chemistry

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Understand the arrangement, movement and energy of the particles in each of the three states of matter: solid, liquid and gas

Solid

Have a fixed shape

Particles are vibrating on the spot

Liquid

Can flow as the particles can move around one another

Takes the shape of the container it is in

Gas

Move rapidly and are independent of one another, colliding with each other and with the walls of the container

Diffuse rapidly and exert pressure on the objects they collide with

Is much less dense than either the solid or the liquid

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Describe how the inter-conversion of solids, liquids and gases are achieved and recall the names used for these inter-conversions

Evaporation: conversion from liquid to gas at room temperature

Boiling: conversion from liquid to gas at boiling point

Sublimation: conversion of a solid to a gas or vice versa without passing through the liquid phrase

E.g. iodine, naphthalene and solid carbon dioxide (dry ice)

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Describe the changes in arrangement, movement and energy of particles during these inter-conversions.

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Describe simple experiments leading to the idea of the small size of particles and their movement including:

i. Dilution of coloured solutions

Dilution of CuSO4 solution

ii. Diffusion experiments

Diffusion is the movement of particles from a region of high concentration to a region of low concentration.

Factors that affect diffusion:

Temperature

o In hot water, rate of diffusion is faster as the molecules gain more kinetic energy

Density

o Lighter particles travel faster than heavier particles

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Understand the terms atom and molecule

A molecule

Is a particle of matter composed of two or more atoms held together in a particular arrangement by strong

chemical bonds

Molecules have a neutral electrical charge that is generally stable. Examples of molecules include molecules of

water (H O) and oxygen (O )

An atom

Is the smallest particle of an element that has all the properties of that element and is a fundamental piece of

matter; it is made up of three fundamental particles (neutrons, protons and electrons).

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Understand the differences between elements, compounds and mixtures

An element

Is a substance that is made entirely from one type of atom

For example, the element hydrogen is made from atoms containing a single proton and a single electron. If you

change the number of protons an atom has, you change the type of element it is

A mixture

Is a substance made by combining two or more different materials in such a way that no chemical reaction

occurs

A mixture can usually be separated back into its original components

A chemical compound

Is a substance composed of two or more different elements chemically bonded together in a fixed proportion by

mass When a compound is formed from its components, a chemical change takes place through chemical reactions.

Elements form compounds to become more stable, which happens when the maximum numbers of possible

electrons are in the outermost energy level (normally two or eight valence electrons)

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Describe techniques for the separation of mixtures, including simple distillation, fractional distillation, filtration, crystallisation and paper chromatography

Decantation:

o Quick method used to separate a mixture of a liquid and a heavier solid

o Allows the solid to sink and settle before pouring out the liquid

o Cannot be used with lighter solids

o Example:

Separating sand and water

Filtration:

o Method used to separate suspensions

o Mixture is poured into a funnel fitted with a piece of filter paper

o Tiny holes in filter paper allow liquid to pass through but solid particles are too large to do so

Residue:

Solid particles that stays on the paper

Filtrate:

Liquid which passes through

o Example:

Separating mud and water

Centrifugation:

o Is used when we want to separate small amounts of suspension

o The suspension of solid in liquid is poured into a centrifuge tube and is spun around very fast in a

centrifuge

o The spinning motion forces the solid to the bottom of the tube

o The liquid can then be poured off from the solid

o Example:

Separating cream from milk to make skimmed milk

Evaporation:

o Is used to separate solutions

o The solution is heated so that the solvent evaporates, and the solid is left behind

o Example:

Obtaining salt from salt water

Crystallization:

o Used to separate dissolved solids from a solution

o By cooling down a hot concentrated solution:

In a heated solvent, more solute can be dissolved than in a cool solvent

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The solution has to be heated to get rid of some water

This makes the solution concentrated

When the solution cools, the solvent can no longer hold as much solute

The excess solute will be separated out as crystals

o Slow evaporation of solution at room temperature:

At room temperature, the solvent will still evaporate

As more solvent evaporates, the solution becomes more concentrated

After the solution is saturated, excess solutes will form crystals

The longer the crystals take to form, the larger they will be as solute particles require time to

arrange themselves in regular shaper in order to form crystals

If dust is exposed to the solution, the crystals will be smaller

o Example:

Obtaining sugar from sugar solution

Distillation:

o Is used to obtain the liquid from a solution after evaporation

o Condenses the hot vapour formed during evaporation by using:

A cold surface

A condenser

This condenses steam more efficiently

Consists of two tubes

o One inside another

o The outer tube contains cool water

o The inner tube contains steam

o The steam can condense easily in the inner tube

o Evaporation + Condensation = Distillation

o Example:

Obtaining water from salty water

Using a separating funnel:

o Is used to separate immiscible liquids

o The liquids form layers

The liquid with the higher density will form the lower layer

o The liquid with the higher density can be separated by removing the stopper and opening the tap.

The lower layer will run through the tap

o Example:

Separate oil and water

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Fractional Distillation

o Is used to separate miscible liquids

o For the separation to work, the liquids in the mixture should have a difference of 10°C in their boiling

points

o The mixture will be heated

o The liquid with the lower boiling point will evaporate, rise up the fractionating column and enter the

condenser.

o The gas will condense and become liquid again in the second beaker

o Example:

Separate ethanol and water (ethanol has a lower boiling point)

Sublimation:

o Used to separate a solid from a solid where one sublimes while the other does not

o The mixture of the two solids is heated

o Only one of the solids will change to vapour

o Example:

Separating iodine from sand (iodine sublimes)

Chromatography:

o Used to identify substances in a mixture

o A mixture is put on a strip of paper, one centimetre away from one of the shorter edge

o The edge with the ink is then dipped in water without putting the ink into the water

o The water will then travel along the strip of paper carried the mixture which will then split into the

different substances

o The substances do not necessarily have to be coloured

Colourless substances can be made to show up by spraying the paper with a locating agent,

which then reacts with each of the colourless substances in order to produce a coloured product

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o The movement of each substance in the mixture depends on:

The solubility of the substance in the solvent

The substance moves with the solvent easily if the substance is very soluble in the

solvent

The absorption of the substance on the filter paper

Some solids are able to attract other substances strongly and hold them on their

surfaces

o This is called adsorption

The substances will not move with the solvent easily if the substance in the mixture is

absorbed strongly by the filter paper

o We call the solids which are able to attract other substances strongly and hold them on their surface

adsorbents

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Recall that atoms consist of a central nucleus, composed of protons and neutrons, surrounded by electrons, orbiting in shells

__________________________________________________________________________________________________

Recall the relative mass and relative charge of a proton, neutron and electron

Particle Mass Charge

Proton 1 +1

Neutron 1 0

Electron Almost 0 (1/1848) -‐1

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Shell/Orbit

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Understand the terms atomic number, mass number, isotopes and relative atomic mass (Ar)

The atomic number

Is the number of protons in the atom, also called proton number; it is the smaller of the two numbers shown in

most periodic tables

The mass number or atomic mass

Is the number of protons and neutrons in the nucleus of an atom; it is the bigger number of the two numbers

shown in most periodic tables

Isotopes

Atoms of the same element which have the same number of protons and electrons but a different number of

neutrons; therefore they have the same atomic number but a different mass number.

Properties of Isotopes:

Isotopes have the same chemical properties because they have the same number of electrons in their outermost

shell.

They have different physical properties e.g. melting point, boiling point, colour, density and rate of diffusion.

Isotopes are used all around the world in agriculture, medicine and even as smoke detectors. Radioactive

Isotopes are used in medicine for diagnosis and treatment of diseases such as cancer. This medicine is called

nuclear medicine. They are used in agriculture as pesticides.

Relative Atomic Mass

The relative atomic mass of an element is the average mass of its atoms compared to an atom of Carbon 12. This

is done as atoms are very small and so it would be very complicated to calculate their average mass.

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Calculate the relative atomic mass of an element from the relative abundances of its isotopes

How to calculate relative atomic mass:

Example for Chlorine:

75% Chlorine 35, 25% Chlorine 37

RAM (Relative Atomic Mass) = 75% * 35 + 25% * 37 = 35.5

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Understand that the Periodic Table is an arrangement of elements in order of atomic number

The Periodic Table

The periodic table is an arrangement of elements in order of increasing atomic number. All metals are on the left

hand side of the step ladder and all non-‐metals are on the right

__________________________________________________________________________________________________

Deduce the electronic configurations of the first 20 elements from their positions in the Periodic Table

The Periodic Table

The vertical columns are called groups and they tell us about the number of electrons in the outermost shell.

There are 8 groups

The horizontal rows are called periods. The periods number tells us the number of shells present around the

nucleus. The first period only has two elements

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Deduce the number of outer electrons in a main group element from its position in the Periodic Table

The Periodic Table


There are 8 groups

All the elements in the same group have the same number of electrons in their outermost shell giving them the

same or very similar chemical properties.

