FOSSIL FUEL -COAL / IGCSE CHEMISTRY / GCSE-CHEMISTRY

Just as coal has formed by the action of heat and pressureon the remains of trees and plants on land over millions of years,so oil and natural gas have formed by the action of heat and pressureon the remains of sea plants and

animals over millions of years.
The remains were buried in sediments which excluded the air (kept out oxygen) and stopped them decaying.
More sediment buried the remains deeper and deeper until pressure and heateventually turned them into coal, oil and natural gas.
They are called fossil fuels because they are buried underground(from Latin fossilis - dug up).
Fossil fuels are a finite resource and non-renewable.
The oil deposits are formed in porous rock sediments.
Porous rock has pores in it. Pores are small holes (see for example sandstone).The small holes allow the oil and natural gas to pass through the rockand rise until they are stopped by a layer of non-porous rock.
Non-porous rock (for example shale) has no holes,and acts as a barrier to prevent the oil and natural gas rising.The oil and natural gas become trapped underground.
The oil is called crude oil (or petroleum, from Latin - rock oil),and has natural gas in it or in a pocket above it trapped by non-porous rock.
Drilling through the rock allows the oil and gas to escape to the surface.
Natural gas is mostly methane (CH4).Crude oil is a mixture of substances (mostly hydrocarbons).

ACID RAIN -IGCSE/GCSE CHEMISTRY NOTES

2) Fossil fuels are burnt on a huge scale - Acid Rain.
Coal (and to a lesser degree Oil and Natural Gas) contain sulphur.When they are burnt the sulfur oxidises (reacts with oxygen)to form sulfur dioxide gas.
sulfur + oxygen sulphur dioxide.
S(s) + O2(g) SO2(g)
Sulfur dioxide gas is acidic, poisonous, and smells of bad eggs.It is removed from the burnt waste gases and used in the contact process.If it gets into the atmosphere it reacts with water and oxygen in the airto form a dilute solution of sulfuric acid.This sulfuric acid is the main pollutant in acid rain.Natural rain is slightly acidic due to dissolved carbon dioxide.Natural rain has a pH of 5·5, acid rain has a pH of
4.The second most important pollutant in acid rain is nitric acid - see next page.
Acid rain kills trees.It runs into rivers and gathers in lakes.Eventually, lakes become too acidic, and plants and fish begin to die.Acid rain reacts with limestone and damages limestone buildings.
Powdered limestone or slaked lime can be added to soils or lakesto make them less acidic. It would be better if we could avoid or reducepollutant gas emissions in the first place.

AMPHOTERIC WATER /IGCSE /GCSE CHEMISTRY

In a sample of water,a very small number of water molecules will form ions.
water hydrogen ion + hydroxide ion.
H2O(l) H+(aq) + OH-(aq)
This ionisation is reversible (shown by the arrow).
The hydrogen ion is acidic. The hydroxide ion is alkaline.Water forms equal amounts of both ions, and so water is neutral.
Compare this reaction with neutralisation

titraction IGCSE/GCSE/ GCE-CHEMISTRY NOTES

In the example below, an acid and an alkali react to make sodium chloride.
hydrochloric acid + sodium hydroxide sodium chloride + water.
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
The burette is filled with hydrochloric acid.A known quantity of alkali (say 50 cm3 sodium hydroxide)is released from a pipette into the conical flask.The tap on the burette is turned opento allow the acid to be added drop by drop into the alkali.The alkali contains an indicator (phenolphthalein)which is pink in an alkali and colourless in an acid.
When enough acid has been added to neutralise the alkalithe indicator changes from pink to colourless.
The titration can be repeated using the same amounts of acid and alkalibut without the indicator.Pure salt crystals which are free from indicatorcan then be crystallised from the neutral solution.
Alternatively a pH meter can be used to find the end point.At neutralisation the pH is 7.
Indicators used for weak Acids or Alkalis.
Indicator
Titration
Colour in Acid
Colour in Alkali
Methyl Orange
Strong Acid+Weak Alkali
Red
Yellow
Phenolphthalein


Strong Alkali+Weak Acid
Colourless
Pink
Either Indicator can be used for a Strong Alkali + Strong Acid.
Universal indicator is not usually used for a titrationbecause it changes gradually giving different colours for a different pH.Methyl orange or phenolphthalein are usedbecause they give a sudden change in colour at neutralisationwhich makes it easier to see the end point of the titration

Ph - acids and bases - igcse/gcse /chemistry notes

pH is a measure of how acidic or how alkaline a solution in water is.The pH scale goes from 1 to 14,with 1 being very strongly acidic,and 14 being very strongly alkaline.A pH of 7 is neutral.
You can measure the pH of a solution using universal indicator.Just as litmus paper will be red for an acid and blue for an alkali,so universal indicator is a mixture of indicatorswhich will give a different colour for a different pH.
Any acid will have a pH of less than 7.Any alkali will have a pH of more than 7.
A strong acid (HCl or H2SO4 or HNO3 )will have a pH of 1 (red).
A weak acid will have a pH of 3 to 4 (orange).Examples of weak acids are ethanoic acid (vinegar),citric acid (lemon juice) and rain water.
Rain water has a natural pH of 5·5 (see carbonic acid).
Water and salts are neutral, pH 7 (green).
A weak alkali (ammonia) will have a pH of 11 to 12 (blue).
A strong alkali (Ca(OH)2or NaOH) will have a pH of 14 (purple).

STRONG AND WEEK ACIDS / IGCSE/GCSE / GCE- CHEMISTRY

Acids and Alkalis
Strong and Weak Acids - Strength and Concentration.
Acids and alkalis can be described as strong or weak.This does not mean the same as concentrated or dilute.
The strength of an acid or alkali depends on how ionised it is in water.
A strong acid or alkali is completely (100%) ionised. For hydrochloric acid
hydrogen chloride (in water) hydrogen ion + chloride ionHCl(aq) H+(aq) + Cl-(aq)
All of the hydrogen chloride moleculesbecome hydrogen ions and chloride ions in water(see examples for other strong acids).
For sodium hydroxide
sodium hydroxide (in water) sodium ion + hydroxide ionNaOH(aq)
Na+(aq) + OH-(aq)
Sodium hydroxide exists as ions both in water and in the solid.(see examples for other strong alkalis).

A weak acid or alkali is only partly (less than 100%) ionised.
For ethanoic acid
ethanoic acid (in water) hydrogen ion + ethanoic ion

CH3CO2H(aq) H+(aq) + CH3CO2-(aq)
Some of the ethanoic acid molecules become ions in waterbut most of them stay as molecules.The reaction is reversible (shown by the arrow).
For ammonia
ammonia + water ammonium ion + hydroxide ion

NH3(g) + H2O(l) NH4+(aq) + OH-(aq)
Some of the ammonia molecules become ions in waterbut most of them stay as molecules.

IONIC EQUATION - NEUTALISATION / IGCSE /GCSE/GSE /O-LEVEL CHEMISTRY NOTES

If we take the reaction betweenhydrochloric acid and sodium hydroxide from the previous page,
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
We can write it in terms of ions,since both the acid and the alkali form ions in water.
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) Na+(aq) + Cl-(aq) + H2O(l)
Chloride ions (Cl-(aq)), and sodium ions (Na+(aq))appear on both sides of the equation.They are spectator ions, they are not changed by the reaction,and so they may be left out of the equation.
This leaves the equation
hydrogen ion + hydroxide ion waterH+(aq) + OH-(aq) H2O(l)
Compare this reaction with the ionisation of water.
This is the reaction that always occurswhen an acid + alkali salt + water.The hydrogen ion of the acid + the hydroxide ion of the alkalicombine to form water,leaving the metal from the alkali and the non-metal from the acidto form a salt solution.
How much acid is needed to neutralise an alkali?

alkalies / IGCSE /GCSE CHEMISTRY NOTES

Alkali is pronounced like alcohol,with 'lie' at the end instead of 'hol'.
An alkali is any substance which produces OH- ions in water.OH- ions are called hydroxide ions
A substance which will neutralise an acid,but does not dissolve in water, is called a base.For example,copper(II) oxide, iron(II) oxide and zinc carbonate are bases,they do not dissolve in Any base which dissolves in water is called an alkali.

