Chemistry Paper 1 AQA NEW SPEC

All substances are made of atoms. An atom is the smallest part of an element that can exist. Atoms of each element are represented by a chemical symbol, eg O represents an atom of oxygen, Na represents an atom of sodium. Atoms are made of a tiny nucleus, surrounded by shells of electrons.

There are about 100 different elements. Elements are shown in the periodic table. Groups in the periodic table are vertical, and periods are horizontal.

Compounds are formed from elements by chemical reactions. Chemical reactions always involve the formation of one or more new substances, and often involve a detectable energy change. Compounds contain two or more elements chemically combined in fixed proportions and can be represented by formulae using the symbols of the atoms from which they were formed. Compounds can only be separated into elements by chemical reactions.

A mixture consists of two or more elements or compounds not chemically combined together. The chemical properties of each substance in the mixture are unchanged.

Mixtures can be separated by physical processes such as filtration, crystallisation, simple distillation, fractional distillation and chromatography. These physical processes do not involve chemical reactions and no new substances are made.

• Filtration separates insoluble solids from liquids, it can be used in purification as well

1) Filter paper is folded into a cone shape and is placed in the filter funnel. 

2) Pour the mixture into the beaker. The insoluble solid would be left on the filter paper because the particles are too big to pass through the filter paper and only the liquid would pass through.

• It separates soluble solids from solutions
• To obtain a sample of pure salt, sodium chloride, from the salt solution following filtration, you would need to separate the sodium chloride in the solution (called the filtrate) from the water. You can do this by evaporating the water from the sodium chloride solution. 
• The best way to do this is by heating it in an evaporating dish on a water bath. Using a water bath is a gentler way of heating than heating the evaporating dish directly on a tripod and gauze.
• Heating should be stopped at the point of crystallisation - when small crystals first appear around the edge of the solution or when crystals appear in a drop of solution extracted from the dish with a glass rod. The rest of the water is then left to evaporate off the saturated solution at room temperature to get a good sample of sodium chloride crsytals. A flat-bottomed crystallisation dish or Petri dish can be used for this final step, to give a large surface area for the water to evaporate from.

• Crystallisation separates a solube solid from a solvent but sometimes you need to collect the solvent itself instead of just letting in evaporate off into the air. For eg. some countries with a lack of fresh water sources purify seawater to obtain usable water.

• A solution is heated and boiled to evaporate the solvent
• The vapour given off then enters a condenser - an outer glass tube with water flowing through it that acts as a cooling 'jacket' around the inner glass tube from the flask
• Here, the hot vapour is cooled and condensed back into a liquid for collection in a recieving vessel and condensed back into a liquid for collection in a receiving vessel. 
• Any dissolved solids will remain in the heated flask

• It is difficult to get pure liquids from mixtures of liquids with similar boiling points by simple distillation, as vapour is given off from each liquid before they actually reach their boiling point. 
• To aid separation, you can add a fractionating column to the apparatus for distillation. This is usually a tall glass column filled with glass beads, fitted vertically on top of the flask being heated.
• The vapours must pass over and between the glass beads in the fractionating column before they reach the condenser. The temperature in the fractionating column is highest at the bottom of the column, getting lower as the vapours rise up.
• The substance with the higher boiling point will condense more readily on the cooler glass beads nearer the bottom of the column and drip back down into the flask beneath
• The substance with the lower boiling point will continue rising and pass over the condenser where it is cool enough to turn back into the liquid state and be collected.
• The boiling point of ethanol is 78°C and water is 100°C. So if the temperature reading can be kept at 80°C, the liquid collected will mainly be ethanol. You can test the difference between the starting mixture of ethanol and water and the distillate collected by applying a lighted splint to a small volume of each in an evaporating dish. Ethanol is a flammable liquid but is not flammable when mixed with an excess of water. The distillate will ignite when a flame is applied, as the ethanol collected should only have a small amount of water present. It burns with a clear blue flame.

• To separate ethanol from a fermented mixture in the alcoholic spirits industry and in the use of ethanol as a biofuel.

New experimental evidence may lead to a scientific model being changed or replaced.

• Before the discovery of the electron, atoms were thought to be tiny spheres that could not be divided.

• The discovery of the electron led to the plum pudding model of the atom. The plum pudding model suggested that the atom is a ball of positive charge with negative electrons embedded in it.
• Thomson did experiments on the beams of particles. They were attracted to a positive charge, showing they must be negatively charged themselves and called them electrons. These electrons must have come from inside atoms in the tub, so Dalton's idea that atoms could not be divided or split had to be revised.
• He knew that atoms themselves carry no overall charge, so any charges in an atom must balance out.

• The results from the alpha particle scattering experiment led to the conclusion that the mass of an atom was concentrated at the centre (nucleus) and that the nucleus was charged. This nuclear model replaced the plum pudding model.
• Geiger and Marsden were doing an experiment with radioactive particles. They were firing dense, positively charged particles - alpha particles - at the thinnest piece of gold foild they could make.
• They expected the particles to pass straight through the gold atoms with their diffuse cloud of positive charge as in Thompson's plum pudding model, but they didn't.
• Rutherford suggest that Thompson's atomic model was not possible. The positive charge must be concentrated at a tiny spot in the centre of the atom.
• Otherwise, the large, positive particles fired at the foil could never be replled back to their source.
• It was proposed that the electrons must be orbiting around this nucleus - centre of the atom - which contains very dense positively charged protons.
• Rutherford and the others gave the conclusion that the nucleus could be divided into smaller particles, each of which has the same charge as a hydrogen nucleus.

