Burning is a reaction of an element with Oxygen The substance formed when an element burn in air is the oxide of the element When hydrogen burns it reacts combines with Oxygen to form the oxide of Hydrogen The oxide of Hydrogen is called water Hydrogen is first dried because a mixture of Hydrogen and air explode The gas is then ignited The products condense on a cold surfaceflask containing a freezing mixture A freezing mixture is a mixture of water and ice


The condensed products are collected in a receiver as a colourless liquid


Tests


a When about 1g of white anhydrous copper II sulphate VI is added to a sample of the liquid it turns to blue This confirms the liquid formed is water


b When blue anhydrous cobalt II chloride paper is dipped in a sample of the liquid it turns to pink This confirms the liquid formed is water


cWhen the liquid is heated to boil its boiling point is 100oC at sea levelone atmosphere pressure This confirms the liquid is pure water


Uses of Hydrogen gas


1 HydrogenationHardening of unsaturated vegetable oils to saturated fatsmargarine


When Hydrogen is passed through unsaturated compounds in presence of Nickel catalyst and about 150oC they become saturated Most vegetable oil is unsaturated liquids at room temperature They become saturated and hard through hydrogenation


2 In weather forecast balloons


Hydrogen is the lightest known gas Meteorological data is collected for analysis by sending hydrogen filled weather balloons to the atmosphere The data collected is then used to forecast weather conditions


3 In the Haber process for the manufacture of Ammonia


Hydrogen is mixed with Nitrogen in presence of Iron catalyst to form Ammonia gas Ammonia gas is a very important raw material for manufacture of agricultural fertilizers


4 In the manufacture of Hydrochloric acid


Limited volumeamount of Hydrogen is burnt in excess chlorine gas to form Hydrogen chloride gas Hydrogen chloride gas is dissolved in water to form Hydrochloric acid Hydrochloric acid is used in picklingwashing metal surfaces


5 As rocket fuel


Fixed proportions of Hydrogen and Oxygen when ignited explode violently producing a lot of energyheat This energy is used to powerpropel a rocket to space


6 In oxyhydrogen flame for welding


A cylinder containing Hydrogen when ignited in pure Oxygen from a second cylinder produces a flame that is very hot It is used to cut metals and welding


Sample revision questions


1 A colourless liquid was added anhydrous copper II sulphate VI which turned blue


aWhy is it wrong to conclude the liquid was pure water


Anhydrous copper II sulphate VI test for presence of water Purity of water is determined from freezingmeltingboiling point


bWrite an equation for the reaction that takes place with anhydrous copper II sulphate VI


Anhydrous copper II sulphate VI  Water  hydrated copper II sulphate VI


CuSO4s  5H2O l  CuSO45H2Os


ciWhich other compound would achieve the same results as anhydrous copper II sulphate VI


Anhydrous cobalt II chlorideCoCl26H2O


iiWrite the equation for the reaction


Anhydrous cobalt II chloride  Water  hydrated cobalt II chloride


CoCl2 s  6H2Ol  CoCl26H2O s


dComplete the equation


i Sulphur VI oxide  Water  Sulphuric VI acid


ii Sulphur IV oxide  Water  Sulphuric IV acid


iii Carbon IV oxide  Water  Carbonic IV acid


iv Nitrogen IV oxide  Water  Nitric V acid


v Phosphorus V oxide  Water  Phosphoric V acid


vi Sodium oxide  Water  Sodium hydroxide


vi Sodium peroxide  Water  Sodium hydroxide


2 Metal B reacts with steam Metal C reacts with cold water Metal A does not react with water


aArrange the metals as they should appear in the reactivity series


B


C


A


bA product residue in D which was brown when hot but turned yellow on cooling during the reaction of metal B was formed Gas E was also evolved Identify


iMetal B LeadPb


iiResidue D Lead II oxidePbO


iiiGas E HydrogenH2


cA portion of product residue in D was added dilute nitric V acid Another portion of product residue in D was added dilute sulphuric VI acid State and explain the observations made


When added dilute nitric V acid D dissolves to form a colourless solution


Lead II Oxide  dilute nitric V acid  Lead II nitrate V  Water


PbO s  2HNO3 aq  Pb NO32 aq  H2O l


When added dilute sulphuric VI acid D does not dissolve A white suspensionprecipitate was formed LeadIIOxide reacts with sulphuricVIacid to form insoluble LeadIIsulphateVI that covercoat unreacted LeadIIOxide stopping further reaction


Lead II Oxide  dilute sulphuric VI acid  Lead II sulphate VI  Water


PbO s  H2SO4 aq  PbSO4 s  H2O l


3 a Hydrogen can reduce copper II Oxide but not alluminium oxide Explain


b When water reacts with potassium metal the hydrogen produced ignites explosively on the surface of water


i What causes this ignition ii Write an equation to show how this ignition occurs


2 In an experiment dry hydrogen gas was passed over hot copper II oxide in a combustion tube as shown in the diagram below


a Complete the diagram to show how the other product substance R could be collected in the laboratory


b Describe how copper could be obtained from the mixture containing copper II oxide


3 The setup below was used to investigate the reaction between metals and water


a Identify solid X and state its purpose


Solid X


Purpose


b Write a chemical equation for the reaction that produces the flame 4 Gas P was passed over heated magnesium ribbon and hydrogen gas was collected as shown in the diagram below


i Name gas P


ii Write an equation of the reaction that takes place in the combustion tube iii State one precaution necessary at the end of this experiment


5 When hydrogen is burnt and the product cooled the following results are obtained as shown in the diagram below


a Write the equation for the formation of liquid Y


b Give a chemical test for liquid Y


Jane setup the experiment as shown below to collect a gas The wet sand was heated before


heating Zinc granules


a Complete the diagram for the laboratory preparation of the gas b Why was it necessary to heat wet sand before heating Zinc granules


a Between N and M which part should be heated first Explain


b Write a chemical equation for the reaction occurring in the combustion tube


8 The setup below was used to investigate electrolysis of a certain molten compound


a Complete the circuit by drawing the cell in the gap left in the diagram


b Write halfcell equation to show what happens at the cathode


c Using an arrow show the direction of electron flow in the diagram above


9 Hydrogen can be prepared by reacting zinc with dilute hydrochloric acid


a Write an equation for the reaction


b Name an appropriate drying agent for hydrogen gas


c Explain why copper metal cannot be used to prepare hydrogen gas


d Hydrogen burns in oxygen to form an oxide


i Write an equation for the reaction


ii State two precautions that must be taken before the combustion begins and at the end of the combustion


e Give two uses of hydrogen gas


f When zinc is heated to redness in a current of steam hydrogen gas is obtained Write an equation for the reaction


g Element Q reacts with dilute acids but not with cold water Element R does not react with dilute acids Elements S displaces element P from its oxide P reacts with cold water Arrange the four elements in order of their reactivity starting with the most reactive


h Explain how hydrogen is used in the manufacture of margarine


10 a The setup below is used to investigate the properties of hydrogen


On the diagram indicate what should be done for the reaction to occur


Hydrogen gas is allowed to pass through the tube for some time before it is lit Explain


iii Write an equation for the reaction that occurs in the combustion tube iv When the reaction is complete hydrogen gas is passed through the apparatus until they cool down Explain


v What property of hydrogen is being investigated


vi What observation confirms the property stated in v above


vii Why is zinc oxide not used to investigate this property of hydrogen gas


11 The set up below was used to collect gas K produced by the reaction between water and


calcium metal


a Name gas K


b At the end of the experiment the solution in the beaker was found to be a weak base Explain why the solution is a weak base


ACIDS BASES AND INDICATORS


INTRODUCTION TO ACIDS BASES AND INDICATORS


1 In a school laboratory


iAn acid may be defined as a substance that turns litmus red


iiA base may be defined as a substance that turns litmus blue


Litmus is lichen found mainly in West Africa It changes its colour depending on whether the solution it is in is basicalkaline or acidic It is thus able to identifyshow whether another substance is an acid base or neutral


iiiAn indicator is a substance that shows whether another substance is a basealkalineacid or neutral


2 Common naturally occurring acids include


3 Most commonly used acids found in a school laboratory are not naturally occurring They are manufactured They are called mineral acids


Common mineral acids include


4 Mineral acids are manufactured to very high concentration They are corrosive causes painful wounds on contact with the skin and attackreacts with garmentsclothesmetals


In a school laboratory they are mainly used when added a lot of water This is called diluting Diluting ensures the concentration of the acid is safely low


5 Bases are opposite of acids Most bases do not dissolve in water


Bases which dissolve in water are called alkalis


Common alkalis include


Common bases which are not alkali include


6 Indicators are useful in identifying substances which lookalike


An acidbase indicator is a substance used to identify whether another substance is alkaline or acidic


An acidbase indicator works by changing to different colors in neutral acidic and alkaline solutionsdissolved in water


Experiment To prepare simple acidbase indicator


Procedure


aPlace some flowers petals in a mortar Crush them using a pestle Add a little sand to assist in crushing


Add about 5cm3 of propanoneethanol and carefully continue grinding


Add more 5cm3 of propanoneethanol and continue until there is enough extract in the mortar


Filter the extract into a clean 100cm3 beaker


bPlace 5cm3 of filtered wood ash soap solution ammonia solution sodium hydroxide hydrochloric acid distilled water sulphuric VI acid sour milk sodium chloride toothpaste and calcium hydroxide into separate test tubes


cPut about three drops of the extract in ato each test tube in b Record the observations made in each case


Sample observations


The plant extract is able to differentiate between solutions by their nature It is changing to a similar colour for similar solutions


iSince lemon juice is a known acid then sulphuric VI and hydrochloric acids are similar in nature with lemon juice because the indicator shows similar colors They are acidic in nature


iiSince sodium hydroxide is a known basealkali then the green colour of indicator shows an alkalinebasic solution


iii Since pure water is neutral then the orange colour of indicator shows neutral solutions


7 In a school laboratory commercial indicators are used A commercial indicator is cheap readily available and easy to store Common indicators include Litmus phenolphthalein methyl orange screened methyl orange bromothymol blue


Experiment


Using commercial indicators to determine acidic basicalkaline and neutral solutions


Place 5cm3 of the solutions in the table below Add three drops of litmus solution to each solution


Repeat with phenolphthalein indicator methyl orange screened methyl orange and bromothymol blue


Sample results


From the table above then the colour of indicators in different solution can be summarized


The universal indicator


The universal indicator is a mixture of other indicator dyes The indicator uses the pH scale The pH scale shows the strength of bases and acids The pH scale ranges from 114These numbers are called pH values


i pH values 1 2 3 shows a substance is strongly acid


ii pH values 4 5 6 shows a substance is a weakly acid


iii pH value 7 shows a substance is a neutral


iv pH values 8 9 10 11 shows a substance is a weak basealkali


v pH values 12 13 14 shows a substance is a strong basealkali


The pH values are determined from a pH chart The pH chart is a multicolored paper with each colour corresponding to a pH valueie


i red correspond to pH 1 2 3 showing strongly acidic solutions


iiOrange yellow correspond to pH 4 5 6 showing weakly acidic solutions


iiiGreen correspond to pH 7 showing neutral solutions


ivBlue correspond to pH 8 9 10 11 showing weakly alkaline solutions


vPurpledark blue correspond to pH 121314 showing strong alkalis


The universal indicator is available as


i Universal indicator paperpH paper


ii Universal indicator solution


When determining the pH of a unknown solution using


i pH paper then the pH paper is dipped into the unknown solution It changesturn to a certain colour The new colour is marchedcompared to its corresponding one on the pH chart to get the pH value


ii universal indicator solution then about 3 drops of the universal indicator solution is added into about 5cm3 of the unknown solution in a test tube It changesturn to a certain colour The new colour is marchedcompared to its corresponding one on the pH chart to get the pH value


