3.3.1 Oxidation and reduction in terms of oxygen

Oxidation is gain of oxygen, reduction is loss

1. Burning magnesium in air forms magnesium oxide, so magnesium gains oxygen and is oxidised: $$2Mg(s)+O_2(g)\to 2MgO(s)$$
   1. Rusting, respiration and the combustion of fuels are other everyday oxidations, because oxygen is added to a substance in each one.
2. Heating black copper(II) oxide in a stream of hydrogen removes its oxygen and leaves pink-brown copper, so copper(II) oxide is reduced.
   1. Removing oxygen from a metal oxide to leave the metal is the most important use of reduction, and it is how iron is obtained from its ore.
3. Both words describe what happens to a whole substance, so the substance that gains oxygen is oxidised and the substance that loses oxygen is reduced.

Oxidation and reduction happen together in oxygen transfer

1. In $CuO+H_2\to Cu+H_{2}O$, the oxygen that leaves copper(II) oxide is gained by hydrogen.
   1. Copper(II) oxide is reduced to copper, while hydrogen is oxidised to water.
2. Oxygen cannot vanish or appear from nowhere, because atoms are conserved in every reaction.
   1. Whenever one substance loses oxygen, another substance in the same reaction must gain it.
3. The name redox is built from the start of the words reduction and oxidation.
4. Heating a metal oxide with carbon is another redox reaction, because carbon takes the oxygen and becomes carbon dioxide: $$2CuO+C\to 2Cu+{CO}_2$$

Identifying the species oxidised and the species reduced

1. Pick out each reactant that takes part in the oxygen transfer and find the product it becomes.
2. A species that ends up with more oxygen has been oxidised.
3. A species that ends up with less oxygen has been reduced.
4. When one reactant is oxygen gas, it is the source of the oxygen, so the species to name as oxidised is the one it combines with.
   1. In the combustion of hydrogen, $2H_2+O_2\to 2H_{2}O$, hydrogen is the species oxidised to water.

The oxygen test works for any oxygen transfer

1. When methane burns, ${CH}_4+2O_2\to {CO}_2+2H_{2}O$, its carbon and hydrogen both gain oxygen, so methane is oxidised.
2. Magnesium keeps burning in carbon dioxide, leaving black specks of carbon: $$2Mg+{CO}_2\to 2MgO+C$$
   1. Carbon dioxide loses its oxygen and is reduced to carbon, while magnesium gains oxygen and is oxidised to magnesium oxide.
3. Magnesium also reacts with steam: $$Mg+H_{2}O\to MgO+H_2$$
   1. Here water loses oxygen and is reduced to hydrogen, while magnesium is oxidised.
4. Heating zinc oxide with carbon, $2ZnO+C\to 2Zn+{CO}_2$, shows zinc oxide being reduced by carbon.

3.3.2 Acids, alkalis and neutralisation

Acids give hydrogen ions and alkalis give hydroxide ions

1. Hydrogen chloride releases hydrogen ions, $H^{+}$, and chloride ions, ${Cl}^{-}$, when it dissolves in water, forming hydrochloric acid.
2. The three common laboratory acids are hydrochloric acid, $HCl$, sulfuric acid, $H_2{SO}_4$, and nitric acid, ${HNO}_3$.
   1. Each of them releases $H^{+}$ ions in water, and these ions give acids their shared reactions.
3. Sodium hydroxide, $NaOH$, dissolves in water to give a solution containing hydroxide ions, ${OH}^{-}$.
   1. Potassium hydroxide and calcium hydroxide are other alkalis, and each supplies ${OH}^{-}$ ions in solution.
4. Dry hydrogen chloride gas does not turn dry indicator paper red, because hydrogen ions form only when it dissolves in water.