__________________________________________________________________________________________________

Calculate relative formula masses (Mr) from relative atomic masses (Ar)

Ar (O) = 16

Mr (O2) = 2 x 16 = 32

Mr (NO2) = (1 x 14) + (2 x 16)

= 14 + 32

= 46

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Understand the use of the term mole to represent the amount of substance

e is its relative atomic mass, or relative formula mass, in grams

Ar (C) = 12

Mr (C) = 12

Mass of one mole of carbon is 12 g

Mass of two moles of carbon is 24g

Mass of one mole of a compound = Mr (g)

One mole of any substance has 6.023 x 10²³ atoms, ions, molecules

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Carry out mole calculations using relative atomic mass (Ar) and relative formula mass (Mr)

n: number of moles (moles)

m: mass (g)

Mr: relative molecular mass (g)

Number of Moles = Mass ÷ Relative Molecular Mass

n = m ÷ Mr

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Write word equations and balanced chemical equations to represent the reactions studied in this specification

Use the state symbols (s), (l), (g) and (aq) in chemical equations to represent solids, liquids, gases and aqueous solutions respectively

Formation Reactions

o Burning elements with oxygen

Carbon + Oxygen Carbon Dioxide

C (s) + O2 (g) CO2 (g)

o Formation of compounds from their elements

Sodium + Chlorine Sodium Chloride

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2 Na (s) + Cl2 (g) 2 NaCl (s)

Combustion Reactions

o Complete combustion will form CO2 and H2O

CH4 + 2 O2 CO2 + 2 H2O

o Incomplete combustion will produce CO and H2O

2 CH4 + 3 O2 2 CO + 4 H2O

Acid Reactions

o Acid + Metal Salt + Hydrogen

o Acid + Carbonate Salt + Water + Carbon Dioxide

o Acid + Alkali Salt + Water

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Understand how the formulae of simple compounds can be obtained experimentally, including metal oxides, water and salts containing water of crystallisation

Weigh the crucible and lid

Weigh the crucible, lid and magnesium

o Mass of magnesium = Mass of the crucible, lid and magnesium Mass of the crucible and lid

Heat the magnesium till it is full oxidized

o When the contents no longer glow as the lid is lifted

Weigh the crucible, lid and magnesium oxide

o Mass of magnesium oxide = Mass of the crucible, lid and magnesium oxide Mass of the crucible and lid

o Mass of oxygen combined = Mass of magnesium oxide Mass of magnesium

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Calculate empirical and molecular formulae from experimental data

The empirical formula of a compound contains the simplest ratio of atoms in that compound

Calculating the empirical formula from masses:

Write down the symbols of each element

Write down the masses of each element

Divide each mass by the atomic mass of each element

Determine the simples whole number ratio

Write down the empirical formula of the compound

Calculating the empirical formula from percentages:

Write down the symbols of each element

Write down the percentage of each element

Write down the masses of each element

o This is done by simply assuming that the compound weighs 100g

Divide each mass by the atomic mass of each element

Determine the simples whole number ratio

Write down the empirical formula of the compound

__________________________________________________________________________________________________

Calculate reacting masses using experimental data and chemical equations

Calculate the number of moles of the given species

o Number of Moles = Mass ÷ Relative Molecular Mass

Calculate the number of moles of the required species

o This can be done by using mole ratio which is the ratio taken from the balancing numbers

Calculate the reactant mass

o Mass = Number of Moles x Relative Molecular Mass

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Carry out mole calculations using volumes and molar concentrations

Molar Volume


V: volume of gas (cm³)

Vr: molar volume (24 000 cm³)

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Number of Moles = Volume of Gas ÷ Molar Volume

n = V ÷ Vr

One mole of any gas has a fixed volume at a given temperature and pressure

The molar volume of any gas is 24 000 cm³ at room temperature and pressure

Room temperature is considered 25ºC

Room pressure is 1 atm

Solutions


C: concentration (mol/L)

V: volume (dm³)

Number of Moles = Concentration x Volume

n = C x V

The number of moles pertains to the amount of solute

The volume pertains to the amount of solution

Calculations only for Gases

To find the volume of a gas, simply use volume ratios taken from the balancing numbers

Example:

2 H2 (g) + O2 (g) 2 H2O (g)

30 cm³ + 10 cm³

Ratio of oxygen to water is 1:2

Volume of oxygen is 10 cm³

Therefore, volume of water is 20 cm³

Volume of gases left at the end of the reaction = 30 cm³

That is because there was 10 cm³ excess of hydrogen

Calculations in Solutions

To find the concentration or volume of a solution, simply use the following equation

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(C1V1) ÷ (C2V2) = n1 ÷ n2

However, the mole ratio on the right hand side of the equation can be taken from the balancing numbers

Example:

NaOH + HCl NaCl + H2O

30 cm³ + 10 cm³

C (NaOH) = ?

V (NaOH) = 0.25 dm³

C (HCl) = 0.1 mol/dm³

V (HCl) = 0.2 dm³

(C1V1) ÷ (C2V2) = n1 ÷ n2

(C1 x 0.25) ÷ (0.1 x 0.2) = 1 ÷ 1

C1 = 0.02 ÷ 0.25

= 0.08 mol/dm³

__________________________________________________________________________________________________

Describe the formation of ions by the gain or loss of electrons

An ion

Is a charged particle which is formed when an atom loses or gains electrons (the number of electrons is not

equal to the number of protons)

o An atom that loses electrons has more protons than electrons and so has a positive overall charge. This

is called a positive ion

o An atom that gains electrons has more electrons than protons and so has a negative overall charge. This

is called a negative ion

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Understand oxidation as the loss of electrons and reduction as the gain of electrons

Oxidation

Is

Loss of electrons

Reduction

Is

Gain of electrons

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Recall the charges of common ions in this specification

Name of the Atom Symbol Formula of the ion

Sodium Na Na

Magnesium Mg Mg²

Aluminium Al Al³

Chlorine Cl Cl

Bromine Br Br

Sulfur S S²

Strontium Sr Sr²

Nitrogen N N³

Helium He He

Iodine I I

Barium Ba Ba²

Caesium Cs Cs

Boron B B³

Carbonate CO

Sulphate SO

Ammonium NH

Nitrate NO

Hydroxide OH

Phosphate PO4³

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Deduce the charge of an ion from the electronic configuration of the atom from which the ion is formed

How to find the charge of an ion

The charge of an ion depends on which group the atom belongs to:

Group 1 +1

Group 2 +2

Group 3 +3

Group 5 -‐3

Group 6 -‐2

Group 7 -‐1

Group 0 no charge (already have full outer shell)

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Explain, using dot and cross diagrams, the formation of ionic compounds by electron transfer, limited to combinations of elements from Groups 1, 2, 3, and 5, 6, 7

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Understand ionic bonding as a strong electrostatic attraction between oppositely charged ions

An Ionic Bond

Is a strong electrostatic force between oppositely charged ions

It is a compound that contains a metallic element and a non-‐metallic element

Is formed either through the gain or loss of electrons

It is formed in order for the atoms to become stable

o They try to obtain an inert gas configuration

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Understand that ionic compounds have high melting and boiling points because of strong electrostatic forces between oppositely charged ions.

Properties of ionic compounds:

1. Most ionic compounds are soluble in water.

2. In solid state they do not conduct electricity due to the presence of ions. However when molten or liquid, they

can conduct electricity due to the movement of ions

3. They have high melting and boiling points because of strong electrostatic forces of attraction between

oppositely charged ions

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Describe the formation of a covalent bond by the sharing of a pair of electrons between two atoms

A covalent bond is formed between two or more non-‐metals by the sharing of electrons

Understand covalent bonding as a strong attraction between the bonding pair of electrons and the nuclei of the atoms involved in the bond

The two atoms involved in bonding show a strong attraction between the bonding pair of electrons and the nuclei of the atoms involved in the bond

Explain, using dot and cross diagrams, the formation of covalent compounds by electron sharing for the following substances:

i. Hydrogen

ii. Chlorine

iii. Hydrogen Chloride

iv. Water

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v. Methane

vi. Ammonia

vii. Oxygen

viii. Nitrogen

ix. Carbon Dioxide

x. Ethane

xi. Ethene

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Recall that substances with simple molecular structures are gases or liquids, or solids with low melting points

Explain why substances with simple molecular structures have low melting points in terms of the relatively weak forces between the molecules

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Properties of Simple Covalent Structures:

Most substances that contain simple covalent molecules have low melting and boiling points and are therefore

liquids or gases at room temperature. This is because the covalent bonds within the molecules are strong but the

bonds between molecules are weak and easy to break.

For example water, oxygen, carbon dioxide, chlorine and hydrogen.

They are also soft and brittle and cannot conduct electricity.

Strong bonds within intramolecular molecules

Weak bonds within intermolecular molecules

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Explain the high melting points of substances with giant covalent structures in terms of the breaking of many strong covalent bonds

In some substances such as sand, diamond and graphite, millions of atoms are joined together by covalent bonds. The

bonds in these substances do not form molecules but vast networks of atoms called giant covalent structures.

All the bonds are covalent, so giant covalent structures have very high melting and boiling points and are usually hard as it requires a lot of energy to break many strong covalent bonds

Sand

Structure of Sand

Sand is made up of the mineral quartz which is silicon dioxide. It has a giant covalent structure made up of silicon and

oxygen atoms.

Diamond

Structure of Diamond

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Properties of Diamond

All the electrons in the outer shell of each carbon atom are involved in forming covalent bonds.

Diamonds cannot conduct electricity because there are no free electrons or ions to carry a charge.

Graphite

Structure of Graphite

Properties of Graphite

In graphite, only three of the four electrons in the outer shell of each carbon atom are involved in covalent bonds.

Graphite is soft and slippery; layers can easily slide over each other as the weak forces of attraction between these

layers can be easily broken. This is why graphite is often used as a lubricant.

Graphite conducts electricity and so it is the only non-metal to do so. This is because each layer has delocalized

electrons from each carbon atom which can carry a charge.

Allotropes of Carbon

Both diamond and graphite are made up of carbon atoms. Different forms of the same element are called allotropes. These allotropes of carbon have different properties because the atoms are bonded in different arrangements which create different giant structures

__________________________________________________________________________________________________

Describe a metal as a giant structure of positive ions surrounded by a sea of delocalized electrons

Structure of metals

The atoms in a pure metal are in tightly-packed layers, which form a regular lattice structure

The outer electrons of the metals atoms separate from the atoms and create a These electrons are delocalized and so are free to move through the whole structure The metal atoms become positively charged ions and are attracted to the sea of electrons This attraction is called metallic bonding

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Explain the malleability and electrical conductivity of a metal in terms of its structure and bonding

Why are metals strong?

Metals will usually be strong and not brittle. This is because when a metal is hit, the layers of metal ions are able to slide

over each other and so the structure does not shatter

The metallic bonds do not break because the delocalized electrons are free to move throughout the structure which also

explains why metals are malleable and ductile

Malleable: easy to shape.

