Longhorn Secondary Chemistry Form 2 PDF

Longhorn Secondary Chemistry Form 2 George Ngaruiya Joan Kimaru Paul Mburu i Published by: Longhorn Publishers (K) Ltd., Funzi Road, Industrial Area, P.O. Box 18033, Nairobi, Kenya. Longhorn Publishers (U) Ltd.,...

Longhorn Secondary Chemistry Form 2 George Ngaruiya Joan Kimaru Paul Mburu i Published by: Longhorn Publishers (K) Ltd., Funzi Road, Industrial Area, P.O. Box 18033, Nairobi, Kenya. Longhorn Publishers (U) Ltd., Plot M220, Ntinda Industrial Area, Jinja Road, P.O. Box 24745, Kampala Uganda. Longhorn Publishers (Tanzania) Ltd., Plot No. 4, Block 37B, Kinondoni, Kawawa Road, P.O. Box 1237, Dar es Salaam, Tanzania. © George Ngaruiya Joan Kimaru Paul Mburu, 2003 All rights reserved. No part of this publication may be reproduced, stored in a retrieval system or transmitted in any form or by any means, electronic, mechanical, photocopying, recording or otherwise without the prior written permission of the Copyright owner. First published 2003, Reprinted 2004, 2006 Second Edition 2007 Corrected 2008 Reprinted 2009 ISBN 978 9966 49 563 0 Illustrations by Timothy Maleche Printed by Printpoint, Changamwe Road, off Enterprise Road, Industrial Area, P.O. Box 30957-00100, Nairobi, ii Kenya.  TABLE OF CONTENTS Unit 1: Structure of the atom and the periodic table................................... 1 1.1 The structure of the atom...................................................................... 1 1.2 Atomic characteristics.......................................................................... 7 1.3 The periodic table................................................................................. 16 1.4 Ion formation........................................................................................ 20 Revision Exercise 1.............................................................................. 41 Unit 2: Chemical families: Patterns in properties....................................... 44 2.1 Introduction........................................................................................... 44 2.2 Group I elements - Alkali metals.......................................................... 44 2.3 Trends in physical properties of alkali metals...................................... 45 2.4 Trends in chemical properties of alkali metals ……………………….49 2.5 Similarities of ions and formulae of compounds formed by alkali metals........................................................................ 55 2.6 Uses of alkali metals............................................................................. 57 2.7 Group II – Alkaline- earth metals......................................................... 58 2.8 Trends in physical properties of alkaline-earth metals......................... 58 2.9 Trends in chemical properties of alkaline-earth metals........................ 59 2.10 Similarities of ions and formulae of compounds formed by alkaline- earth metals.......................................................... 70 2.11 Uses of alkaline earth metals................................................................ 72 2.12 Group VII - Halogens......................................................................... 72 2.13 Trends in physical properties of halogens............................................ 73 2.14 Trends in chemical properties of halogens........................................... 75 2.15 Similarities of ions and formulae of compounds of halogens.............. 83 2.16 Uses of halogens................................................................................... 85 2.17 Group O - Noble gases....................................................................... 86 2.18 Uses of noble gases............................................................................... 87 2.19 Properties and trends across the period three....................................... 88 iii  2.20 Trends in physical properties of elements across period three............. 88 2.21 Trends in chemical properties of elements elements across period three............................................................... 90 Revision Exercise 2.............................................................................. 103 Unit 3: Structrue and bonding....................................................................... 105 3.1 Ionic bonding........................................................................................ 105 3.2 Covalent bonding.................................................................................. 110 3.3 Co-ordinate (dative) bonding................................................................ 113 3.4 Structure of covalent compounds......................................................... 114 3.5 Metallic bonding................................................................................... 117 3.6 Hydrogen bonding................................................................................ 119 3.7 Types of bonds across period three....................................................... 120 3.8 Applications.......................................................................................... 121 Revision Exercise 3.............................................................................. 123 Unit 4: Salts...................................................................................................... 126 4.1 Methods of preparing salts.................................................................... 127 4.2 Solubility of salts.................................................................................. 149 4.3 Action of heat on salts.......................................................................... 152 Revision Exercises 4............................................................................. 162 Unit 5: Effect of an electric current on substances....................................... 165 5.1 Conduction of electricity...................................................................... 165 5.2 Electrolysis........................................................................................... 168 Revision Exercises 5............................................................................ 179 Unit 6: Carbon and its compounds................................................................. 180 6.1 Forms of carbon.................................................................................... 180 6.2 Chemical properties of carbon.............................................................. 186 6.3 Carbon(IV) oxide.................................................................................. 189 6.4 Carbon(II) oxide................................................................................... 197 6.6 Carbonates and hydrogencarbonates.................................................... 204 iv  6.6 Production and manufacture of sodium carbonate............................... 210 6.7 Importance of carbon and its oxides..................................................... 213 Revision Exercise 6.............................................................................. 221 Self check test papers........................................................................................ 225 Self check test paper 1........................................................................................ 225 Self check test paper 2........................................................................................ 228 Self check test paper 3........................................................................................ 231 Self check test paper 4........................................................................................ 234 Self check test paper 5........................................................................................ 237 Self check test paper 6........................................................................................ 239 Glossary............................................................................................................. 242 Appendix 1:The Periodic Table......................................................................... 