Neutralisation forms a salt and water

1. Metal oxides and metal hydroxides, such as copper(II) oxide and sodium hydroxide, act as bases.
   1. Sodium hydroxide dissolves in water, so it is also an alkali, while copper(II) oxide is insoluble and is not.
2. Every neutralisation follows the same word equation: acid + base $\to$ salt + water
3. The salt takes its positive ion from the base and its negative ion from the acid.
   1. Sodium hydroxide and hydrochloric acid give sodium chloride: $$NaOH(aq)+HCl(aq)\to NaCl(aq)+H_{2}O(l)$$
   2. Copper(II) oxide and sulfuric acid give copper(II) sulfate: $$CuO(s)+H_2{SO}_4(aq)\to {CuSO}_4(aq)+H_{2}O(l)$$

Hydrogen ions and hydroxide ions react to form water

1. Hydrochloric acid contains hydrogen ions and chloride ions, and sodium hydroxide solution contains sodium ions and hydroxide ions.
2. When the solutions mix, each $H^{+}$ ion joins an ${OH}^{-}$ ion to make a water molecule: $$H^{+}(aq)+{OH}^{-}(aq)\to H_{2}O(l)$$
3. The sodium and chloride ions do not react, so they are left dissolved in the water as sodium chloride.
4. Every acid gives $H^{+}$ ions and every alkali gives ${OH}^{-}$ ions, so the same equation describes every aqueous neutralisation.
   1. Nitric acid with potassium hydroxide, or sulfuric acid with sodium hydroxide, produces a different salt, but the reaction that forms the water is the same.
5. When all the $H^{+}$ ions have reacted with an equal number of ${OH}^{-}$ ions, the solution is neutral.

Carbonates and some metals also react with acids

1. An acid and a metal carbonate react to form a salt, water and carbon dioxide.
   1. Calcium carbonate reacts with hydrochloric acid like this: $${CaCO}_3(s)+2HCl(aq)\to {CaCl}_2(aq)+H_{2}O(l)+{CO}_2(g)$$
   2. The carbon dioxide makes the mixture fizz, and it turns limewater milky.
2. Metals above hydrogen in the reactivity series, such as magnesium, zinc and iron, react with acids to form a salt and hydrogen.
   1. Magnesium reacts with hydrochloric acid like this: $$Mg(s)+2HCl(aq)\to {MgCl}_2(aq)+H_2(g)$$
3. Copper, silver and gold are below hydrogen, so they do not react with dilute acids.

Writing balanced equations for acid reactions

1. First decide the type of reaction, because it fixes the products: a base gives a salt and water, a carbonate adds carbon dioxide and a metal gives hydrogen in place of water.
2. Name the salt from the metal and the acid, so zinc and sulfuric acid give zinc sulfate.
3. Write the salt formula from the ion charges, so ${Mg}^{2+}$ with ${Cl}^{-}$ gives ${MgCl}_2$.
4. Balance the equation by changing the large numbers in front of formulae, never the small numbers inside them.
5. Check that every element, including oxygen and hydrogen, has the same number of atoms on both sides.

3.3.3 The pH scale and measuring pH

The pH scale measures acidity and alkalinity

1. pH values usually run from 0 to 14, and every aqueous solution has a place on this one scale.
2. A solution with a pH below 7 is acidic, and the lower the pH, the more acidic it is.
3. A solution with a pH of 7 is neutral, like pure water.
4. A solution with a pH above 7 is alkaline, and the higher the pH, the more alkaline it is.
5. Placing solutions on the same scale shows their relative acidity, so a solution of pH 2 is more acidic than one of pH 5.

Universal indicator shows an approximate pH through its colour

1. A few drops of universal indicator added to a solution give a colour that depends on the pH.
2. Acidic solutions turn it red, orange or yellow, with red for the lowest pH.
3. A neutral solution turns it green.
4. Alkaline solutions turn it blue or purple, with purple for the highest pH.
5. The colour is matched against a colour chart to read off the pH, usually to the nearest whole number.
6. Indicator paper soaked in universal indicator works in the same way when a drop of solution is placed on it.
   1. A glass rod transfers one drop to the paper, so the rest of the sample is not contaminated.

A pH probe gives a more precise reading

1. Before use, the probe is calibrated in solutions of known pH, such as pH 4 and pH 7, so its readings are accurate.
2. Rinse the probe with distilled water before it goes into each new sample, so no solution is carried over.
3. Dip the tip of the probe into the sample and wait until the reading is steady before recording it.
4. A probe gives the pH to one or two decimal places, such as pH 6.4, while a colour chart gives whole numbers.
   1. This makes a probe better for spotting small changes, such as the pH of a pond water sample shifting over a week.
5. A probe also avoids judging a colour, which varies from person to person and does not work in a coloured solution such as blackcurrant squash.