Experiment To determine the pH value of some solutions


aPlace 5cm3 of filtered wood ash soap solution ammonia solution sodium hydroxide hydrochloric acid distilled water sulphuric VI acid sour milk sodium chloride toothpaste and calcium hydroxide into separate test tubes


bPut about three drops of universal indicator solution or dip a portion of a piece of pH paper into each Record the observations made in each case


cCompare the colour in each solution with the colors on the pH chart provided Determine the pH value of each solution


Sample observations


Note


1 All the mineral acids Hydrochloric sulphuric VI and nitric V acids are strong acids


2 Two alkalissoluble bases sodium hydroxide and potassium hydroxide are strong basesalkali Ammonia solution is a weak basealkali All other bases are weakly alkaline


3 Puredeionized water is a neutral solution


4 Common saltsodium chloride is a neutral salt


5 When an acid and an alkalibase are mixed the final product has pH 7 and is neutral


Properties of acids


aPhysical properties of acids


1 Acids have a characteristic sour taste


2 Most acids are colourless liquids


3 Mineral acids are odorless Organic acids have characteristic smell


4 All acids have pH less than 7


5 All acids turn blue litmus paper red methyl orange red and phenolphthalein colourless


6 All acids dissolve in water to form an acidic solution Most do not dissolve in organic solvents like propanone kerosene tetrachloromethane petrol


bChemical properties of acids


1 Reaction with metals


All acids react with reactive metals to form a salt and produce evolve hydrogen gas


Metal  Acid  Salt  Hydrogen gas


Experiment reaction of metals with mineral acids


aPlace 5cm3 of dilute hydrochloric acid in a small test tube Add 1cm length of polished magnesium ribbon Stopper the test tube using a thump Light a wooden splint Place the burning splint on top of the stoppered test tube Release the thump stopper Record the observations made


bRepeat the procedure in a above using Zinc granules iron filings copper turnings aluminum foil in place of Magnesium ribbon


cRepeat the procedure in a then b using dilute sulphuric VI acid in place of dilute hydrochloric acid


Sample observations


i effervescencebubbles producedfizzing in all cases except when using copper


ii Colourless gas produced in all cases except when using copper


iii Gas produced extinguishes a burning wooden splint with an explosionpop sound


Explanation


Some metals react with dilute acids while others do not Metals which react with acids produce bubbles of hydrogen gas Hydrogen gas is a colourless gas that extinguishes a burning splint with a pop sound This shows acids contain hydrogen gas


This hydrogen is displacedremoved from the acids by some metals like Magnesium Zinc aluminium iron and sodium


Some other metals like copper silver gold platinum and mercury are not reactive enough to displaceremove the hydrogen from dilute acids


Chemical equations


1 Magnesium  Hydrochloric acid  Magnesium chloride  Hydrogen


Mgs  2HCl aq  MgCl2 aq  H2g


2 Zinc  Hydrochloric acid  Zinc chloride  Hydrogen


Zns  2HCl aq  ZnCl2 aq  H2g


3 Iron  Hydrochloric acid  Iron II chloride  Hydrogen


Fes  2HCl aq  FeCl2 aq  H2g


4 Aluminium  Hydrochloric acid  Aluminium chloride  Hydrogen


2Als  3HCl aq  AlCl3 aq  3H2g


5 Magnesium  Sulphuric VI acid  Magnesium sulphate VI  Hydrogen


Mgs  H2SO4 aq  MgSO4 aq  H2g


6 Zinc  Sulphuric VI acid  Zinc sulphate VI  Hydrogen


Zns  H2SO4 aq  ZnSO4 aq  H2 g


7 Iron  Sulphuric VI acid  Iron II sulphate VI  Hydrogen Fes  H2SO4 aq  FeSO4 aq  H2g


8 Aluminium  Sulphuric VI acid  Aluminium sulphate VI  Hydrogen


2Als  3H2SO4 aq  Al2 SO43 aq  3H2 g


2 Reaction of metal carbonates and hydrogen carbonates with mineral acids


All acids react with carbonates and hydrogen carbonates to form salt water and produce evolve carbon IV oxide gas


Metal carbonate  Acid  Salt  Water CarbonIVoxide gas


Metal hydrogen carbonate  Acid  Salt  Water  Carbon IV oxide gas


Experiment reaction of metal carbonates and hydrogen carbonates with mineral acids


aPlace 5cm3 of dilute hydrochloric acid in a small test tube Add half spatula full of sodium carbonate Stopper the test tube using a cork with delivery tube directed into lime water Record the observations made Test the gas also with burning splint


bRepeat the procedure in a above using Zinc carbonate Calcium carbonate copper carbonate sodium hydrogen carbonate Potassium hydrogen carbonate in place of Sodium carbonate


cRepeat the procedure in a then b using dilute sulphuric VI acid in place of dilute hydrochloric acid


Set up of apparatus


Sample observations


i effervescencebubbles producedfizzing in all cases


ii Colourless gas produced in all cases


iii Gas produced forms a white precipitate with lime water


Explanation


All metal carbonatehydrogen carbonate reacts with dilute acids to produce bubbles of carbon IV oxide gas Carbon IV oxide gas is a colourless gas that extinguishes a burning splint When carbon IV oxide gas is bubbled in lime water a white precipitate is formed


Chemical equations


1 Sodium carbonate Hydrochloric acid  Sodium chloride  Carbon IV Oxide Water


Na2CO3s  2HCl aq  2NaCl aq  H2Og  CO2 g


2 Calcium carbonate Hydrochloric acid  Calcium chloride  Carbon IV Oxide Water


CaCO3s  2HCl aq  CaCl2 aq  H2Og  CO2 g


3 Magnesium carbonate Hydrochloric acid Magnesium chloride  Carbon IV Oxide Water


MgCO3s  2HCl aq  MgCl2 aq  H2O g  CO2 g


4 Copper carbonate Hydrochloric acid Copper II chloride  Carbon IV Oxide Water


CuCO3s  2HCl aq  CuCl2 aq  H2O g  CO2 g


5 Copper carbonate Sulphuric VI acid Copper II sulphate VI  Carbon IV Oxide Water


CuCO3s  H2SO4 aq  CuSO4 aq  H2O g  CO2 g


6 Zinc carbonate Sulphuric VI acid Zinc sulphate VI  Carbon IV Oxide Water


ZnCO3s  H2SO4 aq  ZnSO4 aq  H2O g  CO2 g


7 Sodium hydrogen carbonate Sulphuric VI acid Sodium sulphate VI  Carbon IV Oxide Water


NaHCO3s  H2SO4 aq  Na2SO4 aq  H2O g  CO2 g


8 Potassium hydrogen carbonate Sulphuric VI acid Potassium sulphate VI  Carbon IV Oxide Water


KHCO3s  H2SO4 aq  K2SO4 aq  H2O g  CO2 g


9 Potassium hydrogen carbonate Hydrochloric acid Potassium chloride  Carbon IV Oxide Water


KHCO3s  HCl aq  KCl aq  H2O g  CO2 g


10 Sodium hydrogen carbonate Hydrochloric acid Sodium chloride  Carbon IV Oxide Water


NaHCO3s  HCl aq  NaCl aq  H2O g  CO2 g


3 Neutralization by basesalkalis


All acids react with bases to form a salt and water only The reaction of an acid with metal oxideshydroxides bases to salt and water only is called neutralization reaction


Since no effervescencebubblingfizzing take place during neutralization


i The reaction with alkalis requires a suitable indicator The colour of the indicator changes when all the acid has reacted with the soluble solution of the alkali metal oxides hydroxides


ii Excess of the base is added to ensure all the acid reacts The excess acid is then filtered off


Experiment 1 reaction of alkali with mineral acids


iPlace about 5cm3 of dilute hydrochloric acid in a boiling tube Add one drop of phenolphthalein indicator Using a dropperteat pipette add dilute sodium hydroxide dropwise until there is a colour change


iiRepeat the procedure with dilute sulphuric VI acid instead of hydrochloric acid


iiiRepeat the procedure with potassium hydroxide instead of sodium hydroxide


Sample observation


Colour of phenolphthalein change from colourless to pink in all cases


Explanation


Basesalkalis neutralize acids Acids and basesalkalis are colourless A suitable indicator like phenolphthalein change colour to pink when all the acid has been neutralized by the basesalkalis Phenolphthalein change colour from pink to colourless when all the basesalkalis has been neutralized by the acid


Chemical equation


Sodium oxide  Hydrochloric acid  Sodium chloride  Water


Na2Os  HCl  NaClaq  H2Ol


Potassium oxide  Hydrochloric acid  Potassium chloride  Water


K2Os  HCl  KClaq  H2Ol


Sodium hydroxide  Hydrochloric acid  Sodium chloride  Water


NaOHs  HCl  NaClaq  H2Ol


Ammonia solution  Hydrochloric acid  Ammonium chloride  Water


NH4OHs  HCl  NH4Cl aq  H2O l


Potassium hydroxide  Hydrochloric acid  Potassium chloride  Water


KOHs  HCl  KClaq  H2Ol


Sodium hydroxide  sulphuric VIacid  Sodium sulphateVI  Water


2NaOHs  H2SO4  Na2SO4 aq  2H2O l


Potassium hydroxide  sulphuric VI acid  Potassium sulphate VI  Water


2KOHs  H2SO4  K2SO4 aq  2H2O l


Ammonia solution  sulphuric VI acid  Ammonium sulphate VI  Water


2NH4OHs  H2SO4  NH42SO4 aq  2H2O l


Magnesium hydroxide  sulphuric VI acid  Magnesium sulphate VI  Water


Mg OH2s  H2SO4  MgSO4 aq  2H2Ol


Magnesium hydroxide  Hydrochloric acid  Magnesium chloride  Water


Mg OH2s  HClaq  MgCl2 aq  2H2Ol


CHEMISTRY FORM TWO NOTES


INTRODUCTION TO ELECTROLYSIS ELECTROLYTIC CELL


1 Electrolysis is defined simply as the decomposition of a compound by an electric currentelectricity


A compound that is decomposed by an electric current is called an electrolyte Some electrolytes are weak while others are strong


2 Strong electrolytes are those that are fully ionizeddissociated into many ions Common strong electrolytes include


i All mineral acids


ii All strong alkalissodium hydroxidepotassium hydroxide


iii All soluble salts


3 Weak electrolytes are those that are partiallypartly ionizeddissociated into few ions


Common weak electrolytes include


i All organic acids


ii All bases except sodium hydroxidepotassium hydroxide


iiiWater


4 A compound that is not decomposed by an electric current is called nonelectrolyte


Nonelectrolytes are those compounds substances that exist as molecules and thus cannot ionizedissociate into any ions