• He noticed that the light given out when atoms were heated only had specific amounts of energy.
• He suggested that the electrons must be orbiting the nucleus at set distances, in fixed energy levels or shells.
• The energy must be given out when excited electrons fall from a high to a low energy level.
• Bohr matched his model to the energy values observed.
• Niels Bohr adapted the nuclear model by suggesting that electrons orbit the nucleus at specific distances. The theoretical calculations of Bohr agreed with experimental observations.

• Scientists at the time speculated that there were two types of sub-atomic particles inside the nucleus. They had evidence of protons but a second sub-atomic particle in the nucleus was also proposed to explain the missing mass that had been noticed in atoms.
• These neutrons must have no charge and have the same mass as a proton.
• Because neutrons have no charge, it was very difficult to detect them in experiments. It was not until 1932 that James Chadwick did an experiment that could only be explained by the existence of neutrons.
• The experimental work of James Chadwick provided the evidence to show the existence of neutrons within the nucleus. This was about 20 years after the nucleus became an accepted scientific idea.

Plum pudding model:

• ball of positive charge (spread throughout)
• electrons spread throughout (embedded in the ball of positive charge)
• no empty space in the atom
• mass spread throughout

Nuclear model: 

• positive charge concentrated at the centre
• electrons outside the nucleus
• most of the atom is empty space
• mass concentrated at the centre

Similarities: 

• both have positive charges
• both have (negative) electrons
• neither has neutrons

• In an atom, the number of electrons is equal to the number of protons in the nucleus.

• Proton - +1
• Neutron - 0
• Electron - -1

Proton number

Protons and neutrons.

• Atoms are very small, having a radius of about 0.1 nm (1 x 10-10 m).
• The radius of a nucleus is less than 1/10 000 of that of the atom (about 1 x 10-14 m).
• Almost all of the mass of an atom is in the nucleus.

• Proton - 1
• Neutron - 1
• Electron - Very small

• Atoms of the same element can have different numbers of neutrons and the same number of protons; these atoms are called isotopes of that element.

RAM = sum of (isotope abundance x isotope mass number) / sum of abundances of all of the isotopes.

• The relative atomic mass of an element is an average value that takes account of the abundance of the isotopes of the element.  

• 1st shell : 2
• 2nd shell: 8
• 3rd shell: 8

The elements in the periodic table are arranged in order of atomic (proton) number and so that elements with similar properties are in columns, known as groups. The table is called a periodic table because similar properties occur at regular intervals. Elements in the same group in the periodic table have the same number of electrons in their outer shell (outer electrons) and this gives them similar chemical properties.

• Before the discovery of protons, neutrons and electrons, scientists attempted to classify the elements by arranging them in order of their atomic weights. 
• The early periodic tables were incomplete and some elements were placed in inappropriate groups if the strict order of atomic weights was followed.

Mendeleev overcame some of the problems by leaving gaps for elements that he thought had not been discovered and in some places changed the order based on atomic weights. Elements with properties predicted by Mendeleev were discovered and filled the gaps. Knowledge of isotopes made it possible to explain why the order based on atomic weights was not always correct. 

• both tables have more than one element in a box
• both have similar elements in the same column
• both are missing the noble gases
• both arranged elements in order of atomic weight

• Newlands did not leave gaps for undiscovered elements
• Newlands had many more dissimilar elements in a column
• Mendeleev left gaps for undiscovered elements
• Mendeleev changed the order of some elements (eg Te and I) 

• Mendeleev predicted properties of missing elements
• Elements with properties predicted by Mendeleev were discovered
• Mendeleev’s predictions turned out to be correct
• Elements were discovered which fitted the gaps

Atoms with a charge.

Elements that react to form positive ions are metals.

Elements that do not form positive ions are non-metals.

The majority of elements are metals. Metals are found to the left and towards the bottom of the periodic table. Non-metals are found towards the right and top of the periodic table.

The number of electrons in their outermost shells.

The number of shells an element has.

The elements in Group 0 of the periodic table are called the noble gases. They are unreactive and do not easily form molecules because their atoms have stable arrangements of electrons. The noble gases have eight electrons in their outer shell, except for helium, which has only two electrons. They are unreactive non-metal gases.

The boiling points of the noble gases increase with increasing relative atomic mass (going down the group).

The elements in Group 1 of the periodic table are known as the alkali metals and have characteristic properties because of the single electron in their outer shell. 

In Group 1, the reactivity of the elements increases going down the group. This is because all of the elements in group 1 have 1 electron in its outermost shell. Due to electron shielding and the weak electrostatic force of attraction between the positive nucleus and the negative outermost electron, it would be easier to lose an electron. Going down group 1, the reactivity increases because the number of shells increases which continues to weaken the electrostatic force of attraction between the nucleus and outermost electron, therefore making it easier to lose an electron. The more shells of electrons a group 1 metal has, the more reactive it is. This means that group 1 metals become more reactive down the group. The melting and boiling points decrease down the group.

• When lithium is added to water, lithium floats. It fizzes steadily and becomes smaller, until it eventually disappears.

• lithium + water → lithium hydroxide + hydrogen

  2Li(s) + 2H₂O(l) → 2LiOH(aq) + H₂(g)

• When sodium is added to water, the sodium melts to form a ball that moves around on the surface. It fizzes rapidly, and the hydrogen produced may burn with an orange flame before the sodium disappears.