Acids and Alkalis
Arrhenius, Lowry and Brønsted.
Arrhenius defined an acid as a substance which produces hydrogen ionswritten H+(aq) in water
Lowry and Brønsted defined an acid as a proton donorand a base as a proton acceptor.
If you look at the reaction below
hydrochloric acid + water hydroxonium ion + chloride ionHCl(aq) + H2O(l) H3O+(aq) + Cl-(aq)
According to Arrhenius, hydrochloric acid is an acidbecause it produces hydrogen ions in water(hydrogen ions in water become hydroxonium ions).
According to Lowry and Brønstedhydrochloric acid is an acid because it is a proton donor.A proton is a hydrogen ion.A proton donor is a substance which gives a hydrogen ion away.If you look at the reaction abovehydrochloric acid gives a hydrogen ion to water.
A base is a proton acceptor.This means that a base will gain a hydrogen ion.Water is a base when it is put with hydrochloric acidbecause water will gain a hydrogen ion to become H3O+.
acid + base acid + base

HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq)
On the right side of the arrow, H3O+ is an acidbecause it can give away a hydrogen ion to become H2O.Cl- is a base because it can gain a hydrogen ion to become HCl.
Is water always a base according to Lowry and Brønsted?

activation energy/ collision theroy /igcse/gcse chemistry notes

Collision theory

it says that a chemical reaction can only occurbetween particles when they collide (hit each other).Particles may be atoms, ions or molecules

activation energy

There is a minimum amount of energywhich colliding particles need in order to react with each other.If the colliding particles have less than this minimum energythen they just bounce off each other and no reaction occurs.This minimum energy is called the activation energy.
The faster the particles are going, the more energy they have.Fast moving particles are more likely to react when they collide.You can make particles move more quickly by heating them up

HOW FAR, HOW FAST? ENERGETICS/igcse/ gcse / chemistry notes

HOW FAR, HOW FAST? ENERGETICS
Energy changes in reactionsIn every chemical reaction bonds must be broken in the reactants and new bonds must be formed in the products.
For example in the Haber Process nitrogen reacts with hydrogen to produce ammonia, NH3.
Initially bonds must be broken between the nitrogen atoms in the nitrogen molecule and the hydrogen atoms in the hydrogen molecule.


As energy must be put in to break the bonds and energy is released when the new bonds form this means that most chemical reactions involve energy changes.
When we do reactions in the laboratory the apparatus is usually open to the atmosphere so, if the reaction produces a gas, it is allowed to escape into the air. This means that the reactions are occurring under conditions of constant pressure.
In this case, when a gas is produced in a reaction, it must push back the surrounding air to escape from the apparatus. This uses up some energy. The energy changes we measure under these conditions take account of all these factors and are know as enthalpy changes. They are given the symbol .
The energy changes are mainly in the form of heat though other forms of energy can be involved (eg. burning of magnesium releases a lot of light energy as well as heat).

When energy is lost (or given out) in a reaction it is called an exothermic reaction. As energy is given out in an exothermic reaction the mixture gets hotter and the temperature rises. This can be seen in the reaction between the magnesium and oxygen, as the temperature of the air around the burning magnesium gets very hot.
When energy is taken in by a reaction it is called an endothermic reaction. As energy is taken in the reaction mixture gets colder and its temperature falls.
Sign convention and energy diagramsEnthalpy changes are usually measured in kilojoules (or joules). The symbol for kilojoules is kj.
In enthalpy changes that are exothermic the energy has been lost in the reaction. It is therefore given a negative sign.
In changes that are endothermic the energy has gained by the reaction mixture so it is given a positive sign.


In energy diagrams the energy changes are shown as follows:
This is an example of an energy diagram showing an exothermic reaction.


Standard enthalpy changesSo that scientists can compare data from experiments a set of standard conditions have been defined. These are
Pressure of 1 atmosphere (or 101kPa or 101,000Pa using SI units)
Temperature of 25oC (or 298K)
All solutions have a concentration of 1M (1 mol.dm-3)Standard enthalpy changes are ones that carried out under standard conditions. The symbol used for a standard enthalpy change is---
A so-called thermochemical equation summarises all the information for a reaction.
It shows

1. The amounts in moles of the reactants and products.2. The quantities of energy.3. The conditions of temperature and pressure.

An example is
If half the amounts are used, then the energy change is halved as well.
It is important to include the state symbols (s) for solids, (1) for liquids, (g) for gases and (aq) for aqueous solutions in all equations connected with energy changes.

carbon chemistry -igcse/ gcse/ gce olecvel notes

CARBON CHEMISTRY
Carbon chemistry is also called organic chemistry. This is because most carbon compounds are obtained from living organisms.
You will spend a lot of time on this topic and this is because there are more compounds of carbon than all the other elements put together! This is because carbon is the only element than can form strong bonds to other carbon atoms. There is no limit to the numbers of carbons that can be joined together, and they can join together in an unlimited number of different patterns.
Before we look at organising and naming all these different compounds, let us first find out where we find these compounds.
1. Fossil Fuels.
Then test your understanding of this topic using:
You should now be an expert on the sources of carbon compounds! Crude oil is however a complex mixture of up to one hundred compounds and must be purified before we can use it. This is done using a process known as fractional distillation.
2. Fractional Distillation.
Then read through the notes on the purification of crude oil. You can then try the test at:

work sheet / chemical bonding / igcse/ gcse/ gce olevel

Checkpoint.
Q1. Look at the ten compounds below, and decide whether the bonding will be ionic, covalent or metallic. You may need a periodic table to help you decide which elements are metals and non-metals.
a. potassium chloride (KCl)
b. carbon dioxide (CO2)
c. hydrogen chloride (HCl)
d.carbon dioxide (CO2)
e. iron oxide (Fe2O3)
f. magnesium iodide (MgI2)
g. fluorine gas (F2)
h. brass (a mixture of zinc and copper)
i. ammonia (NH3)
j. calcium carbonate (CaCO3)
Q2. Which of the compounds above are molecules?

Kinetic Theory igcse/gcse/ olevel chemistry notes

Kinetic Theory
Three states of matter – solids, liquids and gases.
Matter made up of very small particles in constant motion.
In solids the particles are packed very close together. They vibrate about fixed positions and have strong forces of attraction between them.
Solids :
have a high density
can not be compressed
do not flow
have a fixed shape
have a fixed size

In liquids the particles are close together but not as close as they are in solids. They can move around in any direction and are not fixed in position. The forces of attraction between them are still quite strong but, again, not as strong as in solids.
Liquids :
have a medium density
can not be compressed
can flow
have the shape of their container
have a fixed size

In gases the particles are very far apart with large distances between them. They move around very quickly in all directions and the forces of attraction between them are very, very weak.
Gases :
have a very low density
can be compressed
can flow
have the shape of their container
have the size of their container

You can change the state of a substance by heating or cooling it.

When the change is from a solid to a liquid it is called melting.
When the change is from a liquid to a gas it is called evaporating.
When the change is from a liquid to a solid it is called freezing.
When the change is from a gas to a liquid it is called condensing.
When a solid is heated it changes to a liquid and then a gas. A graph of temperature against time for this process would look like this:

KEY
A:
Over portion A the particles in the solid are vibrating more and more as they gain the heat energy so the temperature rises.
B:
Over portion B, as the solid changes into a liquid, all the energy is being used to overcome the strong forces of attraction between the particles and separate them so the temperature does not rise
C:
Over portion C the particles move around faster as they gain energy so the temperature rises.
D:
Over portion D the heat energy is all being used to separate the particles and overcome the forces of attraction between them so again the temperature does not rise as the liquid changes to a gas
E:
Over part E the temp rises as the particles move around faster

state symbol -IGCSE/ GCSE NOTES

State symbolsThe state symbols are put in a balanced equation to show whether something is a solid, liquid, gas or dissolved in water (aqueous solution).
The symbols for these are:
state
Symbol
Solid
(s)
Liquid
(l)
Gas
(g)
Aqueous
(aq)
Magnesium + oxygen --> magnesium oxide
2Mg{s} + O2{g} --> 2MgO{s}
hydrochloric acid + calcium carbonate --> calcium chloride + carbon dioxide + water
2HCl (aq) + CaCO3 ---> CaCl2(aq) + CO2(aq) + H20(l)
Balanced equations and ionic equations