Common nonelectrolytes include


i Most organic solvents eg petrolparaffinbenzenemethylbenzeneethanol


ii All hydrocarbons alkanes alkenesalkynes


iiiChemicals of life eg proteins carbohydrates lipids starch sugar


5 An electrolytes in solid state have fused joined ions and therefore do not conduct electricity but the ions cations and anions are free and mobile in molten and aqueous solution dissolved in water state


6 During electrolysis the free ions are attracted to the electrodes


An electrode is a rod through which current enter and leave the electrolyte during electrolysis


An electrode that does not influencealter the products of electrolysis is called an inert electrode


Common inert electrodes include


iPlatinum


iiCarbon graphite


Platinum is not usually used in a school laboratory because it is very expensive Carbon graphite is easilyreadily and cheaply available from used dry cells


7 The positive electrode is called Anode The anode is the electrode through which current enter the electrolyteelectrons leave the electrolyte


8 The negative electrode is called Cathode The cathode is the electrode through which current leave the electrolyte  electrons enter the electrolyte


9 During the electrolysis free anions are attracted to the anode where they lose donate electrons to form neutral atomsmolecules ie


Ml  Ml  e for cations from molten electrolytes


Ms  Maq  e for cations from electrolytes in aqueous state  solution  dissolved in water


The neutral atoms molecules form the products of electrolysis at the anode This is called discharge at anode


10 During electrolysis free cations are attracted to the cathode where they gain acceptacquire electrons to form neutral atomsmolecules


X aq  2e  Xs for cations from electrolytes in aqueous state  solution  dissolved in water


2X l  2e  X l for cations from molten electrolytes


The neutral atoms molecules form the products of electrolysis at the cathode This is called discharge at cathode


11 The below set up shows an electrolytic cell


12 For a compound salt containing only two ionbinary salt the products of electrolysis in an electrolytic cell can be determined as in the below examples


a To determine the products of electrolysis of molten Lead II chloride


iDecomposition of electrolyte into free ions


PbCl2 l  Pb 2 l  2Cll


Compound decomposed into free cation and anion in liquid state


iiAt the cathodenegative electrode


Pb 2 l  2e  Pb l


Cation  Pb 2 gains  accepts  acquires electrons to form free atom


iiiAt the anodepositive electrode


2Cll  Cl2 g  2e


Anion  Cl donatelose electrons to form free atom then a gas molecule


ivProducts of electrolysis therefore are


I At the cathode grey beads solid lead metal


IIAt the anode pale green chlorine gas


b To determine the products of electrolysis of molten Zinc bromide


ZnBr2 l  Zn 2 l  2Brl


Zn 2 l  2e  Zn l


Cation  Zn2 gains  accepts  acquires electrons to form free atom


2Brl  Br2 g  2e


Anion  Br donatelose electrons to form free atom then a liquid molecule which change to gas on heating


I At the cathode grey beads solid Zinc metal


IIAt the anode red bromine liquid  redbrown bromine gas


c To determine the products of electrolysis of molten sodium chloride


NaCl l  Na l  Cll


iiAt the cathodenegative electrode


2Na l  2e  Na l


Cation  Na gains  accepts  acquires electrons to form free atom


iiiAt the anodepositive electrode


2Cll  Cl2 g  2e


I At the cathode grey beads solid sodium metal


IIAt the anode pale green chlorine gas


d To determine the products of electrolysis of molten Aluminum III oxide


Al2O3 l  2Al 3 l  3O2l


4Al 3 l  12e  4Al l


Cation  Al 3 gains  accepts  acquires electrons to form free atom


6O2l  3O2 g  12e


Anion 6O2 donatelose 12 electrons to form free atom then three gas molecule


I At the cathode grey beads solid aluminum metal


IIAt the anode colourless gas that relightsrekindles glowing splint


13 In industries electrolysis has the following usesapplications


aExtraction of reactive metals from their ores


Potassium sodium magnesium and aluminum are extracted from their ores using electrolytic methods


bPurifying copper after extraction from copper pyrites ores


Copper obtained from copper pyrites ores is not pure After extraction the copper is refined by electrolyzing copper II sulphate VI solution using the impure copper as anode and a thin strip of pure copper as cathode Electrode ionization takes place there


iAt the cathode Cu2 aq  2e  Cus Pure copper deposits on the strip


iiAt the anode Cus Cu2 aq  2e impure copper erodesdissolves


cElectroplating


The label EPNS Electro Plated Nickel Silver on some steelmetallic utensils mean they are platedcoated with silver andor Nickel to improve their appearance add their aesthetic value and preventslow corrosionrusting of iron Electroplating is the process of coating a metal with another metal using an electric current During electroplating the cathode is made of the metal to be coatedimpure


Example


During the electroplating of a spoon with silver


i The spoonimpure is placed as the cathode negative terminal of battery


ii The pure silver is placed as the anode positive terminal of battery


iii The pure silver erodesionizesdissociates to release electrons


Ags Ag aq  e impure silver erodesdissolves


iv silver Agions from electrolyte gain electrons to form pure silver deposits  coat cover the spoonimpure


Ag aq  e Ags pure silver deposits coatcover on spoon


CHEMICAL BONDING AND STRUCTURE


CHEMICAL BONDING


A chemical bond is formed when atoms of the same or different elements share gain donate or delocalize their outer energy level electrons to combine during chemical reactions in order to be stable


Atoms have equal number of negatively charged electrons in the energy levels and positively charged protons in the nucleus


Atoms are chemically stable if they have filled outer energy level An energy level is full if it has duplet 2 or octet 8 state in outer energy level


Noble gases have duplet octet All other atoms try to be like noble gases through chemical reactions and forming molecules


Only electrons in the outer energy level take part in formation of a chemical bond There are three main types of chemical bonds formed by atoms


i Covalent bond


ii ionicelectrovalent bond


iii Metallic bond


iCOVALENT BOND


A covalent bond is formed when atoms of the same or different element share some or all the outer energy level electrons to combine during chemical reactions in order to attain duplet or octet


A shared pair of electrons is attracted by the nucleus protons of the two atoms sharing


Covalent bonds are mainly formed by nonmetals to form molecules A molecule is a group of atoms of the same or different elements held together by a covalent bond The number of atoms making a molecule is called atomicity Noble gases are monatomic because they are stable and thus do not bond with each other or other atoms Most other gases are diatomic


The more the number of electrons shared the stronger the covalent bond


A pair of electrons that do not take part in the formation of a covalent bond is called a lone pair of electrons


Mathematically the number of electrons to be shared by an atom is equal to the number of electrons remaining for the atom to be stableattain dupletoctet have maximum electrons in outer energy level


The following diagrams illustrate the formation of covalent bonds


a Hydrogen molecule is made up of two hydrogen atoms in the outer energy level each requiring one electron to have a stable duplet


To show the formation of covalent bonding in the molecule then the following datainformation is required


Symbol of atomelement taking part in bonding H H


Number of protonselectrons 1 1


Electron configurationstructure 1 1


Number of electron in outer energy level 1 1


Number of electrons remaining to be stableshared 1 1


Number of electrons not shared lone pairs 0 0


Diagram method 1


Diagram method 2


H H


Note


After bonding the following intermolecular forces exist


i The attraction of the shared electrons by both nucleus protons of the atoms


ii The repulsion of the nucleus of one atom on the other


iii Balance of the attraction and repulsion is maintained insideintermolecularwithin the molecule as follows


P1 P1


ivProtons P1 from nucleus of atom 1 repel protons P2 from nucleus of atom 2


vElectron E1 in the energy levels of atom 1 repel electron E2 in the energy levels of atom 2


vi Protons P1 from nucleus of atom 1 attract electron E2 in the energy levels of atom 2


vii protons P2 from nucleus of atom 2 attract electron E2 in the energy levels of atom 2


b Fluorine chlorine bromine and iodine molecules are made up also of two atoms sharing the outer energy level electrons to have a stable octet


i Fluorine


Symbol of atomelement taking part in bonding F F


Number of protonselectrons 9 9


Electron configurationstructure 27 27


Number of electron in outer energy level 7 7


Number of outer electrons not shared 3lone pairs 6 6


Diagram method 1


Diagram method 2


ii Chlorine


Symbol of atomelement taking part in bonding Cl Cl


Number of protonselectrons 17 17


Electron configurationstructure 287 287


iii Bromine


Symbol of atomelement taking part in bonding Br Br


Number of protonselectrons 35 35


Electron configurationstructure 28187 28187


Number of electrons remaining to be stableshared 1 1


iv Iodine


Symbol of atomelement taking part in bonding I I


Number of protonselectrons 53 53


Electron configurationstructure 2818187


c Oxygen molecule is made up of two atoms sharing each two outer energy level electrons to have a stable octet as shown below


Symbol of atomelement taking part in bonding O O


Number of protonselectrons 8 8


Electron configurationstructure 26 26


Number of electron in outer energy level 6 6


Number of electrons remaining to be stableshared 2 2


Number of outer electrons not shared 2lone pairs 4 4


Diagram method 1


Diagram method 2


d Nitrogen and phosphorus molecule is made up of two atoms sharing each three outer energy level electrons to have a stable octet as shown below


i Nitrogen


Symbol of atomelement taking part in bonding N N


Number of protonselectrons 7 7


Electron configurationstructure 25 25


Number of electron in outer energy level 5 5


Number of electrons remaining to be stableshared 3 3


Number of outer electrons not shared 3lone pairs 2 2


ii Phosphorus


Symbol of atomelement taking part in bonding P P


Number of protonselectrons 15 15


Electron configurationstructure 285 285


e Water molecule is made up of hydrogen and oxygen Hydrogen requires sharing one electron with oxygen to be stableattain duplet Oxygen requires to share two electrons to be stableattain octet Two hydrogen atoms share with one oxygen atom for both to be stable as shown below


Symbol of atomselements taking part in bonding O H


Number of protonselectrons 8 1


Electron configurationstructure 26 1


Number of electron in outer energy level 6 1


Number of electrons remaining to be stableshared 2 1


Number of electrons not shared 2Oxygen lone pairs 4 0


Diagram method 1


Diagram method 2


f Ammonia molecule is made up of Hydrogen and Nitrogen Hydrogen requires sharing one electron with Nitrogen to be stableattain duplet Nitrogen requires to share three electrons to be stableattain octet Three hydrogen atoms share with one nitrogen atom for both to be stable as shown below


Symbol of atomselements taking part in bonding N H


Number of protonselectrons 7 1


Electron configurationstructure 25 1


Number of electron in outer energy level 5 1


Number of electrons remaining to be stableshared 3 1


Number of electrons not shared 1Nitrogen lone pairs 2 0


g Carbon IV oxide molecule is made up of carbon and oxygen Carbon requires to share four electrons with oxygen to be stableattain octet Oxygen requires to share two electrons to be stableattain octet Two oxygen atoms share with one carbon atom for both to be stable as shown below


Symbol of atomselements taking part in bonding O C


Number of protonselectrons 8 6


Electron configurationstructure 26 24


Number of electron in outer energy level 6 4


Number of electrons remaining to be stableshared 2 4


2lone pairs from each Oxygen atom 2 0


h Methane molecule is made up of hydrogen and carbon Hydrogen requires sharing one electron with carbon to be stableattain duplet Carbon requires sharing four electrons to be stableattain octet Four hydrogen atoms share with one carbon atom for both to be stable as shown below