252 Appendix 2: Safety symbols.............................................................................. 253 v UNIT 1: Structure of the atom and the periodic table 1.1 The structure of the atom Names and symbols of the first twenty elements of the periodic table In Form I, we learnt that there are about 119 known elements. We have also learnt names and symbols of some common elements that can be arranged according to their behaviour under different conditions. In this class, we shall learn how to arrange elements according to their behaviour. Continue to practise the use of chemical symbols particularly those of the first 20 elements of the periodic table. These are listed in Table 1.0 below. Table 1.0: Names and symbols of the first twenty elements of the periodic Table Name of element Chemical symbol 1 Hydrogen H 2 Helium He 3 Lithium Li 4 Beryllium Be 5 Boron B 6 Carbon C 7 Nitrogen N 8 Oxygen O 9 Fluorine F 10 Neon Ne 11 Sodium Na 12 Magnesium Mg 13 Aluminium Al 14 Silicon Si 15 Phosphorus P 16 Sulphur S 17 Chlorine Cl 18 Argon Ar 19 Potassium K 20 Calcium Ca 1  Simple structure of the atom Atoms are very small particles of an element, but they contain even smaller particles called sub-atomic particles. There are three kinds of sub-atomic particles. These are the protons, electrons and neutrons. The protons and neutrons are found in the nucleus. The nucleus is the middle part of the atom. It is a very dense and extremely small part of the atom. Outside the nucleus is a much larger region of the atom where electrons occur in energy levels at different distances from the nucleus of the atom. The electrons can be imagined as circulating the nucleus, at great speeds in those energy levels. We can represent the structure of an atom simply as follows: nucleus (protons and neutrons) region where cloud of electrons is likely to be found Fig. 1.1: Atom showing nucleus and electrons Electrons are negatively charged (–ve), protons are positively charged (+ve) and neutrons have no charge. Therefore the nucleus has a positive charge because of the protons as illustrated in Fig. 1.1. If we magnify the nucleus, we can show the sub-atomic particles (protons and neutrons) as in Fig. 1.2. nucleus neutron proton Fig.1.2: Nucleus containing protons and neutrons Since protons have similar charge (+ve) they repel each other. However, the repulsion is minimised by the presence of the neutrons as illustrated in Fig. 1.3. 2  The nucleus It is at the centre of the atom. It contains protons and neutrons. It has a positive charge because of the protons. The whole mass of the atom is concentrated in the nucleus. However, it is tiny compared to the size of the atom. The electrons They move around the nucleus. They are negatively charged. They are tiny, but cover a lot of space. They occupy the energy levels. Fig. 1.3: Atomic structure showing protons and neutrons in the nucleus and electrons in the energy level In a neutral atom: Number of protons = Number of electrons This means that each proton (+ve) has an electron (_ve) as a partner. Table 1.1 Particles present in an atom Sub-atomic particle Mass Charge Where found in atom Proton 1 + ( positive) inside nucleus Neutron 1 0 ( neutral) inside nucleus 1 Electron – ( negative) outside nucleus 1840 The structure of the atom In a well planned town or city, houses are constructed according to the plan. The house rent also varies. Those people who live in some neighbourhoods of the city are wealthy people and therefore pay a lot more than those who live in other areas. 3  Atomic “city” These are found in a country where people live according to the following city plan. 1. There are three suggested types of houses that are allowed in each neighbourhood. Only 2 people are allowed to occupy any house. 1st neighbourhood - has one low cost stone house only 2nd neighbourhood - has four middle cost stone houses only 3rd neighbourhood - has one high cost stone house only 2. All houses must be built away from city/town centre. How many people are in: (a) Low cost houses in the 1st neighbourhood? (b) Middle cost houses in the 2nd neighbourhood? (c) What is the maximum number of high cost houses expected to be in the 3rd neighbourhood? We notice that the further the houses are from the city centre, the more expensive they are. The cheapest houses are in the first neighbourhood. So the people who live in the 3rd neighbourhood are very wealthy. From our atomic city vocabulary list, we mentioned that the houses are settled according to wealth and ability to pay higher. This implies that the 3rd neighbourhood has wealthy people. The 2nd has poor people and the 1st has the poorer people. In other words the wealth increases as we move from the city centre. To understand the atomic structure, let us look at a suggested analogy of atomic ‘city’. The city plan (vocabulary list) is as follows: City = Atom People = Electrons Neighbourhoods = Energy levels Wealth = Energy City centre = Nucleus Look at Fig. 1.4. What would you get if you completed the dotted lines? The neighbourhoods are represented by energy levels where an energy level is the path of an electron around a nucleus in an atom. 4 3rd 2nd 1st neighbourhoods City/town centre Fig. 1.4: Atomic ‘city’ showing the neighbourhoods. Activity 1.0 Complete the dotted lines in Fig. 1.4 to form circles. Represent each number of people in each type of house by dots ( ) or crosses (x), and distribute them in each neighbourhood. Compare your diagram with Fig. 1.5 which represents a magnesium atom. Fig. 1.5: Structure of a magnesium atom The various energy levels in an atom are represented by a series of circles sharing the same centre (nucleus), separated from each other by roughly equal distances. The nucleus of the atom is at the centre of the circles. The electrons in the energy levels are represented by dots ( ) or crosses (x). The energy levels are labelled 1st, 2nd, 3rd, 4th and so on starting from the one nearest to the nucleus as shown in Fig. 1.6. 3rd 2nd 1st nucleus (protons and neutrons) Fig. 1.6: Energy levels 5  The electrons that occupy the 1st energy level have lower energy than those in the 2nd energy level. Subsequently those in the 2nd energy level have lower energy than those in the 3rd energy level and so on. The 1st energy level usually has a maximum of two electrons while the 2nd energy level has a maximum of eight (8) electrons. The maximum number of electrons that an energy level can hold are given by 2n2. Where n = number of energy level. 1st, 2nd, 3rd, etc. This means that when we draw the atomic structure, we should never put more than two electrons in the 1st energy level or more than 8 in the 2nd level. The third energy level is usually not full but can accommodate one to eight electrons. Third energy level holds a maximum of 8 electrons for the first 20 elements. The first twenty elements of the period table have 1 to 20 protons and 1 to 20 electrons. We can summarise the electrons in each energy level as shown in Table 1.2. Table 1.2 The electron arrangements of the first twenty elements Element Symbol Number 1st Energy 2nd 3rd Energy 4th Energy Electron of protons level Energy level level arrangement level Hydrogen H 1 1 Helium He 2 2 Lithium Li 3 2.1 Beryllium Be 4 2.2 Boron B 5 2.3 Carbon C 6 2.4 Nitrogen N 7 2.5 Oxygen O 8 2.6 Fluorine F 9 2.7 Neon Ne 10 2.8 Sodium Na 11 2.8.1 Magnesium Mg 12 2.8.2 Aluminium Al 13 2.8.3 Silicon Si 14 2.8.4 Phosphorous P 15 2.8.5 Sulphur S 16 2.8.6 Chlorine Cl 17 2.8.7 Argon Ar 18 2.8.8 Potassium K 19 2.8.8.1 Calcium Ca 20 2.8.8.2 6  Arrangement of electrons in an atom helps us to explain the patterns in properties of the elements. These properties are the basis of the periodic table which we will discuss later. Now let us illustrate the electron arrangement of the following atoms: Hydrogen, helium, lithium, fluorine, neon, sodium and chlorine respectively. H1 He 2 Li 2.1 F 2.7 Ne 2.8 Na 2.8.1 Cl 2.8.7 Fig. 1.7: Atomic structure of H, He, Li, F, Ne, Na and Cl Draw diagrams to show the electron arrangement of the following atoms: K, Ar, Ca and S. 1.2 Atomic characteristics Atomic number and mass number We have learnt