Ionic equations only show ions which change in a reaction and ignore those which do not change. E.g.word equation
hydrochloric acid + sodium hydroxide --> sodium chloride + water
balanced chemical equation
HCl(aq) + NaOH(aq) ---> NaCl(aq) + H2O(l)
ionic equationH+(aq) + OH-(aq) ---> H2O(l)
E.g. in the electrolysis of sodium chlorideCl-(aq) --> Cl(g) + e-2Cl(g) --> Cl2(g)
Working out formulae from reacting masseselements reacting magnesium chlorinesymbols of elements Mg Clmasses reacting (from experiment) 2.4g 7.1gmolar mass (look up relative atomic 24g/mol 35.5g/molmass in periodic table)amounts (amount = mass/molar mass) 2.4g/24g/mol 7.1g/35.5g/mol = 0.1mol 0.2molratio of atoms (divide by smallest) 1 : 2formula MgCl2
Work out formulae of compounds formed when the following react:56g of iron and 32g of sulphur (Fe =56, S =32)2g of hydrogen and 16g of oxygen (H=1, O=16)14g of lithium and 16g of oxygen (Li=7)32g of copper and 8g of oxygen (Cu=64)6.4g of copper and 0.8g of oxygen.
C2.08 Calculating reacting masses using equationsYou can work out ratio of the masses of products and reactants by simply multiplying the number of moles shown in the equation by the formula mass of each substance.
Example 1: What mass of magnesium oxide can be made from 12g of magnesium? Relative atomic masses are Mg =24, O = 16. equation 2Mg(s) + O2(g) --> 2MgO(s) formula 2*24 1(16*2) 2(24+16)masses =48 =32 =80 reacting 48g of Mg forms 80g of MgOmasses 1g of Mg forms 80/48 g of MgO 12g of Mg forms 12*80/48 g of MgO = 20g
Example 2: What mass of magnesium oxide can be made from 12g of magnesium? equation 2Mg(s) + O2(g) --> 2MgO(s) amounts 2 moles 1 mole 2 moles masses 2*24 1{16*2} 2{24+16} =48g =32g =80g so 48g Mg forms 80g MgO 1g Mg forms 80/48 g MgO 12g Mg forms 12*80/48 g MgO = 20g Also note that the ratio of amounts of reactants and products in the equation above can be written as:
Amount of Mg/amount of O2 =2/1 Or Amount of O2/amount of MgO = 1/2
problem
You decide to travel from London to Delhi for a holiday and hire your own Airbus 319 jet. You fly the 2183 miles to Cairo first and make 9.4 tons of carbon dioxide. After seeing the Pryramids you fly the 1018 miles to Riyadh making 4.5 tons of carbon dioxide. After a brief stop in the Saudi Capital you fly on the 1900 miles to Delhi making 8.1 tons of carbon dioxide. Finally you return to London travelling 4171 miles and releasing 22 tons of carbon dioxide. The jet burns the fuel kerosine C15H32 in the reaction
C15H32 + 23O2 --> 15CO2 +16H2O(a)
How much carbon dioxide was released into the atmosphere during the trip?(b) What mass of kerosine was burnt?(c) What effect does the trip have of the environment?

formulae and equations

Formulae can be worked out from valency.
Valency
atom or ion with that valency
valency 1
hydrogen, group 1 e.g. sodium and potassium, group 7e.g. fluorine and chlorine, ammonium NH4+, hydroxide OH-, nitrate NO3-
valency 2
group 2 e.g. magnesium and calcium, group 6, sulphate SO42-, carbonate CO32-, copper , II Cu2+ iron II Fe2+
valency 3
group 3 e.g. aluminium, group 5, phosphate PO43-, iron III Fe3+
valency 4
group 4 e.g. carbon
naming of ionic compounds
aluminium oxide
valency 3 2
symbols Al O
formula Al2O3

Use valency to work out the formulae of the following compounds: sodium chloride, potassium bromide, magnesium oxide, calcium sulphide, aluminium nitride, calcium iodide, lithium oxide, aluminium chloride, aluminium sulphide, magnesium nitride.

Calculating relative formula massAdd up the relative atomic mass (found in periodic table) of each atom in the compounsds

e.g. Al203 relative atomic masses of Al = 27, O = 16 (found in periodic table). The formula shows 2 atoms of aluminium and 3 atoms of oxygen so:formula mass of = (2*27) + (3*16) =54 + 48 = 102

Work out the relative formula masses of the following: MgO, FeS, O2, H2O, CaBr2, Na2S, CaCO3, NaOH, HCl, (NH4)2SO4. Relative atomic masses Mg=24, O=16, Fe=56, S=32, Ca=40, Br=80, C=12, Na=23, H=1, Cl=35.5.
Simple balanced equationsIt is possible to write balanced equations for reactions. For example substances such as hydrogen and magnesium combine with oxygen. One method to write them is:

Write a word equation first.

Magnesium + oxygen --> magnesium oxide

Write in the formulae of the substances used.

Mg + O2 --> MgO

Balance the equation so that each element has the same number of atoms on each side.

2Mg + O2 --> 2MgO

sodium + oxygen --> sodium oxide

4Na{s} + O2{g} --> 2Na2O{s}(word equation)

hydrogen + oxygen ---> water(formulae)

H2 + O2 -----> H2O(balance) 2H2 + O2 -----> 2H2O

Representing reactions-IGSCE /GCSE CHEMISTRY NOTES

Word equations
If a reaction occurs between magnesium and oxygen, magnesium oxide is produced, here is the word equation for this reaction: -
magnesium + oxygen --> magnesium oxide
Some other examples are:
hydrochloric acid + calcium carbonate --> calcium chloride+ carbon dioxide+ water
sodium + water --> sodium hydroxide + hydrogen
hydrochloric acid + sodium hydroxide --> sodium chloride + water
Write word equations for the reactions in which the following compounds form from a halogen and another suitable element:
hydrogen fluoride, hydrogen chloride, iron III chloride, iron III bromide, sodium chloride, copper chloride.

Formulae
The formula of an element or compound is simply the symbol of each element present and numbers to show how many atoms are present. Carbon dioxide has the formula CO2. This means that it has one carbon atom and two oxygen atoms in each molecule.

substance formula substance formula
methane CH4 bromine Br2
ethane C2H6 hydrogen H2
propane C3H8 ethanol C2H5OH

diamond and graphite / IGCSE/ GCSE CHEMISTRY NOTES

Giant structures with covalent bonds

In giant covalent structures all the atoms are bonded to each other by strong covalent bonds so they have very high melting and boiling points. They do not usually conduct electricity even if in the liquid state. Diamond and graphite are two examples, which are made from carbon atoms. These two different types of the same element are called allotropes.
Diamond: Each carbon atom forms four covalent bonds in a very rigid giant covalent structure.
Diamond:

Graphite: exists as layers of carbon atoms each held in place by three strong covalent bonds. Each layer is held to the one above it by weak bonds.
Physical properties of giant covalent structures

Giant molecular structures have very high melting points because all atoms are held firmly in place by strong covalent bonds. In graphite each carbon atom is held in place by three strong covalent bonds which gives graphite a high melting point. In diamond 4 strong covalent bonds holds each atom in place. This also gives diamond a very high melting point. The four bonds make diamond very hard. Graphite has weak bonds between layers so the layers slip over each other making graphite soft.
They do not usually conduct electricity even when molten because there are no charged particles to carry the current. There are free electrons between layers in graphite so it conducts electricity.
Explaining differences between propertiesSimple molecular substances like water have weak bonds between molecules so melt at low temperatures because little energy is needed to separate the molecules. Giant covalent structures like diamond have strong covalent bonds holding each atom in place. They melt at high temperatures because a lot of energy is needed to break these strong bonds.

chemicalk bonding- ionic bonding/ covalent bonding / igcse / gcse notes

Bonding
Elements forming compounds with chemical bonds
Electron transfer and ionic bondsAtoms have no charge. A charged particle is called an ion. If an atom loses an electron, it becomes a positively charged (+) ion. An ion that is positively charged is known as a cation. If an atom gains an electron, it becomes a negatively charged (-) ion. An ion that is negatively is known as an anion. The negative and positive ions attract each other to form an ionic bond.

Complete the gaps in the text below:_____ have no charge. A charged particle is called an ___. If an atom loses an ________, it becomes a positively charged (+) ion. An ion that is positively charged is known as a ______. If an atom gains an electron, it becomes a negatively charged (-) ion. An ___ that is negatively is known as an anion. The negative and positive ions attract each other to form an _____ bond.

The formation of sodium and chloride io

Draw atoms and ions for lithium, potassium, fluorine, magnesium, oxygen, sulfur and alumin

b Draw diagrams of ionic bonding in LiF, KF, LiCl, NaF, MgCl2, AlF3, MgO, MgS, Na2O and Al2O3.
Physical properties of giant ionic structuresIonic bonds form when metal and non-metal atoms join. A substance with ionic bonding has an ionic structure. Each ion is firmly held in place by strong ionic bonds so they have high melting and boiling points. If melted, charged ions become free to carry an electric current. The ions also become free if dissolved in water so solutions are also electrolytes. The solids are insulators because the ions are not free to move and cannot carry a current. Sodium chloride NaCl, and magnesium oxide MgO are good examples.
Task C1.10 Pick out the substances which are (a) ionic (b) have covalent bonds (c) have high melting points (d) conduct electricity when molten: sodium chloride, sulfur dioxide, magnesium oxide, iron fluoride, carbon dioxide, NaBr, H2O, NH3, Al2O3, KCl.
Covalent bonds and electron sharingNon-metal atoms join using covalent bonds. When a covalent bond is formed, atoms share their electrons. The atoms then have full shells. One covalent bond needs one shared electron from each atom. Each atom involved has to make enough covalent bonds to fill up its outer shell. Sharing electrons is called covalent bonding. Below is a diagram to show hydrogen gas (H2).
Dot and cross diagrams


Draw atoms of F, H, O, N and C.Draw a dot and cross diagram for fluorine F2, hydrogen fluoride HF, water H2O, ammonia NH3, methane CH4, oxygen O2, nitrogen N2, ethene C2H4

Physical properties of simple molecular substances

Simple Molecular Substances have low melting and boiling points and most are gases or liquids at room temperature. This is because of weak forces between the molecules. Molecular substances do not conduct electricity, because there are no ions. E.g. water.Draw diagrams to show how the molecular structures for the following might look: fluorine F2, hydrogen fluoride HF, water H2O, ammonia NH3, methane CH4, oxygen O2, nitroge

WORK SHEET / ELECTROLYSIS / IGCSE / GCSE / OXIDATION AND READUCTION

1.Name the process of the industrial extraction of oxygen
……………………………………………

What is oxidation?……………………………………….