Symbol of atomselements taking part in bonding C H


Number of protonselectrons 6 1


Electron configurationstructure 24 1


Number of electron in outer energy level 4 1


Number of electrons remaining to be stableshared 4 1


Number of electrons not shared  No lone pairs 0 0


Diagram method 1


Diagram method 2


i Tetrachloromethane molecule is made up of chlorine and carbon Chlorine requires sharing one electron with carbon to be stableattain octet Carbon requires sharing four electrons to be stableattain octet Four chlorine atoms share with one carbon atom for both to be stable as shown below


Symbol of atomselements taking part in bonding C Cl


Number of protonselectrons 6 17


Electron configurationstructure 24 287


Number of electron in outer energy level 4 7


3lone pairs from each Chlorine atom 24 electrons 0 6


j Ethane molecule is made up of six hydrogen and two carbon atoms Hydrogen requires to share one electron with carbon to be stableattain duplet Carbon requires to share four electrons to be stableattain octet Three hydrogen atoms share with one carbon atom while another three hydrogen atoms share with a different carbon atom The two carbon atoms bond by sharing a pair of the remaining electrons as shown below


Number of electrons not shared No lone pairs 0 0


k Ethene molecule is made up of four hydrogen and two carbon atoms Hydrogen requires to share one electron with carbon to be stableattain duplet Carbon requires to share four electrons to be stableattain octet Two hydrogen atoms share with one carbon atom while another two hydrogen atoms share with a different carbon atom The two carbon atoms bond by sharing two pairs of the remaining electrons as shown below


l Ethyne molecule is made up of two hydrogen and two carbon atoms Hydrogen requires to share one electron with carbon to be stableattain duplet Carbon requires to share four electrons to be stableattain octet One hydrogen atoms share with one carbon atom while another hydrogen atoms share with a different carbon atom The two carbon atoms bond by sharing three pairs of the remaining electrons as shown below


Number of protonselectrons 6 1


Electron configurationstructure 24 1


j Ethanol molecule is made up of six hydrogen one Oxygen atom two carbon atoms


Five Hydrogen atoms share their one electron each with carbon to be stableattain duplet One Hydrogen atoms share one electron with Oxygen for both to attain dupletoctet


Each Carbon uses four electrons to share with Oand Hattains octetduplet


NB Oxygen has two lone pairs


j Ethanoic molecule is made up of four hydrogen two Oxygen atom two carbon atoms


Three Hydrogen atoms share their one electron each with carbon to be stableattain duplet One Hydrogen atoms share one electron with Oxygen for both to attain dupletoctet


NB Each Oxygen atom has two lone pairs


By convention as a rule a


i Single covalent bond made up of two shared a pair electrons is represented by a dash


ii Double covalent bond made up of four shared two pairs electrons is represented by a double dash


iii Triple covalent bond made up of six shared three pairs electrons is represented by a triple dash


The representation below show the molecules covered in a to k above


Hydrogen moleculeH2 HH


Fluorine moleculeF2 FF


Chlorine moleculeCl2 ClCl


Bromine moleculeBr2 BrBr


Iodine moleculeI2 II


Oxygen moleculeO2 OO


Nitrogen moleculeN2 NN


Phosphorus moleculeP2 PP


Water molecule H2O HOH


j Ammonia moleculeNH3 HNH


H


k Carbon IV oxide molecule CO2 OCO


l Methane molecule CH4 HCH


Cl


m Tetrachloromethane molecule CCl4 ClCCl


H H


n Ethane molecule C2H6 HCCH


p Ethene molecule C2H4 HCCH


q Ethyne molecule C2H6 HCCH


Dative coordinate bond


A dativecoordinate bond is a covalent bond formed when a lone pair of electrons is donated then shared to an electrondeficient speciesionatom


During dativecoordinate bonding all the shared pair of electrons are donated by one of the combiningbonding species ionatom


Like covalent bonding coordinate dative bond is mainly formed by nonmetals


Illustration of coordinate dative bond


a Ammonium ion NH4


The ammonium ion is made up of ammonia NH3 molecule and hydrogen H ion H ion has no electrons NH3 is made up of covalent bonding from Nitrogen and Hydrogen One lone pair of electrons is present in Nitrogen atom after the bonding This lone pair is donated and shared with an electrondeficient H ion


Diagram method 1


Diagram method 2


b Phosphine ion PH4


The Phosphine ion is made up of Phosphine NH3 molecule and hydrogen H ion H ion has no electrons PH3 is made up of covalent bonding from Phosphorus and Hydrogen One lone pair of electrons is present in Phosphorus atom After the bonding this lone pair is donated and shared with the electrondeficient H ion


c Hydroxonium H3O ion


The hydroxonium ion is made up of water H2O molecule and hydrogen H ion H ion has no electrons The H2O molecule is made up of covalent bonding from Oxygen and Hydrogen One lone pair of electrons out of the two present in Oxygen atom after the bonding is donated and shared with the electrondeficient H ion


d Carbon II oxide CO


Carbon II oxide is made up of carbon and Oxygen atoms sharing each two outer electron and not sharing each two electrons Oxygen with an extra lone pair of electrons donates and share with the carbon atom for both to be stable


e Aluminum III chloride AlCl3Al2Cl6


Aluminum III chloride is made up of aluminum and chlorine One aluminum atom shares its outer electrons with three separate chlorine atoms All chlorine atoms attain stable octet but aluminum does not Another molecule of aluminum chloride shares its chlorine lone pair of electrons with the aluminum atom for both to be stable This type of bond exists only in vapour phase after aluminum chloride sublimes


A dativecoordinate bond is by convention represented by an arrow  heading from the donor of the shared pair of electrons


Below is the representation of molecules in the above examples


a Ammonium ion


H


H NH


b Phosphine ion H


H PH


c Hydroxonium ion


H OH


d Carbon II oxide OC


d Aluminum III chloride Cl Cl Cl


Al Al


Cl Cl Cl


iiIONICELECTROVALENT BOND


An ionicelectrovalent bond is extreme of a covalent bond


During ionicelectrovalent bonding there is complete transfer of valence electrons to one electronegative atom from an electropositive atom


All metals are electropositive and easilyreadily donatelose their valence electrons


All nonmetals are electronegative and easilyreadily gainacquire extra electrons


Ionicelectrovalent bonding therefore mainly involves transfer of electrons from metalmetallic radical to nonmetallic radical


When an electropositive atom donates loses the valence electrons it forms a positively charged cation to attain stable octetduplet


When an electronegative atom gains acquires extra valence electrons it forms a negatively charged anion to attain stable octetduplet


The electrostatic attraction force between the stable positively charged cation and the stable negatively charged anion with opposite charges constitute the ionic bond


Like in covalentdativecoordinate bonding only the outer energy level electrons take part in the formation of ionicelectrovalent bond


Like in covalentdativecoordinate bonding the more electrons taking part  involved in the formation of ionicelectrovalent bond the stronger the ionic electrovalent bond Illustration of ionic electrovalent bond


a Sodium chloride NaCl


Sodium chloride NaCl is formed when a sodium atom donates its outer valence electrons to chlorine atom for both to attain stable octet


Symbol of atomselements taking part in bonding Na Cl


Number of protonselectrons 11 17


Electron configurationstructure 281 287


Number of electron in outer energy level 11 7


Number of electrons donated and gained to be stable 1 1


New electron configurationstructure 28 28


Symbol of cationanion after bonding Na Cl


Diagram


b Magnesium chloride MgCl2


Magnesium chloride MgCl2 is formed when a magnesium atom donate its two outer valence electrons to chlorine atoms Two chlorine atoms are required to gain each one electron All the ions cations and anions attain stable octet


Symbol of atomselements taking part in bonding Mg Cl


Electron configurationstructure 282 287


Number of electron in outer energy level 2 7


Number of electrons donated and gained to be stable 2 1


Symbol of cationanion after bonding Mg2 Cl


c Lithium oxide Li2O


Lithium oxide Li2O is formed when a Lithium atom donate its outer valence electrons to Oxygen atom Two Lithium atoms are required to donatelose each one electron and attain stable duplet Oxygen atom acquires the two electrons and attain stable octet


Symbol of atomselements taking part in bonding Li O


Number of protonselectrons 3 8


Electron configurationstructure 21 26


Number of electron in outer energy level 1 6


Number of electrons donated and gained to be stable 1 2


New electron configurationstructure 2 28


Symbol of cationanion after bonding Li O2


Diagram


d Aluminum III oxide Al2O3


Aluminum III oxide Al2O3 is formed when a Aluminum atom donate its three outer valence electrons to Oxygen atom Two Aluminum atoms are required to donatelose each three electron and attain stable octet Three Oxygen atoms gain acquire the six electrons and attain stable octet


Symbol of atomselements taking part in bonding Al O


Number of protonselectrons 13 8


Electron configurationstructure 283 26


Number of electron in outer energy level 3 6


Number of electrons donated and gained to be stable 3 2


New electron configurationstructure 28 28


Symbol of cationanion after bonding Al3 O2


e Calcium oxide CaO


Calcium oxide CaO is formed when a Calcium atom donate its two outer valence electrons to Oxygen atom Both attain stable octet


Symbol of atomselements taking part in bonding Ca O


Number of protonselectrons 20 8


Electron configurationstructure 2882 26


Number of electron in outer energy level 2 6


Number of electrons donated and gained to be stable 2 2


New electron configurationstructure 288 28


Symbol of cationanion after bonding Ca2 O2


Diagram


Some compounds can be formed from ionicelectrovalent covalent and dativecoordinate bonding within their atomsmolecules


a Formation of ammonium chloride


Ammonium chloride is formed from the reaction of ammonia gas and hydrogen chloride gas Both ammonia and hydrogen chloride gas are formed from covalent bonding During the reaction of ammonia and hydrogen chloride gas to form Ammonium chloride


ammonia forms a dativecoordinate bond with electron deficient H ion from Hydrogen chloride to form ammonium ionNH4ion


the chloride ion Cl and ammonium ionNH4ion bond through ionic  electrovalent bond from the electrostatic attraction between the oppositeunlike charges


b Dissolutiondissolving of hydrogen chloride


Hydrogen chloride is formed when hydrogen and chlorine atoms form a covalent bond Water is formed when hydrogen and Oxygen atoms also form a covalent bond When hydrogen chloride gas is dissolved in water


water molecules forms a dativecoordinate bond with electron deficient H ion from Hydrogen chloride to form hydroxonium ionH3Oion


the chloride ion Cl and hydroxonium ionH3Oion bond through ionic  electrovalent bond from the electrostatic attraction between the oppositeunlike charges


c Dissolutiondissolving of ammonia gas


Ammonia gas is formed when hydrogen and Nitrogen atoms form a covalent bond Water is formed when hydrogen and Oxygen atoms also form a covalent bond When Ammonia gas is dissolved in water


ammonia forms a dativecoordinate bond with electron deficient H ion from a water molecule to form ammonium ionNH4ion


the hydroxide ion OH and ammonium ionNH4ion bond through ionic  electrovalent bond from the electrostatic attraction between the oppositeunlike charges


iiiMETALLIC BOND


A metallic bond is formed when metallic atoms delocalize their outer electrons in order to be stable