Ductile: can be drawn into wires

__________________________________________________________________________________________

Understand an electric current as a flow of electrons or ions

An electric current is a flow of electrons or ions

__________________________________________________________________________________________

Understand why covalent compounds do not conduct electricity

Covalent compounds do not conduct electricity due to the absence of free electrons as electrons are shared

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Understand why ionic compounds conduct electricity only when molten or in solution

Properties of ionic compounds:

1. In solid state they do not conduct electricity due to the presence of ions. However when molten or liquid, they

can conduct electricity due to the movement of ions

__________________________________________________________________________________________

Describe simple experiments to distinguish between electrolytes and non-electrolytes

__________________________________________________________________________________________________

Recall that electrolysis involves the formation of new substances when ionic compounds conduct electricity

In electrolysis, the substance that the current passes through and splits up is called the electrolyte

The electrolyte contains positive and negative ions:

Anions (negative ions) move to the anode (positive electrode) and lose electrons (oxidation)

Cations (positive ions) move to the cathode (negative electrode) and gain electrons (reduction)

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Describe simple experiments for the electrolysis, using inert electrodes, of molten salts such as lead (I I ) bromide

If you pass electricity through the molten salts:

Lead accumulates at the negative electrode

Bromine accumulates at the positive electrode

__________________________________________________________________________________________

Write ionic half-equations representing the reactions at the electrodes during electrolysis

Lead Bromide Lead + Bromine

PbBr (l) Pb (l) + Br (g)

At the negative electrode: At the positive electrode:

Reduction Oxidation

Pb + 2 e Pb Br2 + 2 e

__________________________________________________________________________________________

Understand the terms group and period

The Periodic Table


There are 8 groups.

The horizontal rows are called periods. The periods number tells us the number of shells present around the

nucleus. The first period only has two elements.

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Recall the positions of metals and non-metals in the Periodic Table

All metals are on the left hand side of the step ladder and all non-‐metals are on the right

__________________________________________________________________________________________________

Explain the classification of elements as metals or non-metals on the basis of their electrical conductivity and the acid-base character of their oxides

Electrical conductivity of metals is high as the electrons are free to conduct electricity.

Non-‐

All metals form oxides which are basic in nature.

All non-‐metals form oxides which are acidic in nature.

Examples:

Sodium + Oxygen Sodium Oxide

Sodium Oxide + Water Sodium Hydroxide (basic)


Carbon Dioxide + Water Carbonic Acid (acidic)

__________________________________________________________________________________________

Understand why elements in the same group of the Periodic Table have similar chemical properties

All the elements in the same group have the same number of electrons in their outermost shell giving them the same or

very similar chemical properties

This is mainly because the atoms of the elements gain/lose electrons in a similar manner

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Recall the noble gases (Group 0) as a family of inert gases and explain their lack of reactivity in terms of their electronic configurations

Group 0 (Inert Gases)

Helium-‐ Hot air balloons

Neon-‐ Advertising signs/neon lights

Argon-‐ Bulbs

Krypton-‐ Advertising signs

Xenon-‐ Flash gun in cameras

of electrons which make them very unreactive which is why they are the inert gases. They are stable due to their full outer shell of electrons and so elements

__________________________________________________________________________________________

Describe the reactions of these elements with water and understand that the reactions provide a basis for their recognition as a family of elements

Reactivity of Alkali Metals

Sodium

Observation of Sodium in water:

o It was fizzing.

o It was floating on the surface of water.

o Saw fumes/smoke/a gas was given off.

o It became smaller and smaller in size.

o It was moving around on the surface of water.

o After the reaction the resultant solution turned the universal indicator blue.

Conclusion:

o Sodium is very reactive.

o It has a very low density since it floats on the water.

o The gas produced was hydrogen

Test for Hydrogen

Place a lighted splint in a test-‐tube of gas

If you hear a squeaky pop, hydrogen is present

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o Alkali Metals in water produce alkaline solutions which turn universal indicator solution blue

Chemical Reaction:

o 2Na + 2H O 2NaOH + H

o Sodium + Water Sodium Hydroxide + Hydrogen

Potassium

Observation of Potassium in water:

o The reaction with Potassium was much more vigorous than Sodium.

o Potassium was skidding on the surface of water.

o A gas was given off.

o It was fizzing.

o Potassium was floating on the surface of water and getting smaller in size.

o It burnt with a lilac flame.

o When universal indicator was added to the resulting solution, it turned blue.

Conclusion:

o Potassium is very reactive, even more reactive than Sodium and it also burns with a lilac flame (purple

flame)

o It has a very low density since it floats on the water.

o The gas produced was hydrogen

Test for Hydrogen


If you hear a squeaky pop, oxygen is present

o Alkali Metals in water produce alkaline solutions which turn universal indicator solution blue

Chemical Reaction:

2K + 2H O 2KOH + H

Potassium + Water Potassium Hydroxide + Hydrogen

__________________________________________________________________________________________

Recall the relative reactivities of the elements in Group 1

Group 1 (Alkali Metals)

Li Lithium

Na -‐ Sodium

K -‐ Potassium

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Rb -‐ Rubidium

Cs -‐ Caesium

Fr Francium

Reactivity of Group 1 Elements

The reactivity of group 1 elements increase as you go down the group as the size of the atom increases and

there are also more shells around the nucleus so the outermost electron has very little nuclear attraction and it

can easily be lost in chemical reactions

This group is also the most reactive group as it has one electron in the outermost shell and so it is very unstable

and needs to react

__________________________________________________________________________________________ Recall the colours and physical states of the elements at room temperature

Fluorine is pale yellow (gas)

Chlorine is greenish yellow (gas)

Bromine is reddish brown (liquid)

Iodine is purple (solid)

Astatine is black (solid)

__________________________________________________________________________________________ Make predictions about the properties of other halogens in this group

All Halogens are:

Non-‐metals and so do not conduct electricity

Brittle and crumbly when solid

Poisonous and smelly

__________________________________________________________________________________________

Understand the difference between hydrogen chloride gas and hydrochloric acid

The difference between hydrogen chloride gas and hydrochloric acid:

Hydrogen Chloride ions and so it is not acidic in nature (Hydrogen chloride gas

does not turn blue litmus red)

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However Hydrochloric Acid is produced when hydrogen chloride gas is dissolved in water. Hydrochloric Acid

dissociated to give H ions which are responsible for the acidic nature (they turn blue litmus paper red) making

Hydrochloric Acid, acidic in nature

HCl (g) + H O = HCl (aq)

H Cl

(Dissociation is the temporary or reversible process in which a molecule or ion is broken down into smaller molecules or

ions)

__________________________________________________________________________________________

Explain, in terms of dissociation, why hydrogen chloride is acidic in water but not in methylbenzene

HCl in Water

o Particles dissociate to produce H ions

Resulting solution is acidic in nature

HCl in Methylbenzene

o Particles do not dissociate to produce H ions

Resulting solution is not acidic in nature

__________________________________________________________________________________________

Recall the relative reactivities of the elements in Group 7

Reactivity of Group 7 Elements

In group 7 (halogens) the reactivity of the elements decreases as you go down the group (Fluorine is the most

reactive and Iodine is the least).

This is because as you go down the group, the atomic size increases and so it is difficult for the atom to attract

electrons as the shieldy effect increases and so the nuclear attraction decreases. Therefore Fluorine is the most

reactive halogen and iodine is the least reactive halogen.

Fluorine water will be able to displace all the other halogens and undergo all displacement reactions

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Describe experiments to show that a more reactive halogen will displace a less reactive halogen from a solution of one of its salts

As the colour of the solution is taken from the atom and not the ion, we are able to see, which element gets displaced

Fluorine + Potassium Bromide Potassium Fluoride + Bromine

o Solution turns reddish brown

Therefore, bromine got displaced

Therefore, fluorine is more reactive than bromine

Chlorine + Potassium Iodide Potassium Chloride + Iodine o Solution turns purple

Therefore, iodine got displaced

Therefore, chlorine is more reactive than iodine

__________________________________________________________________________________________

Understand these displacement reactions as redox reactions.

Oxidation

Is

Loss of electrons

Reduction

Is

Gain of electrons

__________________________________________________________________________________________

Recall the gases present in air and their approximate percentage by volume

Gas Amount in Air (%)

Nitrogen 78.1

Oxygen 21.0

Argon 0.9

Carbon Dioxide 0.04

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Describe how experiments involving the reactions of elements such as copper, iron and phosphorus with air can be used to determine the percentage by volume of oxygen in air

Showing that air contains about one-‐fifth oxygen

Using Copper

2 Cu (s) + O2 (g) = 2 CuO (s)

Using Iron

4 Fe (s) + 3 O2 (g) = 2 Fe2O3 (s)

Using Phosphorous

Phosphorous shoulders in air to produce two different phosphorous oxides

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Describe the laboratory preparation of oxygen from hydrogen peroxide

Making Oxygen in the Lab

Catalytic Decomposition

Splitting up using a catalyst

Hydrogen Peroxide Manganese Oxide Water + Oxygen

2 H2O2 (aq) MnO2 2 H2O (l) + O2 (g)

Test for Oxygen

Place a glowing splint into a test tube of gas

If the splint re-‐ignites, oxygen is present

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Describe the reactions with oxygen in air of magnesium, carbon and sulfur, and the acid base character of the oxides produced

Burning Elements in Oxygen

Burning Magnesium

o Burns in air with a bright white flame

o Gives a white, powdery ash of magnesium oxide

o Extremely bright in pure oxygen

Magnesium + Oxygen Magnesium Oxide

2 Mg (s) + O2 (g) 2 MgO (s)

Burning Sulfur

o Burns in air with a tiny, almost invisible, blue flame

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o Poisonous, colourless sulfur dioxide gas is produced

o Bright blue flame in oxygen

Sulfur + Oxygen Sulfur Dioxide

S (s) + O2 (g) 2 SO2 (g)

Burning Carbon

o Burns when heated very strongly in air or oxygen

o Burns with a small yellow-‐range flame

o Sometimes produces sparks

o Colourless carbon dioxide gas is produced


C (s) + O2 (g) CO2 (g)