how to draw atomic structure showing the nucleus and arrangement of electrons in various energy levels. It is important at this stage to know how to calculate the number of protons and neutrons inside the nucleus. This will enable us to show the number of these sub-atomic particles when we draw atomic structure diagrams. Atomic number and mass number are two simple numbers that tell us something about an atom. Note the following points: The atomic number tells us how many protons are there in a nucleus. It is denoted by letter Z. 7  It also tells us the number of electrons in an atom. The atomic number Z The number of protons The number of of an element = in the nucleus of = electrons in its atom an atom The mass number is the sum of protons and neutrons, therefore it is always bigger than the atomic number. Mass number is denoted by letter A. Roughly the mass number is double the atomic number. To get the number of neutrons, we just subtract the atomic number from the mass number, that is A – Z. Study Table 1.3 which shows the relationship between atomic number, number of neutrons and mass number for some elements. Table 1.3 Relationship between the number of protons, neutrons and mass number of some elements Symbol Number of protons: Neutrons Mass Atom number, A Atomic number, Z Hydrogen H 1 0 1 Carbon C 6 6 12 Nitrogen N 7 7 14 Sodium Na 11 12 23 Chlorine CI 17 18 35 From Table 1.3, we notice that mass number, A, is a sum of protons and neutrons i.e Mass number A = number of protons + number of neutrons. If we represent neutrons by N, we see that A = Z + N Using letters A, Z and N, how can we get the number of neutrons N, in the nucleus of an atom? The answer to this question is N = A – Z Example Calculate the number of neutrons in a chlorine atom given that the atomic number, Z = 17 and mass number, A = 35 The number of neutrons =A–Z = 35 – 17 = 18 Therefore, a chlorine atom has 18 neutrons. Remember that since the number of protons and electrons are equal in an atom, once you know the atomic number which represents the number of protons, you automatically know 8 the number of electrons in the atom. Suppose you are required to draw the atomic structure of an atom showing the sub-atomic particles as in the following example: A certain element X has atomic number 6 and mass number 14. Draw the atomic structure of element X showing all the sub-atomic particles. The sub-atomic particles to show are protons, neutrons and electrons. Atomic number = number of protons = number of electrons = 6 Number of neutrons = mass number – atomic number or A - Z = 14 – 6 = 8 neutrons This means that we will show 6 protons and 8 neutrons in the nucleus. The 6 electrons will be indicated in the energy levels as shown in Fig. 1.8. 6p 8n Fig. 1.8: Atomic structure of X Exercise Now draw the atomic structure of an atom Y showing the sub-atomic particles given the following information: Atomic number, Z = 8 Mass number, A = 17 Usually, the atomic number, Z, and mass number A, of an atom of an element X can be written alongside the symbol of that element, one as a superscript and the other a subscript as shown below: Mass number, → A Element, → X Atomic number, → Z So the symbol for an atom of sodium would be written as Mass number 23 Na 11 Atomic number 9  Note: The top number is referred to as the superscript and bottom number as the subscript. We can also represent the sub-atomic particles using symbols with the mass of the particle as a superscript and the charge as a subscript as follows: 1 Proton – 1 p 1 Neutron – 0 n 0 Electron – –1 e An electron has zero mass and charge –1. Can you tell what the letters p and n stand for? Isotopes We are all familiar with charcoal. May be some of us know how charcoal is formed. Burning of charcoal should be discouraged because cutting down trees to make charcoal interferes with our water catchment areas. What is the colour of charcoal? What is charcoal in chemistry? Charcoal is black in colour. In chemistry charcoal is simply carbon. A piece of charcoal has millions and millions of carbon atoms. Let us represent a piece of charcoal as in Fig. 1.9. A piece of charcoal Carbon atoms Fig. 1.9: Millions of carbon atoms that make charcoal Carbon atoms exist in two forms represented as 126 C and 146C The two atoms can be represented using diagrams as shown in Fig. 1.10. 10  nucleus neutrons protons Carbon-12 Carbon-14 Fig. 1.10: Carbon isotopes What difference do you notice between the two carbon atoms? The bottom numbers are the same but the top numbers are different. This means that they have the same atomic number but different mass numbers. What causes the difference in mass number? Calculate the sub-atomic particles for each of the carbon atoms. You will notice that the different masses are caused by 2 extra neutrons in the nucleus of carbon-14. Such atoms are called isotopes. Isotopes are atoms of the same element, which have the same number of protons in the nucleus (atomic number) but different mass numbers. Some elements are made up of just one type of atom, while others exist as mixtures of isotopes. Some of these elements are listed in Table 1.4. Table 1.4 Examples of isotopes (a) Hydrogen Hydrogen Deuterium (0.01%) Tritium HYDROGEN (99.99%) Heavy hydrogen (Trace) Symbols of iso- topes of hydrogen 1 1 H H 2 1 3 1 H No. of protons 1 1 1 No. of neutrons 0 1 2 No. of electrons 1 1 1 Mass number (p + n) 1 2 3 11  (b) Carbon CARBON Carbon-12 Carbon -13 Carbon -14 (98.9%) (1.1 %) (trace) Symbols of car- bon isotopes 12 6 C 13 6 C 14 6 C Protons 6 6 6 Neutrons 6 7 8 Electrons 6 6 6 Mass number (p + n) 12 13 14 (c) Chlorine Chlorine -35 Chlorine -37 CHLORINE (75 %) (25%) Symbols of chlorine 35 Cl 37 Cl isotopes 17 17 Protons 17 17 Neutrons 18 20 Electrons 17 17 Mass number (p + n) 35 37 (d) Oxygen OXYGEN Oxygen -16 Oxygen -17 Oxygen -18 Symbols of oxygen 16 O 17 O 18 O isotopes 8 8 8 Protons 8 8 8 Neutrons 8 9 10 Electrons 8 8 8 Mass number 16 17 18 Relative Atomic Mass Atoms are so tiny that their masses cannot be measured on a balance or scale directly. We cannot weigh an atom! Hydrogen is the lightest known element. 12 A hydrogen atom is therefore the lightest. Originally, the mass of a hydrogen atom was taken as the standard reference atom. Its atomic mass was arbitrarily fixed as one unit i.e H = 1. The mass of any other element was found by comparing its mass with that of hydrogen. The idea was to find out how many times the atom of another element is as heavy as one atom of hydrogen hence relative atomic mass (R.A.M) with a symbol (Ar). This can be expressed mathematically as follows. mass of 1 atom of the element Relative Atomic Mass = mass of 1 atom of hydrogen For instance an oxygen atom (O), has a mass of 16. This means one oxygen atom is 16 times heavier than one atom of hydrogen. Sometimes the symbol Ar is used for R.A.M with the symbol of the atom in parenthesis after the symbol. For example, Ar (O) means relative atomic mass of oxygen. Many changes have occurred since the hydrogen scale was introduced. In the twentieth century, the hydrogen scale was replaced by oxygen as the standard since oxygen combines with most elements. But later the oxygen scale was found to be unsatisfactory because oxygen has several isotopes which would have different masses depending on which oxygen isotope was being considered (see Table 1.4) (d). In 1961 the International Union of Pure and Applied Chemistry (IUPAC) recommended as the standard the most abundant of the carbon isotopes, carbon - 12 ( 6 C ). It has an abundance of 98.9%. 12 Nowadays, the atomic mass of an element is measured by comparing it with 1 12 of the mass of one atom of a carbon-12 ( 126 C ). One atom of 6 C isotope is taken to have a mass of 12.00 units. 12 1 of 12 the mass of one atom of carbon – 12 = 1.00 units Average mass of one atom of the element Ar of an element = 1 12 × mass of one atom of carbon-12 The relative atomic mass (Ar) of an element, is defined as the average mass 1 of one atom of the element compared with 12 of the mass of one atom of carbon - 12. 