FUNDAMENTALS OF CHEMISTRY by PolisasLib3

FATHMAH BINTI SALIM

JABATAN TEKNOLOGI MAKANAN POLITEKNIK SULTAN HAJI AHMAD SHAH

FUNDAMENTAL OF CHEMISTRY

FATHMAH BINTI SALIM

POLITEKNIK SULTAN HAJI AHMAD SHAH

FUNDAMENTAL OF CHEMISTRY

FATHMAH BINTI SALIM

Published by POLITEKNIK SULTAN HAJI AHMAD SHAH SEMAMBU 25350 KUANTAN eISBN: 978-967-XXXX-XX-X

Copyright ©2022, by Politeknik Sultan Haji Ahmad Shah Materials published in this book under the copyright of Politeknik Sultan Haji Ahmad Shah. All rights reserved. No part of this publication may be reproduced or distributed in any form or by means, electronic, mechanical, photocopying, recording, or otherwise or stored in a database or retrieval system without the prior written permission of the publishers.

PREFACE

Fundamentals of chemistry is essential and vital knowledge as a foundation in order to bear a fruitful superstructure for student entire career. It plays a leading role in shaping the future of research and development. This book contains three chapter which is Matter, Periodic Table and Chemical Bonding. Terms and basic concepts that help in understanding chemistry will be discussed in this chapter.

TABLE OF CONTENT

CHAPTER

TOPIC

PAGE

MATTER

1 - 44

PERIODIC TABLE

45 - 85

CHEMICAL BONDING

86 - 109

MATTER

UNIT 1A

OBJECTIVES GENERAL OBJECTIVE : To understand the concept of matter, the particle theory of matter, elements and compounds; to analyse the arrangement of particles in solids, liquids and gases as well as to understand the changes of state and to know the processes involved. SPECIFIC OBJECTIVES : At the end of this unit you will be able to : ✓ Explain the meaning of matter. ✓ Explain the fundamental particles such as atoms, molecules and ions. ✓ Differentiate between elements and compounds. ✓ Compare the arrangement of particles in solids, liquids and gases. ✓ Explain the changes of state that occur in heating and cooling. ✓ State the processes involved in the changes of state.

INPUT 1A.1 THE MEANING OF MATTER Matter refers to any material that has mass and occupies space. Matter is the construction unit of all material around us, whether living or nonliving. According to the theory of conservation of matter, matter cannot be destroyed but can be changed from one state to another. As an example, when candle wax (an example of matter) is burnt, smoke, carbon dioxide, and water are formed. All the products of combustion of wax are matter. Matter can exist in three physical states, namely the solid, liquid or gaseous states at different temperatures, as shown in diagram 1a.1

Diagram 1a.1 Three physical states (solid, liquid and gaseous states). 1A.2 PARTICLE THEORY OF MATTER According to this theory all matter consists of tiny and discrete particles that are constantly in motion. These particles are held together to different degrees by forces of attraction. The particles in a certain substance may be atoms, molecules or ions.

Atoms • •

are neutral (that is, not electrically charged) particles. Examples are atoms of copper, hydrogen and carbon, are the most basic units for any element, for example the oxygen atom is the most basic unit of the element of oxygen, 2

• •

are capable of taking part in chemical reactions, for instance a carbon atom can combine with an oxygen atom to form carbon dioxide, and can exist independently, for example the sodium atom and potassium atom.

Molecules • • •

•

are neutral (that is, not electrically charged) particles. Examples are oxygen gas and nitrogen gas, consist of two or more atoms, of the same kind (for instance the oxygen molecule, O2) or of different kinds (as in the case of the carbon dioxide molecule, CO2), can be formed from atoms in small groups (for example simple molecules such as the oxygen molecule, O2 and the ethanol molecule, C2H5OH) or big groups (for example the big and complex protein molecules, which consist of millions of different atoms), and can exist independently, for example the nitrogen molecule (N2) and the glucose molecule (C6H12O6).

Ions • are electrically charged particles which are formed from electrically neutral atoms or molecules through the loss or gain of one or more electrons. •

Positively charged ions are formed when electrically neutral atoms or molecules lose electrons. They are known as cations. Example: Na sodium atom

•

→

Na+ sodium cation

Negatively charged ions are formed when electrically neutral atoms or molecules gain electrons. They are known as anions.

Example: Cl

+ 1e

→

chlorine atom

Clchloride anion

Some elements are found in 2 or more different forms called allotropes. Did you know that coal and diamond are allotropes of carbon? The only difference between them is in the arrangement of atoms, but the difference in price in the market amounts to thousands of RM!!! Maybe you could create jewellery from coal !!! Try to think of a way… Photo 1a.1: Diamond and carbon

1A.3

ELEMENTS AND COMPOUNDS Matter can be classified as elements and compounds. An element is the simplest form of a chemical substance and cannot be broken down any further into other substances regardless of any chemical or physical process.

Elements can be divided into • • •

Metallic elements, such as sodium (Na) and zinc (Zn) Non-metallic elements, such as hydrogen (H2) and chlorine (Cl2) and Semi-metallic elements, such as boron (B) and silicon (Si).

Elements can exist as solids, liquids or gases depending on temperature and pressure. Compounds are pure substances which are formed by chemically combining two or more elements.

Compounds can be divided into two kinds, namely •

Covalent compounds which are made up of particles in the form of molecules, for example water (H2O), carbon dioxide (CO2) and ethanol (C2H5OH)

•

Ionic compounds which are made up of charged particles (cations and anions) for example sodium chloride (Na+Cl-), magnesium oxide (Mg2+O2-) and copper (II) sulphate (Cu2+SO42-)

Test your understanding by carrying out learning activity 1A.

ACTIVITY 1A

Question 1.1 What do you understand by matter? Question 1.2 Give two (2) properties of atoms, molecules and ions. Question 1.3 State two (2) examples of a metallic element, a non-metallic element and a semi-metallic element. Question 1.4 Define the terms ‘element’ and ‘compound’.

FEEDBACK 1A

Answer 1.1 Matter is anything that has mass and occupies space. Answer 1.2 Two (2) properties of atoms: • Electrically neutral, for example the copper atom • Smallest part of an element, for example the oxygen atom • Can take part in chemical reactions, for example a carbon atom combines chemically with two oxygen atoms to form carbon dioxide; or • Can exist by itself, for example the sodium atom. Two (2) properties of molecules: • electrically neutral, for example oxygen gas • may consist of atoms in small groups, for example the oxygen molecule, or in large groups as for proteins ; or • may exist in separate units, for example the nitrogen molecule. Two (2) properties of ions : • electrically charged particles • formed through a process of loss or gain in electrons Answer 1.3 Two (2) examples of metallic elements are sodium and zinc. Two (2) examples of nonmetallic elements are hydrogen and chlorine. Two (2) examples of semi-metallic elements are boron and silicon. Answer 1.4 Elements are the simplest form of chemical substances and cannot be broken down into other substances, with whatever chemical or physical process. Compounds are pure substances formed when two or more elements combine chemically.

INPUT 1a.4 THE ARRANGEMENT OF PARTICLES IN SOLIDS, LIQUIDS AND GASES Matter exists in three possible states or phases, the solid, liquid and gaseous states at various temperatures and pressure. The physical state of a substance at any particular temperature or pressure is influenced by the properties of its particles such as: (a) arrangement, freedom and speed of motion, (b) energy content and (c) forces of attraction between the particles. These properties influence the physical properties of the substance, such as rate of diffusion, density, compressibility and melting point as well as boiling point. A comparison of the three physical states of matter is made in Table 1a.1 Table 1a.1 Comparison of the three physical states of matter Property

Solid

Liquid

Gas

Arrange -ment of particles

Densely packed and ordered arrangement of particles: Particles held in fixed positions

Arrangement of particles less ordered but particles still in contact.

Particles unarranged and widely spaced.

Property Forces of attraction between particles Freedom of motion

Energy content Rate of diffusion of particles Density

Solid Very strong and cannot be overcome because of low energy content of particles No free motion, only vibration and rotation (no translational motion)

Lowest Lowest because particles have no free motion.

Highest because particles are arranged and packed closely together Compress Almost -ibility incompressible because particles are packed and arranged. Melting Highest because of and strong forces of boiling attraction between points particles.

Liquid Moderately strong and easily overcome because particles at higher energy state. Free motion but limited to low speed and within a small area. Liquid particles show vibrational, rotational and translational motion. Moderately high

Gas Very weak.

Moderately high because particles have free, though limited, motion. Moderately high because particles are quite close together

Highest because particles move freely at high speed. Lowest because particles are not arranged and are far apart.

Poor compressibility because of lack of space between particles. Low because of weaker forces of attraction between particles.

Easily compressed because much space is available between particles.

Move freely at high speed, besides vibrating and rotating.

Highest

Very low because very weak forces of attraction between particles.

1.5

CHANGES OF STATE Matter can change from one physical state to another when it is heated or cooled and its temperature changes. For example: (a) When ice (solid) is heated to 0C, it begins to melt. This continues until all of it changes to water (liquid). If heating is continued the temperature eventually reaches 100C and the water begins to boil and changes to steam (gas). (b)

Conversely, when steam is allowed to cool, it condenses, becoming water at 100C till all the steam has condensed. If the water is then allowed to cool, it eventually freezes, changing back into ice at 0C. heated

Ice (Solid) cooled

heated Water (Liquid) cooled

Steam (Gas/Vapour)

In a change of state, a transfer of energy takes place; heat energy is either absorbed or released. Table 1a.2 gives the names of processes involved in a change of state, as well as the transfer of energy that occurs.

Table 1a.2 Names of processes in a change of state and transfer of energy involved Change of state Name of Transfer of energy process Solid

Liquid Melting

Heat energy is absorbed

Liquid

Solid

Freezing

Heat energy is released

Liquid

Gas

Boiling / Evaporation

Heat energy is absorbed

Gas

Liquid

Condensation

Heat energy is released

Solid

Gas

Sublimation

Heat energy is absorbed

Gas

Solid

Sublimation

Heat energy is released

Water in gaseous state

Water in solid state

Water vapour Ice Freezing

Condensation

Melting Evaporation Water Water in liquid state

Diagram 1a.2: Processes involved in changes of state

ACTIVITY 1B

TEST YOUR UNDERSTANDING BEFORE GOING ON TO THE NEXT INPUT …! DO CHECK YOUR ANSWERS AGAINST THE FEEDBACK SECTION ON THE FOLLOWING PAGE. Question 1.5 State the factors that determine the physical state of matter. Question 1.6 Draw the arrangement of particles in the solid, liquid and gaseous states. Question 1.7 Fill in the blanks below with the properties of the three physical states of matter: PROPERTY

SOLID

LIQUID

Energy content

Melting and boiling points

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GAS

FEEDBACK 1B

Answer 1.5 Factors that determine the physical state of matter are : • Arrangement of particles • Freedom and speed of motion of particles • Energy content of particles • Forces of attraction between particles

Answer 1.6 PROPERTY

SOLID

LIQUID

GAS

Arrangement of particles

Answer 1.7 PROPERTY Energy content

Melting and boiling points

SOLID

LIQUID

GAS

Lowest

Moderately high

Highest

Highest because of strong forces of attraction between particles.

Low because of weaker forces of attraction between particles.

Very low because very weak forces of attraction between particles.

SELF EVALUATION

YOU ARE CLOSE TO SUCCESS. PLEASE TRY ALL THE QUESTIONS IN THIS SELF EVALUATION AND CHECK YOUR ANSWERS AGAINST THE FEEDBACK SECTION THAT HAS BEEN PREPARED. IF ANY PROBLEM ARISES, DO DISCUSS IT WITH YOUR LECTURER.

Question 1 Explain the concept of conservation of matter, using one (1) appropriate example. Question 2 Match the following statements with the most accurate answers: A B

Particle theory of matter Cation

Anion

Covalent compound

Positively charged because of loss of electrons from an atom or molecule Consists of molecules, for example water All matter consists of tiny, discrete particle that are constantly moving Negatively charged due to gain of electrons by an atom or molecule

Question 3 Complete the following diagram : P

Q R solid

S liquid

Vapour / gas

Question 4 Explain the following factors regarding the arrangement of particles in a gas, which influence its physical properties. (a) rate of diffusion of particles (b) compressibility

Question 5 Find the answers to the following questions in the box below: 1. According to the Law of _____________ , different gases have the same volume and pressure. 2. A _________________ is a pure substance formed from 2 or more elements that combine chemically. 3. One of the nonmetallic elements. 4. The one who first observed Brownian Motion. 5. This formula shows the simplest ratio of atoms of the elements in a compound. 6. A metal that is more reactive than copper. 7. Random motion can be described as ____________ . 8. The gas that gives a “pop” sound when a lighted splinter is introduced. 9. The function of anhydrous calcium chloride in the experiment to determine the empirical formula of a compound. 10. A black solid that undergoes sublimation. 11. A gas that diffuses very rapidly. 12. A gas that turns lime water chalky. 13. Matter is anything that has mass and occupies ___________. 14. Beside speed and freedom of motion, kinetic energy affects _______. 15. The physical state in which forces of attraction between particles are strongest. 16. Defined as having atoms of the same element but differing in the number of neutrons in each atom. 17. An example of a semimetal. 18. According to the current theory of matter, matter consists of tiny, ________ particles. 19. Its molecular formula is C6H6. 20. The change from solid to gas is called __________________. 21. The change from gas to liquid is known as ___________________.

