CHEMICAL CHANGE: ACIDS AND BASES Learner Note: The different models are important in this chapter. General equations of acids reacting are important, these are often asked again in Grade 12, especially the reactions of an acid and carbonates and hydrogen carbonates. Ensure that you understand neutralisation and associated terms and indicators. SECTION A: TYPICAL EXAM QUESTIONS Question 1: 10 minutes (Understand the difference between strong and weak acids and bases as well as concentrated and dilute acids and bases. Strong refers to reactivity, concentration refers to amount of water added dilution.) 1.1 Hydrochloric acid is a strong acid. What does this mean? (2) 20g of calcium carbonate reacts with excess dilute sulphuric acid. (Know common formulae of acids and bases and how to balance reactions.) 1.2 Provide a balanced equation for this reaction. (3) 1.3 Calculate the volume of carbon dioxide produced at STP. (6) 1.4 What is a standard solution? (2) (This question is frequently asked) 1.5 What is a Bronsted-Lowry acid? (2) (15) Question 2: 20 minutes (Taken from DoE Exemplar 2007) The stomach secretes gastric juice, which contains hydrochloric acid. The gastric juice helps with digestion. Sometimes there is an over production of acid, leading to heartburn and indigestion. Ant-acids, such as milk of magnesia, can be taken to neutralize the excess acid. Milk of magnesia is only slightly soluble in water and has a formula Mg(OH) 2. (Acids neutralize bases to produce a salt and water.) 2.1 Write a balanced chemical reaction to show how the ant-acid reacts with the acid.(3) 2.2 The directions on the bottle recommend that children under the age of 12 years take one teaspoon of milk of magnesia, where as adults can take two teaspoons. Briefly explain why the dosages are different. (2) 2.3 Why is it not recommended to take an overdose of ant-acid in the stomach? Refer to the hydrochloric acid concentration in the stomach. (2) (Bases are opposites of acids, bases can be as harmful as acids.) In an experiment, 25 cm 3 of a standard solution of sodium carbonate of concentration 0,1 mol.dm -3 was used to neutralise 35 cm 3 of a solution of hydrochloric acid. 2.4 Write a balanced equation for this reaction. (4) 2.5 Calculate the concentration of the acid. (5) (16) Page 1 of 10
Learner Note: The reactions of acids with bases (hydroxides), carbonates and hydrogen carbonates must be well understood. These equations are needed for the writing of balanced equations to calculate concentrations in neutralisation reactions. SECTION B: SOLUTIONS AND HINTS Question 1: (Learn the difference between strong and weak well!) 1.1 A strong acid dissociates almost completely in water, has a strong tendency to donate a proton 1.2 CaCO 3 + H 2 SO 4 CaSO 4 + H 2 O + CO 2 (All equations must be balanced before doing calculations.) 1.3 n CaCO3 = m = 20 = 0,2 mol Mr 100 n CO2 = 0,2 mol V = n x Vm = 22,4 x 0,2 = 4,48 dm 3 1.4 A solution of known concentration 1.5 A proton donor Question 2: 2.1 Mg(OH) 2 + HCl 2 H 2 O + MgCl 2 2.2 A greater volume of acid in the adult stomach requires a greater amount of ant-acid to neatralise it 2.3 The hydrochloric acid is required for digestion if the concentration is too low it can slow down the digestive process. If too much ant-acid is present it is a base, it can have a harmful effect on the stomach as it is naturally acidic. (The volume and concentration of an acid is important during neutralisation.) 2.4 Na 2 CO 3 + 2 HCl 2 NaCl + H 2 O + CO 2 na cava 2.5 nb cbvb (Learn this formula, neutralisation relies on the ratio of acid to base or vice versa do you see the ratio in the equation?) 2 ca35 1 0,1 25 = 0,14 mol.dm -3 Page 2 of 10
Learner Note: It is important to understand this section well. There is always a question on this section, practise the calculations. SECTION C: ADDITIONAL CONTENT NOTES TYPES OF REACTIONS Acid-Base Reactions Acid ph less than 7 Strong acid ph very close to 0 Base ph greater than 7 Strong base ph very close 14 Neutral ph 7 Concentration: measure of the amount of water in an acid or base. Examples of acids H 2 SO 4, HNO 3, HCl, etc Examples of bases- NaOH, KOH, etc The Models 1. Lowry-Bronstead model (proton transfer model) Acid proton donor (giver of protons) Base proton acceptor (taker of protons) Protons H + ions An acid base reaction, or proton transfer reaction is called protolysis. E.g. : HCl + H 2 O H 3 O + + Cl - H 3 O + ion polonium ion 2. The Arrhenius model Acids liberate hydrogen ions (dissociate) (H+) when dissolved in water. Bases form hydroxide ions (OH-) ions when dissolved in water. HNO 3 dissociates : HNO 3 H + - + NO 3 NaOH dissociates : NaOH Na + + OH - Acid Base Strength Strong acids ionize almost completely in solution and form a high concentration of hydrogen ions. Equilibrium lies far to the right The equilibrium constant for the above reaction Ka = [H 3 O + ] [NO 3 - ] [HNO 3 ] [H 3 O + ] [NO 3 - ] is greater than [HNO 3 ], therefore the Ka value is high, and this implies that strong acids have a high Ka value. Page 3 of 10