13  Note: Relative atomic mass has no units. It is a ratio of two masses. The relative atomic masses are not whole numbers like mass numbers. This is because the abundance of isotopes of an element is different. For example a sample of chlorine is a mixture of two isotopes, 35Cl and 37Cl in the ratio of 3 : 1 respectively, i.e 75% is 35Cl and only 25% is 37Cl. Calculation of the Relative Atomic Mass As mentioned earlier many elements naturally consist of a mixture of isotopes. The abundance in which the isotopes occur in an element differs in different elements. This is why the term ‘average’ mass of one atom is used in the definition above. The proportions (abundance) in which the isotopes occur in an element may be stated as: - a ratio - percentage of the total - a fraction of the total. Example 1: You are given ratio abundance Chlorine consists of two isotopes, chlorine – 35 and chlorine – 37 in the ratio 3 : 1. Calculate the relative atomic mass (R.A.M) of chlorine. Solution Suppose the sample contains 4 atoms of chlorine, in the ratio 3 : 1, respectively, then 3 atoms will each have a mass of 35 and 1 atom will have a mass of 37. the total mass of 35Cl = 35 × 3 While the total mass of 37Cl = 37 × 1 Therefore, the average mass of chlorine atoms will be Total mass of all atoms (35 × 3) + (37 × 1) = Total number of atoms 4 The R.A.M = 35.5 Note: R.A.M has no units. It is not the mass number since chlorine has 2 mass numbers; 35 and 37. The R.A.M is also nearer to the mass number of the more abundant isotope, i.e, 35Cl. 14 Example 2: You are given percentage abundance Silicon (Z = 14) consists of three isotopes: silicon – 28, 92.2%, silicon – 29, 4.7% and silicon - 30, 3.1%. Find the relative atomic mass of silicon. Solution Percentage abundance simply means that if we have 100 atoms of an element called silicon, 92.2 atoms will each have a mass of 28 28 × 92.2 Therefore, the total mass of these = 100 4.7 atoms will be silicon – 29 29 × 4.7 Therefore, the total mass of these = 100 3.1 atoms will be silicon – 30 3.1 × 30 Therefore, the total mass of these = 100 The average mass of a silicon atom is 28 × 92.2 4.7 × 29 3.1 × 30 total mass of all atoms = + + 100 100 100 Therefore, R.A.M of silicon = 28.1 Example 3: Abundance is given as a fraction of the total 9 1 A sample of an element X consists of of 168X and of 188 X 10 10 Show that the relative atomic mass of X would be 16.2 Solution 9 1 R.A.M = (16 × ) + (18 × ) = 16.2 10 10 15 1.3 The periodic table We have learnt about the structure of the atom and how the electrons are arranged in various energy levels. We are now going to use that knowledge to build up the periodic table. Activity 1.1 Build up of the periodic table Materials required 2 graph papers. Pens or pencils Rulers Pair of scissors or razor blades Glue stick Procedure 1. Draw twenty squares on the graph papers (4 cm × 4 cm). 2. Cut out the squares. 3. On the top left hand corner of each square, label A on one side and B on the other side as illustrated in Tables 1.5 and 1.6. 4. On side A, write the electron arrangement of the element only. For example if it is magnesium, just write 2.8.2. Table 1.5. 5. On side B, write the actual symbol of the element at the centre of the 24 paper. Include the atomic and mass numbers of the elements e.g. 12 Mg as in Table 1.6. Write the name of each element below the symbol. 7. With side A facing up, shuffle or mix the squares or rectangles randomly. 8. Take a piece of manila paper and draw a table of eight columns (up- down) and four rows (left to right) as shown in Table 1.7. The squares should be 5 cm × 5 cm. 9. Place the pieces of paper, with side A facing up on the squares on the manila paper as follows. (i) Those with one energy level only, place them in the first row. If it has 1 electron in the outermost energy level place it in column I. If it has 2 electrons in the outermost energy level place it in column II. 16 6.2 The periodic table Table 1.5 Side A A A A A A 1 2 2.1 2.2 2.3 A A A A A 2.4 2.5 2.6 2.7 2.8 A A A A A 2.8. 1 2.8. 2 2.8. 3 2.8. 4 2.8. 5 A A A A A 2.8. 6 2.8. 7 2.8. 8 2.8. 8. 1 2.8. 8. 2 Table 1.6 Side B B 1 B B B B 1 H 4 2 He 7 3 Li 9 4 Be 11 3 B Hydrogen Helium Lithium Beryllium Boron B B 14 B 16 B 19 B 20 12 6 C 7 N 8O 9F 10 Ne Carbon Nitrogen Oxygen Fluorine Neon B 23 B 24 B 27 B 28 B Na 11 12 Mg 13 Al 14 Si 31 15 P Sodium Magnessium Aluminium Silicon Phosphorus B 32 B 35 B 40 B 39 B 40 16 s 17 Cl 18 Ar 19K 20 Ca Sulphur Chlorine Argon Potassium Calcium (ii) If it has two energy levels, place it in row two. If there is only 1 electron in the outermost energy level, place it in column I; if it has 2 electrons in the outermost energy level, place it in column II and so on. (iii) Place those with three energy levels in row three, and those with four energy levels in row number four. 17  Always place them under the correct columns according to the number of electrons in the outermost energy level. (iv) Once you have arranged your papers neatly do not disturb them. (v) Turn each piece without changing its position to side B. Arrange them neatly. (vi) Now compare the pattern with an actual periodic table of elements. see appendix 1. (vii) Paste your pieces of papers on the manila paper. Table 1.7 Column Column Column Column Column Column Column Column I II III IV V VI VII VIII Row one Row two Row three Row four The periodic table The periodic table is organised as a big grid. In a grid, there are rows and columns. The periodic table has rows and columns too and they each mean something different. The elements are placed in specific columns because of the way they behave. They have similar chemical properties because they have the same number of electrons in their outermost energy level. In a periodic table, each of the rows is called a period. PERIODS Fig. 1.11 Periods in the periodic table 18 Elements which are located in the same row or period have something in common. All the elements in a period have the same number of energy levels. If we want to know in which period to find an element we count the number of energy levels. If it has 4 energy levels; it is in the 4th period. Every element in the top row (the first period) has one energy level. All of the elements in the second row (the second period) have two energy levels. The maximum number of energy levels is seven representing period seven. The periodic table has a special name for its columns too. A column in the periodic table is called a group. The elements in a group have the same number of electrons in their outermost energy level. GROUPS Fig. 1.12 Groups in the periodic table An element in the first column belongs to group I. It has one electron in its outermost energy level. An element in the second column is in group II and has two electrons in the outermost energy level. As we keep counting the groups, we know how many electrons are in the outermost energy level. Therefore, the group number also is the number of electrons in the outermost energy level. Electron arrangement and the periodic table Electron arrangement short hand Given the atomic number of an element, we can write the electron arrangement in shorthand. We shall illustrate this in the following exercise. The atomic numbers of X, Y and Z are 12, 13 and 17 respectively. Write the electron arrangement of each. Solution. X 2.8.2 Y 2.8.3 Z 2.8.7 19  All these three elements are in period 3 because each has three energy levels. X is in group two, Y in group 3 and Z in group 7, following the number of electrons in their outermost energy levels. From activity 1.1, the order of the elements according to their electron arrangement should be as in Table 1.8 below. Table 1.8 Electron arrangement of the first 20 elements of the periodic table Group I Group II Group III Group IV Group V Group VI Group VII Group VIII A A Period 1 1 2 A A A A A A A A Period 2 2.1 2.2 2.3 2.4 2.5 2.6 2.7 2.8 A A A A A A A A Period 3 2.8.1 2.8.2 2.8.3 2.8.4 2.8.5 2.8.6 2.8.7 2.8.8 A A Period 4 2.8.8.1 2.8.8.2 Check in the periodic table of elements in Appendix 1. Where did you place hydrogen and helium in your build up of periodic table? Can you suggest why there is a difference in the one you built up and the ones in the textbook? Hydrogen is a unique element. It is placed in Group I as its atoms have one electron in the outermost energy level. It also fits in Group VII because its atoms have one electron short of a fully filled first energy level. Helium is different from all other elements. It has two electrons in its single energy level which can accommodate a maximum of 2 electrons. Therefore, its outermost energy level is full of the maximum electrons it can possibly accommodate. Later in this unit we will learn more about these Group VIII elements whose outermost energy levels are full with the maximum number of electrons. 