Metals delocalize their outer electrons to form positively charged cation


The electrostatic attraction force between the metallic cation and the negatively charged electrons constitute the metallic bond


The more delocalized electrons the stronger the metallic bond


Illustration of ionic electrovalent bond


a Sodium Na is made of one valence electron The electron is donated to form Na ion The electron is delocalized free within many sodium ions


Symbol of atomselements taking part in bonding Na Na Na


Number of protonselectrons 11 11 11


Electron configurationstructure 281 281 281


Number of electron in outer energy level 1 1 1


Number of electrons delocalizedfree within 1 1 1


New electron configurationstructure 28 28 28


Symbol of cation after metallic bonding Na Na Na


Diagram


threeMetallic cations attract


three freedelocalized electrons


b Aluminium Al is made of three valence electron The three electrons are donated to form Al3 ion The electrons are delocalized free within many aluminium ions


Symbol of atomselements taking part in bonding Al Al Al


Number of protonselectrons 13 13 13


Electron configurationstructure 283 283 283


Number of electron in outer energy level 3 3 3


Number of electrons delocalizedfree within 3 3 3


Symbol of cation after metallic bonding Al3 Al3 Al3


nine freedelocalized electrons


cCalcium Ca is made of two valence electronThe two electrons are donated to form Ca2 ionThe electrons are delocalized free within many Calcium ions


Symbol of atomselements taking part in bonding Ca Ca Ca


Number of protonselectrons 20 20 20


Electron configurationstructure 2882 2882 2882


Number of electron in outer energy level 2 2 2


Number of electrons delocalizedfree within 2 2 2


New electron configurationstructure 288 288 288


Symbol of cation after metallic bonding Ca2 Ca2 Ca2


Diagram


threeMetallic cations attract


six freedelocalized electrons


d Magnesium Mg is made of two valence electron The two electrons are donated to form Mg2ionThe electrons are delocalized free within many Magnesium ions


Symbol of atomselements taking part in bonding Mg Mg


Number of protonselectrons 12 12


Electron configurationstructure 282 282


Number of electron in outer energy level 2 2


Number of electrons delocalizedfree within 2 2


New electron configurationstructure 28 28


Symbol of cation after metallic bonding Mg2 Mg2


twoMetallic cations attract


four freedelocalized electrons


eLithium Li is made of one valence electron The electron is donated to form Li ion The electron is delocalized free within many Lithium ionsie


Symbol of atomselements taking part in bonding Li Li Li Li


Number of protonselectrons 3 3 3 3


Electron configurationstructure 21 21 21 21


Number of electron in outer energy level 1 1 1 1


Number of electrons delocalizedfree within 1 1 1 1


New electron configurationstructure 21 21 21 21


Symbol of cation after metallic bonding Li Li Li Li


Diagram


fourMetallic cations attract


BCHEMICAL STRUCTURE


Chemical structure is the patternarrangement of atoms after they have bonded There are two main types of chemical structures


i Simple molecular structure


ii Giant structures


iSimple molecular structure


Simple molecular structure is the pattern formed after atoms of nonmetals have covalently bonded to form simple molecules


Molecules are made of atoms joined together by weak intermolecular forces called Vanderwaals forcesThe Vanderwaals forces hold the molecules together while the covalent bonds hold the atoms in the molecule


Illustration of simple molecular structure


aHydrogen moleculeH2


Hydrogen gas is made up of strong covalent bondsintramolecular forces between each hydrogen atom making the molecule Each molecule is joined to another by weak Vanderwaals forces intermolecular forces


bOxygen moleculeO2


Oxygen gas is made up of strong covalent bondsintramolecular forces between each Oxygen atom making the molecule Each molecule is joined to another by weak Vanderwaals forces intermolecular forces


Strong intramolecular forcescovalent bond


OO OO OO OO


weak intermolecular


OO OO OO OO forcesvanderwaals forces


c Iodine moleculeI2


Iodine solid crystals are made up of strong covalent bondsintramolecular forces between each iodine atom making the molecule Each molecule is joined to another by weak Vanderwaals forces intermolecular forces


I I I  I I  I I  I


I  I I  I I  I I  I forcesvanderwaals forces


dCarbonIV oxide moleculeCO2


Carbon IV oxide gas molecule is made up of strong covalent bondsintramolecular forces between each Carbon and oxygen atoms making the molecule Each molecule is joined to another by weak Vanderwaals forces intermolecular forces


OCO OCO OCO


OCO OCO OCO forcesvanderwaals forces


The following are the main characteristic properties of simple molecular structured compounds


aState


Most simple molecular substances are gases liquid or liquids or solid that sublimes or has low boilingmelting points at room temperature 25oC and pressure atmospheric pressure


Examples of simple molecular substances include


all gases eg Hydrogen oxygen nitrogen carbon IV oxide


Petroleum fractions eg Petrol paraffin diesel wax


Solid nonmetals eg Sulphur Iodine


Water


b Low meltingboiling points


Melting is the process of weakening the intermolecular vanderwaal forces of attraction between the molecules that holding the substancecompound


Note


iMelting and boiling does not involve weakeningbreaking the strong intramolecular forcecovalent bonds holding the atoms in the molecule


ii Melting and boiling points increase with increase in atomic radiussize of the atoms making the molecule as the intermolecular forces  vanderwaal forces of attraction between the molecules increase eg


Iodine has a higher meltingboiling point than chlorine because it has a higher bigger atomic radiussize than chlorine making the molecule to have stronger intermolecular force vanderwaal forces of attraction between the molecules than chlorine Iodine is hence a solid and chlorine is a gas


cInsoluble in watersoluble in organic solvents


Polar substances dissolve in polar solvents Water is a polar solvent Molecular substances do not thus dissolve in water because they are nonpolar They dissolve in nonpolar solvents like methylbenzene benzene tetrachloromethane or propanone


d Poor conductors of heat and electricity


Substances with free mobile ions or free mobiledelocalized electrons conduct electricity Molecular substances are poor conductors of heatelectricity because their molecules have no free mobile ionselectrons This makes them very good insulators Hydrogen bonds


A hydrogen bond is an intermolecular force of attraction in which a very electronegative atom attracts hydrogen atom of another molecule


The most electronegative elements are Fluorine Oxygen and Nitrogen Molecular compounds made up of these elements usually have hydrogen bonds


Hydrogen bonds are stronger than vanderwaals forces but weaker than covalent bonds Molecular compounds with hydrogen bonds thus have higher meltingboiling points than those with vanderwaals forces


Illustration of Hydrogen bonding


aWater molecule


During formation of covalent bond the oxygen atom attractpull the shared electrons more to itself than Hydrogen creating partial negative chargesin Oxygen and partial positive chargesin Hydrogen


Two molecules attract each other at the partial charges through Hydrogen bonding


The hydrogen bonding in water makes it


ia liquid with higher boiling and melting point than simple molecular substances with higher molecular mass eg Hydrogen sulphide as in the table below


Influence of Hbond in water H2O in comparison to H2S


iihave higher volume in solid ice than liquid water and thus ice is less dense than water Ice therefore floats above liquid water


bEthanol molecule


Like in water the oxygen atom attractspulls the shared electrons in the covalent bond more to itself than Hydrogen


This creates a partial negative charge  on oxygen and partial positive charge on hydrogen


Two ethanol molecules attract each other at the partial charges through Hydrogen bonding forming a dimmer


A dimmer is a molecule formed when two molecules join together as below


Hydrogen bonds covalent bonds


R1 O H  O


H  R2


R1 and R2 are extensions of the molecule


For ethanol it is made up of CH3CH2  to make the structure


CH3CH2 O H  O


H  CH2CH3


bEthanoic acid molecule


Like in water and ethanol above the oxygen atom attractspulls the shared electrons in the covalent bond in ethanoic acid more to itself than Hydrogen


This creates a partial negative charge on oxygen and partial positive charge on hydrogen


Two ethanoic acid molecules attract each other at the partial charges through Hydrogenbonding forming a dimer


R1 C O H  O


O  H O  C R2


R1 and 2 are extensions of the molecule


For ethanoic acid the extension is made up of CH3  to make the structure


CH3 C O H  O


O  H O  C CH3


Ethanoic acid like ethanol exists as a dimer


Ethanoic acid has a higher meltingboiling point than ethanol This is because ethanoic acid has twomore hydrogen bond than ethanol


d Proteins and sugars in living things also have multiplecomplex hydrogen bonds in their structures


ii Giant structure


This is the pattern formed after substances atoms ions bond to form a long chain network


Giant structures therefore extend in all directions to form a pattern that continues repeating itself


There are three main giant structures


a giant covalentatomic structure bgiant ionic structure


cgiant metallic structure


a giant covalentatomic structure


Giant covalentatomic structure is the pattern formed after atoms have covalently bonded to form long chain pattern consisting of indefinite number of atoms covalently bonded together


The strong covalent bonds hold all the atoms together to form a very well packed structure Examples of substances with giant covalentatomic structure include


i carbondiamond


ii carbongraphite


iiisilicon


iv siliconIV oxidesand


Carbongraphite and carbondiamond are allotropes of carbon


Allotropy is the existence of an element in more than one stable physical form at the same temperature and pressure


Allotropes are atoms of the same element existing in more than one stable physical form at the same temperature and pressure


Other elements that exhibitshow allotropy include


Sulphur as monoclinic sulphur and rhombic sulphur


Phosphorus as white phosphorus and red phosphorus


The structure of carbondiamond


Carbon has four valence electrons The four valence electrons are used to form covalent bonds During the formation of diamond one carbon atom covalently bond with four other carbon atoms


C C


x x


x C x  C x C x C  C C C


After the bonding the atoms rearrange to form a regular tetrahedral in which one carbon is in the centre while four are at the apexcorners


C


This pattern repeats itself to form a long chain number of atoms covalently bonded together indefinitely The pattern is therefore called giant tetrahedral structure It extends in all directions where one atom of carbon is always a centre of four others at the apexcorner of a regular tetrahedral


The giant tetrahedral structure of carbondiamond is very wellclosely packed and joinedbonded together by strong covalent bond


This makes carbondiamond to have the following properties


a High meltingboiling point


The giant tetrahedral structure is very well packed and joined together by strong covalent bonds


This requires a lot of energyheat to weaken for the element to melt and break for the element to boil


b High density


Carbon diamond is the hardest known natural substance


This is because the giant tetrahedral structure is a very well packed patternstructure and joined together by strong covalent bonds


This makes Carbon diamond be used to make drill for drilling boreholesoil wells


The giant tetrahedral structure of carbon diamond is a very closely packed pattern structure such that heat transfer by conduction is possible This makes carbon diamond a good thermal conductor


c Poor conductor of electricity


Carbondiamond has no freedelocalized electrons within its structure and thus do not conduct electricity


d Insoluble in water


Carbondiamond is insoluble in water because it is nonpolar and do not bond with water molecules


e Is abrasiveRough


The edges of the closely well packed patternstructure of Carbondiamond make its surface roughabrasive and thus able to smoothen cut metals and glass


f Have characteristic luster


Carbondiamond has a high optical dispersion and thus able to disperse light to different colours This makes Carbondiamond one of the most popular gemstone for making jewellery


The structure of carbongraphite


During the formation of graphite one carbon atom covalently bond with three other carbon atoms leaving one freedelocalized electron


