C2

 1.1 - Chemical Bonding 

• Elements react to form compounds by gaining or losing electrons, or by sharing electrons
• Atoms of metals in Group 1 combine with atoms of non-metals in Group 7 by transferring electrons to form ions that have the electronic structures of noble gases.

1.2 - Ionic Bonding

• Ionic compounds are held together by strong forces between the oppositely charged ions. This is called ionic bonding.
• The ions form a giant structure or lattice. The strong forces of attraction act throughout the lattice.
• We can represent atoms and ions using dot and cross diagrams.

1.3 - Formulae of ionic compounds

• The charges on the ions in an ionic compund always cancel each other out. 
• The formula of an ionic compound shows the ratio of ions present in the compound.
• Sometimes we need brackets to show the ratio of ions in a compound, e.g. magnesium hydroxide, Mg(OH)2. 

1.4 - Covalent bonding

• A covalent bond is formed when two atoms share a pair of electrons.
• The number of covalent bonds an atom forms depends on the number of electrons it needs to achieve a stable electronic structure.
• Many substances containing covalent bonds consist of simple molecules, but some have giant covalent structures.

1.5 - Metals

• The atoms in metals are closely packed together and arranged in regular layers.
• The electrons in the highest energy level are delocalised. the strong electrostatc forces between these electrons and the positively charged metal ions hold the metal together. (H)

2.1 - Giant ionic structures

• Ionic compounds have high melting points and they are all solids at room temperature. 
• Ionic compounds will conduct electricity when we melt them or dissolve them in water. Their ions can then move freely and can carry charge through the liquid.

2.2 - Simple Molecules 

• Substances made up of simple molecules have low melting points and boiling points.
• simple molecules have no overall charge, so they cannot carry electrical charge and do not conduct electricity.
• The weak intermolecule forces between simple molecules are why substances made of simple molecules have low melting and boiling points. 

2.3 - Giant covalent structures

• Covalently bonded substances with giant structures have very high melting points.
• Diamond is a form of carbon whose atoms each form four covalent bonds.
• Graphite is another form of carbon where the carbon atoms form layers that can slide over each other.
• Graphite can conduct electricity because of the delocalised electrons in its structure. (H)
• Carbon also exists as fullerenes. (H)

Fullerenes are large molecules formd from hexagonal rings of carbon atoms. The rings join together to form cage-like shapes with different numbers of carbons atoms, some of which are nano-sized.  Scientists are finding many applications for fullerenes, including drug delivery into the body, lubricants, catalysts and reinforcing materials.

2.4 - Giant metallic structures

• When we bend and shape metals the layers of atoms in the giant metallic structure slide over each other.
• Alloys are mixturesof metals and are harder and than pure metals because the layers in the structures are disorted.
• If a shape-memory alloy, is deformed, it can return to its original shape on heating.
• Delocalised electrons in metals can enable metals to conduct heat and electricity well. (H)

2.5 - The properties of polymers

• The properties of polymers depend on the monomerrs used to make them.
• Changing reaction conditions can also change the properties of the polymer that is produced.
• Thermosoftening polymers soften or melt easily when heated.
• Thermosetting polymers do not soften or melt when heated.

2.6 - Nanoscience

• Nanoscience is the study of small particles that are between 1 and 100 nanometres in size.
• Nanoparticles behave differently from the bulk materials they are made from.
• Developments in nanoscience are exciting but will need more research into possible issues that might arise from increased use.

3.1 The mass of atoms

• The relative mass of protons and neutrons is 1.
• The atomic number of an atom is its number of protons (which equals its number of electrons).
• The mass number of an atom is the total number of protons and neutrons in its nucleus.
• Isotopes are atoms of the same element with different numbers of neutrons.

3.2 Masses of atoms and moles

• We use relative atomic masses to compare the masses of atoms.
• The relative atomic mass of an element is an average value for the isotopes of an element.
• We work out the relative formula mass of a compound by adding up the relative atomic masses of the elements in it.
• One mole of any substance is its relative formula mass in grams.

3.3 Percentages and formulae

• The relative atomic masses of the elements in a compound and its formula can be used to work out its percentage composition.
• We can calculate empirical formulae given the masses or percentage composition of elements present.

3.4 Equations and calculations

• Balanced symbol equations tell us the number of moles of subtances involved in a chemical reaction.
• We can use balanced symbol equations to calculate the masses of reactants and products in a chemical reaction.

3.5 The yield of a chemical reaction

• The yield of a chemical reaction describe how much product is made.
• The percentage yield of a chemical reaction tells us how much product is made compared with the maximum amount that could be made.
• It is important to maximise yield and minimise energy wasted to conserve the Earth's limited resources and reduce pollution.