Metal and Non-‐Metal Oxides

those that do, tend to form alkaline solutions

o Magnesium Oxide + Water Magnesium Hydroxide

o MgO (s) + H2O (l) Mg(OH)2 (s and aq)

Non-‐metal oxides often react with water to form acidic solutions common exceptions are water and carbon

monoxide

o Water + Sulfur Dioxide Sulfurous Acid

o H2O (l) + SO2 (g) H2SO3 (aq)

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Describe the laboratory preparation of carbon dioxide from calcium carbonate and dilute hydrochloric acid

Making Carbon Dioxide in the Lab

Calcium Carbonate + Hydrochloric Acid Calcium Chloride + Carbon Dioxide + Water

CaCO3 (s) + 2 HCl (aq) CaCl2 (aq) + CO2 (g) + H2O (l)

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Test for Carbon Dioxide

o Add some lime water to a test tube filled with gas

o Shake the test tube

o If the solution turns milky, carbon dioxide is present

Calcium Hydroxide + Carbon Dioxide Calcium Carbonate + Water

Ca(OH)2 (aq) + CO2 (g) CaCO3 (s) + H2O (l)

Calcium Carbonate + Carbon Dioxide + Water Calcium Hydrgencarbonate

CaCO3 (s) + CO2 (g) + H2O (l) Ca(HCO3)2 (aq)

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Describe the formation of carbon dioxide from the thermal decomposition of metal carbonates such as copper (I I ) carbonate

CuCO3 CuO + CO2

Copper Carbonate Copper Oxide + Carbon Dioxide

Green Black

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Recall the properties of carbon dioxide, limited to its solubility and density

Properties of Carbon dioxide

Colourless

Odourless

Denser than air

Slightly soluble in water

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Explain the use of carbon dioxide in carbonating drinks and in fire extinguishers, in terms of its solubility and density

Uses of Carbon Dioxide

Used in carbonated drinks

o As it dissolves in water under pressure

Used in fire extinguishers

o As the dense gas sinks into the flames and prevents any more oxygen from reaching them

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Recall the reactions of carbon dioxide and sulfur dioxide with water to produce acidic solutions

Carbon Dioxide

o Water + Carbon Dioxide Carbonic Acid

o H2O (l) + CO2 (g) H2CO3 (aq)

Sulfur Dioxide

o Water + Sulfur Dioxide Sulfurous Acid

o H2O (l) + SO2 (g) H2SO3 (aq)

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Recall that sulfur dioxide and nitrogen oxides are pollutant gases which contribute to acid rain, and describe the problems caused by acid rain

Non-‐Metal Oxides and the Environment

Acid rain is caused when water and oxygen in the atmosphere react with sulfur dioxide to produce sulfuric acid, or with

various oxides of nitrogen to give nitric acid

Sulfur dioxide mainly comes from power stations and factories burning fossil fuels

Oxides of nitrogen are produced from motor vehicles

Problems:

Erosion of limestone and metals

Kills fish in lakes

Contribute to the death of plants

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Describe the reactions of dilute hydrochloric and dilute sulfuric acids with magnesium, aluminium, zinc and iron

Bubbles of gas released

Hydrochloric Acid

o Magnesium + Hydrochloric Acid Magnesium Chloride + Hydrogen

o Aluminium + Hydrochloric Acid Aluminium Chloride + Hydrogen

o Zinc + Hydrochloric Acid Zinc Chloride + Hydrogen

o Iron + Hydrochloric Acid Iron Chloride + Hydrogen

Sulfuric Acid

o Magnesium + Sulfuric Acid Magnesium Sulfate + Hydrogen

o Aluminium + Sulfuric Acid Aluminium Sulfate + Hydrogen

o Zinc + Sulfuric Acid Zinc Sulfate + Hydrogen

o Iron + Sulfuric Acid Iron Sulfate + Hydrogen

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Describe the combustion of hydrogen

Test for Hydrogen


If you hear a squeaky pop, oxygen is present

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Describe the use of anhydrous copper (I I ) sulfate in the chemical test for water

Water turns white anhydrous copper (II) sulfate blue

Anhydrous copper (II) sulfate lacks water of crystallization and is white. Dropping water onto it replaces the water of crystallization, and turns it blue

CuSO4 (s) + 5 H2O (l) CuSO4 5 H2O (s)

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Describe a physical test to show whether water is pure.

Pure water freezes at exactly 0º C and boils at exactly 100º C at 1 atmospheric pressure

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Recall that metals can be arranged in a reactivity series based on the reactions of the metals and their compounds: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver and gold

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When metals react in similar ways, we say that they have similar chemical properties

When metals react with water, the gas formed burns with a squeaky pop o When metals react with water, they always form a metal hydroxide and hydrogen gas o Metal + Water Metal Hydroxide + Hydrogen

Metal + Cold Water Metal Hydroxide + Hydrogen

If the metal is more reactive than aluminium Heated Metal + Steam Metal Oxide + Hydrogen

If the metal is more reactive than tin

When metals react with water, the solution is alkaline o When metals react with water, the solution turns red litmus paper blue

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Describe how reactions with water and dilute acids can be used to deduce the following order of reactivity: potassium, sodium, lithium, calcium, magnesium, zinc, iron, and copper

When metals react with acid, the gas formed burns with a squeaky pop o When metals react with acid, they always form a metal salt and hydrogen gas o Metal + Water Salt + Hydrogen

If the metal is at least as reactive as lead

Potassium Most Reactive

Sodium

Lithium

Calcium

Magnesium

Aluminium

Carbon

Zinc

Iron

Tin

Lead

Hydrogen

Copper

Mercury

Silver

Gold

Platinum Least Reactive

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Hydrochloric acid makes metal chloride

Sulphuric acid make metal sulphate

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Deduce the position of a metal within the reactivity series using displacement reactions between metals and their oxides, and between metals and their salts in aqueous solutions

When metals react with oxygen, they always form a metal oxide

Metal + Oxygen Metal Oxide

Competition for Oxygen

Involves the reaction of a metal with the oxide of another metal o This results in the two metals competing for the oxygen

The more reactive metal finishes up with the oxygen (as a metal oxide) If the more reactive metal starts as the oxide, then no reaction takes place

Potassium Most Reactive

Sodium

Lithium

Calcium

Magnesium

Aluminium

Carbon

Zinc

Iron

Tin

Lead

Hydrogen

Copper

Mercury

Silver

Gold

Platinum Least Reactive

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Extracting Metal with Oxygen

Consists of two competition reactions o A metal oxide is reacted with charcoal

If the charcoal (carbon) is more reactive, it will remove the oxygen from the metal oxide and leave a trace of metal in the reaction vessel

Displacing Metals from Solution

An ionic solution is collected in a test tube o A metal is placed in the solution

If the metal is more reactive than the metallic element in the ionic solution, it displaces the metallic element

The electrons move from the metal to the ions o The metal atoms become ions o The metallic ions become atoms

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Word Equation

Copper (II) Sulphate + Iron Copper + Iron Sulphate

Chemical Equation

CuSO4 (aq) + Fe (s) Cu (s) + FeSO4 (aq)

Ionic Equation

Cu² (aq) + SO4² (aq) + Fe (s) Cu (s) + SO4² (aq) + Fe² (aq)

Spectator Ions

Sulphate -‐ SO4² (aq)

Ionic Equation with Spectator Ions

Cu² (aq) + Fe (s) Cu (s) + Fe² (aq)

Ion-‐Electron Equations

Oxidation

The iron atoms lose electrons to form iron ions

Fe (s) Fe² (aq) + 2 Electrons

Reduction

The copper ions gain electrons to form copper atoms

Cu² (aq) + 2 Electrons Cu (s)

The iron atoms lose electrons which are gained by the copper ions

If metal atoms and metal ions (in a metal salt) are mixed together, the more reactive metal will always end up as metal ions and the less reactive metal will always end up as metal atoms

Reactive Metal + Less Reactive Metal Salt Less Reactive Metal + Reactive Metal Salt

Elements near the top of the electrochemical series lose electrons and form ions very readily

Elements near the bottom of the electrochemical series stay as atoms or if they are ions, they gain electrons very readily for form atoms

The more reactive metal displaces the less reactive metal

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Understand oxidation and reduction as the addition and removal of oxygen respectively

Any metal element reacting to form a compound is an example of oxidation

When a metal reacts, the only thing it can do is lose electrons and become a metal oxide

Loss of electrons is called oxidation

When metals react, we say they are oxidised

Gain of electrons is called reduction

Reduction and oxidation reactions always take place together

Hydrogen being displaced from a solution of sulfuric acid by zinc

Oxidation

o Zn Zn² + 2e

Reduction

o 2H + 2e H2

REDOX

o 2H + Zn H2 + Zn²

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Understand the terms redox, oxidising agent and reducing agent

REDOX (Reduction-‐Oxidation) o Reversible chemical reaction in which one reaction is an oxidation and the reverse is a reduction

Oxidising Agent o A substance that gains electrons in a redox chemical reaction

The oxidizing agent becomes reduced in the process

Reducing Agent o A substance that loses electrons in a redox chemical reaction

The reducing agent becomes oxidised in the process

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Recall the conditions under which iron rusts

Iron rusts in the presence of oxygen and moisture

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Describe how the rusting of iron may be prevented by grease, oil, paint, plastic and galvanising

A coating on the surface of an iron object will protect it from rusting by preventing the contact between the iron, oxygen and water.

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Galvanising o The process of coating iron and steel objects in a layer of zinc

Tin-‐Plating o The process of coating iron and steel object in a layer of tin

Electroplating o The process of applying a metal (most often gold) to adhere to the surface of another metal using

electrical current

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Understand the sacrificial protection of iron in terms of the reactivity series.

The galvanized iron does not rust even after the zinc coating has been broken because electrons always flow from more reactive metals to less reactive metals. Therefore, all the zinc must get oxidised before the iron can start getting oxidised.

The tin-‐plated iron rusts fast after the tin coating has been broken because electrons always flow from more reactive metals to less reactive metals. Therefore, the iron will get oxidised before the tin gets oxidised.