1.4 Ion formation Group VIII elements are described as being stable. This means that their outermost energy levels have the maximum number of electrons that they can accomodate as illustrated in Table 1.9. 20 Table 1.9 Electron arrangement of some Group VIII elements Element Symbol Atomic Number Electron arrangement Helium He 2 2 Neon Ne 10 2.8 Argon Ar 18 2.8.8 Octet rule The process of achieving a stable electron arrangement of 2 or 8 electrons in the outermost energy level is known as duplet e.g. in helium or octet rule respectively. In order to achieve a more stable outermost electron arrangement of 2 or 8 electrons, electrons must be gained or lost in the outermost energy level. When an atom gains or loses an electron or electrons, the particle formed is charged either negatively (–ve) or positively (+ve). Such a particle is called an ion. Formation of lithium and magnesium ions Let us consider the formation of lithium ion. A lithium atom has an electron arrangement of 2.1. It has only one electron in the outermost energy level. For it to be stable, it can either lose one electron to have an electron arrangement of 2 or gain seven electrons to attain an arrangement of 2.8. More energy will be required for the atom to gain 7 electrons than to lose 1 electron Fig. 1.13. Therefore, it is easier for a lithium atom to lose 1 electron. + 3p 3p + e– 4n 4n Fig. 1.13: Formation of lithium ion 21  Lithium atom Lithium ion Electron arrangement 2.1 Electron arrangement 2 Nucleus has - 3 protons Nucleus has - 3 protons (positively charged) (positively charged) - 4 neutrons (neutral) - 4 neutrons (neutral) Outside nucleus - 3 electrons Outside nucleus - 2 electrons (negatively charged) (negatively charged) Net charge +3-3=0 Net charge +3-2=+1 Symbol of the atom Li Symbol of the ion Li+1 but represented as Li+ Remember lithium has 3 protons in the nucleus which are positively charged and 4 neutrons which are neutral. The charge in the nucleus is +3. The electrons in the two energy levels are three and are negatively charged. This gives charge of -3 outside the nucleus. The net charge of the atom is +3 – 3 = 0. That is why a lithium atom is neutral. When a lithium atom loses one electron, the charge of the nucleus remains as +3 whereas outside the nucleus the charge reduces to –2. The net charge then becomes +3 –2 = +1. Lithium ion is then represented as Li+ but not Li+1. Let us try again with magnesium For magnesium, it is easier to lose two electrons than gain six electrons. Therefore a magnesium ion is formed by loss of two electrons. 2+ 12p 12p 12n 12n + 2e– Fig. 1.14: Formation of magnesium ion, Mg2+ Magnesium atom Magnesium ion Electron arrangement 2.8.2 Electron arrangement 2.8 Nucleus – 12 protons (positively charged) Nucleus – 12 protons (positively charged) – 12 neutrons (neutral) – 12 neutrons (neutral) Outside the nucleus – 12 electrons Outside the nucleus – 10 electrons (negatively charged) (negatively charged) Net charge: +12 - 12 = 0 Net charge +12 - 10 = +2 Symbol of the atom: Mg Symbol of the ion: Mg2+ (but not Mg+2 or Mg++) 22 Li+ and Mg2+ are called ions. Ions are electrically charged atoms. Li+ and Mg2+ are positively charged ions. They are also called cations. Cations are formed when atoms lose one or more electrons to form positively charged ions. How are sodium and aluminium ions formed? Follow the same procedure as for lithium and magnesium to work this out. Formation of fluoride and sulphide ions A fluorine atom has the electron arrangement of 2.7. For fluorine atom to attain a stable electron arrangement of 2.8, it will either lose seven electrons or gain one electron respectively. Much more energy will be required for fluorine to lose seven electrons than to gain one electron. For this reason it is easier for fluorine to gain one electron than to lose the seven electrons. Therefore, a fluorine atom forms a fluoride ion by gaining one electron. 9p 9p 10n + e– 10n Fig. 1.15: Formation of fluoride ion, F– Flourine atom Flouride ion Electron arrangement 2.7 Electron arrangement 2.8 Nucleus – 9 protons (positively charged) Nucleus – 9 protons (positively charged) – 10 neutrons (neutral) – 10 neutrons (neutral) Outside the nucleus – 9 electrons Outside the nucleus – 10 electrons (negatively charged) (negatively charged) Net charge + 9 – 9 = 0 Net charge –10 + 9 = -1 Symbol of atom: F Symbol of the ion: F— Let us try with sulphur. Sulphur atom has the electron arrangement of 2.8.6. For sulphur, it is easier to gain two electrons than to lose six electrons. 23  2– 16p 16p 16n 16n + 2e- Fig. 1.16: Formation of a sulphide ion, S2– Sulphur atom Sulphide ion Electron arrangement 2.8.6 Electron arrangement 2.8.8 Nucleus –16 protons (positively charged) Nucleus –16 protons (positively charged) –16 neutrons (neutral) –16 neutrons (neutral) Outside the nucleus –16 electrons Outside the nucleus – 18 electrons (negatively charged) (negatively charged) Net charge +16-16 = 0 Net charge -18 +16 = –2 Symbol of atom: S Symbol of the ion: S2– F– and S2– are ions. They are electrically charged. F– and S2– are negatively charged ions. Negatively charged ions are called anions. They are formed when a neutral atom gains one or more electrons. How is a chloride ion formed? Follow the same procedure as for fluorine and sulphur to show the formation of a chloride ion. Table 1.10 Electron arrangement of some ions Symbol of Electron Electron Symbol Atom arrangement of the arrangement atom of ion atom of the ion Lithium Li 2.1 Li+ 2 Sodium Na 2.8.1 Na+ 2.8 Fluorine F 2.7 F– 2.8 Chlorine Cl 2.8.7 Cl– 2.8.8 Aluminium Al 2.8.3 Al3+ 2.8 Magnesium Mg 2.8.2 Mg2+ 2.8 Sulphur S 2.8.6 S2– 2.8.8 24 Ionisation energy We have seen from the structure of atoms that protons in the nucleus which are positively charged (+ve) attract electrons which are negatively charged (–ve) and located in the energy levels. Therefore, in order to remove electron, we must overcome this force of attraction. In other words we must supply energy to pull off the electron(s). The energy supplied is called ionisation energy. The unit for this energy is called joule(J) which can be converted to kilojoules(kJ). Ionisation energy is defined as the energy required to remove electron(s) from an atom in gaseous state to produce an ion. The equation for the loss of an electron is represented as: M(g) → M+(g) + e– M-represents a metal atom If an atom loses 2 electrons the equation is as follows: M(g) → M2+(g) + 2e– Note the charge on the ion is the same as the number of electrons lost. If an atom loses three electrons, then the charges will be 3+ M(g) → M3+(g) + 3e– Electron affinity We have seen that non-metals gain electron(s) to gain stability as illustrated in Table 1.10. Remember that electrons are negatively charged. Therefore when an electron tries to enter the outermost energy level, it will be repelled by the electrons which are already there. So “force” is required to put the electrons into the energy level. This “force” is a form of energy. This energy is known as electron affinity. The ions formed are always negatively charged. For example: Cl + e– → Cl– S + 2e– → S2– Oxidation numbers We mentioned earlier, an atom will gain or lose electrons according to the number of electrons in the outermost energy level. In the process the atom becomes an ion. The oxidation number of an ion is simply the charge on the ion. + or – sign is written before the number unlike the charge. See Table 1.11 25  Table 1.11 Oxidation numbers Ion Oxidation number H+ +1 + Na +1 K+ +1 Ca2+ +2 