S A M M O N I A R I O D I N E O Z C X M

B A D K J D S O I K M O S P S I Y C D M

S O L I D G A L S X B R O M S Q Z O O S

L J F D V P J R O B E R T B R O W N J I

Z I G Z A G M U F W N M E B E N J D J C

H L R L N L A A Z E Z Y M P D E S E R A

H Y D R O G E N F M E D P I I S J N K R

M K O Q B J E G U K N N E I S U K S B B

C A X F A S D K D M E K R O C B B A E O

K N O O D P O O K L I T A S R L D T N N

S J N S I R O N M E L J T O E I P I O D

B B M K D P U K J J D O U O T M R O W I

D A B S O R B S W A T E R J E A G N A O

R A O D K J D S M K S I E O N T K J D X

M A V O G A D R O S K J K W Z I P S C I

S C D Z X L S S I L I C O N M O N Z F D

P K H W Z Z B S I C O M P O U N D D Z E

A E D H E L I U M A T T J M E N N E D A

C S X Z Q B U H L K D S G C Y L O A S I

E D P E M P I R I C A L F O R M U L A G

FEEDBACK

Have you tried the exercises? If so please check your answers. Answer 1 According to the concept of conservation of matter, matter cannot be destroyed but can be transformed from one physical state to another. For example, when wax (an example of matter) is burnt, smoke, carbon dioxide, water etc. are produced. All these products of combustion are still matter. Matter can exist in three physical states, namely solid, liquid or gas at various temperatures.

Answer 2 A

Particle theory of matter Cation

Anion

Covalent compound

Answer 3 P = Sublimation Q = Condensation R = Melting S = Evaporation T = Freezing U= Condensation Answer 4 (a) rate of diffusion of particles = fastest because the particles in a gas are free to move at high speed. (b) compressibility = highly compressible because particles in a gas are far apart.

Answer 5

S A M M O N I A R I O D I N E O Z C X M

B A D K J D S O I K M O S P S I Y C D M

S O L I D G A L S X B R O M S Q Z O O S

L J F D V P J R O B E R T B R O W N J I

Z I G Z A G M U F W N M E B E N J D J C

H L R L N L A A Z E Z Y M P D E S E R A

H Y D R O G E N F M E D P I I S J N K R

M K O Q B J E G U K N N E I S U K S B B

C A X F A S D K D M E K R O C B B A E O

K N O O D P O O K L I T A S R L D T N N

S J N S I R O N M E L J T O E I P I O D

B B M K D P U K J J D O U O T M R O W I

Congratulations! You have successfully completed all the learning activities in Unit 1A. Carry on with unit 1B.

D A B S O R B S W A T E R J E A G N A O

R A O D K J D S M K S I E O N T K J D X

M A V O G A D R O S K J K W Z I P S C I

S C D Z X L S S I L I C O N M O N Z F D

P K H W Z Z B S I C O M P O U N D D Z E

A E D H E L I U M A T T J M E N N E D A

C S X Z Q B U H L K D S G C Y L O A S I

E D P E M P I R I C A L F O R M U L A G

UNIT1B

MATTER

OBJECTIVES GENERAL OBJECTIVES: To understand and analyse relative atomic mass, understand and apply relative molecular mass, know mole concept as well as chemical formula SPECIFIC OBJECTIVES: At the end of this unit, students should be able to : • Explain relative atomic mass • Compare relative atomic mass based on hydrogen, oxygen and carbon atom • Explain and calculate relative molecular mass • State mole as a unit of quantity, mass and molar volume • State the meaning of chemical formula • Differentiate empirical formula and molecular formula • Write chemical formulae of ionic compounds and hydrated compounds

INPUT 1B.1 RELATIVE ATOMIC MASS An atom of any element is extremely small. It is impossible to measure or weigh an atom in gram. Chemists compared how heavy one atom is to another atom which is taken as the standard. The mass of an atom when compared to another is known as the relative atomic mass. Relative atomic mass is not the actual mass but is just a value of comparison. Thus relative atomic mass does not have any unit. 1B.1.1 COMPARISON OF RELATIVE ATOMIC MASS BASED ON HYDROGEN, OXYGEN AND CARBON ATOM. A. Based on hydrogen scale

In the beginning of the 19th century, hydrogen was chosen by scientists as the standard because it is the lightest. The atomic mass of a hydrogen atom is assigned a value of 1 unit. Thus the relative atomic mass of hydrogen is represented as H = 1. Based on the hydrogen scale (H=1), relative atomic mass (R.A.M.) of an element is defined as the number of times one atom of the element is heavier than one atom of hydrogen, that is, R. A. M. of an element

Mass of 1 atom of the element = Mass of 1 atom of hydrogen

Later on, hydrogen atom was considered to be unsuitable to be used as the standard and was replaced by oxygen (O = 16). This was because

(a)

Almost all elements combine with oxygen to form pure compounds which can be analysed more accurately. Not many elements react with hydrogen.

(b)

Inaccuracies that arose in the chemical analysis when hydrogen (H = 1) was used.

B. Based on oxygen scale The discrepancies disappear when oxygen atom (O=16) is used. Based on oxygen scale (O=16), relative atomic mass (R.A.M.) of an element is defined as the number of times one atom of the element is heavier than 1/16 oxygen atom, R. A. M. of an Mass of 1 atom of that element element = 1/16 x Mass of 1 atom of oxygen

However, relative atomic mass based on oxygen scale was not satisfactory too because the physicists and the chemists were using two sets of scales that differ from one another because of the existence of 3 isotopes of oxygen such as 168 O , 178 O and 188 O C. Based on carbon scale In 1961, the International Union of Pure and Applied Chemistry, (IUPAC) begin to use carbon-12 isotope (12C=12.000) as the standard to replace oxygen atom. Carbon-12 was chosen as the standard because scientists found that carbon-12 was the most abundant amongst the other isotopes. The relative atomic mass of atoms or molecules were determined using the mass spectrometer with carbon-12 as the standard. The mass of a carbon -12 atom is assigned a value of 12.000 units. Based on the scale carbon-12=12.000, the relative atomic mass of an element is defined as the number of times one atom of the element is heavier than 1/12 of the mass of a carbon -12 atom.

R. A. M of an element

Mass of one atom of the element 1/12 x mass of one carbon-12 atom

Example. Mass of one atom of nitrogen is 14 times heavier than 1/12 of the mass of a carbon -12 atom. Thus the relative atomic mass of nitrogen is 14. The relative atomic mass based on C-12 scale is used till today. Relative atomic mass of any element can be obtained from the Periodic Table. 1B.2 RELATIVE MOLECULAR MASS The relative molecular mass (R.M.M.) of an element or a compound is defined as the number of times one molecule of the compound is heavier than 1/12 of the mass of a carbon -12 atom (12C=12.000), that is R. M. M. of mass of 1 molecule of the element or compound an element = 1/12 x mass of 1 atom of carbon-12 or a compound

Example, relative molecular mass of oxygen gas (O2) is 32. This means one molecule of oxygen gas is 32 times heavier than 1/12 of the mass of a carbon-12 atom. For example, a molecular compound, carbon dioxide (CO2) has a relative molecular mass of 44. This means one molecule of carbon dioxide gas is 44 times heavier than 1/12 of the mass of a carbon -12 atom. For an ionic compound such as sodium chloride (NaCl), magnesium oxide (MgO) and potassium carbonate (K2CO3), the term relative formula mass (R.F.M.) is used. This is because an ionic compound is made up of ions and not molecules.

Relative molecular mass or relative formula mass can be determined by the following method: (a) (b) (c)

Determine the chemical formula of the particular compound Find the relative atomic mass of each element in that chemical formula Find the sum of relative atomic masses of all the atoms as represented by the chemical formula

Table 1B.1 shows relative molecular mass or relative formula mass of some substances. Table 1B.1: Relative molecular mass / relative formula mass of some substances. Substance

Relative molecular (formula) mass

Bromine

Chemical formula Br2

Water

H2O

2H + O = (2 x 1) + 16 = 18

Sulphuric acid H2SO4 Anhydrous Copper(II) sulphate Hydrated magnesium sulphate

CuSO4

MgSO4.7H2O

Br + Br = 80 + 80 = 160

2H + S + 4O = (2 x 1) + 32 + (4 x 15) = 98 Cu + S + 4O = 64 + 32 + (4 x 16) = 160

Mg + S + 4O + 7H2O = 24 + 32 + (4 x 16) + (7 x 18) = 246

ACTIVITY 1B

Test your understanding by doing the following questions. Question 1B.1 Explain the definition relative atomic mass (R.A.M.) based (a) Hydrogen Scale (b) Oxygen Scale Question 1B.2 What is the definition of relative atomic mass of an element based on the IUPAC system as agreed in 1961. Question 1B.3 Find the relative molecular mass for: (a) CuSO4 (b) C6H12 O6 (c) Br2

Responses 1B

Answers 1B.1 (a) Based on hydrogen scale, relative atomic mass (R.A.M.) of an element is defined as the number of times one atom of the element is heavier than one atom of hydrogen. (b) Based on oxygen scale, relative atomic mass (R.A.M.) of an element is defined as the number of times one atom of the element is heavier than 1/16 of an oxygen atom. Answers 1B.2 R.A.M of an element = Mass 1 atom of the element 1/12 x mass 1 atom of carbon Answers 1B.3 (a) CuSO4 = 64 + 32 + (4 x 16) = 160 (b) C6 H12 O6 = (6 x 12) + (12 x 1) + (6 x 16) = 180 (c) Br2 = 80 + 80 = 160

If you failed to plan. you’ve planned to fail.

INPUT 1B.3 MOLE CONCEPT One mole (1 mole) of a substance is the quantity of the substance that contains the same number of particles as there are in 12.000 g of carbon -12. The number of particles is 6.02 x 1023. This means that one mole of any substance is the amount of substance that contains 6.02 x 1023 particles. This number, 6.02 x 1023 is called the Avogadro number or the Avogadro constant (NA) The particles in matter can be atoms, ions or molecules Do you know that each of these substances is 1 mole, even though the mass of the substances are different? Each substance contains the same number of atoms that is 6.02 × 1023. Pictures 2.1:Clockwise: sulphur, iron, copper, carbon and mercury. 1B.3.1 STATING MOLE AS NUMBER OF PARTICLES, MASS AND VOLUME OF GAS A. MOLE AS NUMBER OF PARTICLES

One mole of any substance is the quantity of the substance that contains 6 x 1023 particles. 1 mole = 6 x 1023 atoms, molecules or ions (depends on the kind of particles that is used to make the particular substance) Avogadro number = 6 x 1023 mole-1 Mole atom of an element: 26

(a)

1 mole atom of any element will contain the same number of atoms that is 6 x 1023 atoms.

Example 1B.3.1.1 1 mole Hydrogen (H)

contains the same number of atoms that is6 x 1023 atoms respectively.

1 mole Oxygen (O) 1 mole Copper (Cu)

(b)

One mole of an element is represented by the symbol of that particular element. To indicate quantities other than one, the numerical value is written before the symbol.

Example 1B.3.1.2

Magnesium metal 1 mole of magnesium atoms is represented by Mg, 2 mole of magnesium atoms are represented by 2Mg. Mole molecule of a covalent compound: (a)

1 mole molecule of any covalent compound will contain the same number of molecules, that is 6 x 1023 molecules.

Example 1B.3.1.3 1 mole hydrogen gas (H2) 1 mole oxygen gas (O2)

contains the same number of molecules (6 x 1023 molecules) respectively

1 mole ethanol (C2H5OH)

(b)

1 mole molecule of a covalent compound is represented by the chemical formula of the substance. To indicate quantities other than one, the numerical value is written before the chemical formula.

Example 1B.3.1.4

Water (H2O) 1 mole of water molecules is represented by H2O. 2 moles of water molecules are represented by 2H2O Relationship molecules

between

mole

atoms

and

mole

A molecule is made up of two or more atoms, so the number of moles of atoms in a compound depends on the number of atoms in each molecule. Example 1B.3.1.5 1 molecule of carbon dioxide (CO2) consists of 1 C atom + 2 O atoms = 3 atoms Hence 1 mole of CO2 molecules = 1 mole of C atoms + 2 mole O atoms = 3 moles of C and O atoms B. MOLE AS QUANTITY OF MATTER

Conversion of number of moles of a substance toits mass (a)

The mass of 1 mole atom of an element is the relative atomic mass (R.A. M.) of that element expressed in gram.(g) Mass of 1 mole atom of any element = Relative atomic mass (R.A. M.) of that element expressed in gram.

Example 1B.3.1.6 (i) Relative atomic mass of carbon = 12 Mass of 1 mole carbon atoms = 12 g Mass of 2 mole carbon atoms

= 2 x 12 = 24 g

Mass of 0.25 mole atom carbon = 0.25 x 12 = 3 g

(ii)

1 mole of an element contains 6 x 1023 particles, so 1 mole atom an element = 6 x 1023 atoms = Relative atomic mass of the element in grams

(iii)

The relationship between the number of moles of a substance and mass is therefore

Number of mole atom

mass in gram relative atomic mass

Table 1B.2 shows the number of atoms in and the mass of 1 mole for some elements. Table 1B.2 Number of atoms and mass of 1 mole of some elements Element

Relative Atomic Mass Hydrogen 1 Magnesium 24 Sulphur 32 Lead 207

Number of atoms in per mole 6 x 1023 atom H 6 x 1023 atom Mg 6 x 1023 atom S 6 x 1023 atom Pb

Mass of 1 mole atom 1 24 32 207

Mass and mole molecule (a)

The mass of 1 mole molecule of a compound is the relative molecular mass of the compound expressed in grams. Mass of 1 mole molecule=

Example 1B.3.1.7 Relative molecular mass of ammonia NH3 = N + 3H RMM = 14 + (3 x 1) = 17

relative molecular mass of the compound expressed in grams.