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Describe simple tests for the cations:

i. Li+, Na+, K+, Ca2+ using flame tests

Flame tests are used to show the presence of certain metal ions in a compound

A platinum or nichrome wire is cleaned by dipping it into concentrated hydrochloric acid and then holding it in a hot Bunsen flame

The wire is dipped back into the acid, then into a tiny sample of the solid you are testing, and back into the flame

o Li Red flame

o Na Strong, persistent orange flame

o K Lilac (pink) flame

o Ca² Orange-‐red (brick red) flame

ii. NH4+ using sodium hydroxide solution and identifying the ammonia evolved

Sodium hydroxide reacts with ammonium salts (either solid or in solution) to produce ammonia gas

o In the cold There is just enough ammonia gas produced for you to be able to smell it

o When warmed You can test the gas coming off with a piece of damp red litmus paper

Ammonia is alkaline and turns the litmus paper blue

NH4 (s or aq) + OH (aq) NH3 (g) + H2O (l)

NH4Cl (s) + NaOH (aq) NaCl (aq) + NH3 (g) + H2O (l)

No precipitate, but a smell of ammonia

iii. Cu2+, Fe2+ and Fe3+ using sodium hydroxide solution

Cu² (aq) + 2OH (aq) Cu(OH)2 (s)

CuSO4 (aq) + 2NaOH (aq) Cu(OH)2 (s) + Na2SO4 (aq)

If copper (II) ions are present, a blue precipitate is formed

Fe² (aq) + 2OH (aq) Fe(OH)2 (s)

FeSO4 (aq) + 2NaOH (aq) Fe(OH)2 (s) + Na2SO4 (aq)

If iron (II) ions are present, a green precipitate is formed

Fe³ (aq) + 3 OH (aq) Fe(OH)3 (s)

FeCl3 (aq) + 3 NaOH (aq) Fe(OH)3 (s) + 3NaCl (aq)

If iron (III) ions are present, an orange-‐brown precipitate is formed

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Describe simple tests for the anions:

i.

Make a solution of your suspected chloride, bromide or iodide

Add enough nitric acid to make the solution acidic o This is to remove the other substances which might also produce precipitates with silver nitrate solution

Then, add some silver nitrate solution

Ag (aq) + Cl (aq) AgCl (s)

A white precipitate (of silver chloride) shows the presence of chloride ions

Ag (aq) + Br (aq) AgBr (s)

A pale cream precipitate (of silver bromide) shows the presence of bromide ions

Ag (aq) + I (aq) AgI (s)

A yellow precipitate (of silver iodide) shows the presence of iodide ions

ii.

Dilute hydrochloric acid reacts with a sulphate solution to produce a white precipitate

Make a solution of the suspected sulphate

Add enough hydrochloric acid to make the solution acidic o This is to remove the other substances which might also produce precipitates with barium chloride

solution

Then, add some barium chloride solution

If a white precipitate is produced, the suspected solution contains sulphate ions

Ba² (aq) + SO4² (aq) BaSO4 (s)

iii.

Dilute hydrochloric acid reacts with a solid carbonate to produce carbon dioxide

In a test tube, add a little dilute hydrochloric acid to the suspected solid carbonate

Look for bubbles of gas produced in the cold

Test the gas with lime water to show that it is carbon dioxide

If the limewater turns milky, the suspected solid contains carbonate ions

CO3² (s) + 2H (aq) CO2 (g) + H2O (l)

ZnCO3 (s) + 2HCl (aq) ZnCl2 (aq) + CO2 (g) + H2O (l)

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Describe simple tests for the gases:

i. Hydrogen

Place a lighted splint into a test tube of gas

If you hear a squeaky pop, hydrogen is present

ii. Oxygen

Place a glowing splint into a test tube of gas

If the glowing splint re-‐ignites, oxygen is present

iii. Carbon Dioxide

Add some of the gas to a test tube of lime water

Shake the test tube

If the lime water turns milky, carbon dioxide is present

iv. Ammonia

Add hydrogen chloride gas to a test tube of the solution

If a white smoke is released, ammonia is present

o The white smoke is ammonium chloride

Ammonia turns damp red litmus paper blue

v. Chlorine

Chlorine is the only gas which has a bleaching effect

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Explain the terms homologous series, hydrocarbon, saturated, unsaturated, general formula and isomerism

Homologous Series

o A family of chemical compounds which have the same general formula and similar chemical properties

but show a gradual change in physical properties such as melting point and boiling point

o Successive members differ by CH2

Hydrocarbons

o Are organic compounds which contain only hydrogen and carbon

Saturated

o The molecule has no carbon to carbon double bonds

The molecule only has carbon to carbon single bonds

o The molecule has the maximum number of atoms

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No more atoms can be added to the molecule

Unsaturated

o The molecule has at least one carbon to carbon double bond

o Atoms can still be added to the molecule

General Formula

o Is a way of expressing information about the atoms that constitute a particular chemical compound

o For the Homologous Series it shows the relationship between the number of C atoms and H atoms in the

compounds

Isomerism

o Is the phenomenon whereby certain compounds, with the same molecular formula, exist in different

forms owing to their different arrangement of atoms

Structural Isomers

o Have different structural formulae because their atoms are linked together in different ways

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Recall that alkanes have the general formula CnH2n+2

Alkanes

Simplest family of organic compounds

Saturated compounds

o Compounds with only single carbon to carbon bonds

Names of all members end in ane

o Methane

o Ethane

o Propane

o Butane

o Pentane

o Hexane

o Heptane

o Octane

General formula

o CnH2n+2

Are all flammable

Show a gradual change in melting and boiling point

o First four members are gases at room temperature

o Next thirteen members are liquids at room temperature

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All the other members are solids at room temperature

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Draw displayed formulae for alkanes with up to five carbon atoms in a molecule, and name the straight-chain isomers

Number

of

Carbon

Atoms

Number

of

Hydrogen

Atoms

Name Molecular

Formula Full Structural Formula Shortened Structural Formula

1 4 Methane CH4

CH4

2 6 Ethane C2H6

CH3CH3

3 8 Propane C3H8

CH3CH2CH3

4 10 Butane C4H10

CH3CH2CH2CH3

5 12 Pentane C5H12

CH3CH2CH2CH2CH3

6 14 Hexane C6H14

CH3CH2CH2CH2CH2CH3

7 16 Heptane C7H16

CH3CH2CH2CH2CH2CH2CH3

8 18 Octane C8H18

CH3CH2CH2CH2CH2CH2CH2CH3

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C4H10

Butane

Methylpropane

C5H12

Pentane

2 Methylbutane

2, 2 Dimethylpropane

The alkanes have covalent bonding

They have a simple molecular structure

Their melting and boiling points will be low because only the intermolecular forces of attraction between their

molecules are broken when they melt or boil

Both, the melting point and boiling points increases as the size of the molecule increases

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Recall the products of the complete and incomplete combustion of alkanes

Combustion (or burning) is the process which takes place in the presence of oxygen to produce carbon dioxide and

water vapour

Burning Methane

o Observation:

Soot collected inside the funnel

Anhydrous cobalt chloride paper turned from blue to pink

Anhydrous copper sulphate turns from colourless to blue

Limewater turns cloudy

o Methane + Oxygen = Carbon Dioxide + Water Vapour

o CH4 + 2 O2 = CO2 + 2 H2O

Burning a Candle

o Wax + Oxygen = Carbon Dioxide + Water Vapour

Burning Hexane

o Hexane + Oxygen = Carbon Dioxide + Water Vapour

o 6 C6H14 + 19 O2 = 12 CO2 + 14 H2O

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Incomplete Combustion

Takes place when there is not enough oxygen present

CH4 + 1.5 O2 = CO + 2 H2O

CH4 + O2 = C + 2 H2O

Carbon Monoxide

Is a toxic gas because it combines with haemoglobin in the blood and prevent oxygen from reaching the cells

o Thus, it causes respiratory disorders

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Recall the reaction of methane with bromine to form bromomethane in the presence of UV light.

Halogenation

Is the replacement of one or more hydrogen atoms in an organic compound by a halogen atom

This reaction does not take place in the dark but only in the presence of UV light

o When excess of methane is reacted with chlorine, the products are chloromethane and hydrogen

chloride

o When excess of methane is reacted with bromine, the products are bromomethane and hydrogen

bromide

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Bromination of Methane

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Recall that alkenes have the general formula CnH2n

Alkenes

Unsaturated compounds

o Compounds with at least on carbon to carbon double bond

Names of all members end in ene

o Ethene

o Propene

o Butene

o Pentene

o Hexene

o Heptene

o Octene

General formula

o CnH2n

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Draw displayed formulae for alkenes with up to four carbon atoms in a molecule, and name the straight-chain isomers Number

of

Carbon

Atoms

Number

of

Hydrogen

Atoms

Name Molecular

Formula Full Structural Formula Shortened Structural Formula

2 4 Ethene C2H4

CH2=CH2

3 6 Propene C3H6

CH3CH=CH2

4 8 But-‐1-‐ene C4H8

CH3CH2CH=CH2

4 8 But-‐2-‐ene C4H8

CH3CH=CHCH3

5 10 Pent-‐1-‐

ene C5H10

CH3CH2CH2CH=CH2

5 10 Pent-‐2-‐

ene C5H10

CH3CH2CH=CHCH3

6 12 Hex-‐1-‐ene C6H12

CH3CH2CH2CH2CH=CH2

6 12 Hex-‐2-‐ene C6H12

CH3CH2CH2CH=CHCH3

6 12 Hex-‐3-‐ene C6H12

CH3CH2CH=CHCH2CH3

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Describe the addition reaction of alkenes with bromine, including the decolorising of bromine water as a test for alkenes

Test for Unsaturation

Add a few drops of bromine water to a test tube containing hydrocarbons

Place a stopper in the test tube and shake

o Saturated compounds show no colour change

o Unsaturated compounds turn bromine water from orange to colourless

Alkenes are more reactive than alkanes and cycloalkanes because of the carbon to carbon double bond

When an alkene reacts, this double bond can split open, allowing other substances to add on to the alkene

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Describe the use of the indicators litmus, phenolphthalein and methyl orange to distinguish between acidic and alkaline solutions

Indicators Colours in Acid Colours in Base

Litmus Blue -‐ Red Red -‐ Blue

Phenolphthalein Colourless -‐ Milky/Cloudy Colourless -‐ Pink

Methyl Orange Orange -‐ Red Orange -‐ Yellow

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Understand how the pH scale, from 0-14, can be used to classify solutions as strongly acidic, weakly acidic, neutral, weakly alkaline or strongly alkaline

PH scale

1 2 3 4 5 6 7 8 9 10 11 12 13 14

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Describe the use of universal indicator to measure the approximate pH value of a solution

The Universal Indicator is a mixture of indicators. Each universal indicator colour is given a PH value so that you can measure the approximate PH value of the solution. It is more accurate than other indicators as you can tell if the

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Define acids as sources of hydrogen ions, H+, and alkalis as sources of hydroxide ions, OH¯

Properties of Acids, Alkalis and Bases

ACIDS

Taste sour

Corrode metals

Are electrolytes

React with bases to form salt and water

pH is less than 7

Turns blue litmus paper to red

o All acids produce H (hydrogen) ions in solution e.g.