Choosing a pH method by the precision needed

1. Universal indicator solution is quick and cheap and shows the approximate pH of a colourless solution.
2. Indicator paper suits small samples, such as a single drop taken from a pipette.
3. A pH probe suits a precise value, a coloured solution or a series of readings over time.
4. Whichever method is used, a clean container and a fresh sample stop contamination from changing the result.
5. Samples that are being compared are all tested by the same method, so the comparison is fair.

3.3.4 Oxidation and reduction in terms of electrons

Oxidation is loss and reduction is gain of electrons

1. The memory aid OIL RIG stands for oxidation is loss, reduction is gain, where the loss and gain refer to electrons.
2. When a magnesium atom forms a magnesium ion, it loses two electrons, so it is oxidised: $$Mg\to {Mg}^{2+}+2e^{-}$$
3. When a copper(II) ion forms a copper atom, it gains two electrons, so it is reduced: $${Cu}^{2+}+2e^{-}\to Cu$$
4. Losing negative electrons makes a species more positive, while gaining electrons makes it less positive or more negative.
5. These definitions extend the oxygen definitions, so reactions with no oxygen at all can still be redox reactions.

Redox reactions transfer electrons from one species to another

1. When magnesium is added to copper(II) sulfate solution, a brown coating of copper forms and the blue colour fades: $$Mg(s)+{Cu}^{2+}(aq)\to {Mg}^{2+}(aq)+Cu(s)$$
   1. Each magnesium atom transfers two electrons to a copper(II) ion.
2. Magnesium atoms lose electrons, so magnesium is oxidised.
3. Copper(II) ions gain electrons, so the copper(II) ions are reduced.
4. The sulfate ions do not change, so they are spectator ions and are left out of the equation.
5. The number of electrons lost always equals the number gained, because electrons are not created or destroyed.

Comparing charges shows which species lost or gained electrons

1. Write the charge on each species on both sides of the equation, counting an uncharged atom or molecule as zero.
2. A charge that becomes more positive, such as $0$ to $2+$ or $1-$ to $0$, shows electron loss and oxidation.
3. A charge that becomes less positive, such as $2+$ to $0$ or $0$ to $1-$, shows electron gain and reduction.
4. In ${Cl}_2+2{Br}^{-}\to 2{Cl}^{-}+{Br}_2$, each chlorine atom goes from $0$ to $1-$, so chlorine is reduced.
   1. Each bromide ion goes from $1-$ to $0$, so the bromide ions are oxidised to bromine.

Electron definitions explain burning and electrode reactions

1. When magnesium burns, $2Mg+O_2\to 2MgO$, each magnesium atom loses two electrons to become ${Mg}^{2+}$.
   1. Each oxygen atom gains two electrons to become $O^{2-}$, so oxygen is reduced while magnesium is oxidised.
2. Here the two definitions agree, because magnesium gains oxygen and also loses electrons.
3. In electrolysis, positive ions gain electrons at the cathode, so they are reduced.
   1. Negative ions lose electrons at the anode, so they are oxidised.

3.3.5 Dilute and concentrated, weak and strong acids

Dilute and concentrated describe how much acid is dissolved

1. The concentration of an acid is the amount of acid in each cubic decimetre of solution, measured in ${mol/dm}^3$.
   1. Hydrochloric acid of $2.0{mol/dm}^3$ holds twice as much acid in each ${dm}^3$ as hydrochloric acid of $1.0{mol/dm}^3$.
2. Adding water dilutes an acid, because the same amount of acid is spread through a larger volume.
   1. Making $100{cm}^3$ of acid up to $1000{cm}^3$ with water makes its concentration ten times smaller.
3. Dilute and concentrated are comparative words, so $0.1{mol/dm}^3$ acid is dilute compared with $5{mol/dm}^3$ acid.