Therefore Mass of 1 mole ammonia molecules Mass of 2 mole ammonia molecules

(b)

(c)

= 17 g = 2 x 17 g = 34 g

1 mole of molecules contains 6 x 1023 molecules. The mass of 1 mole of molecules = mass of 6 x 1023 molecules = relative molecular mass in grams. Thus Number of moles =

______mass in grams____________ relative molecular mass in grams

Table 1B.3 shows the number of molecules and mass of one mole of molecules for some substances. Table 1B.3 Number of molecules and mass of one mole molecules of some substances. Substance Formula s (element or compoun d) Hydrogen H2

Relative Molecular Mass

Oxygen

Sodium chloride Sulphur dioxide

NaCl

58.5

SO2

Number of molecules in 1 mole

6 x 1023 molecules 6 x 1023 molecules 6 x 1023 molecules 6 x 1023 molecules

Mass of 1 mole molecul e in gram (g) 2 32 58.5 64

Number of moles of ions and mass of 1 mole of an ionic compound (a)

The total number of moles of ions in 1 mole of an ionic compound would depend on the ratio of positive ions (cations) to negative ions (anions) in the compound.

Example 1B.3.1.8

Formula of sodium chloride is NaCl (or Na+Cl-). The ratio of Na+ ions :Cl- ions is 1 : 1 1 mole of sodium chloride will contain 1 mole Na+ ions and 1 mole Cl- ions, a total of 2 moles of ions in 1 mole of NaCl. Na+ + Cl1 mole ion 1 mole ion

NaCl 1 mole sodium chloride

(b)

The mass of 1 mole of ionic compound is the relative formula mass (J.F.R.) of the ionic compound expressed in grams (g)

Number of mole =

mass in grams_____ relative formula mass

NUMBER OF MOLES OF GAS AND ITS VOLUME

According to Avogadro’s principle, different gases will have the same volume at the same temperature and pressure. 1 mole molecule of any gas contains the same number of molecules. Hence any gas will contain the same number of molecules at a particular temperature and pressure. The volume occupied by one mole molecule of any gas at a particular temperature and pressure is called the molar volume. In general,

1 molar volume

= 22.4 dm3 at standard temperature and pressure (s.t.p.) of any gas Or 3 = 24 dm at room temperature and pressure

Examples of calculations on the mole concept Example 1 Calculate the number of mole of molecules and mole of atoms in the following: (a) 40 g bromine gas (b) 0.3 g ethane gas C2H6 (Relative atomic mass: H = 1; C = 12; O = 16 and Br =80) Solution (a) Relative molecular mass of Br2 gas = 2 x 80 = 160 Mass 1 mole molecule Br2 gas = 160 g Number of mole molecule found in 40 g of Br2 gas = 40 / 160 = 0.25 mole molecule Br2 = 0.5 mole atom Br (because 1 molecule of Br 2 consists of 2 atoms of bromine). (b)

Relative molecular mass of ethane gas, C2H6 = 30 Mass of 1 mole molecule of ethane gas = 30 g Number of mole molecule found in 0.3 g ethane gas = 0.3/30 = 0.01 mole molecule ethane gas =0.01 x 8 mole atom C and H (because 1 molecule ethane is made up of 8 atoms C and H) = 0.08 mole atom.

Example 2 Calculate mass in gram and volume in dm3 at s. t. p. for (a) 2.0 mole ammonia gas, and (b) 0.125 mole sulphur dioxide gas. (Relative atomic mass: H = 1; N = 14; O = 16; S = 32; molar volume of a gas = 22.4 dm3 at s.t.p.)

Solution (a) Mass of 1 mole ammonia gas, NH3 = (14 + 3 x 1) g = 17 g 2 mole NH3 gas = 2 x 17 g = 34 g Volume of 2 mole NH3 gas = 2 x molar volume of gas = 2 x 22.4 dm3 at s. t. p. = 44.8 dm3 at s. t. p. (b)

Mass of 1 mole sulphur dioxide, SO2 gas = 32 + (2 x 16) = 64 g Mass of 0.125 mole SO2 gas = 0.125 x 64 = 8 g Volume 0.125 mole SO2 gas = 0.125 x molar volume of gas = 0.125 x 22.4 dm3 at s.t.p = 2.8 dm3 at s.t.p

1B.4 CHEMICAL FORMULAE 1B.4.1 DEFINITION OF CHEMICAL FORMULA

The Chemical Formula of an element shows the symbol representing that element and the number of atoms in the element. Example chemical formula of magnesium chemical formula of chlorine

Mg Cl2

Chemical Formula of a compound shows the symbols of all elements in the compound. It also gives the number of atoms of each element in one molecule of the compound. Example chemical formula of sulphuric acid H2SO4 In 1 molecule of sulphuric acid there are 2 hydrogen atoms, one sulphur atom and 4 oxygen atoms.

In general, a chemical formula indicates:

(a)

Qualitatively: types of element(s) present in the compound and

(b)

Quantitatively: relative numbers of atoms of each element present in the molecule, indicated by subscripts* (* number written at the bottom right hand corner of the symbol of the element)

For example, the chemical formula of oxygen gas (O2) shows

• qualitatively that the substance is made up of the element oxygen (O) and • quantitatively, that ✓ 1 molecule of oxygen is made up of 2 oxygen atoms. ✓ 1 mole of oxygen gas (O2) contains 2 moles of oxygen atoms. ✓ Mass of 1 mole of oxygen gas is the same as the relative molecular mass in grams and ✓ 1 molar volume of oxygen gas is 22.4 dm3 at s.t.p. or 24 dm3 at room temperature Similarly, with other compounds such as: (a)

Chemical Formula of glucose C6H12O6 shows i) 1 molecule of glucose is made up of 6 atoms of carbon (C)12 atoms of hydrogen (H) and 6 atoms of oxygen (O), ii) 1 mole molecule of glucose contains 6 mole atoms of C, 12 mole atoms of H and 6 mole atoms of O, iii) Mass of 1 mole molecule of glucose is the relative molecular mass in gram, that is 6C + 12H + 6O = (6 x 12) + (12 x 1) + (6 X16) 34

(b)

= 180 g Chemical Formula of ionic compound copper(II) nitrate Cu(NO3)2 shows

(i) it contains the element copper(Cu), nitrogen (N) and oxygen(O), (ii) it is made up of copper (II) ions Cu2+ and nitrate ions (NO3-) that combines in the ratio 1 : 2 (iii) 1 mole copper(II) nitrate contains 1 mole ion Cu2+ and 2 mole NO3-ions (iv) mass of 1 mole copper (II) nitrate is the relative formula mass in unit gram = Cu + 2 (N + 3O) = 64 + 2 (14 + (3 x 16) g = 188g 1B.4.2 EMPIRICAL FORMULA AND MOLECULAR FORMULA

Empirical formula is the chemical formula which shows the smallest ratio of the atoms of the elements that combine to form the compound. Molecular formula is the chemical formula which shows the actual number of the atoms of the elements that combine to form the compound

The difference between the two types of chemical formula can be seen in the following example Example : propene C3H6 (a) Molecular formula C3H6 shows 1 molecule of propene is made up of 3 atoms of C and 6 atoms of H (b)

Empirical formula CH2 (or C1H2) shows the simplest ratio of atoms number of atom C : H = 1 : 2 ratio C : H = 3: 6 =1:2 35

There are some compounds which have molecular formulae that are multiples of the empirical formula, that is molecular formula = (empirical formula)n where n = a whole number. Example 1B.4.1.1 Molecular formula glucose = (Empirical formula glucose)n, where n = 6

= (CH2O)6

= C6H12O6

If relative molecular mass of a substance is known, the molecular formula can be determined.

Example 1B.4.1.2 If empirical formula of benzene is CH and relative molecular mass = 78 (CH)n = 78 where n = a whole number nC+nH = 78 12n + n = 78 (R.A.M. : C = 12, H = 1) 13n = 78 n = 78 = 6 13

The molecular formula of benzene is (CH)6 that is C6H6. Not all compounds have empirical formula and molecular formula that are different. Some of the compounds especially ionic compounds such as magnesium oxide, ammonia, copper (II) sulphate and calcium carbonate have empirical formula and molecular formula that are the same. Empirical formula of a substance can be determined from (a) Percentage composition or (b) Mass of each component element in the compound

The following steps are used: o Ensure relative atomic mass of each element. o Divide % composition or mass composition of each element by its relative atomic mass to get number of moles. o Calculate simplest whole number ratio for each element and then determine the empirical formula

Example 1B.4.1.3 A nitrogen compound contains 82.3% nitrogen, and the rest of it is hydrogen. What is the empirical formula of the compound? (Relative atomic mass: H = 1; N = 14) Solution

% mass Mass in 100 g Number of mole atom = mass R.A.M Simplest Ratio

Element N 82.3 82.3 g 82.3 14 = 5.88

Element H 17.7 17.7 g 17.7 1 = 17.7

5.88 5.88 = 1

Empirical formula

17.7 5.88 = 3 N1H3 or NH3

Example 1B.4.1.4 Zinc reacts with iodine to form zinc iodide. In an experiment, 3.25 g of zinc produces 15.95 g of zinc iodide. Find the empirical formula of zinc iodide. (Relative atomic mass: Zn = 65; I = 127) Solution Element Zn 3.25 g

Mass Number of mole atom = Mass R.A.M. Simplest ratio

3.25 65

= 0.05

Element I 15.95 – 3.25 = 12.7 g 12.7 12.7 = 0.1

0.05 0.05 = 1

Empirical formula

0.1 0.05 = 2 Zn1I2 or ZnI2

1B.4.2 FORMULA OF IONIC COMPOUND

Ionic compound is made up of positive ions (cations) and negative ions (anions) in a certain ratio. Table 1B.4 Positive and negative ions Cations Charge Charge Charge Charge 2+ 3+ 4+ 1+ H+ Ca 2+ Al3+ Pb4+ Na+ Mn2+ Fe3+ Sn4+ + 2+ 3+ K Mg Cr Mn4+ Cu+ Hg2+ + Ag Zn2+ + NH4 Ni2+ Cu2+ Ba2+ Fe2+ Pb2+

Anions Charge Charge Charge 1232Cl CO3 PO432Br SO4 N3IO2OH HCO3NO3-

To determine chemical formula of ionic compounds 1. A chemical compound is always electrically neutral, the total positive charge of the cation must be equal to the total negative charge of the anion. 2. When writing the chemical formula of an ionic compound, sum of the positive charges must be equal to the sum of the negative charges.

Steps To Write The Chemical Formula Of An Ionic Compound

a) b) c)

1B.4.3

Steps Write the name of the ionic compound Write the formulae of the ions and their charges. Write the numerical charge of the cation next to the anion as a subscript and the numerical charge of the anion next to the cation as a subscript Write the chemical formula of the ionic compound without the charges

Example 1 Example 2 Calcium Zinc nitrate chloride Ca 2+ and Cl- Zn2+ and NO3Ca 2+ 1 CaCl2

Cl2

Zn2+ (NO3)1 2 Zn (NO3 )2

FORMULA OF HYDRATED SALT 1.

Hydrated salt is an ionic compound that has water of crystallization combined with its anhydrous salt.

2. The chemical formula of hydrated salt can be represented by writing the formula of anhydrous salt followed by nH2O where n is the number of moles of water of crystallization that has combined with one mole of its anhydrous salt. Example : Formula of hydrated copper sulphate is CuSO4.5H2O. This means there are 5 moles of water of crystallization (5H2O) that has combined with one mole of anhydrous copper (II) sulphate .

ACTIVITY2B

Fill in the empty spaces with suitable answers. Question 1B.4 _______________ of any substance contains _____________ particles . Question 1B.5 1 mol molecule of a covalent compound will contain_______________ molecules. Question 1B.6 Mass of one mole atom of an element = _______________________________. Question 1B.7 The chemical formula for magnesium carbonate is _______________ Question 1B.8 _____________ is a chemical formula that shows the actual number of elements that combine to form the compound.

atoms of the

RESPONSES

Answers 1B.4 One mole of any substance contains 6 x 1023 particles . Question 1B.5 1 mol molecule of a covalent compound will contain 6 x 1023 molecules. Question 1B.6 Mass of one mole atom of an element = = Relative atomic mass of the element in grams.

Question 1B.7 The chemical formula for magnesium carbonate is Mg CO3 Question 1B.8 Molecular Formula is a chemical formula that shows the actual number of atoms of the elements that combine to form the compound.

