.

In ionic bonding metals donate their valence electrons to non-metals such that both

attain a stable noble gas configuration according to the octet rule. The octet rule

states that atoms tend to combine by gain, loss, or sharing of electrons so that the

outer energy level of each atom holds or shares four paired electrons.

Example: the formation of NaCl

2,8,1 2,8,7 2,8 2,8,8

1s22s

22p

63s

1 1s

22s

22p

63s

23p

5 1s

22s

22p

6 1s

22s

22p

63s

23p

6

[Ne]3s1[Ne]3s

23p

5[Ne] [Ar]

Note that the configurations of the ions formed are also written. Thus, just as configurations

can be written for atoms they can be written for ions as well.

Ionic solids are stable, high – melting substances. They are non conductors when in solid

state but are strong electrolytes when molten or in aqueous solution.

6.2.3 Covalent bonding

A covalent bond occurs when two atoms of non-metals contribute one electron each to bond

formation. Covalent bonding is therefore based upon electron sharing and is the attraction

between two atoms that share electrons.

2,8,7 2,8,7 2,8,8; 2,8,8

This is an example of a single bond, where two atoms are held together by sharing two

electrons. Double, triple, or multiple covalent bonds are possible by sharing 4, 6 or more

electrons by two atoms. Quadruple bonds are rarely formed.

N N

N N

8 electrons 8 electrons

This is a triple bond

N N

NaCl

x x

x x

xxx Cl

x x

x x

xxxNa

++

Cl Cl

x x

x x

xx

xCl

Cl

x x

x x

xx Clor Cl Clor

Cl

x x

x x

xx x

O C O O C O

8 electrons 8 electrons

8 electrons

The structural formula of CO2 is O C O

Covalent compounds could exist as gas, liquid or solid. Generally they have low melting

points but the melting points could be unusually high depending on the molecular mass and

on the strength of intermolecular bonds that exist between them. (See your CHE101 notes).

6.2.4 Variation in bond types

Non-polar and polar covalent bonds:

In molecules like H2, Cl2, Br2, F2 and N2 the bond electrons are shared equally. Such

molecules are said to be non-polar.

Covalent compounds where the bond electrons are not shared equally are said to be polar.

A polar molecule is a dipole- a pair of opposite charges of equal magnitude at a specific

distance from each other. The entire molecule is electrically neutral.

Examples are HF and H2O

H

F polar covalent bond

Bonding continua

The non-polar covalent bond is at one end of a continuum of variation in the polarity of

bonds. The other end is the completely ionic bond. Bonding in most compounds, either ionic

or covalent, is not 100% covalent or 100% ionic. In most compounds the bond is in between-

anywhere along the continuum of bonding from 100% covalent to 100% ionic. There is not

likely to have a compound that is 100% ionic.Most ionic compounds are polarized, whereby

some covalency is introduced into a bond that is presumed to be ionic. The cation attracts

electron density of the anion, causing it to be somewhat unsymmetrical, or polarized.

Polarization

The polarization of an ion is the distortion of its electron cloud by an ion of opposite charge.

Cations, being small and having high charge density, tend to attract the electron

clouds of anions.

The smaller the cation, the higher its charge density and the greater its polarizing

ability and the more it draws electron density into the region between itself and the

anion. This results in increase in covalent character of the bond.

+-

Idealized 100% ionic bond

+ -

Polarized, partly covalent, ionic bond The charge-to-size ratio of a cation (obtained by dividing the charge by its ionic radius) is a

relative measure of the polarizing ability of a cation.

The larger the charge-to-size ratio of an ion, the greater the degree of covalent

character in its bonds. The ratios of charge-to-radii of Be2+

and Ca2+

is 2/0.035 (= 57)

and 2/0.099 (= 20) respectively. This shows that Be2+

has a greater ability to polarize

an anion than Ca2+

. Most Be compounds, as you will see later, are highly covalent.

6.3 LEWIS SYMBOLS: ELECTRON BOOK-KEEPING

6.3.1 Binary compounds

Lewis symbols are used to show the valence electrons of an atom or ion. To draw the Lewis

symbol for an element we write its chemical symbol surrounded by a number of dots (or

other similar symbols) which represent the atom‘s valence electrons, e.g. Lithium (1s22s

1) has

one valence electron.

Li x

Other members of Period 2 elements have

NeFONCBe BLix

xx x x xx

xxx

xx

xx

x

x

xxx

xx

xxx

xx

xx

xx

x xxxxx

Thus, the group number for the representative elements is equal to the number of valence

electrons.