Al3+ +2 Cl– –1 F– –1 O2– –2 Why is the oxidation number of H, +1, yet it is a non-metal? Sometimes a hydrogen atom loses an electron like a metal. During other reactions it gains an electron to achieve the helium electron arrangement. Table 1.12 Periodic table with some common oxidation numbers I II III IV V VI VII VIII H He +1 0 Li Be O F Ne +1 +2 –2 –1 0 Na Mg Al S Cl Ar +1 +2 +3 –2 –1 0 K Ca +1 +2 From the periodic table above there are similarities of the oxidation number of elements that belong to the same group. For example, Group I elements have oxidation of +1 whereas group seven elements have oxidation numbers of -1 etc. Valencies The number of electrons an atom requires to attain the stable noble gas electron arrangement is known as valency or combining power of the atom or group of atoms. A group of atoms which occur in compounds but cannot exist on its own is called a radical, for example sulphate, SO 42– , hydroxide, OH–, etc. 26  The oxidation number has a negative or positive sign unlike valency which does not have. In metals the valency is just the number of electrons in the outermost energy level. For non-metals it is the difference between group 8 and group number of the elements, e.g. the valency of oxygen is 8 – 6 = 2, valency of phosphorous is 8 – 5 = 3. Table 1.13 Valencies of the first twenty elements in the periodic table Atomic Element Symbol Electron To gain stability of a noble gas Valency Number arrangement atoms gain or lose electrons 1 Hydrogen H 1 Not stable, can lose 1e– 1 2 Helium He 2 Stable, cannot lose or gain 0 3 Lithium Li 2.1 Not stable, loses 1e– 1 4 Beryllium Be 2.2 Not stable, loses 2e– 2 5 Boron B 2.3 Not stable, loses 3e– 3 6 Carbon C 2.4 Not stable, loses or gains 4e– 4 7 Nitrogen N 2.5 Not stable, easy to gain 3e– 3 8 Oxygen O 2.6 Not stable, easy to gain 2e– 2 9 Fluorine F 2.7 Not stable, easy to gain 1e – 1 10 Neon Ne 2.8 Stable, cannot gain or lose electron 0 11 Sodium Na 2.8.1 Not stable, loses 1e– 1 12 Magnesium Mg 2.8.2 Not stable, loses 2e– 2 13 Aluminium Al 2.8.3 Not stable, loses 3e– 3 14 Silicon Si 2.8.4 Not stable, loses 4 or gains 4e– 4 15 Phosphorous P 2.8.5 Not stable, gains 3e– 3 or 5 16 Sulphur S 2.8.6 Not stable, gains 2e– 2 17 Chlorine Cl 2.8.7 Not stable, loses 1e– 1 18 Argon Ar 2.8.8 Stable (octet) cannot gain or lose electrons 0 19 Potassium K 2.8.8.1 Not stable, loses 1e– 1 20 Calcium Ca 2.8.8.2 Not stable, loses 2e– 2 Table 1.14 Valencies of the first twenty elements in the periodic table I II III IV V VI VII VIII H He 1 0 Li Be B C N O F Ne 1 2 3 4 3 2 1 0 Na Mg Al Si P S Cl Ar 1 2 3 4 3 or 5 2 1 0 K Ca 1 2 27  Note that valency corresponds to the group number for metals. For non-metals subtract group number from 8. Table 1.15 Valencies of other common elements Name of metal Symbol Valency Zinc Zn 2 Iron Fe 2 or 3 Tin Sn 2 or 4 Lead Pb 2 or 4 Copper Cu 1 or 2 Silver Ag 1 Barium Ba 2 Table 1.16 Valencies of common radicals Valency 1 Valency 2 Valency 3 Radical Formula Radical Formula Radical Formula Ammonium NH4+ Carbonate CO2–3 Phosphate PO3– 4 Hydroxide OH– Sulphate SO42– Nitrate NO3– Sulphite SO2–3 Chloride Cl– Hydrogencarbonate HCO3– Hydrogensulphate HSO–4 We already mentioned that valency is combining power of an element or radical. However some elements have varying valencies, for example, copper can have valency of 1 or 2. The valencies of these elements are indicated in Roman numbers in brackets after the name of the element when naming compounds. This is illustrated in Table 1.17. 28 Table 1.17 Valencies of elements in compounds Compound Element Valency Copper(I) oxide Copper 1 Copper(II) oxide Copper 2 Iron(II) sulphate Iron 2 Iron(III) chloride Iron 3 Sulphur(IV) oxide Sulphur 4 Sulphur(VI) oxide Sulphur 6 Carbon(IV) oxide Carbon 4 Carbon(II) oxide Carbon 2 Writing chemical formulae Every subject has its own special language as you might have realized by now. For example, in Chemistry we have substances with names like sodium chloride. However, we have a short notation of writing names called formula. The formula of a compound shows the atoms present in the compound in their simplest ratio. The formula of sodium chloride is NaCl. The formula for water is H2O. Why then do we write the number 2 in the formula for water whereas we do not in sodium chloride? These are some of the questions we shall answer in this section as we discuss how to write chemical formulae. Activity 1.2 Writing chemical formulae using a game of cards Procedure 1. Get pieces of plain paper. 2. For group I elements or cations, with valency 1, cut the pieces of paper into the following shape but make sure that they are of the same size. Indicate the symbol and the name of the element. only one projection representing valency 1 symbol and name Fig. 1.17: Shape representing a cation with valency 1 29  3. For group II elements or cations with valency 2, cut the paper such that it has two projections as shown in the figure below. Two projections symbol and name representing valency 2 Fig. 1.18: Shape representing a cation with valency 2 4. For group III elements or cations with valency 3, your piece of paper will have three projections as shown in Fig. 1.19. Three symbol and name projections representing valency 3 Fig. 1.19: Shape representing a cation with valency 3 5. For group VII elements or anions with valency 1, cut the piece of paper so that it has one groove as shown in Fig. 1.20 One groove representing symbol and name valency 1 Fig. 1.20: Shape of anion with valency 1 6. Group VI elements or anions with valency 2, will be represented by pieces of paper with two grooves. Two grooves representing symbol and name valency 2 Fig. 1.21: Shape representing an anion with valency 2 30 7. For elements with valency 3, cut out three grooves. Three grooves symbol and name representing valency 3 Fig. 1.22: Shape representing an anion with valency 3 Note: It is important to take care of the following point before you play this game. When writing the formula, metals are always written first. For example it is not correct to write ClNa or Br2Ca even though it gives us the same information. The correct way is NaCl and CaBr2 respectively. Examples (a) Write the chemical formula of sodium sulphate. From the above activities, get the sodium and the sulphate cards as in step 1. Step 1 Na Sodium SO4 Sulphate Fig. 1.23: Sodium and sulphate cards Step 2 Fix the two cards together such that the protruding side of the sodium fits into the groove of the sulphate. Na Sodium SO4 Sulphate Fig. 1.24: Joined sodium and sulphate cards Notice that one groove of the sulphate is still free. For it to be complete, you need another card of sodium. 31  Step 3 Fix one other card of sodium Na Sodium SO4 sulphate Na Sodium Na Sodium SO4 sulphate Na Sodium Fig. 1.25: Complete joined cards representing sodium sulphate You can see that we need two cards of sodium and one card of sulphate. The chemical formula of sodium sulphate is therefore Na2SO4. (b) Write the formula of potassium hydroxide. K OH Potassium hydroxide K OH Potassium hydroxide Fig. 1.26: Representation of potassium hydroxide with the cards We need one potassium card and one hydroxide card. So the formula of potassium hydroxide is KOH. Now let us use another simple method of writing chemical formulae. Always remember that to write a correct formula we must write down the; (a) correct symbol of the element or radical (b) correct valency of the symbol or radical. Let us write the formula for sodium sulphate again following the steps below: 32 Step 1 Write the symbols of the elements and radical. Na SO4 Step 2 Write the valencies of the element and radical above and to the right side of each. Na1 SO24 Step 3 Exchange the valencies by writing them below the symbols as shown by the arrows. 