Strong and weak describe how fully an acid ionises

1. When an acid dissolves, its molecules split up to release hydrogen ions, and this process is called ionisation.
2. Every HCl molecule in hydrochloric acid ionises, so it is a strong acid: $$HCl(aq)\to H^{+}(aq)+{Cl}^{-}(aq)$$
   1. Sulfuric acid and nitric acid are also fully ionised in water.
3. Only a small fraction of the molecules in ethanoic acid, the acid in vinegar, ionise, so it is a weak acid: $${CH}_{3}COOH(aq)\rightleftharpoons {CH}_3{COO}^{-}(aq)+H^{+}(aq)$$
   1. The reversible arrow shows that most ethanoic acid stays as whole molecules, with ions forming and joining back together all the time.
4. Citric acid and carbonic acid are other examples of weak acids.

An acid's strength and concentration are separate ideas

1. A concentrated strong acid, such as $5{mol/dm}^3$ hydrochloric acid, has a lot of acid in each ${dm}^3$ and all of it is ionised.
2. A dilute strong acid, such as $0.01{mol/dm}^3$ hydrochloric acid, has little acid in each ${dm}^3$ but all of it is still ionised.
3. A concentrated weak acid, such as $5{mol/dm}^3$ ethanoic acid, has a lot of acid in each ${dm}^3$ but only a small fraction of it is ionised.
4. A dilute weak acid, such as vinegar, has little acid in each ${dm}^3$ and only a small fraction of it is ionised.
5. Diluting an acid lowers its concentration, but a weak acid stays weak and a strong acid stays strong.

At equal concentration, strong acids release more hydrogen ions

1. Hydrochloric acid and ethanoic acid, both at $1{mol/dm}^3$, contain the same amount of acid in each ${dm}^3$.
2. Hydrochloric acid is fully ionised, so it has a much higher concentration of hydrogen ions.
   1. Its pH is therefore lower, about pH 0 compared with about pH 2 for the ethanoic acid.
3. Magnesium ribbon fizzes more vigorously in the hydrochloric acid, because more hydrogen ions are available to react with it.
4. These differences come from the degree of ionisation, since the two acids have the same concentration.

3.3.6 Hydrogen ion concentration and pH

Hydrogen ion concentration decides the pH

1. The square brackets in $[H^{+}]$ stand for the concentration of hydrogen ions, measured in ${mol/dm}^3$.
2. The higher the $[H^{+}]$, the lower the pH and the more acidic the solution.
3. The lower the $[H^{+}]$, the higher the pH.
4. Pure water contains a tiny concentration of hydrogen ions, $1\times 10^{-7}{mol/dm}^3$, which gives it pH 7.
   1. For whole-number values the pH matches the power of ten, so $[H^{+}]=1\times 10^{-3}{mol/dm}^3$ gives pH 3.

Hydrogen ions in neutral, acidic and alkaline solutions

1. In a neutral solution at pH 7, the concentrations of hydrogen ions and hydroxide ions are equal.
2. An acidic solution has more hydrogen ions than hydroxide ions, so its pH is below 7.
3. An alkaline solution has fewer hydrogen ions than hydroxide ions, so its pH is above 7.
   1. Even at pH 12 some hydrogen ions are present, at $1\times 10^{-12}{mol/dm}^3$.
4. Adding acid raises the $[H^{+}]$ and lowers the pH, while adding alkali removes hydrogen ions and raises the pH.

Each tenfold concentration increase lowers pH by one

1. Multiplying $[H^{+}]$ by ten lowers the pH by one.
   1. Going from $1\times 10^{-4}$ to $1\times 10^{-3}{mol/dm}^3$ gives ten times more hydrogen ions, and the pH falls from 4 to 3.
2. Dividing $[H^{+}]$ by ten raises the pH by one.
   1. Diluting a strong acid tenfold with water divides its $[H^{+}]$ by ten, so pH 1 becomes pH 2.
3. A change of two pH units means a factor of $10\times 10=100$ in $[H^{+}]$, and three units means a factor of $1000$.

pH differences give the hydrogen ion concentration factor

1. Subtract the smaller pH from the larger one to find the number of steps.
2. Raise ten to that power to find the factor, so a difference of three gives $10^3=1000$.
3. The solution with the lower pH has the higher hydrogen ion concentration.
4. The same method works in reverse, so a solution with $1000$ times fewer hydrogen ions has a pH three units higher.