What is reduction?…………………………

Give example for oxidizing

agent…………………………………………….

Give example for reducing agent…………………

2.Complete these reactions
( balance )S+ O2__________
Fe + O2__________
Mg + O2 _____________
Cu + O 2_________
C +O 2 ___________
CH4 + O2_____________

3.Describe an experiment to determine the % of oxygen in air (copper turning )

4.Give the reaction of sulphur dioxide with water…………………………………………….

Give two equations for the reactions in contact process ( impor)
…………………………………………………………….……………………………………………………….

5.Which one is reversible reaction?……………………………………………


6.Describe the laboratory preparation of carbon dioxideGive the equation for reaction……………………………………………………………………………….


7.Give the uses of carbon dioxide…………………………………

8.What is dry ice?……………………………

9.Give two uses of dry ice………………………………………..
10.Give the reaction of nitrogen with oxygen
( 2 equations)………………………………………
11.Give the conditions needed for the rusting of iron

12.Give method for the prevention of rusting

13.How this method work in preventing rusting of ironExplain sacrificial protection. Give example.

14.Sodium can be prepared by electrolysis only were as iron fro blast furnace give your answer (in terms of reactivity)

15.Give the allotropes of sulphur

Name the process by which nitrogen is industrially

extracted………………………………………………………………..

Give one use of nitrogen……………………………………….

16.Why nitrogen is used for protecting dry foods……………………………………………….

Describe the laboratory preparation of ammonia

17.Give the equation for the preparation………………………………………………………………..

18.Give physical properties of ammonia………………………………………………………………………………………

19.ammonia dissolve in water forms …………………………………………
give the equation………………
…………………………
20.Give the uses of ammoniato prepare nitric acidto prepare fertilizer

Give the reaction between sulphuric acid and ammonia ( important)

NH 3 + H2SO4 ----------- (NH4)2 SO4 ( AMMONIUM SULPHATE)


21.Give the reaction between NITRIC acid and ammonia ( important)

NH3+ HNO3 ------- NH 4 NO3( AMMONIUM NITRATE)

22.What are the elements present in NPK FERTILIZER?………………………………………………………..

23.Describe the chemical test for water……………………………………………………………..

24.Describe the physical test for water………………………………………………………………


25.Describe the laboratory preparation of hydrogen

26.Name the gas produced when a metal react with dilute acid…………………………………………………….

27.Give the test for this gas…………………………………………………….

28.Give word equation for the reaction between hydrogen and oxygen…………………………………………………..

29.Give balanced equation for the reaction between hydrogen and

oxygen…………………………………………………..

30.Give one example for transition metal which is used as catalyst…………………………………………………….

31.Iron react with chlorine to form a yellow( brown) color solid
Name it…………………………….

Give equation for the reaction…………………………………………………..

Give the color of ion with sodium hydroxide

Fe2+ …………………Fe3+…………………Cu2+ ………………

32.Name the compounds……………………………………………………..Write down the action of iron over steam.

33.Give the balanced equation ( notes)…………………………………………………………………..

34.Write down the action of iron hydrogen chloride gas Give the balanced equation
ColorCopper oxide-redCopper hydroxide – blueCopper nitrate –greenCopper sulfate- blueCopper carbonate-blackCopper chloride- green


35.Write down the reaction of ammonia with copper (II) ions ( important)………………………………………………………………………………


36.Give the properties of transitional metals( answer given below)
1) variable valency2) colored compounds3) catalytic property


work sheet 2
1,Give example for Fe shows variable valency…………………………………………………………………….


What is reactivity series?


2.Give the order…………………………………………………………………………………


3.What do you know about the reactivity of and displacement reaction?………………………………………………………………………….


4.A copper dipped in silver nitrate solution .5.What is observation that you seeAfter 2hours……………………………………………………..

6.Iron nail dipped in copper sulphate ……………………………………


7.Silver nail dipped ii copper sulphate……………………………………


8.Copper nail in iron sulphate……………………………..(Important portion)Cat ionFlame colorLithium(Li+Sodium (Na+)Potassium( k+)Calcium( ca2+)test for gases oxygen------------------carbon dioxide----------------hydrogenchlorine-----------------------Sulpher dioxide----------------------ammonia--------------------------Test for anionsChloride(Cl-)-----------------------Bromide(Br ---------------------------Iodide (I-)-----------------------Sulphate(so42+)---------------------Sulphite( S2_)-----------------Carbonate(CO32_)------------------Test for cat ionAmmonia (NH4+)-----------------Copper (Cu2+)---------------Iron( Fe2+)-----------------

9.The pure iron get from the furnace has little use only?

10.The pure iron is rustingGive good methods to prevent rusting………………………………….

11.Give the chemical name of rust.………………………..

12.Why not cool air is blown from bottom.Temperature in side furnace get reduced

CHEMISRTY OF ELEMENTS / WORK SHEET / IGCSE / GCSE / HALOGENS /PERIODIC TABLE

Revision sheet

Chemistry of elements1.

In the periodic table the elements are arranged according to the ……………………………. Number

2. In the periodic table the elements are arranged according to the …………………………….

No of particle in the nucleus

3. Give any two property of metal………………………………………………………………………………….

4. Give any two properties of non metal……………………………………...

5. NAME THE FIRST GROUP ELEMENTS……………………………………………..

6. Name the second group elements…………………………………………………….

7. Name the seventh group elements………………………………….

8. Group 1 element are always has a …………………. Charge

9. Group 2 element are always has a …………………. Charge

9. Group 7 element are always has a …………………. Charge

10. Group 6 element are always has a …………………. Charge

11. Element A put in water to produce a gas BAnd an alkali C

12.It will give lilac color to the flameName the element A ……………………

13.Name the gas B ………………………………………….

14.NAME THE ALKALI C ………………………………………

15.Write three things that you see in this experiment

16.What is color change shown by the universal indicator in the solution?………………………………………….

17. Write down the word equation for the above reaction…………………………………………………..

18. Write down the balanced equation for the above reaction…………………………………………………..

19. What is the effect of blue litmus red litmus in the beaker…………………………………………

20. How will you test the gas evolved………………………………

21. Name a non metal for the same period as lithium ………………………..

22. Name another metal in the same period as lithium……………………………..

23. What is similar in the arrangement of electron in alkali metals?…………………………………………….

24. Why diamond and graphite have similar properties………………………………………….

25. Why diamond and graphite have different properties………………………………………….

26.why graphite conducts electricity but diamond not. Why?………………………………………………………………………………………..

27. Give the equation for the burning of graphite in air…………………………………………………………..

28.give the word equation for the burning of magnesium……………………………………………………..

29. give the( halogen) colour and Physical state

30. How will you convert the hydrogen chloride gas into hydrochloric acid

31. A solution in which the hydrochloric acid is dissolved is tested with blue litmus. What do you observe?

32. Hydro choric acid is dissolved in methyl benzene. What do the change in the blue litmus?

33. Describe the laboratory preparation of chlorine from hydro choric acid ( important)…
…………………………………………………………………………………………………..

34.Name the oxidizing agent in this reaction……………………………………

35. Give the test for chlorine…………………………………………………………….

36. Chlorine and bromine have similar properties why?……………………………………………………………………………….

37. Chlorine is passed through a solution of potassium bromide

Name the type of reaction………………………………………….

the solution became red in color. Why?…………………………………………………………………………..

38.Explain why no displacement reaction when we add bromine to potassium chloride …………………………………………………………………………………………
Write the ionic equation for the reaction……………………………………………….

39.Chlorine is bubbled through the potassium iodide solution.
What will be the color of theFinal solution ……………………………………………..

Write down the ionic half equation for the reaction …………………………………

40.What is the color change during the below reaction……………………………………………………H2 + Cl 2___________ 2HCl

reactivity series/ SALT ANAYSIS / IGCSE /GCSE WORK SHEET

work sheet 2


1.Give example for Fe shows variable
valency…………………………………………………

What is reactivity series?

2.Give the order…………………………………

3.What do you know about the reactivity of and displacement reaction?

………………………………………

4.A copper dipped in silver nitrate solution .

5.What is observation that you seeAfter 2hours…………………………

6.Iron nail dipped in copper sulphate ……………………………………

7.Silver nail dipped ii copper sulphate……………………………………

8.Copper nail in iron sulphate……………………………..
(Important portion)

Cat ionFlame color

Lithium(Li+Sodium (Na+)Potassium( k+)Calcium( ca2+)

test for gases oxygen------------------

carbon dioxide----------------

hydrogenchlorine-----------------------

Sulpher dioxide----------------------

ammonia--------------------------

Test for anionsChloride(Cl-)-----------------------

Bromide(Br ---------------------------

Iodide (I-)-----------------------

Sulphate(so42+)---------------------

Sulphite( S2_)-----------------

Carbonate(CO32_)------------------

Test for cat ionAmmonia (NH4+)----------

Copper (Cu2+)---------------

Iron( Fe2+)-----------------

9.The pure iron get from the furnace has little use only?