C C


x x


x C x  C x C x  C C x freedelocalized electron


After the bonding the atoms rearrange and join together to form a regular hexagon in which six carbon atoms are at the apexcorners


The regular hexagon is joined to another in layers on the same surface by vanderwaals forces


Each layer extends to form a plane in all directions


The fourth valence electron that does not form covalent bonding is freemobile delocalized within the layers


This structurepattern is called giant hexagonal planar structure


The giant hexagonal planar structure of carbongraphite is closely packed and joinedbonded together by strong covalent bonds This makes carbongraphite to have the following properties


a High meltingboiling point


The giant hexagonal planar structure of carbongraphite is well packed and joined together by strong covalent bonds


This requires a lot of energyheat to weaken for the element to melt and break for the element to boil


b Good conductor of electricity


Carbongraphite has freedelocalized 4th valence electrons within its structure and thus conducts electricity


c Insoluble in water


Carbongraphite is insoluble in water because it is nonpolar and do not bond with water molecules


d Soft


Layers of giant hexagonal planar structure of carbon graphite are held together by vanderwaals forces


The vanderwaals forces easily break when pressed and reform back on releasingreducing pressureforce thus making graphite soft


e Smooth and slippery


When pressed at an angle the vanderwaals forces easily break and slide over each other making graphite soft and slippery


It is thus used as a dry lubricant instead of oil


fSome uses of carbongraphite


1 As a dry lubricant carbon graphite is smooth and slippery and thus better lubricant than oilOil heat up when reducing friction


2 Making Leadpencils When pressed at an angle on paper the vanderwaals forces easily break and slide smoothly over contrasting background producing its characteristic black background


3 As moderator in nuclear reactors to reduce the rate of decaydisintegration of radioactive nuclidesatomsisotopes


4 As electrode in drywet cellsbattery carbon graphite is inert and good conductor of electricity Current is thus able to move from one electrodeterminal to the other in dry and wet cellsbatteries Carbon graphite is also very cheap


b giant ionic structure


Giant ionic structure is the pattern formed after ions have bonded through ionicelectrovalent bonding to form a long chain consisting of indefinite number of ions


The strong ionicelectrovalent bond holds all the cations and anions together to form a very well packed structure


Substances with giant ionic structure are mainly crystals of salts eg sodium chloride Magnesium chloride Sodium iodide Potassium chloride copper II sulphateVI


The structure of sodium chloride


Sodium chloride is made up of sodium Na and chloride Clions


Sodium Na ion is formed when a sodium atom donate loosedonate an electron Chloride Cl ion is formed when a chlorine atom gain acquire an extra electron from sodium atom


Many Na and Cl ions then rearrange such that one Na ion is surrounded by six Cl ions and one Cl ion is surrounded by six Na ions


The pattern formed is a giant cubic structure where Cl ion is sand witched between Na ions and the same to Na ions


This pattern forms a crystal


A crystal is a solid form of a substance in which particles are arranged in a definite pattern regularly repeated in three dimensions


The giant cubic structurecrystal of sodium chloride is as below


The giant cubic structurecrystal of sodium chloride is very well packed and joined by strong ionicelectrovalent bonds This makes sodium chloride and many ionic compounds to have the following properties


a Have high melting boiling points


The giant cubic lattice structure of sodium chloride is very closely packed into a crystal that requires a lot of energyheat to weaken and meltboil This applies to all crystalline ionic compounds


b Are good conductors of electricity in molten and aqueous state but poor conductor of electricity in solid


Ionic compounds have fused ions in solid crystalline state


On heating and dissolving in water the crystal is broken into free mobile ions Na and Cl ions


The free mobile ions are responsible for conducting electricity in ionic compounds in molten and aqueous states


cSoluble in water


Ionic compounds are polar and dissolve in polar water molecules


On dissolving the crystal breaks to free the fused ions which are then surrounded by water molecules


b giant metallic structure


This is the pattern formed after metallic atoms have bonded through metallic bond


The pattern formed is one where the metallic cations rearrange to form a cubic structure


The cubic structure is bound together by the free delocalized electrons that move freely within


The more delocalized electrons the stronger the metallic bond


The structure of sodium and aluminium


Sodium has one valence electrons


Aluminium has three valence electrons


After delocalizing the valence electrons the metal cations Na and Al3 rearrange to the apex corners of a regular cube that extend in all directions


The delocalized electrons remain free and mobile as shown below


The giant cubic structure makes metals to have the following properties


a Have high meltingboiling point


The giant cubic structure is very well packed and joinedbonded together by the free delocalized electrons


The more delocalized electrons the higher the meltingboiling point


The largerbigger the metallic cation the weaker the packing of the cations and thus the lower the meltingboiling point eg


i Sodium and potassium have both one valence delocalized electron


Atomic radius of potassium is largerbigger than that of sodium and hence less well packed in its metallic structure


Sodium has therefore a higher meltingboiling point than potassium


ii Sodium has one delocalized electron


Aluminium has three delocalized electrons


Atomic radius of sodium is largerbigger than that of aluminium and hence less well packed in its metallic structure


Aluminium has therefore a higher meltingboiling point than sodium because of the smaller well packed metallic Al3 ions and bondedjoined by morethree delocalized electrons


The table below shows the comparative meltingboiling points of some metals


b Good electrical and thermal conductorelectricity


All metals are good conductors of heat and electricity including Mercury which is a liquid


The mobile delocalized electrons are free within the giant metallic structure to move from one end to the other transmitting heatelectric current


The more delocalized electrons the better the thermalelectrical conductivity


High temperaturesheating lowers the thermalelectrical conductivity of metals because the delocalized electrons vibrate and move randomly hindering transfer of heat


From the table above


Compare the electrical conductivity of


iMagnesium and sodium


Magnesium is a better conductor than sodium


Magnesium has moretwo delocalized electrons than sodium The more delocalized electrons the better the electrical conductor


iiPotassium and sodium


Potassium is a better conductor than sodium


Potassium has biggerlarger atomic radius than sodium The delocalized electrons are less attracted to the nucleus of the atom and thus more free mobile and thus better the electrical conductor


c Insoluble in water


All metals are insoluble in water because they are non polar and thus do not bond with water


Metals higher in the reactivityelectrochemical series like Potassium sodium Lithium and Calcium reacts with cold water producing hydrogen gas and forming an alkaline solution of their hydroxidesie


2Ks  2H2Ol  2KOHaq  H2g


2Nas  2H2Ol  2NaOHaq  H2g


2Lis  2H2Ol  2LiOHaq  H2g


Cas  2H2Ol  CaOH2aq H2g


Heavy metal like Magnesium Aluminium Iron Zinc and Lead react with steamwater vapour to produce hydrogen gas and form the corresponding oxide


Mgs  H2Og  MgOs  H2g


Fes  H2Og  FeOs  H2g


Zns  H2Og  ZnOs  H2g


Pbs  H2Og  PbOs  H2g


2Als  3H2Og  Al2O3s  3H2g


Metals lower in the reactivityelectrochemical series than hydrogen like copper Mercury Gold Silver and Platinum do not react with watervapour


d Shiny metalliclustre


All metals have a shiny grey metallic luster except copper which is brown


When exposed to sunlight the delocalized electrons gain energy they vibrate on the metal surface scattering light to appear shiny


With time most metals corrode and are covered by a layer of the metal oxide


The delocalized electrons are unable to gain and scatter light and the metal surface tarnishesbecome dull


e Ductile and malleable


All metals are malleable can be made into thin sheet and ductile can be made into wire


When beatenhitpressed lengthwise the metallic cations extend and is bound bonded by the freemobile electrons to form a sheet


When beatenhitpressed lengthwise and bredthwise the metallic cations extend and is bound bonded by the freemobile electrons to form a wirethin strip


f Have high tensile strength


Metals are not brittle The free delocalized electrons bind the metal together when it is bent coiled at any angle


The metal thus withstand stresscoiling


g Form alloys


An alloy is a uniform mixture of two or more metals


Some metals have spaces between their metallic cations which can be occupied by another metal cation with smaller atomic radius


Common alloys include


BrassZinc and Copper alloy


BronzeCopper and Tin alloy


German silver


Summary of Bonding and structure


C PERIODICITY OF BONDING AND STRUCTURE


The periodic table does not classify elements as metals and nonmetals The table arranges them in terms of atomic numbers


However based on structure and bonding of the elements in the periodic table


ithe top right hand corner of about twenty elements is nonmetals


iileft of each nonmetal is an element which shows characteristics of both metal and nonmetal


These elements are called semimetalsmetalloids They include Boron silicon Germanium Arsenic and Terullium


iiiall other elements in the periodic table are metal


ivHydrogen is a nonmetal with metallic characteristicproperty of donatinglosing outer electron to form cationH ion


v bromine is the only known natural liquid nonmetal element at room temperature and pressure


vi mercury is only known natural liquid metal element at room temperature and pressure


vii Carbongraphite is a semi metalsmetalloids Carbondiamond is a pure nonmetal yet both are allotropes of carbon same element


a Sketch of the periodic table showing metals metalloid and nonmetals


Metals Metalloids Nonmetals


bPeriodicity in the physical properties of elements across period 2 and 3


Study table I and II below


Table Iperiod 2


Table II period 3


From table I and II above


1 Explain the trend in atomic radius along across a period in the periodic table


Observation


Atomic radius of elements in the same period decrease successively acrossalong a period from left to right


Explanation


Acrossalong the period from left to right there is an increase in nuclear charge from additional number of protons and still additional number of electrons entering the same energy level


Increase in nuclear charge increases the effective nuclear attraction on the outer energy level pulling it closer to the nucleus successively across the period eg


iFrom the table 1and 2 above atomic radius of Sodium 0157nM is higher than that of Magnesium0137nM This is because Magnesium has more effective nuclear attraction on the outer energy level than Sodium hence pulls outer energy level more nearer to its nucleus than sodium


iiThe rate of decrease in the atomic radius become smaller as the atom become heavier eg Atomic radius of Magnesium from sodium falls by0157nM 0137nM 002


Atomic radius of Chlorine from sulphur falls by0104nM 0099nM 0005


This is because gainingadding one more proton to 11 already present cause greater proportional change in nuclear attraction power to magnesium than gainingadding one more proton to 16 already present in sulphur to chlorine


iiiPeriod 3 elements have more energy levels than Period 2 elements They have therefore biggerlarger atomic radiussize than corresponding period 2 elements in the same group


2 Explain the trend in ionic radius alongacross a period in the periodic table


Ionic radius of elements in the same period decrease successively acrossalong a period from left to right for the first three elements then increase drastically then slowly successively decrease


Acrossalong the period from left to right elements change form electron donorslosers reducing agents to electron acceptors oxidizing agents


iAn atom form stable ion by either gainingacquiring accepting extra electron or donatinglosing outer electrons


iiMetals form stable ions by donatinglosing all the outer energy level electrons and thus also the outer energy level ie


Sodium ion has one less energy level than sodium atom The ion is formed by sodium atom donatinglosing all the outer energy level electron and thus also the outer energy level making the ion to have smaller ionic radius than atom


iiiIonic radius therefore decrease acrossalong the period from Lithium to Boron in period 2 and from Sodium to Aluminium in period 3This is because the number of electrons donatedlost causes increased effective nuclear attraction on remaining electrons energy levels


iv Nonmetals form stable ion by gainingacquiringaccepting extra electron in the outer energy level The extra electrons increases the repulsion among electrons and reduces the effective nuclear attraction on outer energy level The outer energy level therefore expandenlargeincrease in order to accommodate the extra repelled electrons The more electrons gainedacceptedacquired the more repulsion and the more expansion to accommodate them and hence biggerlarger atomic radius eg