Recall the octet rule, which states that when atoms react, they tend to achieve an outer shell

having eight electrons (the noble gas configuration). The octet rule applies only to the

representative elements (the first 18 elements in the Periodic Table). In general, it does not

work with transition metals.

Lewis symbols in ionic bonding, e.g. NaCl (See page 36)

Lewis symbols in covalent bonding, e.g. Cl2 (See page 36)

Chlorine gas is made up of two chlorine atoms hence chlorine is a diatomic molecule. The

structural formula is Cl—Cl.

In most of the compounds in which they occur carbon forms four covalent bonds, nitrogen

forms three and oxygen forms two:

C

H

HH

H

N

H

HH O

H

H

or

C

H

HH

H

N

H

HH O

H

H

x x

x xx

xxx

x

6.3.2 Multiple Bonds

The bond produced by the sharing of one pair of electrons between two atoms is called a

single bond. There are, however, many molecules in which more than one pair of electrons

are shared between atoms, e.g. N2 (See pages36-37)

6.3.3 When the octet rule fails: exceptions to the octet rule:

The atoms in some molecules cannot obey the octet rule because there are too few or too

many electrons.

This happens when the atom forms more than four bonds, e.g. PCl5 and SF6 and also in

CIF3. There are five P—Cl bonds, six S—F bonds and three Cl—F bonds respectively.

Since each covalent bond requires the sharing of a pair of electrons, phosphorus and

chlorine have 10 electrons each and S has 12. The Lewis formulae are

P

Cl

Cl

Cl

Cl

Cl

S

F

F

F

F

F

Cl

F

F

F

extremely stable10e in valence shell

F

Elements in Period 2 such as carbon or nitrogen, never exceed an octet simply because

their valence shells, having n = 2, can hold a maximum of only 8 electrons.

Elements in Periods 3 and above, however, sometimes do exceed an octet, because their

valence shells can hold more than 8 electrons.

Period 3, n = 3 hold maximum 18 electrons

Period 4, n = 4 hold maximum 32 electrons

In some molecules (but not many), an atom has less than an octet. The most common

examples are compounds of Be and B.

ClBe Cl Be Cl

ClB Cl B Cl+ 3

+ 2

4e around Be

6e around B

Compounds with unpaired electrons: Some compounds exist in which one or more

electrons remain unpaired. In most the total of the valence electrons of the central atom

and the atom bonded to it is an odd number, e.g. ClO2 has a total of 19 valence electrons

(6 from each of the two oxygen atoms and 7 from the chlorine atom.) The compound is

paramagnetic, corresponding to 1 unpaired electron.

A reasonable Lewis structure is

O Cl O A gas

This compound is unstable and explosive.

6.3.4 Lewis structures for molecular compounds and polyatomic ions

To write a Lewis structure one must know the arrangement of the atoms, i.e. which atoms are

directly bonded to each other? Is SO2 arranged like SOO or OSO? These

arrangements can be predicted using the following guidelines:

Write the correct arrangement of the atoms.

Smaller, more electronegative non-metal atoms surround larger, less electronegative

non-metal atoms.

Oxygen, hydrogen, and/or halogen atoms often surround a central metal or non-metal

atom in a symmetrical arrangement.

Carbon atoms are usually bonded to each other.

Oxygen atoms are bonded to each other only in peroxides (or superoxides).

In most acids, such as H2SO4 and in many other compounds that contain oxygen and

hydrogen atoms, the hydrogen atoms are all bonded to oxygen atoms.

OHOSO

HO OO

HON O

OClO

O

OCO

Count the total number of valence electrons by

adding the number of valence electrons

subtracting one electron for each unit of positive charge or adding one electron for

each unit of negative charge

(It helps to know the group of the elements to fix the number of valence electrons.)

SOH HNO ClO CO E.g.42342

Assign electrons to each covalent bond. E.g. How many dots, representing electrons, must

appear in the Lewis structures of SO3, NO3- and NH4

-

SO3 : S (Group VIA) contributes 6 x 1 electrons = 6

O (Group VIA) contributes 6 x 3 electrons = 18

24e-

NO3

- : N (Group VA) contributes 5 x 1 electrons = 5e-

O (Group VIA) contributes 6 x 3 electrons = 18e-

Add another electron for the –1 charge = 1e-

24e-

NH4

: N (Group VA) contributes 5 x 1 electrons = 5e-

H (Group IA) contributes 1 x 4 electrons = 4e-

Subtract one electron for the +1 charge = 1e-

8e-

Distribute the remaining electrons so that each atom has the appropriate number of non-

bonded electrons:

for period 2 elements (except Be and B) this is the number of electrons needed so that

each atom is surrounded by an octet (i.e. 8 electrons);

for Period 3 and beyond (except Al) this is often the number of electrons needed to

complete an octet, although extra electrons can also be placed around atoms of these

elements when they are the central atoms in compounds, (atoms bonded to a central

atom usually obey the octet rules), e.g. Write the Lewis structure for the SO4

-2 ion.