10.The pure iron is rustingGive good methods to prevent rusting……

11.Give the chemical name of rust.………………………..

12.Why not cool air is blown from bottom.Temperature in side furnace get reduced

IGCSE CHEMISTRY NOTES / WORK SHEETTS

Chemistry in society -EXTRACTION OF METALS ( work sheet)

1.Explain the extraction of iron ( draw and label the diagram)


2.Name the apparatus for the production of iron……………………………………………….

3.Give the raw material ( charge) for the production of iron……………………………………………………………………..

4.Give the reactions that take place in the blast furnace

5.Give the to reducing agents in the blastfurnace

6.Name the slag in theblastfurnace

7.Give the chemical formula of the slag

8.Name the main impurity present in the iron ore

9.Explain high temperature is produced in the furnace

11.What material is used for making the furnace? Why?

10.Give two ways the carbon dioxide is gas produced

11.Name the rock (ore) iron

12.Why hot air is blown for bottom

15.Name the iron produced in the blast furnace

16.What main impurity is present in the pig iron?

17.The impure iron produced in the blast furnace is not use full why?

18.Name the process by which we make steel

19.What are alloys

chemical calculation - IGCSE NOTES / BALANCING EQUATION

Word equations

If a reaction occurs between magnesium and oxygen, magnesium oxide is produced, here is the word equation for this reaction: - magnesium + oxygen --> magnesium oxide Some other examples are:

hydrochloric acid + calcium carbonate --> calcium chloride+ carbon dioxide+ water


sodium + water --> sodium hydroxide + hydrogen

hydrochloric acid + sodium hydroxide --> sodium chloride + water


Task

1.Write word equations for the reactions in which the following compounds form from a halogen and another suitable element: hydrogen fluoride, hydrogen chloride, iron III chloride, iron III bromide, sodium chloride, copper chloride.

2. State the name and the number of atoms of each element in the formulae above. Formulae can be worked out from valency.
Valency
atom or ion with that valency


1 hydrogen, group 1 e.g. sodium and potassium, group 7e.g. fluorine and chlorine, ammonium NH4+, hydroxide OH-, nitrate NO3-

2 group 2 e.g. magnesium and calcium, group 6, sulphate SO42-, carbonate CO32-, copper , II Cu2+ iron II Fe2+

3.group 3 e.g. aluminium, group 5, phosphate PO43-, iron III Fe3+

4 group 4 e.g. carbon

name aluminium oxidevalency 3 2symbols Al Oformula Al2O3

Task : Use valency to work out the formulae of the following compounds:

sodium chloride, potassium bromide, magnesium oxide, calcium sulphide, aluminium nitride, calcium iodide, lithium oxide, aluminium chloride, aluminium sulphide, magnesium nitride.


4.Calculating relative formula massAdd up the relative atomic mass (found in periodic table) of each atom in the compound.e.g.
Al203 relative atomic masses of Al = 27, O = 16 (found in periodic table). The formula shows 2 atoms of aluminium and 3 atoms of oxygen so:formula mass of = (2*27) + (3*16) =54 + 48 = 102

5. Work out the relative formula masses of the following: MgO, FeS, O2, H2O, CaBr2, Na2S, CaCO3, NaOH, HCl, (NH4)2SO4. Relative atomic masses Mg=24, O=16, Fe=56, S=32, Ca=40, Br=80, C=12, Na=23, H=1, Cl=35.5.



Simple balanced equationsIt is possible to write balanced equations for reactions.

For example substances such as hydrogen and magnesium combine with oxygen. One method to write them is:Write a word equation first.Magnesium + oxygen --> magnesium oxideWrite in the formulae of the substances used.

Mg + O2 --> MgO

Balance the equation so that each element has the same number of atoms on each side.

2Mg + O2 --> 2Mg

Osodium + oxygen --> sodium oxide

4Na{s} + O2{g} --> 2Na2O{s}

(word equation) hydrogen + oxygen ---> water(formulae)

H2 + O2 -----> H2O(balance) 2H2 + O2 -----> 2H2O



6.Balanced equations and ionic equations
Ionic equations only show ions which change in a reaction and ignore those which do not change. E.g.word equation

hydrochloric acid + sodium hydroxide --> sodium chloride + water

balanced chemical equationHCl(aq) + NaOH(aq) ---> NaCl(aq) + H2O(l)

ionic equation

H+(aq) + OH-(aq) ---> H2O(l)

E.g. in the electrolysis of sodium chlorideCl-(aq) --> Cl(g) + e-2Cl(g) --> Cl2

7 .Working out formulae from reacting masseselements reacting magnesium chlorinesymbols of elements Mg Clmasses reacting (from experiment)

8.1gmolar mass (look up relative atomic 24g/mol 35.5g/molmass in periodic table)amounts (amount = mass/molar mass) 2.4g/24g/mol 7.1g/35.5g/mol = 0.1mol 0.2molratio of atoms (divide by smallest) 1 : 2formula MgCl2


Task :Work out formulae of compounds formed when the following react:56g of iron and 32g of sulphur (Fe =56, S =32)2g of hydrogen and 16g of oxygen (H=1, O=16)14g of lithium and 16g of oxygen (Li=7)32g of copper and 8g of oxygen (Cu=64)6.4g of copper and 0.8g of oxygen.
C2.08

Calculating reacting masses using equationsYou can work out ratio of the masses of products and reactants by simply multiplying the number of moles shown in the equation by the formula mass of each substance.
Example 1: What mass of magnesium oxide can be made from 12g of magnesium? Relative atomic masses are Mg =24, O = 16. equation 2Mg(s) + O2(g) --> 2MgO(s) formula 2*24 1(16*2) 2(24+16)masses =48 =32 =80 reacting 48g of Mg forms 80g of MgOmasses 1g of Mg forms 80/48 g of MgO 12g of Mg forms 12*80/48 g of MgO = 20g


Example 2: What mass of magnesium oxide can be made from 12g of magnesium? equation 2Mg(s) + O2(g) --> 2MgO(s) amounts 2 moles 1 mole 2 moles masses 2*24 1{16*2} 2{24+16} =48g =32g =80g so 48g Mg forms 80g MgO 1g Mg forms 80/48 g MgO 12g Mg forms 12*80/48 g MgO = 20g Also note that the ratio of amounts of reactants and products in the equation above can be written as:
Amount of Mg/amount of O2 =2/1 Or Amount of O2/amount of MgO = 1/2



You decide to travel from London to Delhi for a holiday and hire your own Airbus 319 jet. You fly the 2183 miles to Cairo first and make 9.4 tons of carbon dioxide. After seeing the Pryramids you fly the 1018 miles to Riyadh making 4.5 tons of carbon dioxide. After a brief stop in the Saudi Capital you fly on the 1900 miles to Delhi making 8.1 tons of carbon dioxide. Finally you return to London travelling 4171 miles and releasing 22 tons of carbon dioxide. The jet burns the fuel kerosine C15H32 in the reactionC15H32 + 23O2 --> 15CO2 +16H2O(a) How much carbon dioxide was released into the atmosphere during the trip?(b) What mass of kerosine was burnt?(c) What effect does the trip have of the environment?

PERIODIC TABLE-IGCSE /GCSE NOTES

PERIODIC TABLE



The elements in the periodic table are arranged according ti the increasing atomic number.
In the peridoic table there are eight main groups and seven periods. The elements in the same groups have similar properties because they have the same number of electrons in their outer shell.
The group number of the elements is the same as the number of the outer electrons. 3/4 of the periodic table is filled with metals.
There are other elements called the semi metals whos hare both the properties of metals and non metals.They conduct like metals and are brittle like non-metals.

GROUP 1

They are called alkali metals because they react with water to form strong alkali.

The increase in number of shells and atomic size causes increase ub the reactivity down the group because bigger the atom, the less its attractive force towards the nucleus.
Group 1 metals have one electron in the outermost shell which is why they are called group 1.
They are less dense than group 2 metals because they have only one electron in the outermost shell and they're stored under oil to prevent them from reacting iwth oxygen and water vapour form at the atnoshpere as it is dangerous.


Reaction with water

Sodium is a strong alkali, when it reacts with water :

1) It floats and melts.
2) Hydrogen gas is produced.
3) It forms sodium hydroxide.

Sodium + water ---------------> Sodium hydroxide + hydrogen
Na + H2O ---------------------> NaOH + H2
Na + HOH -------------> NaOH + H2

Potassium when reacting with water :

1) It floats.
2) It catches fire and burns with a lilac flame.
3) It forms potassium hydroxide and hydrogen.