Weak acids ionize only partially in solution and form a low concentration of hydrogen ions. Example CH 3 COOH CH 3 COO - + H + Equilibrium lies far to the left The equilibrium is constant for the above reaction Ka = [H + ] [CH 3 COO - ] CH 3 COOH] [H + ] [CH 3 COO - ] is less than [CH 3 COOH], therefore the Ka value is low, and this implies that weak acids have a low Ka value. The same rule applies to bases, strong bases have a high Kb value i.e. the equilibrium lies far to the right because they ionise almost completely and weak bases have a low Kb value i.e. the equilibrium lies to the left because they ionize only partially. Acid-base pairs The liberation of a proton can be represented by the equation: HA A - + H + acid proton The reverse reaction is: H + + A - HA The negative ion, A -, takes up a proton and is therefore a proton acceptor or base. A - + H + HA base proton In a closed system, both these reactions occur and an acid-base equilibrium is established. HA A - + H + Acid base proton HA (the acid) and A - (the base) differs from one another in respect to one proton : one is obtained from the other by the removal or addition of a single proton. Such an acid and a base is called an acid-base pair. A - is called the conjugate base of the acid HA; HA is called the conjugate acid of the base A - A strong acid has a weak conjugate base, and a strong base has a weak conjugate acid. Page 4 of 10
Ampholytes Substances that can react both as an acid and as a base are amphiprotic and are called ampholytes. H 2 O is an example : As an acid : H 2 O H + + OH - As a base : H 2 O + H + H 3 O + Other examples : HSO - 4, HCO - 3, H 2 PO - 4, etc Protolytic reactions Any reaction in which a proton transfer takes place is called a protolytic reaction. Conjugate acid-base pairs form when protolysis takes place. Example 1 : HCl + NH 3 NH 4 + + Cl - acid conjugate base acid conjugate base Strong acids dissolved in water : HCl + H 2 O H 3 O + + Cl - If an acid donates one proton, e.g. HCl monoprotic. If an acid donates two protons, e.g. H 2 SO 4 diprotic. If an acid donates three protons, e.g. H 3 PO 4 triprotic. Acids that have more than one proton to donate are called polyprotic acids. Please Note : When an acid is strong, its conjugate base is weak and when a base is strong, its conjugate acid is weak. Protolysis (Hydrolysis) of salts The salt of a strong acid and a strong base forms a neutral solution. - the metal chlorides, metal sulphates and metal nitrates which dissolve in water, do not undergo protolysis in water because they are all weak conjugate bases of strong acids. Example NaCl + H 2 O Na + + Cl - The salt of a strong base and a weak acid dissolves in water to form a basic solution. Page 5 of 10
Example : NaHCO 3 Na + + HCO 3 - HCO 3 - + H 2 O H 2 CO 3 + OH - Strong base The salt of a strong acid and a weak base dissolves in water to form an acidic solution. Example NH 4 Cl NH 4 + + Cl - NH 4 + + H 2 O H 3 O + + NH 3 Strong acid Hydrolysis (protolysis) of Water H 2 O + H 2 O H 3 O + + OH - The equilibrium constant for the above reaction, K = [H 3 O + ] [OH - ] [H 2 O] [H 2 O] [H 2 O] = 1, therefore the equilibrium constant for water : Kw = [H 3 O + ] [OH - ] Kw is known as the ionic product of water and has an experimentally determined value of 1,0 x 10-14 at 25 o C. The ph scale The function ph is defined as similarly poh = -log [OH - ] ph + poh = 14 ph = -log[h 3 O + ] or ph = log [H + ] and because [H 3 O + ] [OH - ] = 1,0x10-14, it follows therefore that The relationship between [H 3 O + ] and ph is shown on the ph scale below: (Please note that it is not necessary to give units with ph.) Page 6 of 10