• sodium + water → sodium hydroxide + hydrogen

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

• When potassium is added to water, the metal melts and floats. It moves around very quickly on the surface of the water. The hydrogen ignites instantly. The metal is also set on fire, with sparks and a lilac flame. There is sometimes a small explosion at the end of the reaction.

• potassium + water → potassium hydroxide + hydrogen

  2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g)

• The hydroxides formed in all of these reactions - lithium, sodium and potassium with water - dissolve in water to form alkaline solutions. These solutions turn universal indicator purple, showing they are strongly alkaline. Strong alkalis are corrosive, so care must be taken when they are used - for example, by using goggles and gloves.

• If a piece of hot lithium is lowered into a jar of chlorine, white powder is produced and settles on the sides of the jar. This is the salt lithium chloride.

• lithium + chlorine → lithium chloride

  2Li(s) + Cl₂(g) → 2LiCl(s)

• If a piece of hot sodium is lowered into a jar of chlorine, the sodium burns with a bright yellow flame. Clouds of white powder are produced and settle on the sides of the jar. This is the salt sodium chloride.

• The reaction of sodium with chlorine is similar to the reaction with lithium, but more vigorous.

• sodium + chlorine → sodium chloride

  2Na(s) + Cl₂(g) → 2NaCl(s)

• Potassium reacts more violently with chlorine than sodium does.

• potassium + chlorine → potassium chloride

  2K(s) + Cl₂(g) → 2KCl(s)

• Lithium, sodium and potassium are easily cut with a blade.
• The freshly cut surfaces are silvery and shiny, but quickly turn dull as the metal reacts with oxygen in the air.
• The group 1 metals react vigorously with oxygen to form metal oxides. This forms a layer of oxides on the shiny surface.
• Lithium burns with a red flame, sodium with a yellow-orange flame, and potassium burns with a lilac flame.

• The reaction between an alkali metal and water also produces a metal hydroxide.
• They must be stored under oil to keep air and water away from them. group 1 elements form alkaline solutions when they react with water, which is why they are called alkali metals.

• The elements in Group 7 of the periodic table are known as the halogens and have similar reactions because they all have seven electrons in their outer shell.
• The halogens are non-metals and consist of molecules made of pairs of atoms.

• Group 1 elements don't need much energy to lose their one outer electron to form a full outer shell, so they readily form +1 ions.
• They form ionic compounds which are usually white solids that dissolve in water to form colourless solutions.

• They can form molecular compounds
• They can share electrons via covalent bonding with other non-metals as a way to achieve a full outermost shell
• The compounds that form when halogens react with non-metals all have simple molecular structures

• Flourine - poisonous yellow gas
• Chlorine - green gas - sterilising water
• Bromine - red-brown volatile liquid/orange liquid - making pesticides and plastics
• Iodine - dark grey crystalline solid or a purple vapour/grey solid - sterilising wounds - can form a purple vapour when warmed.

• They are all diatomic which means that the two atoms are chemically bonded together.

• Halides form when a group 1 metal reacts with a group 7 non-metal.
• This is because halides have eight electrons in their outermost shell.

• In Group 7, the further down the group an element is the higher its relative molecular mass, melting point and boiling point.
• In Group 7, the reactivity of the elements decreases going down the group.
• A more reactive halogen can displace a less reactive halogen from an aqueous solution of its salt.

• they have low melting and boiling points compared to most other metals
• they are very soft and can be cut easily with a knife
• they have low densities (lithium, sodium and potassium will float on water)
• they react quickly with water, producing hydroxides and hydrogen gas
• their hydroxides and oxides dissolve in water to form alkaline solutions

As you go down the group:

• their melting points decrease
• their densities increase
• they become softer
• they become more reactive

• they are non-metals
• they have low melting and boiling points
• they are brittle when solid
• they are poor conductors of heat and electricity
• they have coloured vapours
• their molecules each contain two atoms (they are diatomic)

Melting point and boiling point

The halogens have low melting points and boiling points. This is a typical property of non-metals. You can see from the graph that fluorine, at the top of Group 7, has the lowest melting point and boiling point in the Group. The melting points and boiling points then increase as you go down the Group.

State at room temperature

Room temperature is usually about 20°C. At this temperature, fluorine and chlorine are gases, bromine is a liquid, and iodine and astatine are solids. You should remember this trend down the periodic table - the top two elements are gases, the bottom two are solids and the middle element is liquid.

Colour

The halogens become darker as you go down the group. Fluorine is very pale yellow, chlorine is yellow-green and bromine is red-brown. Iodine crystals are shiny purple-black but easily turn into a dark purple vapour when they are warmed up.

• Many transition elements have ions with different charges, form coloured compounds and are useful as catalysts. 

• There are three types of strong chemical bonds: ionic, covalent and metallic.
• For ionic bonding the particles are oppositely charged ions. Ionic bonding occurs in compounds formed from metals combined with non-metals. 
• For covalent bonding the particles are atoms which share pairs of electrons. Covalent bonding occurs in most non-metallic elements and in compounds of non-metals. 
• For metallic bonding the particles are atoms which share delocalised electrons. Metallic bonding occurs in metallic elements and alloys. 