Hard Work is key to Success

SELF ASSESSMENT

Question 1 Relative atomic mass of element-elements P, Q, R and S is 23, 40, 35.5 and 16 respectively. P and Q are transition metals, R and S are non metals. Avogadro constant is 6 x 10 23 mole– 1. Calculate the number of atoms found in the following substances (a) 9.20 g element P (b) 16.00 g element Q (c) 28.4 g element R Question 2 Explain how you can calculate relative molecular mass and relative formula mass. Question 3 A compound is formed when 0.63g hydrogen combines with and 4.68g oxygen. What is the empirical formula of the compound? Question 4 A hydride of metal M( Relative atomic mass = 40) contains 0.25g hydrogen and 5g metal M. Find the empirical formula of the hydride ( H=1). If the relative molecular mass is 168, find the molecular formula. Question 5 Find the number of moles in the following substances: (a) 49.04 g sulphuric acid (b) 20 g C12 H22 O11

Question 6 Formula of hydrated copper (II) sulphate is CuSO4. 5 H2O. Find the number of moles for the following: (a) water of crystallization (b) anhydrous copper (II) sulphate

RESPONSES

Answers 1 (a) 2.4 x 10 23 atom P (b) 2.4 x 10 23 atom Q (c) 4 .8 x 10 24 atom R Answers 2 i. Find the chemical formula form the substance ii. Find the relative atomic mass for each of element in the chemical formula. iii. Find the sum of the relative atomic mass of all the elements as represented in the chemical formula. Answer 3

Number of mole Simple Ratio

Hydrogen 0.63/1 =0.63 0.63/ 0.2925 =2

Oxygen 4.68/16 =0.2925 0.2925/ 0.2925 =1

Atom M 5/40 = 1/8 1

Atom H 0.25/1 = 1/4 2

Formula: H2O Answer 4

Number of mole Simple Ratio

Empirical formula hydride: MH2 Molecular formula = (MH2)n = 168 42n = 168 n =4 Therefore the molecular formula is M4H8 Answers 5 (a) 49.04 g/ 98 = 0.5 mole sulphuric acid (b) 20 g/ 342 = 0.06 mole C12 H22 O11

Answers 6 (c)

5 moles water of crystallization (d) 1 mole of anhydrous copper (II) sulphate

This is not the end yet. Go on to the following chapters…..

UNIT 2A

PERIODIC TABLE

OBJECTIVES GENERAL OBJECTIVE: To know and understands the classification of elements in the periodic table. To know and understand the Bohr model of the atom, modern atomic structure, proton number and nucleon number. To understand and analyse the concept of isotopes and electron configuration. SPECIFIC OBJECTIVES: At the end of this unit you should be able to: ✓ Explain the classification of elements in the periodic table and state the importance of the periodic table of the elements in chemistry. ✓ Describe the concept of the basic Bohr atom. ✓ Explain modern atomic structure. ✓ Differentiate between proton number and nucleon number. ✓ Explain the definition of an isotope. ✓ Compare the properties of the isotopes of an element. ✓ Differentiate between isotopes that are stable and those that are unstable. ✓ Describe electron configuration. ✓ Give an example of the electron configuration in an atom. ✓ Identify a valence electron. ✓ Explain the connection between electron configuration and position of an element in the periodic table

INPUT 2A.1 THE CLASSIFICATION OF ELEMENTS IN THE PERIODIC TABLE The periodic table is a table that lists all known elements. All these elements are arranged in a particular order. The table holds approximately 109 elements which are arranged according to increasing proton number. These elements are arranged in specific groups and periods. Vertical columns are called groups and horizontal rows are called periods. Did you know that most of the elementsexist naturally and can be found in the earth’s crust, mantel and core? 99.7% of the earth‘s crust consists of oxygen,silica, aluminium, iron, calcium, magnesium,sodium, potassium,titanium, hydrogen, phosphorus and manganese.

Earth’s crust

Core

The mantel is made up of iron and nickel. The core consists of hot iron, carbon, silica and sulphur.

Figure 3.1: Cross-section of earth

2A.1.1 ELEMENTS IN THE MODERN PERIODIC TABLE

The periodic tablehas 18 groups, startingwith Group I, then Group II toGroup XVIII. Four of these groups have special names: Group I - Alkaline Metals Group II - AlkalineEarth Metals Group XVII - Halogens Group XVIII - NobelGases The elements in the same group always have the number of electrons in the outermost electron shell of their atoms. The modern periodic tablehas one block of elements right in the middle of the table. This block consists of the elements that have been called the Transition Elements. The periodic table has seven periods numbered from 1 to 7. Period 1 has only2 elements:hydrogenand helium. Period 2 and 3 each has 8 elements. The first three periods are called short periods. Periods 4 and 5 each have 18 elements while Period 6 actually has 32 elements. The elements with proton numbersfrom 58 to 71 are listed separately in the lower part of the periodic table. This seriesof elements is called the Lanthanide Series. Period 7 has 19 elements.The elements with proton numbers 90 to 103 are separately listed andare named the Actinide Series. 2A.1.2 THE IMPORTANCE OF THE PERIODIC TABLEIN CHEMISTRY

The periodic tableis very importantin chemistry because (a) we can study the elements and their compounds more systematically. (b) we can roughly predict the properties of unknown elements, if their positions in the periodic table are known. (c) we can also see the relations between elements in different groups. (d) the properties of the elements can be better understood and be more easily remembered.

Photo 3.1:Gold coins made from pure gold. Currently used as currency.

Table 2A.1 The Modern Periodic Table

2A.2 MODEL OF ATOMIC STRUCTURE 2A.2.1 THE BOHR MODEL OF THE ATOM

In the year 1913, Neils Bohr improved upon Rutherford’s atomic model. He suggestedthat: (a) electrons are not randomly distributed around the nucleus, but circle the nucleus within shells at various distances from the nucleus of an atom, and (b)

each shell represents a particular energy level of its electrons, and can be filled with only a specific number of electrons. (Figure 2A.1) electron shell

electron

nucleus

electron shell

Figure 2A.1 Bohr Model of the Atom

With the Bohr model, the atom can be pictured as being like a mini solar system.Until the year 1932, the atom of an element was thought to consist of protonsand electrons only. This assumption was corrected by Sir James Chadwickin 1932, when he discovered the neutron. The neutron has been found to be (a) electrically neutral; its motion is not deflected in a magnetic or electrical field, and (b) of similar mass as a proton, which is 1 unit. 51

With the discovery of the neutron, the problem of the difference between atomic mass values calculated according to the BohrRutherford model and actual,measured values was solved. 2A.2.2 MODERN THEORY OF ATOMIC STRUCTURE

The atom is now thought to be made up of : •

•

a tiny nucleus in the centre, which contains protonswhich are positively charged, and neutrons which are neutral. Almost all the mass of an atom is to be found concentrated the dense nucleus. Electrons in constant motion around the nucleus, within shells representing specific energy levels. In a neutral atom, the number of protons is equal to the number of electrons. Empty space between nucleus and electron shells nucleus of protons and neutrons Simplified:

electrons circling nucleus

Figure 2A.2

Modern model of atom

Subatomic Particles :

Scientists have found that an atom consists of three different subatomicparticles, the proton, electron and neutron.Table 2A.2 gives details on these particles.

Table 2A.2 Three subatomic particles in an atom Particle Symbol Mass (g) Relative Charge Location in Mass atom -24 Proton p 1.67 x 10 1 +1 In nucleus -24 Neutron n 1.67 x 10 1 0 In nucleus -28 Electron e 9.11 x 10 1/1840 -1 Constant motion around nucleus 2A.3 PROTON NUMBERANDNUCLEON NUMBER The Proton number (Z) of an element is the total number of protons in the nucleus of one atom of that element. The Nucleon number (A) of an element is the total number of protons and neutrons inyang terdapat in the nucleus of one atom of that element.

Nucleon number= Number of protons + Number of Neutrons (A) (p) (n)

Since the Proton number (Z) = Number of protons (p), Nucleon number (A) =Proton number (Z) + Number of neutrons (n)

And so

A =

Z + n

Generally, the proton number and nucleon number of a particular element can be shown to gether with the symbol for the element, as shown in the following: 𝐴 𝑍X

This shows that each atom of element X has Z protons (and alsoZ electrons) and (A – Z) neutrons. The relationship between nucleon number and proton number is shown in Table 3.3

Table 2A.3 Relationship betweennucleon number and proton number Element Symbol Mass Atomic Number of A ( Z X) No. No. (Z) Protons Electron Neutron s s (A) 1 Hydrogen 1 1 1 1 0 1H 4 Helium 4 2 2 2 2 2 He 7 Lithium 7 3 3 3 4 3 Li 12 Carbon 12 6 6 6 6 6C Fluorin

19 9

Neon

20 10

2A.4 THE ISOTOPE The atoms of an element may have the same number of protons, but different number of neutrons. This causes the atoms to differ slightly in their mass. Such atoms, belonging to the same element but having different mass are called isotopes. 2A.4.1 THE DEFINITION OF AN ISOTOPE

An isotopecan be defined as one of two or more atoms that have the same proton number (that is, they belong to the same element) but different numbers of neutrons.

As an example, hydrogenhas three isotopes: 1 1

light hydrogen

2 1

3 1

deuterium

tritium

6e 6p 5n

Z : A : Mass of isotope : Symbol : C 6

THREE ISOTOPES OFHYDROGEN 6e 6p 6n

6 6 11 11.011 433 11 C 6

6e 6p 8n

6 12 12.000 000 12

14 14.003 242 14 C 6

THREE ISOTOPES OFCARBON (Other isotopes of carbonare known to exist.)

2A.4.2 COMPARISON OF THE PROPERTIES OF ISOTOPES

Table 2A.4 shows a comparison of isotopes of an element. Table 2A.4 Comparison of isotopes Similarities (a) Number of protons (b) Proton number (Z) (b) Number and configuration of electrons (c) Chemical properties

2A.4.1

Differences (a) Number of neutrons (b) Nucleon number (A) (c) mass (relative isotope mass) (d) Physical properties

STABLE AND UNSTABLE ISOTOPES

Of all the known isotopes, only a small proportion is stable.The rest which are unstable undergo radioactive decay.Unstable isotopes are radioactive and emit highly dangerous radiation. However these isotopes have been used in diversefields such as medicine, agriculture, food preservation and energy generation.

2A.5 ELECTRON CONFIGURATION 2A.5.1 ELECTRON CONFIGURATION IN AN ATOM

According to the Bohr model of the atom, electrons move around the nucleus according to specific and discrete energy levels, at specific distances from the nucleus. Each energy level around the nucleus is referred to as a shell and given a symbol such as K, L, M, N …. to represent the first, second, third, fourth shell … in order of increasing distance from the nucleus. The nearer an electron shell is to the nucleus, i. the stronger is the electrostatic force of attraction between the electronsand the protons in the nucleus, and so ii. the more stableare the electrons or the lower are their energy level. On the contrary, the further the shell is from the nucleus, the higher is the energy level of its electrons. A particular shell or energy level can be occupied by a specific number of electrons only, for example the first shell can be occupied by only two electrons, the second shell, eight electronsandso on. Generally, the maximum number of electrons that be placed in a particular shell follows the formula of 2n2wheren is the shell number: 1, 2, 3, 4 …. as summarised in Table 2A.5 and Figure 2A.3. Table 2A.5 The maximum number of electrons in each electron shell Shell No., n 1 2 3 4 Symbol for K L M N shell 2n2 2 x 12 2 x 22 2 x 32 2 x 42 Number of 2 8 18 32 Electrons Energy level Lowest increasing

nucleus

electron shells (or energy levels)

Figure 2A.3Total number of electrons in different energy shellsof an atom In the arrangement of electrons of an atom, the first shell to be filled is the one with the lowest possible energy, in other words, the shell that is nearest to the nucleus. As an illustration, each sodium atom (proton number 11) has 11 electrons. The first shell is therefore filled with 2 electrons, the second shell with 8 electronsandthe third shell with the last remaining electron. The arrangement of electrons in an atom can be pictured as a container to represent each electron shell surrounding the nucleus, as shown in Figure 2A.4. It can also be written in the form of numbers within brackets, which then represents an atom’s electron configuration.

Figure 2A.4 Electron configuration in the atoms of several elements

2A.5.2 VALENCE ELECTRONS

The outermost shell which has electrons and which is also the furthest from the nucleus, is called the valence shell.

The electrons in this shell are known as valence electrons.Because they are furthest from the nucleus and electrostatic forces of attraction are correspondingly weakest,valence electrons are easily released from the atom. These are the electrons that are involved in the formation and breaking of bonds in chemical reactions. As such the number of valence electrons in the atom determines the chemical properties of that element. Atomswhich have the maximum number of valenceelectrons, for instance 2 valence electrons in the helium (He) atom and 8 valence electrons in the atoms of the other noble gases, are especially stable chemically and show very low reactivity. Different elements with atoms that have the same number of valence electrons (1 to 8) demonstrate similar chemical properties. Having the maximum of 2 valence electrons in the singular shell of the helium (He) atom causes it to be highly stable chemically. This is referred to as the duplet configuration.

Those atoms that have 8 valence electrons in the outermost shell are said to have the octet configuration, as for all the noble gases other than helium. (Figure 3.5)

Figure 2A.5

Octet electron configuration in atoms of noble gases 58

2A.5.3 ELECTRON CONFIGURATION ANDTHE PERIODIC TABLE

Table 2A.5 Electron configuration of some elements in the periodictable Group I Element Li Electron 2.1 configuration Element Na Electron 2.8.1 configuration

II Be 2.2

XIII B 2.3

XIV C 2.4

XV N 2.5

XVI O 2.6

XVII F 2.7

XVIII Ne 2.8

Mg 2.8.2

Al 2.8.3

Si 2.8.4

P 2.8.5

S 2.8.6

Cl 2.8.7

Ar 2.8.8

All elements in a group share similar chemical properties. These elements also have the same number of electrons in the outermost orbit. This means that the chemical properties of an element are influenced by the number of electrons in the outermost orbit of its atoms. Elements in the periodic table are arranged by their proton number in increasing order. The proton number is the number of protons found in one particle. An atom, since it is electrically neutral, always has the same number of electrons and protons. And so it is correct to say that the number of electrons in each atom of the elements increasefrom left to right across a period in the periodic table. With the exception of groups III – XII (transition metals), the number within the unit's place of the group number identifies the number of electrons in the outermost orbit of an atom in that group. The number of electrons in the inner orbits does not influence the chemical properties of an element. The inner electrons usually affect only the physical properties or reactivity of the element. The valencyof an element is the number of electronsthat can be received or released by one atom of the element, to achieve thestable electron configuration of a noble gas.