The skeletal structure is given below:

O O

S

O O

Total number of electrons = 6e- (from S) + 24e

- (from four O atoms)

= 30e-

Add 2e- for the –2 charge 2e

-

32e-

Electron distribution in the skeletal structure is done by placing a pair in each bond:

The final structure has 32 dots and each atom obeys the octet rule. The final structure

is therefore

S OO

O

O

2_

O

O2

_

or

O

SO

When, sometimes, you find that there are either too few electrons to complete the octets

of all the atoms or there are electrons left over after all the octets have been filled, apply

the following rules:

When there are insufficient electrons to give each atom an octet, create a multiple

bond (recall that Be, B and Al are exceptions.) (Multiple bonds can be written

between C, N, O, S, Se and P atoms);

when electrons are left over, they are always placed on the central atom, in pairs, e.g.

Write the Lewis structure for the air pollutant SO3

O

O

SO

O

O

O S

Note that there is no octet around S. We cannot add more dots, thus we move a pair of

electrons that solely belong to one of the O atoms to a SO bond so that they can also

be counted as belonging to S, i.e. a double bond is placed between S and one of the

oxygens.

O

O

O S

O

O

O S

O

O

O Sor

6.3.5 Resonance

N N O or N N ON2O can be written as

O O O or O O OO3

C OCO32-

O

O

C O

O

O

C O

O

O2- 2- 2-

or or

All these possible structures are accepted to be correct to a certain extent, but none of them

adequately represents the correct structure of each of the species. Lewis structure is not

completely adequate to explain these possibilities. All the bond lengths are equal, even

though the single bonds are expected to be longer than the double bonds. The real structure

of each of the examples above is a mixture of all the possible structures shown.

orBenzene

The ‗mixing‘ of structures is called resonance and the resulting structure is called a resonance

hybrid. The actual molecule is always the same. There is only electron redistribution.

Sometimes a dashed or dotted line in a single structure is used to indicate resonance, e.g.

O O O

or

6.3.6 Rules for writing resonance structure

The sequence of atoms in each resonance structure must be the same, i.e. the same atoms

must be connected to the same other atoms, e.g.

are not resonance structures for the same compound.

All resonance structures for the same molecule must have the same total number of

valence electrons.

Example: Write the possible resonance forms of nitryl chloride, NO2Cl, in which

nitrogen is the central atom.

The structure is

Total number of valence electrons = 24. With 6 electrons in the three single bonds, 18

electrons remain to be distributed. For complete octets on each atom, 20 electrons are

required (7 for Cl, 5 for N, 6 each for the O atoms), two more than we have. Therefore,

N C O H and H N C O

O N O

Cl

one double bond must be used. This has to be NO double bond (NCl double bond is

not possible). This gives two resonance forms.

Cl

N

OO

Cl

N

OO

Cl N

O

O

or

6.4 Shapes of molecules (Valence Shell Electron Pair Repulsion Theory – VSEPR

Theory)

VSEPR theory uses repulsion between electron pairs as the basis for predicting molecular

geometry.

To apply this theory, molecules and ions are classified according to the number of bonding

electron pairs and lone pairs surrounding central atoms. The predicted molecular geometry is

the one that places the atoms or groups bonded to the central atom, as well as the lone pairs,

as far apart as possible.

Let us examine a molecule ABn where A is the central atom.

For n=1 AB can only be linear, whether or not electron pairs are present. AB, e.g.

HCl

H Cl

For n=2 AB2 contains two shared electron pairs. Obviously they achieve the greatest

distance from each other if the AB bonds are on opposite sides in a linear

molecule, BAB.

For n=3 AB3 BAB angles of 120 place the bonding electron pairs at the greatest

distance from each other, i.e. a planar molecule, with A at the centre of an

equilateral triangle; and B at the corners.

A

B

B

B For n=4 AB4 is tetrahedral with bond angles at 109.47.