• When a metal atom reacts with a non-metal atom electrons in the outer shell of the metal atom are transferred.
• Metal atoms lose electrons to become positively charged ions.
• Non-metal atoms gain electrons to become negatively charged ions.
• The ions produced by metals in Groups 1 and 2 and by non-metals in Groups 6 and 7 have the electronic structure of a noble gas (Group 0).
• The electron transfer during the formation of an ionic compound can be represented by a dot and cross diagram, eg for sodium chloride.
• Electrons are transfered between the metal and the non-metal atoms.

• Ionic bonding is between a metal and a non-metal. Because of this, there are electrostatic forces of attraction between them.
• A metal forms a positive ion and a non-metal forms a negative ion. 

• An ionic compound is a giant structure of ions.
• Ionic compounds are held together by strong electrostatic forces of attraction between oppositely charged ions.
• These forces act in all directions in the lattice and this is called ionic bonding. 
• They have strong intermolecular forces of attraction acting in all directions.

• Ionic compounds have regular structures (giant ionic lattices) in which there are strong electrostatic forces of attraction in all directions between oppositely charged ions.
• These compounds have high melting points and high boiling points because of the large amounts of energy needed to break the many strong bonds.
• When melted or dissolved in water, ionic compounds conduct electricity because the ions are free to move and so charge can flow. 

• there are no gaps/sticks between the potassium ions and sulfide ions 

Physical

Transition elements

• high melting points
• high densities
• strong
• hard

Group 1

• low melting points
• low densities
• soft

Chemical

Transition elements

• low reactivity/react slowly (with water or oxygen)
• used as catalysts
• ions with different charges
• coloured compounds

Group 1

• very reactive/react (quickly) with water/non-metals
• not used as catalysts
• white/colourless compounds
• only forms a +1 ion

• When non-metal atoms share pairs of electrons, they form covalent bonds. These bonds between atoms are strong.
• Covalently bonded substances may consist of small molecules. 
• Some covalently bonded substances have very large molecules, such as polymers.
• Some covalently bonded substances have giant covalent structures, such as diamond and silicon dioxide.

• Metals consist of giant structures of atoms arranged in a regular pattern.
• The electrons in the outer shell of metal atoms are delocalised and so are free to move through the whole structure.
• The sharing of delocalised electrons gives rise to strong metallic bonds.

• The three states of matter are solid, liquid and gas.
• Melting and freezing take place at the melting point, boiling and condensing take place at the boiling point.
• The three states of matter can be represented by a simple model.
• In this model, particles are represented by small solid spheres.
• Particle theory can help to explain melting, boiling, freezing and condensing.

• The amount of energy needed to change state from solid to liquid and from liquid to gas depends on the strength of the forces between the particles of the substance.
• The nature of the particles involved depends on the type of bonding and the structure of the substance.
• The stronger the forces between the particles the higher the melting point and boiling point of the substance. 

• Limitations of the simple model above include that in the model there are no forces, that all particles are represented as spheres and that the spheres are solid.

•  (s), (l) and (g), with (aq) for aqueous solutions.

• Substances that consist of small molecules are usually gases or liquids that have relatively low melting points and boiling points.
• These substances have only weak forces between the molecules (intermolecular forces).
• It is these intermolecular forces that are overcome, not the covalent bonds, when the substance melts or boils.
• The intermolecular forces increase with the size of the molecules, so larger molecules have higher melting and boiling points.
• These substances do not conduct electricity because the molecules do not have an overall electric charge.

Covalent bonds.

• Polymers have very large molecules.
• The atoms in the polymer molecules are linked to other atoms by strong covalent bonds.
• The intermolecular forces between polymer molecules are relatively strong and so these substances are solids at room temperature. 

• Substances that consist of giant covalent structures are solids with very high melting points.
• All of the atoms in these structures are linked to other atoms by strong covalent bonds.
• These bonds must be overcome to melt or boil these substances.
• Diamond and graphite (forms of carbon) and silicon dioxide (silica) are examples of giant covalent structures.

• Metals have giant structures of atoms with strong metallic bonding.
• This means that most metals have high melting and boiling points.
• In pure metals, atoms are arranged in layers, which allows metals to be bent and shaped.
• Pure metals are too soft for many uses and so are mixed with other metals to make alloys which are harder. 
• Alloys are harder than pure metals because normally, the layers of atoms are able to slide over each other, so the metals can be bent and shaped.
• The different sizes of atoms in an alloy distort the layers in the structure, making it more difficult for the layers to slide over each other.

• Metals are good conductors of electricity because the delocalised electrons in the metal carry electrical charge as they move freely through the metal.
• Metals are good conductors of thermal energy because energy is transferred by the delocalised electrons.

• In diamond, each carbon atom forms four covalent bonds with other carbon atoms in a giant covalent structure, so diamond is very hard, has a very high melting point and does not conduct electricity.

• In graphite, each carbon atom forms three covalent bonds with three other carbon atoms, forming layers of hexagonal rings which have no covalent bonds between the layers.
• In graphite, one electron from each carbon atom is delocalised, which makes graphite similar to normal metals.
• Graphite is soft and slippery because there are only weak intermolecular forces between its layers.

Graphene is a single layer of graphite and has properties that make it useful in electronics and composites.

• Fullerenes are molecules of carbon atoms with hollow shapes.
• The structure of fullerenes is based on hexagonal rings of carbon atoms but they may also contain rings with five or seven carbon atoms.
• The first fullerene to be discovered was Buckminsterfullerene (C60) which has a spherical shape.
• They can be used as lubricants.
• They can be used to cage other molecules so they could be used to deliver drugs to the body.