Potassium + water --------------------> Potassium hydroxide + hydrogen
K + H2O --------------------------------> KOH + H2
K + HOH -------------------------------> KOH + H2



PERIODIC TABLE - IGCSE /GCSE CHEMISTRY REVISION NOTES


Lithium forms Lithium hydroxide and hydrogen gas when reacting with water.

Li + H2O ----------------> LiOH + H2

Rubedium (Rb)

Rb + H2O -----------> Rb (OH) + H2

When rubedium reacts with water :

1) Heat is produced.
2) Explosion takes place.
3) It forms Rubedium hydroxide and hydrogen gas.

Group 2/ halaogens /IGCSE chemistry notes

group 2


* Group 2 elements are known as alkaline earth metals.

* They can form strong alkali and they occur widely on earth crust.

* They are metals.

* They have melting points and densities that are quite low for metals.

* They react with air, water or steam to form oxides and hydroxides.

* They show the same reactivity trend as group 1, increasingly reactive as you go down the group.

* They form compounds in which the metal has a +2 ion.

* They form white or colorless compounds which gives colorless solutions.

* The alkaline earth metals are reactive, but less reactive than the alkali metals.



HALOGENS

* Each of the halogens gives colored vapour.

* Halogens react with hydrogen to form hydrogen halides.

* Halogens react with metals to form metal halides such as iron chloride and sodium chloride, ordinary salt.

* Halogens are oxidizing agents and are themselves reduced.

* Any halogen can displace another that is lower down in group 7.

* We can use ionic equations to show how halogens react.

* Halogens compounds have useful properties.

* They have diatomic molecules X2 .

* They go from gases to liquid to solid as you go down the group.

* They become less reactive towards the bottom of the group.

WORK SHEET - TEST FOR GASES

1. Suggest a method to collect the following gases. Give reason

Chlorine, hydrogen. Carbondixide , ammonia

2. complete the table

Gas ----test ------Result
hydrogen
oxygen
carbondioxide
chlorine
Sulphur di oxide



3. Draw the apparatus used to measure and collect gases





1. Suggest a method to collect the following gases. Give reason

Chlorine, hydrogen. Carbondixide , ammonia

igcse chemisyrty notes -extraction of metals

Oxidation and reduction

Oxidation is a reaction where oxygen is gained.

eg iron + oxygen ---> iron oxide

Reduction is a reaction where oxygen is lost. Eg iron oxide is reduced to form iron.

Classify the following reactions as oxidation or reduction:

burning magnesium,

making aluminium from aluminium oxide,

burning carbon in air,

copper oxide + carbon --> copper + carbon dioxide

carbon monoxide forming carbon dioxide.

Electron loss and gain

Oxidation is also a loss of electrons

eg Fe --> Fe2+ + 2e-

Reduction also the gain of electrons eg (when Fe ions becomes Fe atoms

Identify all of the elements below as being oxidised or reduced in the following reactions:

Cu2+ + 2e- ---> Cu

lead ions forming lead atoms

Na ---> Na+ + e-
Ca + Cl2 --> CaCl2
Fe2+ ---> Fe3+ + e-
Reduction of metal ores
An ore is a material found in the ground which contains a metal. An ore is often a metal oxide mixed with rock . When a metal is extracted its ore is reduced. Metal loses oxygen from its oxide. Examples include: haematite which is mostly iron III oxide
bauxite which is mostly aluminium oxide Al2O3
malachite which is mostly copper carbonate CuCO3
Extraction and position in reactivity series
Uses and properties of aluminium
extraction method that uses a lot of electrical energy. Electrical energy is more expensive than energy from burning carbon
The extraction of aluminiumAluminium is found in the ground in an ore called bauxite. Bauxite is aluminium oxide (Al2O3) with iron oxide impurities.


After purification aluminium oxide is mixed with cryolite to lower the melting point from 2000º to 1000º, which saves money.
This mixture is heated and the molten liquid used as the electrolyte. Both electrodes are made of graphite (carbon). The anode (+ve) is graphite and the cathode (-ve) is a graphite lining to a steel case.
The carbon anodes react with oxygen so have to be replaced.
C + O2 --> CO2

At cathode - positive aluminium ions attracted, gain electrons and become atoms.

Al3+ + 3e- ---> Al

At anode - negative oxide ions attracted, lose electrons and become atoms.2O2- ----> O2 + 4e-
uses and properties of aluminium
overhead power cables-good electrical conductor, low density
drinks cans-Does not react with water
aircraft parts -high strength and low density
Carbon and carbon monoxide for reducing oxides
Carbon and carbon monoxide can both remove oxygen from other compounds so are good for reduction.
They are used to reduce the ores of metals below carbon in the reactivity series.
E.g. zinc, iron, tin and lead.
Iron extraction using the blast furnace Drag and drop labels test on blast furnace
A blast furnace is used in the process of extracting iron. The raw materials iron ore, coke and limestone are put in at the top. Hot air is blasted into this furnace at the bottom making the coke (carbon) burn faster and the temperature rises to about 1500º.
When the coke burns, carbon dioxide is produced. C + O2 ---> CO2 CO2 reacts with the unburnt coke to form carbon monoxide CO CO2 + C ---> 2CO Iron oxide Fe2O3 in the ore is reduced to iron by the reaction with the carbon monoxide. 3CO + Fe2O3 ---> 3CO2 + 2Fe Molten iron is a dense liquid, so runs to the bottom of the furnace and is tapped off.
Limestone CaCO3 helps remove impurities during the extraction by forming calcium oxide

CaO. CaCO3 ---> CaO + CO2



The rock impurities silicon dioxide SiO2 are then removed by the following reaction.

CaO + SiO2 ---> CaSiO3 CaSiO3
is known as slag and can be used in making cement and road building.

The purification of copperVery pure copper is needed for copper wires. Electrolysis is needed to purify copper. The anode is a mass of impure copper and the cathode is pure copper. The electrolyte is sulphuric acid. The impurities drop at the anode as sludge during electrolysis. At anode Cu ---> Cu2+ + 2e-





At cathode Cu2+ + 2e- ---> Cu

only for gcse chemistry -Chemical from calcium carbonate

Thermal decomposition of calcium carbonatecalcium carbonate ---heat---> calcium oxide + carbon dioxideCaCO3(s) ---heat---> CaO(s) +CO2(g)(limestone) (quicklime)

Water and calcium oxidecalcium oxide + water ---> calcium hydroxideCaO(s) + H2O(l) ---> Ca(OH)2(s) (slaked lime) looks flaky and gets hotWith more water the calcium hydroxide dissolves to make limewater Ca(OH)2(aq)
Neutralising soil acidityCalcium oxide and calcium hydroxide are bases. (like all metal oxides and hydroxides)acid + base ---> salt + waterso acid soil is neutralised by calcium oxide or calcium hydroxidee.g. nitric acid + calcium oxide --> calcium nitrate + water

2HNO3(aq) + CaO(s) ---> Ca(NO3)2(aq) +H2O(l)Farmers can add calcium oxide or calcium hydroxide to their acid soil to neutralise it.The neutral soil is better for growing plants.Task C3.25 Write word equations and balanced chemical equations for neutralising the following acids using calcium oxide and calcium hydroxide: nitric acid, sulphuric acid, hydrochloric acid, phosphoric acid (H3PO4)

Uses of calcium carbonateCement is made from calcium carbonate (limestone).limestone --heat--> calcium oxide + carbon dioxide


Calcium oxide + clay ----> cementGlass is made using calcium carbonate.


Calcium carbonate + sand + sodium carbonate ---heat--> glassIron is manufactured using calcium carbonate in the blast furnace.calcium carbonate --heat--> calcium oxide + carbon dioxideCalcium oxide + silicon dioxide ----> calcium silicate (mostly sand) (slag) (floats on top of molten iron)

igcse/ gcse chemistry notes- Chemicals from salt

The electrolysis of aqueous sodium chloride solutionWhen the electrolysis of aqueous sodium chloride takes place, hydrogen and chlorine are given off as gases and sodium hydroxide is left. Aqueous sodium chloride contains hydrogen ions H+ and hydroxide ions OH- (from the water) and sodium ions Na+ and chloride ions Cl-. The positive sodium and hydrogen ions go to the cathode and the negative chloride and hydroxide ions go to the anode. Hydrogen is formed at the cathode and chlorine is formed at the anode.


substance-hydrogen
test -lighted splint
result-squeaky pop

substance-chlorine
test-damp blue litmus
result-turns red then white

substance -sodium hydroxide
test-damp red litmus
result-turns blue/purple

Write a word and a balanced chemical equation for the reaction in which hydrogen burns.


Complete: Chlorine is (acidic/alkaline) because it turns litmus red. Chlorine is a (dye/bleach) because it turns litmus white. This type of reaction is (oxidation/reduction).