3.6 Reversible reactions

• In a reversible reaction the products of the reaction can react to make the original reacants.
• We can show a reversible reaction using the (reversible reaction) sign >/<

3.7 Analysing substances

• Chemical analysis is used to identify food additives.
• Paper chromotography can be used to detect and identify artificial colours.

3.8 Instrumental analysis

• Modern istrumental techniques provide fast, accurate an sensitive ways of analysing chemical substances.
• Compounds in a mixture can be separated using gas chromatography.
• Once separated, compounds can be identified using a mass spectrometer.
• The mass spectrometer can be used to find the relative molecular mass of a compound from its molecular ion peak.

4.1 How fast?

• We can find the rate of a chemical reaction by measuring the amount of reactants used up over time or by measuring the amount of products made over time.
• The gradient or slope of the line on a graph of amount of reactant or product against time tells us the rate of reaction at that time. The steep the gradient, the faster the reaction.

4.2 Collision theory and surface area

• Particles must collide with a certain amount of energy before they can react.
• The minimum amount of energy that particles must have in order to react is called the activation energy.
• The rate of a chemical reaction increases if the surface area of any solid reatants is increased. This icreases the frequency of collisions between reacting particles.

4.3 The effect of temperature

• Reactions happen more quickly as the temperature increases.
• Increasing the temperature increases the rate of reaction because partiles collide more frequently and more energetically.
• At a higher temperature more of the collisions result in a reaction because a higher proportion of particles have energy greater than the activation energy.

4.4 The effect of concentration or presure.

• Increasing the concentration of reactants in solutions increases the frequency of collisions between particles, and so increases the rate of reaction.
• Increasing the pressure of reacting gases also increases the frequency of collisions and so increases the rate of reaction.

4.5 The effect of catalysts

• A catalyst speeds up the rate of a chemical reaction.
• A catalyst is not used up during a chemical reaction.
• Different catalysts are needed for different reactions.

4.6 Catalysts in action

• Catalysts are used in industry to increase the rate of reactions and reduce energy costs.
• Traditional catalysts are often transition metals or their compounds.
• Modern catalysts are being developed in industry which result in less waste and are safer for the environment.

4.7 Exothermic and endothermic reactions

• Energy may be tranfrred to or from the reacting substances in a reaction.
• A reaction in which energy is transferred from the reacting substances to their surroundings is called an exothermic reaction.
• A reaction in which energy is tranferred to the reacting substances from their surroundings is called an endothermic reaction.

4.8 Energy and reversible reactions

• In reversible reactions, the reaction in one direction is exothermic and in the other direction it is endothermic.
• In any reversible reaction, the amount of energy released when the reaction goes in one direction is exactly equal to the energy absorbed when the reaction goes in th opposite direction.

4.9 Using energy transfers from reactions.

• Exothermic changes can be used in hand warmers and self-heating cans.
• Endothermic changes can be used in instant cold packs for sports injuries.

5.1 Acids and alkalis

• When acids are added to water they produce hydrogen ions, H+(aq), in the solution.
• Bases are substances that will neuralise acids.
• Alkali dissolve in water to give hydroxide ions, OH-(aq), in the solution.
• The pH scale shows how acidic or alkaline a solution is.

5.2 Making salts from metals or bases

• When an acid reacts with a base a neutralisation reaction takes place and produces a salt and water.
• Some salts can be made by the reaction of a metal with an acid. This reaction produces hydrogen gas as well as a salt.
• Salts can be crystallised from solutions by evaporating off water.

5.3 Making salts from solutions

• When a soluble salt is made from an alkali and an acid, an indicaor can be used to show when the reaction is complete.
• Insoluble salts can be made by reacting two solutions to produce a precipitate.
• Precipitation is an important way of removing some metal ions from industrial waste water.

5.4 Electrolysis

• Electrolysis splits up a substance using electricity.
• Ionic compounds can only be electrolysed when they are molten or in solution because then their ions are free to move to the electrodes.
• In electrolysis, positive ions move to the negative electrode and negative ions move to the positive electrode.

5.5 Changes at the electrodes

• Negative ions lose electrons and so are oxidised at the positive electrode.
• Positive ions gain electrons and so are reduced at the negative electrode.
• When aqueous solutions are electrolysed, oxygen gas is produced at the positive electrode unless the solution contains halide ions.
• When aqueous solutions are electrolysed, hydrogen gas is produced at the negative electrode unless the solution contains ions of a metal that is less reactive than hydrogen. 