• Carbon nanotubes are cylindrical fullerenes with very high length to diameter ratios.
• Their properties make them useful for nanotechnology, electronics and materials. They can do this without adding weight.
• They have a large surface area to volume ratio.
• They can conduct both electricity and thermal energy.
• High tensile strength - don't break when they're stretched.

• Nanoscience refers to structures that are 1–100 nm in size, of the order of a few hundred atoms.
• Nanoparticles, are smaller than fine particles (PM2.5), which have diameters between 100 and 2500 nm (1 x 10-7 m and 2.5 x 10-6 m).
• Coarse particles (PM10) have diameters between 1 x 10-5 m and 2.5 x 10-6 m.
• Coarse particles are often referred to as dust.
• As the side of cube decreases by a factor of 10 the surface area to volume ratio increases by a factor of 10.

• Nanoparticles may have properties different from those for the same materials in bulk because of their high surface area to volume ratio.
• It may also mean that smaller quantities are needed to be effective than for materials with normal particle sizes.
• Some nanoparticles conduct electricity so they can be used in electric circuits for computer chips.
• Silver nanoparticles have antibacterial properties. They can be added to polymer fibres that are often used to make surgical masks and wound dressings and they can be added to deodorants.
• They can also be used in cosmetics.

• Large surface area to volume ratio makes them effective catalysts

• Nanoparticles in sun creams can be absorbed deeper into the skin

• Nanoparticles in face creams cover better so you have to use less

• Sun creams that use nanoparticles are transparent while traditional ones are white

• So small they can enter the skin and therefore the bloodstream

• Large surface can make them too reactive and explosive in some situations

• Easily become airborne, breathing in can potentially damage the lungs
• They might be toxic to some types of cell, such as skin, bone, brain and liver cells

• The law of conservation of mass states that no atoms are lost or made during a chemical reaction so the mass of the products equals the mass of the reactants.
• This means that chemical reactions can be represented by symbol equations which are balanced in terms of the numbers of atoms of each element involved on both sides of the equation.

• When a reactant or product is a gas
• Some reactions may appear to involve a change in mass but this can usually be explained because a reactant or product is a gas and its mass has not been taken into account.
• For example: when a metal reacts with oxygen the mass of the oxide produced is greater than the mass of the metal or in thermal decompositions of metal carbonates carbon dioxide is produced and escapes into the atmosphere leaving the metal oxide as the only solid product. 

• Whenever a measurement is made there is always some uncertainty about the result obtained.

• The number of atoms, molecules or ions in a mole of a given substance is the Avogadro constant. The value of the Avogadro constant is 6.02 x 10 to the power of 23 per mole.

Moles = mass / mr

• The balancing numbers in a symbol equation can be calculated from the masses of reactants and products by converting the masses in grams to amounts in moles and converting the numbers of moles to simple whole number ratios.

• In a chemical reaction involving two reactants, it is common to use an excess of one of the reactants to ensure that all of the other reactant is used.
• The reactant that is completely used up is called the limiting reactant because it limits the amount of products.

Many chemical reactions take place in solutions. The concentration of a solution can be measured in mass per given volume of solution, eg grams per dm cubed (g/dm cubed ).

Mass/volume = concentration

• the reaction may not go to completion because it is reversible
• some of the product may be lost when it is separated from the reaction mixture
• some of the reactants may react in ways different to the expected reaction.

(Mass of product actually made / maximum theoretical mass of product) x 100

(R.F.M of desired product / sum of R.F.M of all reactants) x 100

• The amount of a product obtained is known as the yield. When compared with the maximum theoretical amount as a percentage, it is called the percentage yield.

• The atom economy (atom utilisation) is a measure of the amount of starting materials that end up as useful products.
• It is important for sustainable development and for economic reasons to use reactions with high atom economy

Mol/volume

1 mole of any gas has a volume of 24dm³ volume (dm³) = moles x 24dm³

• 1dm³ = 1000cm³

• Metals react with oxygen to produce metal oxides. The reactions are oxidation reactions because the metals gain oxygen.

• When metals react with other substances the metal atoms form positive ions.
• The reactivity of a metal is related to its tendency to form positive ions.
• Metals can be arranged in order of their reactivity in a reactivity series.
• The metals potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper can be put in order of their reactivity from their reactions with water and dilute acids.

• K
• Na
• Li
• Ca 
• Mg
• Al 
• Extracted by electrolysis
• C
• Zn
• Fe
• H
• Cu 
• Extracted with reduction by carbon
• Ag
• Au

• Oxidation = gain of oxygen
• Reduction = loss of oxygen

• Metals higher than carbon in the reactivity series have to be extracted using electrolysis, which is expensive
• Metals below carbon in the reactivity series can be extracted by reduction using carbon. For example, iron oxide is reduced in a blast furnace to make iron. This is because carbon can only take the oxygen away from metals which are less reactive than itself is.
• A few metals are so unreactive like gold and silver, but they are just mined in their elemental form.