Uses of sodium chloride, hydrogen, chlorine and sodium hydroxide
substance


uses
sodium chloride
preserve food, flavouring, stops ice on roads because salt lowers freezing point of water to -5oC, manufacture of sodium, chlorine and sodium hydroxide


chlorine
kills bacteria in swimming pools and drinking water, bleach, making hydrochloric acid and solvents.



hydrogen
rocket fuel, making margarine, making ammonia
sodium hydroxide
detergents, bleach, paper, fibres, purifying bauxite


Match the uses above to the explanations below: compounds of carbon and chlorine are good at dissolving grease, hydrogen reacts with chlorine to form hydrogen chloride, vegetable oil reacts with hydrogen to form margarine, nitrogen and hydrogen react together in the Haber process, soap is made from vegetable oil and sodium hydroxide, the structure of wood is broken up by sodium hydroxide, salt is easily detected by the tongue, bacteria cannot survive in salty solutions, by the electrolysis of molten sodium chloride, by the electrolysis of aqueous sodium chloride.

igcse chemistry free notes -Transition metals

The colours of transition metal compoundsAll compounds of transition metals are coloured. Some examples are:copper sulphate - blue, iron III oxide - brown, cobalt chloride - pink.

Transition metals are used as catalysts
manganese (IV) oxide
decomposing hydrogen peroxide

platinum
catalytic converter in car exhaust

nickel
making margarine from vegetable oil and hydrogen

vanadium (V) oxide
making sulphuric acid

platinum with rhodium
making nitric acid


iron
making ammonia


Uses and properties of titanium, iron and copper

iron
Bridge Construction
High Strength

iron
permanent magnets
Magnetic

copper
water pipes
does not react with water

copper
Electrical wires
Good conductor of electricity

titanium
propeller shafts on ships, hip joints
does not react with water

titanium
aircraft engines
high strength, low density, high mp

igcse free chemistry notes -Halogens

Physical properties of halogens

The colour of the Halogens changes from a lighter colour to a dark colour as we go down the group and the melting and boiling points increase as you go down the group.

Fluorine at the top of the group is a yellow gas This changes to a yellow-green gas for chlorine, a red liquid for bromine and finally black solid for iodine at the bottom of the group.

Reactions of halogens with metalsMetal + Halogen --> Metal Halide

Iron reacts slowly with iodine to form iron II iodide. Iron reacts faster with bromine to form iron III bromide.

Iron reacts with chlorine reacts even faster to form iron III chloride.
Other metals like sodium also react.sodium + chlorine ---> sodium chloride

The reaction of halogens with hydrogenHydrogen reacts with halogens to form hydrogen halides. E.g.Hydrogen + chlorine ---> hydrogen chloride

H2(g) +Cl2(g) --> 2HCl(g) (a hydrogen halide)

Hydrogen halides are very soluble in water.

The gas hydrogen chloride forms hydrochloric acid when it dissolves in water. Other hydrogen halides like hydrogen bromide, HBr and hydrogen iodide, HI also dissolve to form acid solutions. An acid solution has a pH less than 7


e.g. 1. An acid turns universal indicator red.

Displacement reactions

Halogens like chlorine are very reactive and displace less reactive halogens like bromine from halides like bromides.

Potassium bromide + chlorine ---> potassium chloride + bromine

2KBr(aq) + Cl2(aq) --> 2KCl(aq) + Br2(aq)

2Br-(aq) + Cl2(aq) --> 2Cl-(aq) + Br2(aq)

Chlorine and bromine also react with potassium iodide showing that the order of reactivity is: Chlorine > bromine > iodine


Uses of halogens and halidesFluorine compounds (fluorides) are put into toothpaste and some drinking water supplies. Fluorides join with tooth enamel and make teeth resist attack by acid which prevents tooth decay. Chlorine is used to in swimming pools and drinking water to kill bacteria. Iodine is used as an antiseptic because it will kill the germs on the skin without damaging it.

IGCSE CHEMISTRY FREE NOTES Alkali metals

Properties of alkali and other metalsThe alkali metals have low melting points low densities and are soft compared to other metals.

For example lithium has a melting point of 180ºC and caesium has a melting point of just 29ºC.

These are low compared to other typical metals such as iron, melting point 1500ºC.

Sodium and potassium are soft and can be cut with a knife but iron is hard.

Alkali metals and waterWhen an alkali metal is put into water, it reacts very vigorously. It moves around the surface fizzing.

Hydrogen is produced which can be tested using a lighted splint. The metal also gets hot and potassium gets hot enough to ignite in its own heat.

The water turns into a hydroxide. Reactivity increases as you go down the group.

2Na(s) + 2H2O(l) ---- 2NaOH (aq) + H2(g

Reactivity of metal with water

lithium
fizzes gently moving slowly
reactive


sodium
fizzes vigorously, moves quickly
very reactive


potassium
violent fizzing, bursts into lilac flame
violently reactive

Properties of alkali metal compoundsAlkali metal compounds are all white solids which are soluble in water.
Alkali metal hydroxides or oxides dissolve in water and are alkalis. An alkali has a pH of over 7, e.g. 14. An alkali turns universal indicator purple. An alkali can neutralise an acid. Whenever an acid is neutralised it produces water and a salt.
sodium hydroxide + hydrochloric acid ---> sodium chloride + wateR (a salt)


NaOH(aq) + HCl(aq) ---> NaCl(aq) + H2O(l)

igcse chemistry free notes-Noble gases

Monatomic nature of noble gasesNoble gases are monatomic. This means they exist only as single atoms. Their atoms cannot combine with other atoms because in their electronic structure, the outer shell is always full and this makes noble gases unreactive


Argon
Light bulbs
Doesn’t react with the metal filament


Helium
Used with O2 for deep sea dives
Low solubility of helium in the blood.

Helium
To inflate the tyres of large aircraft
Non-flammable

Helium
To fill airships and weather balloons
Low density, does not burn
Uses and properties of noble gases


Neon
In advertising signs because it glows red when electricity passes
Conductor of electricity at high voltage

Krypton/Xenon
In lamps used in photographic flash units, in stroboscopic lamps used in lighthouses
Gives out a lot of light when electricity passes through

igcsechemistry free -notes The periodic table

Elements are placed in order of increasing atomic number so hydrogen atomic number 1 is firsthelium atomic number 2 is second etc.

Groups

A group is a column of elements with similar properties in the periodic table. See figure1. e.g all of the following are in the same group and are silver in colour, soft and good conductors of electricity.
lithium
sodium
potassium
rubidium
caesium

The position of groups in the periodic tableThe positions of the alkali metals (group 1), the halogens (group 7), the noble gases (group 0) and the transition metals are shown in figure 1 above.

Task :Label a copy of the periodic table to show a group, a period, the alkali metals, halogens, noble gases, transition metals, metals, and non-metals.

Reactions and electron arrangements

Group 1 elements all have atoms with 1 outer electron. e.g. sodium 2,8,1All of these atoms try to lose 1 electron so all react in the same way.

They all form a positive ion and form ionic compounds. E.g. sodium Na form sodium ion Na+.Group 1 elements all react with oxygen (that is burn) to form oxides. The products like sodium oxide, are all ionic.

Group 0 elements all have atoms with complete outer shells

e.g. helium 2, neon 2,8. None of these elements react with anything.

Task - Name other elements in group 1 and say how they react with oxygen giving a reason. State how argon reacts with oxygen and give a reason.

Trends in group propertiesGroup 1 elements get more reactive with water as you go down the group. e.g.Lithium (gentle fizz) is less reactive than potassium (violent lilac flame).

Group 7 elements get more reactive with iron as you go up the group.Iodine darkens iron when it is heated but chlorine makes iron burst into flames.

igcse free - chemistry notes / Metals and Non-Metal

Properties of Metals
malleable
ductile
sonorous
shiny
conduct heat
conduct electricity
hard and strong
Non Metals
brittle
dull
do not conduct electricity
do not conduct heat
not magnetic
Alkali Metals
very reaction (have to be stored under oil to prevent reaction with air)
form ions with 1+ charve.
soft and can be cut with a knife
potassium and sodium float on water
Trends down the group
atomic size increases (increased no of electrons and shells)
lose outer shell electron to form 1+ charged ions. Therefore reactivity increases as you go down the group. The electron is farther away from the nucleus.
melting points decrease
Reactions:
alkali metals with oxygen in the air
Sodium + Oxygen = Sodium Oxide
4Na + O2 -> 2NaO2
alkali metals with water
gives off hydrogen gas
resulting solution is alkali
Sodium + Water -> Sodium Hydroxide + Hydrogen
2Na + 2H20 -> 2NaOH + H2
alkali metals with halogens
they react rapidly with halogens
Sodium + Chlorine -> Sodium chloride
2Na + Cl2 -> 2NaCl
Halogens
non metals
form ions with 1- charge
very reactive
toxic
as you go down the group the colour of vapour becomes darker
fluorine - yellow
chlorine - green
bromine - red
iodine - violet
Halogens are used as bleaching agents
Trends as you go down the group
atomic size inicreases
reactivity decreases (easier to gain electrons when electron shell is closer to the nucleus)
melting and bioling point increase
Reactions
halogens with hydrogen
Chlorine + Hydrogen -> hydrogen chloride (hydrochloric acid)
Cl2 + H2 -> 2HCl
halogens with metals (forms a salt)
Magesium + Bromine -> Magnesium bromide
2Mg + Br2 -> 2MgBr
Displacement Reactions with Halogens
A more reactive halogen will displace a less reactive one.
Example: chlorine can displace bromine from potassium bromide, but cannot displace fluorine from potassium fluoride.
Cl2 + 2KBr -> 2KCl + Br2 (chlorine displaced bromine)
Br2 + 2KI -> 2KBr + I2 (bromine displaced iodine)
Br2 + 2KF -> 2KF + Br2 (bromine can't displace fluorine)
Noble Gasses
gasses at room temperature
uncreactive
density with increase in atomic size
Argon is used inside light bulbs
Helium is used in hot air ballons
boiling point increases as you go down the group
Transition Metals
hard dense metals
not very reactive
often used as catalysts
have variable valency (different charges, Fe2+ Fe3+)
Some metals form oxides in reaction to oxygen
some metals form basic oxides
some non-metals form acidic oxides
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igcse - free- chemistry notes Atoms, Bonding and the Periodic Table

The group number is related to the number of electrons in the outer shell. e.g. (Mg) Magnesium, in group 'II', has two electrons in the outer shell
electrons are arranged in shells surrounding the nucleus, with the first shell containing 2 electrons and subsequently 8.