Nitrogen ion has three electrons more than Nitrogen atom The outer energy level expandenlargeincrease to accommodate the extra repelled electrons Nitrogen atom thus has smaller atomic radius than the ionic radius of nitrogen ion


v Ionic radius decrease from group IV onwards from left to right This because the number of electrons gained to form ion decrease acrossalong the period from left to right eg Nitrogen ion has biggerlarger ionic radius than Oxygen


3 Explain the trend in melting and boiling point of elements in a period in the periodic table


Observation


The melting and boiling point of elements rise up to the elements in Group IVCarbonSilicon alongacross the period then continuously falls


Explanation


Meltingboiling points depend on the packing of the structure making the element and the strength of the bond holding the atomsmolecules together


Acrossalong the period 2 and 3 the structure changes from giant metallic giant atomiccovalent to simple molecular


iFor metals the number of delocalized electrons increases acrossalong the period and hence stronger metallic bondstructure thus requiring a lot of heatenergy to weaken


The strength of a metallic bond also depends on the atomic radiussize The melting boiling point decrease as the atomic radiussize of metals increase due to decreased packing of larger atoms eg


The melting boiling point of Lithium is lower than that of Beryllium because Beryllium has twomore delocalized electrons and hence stronger metallic structurebond


The melting boiling point of Lithium is higher than that of Sodium because the atomic radiussize Lithium is smaller and hence better packed and hence forms stronger metallic structurebond


iiCarbongraphitecarbondiamond in period 2 and Silicon in period 3 form very well packed giant atomiccovalent structures held together by strong covalent bonds These elements have therefore very high meltingboiling points


Both Carbongraphite carbondiamond have smaller atomic radiussize than Silicon in period 3 and thus higher meltingboiling points due to bettercloser packing of smaller atoms in their well packed giant atomiccovalent structures


ii Nonmetals from group V alongacross the period form simple molecules joined by weak intermolecular vanderwaals force The weak intermolecular vanderwaals force require little energyheat to weaken leading to low meltingboiling points The strength of the intermolecular vanderwaals forces decrease with decrease in atomic radius size lowering the meltingboiling points alongacross the period and raising the meltingboiling points down the groupeg


The melting boiling point of Nitrogen is higher than that of Oxygen This is because the atomic radius size of Nitrogen is higher than that of Oxygen and hence stronger intermolecular vanderwaals forces between Nitrogen molecules


The melting boiling point of Chlorine is higher than that of Fluorine This is because the atomic radius size of Chlorine is higher than that of Fluorine and hence stronger intermolecular vanderwaals forces between Chlorine molecules


iii Rhombic sulphur exists as a puckered ring of S8atoms which are well packed Before melting the ring break and join to very long chains that entangle each other causing the unusually high meltingboiling point of Rhombic sulphur


iv Both sulphur and phosphorus exists as allotropes


Sulphur exists as Rhombicsulphur and monoclinicsulphur Rhombicsulphur is the stable form of sulphur at room temperature and pressure


Phosphorus exists as whitephosphorus and redphosphorus


Whitephosphorus is the stable form of Phosphorus at room temperature and pressure


4 State and explain the trend in density of elements in a period in the periodic table


Observation Density increase upto the elements in group IV then falls acrossalong the period successively


Explanation


Density is the mass per unit volume occupied by matterparticlesatomsmolecules of element


iFor metals the stronger metallic bond and the more delocalized electrons ensure a very well packed giant metallic structure that occupy less volume and thus higher densityThe more the number of delocalized electrons alongacross the period the higher the density eg


iAluminium has a higher density than sodium This is because aluminium has more three delocalized electrons than one sodium thus forms a very well packed giant metallic structure that occupy less volume per given massdensity


iiCarbongraphite carbondiamond and silicon in group IV form a well packed giant atomiccovalent structure that is continuously joined by strong covalent bonds hence occupy less volume per given massdensity


Carbongraphite form a less well packed giant hexagonal planar structure joined by Vanderwaals forces Its density 225gcm3 is therefore less than that of Carbondiamond353gcm3 and silicon233gcm3Both diamond and silicon have giant tetrahedral structure that is better packed Carbondiamond has smaller atomic radiussize than silicon Its density is thus higher because of better packing and subsequently higher density Carbondiamond is the hardest known natural substance by having the highest density


iii For nonmetals the strength of the intermolecular vanderwaals forces decreases with decrease in atomic radiussize alongacross the period This decreases the mass occupied by given volume of atoms in a molecule from group VI onwards eg


Phosphorus has a higher atomic radiussize than chlorine and Argon and thus stronger intermolecularvanderwaals forces that ensure a given mass of phosphorus occupy less volume than chlorine and neon


5 State and explain the trend in thermalelectrical conductivity of elements in a period in the periodic table


Observation


Increase alongacross the period from group I II and III then decrease in Group IV to drastically decrease in group V to VIII O


iMetals have free delocalized electrons that are responsible for thermalelectrical conductivityThermalelectrical conductivity increase with increase in number of delocalized electrons The thermal conductivity decrease with increase in temperatureheating eg Aluminium with three delocalized electrons from each atom in its metallic structure has the highest electrical thermal conductivity in period 3


iiCarbongraphite has also free 4th valency electrons that are delocalized within its layers of giant hexagonal planar structure They are responsible for the electrical conductivity of graphite


iiiSilicon and carbon diamond do not conduct electricity but conducts heat With each atom too close to each other in their very well packed giant tetrahedral structure heat transfer radiate between the atoms The thermal conductivity increase with increase in temperatureheating


iv All other nonmetals are poor nonconductor of heat and electricity They are made of molecules with no free mobile delocalized electrons in their structure


Periodicity of the oxides of elements alongacross period 3


The table below summarizes some properties of the oxides of elements in period 3 of the periodic table


1 All the oxides of elements in period 3 except those of sulphur and chlorine are solids at room temperature and pressure


2 Acrossalong the period bonding of the oxides changes from ionic in sodium oxide magnesium oxide and aluminium oxide show both ionic and covalent properties to covalent in the rest of the oxides


3 Acrossalong the period the structure of the oxides changes from giant ionic structure in sodium oxide magnesium oxide and aluminium oxide to giant atomiccovalent structure in silicon IV oxide The rest of the oxides form simple moleculesmolecular structure


4 Sodium oxide and magnesium oxide are basic alkaline in nature Aluminium oxide is amphotellic in nature shows both acidic and basic characteristics The rest of the oxides are acidic in nature


5 Ionic compoundsoxides have very high meltingboiling points because of the strong electrostatic attraction joining the giant ionic crystal lattice


The meltingboiling points increase from sodium oxide to aluminium oxide as the number of electrons involved in bonding increase increasing the strength of the ionic bondstructure


6 Silicon IV oxide is made of a well packed giant atomiccovalent structure joined by strong covalent bonds


This results in a solid with very high meltingboiling point


7Phosphorus V oxide sulphurIV oxide sulphur VI oxide and dichloride heptoxide exist as simple moleculesmolecular structure joined by weak vanderwaalsintermolecular forces


This results in them existing as low melting boiling point solidsgases


8 Ionic oxide conducts electricity in molten and aqueous states but not in solid


In solid state the ions are fusedfixed but on heating to molten state and when dissolved in water the ions are free  mobile


Sodium oxide magnesium oxide and aluminium oxide are therefore good conductors in molten and aqueous states


9 Covalent bonded oxides do not conduct electricity in solid molten or in aqueous states


This is because they do not have free  mobile ion Phosphorus V oxide sulphur IV oxide sulphur VI oxide and dichloride heptoxide are thus nonconductorsinsulators


10 Silicon IV oxide is a poorweak conductor of heat in solid state This is because it has very closely packed structure for heat to radiate conduct along its structure


11 Electopositivity decrease across the period while electronegativity increase across the period The oxides thus become less ionic and more covalent alongacross the period


12The steady change from giant ionic structure to giant atomic covalent structure then simple molecular structure lead to profound differences in the reaction of the oxides with wateracids and alkalisbases


i Reaction with water


a Ionic oxides react with water to form alkaline solutions eg


ISodium oxide reactsdissolves in water forming an alkaline solution of sodium hydroxide


Chemical equation Na2Os  H2O l  2NaOHaq


II Magnesium oxide slightly slowly reactsdissolves in water forming an alkaline solution of magnesium hydroxide


Chemical equation MgOs  2H2O l  MgOH 2 aq


III Aluminium oxide does reactsdissolves in water


b Nonmetallic oxides are acidic They react with water to form weakly acidic solutions


I Phosphorus V oxide readily reactsdissolves in water forming a weak acidic solution of phosphoric V acid


Chemical equation P4O10 s  6H2O l  4H3PO4 aq


Chemical equation P2O5 s  3H2O l  2H3PO4 aq


II Sulphur IV oxide readily reactsdissolves in water forming a weak acidic solution of sulphuric IV acid


Chemical equation SO2 g  H2O l  H2SO3 aq


Sulphur VI oxide quickly fumes in water to form concentrated sulphuric VI acid which is a strong acid


Chemical equation SO3 g  H2O l  H2SO4 aq


III Dichlorine oxide reacts with water to form weak acidic solution of chloricI acidhypochlorous acid


Chemical equation Cl2O g  H2O l  2HClO aq


IV Dichlorine heptoxide reacts with water to form weak acidic solution of chloricVII acid


Chemical equation Cl2O7 l  H2O l  2HClO4 aq


c Silicon IV oxide does not react with water


It reacts with hot concentrated alkalis forming silicate IV salts eg


Silicon IV oxide react with hot concentrated sodium hydroxide to form sodium silicate IV salt


Chemical equation SiO2 s  2NaOH aq  Na2SiO3 aq  H2O l


ii Reaction with dilute acids


a Ionic oxides react with dilute acids to form salt and water only This is a neutralization reaction eg


Chemical equation Na2Os  H2SO4 aq  Na2SO4 aq  H2Ol


Chemical equation MgOs  2HNO3aq  Mg NO3 2 aq  H2Ol


Chemical equation Al2O3 s  6HClaq  2AlCl3 aq  3H2Ol


Aluminium oxide is amphotellic and reacts with hot concentrated strong alkalis sodiumpotassium hydroxides to form complex sodium aluminateIII and potassium aluminateIII salt


Chemical equation Al2O3 s  2NaOH aq  3H2Ol  2 NaAlOH4 aq


Chemical equation Al2O3 s  2KOHaq  3H2Ol  2 KAlOH4 aq


b Acidic oxides do not react with dilute acids


c Periodicity of the Chlorides of elements alongacross period 3


The table below summarizes some properties of the chlorides of elements in period 3 of the periodic table


1 Sodium Chloride Magnesium chloride and aluminium chloride are solids at room temperature and pressure


SiliconIV chloride phosphorusIIIchloride and disulphur dichloride are liquids PhosphorusVchloride is a solid Both chlorine and sulphur chloride are gases