2 Na14 SO4 Step 4 Write the symbols close together Na2SO4 Other examples (a) Lithium oxide Symbols Li O Valencies 1 2 2 Formula Li1 O Li2O (b) Sodium chloride Symbols Na Cl Valencies 1 1 1 1 Formula Na Cl NaCl 33  (c) Calcium chloride Symbols Ca Cl Valencies 2 1 1 Formula Ca2 Cl CaCl2 (d) Ammonium sulphate Symbols NH4 SO4 Valencies 1 2 2 Formula NH14 SO (NH4)2 SO4 (e) Magnesium hydroxide Symbols Mg OH Valencies 2 1 1 Formula Mg2 OH Mg(OH)2 Note: We write brackets for (d) and (e) because the radical consists of two different elements eg. OH contains oxygen atom and hydrogen atom; NH4 consists of nitrogen and hydrogen atoms. But it is wrong to put bracket for (c) i.e. Ca(Cl)2 as “C” and “l do not represent different elements. (f) Write the formula of lead (IV) oxide. Symbol Pb O Valencies 4 2 4 2 Formula Pb O Note: Before we bring the valencies down, we must divide the valencies by the common factor, 2. PbO2 is the formula 34 (g) Write the formula of iron(III) oxide Here the valencies have no common number which can divide them and give us a whole number. Therefore, we just exchange the 2 valencies as shown Fe3 O The formula is Fe2O3. Use the card game for the above worked examples and to see whether you get the same formulae. Exercises Write the formulae for the following compounds 1. Calcium carbonate 2. Sodium phosphate 3. Ammonium phosphate 4. Lead(II) sulphate 5. Iron(II) sulphate 6. Iron(II) hydroxide 7. Lead nitrate 8. Manganese(IV) oxide 9. Copper(II) hydroxide 10. Iron(III) chloride Writing simple balanced chemical equations Chemical equations are short, clear and accurate descriptions of chemical reactions. A reaction process can be explained using an equation. For example, when oxygen reacts with magnesium ribbon, a white solid of magnesium oxide is formed. We have been using a word equation to describe such a reaction. Thus: Magnesium + oxygen → Magnesium oxide The (+) sign here is not used to mean addition, but in chemistry it is used to mean ‘reacts with’. The → sign is used to indicate formation of a product. Using equal sign instead of an arrow, is wrong. Names of starting substances like magnesium and oxygen in the above example are written on the left side of the arrow; these substances are called reactants. The new substances produced by the chemical reaction are called products and are written on the right side of the arrow. Magnesium + oxygen → Magnesium oxide (Reactants) (Product) 35  Chemical equations have various notations that indicate the physical states of the reactants and products. These notations are very important. Infact, whenever one writes an equation and misses to write them, the equation is not complete. These notations are as follows: Table 1.20 State symbols Physical state Representation of state Description Solid (s) a solid can be a precipitate, suspension, etc Liquid (l) a pure liquid like water, paraffin, etc Aqueous (aq) a solute or liquid dissolved in water Gas (g) a gas or vapour Balancing chemical equations In order for an equation to describe a reaction accurately, the equation must be balanced. Chemical reactions should always follow the law of conservation of mass so that the total mass of reactants must be equal to the total mass of products. Reactants Products Fig. 1.27: Conservation of mass in a balanced chemical reaction Let us consider our previous example of magnesium ribbon reacting with oxygen. We can write an initial equation containing the formulae of reactants and products. Mg + O2 → MgO This equation is not balanced. Count the number of atoms of the reactants and the products. REACTANTS PRODUCT 1 Magnesium atom 1 Magnesium atom 2 Oxygen atoms 1 Oxygen atom The left side (reactants) has two oxygen atoms but the right side (product) has only one oxygen atom. 36  MgO Mg + O 2 Fig. 1.28: Unbalanced equation, more oxygen atoms on the left An equation can be balanced using a number of rules. Let us follow these rules to balance the above equation. Rule number 1 Write the equation using correct formulae for the reactants and products Magnesium + oxygen → Magnesium oxide Mg + O2 → MgO Rule number 2 Count the number of atoms of each element in the reactants and in the products. Check whether they are equal as in Table 1.21. Table 1.21 Balancing number of atoms for each element in the reactants and product Atoms Reactants Product MG 1 1 O 2 1 We notice that the oxygen atoms are not equal. We have 2 atoms on the left and only one on the right. Rule number 3 To make oxygen atoms equal, balance the equation by writing numbers in front of the formula. Remember that 1 is assumed to be there already. Therefore start by inserting 2. If this does not balance, go to 3, 4 until the equation is balanced. Usually we do not go to very big numbers. 2Mg + O2 → 2MgO The number 2 now balances the equation. When you have 2O2 it means two oxygen molecules. The number in front of a formula means, everything following is multiplied by that number. 37  For example: 2Mg means 2 × Mg that is why we have 2 Mg atoms on the left and 2 Mg on the right. This equation 2Mg + O2 → 2MgO means 2 atoms of Mg react with 1 molecule of oxygen (containing 2 atoms) to form 2 molecules of magnesium oxide. Step number 4 Count again the number of atoms of each element on the reactants and product sides. Note that all atoms are balanced as illustrated in Table 1.22. Table 1.22: Balancing number of atoms for each element in the reactants and products Atoms Reactants Product Mg 2 2 O 2 2 Step number 5 Insert the correct state symbols for each substance. 2Mg(s) + O2(g) → 2MgO(s) Let us practice the above steps by writing equations for the following reactions. Example Zinc granules react with dilute hydrochloric acid to produce aqueous zinc chloride and hydrogen gas. Step 1 Write an initial equation using correct chemical formulae of reactants and products. Zinc + hydrochloric acid → zinc chloride + hydrogen Zn + HCl → ZnCl2 + H2 Step 2 Count atoms 38 Table 1.23 Balancing the number of atoms for each element in the reactants and products Atoms Reactants Products ZN 1 1 H 1 2 CL 1 2 Note: If you look at the reactants and products, zinc is already balanced but hydrogen and chlorine are not. We need to have 2 hydrogen atoms and 2 chlorine atoms on the reactants side (left side). Step 3 Insert a number infront of the formula to balance the equation. Zn + 2HCl → ZnCl2 + H2 Note: A number of 2 has been put infront of HCl at the left side to make sure that the balance is achieved. Let us confirm in the next step. Step 4 Count the atoms again. Table 1.24 Balancing the number of atoms for each element in the reactants and products Atoms Reactants Products ZN 1 1 H 2 2 CL 2 2 Step 5 Insert correct state symbols for each substance. Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g) Solid zinc reacts with dilute hydrochloric acid to form aqueous zinc chloride and hydrogen gas. 39  Exercises Write the correct balanced equations for the following reactions. 1. Reaction of copper(II) oxide with hydrogen to produce copper and water. 2. Reaction of solid calcium carbonate with dilute hydrochloric acid to produce aqueous calcium chloride, carbon(IV) oxide and water. 3. Reaction of solid copper(II) oxide with dilute sulphuric acid to produce aqueous copper(II) sulphate and water. 4. Reaction of aqueous aluminium chloride with aqueous sodium hydroxide to produce aluminum hydroxide and aqueous sodium chloride. 5. Reaction of solid copper carbonate with dilute hydrochloric acid to produce aqueous copper(II) chloride, carbon(IV) oxide and water. Projects Construction of cards for chemical formulae. Construction of models of the periodic table. Summary 1. Atoms are the smallest particles of an element. 2. The central part of the atom is the nucleus. 3. The nucleus contains smaller particles called protons and neutrons. 4. The number of protons = the number of electrons. 5. Protons are positively charged while electrons are negatively charged. Neutrons have no charge. 6. Electrons move around the nucleus along energy levels. 7. The atomic number (Z) refers to the number of protons in the nucleus of an atom. 8. The mass number (A) is the sum of protons and neutrons. 9. Number of neutrons = mass number – atomic number. 10. The energy levels have different energies and are at different distances from the nucleus of the atom. 11. The energy level next to the nucleus has the least amount of energy and can accomodate a maximum of two electrons. 12. The other energy levels if it is the outermost, can accomodate a maximum number of eight electrons. 