B

AB

B

B

For n=5 AB5 is triangular bipyramidal

Be

Be

Be

A

Ba

Ba

Ba = axial BBe = equatorial B

<BaABe = 120o

<BeABa = 90o

For n=6 AB6 is octahedral B

BB

B B

A

B

<BAB = 90o

In predicting molecular geometry by VSEPR theory, double and triple bonds are treated like

single bonds, e.g. methanal, H2C=O, is classified as AB3 molecule. In H2C=CH2 the carbon

atom is treated as A of an AB2 situation where one of the B‘s is =CH2. Double bonds take up

more space around A than single bonds because their greater electron density repels the other

bonding electrons. As a result, the bond angles are distorted from the ideal angles above.

H

H

O118o

120o

CC C

H

H

H

H121.3

o

Triangular planar molecule A planar molecule

6.4.1 Rules for using VSEPR Theory

Write the correct Lewis structure for the molecule and determine whether or not there are

any lone electron pairs present on the central atom. If there are no lone pairs, the

molecule is of the ABn type

ABn

n=1

n=2

n=3

Linear

Linear

Trigonal planar

A B

B A B

B A

B

B

n=4 n=5

n=6

Tetrahedral Trigonal bipyramidal

octahedral

B

A

B

B

B

B

B

A

B B

B

BB

BB

A

B

B

Example: Use VSEPR to predict the geometry of

(a) PF5 (b) SO42-

(c) BeCl2(g)

(a) P, Group V, Period 3 has 5 valence electrons. Some of the electrons can occupy vacant

3d orbitals, the P atom can accommodate more than an octet of valence electrons. There

are no lone pairs of electrons in the P atom.

P

F

F

F

F

F

PF5 is an AB5 molecule and should have triangular (trigonal) bipyramid geometry.

b) S is in Group VI (Period 3). There is no lone pair on the S atom (from the Lewis

structure) and therefore it is an AB4 ion. The geometry is tetrahedral.

O

O

O O

S

2-

c) Be has 2 valence electrons and the Lewis structure shows no lone pair electrons or the

central Be atom.

BeCl2 is AB2 type molecule. It is linear. (This structure is for gaseous BeCl2).

6.4.2 VSEPR Theory in molecules with lone-pair electrons on the central atom

Like bonding pairs, lone pairs repel each other. There is also repulsion between lone pairs

and bond pairs. This leads to arrangements in which all electron pairs – lone and bonding –

are arranged as far as possible from each other. Thus bonding and lone pairs assume the

same general arrangements as shown earlier. However, because in molecular geometry we

are looking only at the positions of atoms, replacing an atom with a lone pair changes the

shape of the molecule, e.g. CH4, NH3 and H2O molecules each has 4 electron pairs around a

central atom. Each has a different molecular geometry but all derive their geometry from the

AB4 tetrahedron.

From the angles shown in the above structure it is evident that

lone pairs repel

lone pairs

more

than

lone pairs repel

bonding pairs

more

than

bonding pairs repel

bonding pairs

or

lone pair - lone

pair

repulsion

>

lone pair - bond

pair repulsion

>

bond pair - bond

pair repulsion

NH3 is AB4-type molecule where one of the B‘s is replaced by a lone pair. Such a molecule

is represented as AB3E (E = lone pair). H2O is AB4-type with two B‘s replaced by lone pairs,

hence AB2E2.

Lone pairs repel bond pairs of AB resulting in BAB bond angles reducing. Hence

HCH angle in CH4> HNH in NH3> HOH in H2O.

In general:

Cl ClBe

AB2E - angular/bent AB3E -

AB4E - distorted tetrahedral AB5E

AB2E2 AB3E2

AB4E2 AB2E3

-

- -

- -

T-shape

square planar Linear

square pyramidal

angular/bent(or distorted tetrahedral

B A

BA

B B

B

B

AB

B

B

BA

BB

BB

B A

B

B

AB

B

triangular pyramidal

e.g. NH3, PCl3

e.g. SF4 e.g. ClF5

e.g. H2O, SCl2

BA

B

BB

e.g. XeF4

B

A

B

e.g. XeF2

UNIT 7: Solid state

7.1 Structure and Types of Solids

Binding forces holding particles together in solids may be primary chemical bonds, as in

metals and ionic solids, or they may be secondary van der Waals‘ forces of solids, such as in

ice, paraffin wax and most polymers. In solids, the way the atoms or molecules arrange

themselves contributes to the appearance and the properties of the materials.