2 Acrossalong the period bonding changes from ionic in Sodium Chloride and Magnesium chloride to covalent in the rest of the chlorides


3 Anhydrous aluminium chloride is also a molecular compound Each aluminium atom is covalently bonded to three chlorine atoms


In vapourgaseous phasestate two molecules dimerizes to Al2O6 molecule through coordinatedative bonding


4 Acrossalong the period the structure changes from giant ionic in Sodium Chloride and Magnesium chloride to simple moleculesmolecular structure in the rest of the chlorides


5 Ionic chlorides have very high melting boiling points because of the strong ionic bondelectrostatic attraction between the ions in their crystal latticeThe rest of the chlorides have low melting boiling points because of the weak vanderwaal intermolecular forces


6 Sodium Chloride and Magnesium chloride in molten and aqueous state have freemobile ions and thus good electrical conductors Aluminium chloride is a poor conductor The rest of the chlorides do not conduct because they have no freemobile ions


7 Ionic chloride form neutral solutions with pH 7 These chlorides ionizedissociate completely into free cations and anionsie


Sodium Chloride and Magnesium chloride have pH7 because they are fullycompletely ionizeddissociated into free ions


Chemical equation NaCl s  Naaq  Claq Chemical equation MgCl2 s  Mg2aq  2Claq


8 Acrossalong the period from aluminium chloride hydrolysis of the chloride takes place when reactingdissolved in water


Hydrolysis is the reaction of a compound when dissolved in water


a Aluminium chloride is hydrolyzed by water to form aluminium hydroxide and fumes of hydrogen chloride gas Hydrogen chloride gas dissolves in water to acidic hydrochloric acid Hydrochloric acid is a strong acid with low pH and thus the mixture is strongly acidic


Chemical equation AlCl3 s  3H2Ol AlOH3s  3HClg


bSiliconIV chloride is hydrolyzed by water to form siliconIVoxide and fumes of hydrogen chloride gas Hydrogen chloride gas dissolves in water to acidic hydrochloric acid Hydrochloric acid is a strong acid with low pH and thus the mixture is strongly acidic


Chemical equation SiCl4 l  2H2Ol SiO2s  4HClg


This reaction is highly exothermic producing evolving a lot of heat that cause a rise in the temperature of the mixture


c Both phosphoric V chloride and phosphoric III chloride are hydrolyzed by water to form phosphoric V acid and phosphoric III acid respectively Fumes of hydrogen chloride gas are produced Hydrogen chloride gas dissolves in water to acidic hydrochloric acid Hydrochloric acid is a strong acid with low pH and thus the mixture is strongly acidic


Chemical equation PCl5 s  4H2Ol H3PO4aq  5HClg


Chemical equation PCl3 s  3H2Ol H3PO4aq  3HClg


This reaction is also highly exothermic producing evolving a lot of heat that cause a rise in the temperature of the mixture


d Disulphur dichloride similarly hydrolyzes in water to form yellow deposits of sulphur and produce a mixture of sulphur IV oxide and hydrogen chloride gas Hydrogen chloride gas dissolves in water to acidic hydrochloric acid Hydrochloric acid is a strong acid with low pH and thus the mixture is strongly acidic


Chemical equation 2S2Cl2 l  2H2Ol 3Ss  SO2g  4HClg


D COMPREHENSIVE REVISION QUESTIONS


1The grid below represents periodic table Study it and answer the questions that follow The letters do not represent the actual symbols of the elements


a I Indicate on the grid the position of an element represented by letter N whose electronic configuration of a divalent cation is 288   1 mark


II Name the bond formed between D and H react Explain your answer2 marks


Ionicelectrovalent


D is electropositive thus donates two electrons to electronegative H


III Write an equation for the reaction between B and water  1 mark


Chemical equation 2B s  2H2Ol  2BOHaq  H2 g


IV How do the atomic radii of I and L compare Explain 2 marks


V In terms of structure and bonding explain why the oxide of G has lower melting point than oxide of L 2 marks


b Study the information given below and answer the questions that follow


IWhy is the formula of aluminium chloride given as Al2Cl6 and not AlCl3   1 mark


II Give two chlorides that are liquid at room temperature Give a reason for the answer 2 marks


III Give a reason why Al2Cl6 has a lower melting point than MgCl2 although both Al and Mg are metals 1 mark


IV Which of the chlorides would remain in liquid state for the highest temperature range explain why 2 mark


Kakamega


2 a Study the information given below and answer the questions that follow


i Which elements are nonmetals Give a reason 2mks


ii Explain why the melting point of the oxide of R is higher than that of the oxide of S 2mks


iii Give two elements that would react vigorously with each other Explain your answer 2mks


b Study the information in the table below and answer the questions that follow The letters do not represent the actual symbols of the elements


i What chemical family do the elements A B and C belong 1mk


ii What is meant by the term ionization energy 1mk


iii The 2nd ionization energy is higher that the 1st ionization energy of each Explain


ivWhen a piece of element C is placed in cold water it sinks to the bottom and an effervescence of a colourless gas that burns explosively is produced Use a simple diagram to illustrate how this gas can be collected during this experiment 3mks


3 The grid below represents part of the periodic table The letters do not represent the actual symbols


a Select the most reactive nonmetal 1mk


b Write the formula of the compound consisting of


ID and Z only 2mk


II X and Z


c Select an element that can form an ion of change 2 1mk


d Which element has the least ionization energy Explain 2mks


e Suggest with reasons a likely pH value of an aqueous solution of the chlorine of3mks


T


B


X


f To which chemical family do the following elements belong 2mk


J


V


g An element K has relative atomic mass of 402It has two isotopes of masses 39 and 42 Calculate the relative abundance of each isotope 3mks


4The grid below shows part of the periodic table study it and answer the questions that follow The letters do not represent the true symbols


a Which element forms ions with charge of 2 Explain 2mks


b What is the nature of the oxide formed by C 1mk


c How does the reactivity of H compare with that of E Explain 2mks


dWrite down a balanced equation between B and Chlorine 1mk


e Explain how the atomic radii of F and G compare 1mk


f If the oxides of F and D are separately dissolved in water state and explain the effects of their aqueous solutions on litmus 3mks


5 a The grid below show part of the periodic tableThe letter do not represent the actual symbolsUse it to answer the questions that follow


iSelect the most reactive nonmetal 1mk


iiSelect an element that forms a divalent cation 1mk


iiiElement Z has atomic number 14Show its position in the grid 1mk


ivHow do the atomic radii of U and J compare 2mks


vHow do electrical conductivity of A and Y compare 2mks


viHow does the boiling point of elements K L and M vary Explain 2mks


b The table below gives information on four elements by letters K L M and N Study it and answer the questions that follow The letters do not represent the actual symbols of the elements


a Which two elements have similar properties Explain 2mks


b Which element is a nonmetal Explain 1mk


c Which one of the elements is the strongest reducing agent 1mk


6 The grid given below represents part of the periodic table study it and answer the questions that follow The letters do not represent the actual symbols of the elements


i What name is given to the group of elements to which C and F belong 1mk


ii Which letter represents the element that is the least reactive 1mk


iii What type of bond is formed when B and E react Explain 2mks


ivWrite formula of the compound formed where elements D and oxygen gas react 1mk


v On the grid indicate the a tick  the position of element G which is in the third period of the periodic table and forms G3 ions 1mk


b Study the information in the table below and answer the questions that follow The letter do not represents the actual symbols of the substance


Ii Which substance would dissolve in water and could be separated from the solution by fractional distillation 1mk


ii Which substances is a liquid at room temperature and when mixed with water two layers would be formed 1mk


II Which letter represents a substance that is a gas at room temperature and which can be collected


i Over water 1mk


ii By downward displacement of air Density of air at room temperature  129 x 103 gC 1mk


AATOMIC STRUCTURE


The atom is the smallest particle of an element that take part in a chemical reaction The atom is made up of three subatomic particles


iProtons


iiElectrons


iiiNeutrons


1 The proton is positively charged


2 Is found in the centre of an atom called nucleus


3 It has a relative mass 1


4 The number of protons in a atom of an element is its Atomic number


1 The Electrons is negatively charged


2 Is found in fixed regions surrounding the centre of an atom called energy levelsorbitals


3 It has a relative mass 11840


4 The number of protons and electrons in a atom of an element is always equal


1 The Neutron is neither positively nor negatively charged thus neutral


2 Like protons it is found in the centre of an atom called nucleus


4 The number of protons and neutrons in a atom of an element is its Mass number


Diagram showing the relative positions of protons electrons and neutrons in an atom of an element


Diagram showing the relative positions of protons electrons and neutrons in an atom of Carbon


The table below show atomic structure of the 1st twenty elements


Most atoms of elements exist as isotopes


Isotopes are atoms of the same element having the same number of protonsatomic number but different number of neutronsmass number


By convention isotopes are written with the mass number as superscript and the atomic number as subscript to the left of the chemical symbol of the element ie


mass number


atomic number m n X symbol of element


Below is the conventional method of writing the 1st twenty elements showing the mass numbers and atomic numbers


11H 42He 73Li 94Be 115B 126C


147N 168O 199F 2010Ne 2311Na 2412Mg


2713Al 2814Si 3115P 3216S 3517Cl 4018Ar


3919K 4020C


The table below shows some common natural isotopes of some elements


The mass of an average atom is very small 1022 gMasses of atoms are therefore expressed in relation to a chosen element


The atom recommended is 12C isotope whose mass is arbitrarily assigned as 12000 atomic mass unitsamu


All other atoms are compared to the mass of 12C isotope to give the relative at The relative atomic massRAM is therefore defined as the mass of average atom of an element compared to 112 an atom of 12C isotope whose mass is arbitrarily fixed as 12000 atomic mass unitsamu  ie


RAM  mass of atom of an element


112 of one atom of 12C isotope


Accurate relative atomic masses RAM are got from the mass spectrometer Mass spectrometer determines the isotopes of the element and their relative abundanceavailability


Using the relative abundancesavailability of the isotopes the relative atomic mass RAM can be determined calculated as in the below examples


Chlorine occurs as 75 3517Cl and 25 3717Cl isotopes Calculate the relative atomic mass of Chlorine


Working


100 atoms of chlorine contains 75 atoms of 3517Cl isotopes


100 atoms of chlorine contains 75 atoms of 3717Cl isotopes


Therefore


RAM of chlorine   75100 x 35  25100 x 37  355


Note that


Relative atomic mass has no units


More atoms of chlorine exist as 3517Cl75 than as 3717Cl25 therefore RAM is nearer to the more abundant isotope


Calculate the relative atomic mass of potassium given that it exist as


931 3919K  001 4019K  689 4119K


100 atoms of potassium contains 931 atoms of 3919K isotopes


100 atoms of potassium contains 001 atoms of 4019K isotopes


100 atoms of potassium contains 689 atoms of 4119K isotopes


RAM of potassium  931100 x39  001100 x 40 689 100 x 39 Note that


More atoms of potassium exist as 3919K 931 therefore RAM is nearer to the more abundant 3919K isotope


Calculate the relative atomic mass of Neon given that it exist as


9092 2010Ne  026 2110Ne  882 2210Ne


Working


100 atoms of Neon contains 9092 atoms of 2010Ne isotopes


100 atoms of Neon contains 026 atoms of 2110Ne isotopes