Chapter 13 – Introducing Acids & Bases

1. Week 3, Lesson 3 Chapter 13 – Introducing Acids & Bases

2. Acids • Acids are commonly used in our homes. • Many foods contain acids to enhance flavour or as a preservative. • The sour taste of some food substances is due to the presence of acid. • In industry solutions of acids are used extensively to produce a wide range of products, such as fertilisers, drugs, explosives and plastics. • Acids are also used to clean metal surfaces before use.

3. Acids cont… • The three most commonly used acids in the lab are hydrochloric acid, sulfuric acid and nitric acid. • Acids can cause severe problems if misused. • Eg, acid rain, acid contamination.

4. Properties of Acids • Change the colour of some indicators. • Tend to be corrosive • Taste sour • React with bases • Have a relatively low pH.

5. Bases • Acids react with bases. • In solution, bases sometimes taste bitter and have a slippery feel. • They react with some plant extracts to counteract the effect of acids. • If a base is continually added to a solution of an acid, the properties of the acid slowly disappear. • The same goes for if an acid is added to a basic solution. • The acid and base are said to have neutralised each other. • Bases are effective cleaners because they react with fats and oils to produce water-soluble soaps.

6. Properties of Bases • Turn litmus blue • Have a slippery feel • Are caustic • React with acids • Have a relatively high pH

7. Safety with Acids and Bases • Acids and bases should be treated with caution and you should avoid these solutions coming in contact with your skin and eyes. • You should always wear safety goggles and lab coats. • Concentrated sulfuric acid is a viscous liquid and accidents can happen if water is added to the acid in order to produce dilute solutions. • The ionisation/hydrolysis of sulfuric acid is an energy-releasing reaction. • A common method of neutralising an acid spill is to add sodium hydrogen carbonate powder.

8. Indicators • One of the characteristic properties of acids and bases is their ability to change to colour of certain plant extracts. • Litmus is a dye obtained from lichen. • In the presence of acid, litmus turns red. • Such plant extracts are called indicators. • Indicators are often extracted from plant dyes and are themselves acids or bases. • They change from one colour in acid to another in basic solution. • Common indicators include methyl orange, phenolpthalein and litmus. • Universal indicator is a mixture of many indicators and changes through a range of colours to easily establish the pH of a solution. • The indicators that undergo a single colour change are also used for many analyses. • pH meters are often used in laboratories for more accurate determination of solution acidity and are not affected by coloured solutions.

9. Indicators and pH Ranges

10. Reactions Involving Acids and Bases • Acids react with many metals particularly those found in the main groups of the periodic table, although they also react with several transition metals. • These reactions typically produce a solution of a metal salt and hydrogen gas. • Salts are compounds usually made up of a metal cation and a non-metal anion. • The salt produced will depend on the acid used in the reaction. • Acids also react with many compounds such as metal hydroxides and metal carbonates. • A salt is again produced with each of these, together with water. • In the case of metal carbonates, carbon dioxide is also formed.

11. General Reaction Types involving Acids • There are numerous ways in which acids and bases react. • It is possible to group some reactions together on the basis of the reactants involved and the products formed. • The following are six of the more common reaction types.

12. Acid + Reactive Metal  Salt + Hydrogen • Reactive metals include Ca, Mg, K but not Cu, Ag or Au. • When dilute acids are added to main group metals, and some transition metals, bubbles of hydrogen gas are released, and a salt is formed. • For example, the reaction between dilute hydrochloric acid and zinc metal can be represented by the equation: 2HCl(aq) + Zn(s)  ZnCl2(aq) + H2(g)

13. Cont… • This reaction can also be represented by an ionic equation. • In an aqueous solution the hydrochloric acid is ionised and the zinc chloride is dissociated. • The equation can therefore be written as: 2H+(aq) + 2Cl-(aq) + Zn(s)  Zn2+(aq) + 2Cl-(aq) +H2(g) However the chloride is the only spectator ion. The ionic equation is therefore: 2H+(aq) + Zn(s)  Zn2+(aq) + H2(g)

14. Acid + Metal Hydroxide  Salt + Water • Metal hydroxides include NaOH, Ca(OH)2 and Mg(OH)2. • The hydroxide ions from metal hydroxides readily react with the H+ ion from acid. • The products are a salt and water. • For example, the reaction between solutions of sulfuric acid and sodium hydroxide can be represented as: H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + H2O(l)

15. Cont… • The sulfuric acid is ionised in solution and both sodium hydroxide and sodium sulfate are ionic, and therefore dissociate. • Water however is a covalent molecular substance that does not ionise to a significant extent. • So the equation becomes: 2H+(aq) + SO42-(aq) + 2Na+(aq) + 2OH-(aq) 2Na+(aq) + SO42-(aq) + 2H2O(l) The ionic equation is therefore: H+(aq) + OH-(aq)  H2O(l)

16. Acid + Metal Oxide  Salt + Water • Metal oxides include Na2O, MgO, CaO and ZnO. • Metal oxides are usually basic since they contain the oxide ion. • Water soluble oxides tend to form the hydroxide ion. • The reaction between dilute nitric acid and solid calcium oxide can be represented by the equation: 2HNO3(aq) + CaO(s)  Ca(NO3)2(aq) + H2O(l)

17. Cont… • The calcium oxide is solid, so the ions are not dissociated. The nitric acid is ionised in solution and therefore dissociates. • The equation becomes: 2H+(aq) + 2NO3-(aq) + CaO(s)  Ca2+(aq) + 2NO3-(aq) + H2O(l) The nitrate ions are the spectator ions so the ionic equation is: 2H+(aq) + CaO(s)  Ca2+(aq) + H2O(l)

18. Acid + Metal Carbonate  Salt + Water + Carbon Dioxide • Metal carbonates include Na2CO3, MgCO3 and CaCO3. • Acids reacting with metal carbonates produce carbon dioxide gas with a salt and water. • For example the reaction between a solution of nitric acid and solid magnesium carbonate can be represented by the equation: 2HNO3(aq) + MgCO3(s)  Mg(NO3)2(aq) + H2O(l) + CO2(g) The nitrate ions are the spectator ions so the ionic equation is: 2H+(aq) + MgCO3(s)  Mg2+(aq) + H2O(l) + CO2(g)

19. Acid + Metal Hydrogen Carbonate  Salt + Water + Carbon Dioxide • Metal hydrogen carbonates include NaHCO3, KHCO3 and Ca(HCO3)2. • Acids added to metal hydrogen carbonates (aka bicarbonates) produce a salt, water and carbon dioxide. • For example, the reaction between solutions of hydrochloric acid and sodium hydrogen carbonate can be represented by the equation: HCl(aq) + NaHCO3(aq)  NaCl(aq) + H2O(l) + CO2(g) The sodium and chloride are spectator ions in this reaction so the ionic equation is: H+(aq) + HCO3-(aq)  H2O(l) + CO2(g)

20. Acidic Oxide (non-metal oxide) + Base  Salt + Water • Acidic oxides include SO2, SO3, P4O10 and CO2. • When oxides are added to water they form acidic solution. • The reaction of these with bases produce a salt and water. • For example the reaction between carbon dioxide and a solution of calcium hydroxide can be represented by the equation: CO2(aq) + Ca(OH)2(aq)  CaCO3(s) + H2O(l)

21. Cont… • The table below lists some common acidic oxides and the anions they produce in the reaction.