5.6 The extraction of aluminium

• Aluminium oxide is electrolysed to manufacture aluminium.
• The aluminium oxide is mixed with the molten cryolite to lower its melting point.
• Aluminium forms at the negative electrode and oxygen at the positive lectrode.
• The positive carbon electrodes are replaced regularly as they gradually burn away.

5.7 Electrolysis of brine

• When we electrolyse brine we get three products - chlorine gas, hydrogen gas and sodium hydroxide solution.
• The products are important reactants used in industry.

5.8 Electroplating

• We can electroplate objects to improve their appearance, protect their surface and to use smaller amounts of precious metals.
• The object to be electroplated is made the negative electrode in an electrolysis cell. The plating metal is made the positive electrode. The electrolyte contain ions of the plating metal.

1.1 The early periodic table

• The periodic table of the elements developed as chemists tried to classify the elements. It arranges them in a pattern in which similar elements are grouped together.
• Newlands' table put the elements in order of atomic weight but failed to take account of elements that were unknown at that time.
• Mendeleev's periodic table left gaps for the unknown elements, and so provided the basis for the modern periodic table.

1.2 The modern periodic table

• The atomic (proton) number of an element determines its position in the periodic table.
• The number of electrons in the outermost shell (highest energy level) of an atom determines its chemical properties.
• The group number in the periodic table quals the number of electrons in the outermost shell.
• We can explain trends in reactivity as we go down a group in terms of:
• the distance between the outermost electrons and the nuceus
• the number of occupied inner shells (energy levels) in the atoms.

1.3 Group 1 - the alkali metals

• The elements in Group  of the periodic table are called the alkali metals.
• These metls all react with water to produce hydrogen and an alkaline solution containing the metal hydroxide.
• They form positive ions with a charge of 1+ in reactions to make ionic compounds. Their compounds are usually white or colourless crystals that dissolve in water producing colourless solutions.
• The reactivity of the alkali metal increases going down the group.

1.4 The transition elements

•  Compared with the alkali metals, transition elements have much higher melting points nd densities. They are also stronger and harder, but are much less reactive.
• The transition elements do not react vigourously with oxygen or water.
• Transition elements can form ions with different charges, in compounds that are often coloured.
• Transition elements and their compounds are important industrial catalysts.

1.5 Group 7 - the halogens

• The halogens all form ions with a single negative charge in their ionic compounds with metals.
• The halogens form covalent compounds by sharing electrons with other non-metals.
• A more reactive halogen can dosplace a less reactive halogen from a solution of one of its salts.
• The reactiviy of the halogens decreases going down the group.

2.1 Hard water

• Hard water contains dissolved compounds such as calcium and magnesium salts.
• The calcium and/or magnesium ions in hard water react with soap producing a precipitate called ***.
• Temporary hard water can produce a solid scale when it is heated, reducing the efficiency of heating systems and kettles.
• Hard water is better than soft water for developing and maintaining teeth and bones. It may also help to prevent heart disease.

2.2 Removing hardness

• Soft water does not contain salts that produce s*** or scale.
• Hard water can be softened by removing the salts that produce s*** and scale.
• Temporary hardness is removed from water by heating it. Permenant hardness is not changed by heating.
• The hydrogencarbonate ions in temporary hard water decompose on heating. The carbonate ions formed react with Ca^+(aq) and Mg^+(aq), making precipitates.
• Both types of hard water can be softened by adding washing soda or by using an ion-echange resin to remove calcium and magnesium ions.

2.3 Water treatment

• Water for drinking should contain only low levels of dissolved substances and microbes.
• Water is made fit to drink by filtering it to remove solids and adding chlorine to kill microbes.
• We can make pure water by distillation but this requires large amounts of energy which makes it expensive.

2.4 Water issues

• There are advantages and disadvantages to any type of water treatment.
• Water can be treated to remove hardness, to remove harmful microbes and to improve dental health.

3.1 Comparing the energy released by fuels.

• When fuel and food react with oxygen, energy is released in an exothermic reaction.
• A simple calorimeter can be used to compare the energy released by different fuels or different foods in a school lab.

3.2 Energy transfers in solutions

• We can calculate the energy chage for reactions in solution by measuring the temperature change and using the equation Q=mc^T
• Neutralisation and displacement reactions are both examples of reactions that we can use this technique for.

3.3 Energy level diagrams

• We can show the relative differenece in the energy of reactants nd products on energy level diagrams.
• Catalyst provide a pathway with a lower activation energy so the rate of reaction increases.
• Bond breaking is endothermic and bond making is exothermic.

3.4 Calculations using bond energies

• In an exothermic reaction, the energy released when new bonds are formed is greater than the energy absorbed when bonds are broken.
• In an endothermic reaction, the energy released when new bonds are formed is less than the energy absorbed when bonds are broken.
• We can calculate the overall energy change in a chemical reaction using bond energies.