(a) Example : Table 2A.6 : Valency of several elements in Groups I, II andXIII Group I II XIII Element Li Be Al Electron configuration 2.1 2.2 2.8.3 Number of electrons that 1e 2e 3e can be released New electron 2 2 2.8 configuration Comparison with electron Helium Helium Neon configuration of a noble 2 2 2.8 gas Lithium easily loses 1 electron to achieve the electron configuration of helium. As such, Li, a Group I element, has a valency of 1. Similarly, beryllium in Group II easily loses 2 electrons and achieves the electron configuration of helium. Beryllium has a valency of 2. Aluminium in group XIII tends to release 3 electrons, thereby achieving the electron configuration of neon, and so aluminium is said to have a valency of 3.

(b) Table 2A.7: Valency of several elements in Group XVI and XVII Group Element Electron configuration Number of electronsreceived New electron configuration Comparison with electron configuration of noble gases

XVI O 2.6 2e

XVII Cl 2.8.7 1e

2.8

2.8.8

Neon 2.8

Argon 2.8.8

An oxygen atom tends to receive 2e,to achieve the electron configuration of neon, whereas a chlorine atom receives 1e to achieve the structure of argon. Generally, for elements inGroups I and II, valency and group number have the same values. (Table 3.8) Elements in Groups XIII to XVIII have valencies that are the same as the number in the unit place of the group number. Table 2A.8 : Valency and group number forGroups I, II andXIII Group I II XIII

Valensi 1 2 3

Element Li, Na, K Be, Mg, Ca B, Al, Ga

Ion Li+, Na+, K+ Be2+, Mg2+, Ca2+ B3+, Al3+, Ga3+

Therefore, elements in Group I, II and XIII tend to release electrons, forming positive ions. Generally for Groups XIV to XVII, elements have valency values that are equal to 8 – (group number – 10). (Table 2a.9)

Table 2A.9: Valency and group number for Groups XIV to XVII Group

Element

Valency

XIV

C Si N P O S Cl Br

8 - (14 – 10) = 4 8 - (14 – 10) = 4 8 - (15 – 10) = 3 8 - (15 – 10) = 3 8 - (16 – 10) = 2 8 - (16 – 10) = 2 8 - (17 – 10) = 1 8 - (17 – 10) = 1

XV XVI XVII

Ion orCompound formed CH4 SiH4 NH3 PH3 O2S2CBr-

The valency of elements in Groups XVI and XVII is written in negative form. For example valency of oxygen = -2 and valency of chlorine = -1

UNIT2B

PERIODIC TABLE

OBJECTIVES GENERAL OBJECTIVES : To know and understand the properties of groups and rows in the Periodic Table of Elements. SPECIFIC OBJECTIVES : At the end of this unit students should be able to: ✓ Explain the relationship between electron configuration and the position of the element in the Periodic Table. ✓ State the physical and chemical properties of Group I. ✓ State the physical and chemical properties of Group XVII. ✓ State the physical properties of Group XVIII. ✓ List the uses of noble gases.

✓ Explain the changes in properties of the elements as we go across the period ✓ List the uses of semi metals in microelectronic industry.

INPUT 2B.1 PROPERTIES OF GROUPS IN THE PERIODICTABLE OF ELEMENTS 2B.1.1 ELECTRON CONFIGURATION AND GROUP NUMBER

Elements in the Periodic Table are arranged according to their proton numbers. For atoms, which are electrically neutral because they possess equal numbers of electrons and protons, the proton number of an element also shows the number of electrons in an atom of that element. Table 2B.1 shows the electron arrangement of the elements with proton numbers 1 to 20 in the Periodic Table. Table 2B.1Electron arrangement of the first 20 elements in the Periodic Table Period 1 2 3 4

Group XIV XV

I II XIII XVI XVII XVIII H He 1 2 Li Be B C N O F Ne 2.1 2.2 2.3 2.4 2.5 2.6 2.7 2.8 Na Mg Al Si P S Cl Ar 2.8.1 2.8.2 2.8.3 2.8.4 2.8.5 2.8.6 2.8.7 2.8.8 K Ca 2.8.8.1 2.8.8.2

Table 2B.1 clearly indicates that the group number of an element is determined by the number of valence electrons inan atom of the element. Example: Hydrogen valence electron is 1 group is 1 Sodium valence electron is 1 group is 1 Fluorine valence electrons is 7 group is 17

Table 2B.2 Relationship between number of valence electrons and group number Number of valence electrons Group

8 (except helium)

Based on Table 2B.1, the period number of an element is determined by the number of shells occupied with electrons in an atom of that element. Example: Hydrogen, number of electron filled shells is 1, period number is 1. Lithium,number of electron filled shells is 2, period number is 2. Sodium, number of electron filled shells is 3, period number is 3. Table 2B.3 Relationship between number of shells and period number Number of electron filled shells Period

Hence, the period number of an element is equal to the number of electron filled shells. Elements with the same number of valence electrons are placed in the same group. Elements in the same group will exhibit similar chemical properties. However the physical properties and the reactivity of the elements in a group change gradually down the group as in Figure 2B.1.

Group XIV

Reactivity increases

Metallic nature increases

Group XVII

Reactivity decreases

electronegativity decreases atomic size increases

electrpositivity increases atomic size increases

Group I

Figure 2B.1Trends in reactivity and physical properties of elements in different groups.

The reactivity and physical properties change gradually because of the increase in electrons as we go down the group. Valency of an element refers to the number of electrons that is lost or gained by the element to achieve duplet (same as helium) or octet electron arrangement (as noble gases) that is stable. In general, (a) elements in Group I, II and XIII have a valency that is the same as its group number. These elements have a tendency to lose its valence electrons to form positive ions with a positive charge same as its group number. (b) elements in Group XIV, XV, XVI and XVII have a valency that is equal to 18 minus the Group number. These elements have a tendency to receive the same number of electrons as their valency to form negative ions with a charge of -4, -3, -2 and-1 respectively.

Do you know the process of electron movement to a lower shell produces light? This concept is used in the making of coloured advertisement lights and street lights. Example : mercury is used to produce blue light in lamps, neon is used to orange light while sodium produces yellow light.

Picture photo 4.1: Street light as an application of atomic emission.

2B.1.2 PHYSICAL AND CHEMICAL PROPERTIESOF GROUP 1 ELEMENTS (ALKALI METALS)

Group1 elements are lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr). These elements have one valence electron in each of the atoms. Group 1 elements are also known as the alkali metals. PHYSICAL PROPERTIES

1) All elements in Group1 are solids at room temperature. 2) Like all other metals, elementsinGroup1 (a) Are grey in colour and have shiny silvery surfaces when freshly cut. (b) are good conductors of heat and electricity.

3) Unlike all other metals, elements in Group 1 (a) are soft because the electrostatic force between the valence electron and the nucleus is weak. (b) have low melting and boiling points because the attractive force between atoms are weak so less heat is required to overcome the force. (c) have low densities because the atoms have larger atomic radius compared to other metals having the same number of filled electron shells.

CHANGES IN PHYSICAL PROPERTIES OF ALKALI METALS WITH INCREASE IN PROTON NUMBER. Physical properties of alkali metals change gradually with increase in proton number. When going down the group: (a) atomic radius and atomic size increases because the number of shells occupied with electrons increases down the group. (b) Densities increase because the increase in atomic mass is bigger than the increase in volume down the group. On the other hand, (c) hardness decreases because when the atomic size increases the attractive force between atoms decreases. (d) melting and boiling points decreases because when the atomic size increases the attractive force between atoms decreases. Less heat energy is required to overcome the attractive force. Table 2B.4shows electron arrangement and some physical properties of alkali metals in Group 1. Table 2B.4

Electron arrangement and some physical properties of alkali metals in Group 1

Elements

Proton Number

Nucleon Number.

Electron Arrangement

Lithium, Li Sodium, Na Potassium, K Rubidium, Rb Caesium, Cs

3 11 19 37 55

7 23 39 85 133

2.1 2.8.1 2.8.8.1 2.8.18.8.1 2.8.18.18.8.1

Density (g cm3 ) 0.534 0.970 0.860 1.530 1.90

Melting point (C) 186 98 64 39 29

Boiling point (C) 1609 880 760 700 670

Atomic radius (nm) 0.133 0.157 0.203 0.216 0.235

CHEMICAL PROPERTIES 1) All alkali metals in Group I have one valence electrons each in

its atom. These elements have a tendency to lose its valence electron to form a univalent ion with a charge of +1, as shown below. Li - 1e 2.1 Na - 1e 2.8.1

Li+ 2 Na+ 2.8

K - 1e 2.8.8.1

K+ 2.8.8

In general, M – 1e

M+

2) Hence, alkali metals are

(a) Elements that are very electropositive and very reactive in its reactions with electronegative elements such as halogen. (b) strong reducing agents. REACTONS OF ALKALI METALS 1. Reactions with water

Alkali metals react vigorously with water to produce a strong alkali and hydrogen gas. 2Li(s) + 2H2O(l)

2LiOH(aq) + H2(g)

2Na(s) + 2H2O(l)

2NaOH(aq) + H2(g)

2K(s)+ 2H2O(l)

2KOH(aq) + H2(g)

Or in general, 2M(s)+ 2H2O(l) alkali metal water

2MOH(aq) + H2(g) alkali hydrogen

2. Reaction with chlorine gas

Alkali metals react vigorously with chlorine gas to form white metal chloride. 2Li(s) + Cl2(g)

2LiCl(s)

or in general 2M(s) + Cl2(g)

2MCl(s)

Figure 2B.2 Reaction of alkali metals with chlorine gas burning lithium chlorine e gas

3. Reaction with oxygen gas

When an alkali metal burns in oxygen, it forms a white solid metal oxide. 4Na(s) + O2(g) 4K(s) + O2(g)

2Na2O(s) 2K2O(s)

or in general 4M(s) + O2(g)

2M2O(s)

The white solid metal oxides dissolve in water to produce metal hydroxide solutions which are alkaline. K2O(s) + H2O(l) or in general M2O(s) + H2O(l)

2KOH(aq) 2MOH(aq)

Similarly alkali metals react vigorously with electronegative elements such as Cl2, and O2or electronegative radicals such as OH-, NO3-, SO42-and CO32-to form compounds as shown in Table 2B.5. Table 2B.5 Formulae of ionic compounds of alkali metals Alkali Formula metal of ion M M+ Li Li+ Na Na+ K K+ General formula of compound

Chloride ClLiCl NaCl KCl MCl

Formula of compounds of alkali metals Oxide Hydroxide Carbonate Sulphate O2OHCO32SO42Li2O LiOH Li2CO3 Li2SO4 Na2O NaOH Na2CO3 Na2SO4 K2O KOH K2CO3 K2SO4 M2O MOH M2CO3 M2S04

Nitrate NO3LiNO3 NaNO3 KNO3 MNO3

Compounds of other alkali metals can be predicted.

Example, rubidium (Rb) which is placed under Potassium in Group I is expected to form : i) a chloride with a formula RbCl, ii) an oxide with a formula Rb2O, iii) a hydroxide with a formula RbOH, iv) a carbonate with a formula Rb2CO3, v) a nitrate with a formula RbNO3, vi) a sulphate with a formula Rb2SO4 Many ionic compounds of alkali metals are stable and are not decomposed by heat. However there are some ionic compounds that decompose partially when heated. 70

Li2CO3(s)

Li2O(s)

CO2(g)

2LiOH(s)

Li2O(s)

H2O(l)

2LiNO3(s)

2LiNO2(s)

O2(g)

2NaNO3(s)

2NaNO2(s)

2KNO3(s)

2KNO2(s)

O2(g) +

O2(g)

Change in electropositivity and reactivity of alkali metals: i)

ii.)

Alkali metals are very electropositive and very reactive because these metals have a high tendency to lose its valence electron from its outer shell to achieve stable electron arrangement (duplet/ octet arrangement). Li 2.1

Li+ 2

Na 2.8.1

Na+ 2.8

+ 1e

When going down Group I, the atomic size of alkali metals increases. Hence the single valence electron becomes further away from the nucleus. This causes the force attraction between the nucleus and the valence electron to become weaker. Thus the valence electron can be released more easily down the group. As a result, the electropositivity and reactivity of alkali metals increases down the group as shown in Table 2B.6

Table 2B.6: Changes in electropositivity and reactivity of elements in Group1 Elements

Li Na K Rb Cs Fr

Electron configuration 1 2 3 4 5 6 7 2 1 2 8 1 2 8 8 1 2 8 18 8 1 2 8 18 18 8 1 2 8 18 18 32 8 1

Atomic radius

increases

Tendency to lose valence electron(electropositivity)

Reactivity

increases

2B.1.3 PHYSICAL AND CHEMICAL PROPERTIES OF XVII ELEMENTS (HALOGEN)

Group XVII elements are fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At). These elements are known as the halogens. These elements have 7 valence electrons each which make them unstable. All halogens exist as diatomic covalent molecules with molecular formulae of F2,Cl2,Br2and I2. PHYSICAL PROPERTIESOF HALOGENS

All halogens have its own colour. The colour of these elements becomes darker down the group, for example fluorine is pale yellow, chlorine is greenish yellow, bromine is reddish brown and iodine is purplish black.

All halogens do not conduct electricity and heat.