Hydrochloric acid (HCl)

Sulphuric acid (H SO )

Nitric acid (HNO )

Acetic acid (CH COOH)

ALKALIS

o Are water soluble bases

o All Alkalis produce OH (hydroxide) ions in solution e.g.

Sodium hydroxide (NaOH)

Potassium hydroxide (KOH)

Ammonium hydroxide (NH OH)

Strong Acid Weak Acid Strong Alkali Weak Alkali Neutral

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BASES

o Are any substances that neutralizes an acid

Produce OH ions in water

Tastes bitter and chalky

Are electrolytes

Feel slippery and soapy

React with acids to form salt and water

pH is more than 7

Turns red litmus paper to blue.

o Bases are any substance that neutralizes an acid e.g.

Calcium oxide (CaO)

Magnesium carbonate (MgCO )

Most metal oxides and carbonates

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Predict the products of reactions between dilute hydrochloric, nitric and sulfuric acids; and metals, metal oxides and metal carbonates (excluding the reactions between nitric acid and metals)

Method of preparation Reactants Salt formed Other products

Acid + Alkali Sodium Hydroxide + Nitric Acid Sodium Nitrate Water

Acid + Metal Zinc + Hydrochloric Acid Zinc Chloride Hydrogen

Acid + Metal Carbonate Sodium Carbonate + Hydrochloric Acid Sodium Chloride Water and Carbon

Dioxide

Acid + Base Sulphuric Acid + Copper (II) Oxide Copper (II) sulphate Water

Metal + Acid = Salt + Hydrogen

Carbonate + Acid = Salt + Carbon Dioxide + Water

Acid + Base = Salt + Water

Acid + Alkali = Salt + Water

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Recall the general rules for predicting the solubility of salts in water:

i. All common sodium, potassium and ammonium salts are soluble ii. All nitrates are soluble iii. Common chlorides are soluble, except silver chloride iv. Common sulfates are soluble, except those of barium and calcium v. Common carbonates are insoluble, except those of sodium, potassium and

ammonium

Soluble and Insoluble salts

Soluble Insoluble

All Na, K, NH (Ammonium) salts

All nitrates

All chlorides AgCl (Silver chloride)

PbCl (Lead chloride)

All sulphates CaSO (Calcium sulphate)

BaSO (Barium sulphate)

PbSO (Lead Sulphate)

Na, K, NH carbonates

All other carbonates, hydroxides and oxides Na, K, NH hydroxides and oxides

__________________________________________________________________________________________

Describe how to prepare soluble salts from acids

How to make Soluble Salts

Experiment 1

Metal + Acid = Salt + Hydrogen

1. Add metal to acid and stir.

2. Filter to remove excess metal. (Ensure that all acid is used up)

3. Heat to concentrate the solution. (The evaporating dish should be filled to 2/3rds)

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4. Leave to cool and crystallise. For dry crystals: use paper towels, oven or hairdryer, the faster the cooling of the

solution-‐the smaller the crystals and vice-‐versa

5. Test for hydrogen gas: Put a lighted splint into the test-‐tube. The gas will burn with a squeaky pop.

Do not use this method when:

The metal in the salt is very reactive e.g. sodium

The metal does not react with acids e.g. copper or silver

The salt does not dissolve in water

Experiment 2

Carbonate + Acid = Salt + Carbon Dioxide + Water

1. Add carbonate to acid and stir.

2. Filter to remove excess carbonate.

3. Heat to concentrate the solution. (The evaporating dish should be filled to 2/3rds)

4. Leave to cool and crystallise. For dry crystals: use paper towels, oven or hairdryer, the faster the cooling of the

solution-‐the smaller the crystals and vice-‐versa.

5. Test for Carbon Dioxide: Add limewater to the test tube containing the gas. The limewater will turn milky.


The salt does not dissolve in water

The carbonate does dissolve in water

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Experiment 3

Reacting an oxide (insoluble base) with an acid

Insoluble bases are usually oxide or hydroxides of metals

Acid + Base = Salt + Water

1. Add oxide to acid and stir until no more dissolves

2. Filter to remove extra oxide

3. Heat to concentrate the solution

4. Leave to cool and crystallize


If the base is soluble (an alkali)

If the salt is insoluble in water

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Describe how to prepare insoluble salts using precipitation reactions

Precipitation Method

1. Mix solutions until no more precipitate forms

2. Filter to remove precipitate

3. Wash the precipitate

4. Dry the precipitate

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Describe how to carry out acid-alkali titrations.

Acid-‐Alkali Titration

Acid + Alkali = Salt + Water

1. Fill a burette up to the zero mark with acid

2. Use a pipette to place 25cm of alkali in a beaker or conical flask. Add 2 drops of indicator (phenolphthalein) to

the alkali.

3. Run acid from the burette, a little at a time, until the indicator shows that the solution is neutral. It will turn

colourless from pink. Note the volume of acid added.

4. Repeat 1, 2 and 3 without the indicator.

5. Transfer the neutralised solution to an evaporating basin.

6. Heat to concentrate the solution.

7. Leave to crystallise.

Acids and Alkalis are both electrolytes which mean that they are solutions which conduct electricity

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Recall that chemical reactions in which heat energy is given out are described as exothermic and those in which heat energy is taken in are endothermic

All reactions are exothermic (give out heat) in one direction and endothermic (take in heat) in the other

If the temperature is increased o Equilibrium shifts to decrease the temperature o Equilibrium shifts in the endothermic direction

If the temperature is decreased o Equilibrium shifts to increase the temperature o Equilibrium shifts in the exothermic direction

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Describe simple calorimetry experiments for reactions such as combustion, displacement, dissolving and neutralization in which heat energy changes can be calculated from measured temperature changes

Combustion Displacement, Dissolving and Neutralization

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Uendothermic reactions

When chemical reactions occur, as well as the formation of the products -‐ the chemical change, there is also a heat energy change which can often be detected as a temperature change.

This means the products have a different energy content than the original reactants If the products contain less energy than the reactants, heat is released or given out to the surroundings and the

change is called exothermic. The temperature of the system will be observed to rise in an exothermic change. o Examples:

The burning or combustion of hydrocarbon fuels e.g. petrol or candle wax. The burning of magnesium, reaction of magnesium with acids, or the reaction of sodium with

water The neutralisation of acids and alkalis

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Using hydrogen as a fuel in fuel cells If the products contain more energy than the reactants, heat is taken in or absorbed from the surroundings and

the change is called endothermic. If the change can take place spontaneously, the temperature of the reacting system will fall but, as is more likely, the reactants must be heated to speed up the reaction and provide the absorbed heat.

o Examples: The thermal decomposition of limestone The cracking of oil fractions

The difference between the energy levels of the reactants and products gives the overall energy change for the reaction

At a more advanced level the o -‐ve) for exothermic reactions i.e. heat energy is given out and lost from the system to the

surroundings which warm up. o reactions i.e. heat energy is gained by the system and taken in from

the surroundings which cool down OR, as is more likely, the system is heated to provide the energy needed to effect the change

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Represent exothermic and endothermic reactions on a simple energy level diagram

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Recall that the breaking of bonds is endothermic and that the making of bonds is exothermic

Energy is required to break bonds o Therefore, the breaking of bonds is endothermic

Energy is released when bonds are formed o Therefore, the making of bonds are exothermic

In a chemical reaction you need to put energy in to break bonds in the reactants, you get energy out when new bonds are formed to make the products

If you get out more energy than you have to put in, then overall the reaction is exothermic If you have to put in more energy than you get out, then the reaction is endothermic

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Describe experiments to investigate the effects of changes in surface area of a solid, concentration of solutions, temperature and the use of a catalyst on the rate of a reaction

Surface Area of a Solid

In the reaction between calcium carbonate and dilute hydrochloric acid Hydrochloric Acid + Calcium Carbonate Calcium Chloride + Carbon Dioxide + Water Calcium carbonate may be used in the form of marble chips The reaction rates can be compared using large marble chips, and the same mass of small marble chips The reaction can be followed by plotting the loss of mass against time

Concentration of Solutions

In the reaction between sodium thiosulfate solution and dilute hydrochloric acid Hydrochloric Acid + Sodium Thiosulfate Sodium Chloride + Sulfur Dioxide + Sulfur + Water Solid sulfur (S(s)) is formed in the flask Increasing the concentration of sodium thiosulfate means that the solid sulfur will be produced more quickly

and there will be less time before the cross can no longer be seen

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Temperature

In the reaction between sodium thiosulfate solution and dilute hydrochloric acid Hydrochloric Acid + Sodium Thiosulfate Sodium Chloride + Sulfur Dioxide + Sulfur + Water