3.5 Fuel issues

• Much of the world relies on fossil fuels. However, they are non-renewable and they cause pollution. Alternative fuels need to be found soon.
• Hydrogen is one alternative. It can be burned in combustion engines or used in fuel cells to poer vehicles.

4.1 tests for positive ions

• Most Group 1 and Group 2 metal ions can be identified using flame tests.
• Sodium hydroxide solution can be used to identify different metal ions, depending on the precipitate that is formed.

4.2 Tests for negative ions

• We identify carbonates by adding dilute acid, which produces carbon dioxide gas. The gas turns limewater cloudy.
• we identify halides by adding nitric acid, then silver nitrate solution. This produces a precipitate of silver halide (chlorine = white, bromide = cream, iodide = pale yellow).
• we identify sulfates by adding hydrochloric acid, then barium chloride solution. This produces a white precipitate of barium sulfate.

4.3 Titrations

• A titration is used to measure accurately how much acid and alkali react together completely.
• The point at which an acid-base reaction is complete is called the end point of the reaction.
• We use an indicator to show the end point of the reaction between an acid and an alkali.

4.4 Titration calculations

• Concentrations of solutions can be measured in g/dm^3 or mol/dm^3
• Concentrations can be calculated from the mass of solute dissolved in a known volume of solution.
• The mass of solute in any volume of solution can be calculated from its concentration.
• If the concentration of one of the solutions used in a titration is known, the results of the titration can be used to calculate the concentration of the other solution.

4.5 Chemical analysis

• Scientists working in environmental monitoring, medicine and forensic science all need to analyse substances.
• The results of their analysis are often matched against existing databases to identify substances (or suspects in the case of forensics).

4.6 Chemical equilibrium

• In a reversible reaction the products of the reaction can react to re-form the original reactants.
• In a closed system, equilibriumis achieved when the rates of the forward and reverse reactions are equal.
• Changing the reaction conditions can change the amounts of products and reactants in a reaction mixture at equilibrium.

4.7 Altering conditions

• Changing the pressure can affect reversible reactions involving gases at equilibrium.

     - Increasing the pressure favours the reaction with the smaller number of molecules of gas              formed.

     - Decreasing the pressure favours the reaction with the larger number of molecules of gas                formed.

• Changing the temperature at which we carry out a reversible reaction can change the amount of products formed at equilibrium.

     - Increasing the temperature favours the endothermic reaction.

     - Decreasing the temperature favours the exothermic reaction.

4.8 Making ammonia - the Haber process

• Ammonia is an important chemical for making other chemicals, including fertilisers.
• Ammonia is made from nitrogen and hydrogen in the Haber process.
• The Haber pricess is done using coniions which are chosen to give a reasonable yield of ammonia as quickly as possible.
• Any unreacted nitrogen and hydrogen are recycled in the Haber process.

4.9 The economics of the Haber process

• The Haber process uses a pressure of around 200 atmospheres to increase the amount of ammonia produced.
• Although higher presure would produce more ammonia, they would make the chemical plant too expensive to build and run.
• A temperature of about 450 C is used for the reaction. Although lower temperatures would increase the yield of ammonia, it would be produced too slowly.

5.1 Structures of alcohols, carboxylic acids and esters.

• The homologous series of alcohols contain the -OH functional group.
• The homologous series of carboxylic acids contain the -COOH functional group.
• The homolgous series of esters contains the -COO- functional group.

5.2 Properties and uses of alcohols

• Alcohols are used as solvents and fuels, and ethanol is the main alcohol in alcoholic drinks.
• Alcohols burn in air, forming carbon dioxide and water.
• Alcohols react with sodium to form a solution and give off hydrogen gas.
• Ethanol can be oxidised to ethanoic acid, either by chemical oxidising agents or by the action of microbes. Ethanoic acid is the main acid in vinegar.

5.3 Carboxylic acid and esters.

• Solutions of carboxylic acids have a pH value less than 7. Their acidic solutions react with carbonates, gently fizzing as they release carbon dioxide gas.
• Aqueous solutions of weak acids have a higher pH value than solutions of strong acids with the same concentration.
• Esters are made by reacting a carboxylic acid and an alcohol together with an acid catalyst.
• Esters are volatile compounds used in flavourings and perfumes.

5.4 Organic issues

• Alcohols, carboxylic acids and esters have many uses which benefit society.
• However, some of these substances, such as ethanol and solvents, can be abused.
• In future, the use of biofuels, such as ethanol and esters, could help society as crude oil supplies run out.
• However, future uses of biofuels might conflict with the need to feed the world.