All halogens in gaseous form are pungent and are poisonous.

CHANGES IN PHYSICAL PROPERTIES OF HALOGENS WITH INCREASE IN PROTON NUMBER

When coming down Group XVII, atomic size and the attractive force between the molecules increases gradually which causes the following: (a) (b) (c) (d)

state of matter, The first two elements F2and Cl2are gases, Br2is a liquid and At2is a solid. melting points and boiling points increase, density increases, but solubility in water decreases.

Table 2B.7 Electron configuration and properties of halogens Elements Fluorine, F Chlorine, Cl Bromine, Br Iodine, I Astatine, At

Electron configuration 1 2 3 4 5 6 2 7 2 8 7 2 8 18 7 2 8 18 18 7 2 8 18 32 18 7

Atomic radius

Oxidising ability

increases

decreases

Electronegativity

decreases

Reactivity

decreases

CHEMICAL PROPERTIESOF HALOGEN 1)

All halogens have 7 valence electrons in its atom During a chemical reaction, a halogen atom will gain one electron to form univalent negative ion to attain the stable octet in its electron arrangement, as shown below: Cl 2.8.7

Cl 2.8.8

Br 2.8.18.7

Br 2.8.18.18.8

I 2.8.18.18.7

+ 1e 2.8.18.18.8

–

All halogens have high electronegativities. Electronegativity is a tendency of an element to attract electrons. The electronegativity decreases down the group from chlorine to iodine. As the atomic radius increases down the group the force of attraction between the nucleus and the electrons become weaker. This causes the reactivity of the halogen to decrease down the group.

Halogens are good oxidising agents because each of the atoms gains one electron to achieve stable octet electron arrangement. Oxidising agents are substances that accept electrons (electron acceptors).As reactivity decreases down Group XVII, the oxidising ability also decreases down the group.

REACTIONSOF HALOGENS 1) Reaction with alkali metals

All halogen have high electronegativities. They react vigorously with electropositive elements such as alkali metals (Li, Na, K)to form halide compounds such as sodium fluoride, sodium chloride, sodium bromide and sodium iodide, as represented below: 2Na(s) + Cl2(g)

2NaCl(s)

or in general, 73

2M(s) + X2(g)

2MX(s)

2) Reaction with other metals

Example iron (Fe): When a halogen reacts with heated iron wool (iron), iron halide is formed. Figure 4.3 shows the reaction of chlorine with heated iron which forms a product, Iron (III) chloride 2Fe(s) + 3Br2(g)

2FeBr3(s)

or in general, 2Fe(s) + 3X2(g)

2FeX3(s) iron wool

reaction of bromine with iron

iron woolheated

heated

bromine or iodine

reaction of chlorine with iron heated

Figure 2B.3 Reactionof halogen with iron 3) Reaction with water

When halogen is passed through water, a reaction takes place to form an acidic solution as shown below: X2(g) + H2O(l) HX(aq) + HOX(aq) The solution has a bleaching effect. Example: Chlorine reacts with water to form hydrochloric acid (HCl) and hypochlorous acid (HOCl). The solution is called as chlorine water which actually consists of two types of acids,

hypochlorous and hydrochloric acids. Hypochlorous acid easily dissociate to form oxygen gas and hydrochloric acid. 2HOCl(aq)

2HCl(aq) + O2(g)

The same reaction takes place when other halogens dissolve in water. Br2(g) + H2O(l)

HBr(aq)

HOBr(aq)

Brominewarm waterhydrobromic acidhypobromousacid

I2(g)

H2O(l)

Iodinewarm waterhydroiodicacid

HI(aq) +

HOI(aq)

hypoiodus acid

4) Reaction with strong alkali

Halogen reacts with strong alkali such as sodium hydroxide (NaOH) to form two types of salts. Example When chlorineis passed through sodium hydroxide solution, sodium chloride (NaCl) and sodium chlorate were formed, as shown in the following equations: Cl2(g) + 2NaOH(aq)

NaCl(aq) + NaOCl(aq)+H2O(l)

Chlorine sodium hydroxide

sodium chloride

sodium chlorate

Sodium chlorate has bleaching properties and is usually used to wash clothes. 5) Reaction with non-metals Halogens react with non-metals compounds.

form

covalent

Example 1:Phosphorus When a little amount of white phosphorus is placed in a gas jar filled with chlorine gas, it burns immediately to produce white fumes, phosphorus trichloride (PCl3). 2P(s) + 3Cl2(g)

2PCl3(g)

Bromine and iodine react in the same way. 2P(s) + 3Br2(g)

2PBr3(g) phosphorus tribromide

2P(s) + 3I2(g)

2PI3(g) phosphorus triiodide

Example 2: Hydrogen Halogens react with hydrogen gas at certain pressure to form a hydrogen halide covalent compound (HX where X = halogen). Hydrogen halide dissolves in water to form an acidic solution containing H3O+and X-in it. H2(g) + Cl2(g

2HCl(g), H3O+(aq)+ Cl-(aq)

HCl(g) + H2O(l)

H2(g) + Br2(g)

2HBr(g), H3O+(aq) + Br-(aq)

HBr(g) + H2O(l) H2(g) +

I2(g)

HI(g) + H2O(l)

2HI(g), H3O+ (aq) + I-(aq)

6) Displacement reaction with halogens Electronegativity and reactivity of halogens as an Oxidising agent decreases down the group from fluorine to iodine. F > Cl > Br >

I >

Electronegativity and reactivity decreases

The more reactive halogens will displace the less reactive halogens from its halide solution. Example: Chlorine will displace bromine and iodine from its halide solution Cl2(aq) + 2Br-(aq)

2Cl-(aq) + Br2(aq)

Cl2(aq) + 2I-(aq)

2Cl-(aq) + I2(aq)

Bromine which is less reactive than chlorine but more reactive than iodine will not be able to displace chlorine but will be able to displace iodine from its halide solution. Br2(aq) + 2I-(aq)

2Br-(aq) + I2(aq) 76

2B.1.4

PHYSICAL PROPERTIES OF GROUP XVIII (INERT GASES)

The inert gases is made up of helium(He),neon(Ne), argon(Ar), krypton (Kr), xenon (Xe) and radon (Rn) in Group XVIII of the Periodic Table. Inert gases are also known as noble gases. All inert gases are chemically inert so they are unreactive. They do not react with any element in normal conditions. Inert gases are unreactive because they all have filled electron shells, duplet electron arrangement (as in helium) or octet electron arrangement (as in other inert gases). The stable electron arrangement explains why (a) all inert gases exist as monoatomic gases (He, Ne, Ar, Kr, Xe or Rn) and not as diatomic molecules such as H2, O2, N2 and Cl2. (b) inert gases do not receive, lose or share electrons with other elements, thus they do react with other elements under normal conditions. PROPERTIES OF INERT GASES All inert gases are non-metals which have the following properties: (a) (b) (c) (d)

low melting and boiling points because of the Van der Waals’ forces of attraction is very weak, low densities insoluble in water, do not conduct electricity or heat

CHANGES IN PHYSICAL PROPERTIES WITH INCREASE IN PROTON NUMBER The physical properties of inert gases change gradually with increase in proton number. When coming down Group XVIII, (a) (b)

(c)

radii and atomic size of inert gases increase because the number of electron filled shells increases down the group. density increases because with increase in proton number, mass of the atom increases more rapidly compared to increase in atomic size. melting and boiling points increase gradually as shown in Table 2B.8

Table 2B.8 Inert gases and their properties. Elements

Symbol

Helium Neon Argon Krypton Xenon Radon

He Ne Ar Kr Xe Rn

Electron Configuration 1 2 3 4 5 6 2 2 8 2 8 8 2 8 18 8 2 8 18 18 8 2 8 18 32 18 8

Melting point (C) -272 -249 -189 -157 -112 -71

Boiling point (C) -269 -246 -186 -153 -108 -62

% volume in water 0.0005 0.0018 0.9300 0.0001 0.00001 -

2B.1.5 USES OF INERT GASES

All inert gases exist in small quantities, about 1% in the atmosphere. Most of the inert gases can be obtained by fractional distillation of liquid air. The uses of inert gases are as shown in Table 2B.9 Table 2B.9 Uses of inert gases Inert Gas

Uses

Helium

(a) used to fill meteorological balloons, advertisement balloons and airships. Helium is very light and nonflammable, hydrogen is light but it is flammable (b) a mixture of helium and oxygen is used in oxygen tanks by divers (c) used in welding process to prevent oxidation of the metal by atmospheric air.

Neon

(a) used to fill advertising light bulbs and landing light bulbs at the airport. Neon glows with a reddish – orange light (b)used in electronic devices such as voltage stabilisers Argon (a) used in electric bulbs to prevent oxidation of hot tungsten filament. (b) to supply an inert atmosphere for welding Krypton used to fill high speed photographic flash lamps Xenon Radon

(a) to fill lights used in radars and projectors (b) used in electron tubes and stroboscopic lamps Used to treat cancer

2B.2 PROPERTIES OF A PERIOD IN THE PERIODIC TABLEOF ELEMENTS 2B.2.1 DEFINITION OF PERIOD

A Period is a horizontal row of elements in the Periodic Table. There are a total of seven periods, Period 1 to Period 7 (Figure 2B.4) period 1 period 2 period 3 period 4 period 5 period 6 period 7

transistion series 1 transistion series 1I transistion series 1II

lanthanide series actinide series

Semimetals

Figure 2B.4 Position of periods and other series in the Periodic Table Between Group II and XIII, there are a number of horizontal series of elements which are given special names. (a) Transistion elements series I, II and III is made up of metals that are harder and stronger compared to metals in Group I and II. (b) Lantanide element series with proton number 58 to 71. (c) Actinide 103.

elements

series

with

proton

number

When going across any period from left to right, (a) The proton number increases by one unit steps. (b) The first element has one valence electron, the valence electron then increases by one across the period till a maximum number in the outer shell is achieved. Table 2B.10 shows: (a) Period number of an element is the number of shells occupied with electrons in an atom of the element. For example, the atoms of Period 3 elements have three shells occupied with electrons. (b) The total number of elements found in a period depends on the number of electrons that can be filled in the outer shell of the atoms of the elements. For example there are 8 elements in Period 2 because the outer shell can be filled with a maximum of 8 electrons. Period 4 has 18 elements because the outer most shell can be filled with a maximum of 18 electrons.

Table 2B.10: Changes in physical properties of elements in Period 3. Elements Na Physical properties Electron Arrangement Number of valence electron

2.8.1 1

2.8.2 2

Atomic Size Electronegativity

Metal conductor

Physical States Melting point Boiling point

2B.2.2

2.8.3

2.8.4

2.8.5

2.8.6

Decreases Decreases Increases Semi- metal

Electropositivity Metallic properties Electrical Conductivity

semiconductor

solid Increases Increases

2.8.7

2.8.8

Non metal Non conductor gas Decreases Decreases

CHANGES IN PROPERTIES OF ELEMENTS ACROSS A PERIOD

Electron arrangement and trends of changes in the physical properties of elements across Period 3 is shown in Table 2B.10 When going across a period from left to right, (a)

Atomic size decreases. The proton number increases across period 3 from left to right. This causes the positive charge of the nucleus to increase across the period. Hence the attraction of the nucleus on the electrons increases, which pulls the electrons closer to the nucleus. Therefore the atomic size decreases across period 3.

(b)

Electropositivity decreases. The attractive force of attraction of the nucleus on the valence electron increases As a result it becomes extremely difficult for the atom to lose an electron to form a positive ion.

(c)

Electronegativity increases. The atomic size across the period decreases. The proton number increases across the period. This causes an increase in the strength of the 82

nucleus to attract electrons to increase. Therefore, the electronegativity increases.

2B.1.3

(d)

Metallic properties decreases. The elements change from metals to metalloid and finally to non metals across the period. The elements on the left of the period are metals such as sodium, magnesium and aluminum. Silicon has some metallic and some non metallic properties. It is called a metalloid or semimetal. The elements on the right of the period are non metals.(phosphorus, sulphur, chlorine, and argon)

(e)

Electrical conductivity decreases from left to right. Metals are good conductors of electricity, semi-conductors are weak conductors but their conductivity increases with the addition of boron or phosphorus. Non-metals are nonconductors.

(f)

Melting point and boiling point. They increase from the left but begin to decrease from the middle of the period.

(g)

Physical states. Elements change from solids to liquids, to gases when going across a period. Metals on the left are usually solids while non-metals on the right are usually gases. USES OF SEMI METALS (METALLOID ELEMENTS)IN MICROELECTRIC INDUSTRY

Semi metals are weak conductors of electricity but their conductivity increases when its temperature increases. Semi metals such as silicon (period 3) and germanium (period 4) are used widely as semiconductors in electronic components in computers such as transistors and diodes. When mixed with a small amount of foreign substances such as boron or phosphorus , silicon allows the flow of electrons in one direction but not in the opposite direction.