All atoms want a full outer shell of 8 electrons (like noble gasses)


Ionic Bonding:


Ions are formed by gaining or loosing electrons
is the transfer of electrons from one atom to another, giving the atom a positive of negative charge

The two atoms together have a neutral charge
It occurs mainly when metals bond with non-metals

Examples

Magnesium + Oxygen
Electron Configuration
Mg: 2 --> transfer two electrons from Mg to give both atoms a stable


electron configuration
O: 2,6

Properties of Ionic Compounds
solids at room temperature (strong bonds, opposite charges)
high melting and boiling point (lots of energy required to overcome strong bonds)
don't conduct electrivity in the solid state, since electrons can't move (can in the molten state)
form crystaline solids because of ionic lattice
dissolve readily in water, but not in molecular solvents
Covalent Bonding
occurs when non-metals bond.
They share electrons to get a noble gas structure
Example: Hydrochloric Acid (HCl)
both H and Cl need one electron.
They form a single covalent bond to get a full outer shell
Properties of Covalent Compounds
molecules have no charge
low melting and boiling points. (little attraction between molecules,)
dont attract electricity (don't have charged particles)
liquids of gasses at room temperature
dissolve readily in molecular solvents, but not in polar solvents like water
Isotopes
are atoms of the sama element with:
same number of protons
same number of electrons
different number of neutrons
Examples
126C (6 electrons, 6 protons, 6 neutrons[12-6])
and

146C (6 electrons, 6 protons, 8 neutrons [14-6])

note: top number is the atomic mass

Relative Atomic Mass

average mass of atombased on various isotopes present

Noble Gasses

unreative (inert)because of full outer shell of electrons

igcse chemistry notes -properties of metals and non metals

metals -good conductors of heat and eletricity
non metals-poor conductors of heat and electricity
metals-shiny
non metals -dull
metals-malleable
non metals -brittle
metals -strong
nonmetals -weak
metals -react with oxygen to form basic oxides
nonmetals-react with oxygen to form acidic oxides
metals-usually high density
nonmetals-usually low density
metals -usually high melting points except alkali
nonmetals-usually low melting points except carbon
metals-many react with dilute acid to produce hydrogen
nonmetals -no reaction with dilute acid

igcse chemistry notes -compounds and mixtures

Mixture-No heat or light is given out or absorbed when mixing occurs
compound-Heat and sometimes light is usually given out or absorbed when a compound is made.
mixture-The substance in a mixture can be separated by physical means
compund -the elements in a compound cannot be separated by physical means
mixtures -the properties of a mixture are the average of the properties of thes substances in it
compounds -the properties of a compound are quite different from those of the elements in it
mixtures -the substances in a mixture can be presented in any proportions by mass
the elements in a compound are combined together in definite proportions by mass (we have not proved this yet).

igcse/gcse / o level chemistry notes -Elements and Compounds

Atoms, Elements and Compounds
The Atom
? Smallest particle that can take place in a chemical reaction
? Consists of three sub-atomic particles, electrons, protons, neutrons
Electrons
? Negatively charged, found in energy shells outside the central part of the atom
? Move at very high speeds in orbit
? Have very negligible mass
? Mass is 0.0005 of a proton
Protons
? Are positively charged and found in the nucleus of an atom
? Have a unit mass of one (ie. 1 a.m.u)
Neutrons
? Have no charge at all and are located in the nucleus
? Have a unit mass each
1 atomic mass unit (1 a.m.u) = 1.67 x 10-27 kg)
The Atomic Number is also called the proton number (distinguishes element)
The proton number equals the electron number so that charges balance
The mass number is the total mass of the atom. Calculated with proton + neutron. (Electron neglected)
Nucleon number: total number of particles inside the nucleus (ie. Protons + neutrons)
Energy Shells
? Shells increase as they become full
? The first shell only contains 2 electrons and following shells contain 8 each
Chemical Stability
? Electrons are spaced out regularly at intervals as shown below. The presence of unpaired electrons determines the chemical reactivity of the atom.
? The electrons in the outer shell of an atom are referred to as valence electrons
? If the valence electrons for the outer shell are not enough to fill the shell, then the atom is chemically reactive and unstable. (E.g. Sodium Na+)
? Atoms with fully filled outer shells such as noble gases, are chemically unreactive and stable.
Noble Gases
? Also called inert gases have fully filled outer shelled atoms (E.g. Helium, Neon, Argon, Krypton)
Electronic Configuration
? A group of numbers which show the arrangement of electrons in their various shells
? The numbers represent number of electrons in each shell starting with the innermost shell, separating each number (shell) by a comma
E.g. He: 2, Ne: 2,8, O: 2,6
Isotopes
? Atoms of the same element having different mass #s but the same atomic number (ie. Different neutrons)
? Most elements exist as a mixture of their isotopes
? Hydrogen has 3 isotopes: Protium (0n), Deuterium (1n) and Tritium (2n).
These are called H-1, H-2 and H-3
The two types of isotopes are: (i) Radioactive and (ii) Non-radioactive isotopes
? Radioactive isotopes are isotopes with an unstable nucleus (the general term for radioactive substances since they are all isotopes).
? Non-Radioactive isotopes are stable and unreactive
Medical uses:
? Isotope cobalt-60 emits gamma radiation which can be used to sterilize medical equipment
Industrial Uses:
? Uranium-238 can be used to estimate the ages of rocks
Differences between Elements, Mixtures and Compounds
? An Element is a pure and simple substance which cannot be broken down into any simpler substances other than itself by any ordinary chemical means.

? A Mixture is made up of two or more substances physically mixed together.

The components in a mixture can always be separated by physical means

A compound is a pure substance which is made up of two or more elements chemically combined.
When these combine chemically they lose their identities and the compound takes its own properties.

igcse chemistry notes - atomic structure

The structure of an atom
Particle inside the atom
proton
Neutron
Electron
Atomic number, mass number and relative atomic mass

Atomic number is equal to the number of protons or electrons in an atom.

Mass number is equal to the number of neutrons plus the number of protons in an atom.

Relative atomic mass is the average mass of atom of an element compared to a standard.

eg
12
C
6


Electronic structures of the first twenty elements in the periodic table
Put electrons in shells so that the first is filled before the second etc and
the first shell has a maximum of 2 electrons
the second shell has a maximum of 8 electrons
and the third shell has a maximum of 8 electrons



Electronic structure of elements
1 Hydrogen 1
11 Sodium 2,8,1

2 Helium 2
12 Magnesium 2,8,2
3 Lithium 2,1
13 Aluminium 2,8,3
4 Beryllium 2,2
14 Silicon 2,8,4
5 Boron 2,3
15 Phosphorous 2,8,5
6 Carbon 2,4
16 Sulphur 2,8,6
7 Nitrogen 2,5
17 Chlorine 2,8,7
8 Oxygen 2,6
18 Argon 2,8,8
9 Fluorine 2,7

19 Potassium 2,8,8,1

10 Neon 2,8

20 Calcium 2,8,8,2
make up your own atom





Isotopes

Isotopes are atoms of the same element which have different numbers of neutrons in their nuclei. As these are the same element the atoms all have the same number of protons. For example hydrogen has 3 isotopes. Each atom has 1 proton but a different number of neutrons.
Relative atomic mass of an element
Use relative mass of isotopes and their relative abundance. E.g. Chlorine has two
isotopes with mass numbers 35 and 37.

35 37

75% is Cl,25% is Cl

17 17
Let there be 100 atoms
Total mass of 100 atoms = (75 * 35) + (25 * 37) = 3550
Average mass of an atom (relative atomic mass of chlorine) = Total mass /Number of atoms =3550/100
so relative atomic mass of chlorine = 35.5
home work
draw the electron arrangement of
calcium
lithium
carbon
chlorine