Chemistry 5072

Summary notes

Unit 7 The Chemistry and uses of Acids Bases and Salts7.1 Characteristic Properties of acids and basesAcids An acid is a compound that ionizes and produces hydrogen (H+) ions when

dissolved in water.

A pure acid (no water present) consists of covalent molecules. If a pure acid is dissolve in an organic solvent, such as chloroform or alcohol, it will not show acidic properties as there are no hydrogen ions produced.

The three common inorganic acids are sulfuric acid, nitric acid and hydrochloric acid. A weak organic acid is ethanoic acid (CH3COOH).

Base A base is a metallic oxide or metallic hydroxide. Eg. MgO, CuO, FeO, Mg(OH)2.

Alkali An alkali is a soluble base that ionises and produces hydroxide (OH-) ions when dissolved in water. Hydroxide ions are responsible for alkaline properties.All alkalis are soluble bases. Not all bases are alkalis.Examples of alkalis are aqueous ammonia, NH3 (aq), sodium hydroxide, NaOH and potassium hydroxide, KOH. Aqueous ammonia is a weak alkali while sodium hydroxide is a strong alkali.

pH scale The pH is a measure of how acidic or alkaline a solution is in water. The pH scale is numbered from 0 to 14: 0 7 14

neutral acidity increases alkalinity increase

pH 7 is neutral. (Eg. Pure water) pH less than 7 is acidic. The smaller the pH, the more acidic the

solution, the more hydrogen ions it contains. Eg. Hydrochloric acid, HCl, ethanoic acid, CH3COOH.

pH more than 7 is alkaline. The bigger the pH, the more alkaline the solution, the more hydroxide ions it contains. Eg. Aqueous ammonia, NH3(aq) and sodium hydroxide, NaOH.

1

Universal indicator

Universal indicator is a mixture that gives different colours at different pH. The pH of solution can be measured by dipping a piece of Universal paper in solution, then compare with the colour chart. Need to know colour changes of phenolphthalein, litmus and methyl orange.Universal indicator is made of several organic dyes, that can produce different colour changes, when either an acid or alkali is added together with the indicator in water.

Difference between strong and weak acids

When water is added, the acid molecules ionize to form H+ ions.

In a strong acid, all the acid molecules are completely (100%) ionized when dissolved in water. For example: HCl (aq) H+(aq) + Cl-(aq)You illustrate a strong acid by a single arrow in the chemical equation.

In a weak acid, the acid molecules are only partially or weakly ionized when dissolved in water.

For example: CH3COOH(aq) CH3COO-(aq) + H+(aq)Ethanoic acid

You illustrate a weak acid by a double reversible reaction arrow in the equation. The double arrow indicates that the reaction can take place in the forward and backward directions. CH3COOH(aq) molecules, CH3COO- (aq) ions and H+(aq) ions are present in a solution of ethanoic acid.

A strong acid contains a higher concentration of H+ ions than a weak acid of the same concentration.

2

Reactions of acids

3 important reactions

(a) Acid + metals Acids (HCl and H2SO4) react with reactive metals (eg. Mg, Zn and Fe)

to form a salt and hydrogen only. K, Na and Ca react explosively – not allowed in the lab. Al metal is covered by oxide layer. So reaction is slow at first.

Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)

Copper does not react with dilute acids, because copper is an unreactive metal (see Reactivity series in chapter of Metals).

(b) Acid - carbonate Acids react with metal carbonates to produce a salt, carbon dioxide and water only.

Insoluble carbonates cannot react to produce insoluble salts. Insoluble salts are produced by mixing two (soluble) solutions together.

CaCO3(s) + 2HCl(aq) CaCl2(aq) + H2O(l) + CO2(g)CaCO3(s) + H2SO4 (aq) No reaction as CaSO4 is insoluble.

(c) Acid – alkali (neutralisation reaction) Acids neutralise base/alkalis to form salt and water only.

NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)

State symbols should always be added to every equation.

Reactions of bases

(a) Reaction with ammonium salts to produce of ammonia gas.e.g. 2NaOH(aq) + (NH4)2SO4(aq) Na2SO4(aq) + 2H2O(l) + 2NH3(g)

How do you test for ammonia gas?

Neutralisation reaction

Bases react with acids to form salt and water only. This is neutralisation.The ionic equation for all neutralisation reactions: H+ + OH- H2OThis shows that it is a reaction between hydrogen ions and hydroxide ions to form water.E.g. NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l) CuO (s) + 2HCl (aq) CuCl2 (aq) + H2O (l)Work out how you can obtain the ionic equation stated above by cancelling out the spectator ions. Spectator ions do not take part in the reaction.

Effects of pH on soil

Plants need soil of a suitable pH for good growth. Much soil may be acidic due to acid rain.To reduce acidity, calcium oxide (quicklime) or calcium hydroxide (slaked lime) can be added. The reaction that happens is called neutralisation.

3

Oxides Both metals and non metals can react with oxygen to form oxides.Oxides can be classified as acidic, neutral, basic or amphoteric.

Acidic : Non-metallic oxides - CO2 ,SO2 , NO2, P4O10

Neutral : Non-metallic oxides - CO , NO, H2OBasic : Metallic oxides - Na2O, MgO, CaO etcAmphoteric : Metallic oxides - Al2O3, PbO, ZnO (both acidic and basic

properties) Only oxides of Al, Pb and Zn form amphoteric oxides at O levels. The other metallic oxides are basic oxides.

Acidic oxides dissolve in water to produce an acid.Some basic oxides are soluble in water, eg. Sodium oxide, potassium oxide. The other basic oxides are insoluble in water. Basic oxides react with acids to form salt and water only.Amphoteric oxides react with both acids and alkalis to form salt and water.

7.2 Preparation of salts

Preparation of salts

The method used to prepare a salt depends on the solubility of the salt.

Solubility of salts Solubility of ionic compounds in water.

(a) All Group I and ammonium compounds are soluble.

(b) All nitrates are soluble.

(c) All chlorides are soluble, except AgCl, PbCl2(d) All sulfates are soluble, except PbSO4, CaSO4, BaSO4

(e) All carbonates are insoluble, except Group I carbonates and (NH4)2CO3

(f) All oxides and hydroxides of metals are insoluble, except Group I and ammonium compounds and Ca(OH)2 is slightly soluble.

(g) All lead compounds are insoluble except for lead(II) nitrate and lead(II) ethanoate.

Precipitation method

Insoluble salt: Both reactants must be soluble in water.1. Mix the two aqueous solutions of the reactants together.2. Filter the mixture to collect the residue (precipitate).3. Wash the precipitate with distilled water and dry with filter paper.

4

Titration method Soluble salt: By titration only for salts of K+, Na+ and NH4+ .

1. Titrate 25.0cm3 of alkali with acid, with a suitable indicator, usually methyl orange or phenolphthalein. Do not use universal indicator as there are too many colour changes.

2. Note the volume of acid required for neutralisation.3. Repeat the titration without the indicator.4. Evaporate to obtain a saturated solution, cool and crystallise to obtain

the pure crystals of the salt. Dry the crystals between filter paper.(See chapter 1).

Acid + insoluble base/carbonate/metal

Soluble salt: (NOT by titration, ie. For all other soluble salts which are NOT ammonium, sodium and potassium salts) Acid with insoluble oxide/carbonate/metal.

1. Add excess insoluble base/carbonate/metal to warm acid, with stirring, till no more solid can dissolve.

2. Filter to remove excess insoluble base/carbonate/metal.3. Evaporate the filtrate to saturation.4. Cool the saturated solution for crystals to form.5. Filter the crystals, wash with cold distilled water and dry with filter

paper.