ACTIVITY 2B

State TRUE or FALSE for the statements given below regarding their changes in the physical properties: Question 2B.1 Changes in the physical properties of alkali metals: (a) Density decreases because of increase in proton number. ---------(b) Melting point and boiling point increases because of increase in atomic size. ---------(c) Hardness decreases because of increase in atomic size. ---------(d) Radius and atomic size decreases in proton number. ---------Question 2B.2 Changes in the physical properties of halogens. (a) Melting point and boiling point increases with increase in proton number. --------(b) Solubility increases in water with increase in proton number. --------(c) Density decreases with increase in proton number. --------(d) Physical states of halogens changes from gas for fluorine to solid for iodine. --------Question 2B.3 Changes in the physical properties of the noble gases. (a) Density of the noble gases increases with increase in proton number. --------(b) Melting point and boiling point increases with increase in proton number. ---------(c) Radius and atomic size of the noble gases decreases. ---------84

RESPONSES

Answers 2B.1 Changes in the physical properties of alkali metals . (a) False (b) False (c) True (d) False Answers 2B.2 Changes in physical properties of halogens. (a) True (b) False (c) False (d) True

Answers 2B.3 Changes in physical properties of noble gases. (a) True (b) True (c) False

UNIT 3

CHEMICAL BONDS OBJECTIVES

GENERAL OBJECTIVE : To know and understand the formation of chemical bonds, the stability of noble gases, the ionic bond, the covalent bond, the metallic bond, the hydrogen bond and van Der Waals forces of attraction. SPECIFIC OBJECTIVES : At the end of this unit you will be able to : ✓ Explain the formation of ionic bonds, covalent bonds, metallic bonds, hydrogen bonds and van Der Waals forces of attraction. ✓ Explain the stability of the noble gases. ✓ Give examples of the formation of ionic bonds, covalent bonds, metallic bonds, hydrogen bonds and van Der Waals forces of attraction. ✓ Compare the properties of ionic compounds with those of simple covalent compounds. ✓ Describe the effect of the presence of hydrogen bonds on the physical properties of molecular compounds. ✓ Describe the effect of the presence of van Der Waals forces of attraction on the melting and boiling points of molecular compounds.

INPUT 3.1 THE FORMATION OF COMPOUNDS AND THE STABILITY OF CHEMICAL BONDS Minerals in the earth crust normally exist in the form of silicates, oxides, sulphides and carbonates. This means that most minerals exist as compounds in the earth crust and only several minerals exist in the form of the pure element, as in the case of gold, diamond and platinum. This shows that compounds are more stable that the elements from which they are formed. The higher stability of these compounds is due to the presence of forces of attraction that hold together the particles of the component elements. These forces within the molecules of a compound are called chemical bonds. The higher stability derived from chemical bonding forms the basis of the formation of chemical compounds in nature. As an example, when sodium metal is heated and placed in a jar of chlorine gas, the reactive metal burns with a golden, yellow flame and releases white smoke. The product of this reaction is a white solid, sodium chloride, which is found plentifully in sea water. Sodium +

chloride

Sodium chloride

The physical and chemical properties of a compound are influenced by the kind of chemical bonding involved in its formation from its elements. Chemical bonds can be divided into five types. The two main types of chemical bonds are • the ionic or electrovalent bond • the covalent bond

Three other types of chemical bonds include • the metallic bond • the hydrogen bond • the van der Waals forces Ionic or electrovalent bonds are electrostatic attraction between ions of opposite charge, while covalent bonds involve the sharing of electrons between two atoms, forming a molecule. For metals, the attraction between the positive metal ion and the electron ‘cloud’ or ‘sea’ around it forms the metallic bond. Metallic bonds are found in metals and alloys. Two primary forces of attraction found between molecules are hydrogen bonds and Van der Waals forces. 3.2

STABILITYOF NOBLE GASES The noble gases are elements in Group XVIII of the Periodic Table. These elements are helium, neon, argon, krypton, xenon and radon. The noble gases are highly stable elements and not reactive at all chemically. Table 3.1 shows the electron configuration of all the noble gases.

Table 3.1: Electron configuration of Noble Gases Noble Gas

Electron configuration

Helium (He) Neon (Ne) Argon (Ar) Krypton (Kr) Xenon (Xe) Radon (Rn)

2 2.8 2.8.8 2.8.18.8 2.8.18.18.8 2.8.18.32.18.8

All the atoms of the noble gases, except helium, have eight electrons in the outermost electron shell. This is called an octet electron configuration. The structure of the helium atom, which has only two electrons in the single (full) electron shell, is described as a duplet electron configuration.

The duplet configuration for helium and octet configuration for the other noble gases confers very high stability upon these elements. The noble gas atom does not release, receive or share any electron and so it does not interact with other atoms to form molecules. That is why the noble gases are totally inert chemically and exist only as monoatomic molecules. Compared with the noble gases, other atoms are less stable because they do not have the duplet or octet electron configuration in its outer shell. As such the atoms of these elements are likely to take part in chemical reactions to form molecules or compounds. In the formation of chemical bonds, each atom that is involved undergoes a change in its electron configuration, achieving the noble gas configuration and resulting in increased stability of the compound formed. When atoms bond to form molecules or compounds, they can achieve noble gas configuration by (a) (b) (c)

releasing electrons, receiving electrons, or sharing electrons.

And so atoms may form any one of the following two kinds of chemical bonds: (a) ionic (or electrovalent) bonds (b) covalent bonds

A chemical bond that is formed through electron transfer from a metal atom to a non-metal atom is called an ionic (or electrovalent)bond. In the formation of this ionic bond, the metal atom releases one or more electrons, whereas the non-metal atom receives them. A chemical bond that is formed through the sharing of electrons between two non-metal atoms is known as a covalent bond.

Did you know that this is a computer generated model of a CO2 molecule? The carbon atom is found between the two oxygen atoms. Electrons play an important role in bonding together the carbon and oxygen atoms.

Figure 3.1: A molecule of carbon dioxide 3.3

IONIC OR ELECTROVALENT BONDS 3.3.1 THE FORMATION OF POSITIVEAND NEGATIVE IONS

The atoms of an element are electrically neutral because each atom has equal numbers of protons(positively charged) and electrons (negatively charged). If a neutral atom releases or receives an electron, it becomes an electrically charged particle, called an ion. (Figure 3.1) When an atom releases an electron, it now has more protons than electrons, and a positive ion is formed. On the contrary, if the atom receives an electron, it now has more electrons than protons, and a negative ion is formed.

Figure 3.1 : The formation of a positive ion and a negative ion The formation of an ion from its atom in a reaction is due to the fact that the atom has an electron configuration that is relatively unstable, and so tends to release or receive electrons to form an ion with the octet configuration, which is more stable. In a reaction between a metal and a non-metal, the metal atom releases its electron(s) to form a positive ion while the non-metal atom receives electron(s)to form a negative ion. M - ne Metal Atom A + ne non-metal atom

Mn+ positive ion

Mn+ + AnMn+Anionic compound

Annegative ion

The ions formed have opposite charges (positive and negative charge). The resulting electrostatic force of attraction that holds these ions together is called an ionic bond. The compounds formed by such bonding are categorised as ionic compounds. 3.3.2 THE FORMATIONOF SODIUM CHLORIDE

Sodium chloride is an ionic compound. Sodium chloride is formed when sodium metal reacts with chlorine gas. In this reaction (a) each sodium atom (proton no. 11),which has the relatively unstable electron configuration of 2.8.1, loses its single valence electron to form a sodium ion, Na+. This sodium ion now has the octet electron configuration of 2.8, which is more stable. 91

(b) Meanwhile, the electron released is captured by one chlorine atom (proton no.17), which has the electron configuration of 2.8.7, forming a chloride ion, Cl-.This ion is relatively more stable than the chlorine atom, because it too achieves the octet electron configuration of 2.8.8. (c) Through the transfer of a electron, both sodium and chlorine atoms achieve stable, noble gas electron configurations in the Na+ and Cl-ions. Na+ (as for Ne) 2.8 Na+ + Cl-Na +Cl -

Na - 1e 2.8.1 Cl + 1e 2.8.7

Cl- (as for Ar)

Sodium chloride

2.8.8

The Na+ and Cl- ions are held together by the electrostatic force of attraction between them. In other words, an ionic bond is formed and the resulting compound is sodium chloride, NaCl. The electron transfer that takes place in the reaction between sodium and chlorine to form sodium chloride can be pictured as in Figure 3.2. electron transfer

Na atom

Cl atom

Sodium chloride, NaCl

ion

Figure 3.2 :The formation of sodium chloride

IONIC COMPOUNDSANDTHEIR PROPERTIES

Ionic compounds consist of positive and negative ions in certain ratios, as can be seen in all chloride, carbonate, sulphate and metal nitrates, as well as most metal oxides. In an ionic compound, its postive and negative ions bond in such a ratio that the total positive charge is exactly balanced by the total negative charge. That is why ionic compounds are electrically neutral, as shown in the examples in Table 3.2. Table 3.2 : Several examples of ionic compounds Ionic compounds Potassium Chloride Magnesium oxide Calcium chloride Aluminium oxide Lead(II) nitrate Copper(I) oxide Copper (II) sulphate

Formula Total positive charge KCl K+ = 1+

Total negative charge Cl- = 1-

MgO

Mg2+

O2-

= 2-

CaCl2

Ca2+

2Cl-

= 2-

Al2O3

2Al3+ =

3O2-

= 6-

Pb(NO3)2 Pb2+

2NO3- = 2-

Cu2O

2Cu+

O2-

CuSO4

Cu2+

SO42- = 2-

= 2-

THE PROPERTIES OF IONIC COMPOUNDS

(a)

High melting and boiling points. This is because ionic compounds in the solid state consist of ions held together by electrostatic forces in a giant crystal lattice. Much heat is needed to break this lattice structure and to free the ions as the compound melts.

(b)

Conduct electricity in molten form or aqueous solution (if soluble in water), but not in solid form. In the solid state, the ions are held together by strong ionic bonds and are not free to move. Therefore the solid does not conduct electricity. When melted or dissolved in water, however, those ions are freed to move and so the compound does conduct electricity in these conditions.

(c)

Dissolve in water but not in organic solvents. Water molecules have dipolar properties. When water is added to an ionic compound, its ions are attracted to the water molecules. The electrostatic forces holding the ions together are overcome, the lattice structure breaks down, and the ions are free to mix with the water. Hence the compound dissolves in water. Organic solvents like benzene and acetone are nonpolar solvents, having covalent molecules which are non-polar and are therefore not attracted to the charged ions in the lattice structure. Hence ionic compounds are not soluble in organic solvents.

3.4

COVALENT BONDS The non-metallic elements in groups XIV, XV, XVI and XVII in the Periodic Table do not form ionic compounds among themselves. This is because the relatively high electronegativity of these elements causes their atoms to tend to attract electrons. That is why, when two or more non-metal atoms react chemically, electron transfer does not take place.

To achieve the octet structure with 8 electrons in the outer shell, two atoms of the same or different non-metal element(s) may bond, by sharing electrons (in pairs), between themselves to form diatomic molecules such asH2, O2, N2and Cl2. In other cases they may bond to form part of a molecule, as in the case of the carbon-carbon bond in the ethane molecule (C2H6). Covalent bonds form when a non-metal atom bonds with another nonmetal atom by sharing electrons. The molecule formed is called a covalent molecule. The covalent bond is a strong force of attraction. Table 3.3 shows several examples of covalent molecules formed when atoms of a non-metal bond among themselves. Table 3.3 : Several examples of covalent molecules Combining non metallic atoms Chlorine Chlorine Oxygen Oxygen Nitrogen Nitrogen Sulphur Sulphur Phosphorus Phosphorus Carbon Oxygen Carbon

Chlorine

Nitogen

Oxygen

Sulphur

Oxygen

Covalent molecule Chlorine molecule, Cl2 Oxygen molecule, O2 Nitrogen molecule, N2 Sulphur molecule, S8 Phosphorus molecule, P4 Carbon dioxide molecule, CO2 orcarbon monoxide molecule, CO Tetrachloromethane molecule, CCl4 Nitrogen dioxide molecule, NO2 or nitrogen monoxide molecule, NO Sulphur dioxide molecule, SO2 orsulphur trioxide molecule, SO3

The hydrogen atom also bonds with non-metal atoms to form covalent molecules. Examples include hydrogen chloride (HCl), hydrogen iodide (HI), hydrogen bromide (HBr), water (H2O), ammonia (NH3), hydrogen sulphide (H2S), methane (CH4). 3.4.1 THE FORMATIONOFCOVALENT MOLECULES

In the formation of covalent molecules, each of the reacting nonmetal atoms contributes one, two or three electrons to be shared, to achieve the electron configuration of the noble gases. This forms one or more shared electron pair(s) that hold the atoms together. Such bonds are called covalent bonds.

