Chemical Changes

Acids and Bases

• The Ph scale is a measure of how acidic or alkaline a solution is. The lower the pH, the higher the acidity, and the higher the pH, the more alkaline it is. Neutral substances are pH 7.
• You can measure pH:
• Indicator - a dye that changes colour depending on its position on the pH scale. Some contain a mixture of dyes, so can change over time and give a broader range of pH. These are called wide range indicators.
• pH probe - a pH probe attached to a pH meter can electronically measure pH. The probe is placed in a solution and the pH is displayed numerically on the digital display. It is more accurate than an indicator.
• An acid is a substance that forms aqueous solutions with a pH of less than 7. They form H+ ions in water
• A base is a substance with a pH greater than 7
• An alkali is a base that dissolves in water to form a solution with a pH greater than 7. They form OH- ions in water.
• Reactions are as follows:
• Acid + Base ---> Salt + Water
• In terms of H+ and OH- ions:
• H+ + OH- ---> H2O
• These reactions are called neutralisations, as the products are natural (have a pH of 7)

Strong acids and weak acids

• Acids ionise in aqueous solutions, they form H+ ions.
• Strong acids - Completely ionise in water and all acid particles dissociate to release H+ ions
• Weak acids - Do not fully ionise in water and only a small proportion of acid particles dissociate to release H+ ions
• The ionisation of weak acids is a reversible reaction, as an equilibrium is set up between the dissociated and undissociated acid. The equilibrium lies to the left, as only a few H+ ions are released.
• Reactions of acids involve the H+ ions reacting with other substances. If the concentration of H+ ions is higher, the rate of reaction will be faster, so stronger acids are more reactive than weak acids.
• The pH of an acid or alkali is a measure of the concentration of H+ ions in the solution. On the pH scale
• Factor H+ ion concentration changes by = 10^-x
• X is the difference in pH
• Strength is different to concentrations:
• Acid strength measures the proportion of acid that ionises in water
• Concentration measures the amount of acid in a given volume
• The larger amount of acid in a volume there is, the more concentrated it is. So you can still have an acid that isn't very concentrated (dilute), but is still strong.
• However, pH will decrease if concntration is increasing, regardless of whether it is a strong or weak acid.

Reactions of Acids

• Some metal oxides and metal hydroxides dissolve in water to form alkalis. Alkalis can react with acids in neutralisation reactions. Bases that aren't dissolved in water can also act this way. So:
• Acid + Metal Oxide/Metal Hydroxide ----> Salt + Water
• Metal carbonates are bases, and produce salt and water, with the addition of carbon dioxide, so:
• Acid + Metal Carbonates ----> Salt + Water + Carbon Dioxide
• Examples of metal oxides or hydroxides in neutralisation reactions:
• Hydrochloric acid + copper oxide -----> Copper chloride + Water
• Sulfuric acid + Pottasium hydroxide ---> Pottasium sulfate + Water
• Examples of metal carbonates:
• Hydrochloric acid + Sodium Carbonate ---> Sodium Chloride + Water + Carbon Dioxide
• Sulfuric acid + Calcium carbonate --> Calcium sulfate + Water + Carbon Dioxide

The Reactivity Series

• The reactivity series lists metals in order of their reactivity towards other substances. Reactivity is based on how easily a metal can lose an electron and form a positive ion. The higher up the list, the easier it is to lose the electron. This also means that the higher up the metal is, the more it reacts with acid and water.
• Some metals do this reaction:
• Acid + Metal ---> Salt + Hydrogen
• The speed of the reaction is determined by the rate at which bubbles of hydrogen are given off.
• More reactive metals react explosively, whereas others react less violently, and some simply do not react. The more reactive a metal, the greater the temperature change should be.
• Metals also react with water:
• Metal + Water ---> Metal Hydroxide + Hydrogen
• This only really happens with more reactive metals.
• Reactivity Series:
• Pottasium - Peter
• Sodium - Sells
• Lithium - Little
• Calcium - Children
• Magnesium - Multi
• Carbon - Coloured
• Zinc - Zebras
• Iron - In
• Hydrogen - Heavy
• Copper - Clothes

Separating metals from metal oxides

• Many common metals react with oxygen to form oxides, in a process called oxidation. These metals need to be extracted from these ores using a process called reduction:
• The formation of the ore is a process of Oxidation
• The extraction of the metal is a process called Reduction
• Metals can also be extracted from their ores using carbon. The ore is reduced as oxygen is removed from it, then the carbon gains oxygen so is oxidised. However, metals need to be below carbon in the reactivity series to do this, as carbon can only take the oxygen away from metals that are less reactive than itself.

Redox Reactions

• Remember OIL RIG:
• Oxidation Is Loss in terms of electrons, and Reduction Is Gain in terms of electrons.
• Redox reactions are when reduction and oxidation both occur at the same time, hence REDOX.
• Displacement reactions are an example of a redox reaction, the rule for displacement is:
• A more reactive metal will displace a less reactive metal from its compound
• For example, if you put iron in a solution of copper sulfate, the more reactive iron will displace the less reactive copper from the solution:
• Iron + Copper Sulfate ---> Iron Sulfate + Copper
• In this reaction, the iron is oxidised (has lost electrons), and the copper is reduced (gained electrons).
• In displacement reactions, the metal ion always gains the electrons and the metal atom always loses the elctron.
• Ionic equations focus on the substance that has been oxidised or reduced.

Electrolysis of Aqueous Solutions

• In aqeous solutions, as well as ions from the ionic compounds, there will be hydrogen ions/H+ ions, and hydroxide ions/OH- ions from the water.
• At the cathode, if H+ ions and metal ions are present, hydrogen gas wil be produced if the metal ions form an element that is more reactive than hydrogen. If the element produced is less reactive than hydrogen, the pure metal will form instead.
• At the anode, if OH- ions and halide ions are present, molecules of Chlorine, Bromine and Iodine will be formed. If no halide ions are present, then OH- is discharged and oxygen is formed.
• To test for gases, use these tests:
• Chlorine bleaches litmus paper white
• Hydrogen makes a squeaky pop with a lit splint
• Oxygen will relight a glowing splint