Atoms can be gathered together as an aggregate through a number of different processes,

including condensation, pressurization, chemical reaction, electro-deposition, and melting.

The process usually determines, at least initially, whether the collection of atoms will take to

form of a gas, liquid or solid. The state usually changes as its temperature or pressure is

changed. Melting is the process most often used to form an aggregate of atoms. When the

temperature of a melt is lowered to a certain point, the liquid will form either a crystalline

solid or/and amorphous solid.

Amorphous Solids: A solid substance with its atoms held apart at equilibrium spacing, but

with no long-range periodicity in atom location in its structure is an amorphous solid.

Examples of amorphous solids are glass and some types of plastic. They are sometimes

described as super-cooled liquids because their molecules are arranged in a random manner

somewhat as in the liquid state. For example, glass is commonly made from silicon dioxide

or quartz sand, which has a crystalline structure. When the sand is melted and the liquid is

cooled rapidly enough to avoid crystallization, an amorphous solid called a glass is formed.

Amorphous solids do not show a sharp phase change from solid to liquid at a definite melting

point, but rather soften gradually when they are heated. The physical properties of amorphous

solids are identical in all directions along any axis so they are said to have isotropic

properties.

.

Si

O

O

O

Si

O

O Si

Si

Si

Si

O

O

O

O

O

O

O

O

SiO

O

Si

O

Si

Si

SiO

O

Si O

OO

O

O

Si

O

O

SiO

O

Si

O

Si O

Crystalline Solids: More than 90% of naturally occurring and artificially prepared solids are

crystalline. Minerals, sand, clay, limestone, metals, carbon (diamond and graphite), salts (

NaCl, KCl etc.), all have crystalline structures. A crystal is a regular, repeating arrangement

of atoms or molecules. The majority of solids, including all metals, adopt a crystalline

arrangement because the amount of stabilization achieved by anchoring interactions between

neighboring particles is at its greatest when the particles adopt regular (rather than random)

arrangements. In the crystalline arrangement, the particles pack efficiently together to

minimize the total intermolecular energy. Proper arrangement of component ions or

molecules gives rise to the well-definedVBN shapes of crystals.

•The particles in a crystal are located in a well-defined array called a crystal lattice.

•The smallest portion of the crystal that is repeated in all three directions is the unit cell.

Several types of unit cells are illustrated on the next few slides.

Crystalline Solids: Crystalline solids often fall into one of the following packing categories:

Simple Cubic – The simple cubic unit cell consists of a cube with one atom at each corner.

There is one atom in each unit cell (1/8

th of 8 atoms). Each atom is in contact with six other

atoms (coordination number = 6), and the packing efficiency is 52%.

Body-Centered Cubic – The body-centered cubic unit cell consists of a cube with one atom

at each corner and another atom in the center of the cube. There are two atoms in each unit

cell. The coordination number is 8, and the packing efficiency is 68%.

Face-Centered Cubic – The face-centered cubic unit cell consists of a cube with one atom

at each corner and another atom in the center of each cube face. There are four atoms in each

unit cell. The coordination number is 12, and the packing efficiency is 74%.

Hexagonal Closest Packing – Hexagonal closest packing occurs when layers of atoms stack

on top of each other in such a way that the upper and lower layers (Layer A) are offset from

Layer B by half a sphere, giving a hexagonal unit cell with a coordination number of 12 and a

packing efficiency of 74%.

Cubic Closest Packing – Cubic closest packing occurs when layers of atoms stack on top of

each other in such a way that the upper and lower layers (Layer A and C) are offset from

Layer B and from each other by half a sphere, giving a hexagonal unit cell with a

coordination number of 12 and a packing efficiency of 74%.

Types of Crystalline Solids

Type of Solid Intermolecular Forces Properties Examples

Molecular solids London forces, dipole-

dipole forces, hydrogen

bonds

Soft, low melting

points, non-

conducting

H2O, Br2, CO2,

CH4

Ionic solids Ion-ion forces Brittle, hard, high

melting points

NaCl, KBr,

MgCl2

Nonbonding atomic

solids

London forces Very low melting

points

Ar, Kr, Xe

Metallic atomic solids Metallic bonds Variable hardness

and melting point,

conducting

Na, Zn, Cu, Fe

Covalent network solids Covalent bonds Hard, high melting

points

C (diamond,

graphite), SiO2

(quartz, etc.)