[H 3 O + ] 10 0 10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 10-10 10-11 10-12 10-13 10-14 ph 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 ACIDIC NEUTRAL BASIC Neutralization and Acid base titrations: Neutralization is the reverse of hydrolysis: in hydrolysis water is a reactant, in neutralization water is a product. An acid or base of known concentration (a standard solution) is added to an acid or base of unknown concentration in the presence of an indicator until the indicator changes colour, this is called the point of neutralization or the end point. Volumetric analysis is a method of determining the concentration of solutions by measuring the volumes of solutions which react with each other. Acid-base titrations are volumetric analysis applied to acids and bases. The stage at which the reaction has reached completion is called the end-point or the equivalence point. This is when chemical equivalent quantities of the acid and base react with each other. This process is known as neutralisation. The equivalence point is indicated by the colour change of the suitable indicator. Indicators The choice of an indicator depends on the ph of the solution at the end point of the reaction. Recall : (a) The salt of a strong acid and a strong base forms a neutral solution. (b)the salt of a strong base and a weak acid dissolves in water to form a basic solution. (c)the salt of a strong acid and a weak base dissolves in water to form an acidic solution. Titration of ph range Suitable indicator Strong acid + strong base Strong acid + weak base 4-7 Bromothymol blue, methyl orange, phenolphthalein 3,5 6.5 Methyl orange Weak acid + strong base 7,5 10,5 Phenolphthalein Page 7 of 10
A typical indicator reacts as a weak acid and can be represented as follows : HIn + H 2 O In - + H 3 O + a 1 b 2 b 1 a 2 red blue The acid and its conjugate base are usually differently coloured. When the indicator is added to the acid, the concentration of the H 3 O + ions is high, which shifts the equilibrium to the left, the colour of solution is red. If the acid is neutralized by the base, the concentration of H 3 O + ions is reduced, the equilibrium shifts to the right, the colour will be blue. Standard Solutions A standard is a solution of known concentration. Concentration is defined as number of moles in a given volume C = n V n = m therefore C = m Mr, Mr x V Units: mass : g, volume: dm 3, concentration: mol/dm 3 Calculating the concentration of the unknown in acid-base titrations Consider the following reaction Acid + base salt and water e.g. HCl + NaOH NaCl + H 2 O the number of moles of acid n a = C a V a the number of moles of the base n b = C b V b Therefore n a = C a V a n b C b V b n in this case is the molar ratio in the balanced equation i.e. the coefficient next to the reactant or product. Page 8 of 10
SECTION D: HOMEWORK Learner Note: Many questions are asked on the terminology, concentration and type of indicator used. Learn well and practise the calculations. Question 1: 10 minutes 1.1 What element is found in every acid? (1) Provide 1.2 two examples of weak acids. (2) 1.3 two examples of strong acids. (2) 1.4 one example of a weak base. (2) 1.5 one example of a strong base. (2) (9) Question 2: 10 minutes An amount of 12,3cm 3 of a 0,03mol/dm 3 solution of sulphuric acid reacts with 36,9cm 3 of a sodium hydroxide solution. 2.1 Provide a suitable indicator for this reaction. (1) 2.2 Write a balanced equation for this reaction. (4) 2.3 Calculate the concentration of the sodium hydroxide solution. (5) (10) Question 3: 5 minutes Will the following salts form solutions that are acidic, basic or neutral. 3.1 Na 2 CO 3 (1) 3.2 NH 4 Cl (1) 3.3 NaCl (1) 3.4 Provide reasons for your answers. (3) Question 4: 10 minutes 4. Provide equations for the reactions of sulphuric acid with the following 4.1. sugar (3) 4.2. blue copper sulphate crystals (3) (12) Page 9 of 10
SECTION E: SOLUTIONS TO SESSION 19 HOMEWORK Question 1: 1.1. Heat of reaction - is the difference between the energy of the products and the energy of the reactants. 1.2. Endothermic reaction takes in energy, products have more energy than the reactants 1.3. Activation energy - the energy hill which must be overcome by the addition of this amount of energy before a reaction can take place. Question 2: 2.1. Exothermic 2.2. Endothermic 2.3. Exothermic Question 3: 3.1. the sun 3.2. flame 3.3. flame Question 4: 4.1. X-axis course of reaction Y-axis potential energy 4.2. Eproducts < Ereactants 4.3. activated complex - temporary, unstable, high-energy composition of atoms, which represents a transition state between reactants and the products. 4.4. negative 4.5. exothermic The SSIP is supported by Page 10 of 10