When each bonding atom donates one electron to be shared, one shared pair of electrons holds the two atoms together. This type of bond is called a singe covalent bond.(Figure 3.3) shared that that isis 1 shared electron pair

Single covalent bond

Figure 3.3 :The formation of single covalent bonds When each reacting non-metal atom donates two electrons to be shared, two pairs of electrons are shared, which hold the two atoms together. This type of bond is called a double covalent bond.(Figure 3.4) shared that is shared

2 shared electron pairs

Double covalent bond

Figure 3.4 :The formation of double covalent bonds When each reacting non-metal atom donates three electrons to be shared, three pairs of electrons are shared, which hold the two atoms together. This type of bond is called a triple covalent bond. (Figure 3.5) shared that is

3 shared electron pairs

Triple covalent bond

Figure 3.5 :The formation of a triple covalent bond

An Example Of A Covalent Molecule With Single Covalent Bonds The formation of the chlorine molecule, Cl2 96

The electron configuration of the chlorine atom is 2.8.7. Each atom has 7 valence electrons, which means each atom is one electron short of achieving the electron configuration of a noble gas. As such, two chlorine atoms may bond to achieve this. Each of them donates one electron from the outer shell to be shared. This means one pair of electrons is shared, holding the two atoms together. This is called a single covalent bond. The covalent molecule Cl2 is formed. The figure below shows how the covalent is formed. (Figure 3.6)

shared

Cl atom

Cl atom 2.8.7

Cl2molecule

Figure 3.6 : The formation of a chlorine molecule, Cl2 The formation of this covalent molecule can also be illustrated by using (a) dot-and-cross diagrams to represent electrons in the outer shell of bonding atoms as shown in Figure 3.7,or (b) short lines as in Figure 3.8. shared

Cl atom

Cl2 molecule

Single covalent bond in a chlorine molecule, Cl2

can be illustrated as

Figure 3.7 :Dot-and-cross representation of a single covalent bond in the chlorine molecule Cl2

Figure 3.8 : Short-line representation of a single covalent bond in the chlorine molecule Cl2

An example of a covalent molecule with a double bond

The formation of a molecule of oxygen, O2. The oxygen atom has the electron configuration of 2.6, which means the oxygen atom is two electrons short of achieving a noble gas electron configuration. Two oxygen atoms bond, each donating two electrons to be shared. This means two pair of electrons are shared, holding the two atoms together. The covalent molecule of oxygen, O2 is formed. This is called a double covalent bond. Shared

Shared

O2 molecule Or

O atom

O2 molecule

O atom

Double bond

Figure 3.9 :The formation of the oxygen molecule, O2 The formation of the carbon dioxide molecule, CO2 One carbon atom and two oxygen atoms bond to form one covalent molecule of CO2, as shown in Figure 3.10. Shared

O Atom (2.6)

Shared

C Atom (2.4)

O Atom (2.6)

CO2molecule

O Atom

C Atom

O Atom

Double bond

Figure 3.10 : The formation of the carbon dioxide molecule, CO2

An example of a covalent molecule with triple bonds

The formation of the nitrogen molecule, N2 The nitrogen atom has the electron configuration of 2.5, which means the nitrogen atoms needs three electrons to achieve the electron configuration of the noble gases. Two nitrogen atoms bond, each donating three electrons from the outer shell to be shared. Three electron pairs are shared, holding the two nitrogen atoms together to form the covalent molecule N2. (Figure 3.11).This type of chemical bond is called a triple covalent bond. Shared

N atom (2.5)

N2 molecule

Or Shared

or triple bond

Figure 3.11 :The formation of the nitrogen molecule, N2 Diatomic covalent molecule Elements like hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine and iodine exist as diatomic covalent molecules because the atoms in the molecules achieve the stable noble gas electron configuration. Figure 3.12 shows the covalent bonds in several diatomic covalent molecules.

that is

H2 molecule

N2 molecule

O2 molecule

F2 molecule

Cl2 molecule

Br2 molecule

I2 molecule

Figure 3.12 : Several diatomic covalent molecules

INPUT 3.4.2 COVALENT COMPOUNDS AND THEIR PROPERTIES

Compounds that have atoms which are bonded with covalent bonds in their molecules are called covalent compounds. Most covalent compounds have small, simple molecules although there are some that have giant molecular structures or long chains. (a) In covalent compounds that have small, simple molecules (Figure 3.13), even though the atoms in those molecules are held together by strong, covalent bonds, the molecules are held together by weak van der Waals forces. Therefore, not much heat is required to separate those molecules. This explains why covalent compounds with simple molecules have low melting and boiling points, and normally exist as gases or volatile liquids.

ammonia

ethanol

hydrogen chloride, HCl

water

ethane

carbon dioxide, CO2

propene

In covalent compounds or those that have giant molecules (macromolecules), such as silicon dioxide (SiO2) and diamond (a carbon allotrope) (Figure 3.14), the atoms are bonded with strong covalent bonds to form an infinite, three dimensional lattice structure. This explains why covalent compounds with giant molecules are very hard substances and have high melting and boiling points.

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o oxygen atom ● silicon atom

o carbon atom

Silicon dioxide (SiO2)

Diamond (an allotrope of carbon)

Figure 3.14 :Giant structure of two covalent substances (b) Polymers such as starch, cellulose, proteins and rubber as well as synthetic polymers like polythene, polyvinylchloride and nylon are covalent compounds that have long-chain molecules. Each molecule is made up of many carbon atoms bonded together by covalent bonds, as shown for the polythene molecule (Figure 3.15).

Figure 3.15 :A portion of the polythene molecule Different from covalent compounds that have three dimensional lattice structures as for silicon dioxide and diamond, compounds that have long-chain molecules usually are harder substances have lower melting and boiling points. PHYSICAL PROPERTIES OF COVALENTCOMPOUNDS

(a)

Melting and boiling point Covalent compounds that have small and simple molecules have lower melting and boiling points, whereas compounds that have giant structures have higher melting and boiling points.

(b)

Physical State Covalent compounds that have simple molecules usually exist in the gaseous or liquid states where as compounds

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with giant structures or long-chain molecules exist as solids at room temperature. (c)

Electrical conductivity All covalent compounds consist of neutral atoms or molecules and not ions, and so do not conduct electricity.

(d)

Solubility Covalent compounds that have simple molecules do not dissolve in water but dissolve readily in organic solvents. Covalent compounds that have giant molecular structures do not dissolve in either in water or organic solvents.

3.4.3 COVALENT COMPOUNDSAS SOLVENTS

Almost all covalent compounds that have simple molecules exist as liquids at room temperature, and most of them play an important role as solvents in daily life or industry. Generally solvents can be categorised as water and organic solvents.Even though water is a covalent compound, it consists of simple molecules, H2O, which are dipolar. Water can dissolve almost all substances (except organic substances), including ionic compounds,and is known as a universal solvent. The importance of wateras a solvent is seen in the following: (a) It dissolves digested food substances such as glucose and amino asids to be transported in blood plasma to the whole body. (b) It dissolves the metabolic waste products found in cells, such as urea and uric acid to be isolated and disposed of. (c) It dissolves sugar in the soft drink industry and various salts in the chemical industry. (d) It acts as a cleansing medium in removing dirt when used in cleaning. Organic solvents are covalent compoundsthat exist as liquids at room temperature. The use of various organic solvents in daily life and also in industry is summarised in Table 3.4.

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Table 3.4: The use of various organic compounds as solvents. Organic Solvent (a) Petrol and kerosene (b) Alcohols

(f) Chloroform (trichloromethane) (g) Xylene

3.5

Use Dissolves and removes grease in the workshops. Dissolves iodine smears, andalso act assolventsfor perfumes and in the production of various medicines Dissolves paint and removes left-over paint marks. Used in dry-cleaning to remove stains on clothes. As a solvent for (i) polystyrene,which is used in making children’s toys. (ii) cellulose ethanoate,used in the production of synthetic thread (rayon) (iii) nail varnish, common cement and various types of paint. As a solvent in many plastics. As a solvent for sulphur.

A COMPARISON OF IONIC AND SIMPLE COVALENT COMPOUNDS A comparison of ionic compounds ion and simple covalent compounds is shown in Table 3.5.

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Table 3.5 : Comparison of ionic compounds and simple covalent compounds Ionic compounds 1 Formed from ions with opposite charge. These ions are bonded by strong electrostatic forces. 2

3.6

(Simple) covalent compounds Formed from molecules. These molecules are attracted to one another by weak molecular forces (Van der Waals forces) Have high melting and boiling points Have low melting and boiling (because much energy is needed to points. overcome strong electrostatic forces). Most exist as crystalline solids. Most exist as gases, volatile liquids or solids at room temperature. Formed from the ions of a metal and Formed from non-metal atoms a non metal. (that is, an with non-metal atoms.(bonding electropositive element and of electronegative atoms.) anelectronegative element. Melted compounds or their aqueous Non-electrolytes. solutions are good electrolytes – because of the existence of free ions. Can conduct electricity when melted Do not conduct electricity but not as a solid. This is because in whether in solid physical state or the solid physical state, all the ions when melted.Covalent are held together and cannot move compounds do not have free freely. ions. Dissolves in water. Dissolve in oganic solventssuch as tolueneand tetrachloromethane. (Doesn’t dissolve in water.) Examples : NaCl , MgO, Li3N Examples : CO2, N2, NH3

METALLIC BONDING One particular property of metals cannot be explained from current understanding of ionic or covalent bonds. In covalent and solid ionic compounds, electrons are not free to move according to any applied electric potential. That is why ionic solids and covalent compounds are electrical insulators. In metals, valence electrons (electrons in the outer shell) move freely between atoms in the whole piece of metal. X-ray analysis shows that metal crystals consist of positive ions surrounded by electrons which are not localised.

104

Metallic bonding is therefore defined as electrostatic attraction between positively charged metal ions and a surrounding ’electron sea’. The phenomenon of delocalised electrons explains the high electrical and thermal conductivity found in metals. As such metals conduct electricity both in the solid and liquid states. (Figure 3.16) Positive metal ion e

e ++++ e

+++

e e

delocalised valence electron e e +

e e

Figure 3.16 : Metallic bonding based on Electron-Sea Model The movement of electrons in metals can be explained using the energy band model of metal crystals. Metallic bonding can be found in metals and alloys. 3.7

THE HYDROGEN BOND When a hydrogen atom is bonded to a highly electronegative atom, (such as N, O and F), the resulting covalent bond is highly polarised, for example + ⎯O⎯H The H + atom can form another weak bond with the electron pair from another highly electronegative atom. Because the hydrogen is very small, it can approach other atoms very closely and so form relatively strong bonds with them (even though this type of bond is not as strong as the covalent bond). This type of bond is called the hydrogen bond. Several other instances of hydrogen bonding are shown in Figure 3.17. The dotted lines represent hydrogen bonds.

105

(Strong hydrogen bond)

(Weak hydrogen bond)

Figure 3.17: Several instances of hydrogen bonding Hydrogen bonds are strong enough to bind molecules together and affect the physical properties of a substance, but not its chemical properties 3.7.1 The effects of hydrogen bonding on physical properties of molecular compounds i. Melting and boiling points Hydrogen bonding increases the forces of attraction between molecules and so raises the melting and boiling points of molecular substances. The boiling point of an alcohol is higher than that of an alkane with similarly sized molecules, because the alcohol molecules are bound by hydrogen bonds. As such ethanol has a boiling point of 78C whereas propane has a boiling point of - 42C. In liquid ethanol, the molecules are interlinked with hydrogen bonds. (Figure 3.18)

Figure 3.18: Hydrogen bonds between ethanol molecules ii. Solubility in water Generally, organic compounds are not soluble in water, because they have a molecular structure. However there are several groups of organic compounds which are moderately soluble in 106

water, due to the occurrence of hydrogen bonding between their molecules. Examples are found in the alcohols, sugars, carboxylic acids, amines and phenol. Some of these examples are shown in Figure 3.19.

Propanoic acid in water

Phenol in water

Figure 3.19: Some examples of organic compounds that are soluble in water iii. Dimerisation If benzoic acid is dissolved in benzene, the dissolved acid form particles with the formula of (C6H5.COOH)2. Benzoic acid molecules pair up, linking to each other by hydrogen bonds, to form dimers. (Figure 3.20)

Figure 3.20 : Benzoic acid molecules linked to form pairs. iv. Give the specific shapes of protein molecules The 3-dimensional shape of a protein molecule is in part maintained by hydrogen bonds formed between different points along the chain of the molecule. Protein chains normally have a helical structure and are kept in this shape by hydrogen bonds between C = O and N — H groups found in polipeptide chains. (Figure 3.21)

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Figure 3.21: Hydrogen bonds in a protein helix 3.8

VAN DER WAALS FORCES 3.8.1 FORMATION OF VAN DER WAALS FORCES

The electrons in a molecule vibrate about a mean position. At any one moment, one end of that molecule is a little negative and the opposite end a little positive. The negative charge affects the nearby end of a neighbouring molecule (by repelling its electrons a little). This produces a negative charge on the other end of the neighbouring molecule. This process is repeated in other neighbouring molecules. This results in a temporary dipole being formed in a series of molecules. This temporary dipole causes the molecules to attract one another. This temporary attractive force between molecules is known as van der Waals force. Even though this temporary dipole is quickly lost, new dipoles are constantly formed so that there are always temporary dipoles acting between the molecules. (Figure 3.22) +-

||||||||

+-

||||||||

+-

Figure 3.22 : Temporary dipoles between I2 molecules

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3.8.2 THE EFFECTS OF DER WAALS FORCES ON THE MELTING AND BOILING POINTS OF MOLECULAR COMPOUNDS

Van der Waals forces may be very weak but are actually the main bonding force which hold molecular substances together, in the solid and liquid states (for instance in solid iodine (I2), liquid tetrachloromethane (CCl4) and in liquid hexane). Because this force is relatively weak, molecular substances tend to have low boiling and melting points.

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REFERENCES

1. Ang, C. (2009). A Complete Guide to GCE O Level Pure Chemistry. Singapore: Fairfield Book Publishers. 2. Lim, Y.S., Yip, K.H. (2009). Pre U Text STPM Chemistry. Malaysia: Pearson Malaysia 3. Lim, Y.S., Yip, K.H.(2009). Pre U Text STPM Physical Chemistry. Malaysia: Pearson Malaysia. 4. Norbani Abdullah et. al. (1998). Kimia Fizikal Asas Matrikulasi. Malaysia: Fajar Bakti. 5. Chang, R. (2002). Chemistry. New York: McGraw Hill.

6. Slowinski, E. J., Wolsey, W. C. & Rossi, R. C, (2012). Chemical Principles in the Laboratory. Australia. Brooks/Cole, Cengage Learning. 7. Timberlake, K.C, (2010). Basic Chemistry (3rd edition). United States. Prentice Hall.