All four categories involve packing discrete molecules or atoms into a lattice or repeating

array, though network solids are a special case. The categories are distinguished by the nature

of the interactions holding the discrete molecules or atoms together.

7.2 Structure and Bonding in Metallic Solids

A metallic crystal can be pictured as containing spherical atoms packed together and bonded

to each other equally in all directions. We can model such a structure by packing uniform,

hard spheres in a manner that most efficiently uses the available space. Such an arrangement

is called closest packing. The spheres are packed in layers in which each sphere is

surrounded by six others. In the second layer the spheres do not lie directly over those in the

first layer. Instead, each one occupies an indentation (or dimple) formed by three spheres in

the first layer. In the third layer the spheres can occupy the dimples of the second layer in two

possible ways: They can occupy positions so that each sphere in the third layer lies directly

over a sphere in the first layer [the abab arrangement], or they can occupy positions so that

no sphere in the third layer lies over one in the first layer [the abca arrangement]. The abab

arrangement has the hexagonal unit cell, and the resulting structure is called the hexagonal

close packed (hcp) structure.

abab arrangement: abca arrangement:

The abca arrangement has a face-centered cubic unit cell and the resulting structure is called

the cubic close packed (ccp) structure. Note that in the hcp structure the spheres in every

other layer occupy the same vertical position (ababab . . .), whereas in the ccp structure the

spheres in every fourth layer occupy the same vertical position (abcabca . . .). A

characteristic of both structures is that each sphere has 12 equivalent nearest neighbors: 6 in

the same layer, 3 in the layer above, and 3 in the layer below (that form the dimples).

Molecular Solids

There are three basic types of crystalline solids: molecular, ionic, and atomic solids.

Molecular solids are composed of molecules held in a crystal lattice by intermolecular

forces. They are fairly soft, and have relatively low melting points. They are poor conductors

of heat and electricity. An example of molecular solid is ice (H2O). Each individual water

molecule is held together by intermolecular hydrogen bonding between oxygen atoms of one

molecule and the hydrogen atom of neighbouring molecules as shown below. The extensive

hydrogen bonding in ice also exists in liquid state and is responsible of the relatively high

boiling point of water.

H

OH

H

O

H

H

OH

H

OH

H

O

HH

OH

H

O

H

H

O

H

Atomic Solids

Atomic solids are composed of individual atoms. Nonbonding atomic solids are held

together by London forces. They have very low melting points that increase with increasing

atomic mass. These include the noble gases in their solid form (Ar, mp -189 oC; Xe, mp -112

oC).

Covalent Network Solids (Giant molecules)

This is a relatively rare structure, diamond being probably the best known example. Boron

nitride (BN)n and silicon carbide (SiC)n (carborundum) are similar types of solid. These

solids are non-conducting, indicating that the electrons are less free and more localized than

the electrons in a metal which move easily allowing an electric current to flow through the

lattice.

In network solids, conventional chemical bonds hold the chemical subunits together. The

bonding between chemical subunits, however, is identical to that within the subunits,

resulting in a continuous network of chemical bonds. Two common examples of network

solids are diamond (a form of pure carbon) and quartz (silicon dioxide). In quartz one cannot

detect discrete SiO2 molecules. Instead the solid is an extended three-dimensional network of

...-Si-O-Si-O-... bonding.

Carbon: Carbon exists as a pure element at room temperature in three different forms:

graphite (the most stable form), diamond, and fullerene.

Diamond – The structure of diamond is shown below. The balls represent the carbon atoms

and the sticks represent a covalent bond. In addition, a single stick is drawn to represent a

covalent bond. In the diamond structure, all bonds are single covalent bonds (sigma bonds).

Here, the carbon atoms are not close-packed. Each carbon is surrounded tetrahedrally by four

other carbon atoms. Clearly, each carbon is exerting a tetrahedrally directed force on its

neighbours and such directed forces are operative throughout the whole crystal. Diamond is

found to be a refractory solid, i.e. it has an extremely high melting point, indicating that the

bonding forces are extremely strong.

Notice that diamond is a network solid. The entire solid is an "endless" repetition of carbon

atoms bonded to each other by covalent bonds. (In the display at the right, the structure is

truncated to fit in the display area.)

Graphite – The most stable form of carbon is graphite. Graphite consists of sheets of carbon

atoms covalently bonded together. These sheets are then stacked to form graphite. The

display at the right shows a ball-and-stick representation of graphite. The sheets extended

"indefinitely" in the xy plane, but the structure has been truncated for display purposed.