• Oxidation is Losing, Reduction is Gaining. This refers to the transfer of electrons. If REDuction and OXidation occur at the same time, it is known as a REDOX reaction. 
• Examples of a REDOX reaction:
• Metals reacting with acids: all reactions of metals with acids are redox reactions. For eg. the reaction of ion with dilute sulfuric acid is a redox reaction.
• The iron atoms lose electrons to become iron (II) ions - they are oxidised by the hydrogen ions. Fe -> Fe (2+) + 2 e(-)
• The hydrogen ions gain electrons to become hydrogen atoms - they are reduced by the hydrogen atoms. 2H(+) + 2e(-) -> H2
• The ionic equation for the redox reaction is Fe + 2H(+) -> Fe(2+) + H(2)
• Halogen Displacement Reaction: a more reactive halogen can displace a less reactive halogen. For eg. chlorine displaces bromine from potassium bromide solution.
• The chlorine atoms gain electrons to become chloride ions - they are reduced by the bromide ions. Cl(2) + 2e(-) -> 2Cl (-)
• The bromide ions lose electrons to become bromine ions - they are oxidised by the chlorine atoms. 2Br(-) -> Br(2) +2e(-)
• The ionic equation for this redox reaction is Cl(2) +2Br(-) -> 2Cl(-) + Br(2)

Acid + Metal -> Salt + Hydrogen

Metal + Water -> Metal Hydroxide + Hydrogen 

Acid + Metal Carbonate -> Salt + Water + Carbon dioxide

Acid + Metal Oxide -> Salt + Water

Acid + Metal Hydroxide -> Salt + Water

Acid + Base -> Salt + Water

• Acids are neutralised by alkalis (eg soluble metal hydroxides) and bases (eg insoluble metal hydroxides and metal oxides) to produce salts and water, and by metal carbonates to produce salts, water and carbon dioxide.
• The particular salt produced in any reaction between an acid and a base or alkali depends on: •
• the acid used (hydrochloric acid produces chlorides, nitric acid produces nitrates, sulfuric acid produces sulfates)
• the positive ions in the base, alkali or carbonate.

• Soluble salts can be made from acids by reacting them with solid insoluble substances, such as metals, metal oxides, hydroxides or carbonates.
• The solid is added to the acid until no more reacts and the excess solid is filtered off to produce a solution of the salt.
• Salt solutions can be crystallised to produce solid salts.

• Pick the right acid, plus an insoluble base such as an insoluble hydroxide, metal oxide or carbonate. For eg. dilute hydrochloric acids and copper chloride to make copper crystals.
• Gently warm the dilute acid using a bunsen burner, then turn off the bunsen burner
• Add the insoluble base to the acid a bit at a time until no more reacts as the base is in excess. You'll know when the acid has been neutralised because een after sitrring, the excess solid will just sink to the bottom of the flask. 
• Then filter out the excess solid to get the salt solution. 
• To get pure, solid crystals of the salt, gently heat the solution using a water bath or an electric heater to evapourate some of the water (to make it more concentrated) and then stop heating it and leave the solution to cool. Crystals of the salt should form, which can be filtered out of the solution and then dried. This is called heating to the point of crystallisation.

• Acids produce hydrogen ions (H+) in aqueous solutions.
• Aqueous solutions of alkalis contain hydroxide ions (OH– ).
• The pH scale, from 0 to 14, is a measure of the acidity or alkalinity of a solution, and can be measured using universal indicator or a pH probe.
• A solution with pH 7 is neutral. Aqueous solutions of acids have pH values of less than 7 and aqueous solutions of alkalis have pH values greater than 7.
• In neutralisation reactions between an acid and an alkali, hydrogen ions react with hydroxide ions to produce water.
• This reaction can be represented by the equation:
• H+ (aq) + OH- (aq) -> H(2)0 (l)

• pH scale is a measure of how acidic or alkaline a solution is. It is also a measure of the concentration of H+ ions in the solution.
• The pH of a strong acid is always lower than the pH of a weaker acid if they have the same concentration.
• Lower pH = more acidic
• A neutral substance has a pH of 7
• You can measure the pH of a solution:

1) An indicator is a dye that changes colour depending on whether it's above or below a certtain pH. Some indicators contain a mixture of dyes that mean they gradually change colour over a broad range of pH. These are called wide range indicators and they are usefule for estimating the pH of a solution. 

2) A pH probe attached to a pH meter can also be used to measure pH electronically. The pH probe is placed in the solution you are measuring and the pH is given on a digital display as a numerical value, meaning it's accurate than an indicator. 

• An acid is a substance that forms an aqueous solution with a pH of less than 7. Acids form H+ ions in water. 
• A base is any substance that will react with an acid to form a salt.
• An alkali is a base that dissolves in water to form a solution with a pH greater than 7. Alkalis form OH- ions in water. 

1) Using a pipette and pipette filler, add a set volume of the alkali to a conical flask. Add two or three drops of the indicator.

2) Use a funnel to fill a burette with some acid of known concentration. Make sure you do this below eye level - health hazard - and wear goggles. Record the initial volume of acid in the burette. 

3) Use the burette to add the acid to the alkali a bit at a time, giving the conical flask regular swirls at regular intervals. Swirl slowly until you think you seen the end-point (colour change) is about to be reached.

4) The indicator changes colour when all of the alkali has been neutralised. 

5) Record the final volume of the acid in the burette, and use it, along with the intial reading, to calculate the volume of acid used to neutralise the alkali.

• To increase the accuracy of your results and to spot any anomalous results, you need to do several consistent readings. 
• The first titration you do should be a rough titration to get an approximate idea of where the solution changes colour (end-point).
• You then need to repeat the whole experiment a few times, making sure tha you get the same answer or concordant results - within 0.10 cubic centimetres.
• Finally, calculate a mean of your results, ignoring any anomalous results.