Solid sulfur (S(s)) is formed in the flask Increasing the temperature of sodium thiosulfate means that the solid sulfur will be produced more quickly and

there will be less time before the cross can no longer be seen

Use of a Catalyst

Hydrogen peroxide is stable at room temperature The presence of a catalyst may cause it to decompose Hydrogen Peroxide Oxygen + Water The rate of the reaction can be followed by recording the volume of oxygen produced The catalyst used is Manganese(IV) oxide -‐ MnO2(s) Using more catalyst will show an increase in reaction rate This is because more catalyst will have a greater surface area for the reaction to take place

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Describe the effects of changes in surface area of a solid, concentration of solutions, pressure of gases, temperature and the use of a catalyst on the rate of a reaction

Activation Energy o The amount of energy needed to start a reaction

Catalyst o A substance that increases the rate of a chemical reaction without being used up



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Concentration o The number of molecules of a substance in a given volume

Enzyme o A biological catalyst

Rate of Reaction o The change in the concentration over a certain period of time

Effect of Surface Area on Rate of Reaction o The larger the surface area, the faster the rate of reaction

Effect of Concentration on Rate of Reaction o The higher the concentration of a dissolved reactant, the faster the rate of reaction

Effect of Pressure on Rate of Reaction o As the pressure of gaseous reactants increases, the rate of reaction increases

Effect of Temperature on Rate of Reaction o The higher the temperature, the faster the rate of reaction

Effect of Catalyst on Rate of Reaction o In the presence of a catalyst, the rate of reaction increases

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Understand the term activation energy and represent it on a reaction profile

Activation Energy o The minimum amount of energy needed for the particles to react o Depends on:

The frequency of collisions between particles The energy with which particles collide

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Explain the effects of changes in surface area of a solid, concentration of solutions, pressure of gases and temperature on the rate of a reaction in terms of particle collision theory

Effect of Surface Area on Rate of Reaction o Any reaction involving a solid can only take place at the surface of a solid

If the solid is split into several pieces, the surface area increases

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This means that there is an increased area for the reactant particles to collide with o The smaller the pieces, the larger the surface area

This means more collisions and a greater chance of reaction

Effect of Concentration on Rate of Reaction o The higher the concentration of a dissolved reactant, the faster the rate of reaction

At a higher concentration, there are more particles in the same amount of space

This means that the particles are more likely to collide and therefore more likely to react

Effect of Pressure on Rate of Reaction o As the pressure of gaseous reactants increases, the rate of reaction increases

As the pressure increases, the space in which the gas particles are moving becomes smaller

The gas particles become close together, increasing the frequency of collisions o This means that the particles are more likely to react

Effect of Temperature on Rate of Reaction o The higher the temperature, the faster the rate of reaction

At higher temperatures, particles have more energy

This means they move faster and are more likely to collide with other particles o When the particles collide, they do so with more energy, and so the number of

successful collisions increases

Effect of Catalyst on Rate of Reaction o Catalysts are substances that change the rate of a reaction without being used up in the reaction o Catalyst never produce more product

They just produce the same amount more quickly

Different catalyenergy

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Understand that a catalyst speeds up a reaction by providing an alternative pathway with lower activation energy

Catalysts increase the rate of a reaction by helping break chemical bonds in reactant molecules This effectively means the activation energy is reduced Therefore at the same temperature, more reactant molecules have enough kinetic energy to react compared to

the uncatalysed situation and so the reaction speeds up with the greater chance of a 'fruitful' collision o A catalyst does NOT change the energy of the molecules, it reduces the threshold kinetic energy needed

for a molecules to react Although a true catalyst does take part in the reaction, it does not get used up and can be reused with more

reactants, it may change chemically on a temporary basis but would be reformed as the reaction products also form

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Recall that some reactions are reversible and are indicated by the symbol in equations

Reversible reactions occur when the backwards reaction (products reactants) takes place relatively easily under certain conditions; the products turn back into the reactants

In some reactions, the products of the reaction can react to reform the original reactants such reactions are reversible

(the symbol in an equation shows it is reversible)

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Describe reversible reactions such as the dehydration of hydrated copper (I I ) sulfate and the effect of heat on ammonium chloride

On heating the blue solid, hydrated copper (II) sulphate, steam is given off and the white solid of anhydrous copper (II) sulphate is formed

When the white solid is cooled and water added, blue hydrated copper (II) sulphate is reformed

o Blue Hydrated Copper (II) Sulphate + Heat White Anhydrous Copper(II) Sulphate + Water

o CuSO4.5H2O(s) CuSO4(s) + 5H2O(g)

The dehydration decomposition to give the white solid is the forward reaction and the 're-‐hydration' to reform the blue crystals is the backward reaction

o The 5H2O in the formula of hydrated copper(II) sulphate is called the water of crystallisation and forms part of the crystal structure when copper(II) sulphate solution is evaporated and crystals form

o This crystal structure is broken down on heating and the water is given off The thermal decomposition is endothermic as heat is absorbed to drive off the water The reverse reaction is exothermic i.e. on adding water to white anhydrous copper (II) sulphate the mixture

heats up as the blue crystals reform The reverse reaction is used as a simple chemical test for water i.e. white anhydrous copper (II) sulphate turns

blue

On heating strongly above 340ºC, the white solid ammonium chloride, thermally decomposes into a mixture of two colourless gases ammonia and hydrogen chloride

On cooling the reaction is reversed and solid ammonium chloride reforms

o Ammonium Chloride + Heat Ammonia + Hydrogen Chloride

o NH4Cl (s) +Heat NH3 (g) + HCl (g)

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The thermal decomposition of ammonium chloride is the forward reaction, and the formation of ammonium chloride is the backward reaction

Reversing the reaction conditions reverses the direction of chemical change, typical of a reversible reaction Thermal decomposition means using 'heat' to 'break down' a molecule into smaller ones. The decomposition is endothermic (heat absorbed or heat taken in) and the formation of ammonium chloride is

exothermic (heat released or heat given out)

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Understand the concept of dynamic equilibrium

A dynamic equilibrium exists when a reversible reaction ceases to change its ratio of reactants/products, but substances move between the chemicals at an equal rate, meaning there is no net change

There will be a mixture of all the reactants and products

Both reactions are still reacting, but at the same rate. So the amount of each substance in the equilibrium will stay the same

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Predict the effects of changing the pressure and temperature on the equilibrium position in reversible reactions

Concentration o If the concentration of the reactants is increased, the equilibrium will move to reduce the increased

concentration of the reactants It does this by moving right and turning reactants into products

o If the concentration of the reactants is reduced, the equilibrium will move to increase the reduced concentration of the reactants

It does this by moving left and turning products into reactants

Temperature o If the temperature is increased, the equilibrium will move to reduce the increased temperature

It does this by moving in the direction of the endothermic reaction o If the temperature is reduced, the equilibrium will move to increase the reduced temperature

It does this by moving in the direction of the exothermic reaction

Pressure o If the pressure is increased, the equilibrium will move to reduce the increased pressure

It does this by moving to the side with less gas molecules o If the pressure is reduced, the equilibrium will move to increase the reduced pressure

It does this by moving to the side with more gas molecules

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Explain how the methods of extraction of the metals in this section are related to their positions in the reactivity series

For metals less reactive than carbon

Extract the metal with oxygen

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o A metal oxide is reacted with charcoal If the charcoal (carbon) is more reactive, it will remove the oxygen from the metal oxide and

leave a trace of metal in the reaction vessel

For metals more reactive than carbon

Extract the metal from its ore using electrolysis o Electricity passes through the melted ore separating the metal from the oxygen

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Describe and explain the extraction of aluminium from purified aluminium oxide by electrolysis, including:

Aluminium is more reactive than carbon and hence is extracted from its ore using electrolysis

The ore of aluminium is called bauxite which is impure aluminium oxide

Bauxite is purified, then dissolved in molten cryolite

Electricity is then passed through the melted ore separating the aluminium from the oxygen

i. The use of molten cryolite as a solvent and to decrease the required operating temperature

Cryolite is an ore of aluminium It lowers the melting point of bauxite from over 2000º C to about 900º C

o This saves time, money and energy

ii. The need to replace the positive electrodes

Aluminium is denser than the alumina/cryolite solution and so it falls to the bottom of the cell where it can be tapped off as pure liquid metal

Oxygen is given off at the positive carbon anode Carbon dioxide is also given off at the carbon anode because hot oxygen reacts with the carbon anode to

form carbon dioxide gas.

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The carbon anodes slowly disappear because each molecule of carbon dioxide which is given off takes a little piece of carbon away with it

The carbon anodes need to be replaced when they become too small

iii. The cost of the electricity as a major factor

The one major cost that makes this process more expensive than the extraction of iron is electricity

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Write ionic half-equations for the reactions at the electrodes in aluminium extraction

Aluminium oxide (Al2O3) is an ionic compound

When it is melted the Al³ and O² ions are free to move and conduct electricity

Electrolysis of the alumina/cryolite solution gives aluminium at the cathode and oxygen at the anode

At the anode

Oxidation

6 O² -‐ 12 e 3 O2

At the cathode

Reduction

Al³ + 12e 4 Al

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Describe and explain the main reactions involved in the extraction of iron from iron ore (haematite), using coke, limestone and air in a blast furnace

The main iron ore is called haematite o Haematite is iron (III) oxide -‐ Fe2O3

The iron ore contains impurities, mainly silica (silicon dioxide) o Limestone (calcium carbonate) is added to the iron ore which reacts with the silica to form molten

calcium silicate in the blast furnace The calcium silicate (called slag) floats on the liquid iron

Since iron is below carbon in the reactivity series, iron in the ore is reduced to iron metal by heating with carbon (coke)

o It is actually carbon monoxide which does the reducing in the blast furnace

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Hot air is blasted into the furnace causing coke (carbon) to burn rapidly and raise the temperature to 2000 °C o Carbon + Oxygen Carbon Dioxide + Heat o C (s) + O2 (g) CO2 (g)

The carbon dioxide then reacts with hot carbon to form carbon monoxide o Carbon Dioxide + Carbon Carbon Monoxide o CO2 (g) + C (s) 2 CO (g)