Graphite may also be regarded as a network solid, even though there is no bonding in the z

direction. Each layer, however, is an "endless" bonded network of carbon atoms.

Questions

1. What is the bonding geometry around each carbon?

2. What is the hybridization of carbon in graphite?

3. The layer of the graphite structure consists of a repeating series of rings. How

many carbon atoms are in a ring?

4. What force holds the carbon sheets together in graphite?

5. Graphite is very slippery and is often used in lubricants. Explain why this

property is expected on the basis of the structure of graphite.

6. What is the bonding geometry around each carbon?

7. What is the hybridization of carbon in diamond?

8. The diamond structure consists of a repeating series of rings. How many

carbon atoms are in a ring?

9. Diamond is renowned for its hardness. Explain why this property is expected

on the basis of the structure of diamond.

10. The slipperiness of graphite is enhanced by the introduction of impurities.

Where would such impurities be located and why would they make graphite a

better lubricant?

Silicon Dioxide – Silicon dioxide (SiO2), also called silica, occurs naturally in many forms.

Quartz is essentially pure silicon dioxide. Sand is composed of small quartz fragments. Many

precious gems are quartz containing coloured impurities. Amethyst is quartz coloured red by

the presence of iron(III) ions. Agate and onyx are also quartz containing impurities. Flint is

silica coloured black by carbon.

Quartz has a very complicated crystal structure, which involves interwoven helical chains.

When heated to about 1500o C, quartz changes into the mineral cristobalite, whose structure

is shown at the right in ball-and-stick form. The brown balls represent the silicon atoms and

the red balls represent the oxygen atoms. Cristobalite is pure SiO2. Notice the similarity in

structure between cristobalite and diamond.

Ionic Solids

Ionic solids are made up of cations and anions held together in a crystal lattice by the strong

attractions between opposite charges on. The type of crystal structure formed in ionic solids

is dependent on the relative sizes of the cations and anions, the coordination number of the

ions, and the need to maintain charge neutrality. These substances are hard and brittle, with

high melting points (e.g. KBr, CsCl, NaCl, ZnS, CaF2).

7.3 Energy Effects in Ionic Compounds

Ionic solid forms because the aggregated oppositely charged ions have a lower energy than

the original elements. How strongly the ions attract each other in the solid state is a measure

of the lattice energy – defined as the change in energy that takes place when separated

gaseous ions are packed together to form an ionic solid:

e.g.: K+

(g) + Br-(g) KBr(s)

Since lattice energy is often defined as the energy released when an ionic solid forms from its

ions, the lattice energy has a negative sign; i.e. when the ions are brought together the energy

decreases. Consider the following illustration for the formation of solid lithium fluoride from

its elements:

Li(s) + ½F2(g) LiF(s)

Breaking this reaction into individual steps that make up the sum of the overall reaction is

presented below:

1. Sublimation of solid lithium:

Li(s) Li(g) (Sublimation energy = 161 kJ/mol.)

2. Ionization of lithium atoms to form Li+ ions in the gas phase:

Li(g) Li+

(g) + e- (1

st ionization energy = 520 kJ/mol.)

3. Dissociation of fluorine molecules:

½F2(g) F(g) (Dissociation energy = 154 kJ/mol. or 77kJ for ½mol.)

4. Formation of F- ions from fluorine atoms in the gas phase:

F(g) F-(g) (Electron affinity of fluorine = -328 kJ/mol.)

5. Formation of solid lithium fluoride from the gaseous Li+ and F

- ions:

Li+

(g) + F-(g) LiF(s) (Lattice energy for LiF = -1047 kJ/mol.)

Since the sum of these five processes yields the desired overall reaction, the sum of the

individual energy changes gives the overall energy change (= -617 kJ per mole of LiF). This

process is summarized by the energy diagram above.

E

Li+(g) + ½F2(g)

Li+(g) + F(g)

Li+(g) + F-(g)

Li(g) + ½F2(g)

Li(s) + ½F2(g)

LiF(s)

(1) 161 kJ

(2) 520 kJ

(4) -328 kJ

(5) -1047 kJ-617 kJ

(3) 77 kJ

A large decrease in energy accompanies formation of LiF(s) from its constituent elements

mainly due to the very large negative lattice energy. In fact, note that the energy released

when an electron is added to a fluorine atom to form the F- ion (328 kJ/mol) is not enough to

remove an electron from lithium (520 kJ/mol).

i.e. Li+

(g) + F-(g) Li

+(g) + F

-(g) ; not favoured.