Phenolphthalein:

• Colourless in acids
• Pink in alkalis

Litmus:

• Red in acids
• Blue in alkalis

Methyl Orange:

• Red in acids
• Yellow in alkalis

• A strong acid is completely ionised in aqueous solution. Examples of strong acids are hydrochloric, nitric and sulfuric acids.
• Exam answer: produces H+ / hydrogen ions in aqueous solution (but is) only partially / slightly ionised
• A weak acid is only partially ionised in aqueous solution. Examples of weak acids are ethanoic, citric and carbonic acids.
• For a given concentration of aqueous solutions, the stronger an acid, the lower the pH.
• As the pH decreases by one unit, the hydrogen ion concentration of the solution increases by a factor of 10.

• The ionisation in a weak acid is a reversible reaction, which sets up an equilibrium between the undissociated ions and dissociated acid, Since only a few of the acid particles release H+ ions, the position of equilibrium moves to the left.

• When an ionic compound is melted or dissolved in water, the ions are free to move about within the liquid or solution.
• These liquids and solutions are able to conduct electricity and are called electrolytes.
• Passing an electric current through electrolytes causes the ions to move to the electrodes.
• Positively charged ions move to the negative electrode (the cathode) where they are reduced, and negatively charged ions move to the positive electrode (the anode) where they are oxidised.
• This creates a flow of charge through the electrolyte as ions travel to the electrodes. As ions gain or lose electrons, they form the uncharged element and are discharged from the electrolyte.
• Ions are discharged at the electrodes producing elements.
• This process is called electrolysis.

• When a simple ionic compound (eg lead bromide) is electrolysed in the molten state using inert electrodes, the metal (lead) is produced at the cathode and the non-metal (bromine) is produced at the anode.
• Electrolysis of molten ionic solids form elements.
• Molten ionic compounds can be electrolysed because the ions can move freely and conduct electricity.
• Molten ionic compounds are always broken up into their elements. A good example is molten lead bromide. 
• The electrodes should be made up of an inert material, so they don't react with the electrolyte. 
• Positive metal ions are reduced to the element at the cathode.
• Negative non-metal ions are oxidised to the element at the anode.

Positive is Anode. Negative is Cathode.

• Metals can be extracted from molten compounds using electrolysis.
• Electrolysis is used if the metal is too reactive to be extracted by reduction with carbon or if the metal reacts with carbon.
• Large amounts of energy are used in the extraction process to melt the compounds and to produce the electrical current.
• Aluminium is manufactured by the electrolysis of a molten mixture of aluminium oxide and cryolite using carbon as the positive electrode (anode). It can be extracted from the ore bauxite by electrolysis. Bauxite contains aluminium oxide (Al2O3). Aluminium oxide has a very high melting temperature so it's mixed with a cryolite to lower the melting point.
• The molten mixture contains free ions so it is able to conduct electricity. 
• The positive Al(3+) ions are attracted to the negative electrode where they can pick up three electrons and turn into the neutral aluminium atoms. These then sink to the bottom of the electrolysis tank.
• The negative O(2-) ions are attracted to the positive electrode where they each lose two electrons. The neutral oxygen atoms will combine to form O(2) molecules.
• Negative electrode: Al(3+) + 3(e-) -> Al
• Positive electrode: O(2-) -> O(2) + 4e(-)
• Overall equation: aluminium oxide -> aluminium and oxygen
• Al(2)O(3) -> 4Al + 30(2)

• The ions discharged when an aqueous solution is electrolysed using inert electrodes depend on the relative reactivity of the elements involved.
• At the negative electrode (cathode), hydrogen is produced if the metal is more reactive than hydrogen.
• At the positive electrode (anode), oxygen is produced unless the solution contains halide ions when the halogen is produced.
• This happens because in the aqueous solution water molecules break down producing hydrogen ions and hydroxide ions that are discharged.

• Energy is conserved in chemical reactions.
• The amount of energy in the universe at the end of a chemical reaction is the same as before the reaction takes place.
• If a reaction transfers energy to the surroundings the product molecules must have less energy than the reactants, by the amount transferred.
• An exothermic reaction is one that transfers energy to the surroundings so the temperature of the surroundings increases.
• Exothermic reactions include combustion, many oxidation reactions and neutralisation.
• Everyday uses of exothermic reactions include self-heating cans and hand warmers.
• An endothermic reaction is one that takes in energy from the surroundings so the temperature of the surroundings decreases.
• Endothermic reactions include thermal decompositions and the reaction of citric acid and sodium hydrogencarbonate.
• Some sports injury packs are based on endothermic reactions.

1) Measure the amount of energy released by a chemical reaction insolution by taking the temperature of the reagents (making sure they're the same), mixing them in polystyrene cup and measuring the temperature of the solution at the end of the reaction. 

2) The biggest problem with energy measurements is the amount of energy lost to the surroundings.

3) You can reduce it a bit by putting the polystyrene cup into a beaker of cotton wool to give more insulation, and putting the lid on the cup to reduce energy lost by evaporation. 

4) This method works for neutralisation reactions or reactions between metals and acids, or carbonates and acids. 

5) You can also use the method to investigate what effect different variables have on the amount of energy transferred eg. the mass or concentration of the reactants used.

1) Put 25 centimetres cubed of 0.25 mol/dm cubed of HCl and NaOH in separate beakers.

2) Place the beakers in a water bath set to 25 degrees celsius until they are both at the same temperature.

3) Add the HCl followed by the NaOH to a polystyrene cup with a lid.