Carbon monoxide then reduces iron in the ore to iron metal o Carbon Monoxide + Iron (III) Oxide Carbon Dioxide + Iron o 3 CO (g) + Fe2O3 (s) 3 CO2 (g) + 2 Fe (l)

The temperature where the reduction takes place is above 1500 °C Iron falls to the bottom of the furnace where the temperature is 2000 °C Iron is liquid at this temperature and is tapped off periodically

Limestone is calcium carbonate (CaCO3) and it is added to the blast furnace to remove the impurities in the iron ore

Calcium carbonate is decomposed by heat in the furnace to give calcium oxide (quicklime) and carbon dioxide o This is called thermal decomposition

Calcium Carbonate Calcium Oxide + Carbon Dioxide

CaCO3 (s) CaO (s) + CO2 (g)

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The main impurity is silica (sand or rock) which is silicon dioxide o Silicon dioxide is solid at the furnace temperature and the furnace would become blocked if it

was not removed o Silicon dioxide reacts with calcium oxide to form calcium silicate (called slag) which is liquid in

the furnace o Slag flows to the bottom of the furnace where it floats on the liquid iron and is easily removed

Calcium Oxide + Silicon Dioxide Calcium Silicate

CaO (s) + SiO2 (s) CaSiO3 (l)

The slag (CaSiO3) is allowed to cool until it becomes a solid and is used for road construction

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Explain the uses of aluminium and iron, in terms of their properties

Properties and Uses of Iron

Iron is one of the three magnetic elements (the others are cobalt and nickel) Cast iron is very brittle (it cracks easily) but it has a greater resistance to corrosion than either pure iron or steel

o Cast iron is used for manhole covers on roads and pavements and as engine blocks for petrol and diesel engines

Pure iron is called wrought iron o Wrought iron is malleable and is mainly used in ornamental work for gates

Iron is also the catalyst in the Haber Process The large majority of iron from the blast furnace is made into steel

Properties and Uses of Aluminium

Is strong, malleable and has a low density Is resistant to corrosion Is a good conductor of heat and electricity Can be polished to give a highly reflective surface

Low density and strength make it ideal for construction of aircraft, lightweight vehicles, and ladders o An alloy of aluminium called duralumin is often used instead of pure aluminium because of its improved

properties Easy shaping and corrosion resistance make it a good material for drink cans and roofing materials Corrosion resistance and low density leads to its use for greenhouses and window frames Good conduction of heat leads to its use for boilers, cookers and cookware Good conduction of electricity leads to its use for overhead power cables hung from pylons (low density gives it

an advantage over copper) High reflectivity makes it ideal for mirrors, reflectors and heat resistant clothing for fire fighting

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Recall that crude oil is a mixture of hydrocarbons

Crude Oil

Oil as it comes out of the ground before it has been treated in any way

Is a very complicated mixture of chemical compounds, most of which are alkanes

o Is a mixture of hydrocarbons

Obtained from an oil well

Is a black liquid which cannot be used without first being treated in an oil refinery

Is thought to have been made from the remains of marine plants and animals that dies millions of years ago

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Describe how the industrial process of fractional distillation separates crude oil into fractions

Fractional Distillation

Is a process used to separate a mixture of liquids that have different boiling points

o When a mixture is heated, liquids with a low boiling point evaporate and turn to vapour

o Liquids with a higher boiling point remain as liquid

o The vapour can then be separated from the liquid

Fractional Distillation of Crude Oil

Oil is heated to about 450º C and pumped into the bottom of a tall tower called a fractionating column, where it

vapourizes

The column is very hot at the bottom but much cooler at the top

As the vapourised oil rises, it cools and condenses

o Heavy fractions (containing large molecules) have a high boiling point and condense near the bottom of

the column

o Lighter fractions (containing small molecules) have a lower boiling point and condense further up the

column

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Recall the names and uses of the main fractions obtained from crude oil: refinery gases, gasoline, kerosene, diesel, fuel oil and bitumen

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Describe the trend in boiling point and viscosity of the main fractions

Flammability

Ability to catch fire easily

Viscosity

The property by which a liquid resists movement

The runniness of a liquid

Characteristic First Fraction Last Fraction

Boiling Point Low High

Colour Light/Colourless Dark/Black

Flammability High Flammability Low Flammability

Ease of Evaporation High Low

Viscosity Low High

Molecule Size Small Large

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Recall that incomplete combustion of fuels may produce carbon monoxide and explain that carbon monoxide is poisonous because it reduces the capacity of the blood to carry oxygen

Incomplete Combustion

Takes place when there is not enough oxygen present

CH4 + 1.5 O2 CO + 2 H2O

CH4 + O2 C + 2 H2O

Carbon Monoxide

Is a toxic gas because it combines with haemoglobin in the blood and prevent oxygen from reaching the cells

o Thus, it causes respiratory disorders

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Recall that, in car engines, the temperature reached is high enough to allow nitrogen and oxygen from air to react, forming nitrogen oxides

Acid Rain

Caused by sulfur dioxide and nitrogen oxides

o Sulfur oxide is formed by the burning coal, oil and gas

o Nitrogen oxide is formed by the combination of nitrogen and oxygen in cars

This is because the temperature is very high

Hence, nitrogen and oxygen from the air reacts forming nitrogen oxides

Makes the soil and plants weaker

o By leaching out the potassium, magnesium and calcium ions which the plans need for food

Destroys the roots of trees

o By dissolving insoluble compounds like aluminium to release aluminium ions

Kills fish

o Acids rain poisons fish

o

This prevents intake of oxygen causing them to suffocate

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Recall that fractional distillation of crude oil produces more long-chain hydrocarbons than can be used directly and fewer short-chain hydrocarbons than required

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Describe how long-chain alkanes are converted to alkenes and shorter-chain alkanes by catalytic cracking, using silica or alumina as the catalyst and a temperature in the range of 600-700 °C

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Catalytic Cracking

The breaking down of large alkane molecules into smaller molecules

o Alkanes and alkenes are formed

The alkene molecules are more useful than the alkane molecules

Requires a high temperature and a catalyst

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Recall that an addition polymer is formed by joining up many small molecules called monomers

Polymers

Are very large molecules made up of thousands of small molecules called monomers joined together

E.g.

o Plastics

o Synthetic fibres

Polymerisation

The chemical reaction where monomers link forming polymers

Addition Polymers

Unsaturated molecules can join together to make polymers

During polymerization, the double bonds in the molecules open up, and immediately join with neighbouring

molecules to form a long chain with only single bonds

For this, high pressure and/or a catalyst is used to split the double bond

Repeating Unit

Is the monomer without its double bond

Polymerization Reaction of Polythene

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Draw the repeat unit of addition polymers, including poly(ethene), poly(propene) and poly(chloroethene)

Monomer Name

Ethene

Monomer Molecular Formula

C2H4

Monomer Structure

Polymer Name

Polythene

Polymer Structure

Repeating Unit

Monomer Name

Propene


C3H6

Monomer Structure

Polymer Name

Polypropene

Polymer Structure

Repeating Unit

Monomer Name

Chloroethene or Vinyl Chloride


C2H3Cl

Monomer Structure

Polymer Name

Polychloroethene or PVC

Polymer Structure

Repeating Unit

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Deduce the structure of a monomer from the repeat unit of an addition polymer

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Recall that nitrogen from air, and hydrogen from natural gas or the cracking of hydrocarbons, are used in the manufacture of ammonia

Some of the nitrogen and hydrogen react to form ammonia

At the same time, some of the ammonia breaks down into nitrogen and hydrogen

Nitrogen + Hydrogen Ammonia

N2 (g) + 3 H2 (g) NH3

The hydrogen is obtained from natural gas or the cracking of hydrocarbons

Nitrogen is obtained from the air

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Describe the manufacture of ammonia by the Haber process, including the essential conditions:

Industrial Conditions

High pressure (200 atmospheres)

Quite high temperature (450ºC)

An iron catalyst

i. A temperature of about 450 °C This is an exothermic reaction (gives out heat)

Monomer Repeating Unit

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A low temperature would actually increase the amount or yield of ammonia, but it would be at the cost of a much slower rate of reaction

A higher temperature would have a faster rate of reaction but, unfortunately, the amount of ammonia produced would be lower

In practice, a compromise temperature of 450ºC is used o This gives a reasonable yield reasonably quickly

Any nitrogen or hydrogen that has not been converted into ammonia can be recycled to reduce costs

ii. A pressure of about 200 atmospheres A high pressure us used because it increases the amount of ammonia made

There are four gas molecules on the left-‐hand side of the equation; one nitrogen and three hydrogen molecules

There are only two ammonia molecules on the right-‐hand side

Increasing the pressure encourages the forward reaction which increases the amount of ammonia because there are fewer molecules on the right-‐hand side of the equation

Ideally, the highest possible pressure should be used; however it is too expensive to build a plant which can withstand pressures greater than 200 atmospheres

iii. An iron catalyst An iron catalyst is used to speed up the rate of reaction and so reduce the cost of producing ammonia

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Understand how the cooling of the reaction mixture liquefies the ammonia produced and allows the unused hydrogen and nitrogen to be recirculated

When the gases leave the reactor they are hot and at a very high pressure

Ammonia is easily liquefied under pressure as long as it isn't too hot, and so the temperature of the mixture is lowered enough for the ammonia to turn to a liquid

The nitrogen and hydrogen remain as gases even under these high pressures, and can be recycled

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Recall the use of ammonia in the manufacture of nitric acid and fertilizers

Ammonia can be oxidized to produce nitric acid

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Ammonia gas reacts with oxygen in the air over a hot platinum catalyst o 4 NH3 (g) + 5 O2 (g) 4 NO (g) + 6 H2O (g)

The nitrogen oxide is cooled, and then reacted with water and more oxygen to form nitric acid o 4 NO (g) + 3 O2 (g) + 2 H2O (l) 4 HNO3 (g)

The nitric acid can be neutralized with ammonia to make ammonium nitrate

Ammonia can also be reacted with sulphuric acid to make ammonium sulphate

Ammonium nitrate and ammonium sulphate are popular fertilizers

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