5

Redox Reactions (Brief Notes)

A tip for learning this topic is to be clear about the distinctions between oxidation and reduction processes. Reduction can be approximately viewed as the opposite of oxidation.

Oxidation Oxidation is defined as:1. The gain in oxygen or the loss of hydrogen2. The increase in oxidation state3. The loss of electrons of a substance in a chemical reaction.

Reduction Reduction is defined as:1. The loss of oxygen or the gain in hydrogen2. The decrease in oxidation state3. The gain of electrons of a substance in a chemical reaction.

Examples of oxidation reactions

1. carbon + oxygen gas carbon dioxideC + O2 CO2

As carbon has gained oxygen atoms from oxygen gas molecules, carbon is being oxidized.2. 2CO + O2 2CO2

C in CO has an oxidation no. of +2. C in CO2 has an oxidation no. of +4. Carbon has increased in oxidation number from CO to CO2. Therefore, we say that carbon monoxide has been oxidized to carbon dioxide. Note that we ignore the stoichiometry coefficients in the chemical equations in redox oxidation numbers calculations.3. Cu Cu2+ + 2e Copper has lost 2 electrons to form copper(II) ion. Hence, we say that copper has been oxidized to copper(II) ions.

Examples of reduction reactions

1. Cl2 + H2 2HClChlorine has gained hydrogen atoms to form hydrogen chloride. Hence, we say that chlorine is being reduced.2. Br2 + H2 2HBrBromine has an oxidation number of zero since it is an element. Br in HBr has an oxidation number of -1. Since there is a decrease in oxidation number of bromine to hydrogen bromide, we say that bromine has been reduced to hydrogen bromide.3. Cl2 + 2e 2Cl-

Chlorine has gained two electrons to form 2 chloride ions. Hence, we say that chlorine has been reduced to chloride ions.

6

. Redox reaction refers to both oxidation and reduction occurring in the same chemical reaction.Types of redox reaction1. Metal + Dilute acids2.All combustion reactions ( burning of fuels)3. Displacement reaction. Takes place when a more reactive element displaces a less reactive ion

from its solution. E.g. (a) Chlorine gas displaces iodide ions to become chloride ions and

iodine, I2 molecules. (b) Zinc displaces Cu2+ ions from CuSO4 to form Zn2+ and form

reddish-brown Cu metal.4. Extraction of less reactive metals using Blast furnace. E.g. iron from iron(III) oxide by carbon monoxideThe list is not exhaustive.

Oxidising agent(usually coloured)

When a substance is reduced, it acts as an oxidising agent. Also, an oxidising agent oxidizes another substance. This is usually observed as a colour change. E.g. acidified potassium dichromate(VI), acidified potassium manganate(VII), chlorine

Acidified potassium dichromate (VI) CrO42- Cr3+

Orange to green

Acidified potassium permanganate (VII)

MnO4- Mn2+

Purple to colourless

Reducing agent(usually colourless)

When a substance is oxidised, it acts as a reducing agent. Also, a reducing agent reduces another substance.E.g. potassium iodide, carbon monoxide, hydrogen, metals (high in the reactivity series)Potassium iodide 2I I2

Colourless to brownHydrogen peroxide (H2O2 is both a reducing and oxidising agent)

H2O2 O2

Oxidation states You need to know how to calculate oxidation state (number) .Oxidation is the increase in oxidation number of a species.Reduction is the decrease in oxidation number of a species.

(1) Oxidation number of an element is zero. Eg. Oxidation number of iron, carbon is zero.

(2) Oxidation number of a simple ion equals the charge of the ion, eg. Oxidation number of iron(II) ions, Fe2+, is +2.

(3) Oxygen usually has an oxidation number of -2 in compounds.(4) Hydrogen usually has an oxidation number of +1 in compounds.(5) SO4

2- ion : Oxidation number of sulfur in sulfate = +6 :+6+(-2)4 = -2 (charge on sulfate ion)

7