4) Take the temperature of the mixture every 30 seconds and record the highest temperature.

5) Repeat steps 1-4 using 0.5 mol/dm cubed and then 1 mol/dm cubed of HCl.

• Bonds breaking - endothermic because energy must be supplied, takes in from surroundings. The energy released when bonds break is higher than the energy released when bonds form. 
• In an endothermic reaction, the energy needed to break existing bonds is greater than the energy released from forming new bonds.
• Bonds forming - exothermic because energy is released. In an exothermic reaction, the energy released from forming new bonds is greater than the energy needed to break existing bonds.

The difference between the bond energy of the reactants and the products. Products - reactants.

• Cells contain chemicals which react to produce electricity.
• The voltage produced by a cell is dependent upon a number of factors including the type of electrode and electrolyte.
• A simple cell can be made by connecting two different metals in contact with an electrolyte.
• Batteries consist of two or more cells connected together in series to provide a greater voltage. In non-rechargeable cells and batteries the chemical reactions stop when one of the reactants has been used up.
• Alkaline batteries are non-rechargeable.
• Rechargeable cells and batteries can be recharged because the chemical reactions are reversed when an external electrical current is supplied.

• The two electrodes must be able to conduct electricity and so they are usually always metal.
• The electrolyte is a liquid that contains ions which reacts with the electrodes.
• The chemical reactions between the electrodes and the electrolyte sets up a charge difference between the electrodes.
• If the electrodes are then connected by a wire, the charge is able to flow and electricity is produced.
• A voltmeter can also be connected to the circuit to measure the voltage of the cell.

• Fuel cells are supplied by an external source of fuel (eg hydrogen) and oxygen or air.
• The fuel is oxidised electrochemically within the fuel cell to produce a potential difference.
• The overall reaction in a hydrogen fuel cell involves the oxidation of hydrogen to produce water.
• Hydrogen fuel cells offer a potential alternative to rechargeable cells and batteries.
• OR:
• A fuel cell is an electrical cell that's supplied with fuel and oxygen and uses the energy from the reaction between them to produce electrical energy efficiently. 
• When the fuel enters the cell, it becomes oxidised and sets up a potential differene within the cell. 
• There are a few different types of fuel cells, using different fuels and different electrolytes. One of the types is a hydrogen-oxygen fuel cells. 
• This fuel cell combines hydrogen and oxygen to produce nice clean water and energy.

• Conventional fuels for vehicles such as petrol have a finite supply and they're very polluting
• This has lead to vehicles that use electrical energy becoming more and more popular.
• Batteries are one way of getting clearer energy but hydrogen-oxygen fuel cells might be even better.
• Fuel cell vehicles don't produce as many pollutants as other fuels - no greenhouse gases, nitrogen oxides, sulfur dioxide or carbon monoxide.
• The only by-products are water and heat.
• Electric vehicles don't produce many pollutants either - but their batteries are much more polluting to dispose of than fuel cells because they're highly toxic metal compounds.
• Batteries in electric vehicles are rechargable but there's a limit to how many times they can be recharged before they need replacing. 
• Batteries are now more expensive to make than fuel cells.
• Batteries store less energy than fuel cells and so would need to be recharged more often - which can take a long time.

• The electrolyte is often a solution of potassium hydroxide and the electrodes are usually porous carbon with a catalyst.
• Hydrogen goes into the anode compartment and oxygen goes into the cathode compartment.
• At the anode, the hydrogen loses electrons to produce H+ ions which is oxidation. H(2) -> 2H+ + 2e(-). 
• H+ ions in the electrolyte move to the cathode. 
• At the cathode, oxygen gains electrons from the cathode and reacts with H+ ions from the electrolyte to form water. This is reduction.
• The electrons flow through an external circuit from the anode to the cathode and this is the electric current. 
• The overall reaction is hydrogen plus oxygen which gives water.

• In aqeuous solutions, as well as the ions from the ionic compound, there will be hydrogen ions which are H+ and hydroxide ions which are OH -. These two ions form water.
• Which ions are discharged at the electrodes when the solution is electrolysed will depend on the relative reactivity if all the ions in the solution. 

Cathode:

• at the cathode, if H+ ions and metal ions are present, hydrogen gas will be produced if the metal ions form an elemental metal that is more reactive than hydrogen.
• if the metal ion forms an elemental metal that is less reactive than hydrogen, a solid layer of pure metal will be produced instead, which will coat the cathode. 

Anode: 

• at the anode, if OH- and halide ions are present, molecules of chlorine, bromine or iodine will be produced.
• if no halide ions are present, then the OH- ions from the water will be discharged and oxygen gas and water will be formed. 

• A solution of copper (II) sulfate (CuSO(4)) contains 4 different ions. Cu(2+), SO4(2-), H+ and OH-. 
• Copper metal is less reactive than hydrogen, so at the cathode copper metal is produced and coats the electrode. Cu (2+) + 2e(-) -> Cu.
• There aren't any halide ions present, so at the anode oxygen and water are produced. They oxygen can be seen as bubbles. 4OH- -> O2 + 2H2O + 2e(-).

• A solution of sodium chloride (NaCl) contains four different ions. Na+, Cl-, H+ and OH-. 
• Sodium metal is more reactive reactive than hydrogen, so at the cathode, hydrogen gas is produced. 2H+ + 2e(-) -> H2
• Chloride ions are present in the solution, so at the anode, chlorine gas is produced. 2Cl- -> Cl2 + 2e-.