13. The electrons in the outermost energy level are very important in determining the chemistry of the element. 14. Isotopes are atoms of the same element. 15. Isotopes have the same number of protons and electrons in each atom, but have different numbers of neutrons in their nuclei. 16. Since isotopes have same number of electrons, they have same or nearly same chemical properties. 40  17. Nearly all the mass of an atom is concentrated in the nucleus. The mass of a proton and neutron are approximately equal. The mass of an electron is negligible. 18. The relative atomic mass of an element is the number of times the 1 average mass of one of its atoms is greater than 12 of the mass of one atom of carbon – 12. 19. Some relative atomic masses are not whole numbers because the relative abundance of the various isotopes is not equal. 20. The elements in the periodic table are arranged in order of increasing atomic number. 21. Elements in the same group have the same number of electrons in their outermost energy level. 22. Elements in the same period have the electrons added to the same energy level. 23. The valency is the combining power of an atom or group of atoms (radicals). Revision Exercise 1 1. Here is a symbol of lithium atom: 7 3 Li (a) What is the name given to the superscript (top number)? (b) What is the name given to the subscript (bottom number)? (c) In the lithium atom shown above, calculate the number of protons, neutrons and electrons in the atom. 2. Select the pair which represents two atoms with the same number of neutrons. A. 12C and 24 Mg 6 12 B. 18O and 19F 8 9 20 11 Na C. 23 and 10 Ne 3. Explain briefly why some elements have relative atomic masses which are not whole numbers. 4. An atom is the smallest particle of an element. Name the sub-atomic particles found in an atom and state where they are found. 5. Explain the following terms: 41  (i) Atomic number (ii) Mass number (iii) Isotopes 6. The following symbols refer to isotopes of oxygen. (i) 168 O (ii) 18 8 O What is the number of protons and neutrons in the nucleus of each isotope? 7. A class was asked to select from the list below the elements whose electron arrangements were impossible to write down. Which elements did the class select? Explain your answer. 10 23 11 Na ; B; 6C; 12 6 C; 25 Mg; 26 12 Mg 8. Write down and complete the following table. Atom Mass Number Number Number number of of of protons neutrons electrons F 19 9 Na 12 11 Ne 20 10 C 12 6 16 9. An atom of element X can be represented by the symbol 6 X. (a) State the: (i) number of electrons of atom X. (ii) mass number of X (iii) atomic number of X. (b) Element X contains 90% of 166X and 10% of 186X. Calculate the relative atomic mass of X. 10. Put the following letters A to H in the correct places in the periodic table above to fit these descriptions. 42  Groups I II III IV V VI VII VIII Periods 1 2 3 4 A. An element with 7 protons. B. It belongs to period three and has 4 electrons in the outermost energy level. C. An element with oxidation number +3 and with 13 protons. D. An element that forms its ion by loss of one electron and belongs to period four. E. An element that can be placed in either group I or group VII. F. An element with the highest number of energy levels. G. An element with electronic arrangement of 2.8.8. H. An element with valency of 3. It gains electrons to form ions and belongs to period three. 43 UNIT 2: Chemical families; patterns in properties 2.1 Introduction The physical and chemical properties of elements and their compounds are highly dependent upon the number of electrons that are found in their outermost energy level. When elements are arranged in the periodic table in order of increasing atomic number, a regular change in the outermost electron arrangement is observed. In turn a regular variation of properties is also observed. Thus, as we move across a period certain repeated patterns can be observed. These lead us to make predictions and intelligent guesses about the unknown properties of elements and their compounds. In addition, elements that are in the same group have the same number of electrons in the outermost energy level and have similar chemical properties. These facts lead to the observation of gradual changes in the physical and chemical properties of the elements. In order to illustrate these changes we will consider a number of trends and patterns observed in the periodic table. 2.2 Group I elements - Alkali metals Group 1 elements are metals. They are called alkali metals and are the most reactive group of metals. They are found on the left hand side of the periodic table as illustrated in Fig 2.1. GROUP I II 7 Li Be 3 Lithium 23 11Na Mg Sodium 39 K Ca 19 Potassium 86 Rb Sr 37 Rubidium 133 Cs Ba 55 Caesium 223 Fr Ra 87 Francium Fig 2.1 Alkali metals 44  Note: At this level we are going to study only lithium, sodium and potassium. 2.3 Trends in physical properties of alkali metals Experiment 2.1 To investigate the physical appearance of alkali metals. Apparatus and chemicals tile (ceramic) filter paper knife pair of tongs lithium sodium potassium Procedure 1. Using a pair of tongs, remove a small piece of lithium from the bottle. 2. Place it on the tile. Describe the appearance of the lithium from the bottle? 3. Cut the piece of lithium into two with a knife to expose the inside. What is the colour of the freshly cut part of lithium. 4. Repeat the above procedure with sodium and potassium and answer the same questions. 5. Record your observations in a table like the one shown below. What do you conclude about the appearance of the group I element? Why does the appearance of the metal change after a short while? Table 2.1 Physical properties of alkali metals Appearance of metal Appearance of freshly from bottle cut metal Lithium Sodium Potassium 45  Remove v metal from v Observe bottle (a) Metal in bottle (b) Cut metal (c) Exposed inner part of metal Fig. 2.2 Physical appearance of alkali metals Table 2.2 Physical properties of alkali metals Electron Melting Boiling Thermal Electrical Element Symbol Physical Physical state appearance arrangement point point conductivity conductivity Lithium Li Solid Silvery 2.1 181˚C 134˚C Good Good metal conductor conductor of of heat electricity Sodium Na Solid Silvery 2.8.1 98˚C 883˚C Good Good metal conductor conductor of of heat electricity Potassium K Solid Silvery 2.8.8.1 63˚C 759˚C Good Good metal conductor conductor of of heat electricity State and explain the trends in physical properties of alkali metals. We saw in the previous section that alkali metals have one electron in the outermost energy level. Because of this, alkali metals have very similar physical and chemical properties. They also show characteristic trends in these properties as one goes down the group. Typical metal properties Group I elements show typical metal properties such as: good conductivity of heat good conductivity of electricity high boiling points shiny surface when freshly cut. Other physical properties of alkali metals Low melting points Very low density (they float on water) Very soft and can be cut with an ordinary knife easily. 46 Important trends down the group. The melting points and boiling points generally decrease. The density generally increases. Atomic radius The atomic radius is the distance between the nucleus and the outermost energy level in an atom. It is sometimes called atomic size. v Atomic radius Fig. 2.3 Atomic radius of sodium Of the three group 1 elements, namely, lithium, sodium and potassium, lithium has the smallest atomic radius followed by sodium. Potassium has the largest atomic radius. The atomic radius increases as we go down the group as illustrated in Fig. 2.4. I II III IV V VI VII VIII or O H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Fig. 2.4 Periodic variation of atomic radii of the 1st twenty elements of the periodic table. From Fig 2.4, we notice that the atomic radius generally increases down a group. Why do you think this happens? Before we answer this question let us first look at Fig. 2.5. 47  I II III IV V VI VII VIII or O H