Therefore, the driving force for the formation of an ionic solid such as LiF (unlike in covalent

compounds) results from the strong mutual attractions among the Li+ and F

- ions in the solid

(i.e. the lattice energy is an important factor).

In the structure of the LiF solid (illustrated below), there is alternating arrangement of the Li+

and F- ions. Each Li

+ is surrounded by six F

- ions, and vice versa. The arrangement

maximizes attractions among the oppositely charged ions and minimizes the repulsions

among the like charges. This sort of arrangement is found in all binary ionic solids formed by

an alkali metal (except in caesium salts) and a halogen.

7.4 Calculation Lattice Energy

In discussing the energetics of the formation of solid lithium fluoride, we emphasized the

importance of lattice energy in contributing to the stability of the ionic solid. Lattice energy

can be represented by a modified form of Coulomb‘s law:

𝐿𝑎𝑡𝑡𝑖𝑐𝑒𝑒𝑛𝑒𝑟𝑔𝑦 = 𝑘 𝑄1𝑄2𝑟

where k is a proportionality constant that depends on the structure of the solid and the

electron configurations of the ions, Q1 and Q2 are the charges on the ions, and r is the

shortest distance between the centers of the cations and anions. Note that the lattice energy

has a negative sign when Q1 and Q2 have opposite signs. This result is expected, since

bringing cations and anions together releases energy. Observe that the process will be more

exothermic the greater the charges and as the separation between the ions decrease.

The importance of the charges in ionic solids can be illustrated by comparing the energies

involved in the formation of NaF(s) and MgO(s) as shown in the figure below. These solids

contain the isoelectronic ions Na+, F

-, Mg

2+, and O

2-. Note the following important features:

a. Energy released when the gaseous Mg2+

and O2-

ions combine to form solid MgO is

more than four times greater than that released when the gaseous Na- and F

-ions

combine to form solid NaF.

b. The energy needed for 2 e ionization of Mg (2180 kJ/mol) is much greater than 1 e

ionization of sodium atom (495 kJ/mol).

c. The electron affinity for 2 e addition to gaseous oxygen atom is 737 kJ/mol. This

energy includes an endothermic component (2nd

electron addition) that must be

obtained indirectly, since the O2-

(g) ion is not stable.

Since much more energy is required for the 2nd

ionization of Mg relative to 1st ionization

energy and since addition of an electron to the gaseous O- ion requires a much more energy,

it might have been expected that magnesium oxide would be easier formed from Mg+ and O

-

ions rather than Mg2+

and O2-

ions. The reason for the formation of the ionic solid from Mg2+

and O2-

ions can be explained by looking at the lattice energy. Exothermicity of the lattice

energy for combining gaseous Mg2+

and O2-

ions ions to form MgO(s) is 3000 kJ/mol more

than that for combining gaseous Na+ and F

- into NaF(s). This is a compensation for the

energies consumed in formation of the Mg2+

and O2-

ions.

Variety of factors operates to determine the composition and structure of ionic compounds.

The most important of these factors involve the balancing of the energies required to form

highly charged ions and the energy released when highly charged ions combine to form the

solid.

Assignment: Suggest reason(s) why the ionic solid NaF is formed from Na+ and F

- ions and

not from Na2+

and F2-

ions as for magnesium oxide.

It is convenient to classify solids according to what type of component occupies the lattice

points. This leads to the classifications:

- atomic solids (atoms at the lattice points)

- molecular solids (discrete, relatively small molecules at the lattice points)

- ionic solids (ions at the lattice points).

In addition, atomic solids are placed into the following subgroups based on the bonding that

exists among the atoms in the solid:

- metallic solids

- network solids

- Group 8A solids

In metallic solids, a special type of delocalized non-directional covalent bonding occurs. In

network solids, the atoms bond to each other with strong directional covalent bonds that lead

to giant molecules, or networks, of atoms. In the Group 8A solids, the noble gas elements are

attracted to each other with London dispersion forces. The classification of solids is

summarized in the following table.

Atomic Solids

Molecular

Solids

Ionic

Solids

Metallic Network Group 8A

Occupants of Discrete Ions Metal Non-metal Group 8A

lattice

Points:

molecules

atoms atoms atoms

Bonding: Dipole–dipole

and/or London

dispersion forces

Ionic

Delocalized

covalent

Directional

covalent

(leading to

giant

Molecules)

London

dispersion

forces