1. Chapter 14 – Acids and Bases

2. History of Acids & Bases Vinegar was probably the only known acid in ancient times. Strong acids such as sulfuric, nitric and hydrochloric acids were not discovered until after the 12 th century. Over the years, there have been many attempts to define acids and bases.

3. Old Definitions of Acids and Bases At first, acids and bases were defined in terms of their observed properties such as taste, effects on indicators and reactions with other substances. In the 17 th century, Boyle described the properties of acids in terms of taste, their action as solvents and how they changed colour of certain vegetable materials. He also noticed that alkalis (soluble bases) could reverse the effects of acids. Lavoisier, in the 18 th century, thought that acidic properties were due to the presence of oxygen. In 1810, Davy suggested that the acid properties of substances were associated with hydrogen and not oxygen. In 1887, Arrhenius defined acids as substances that produced hydrogen ions (H + ) in water while bases produced hydroxide ions (OH - ) in water. According to his theory, when acids and bases react together, the H + and OH - form water according to the equation: H + + OH -  H 2 OArrhenius called this a neutralisation reaction.

4. Definitions cont… There were, however, limitations to these theories. Arrhenius’ definition for example was restricted to acids and bases in water. One of the more useful definitions used today was first proposed by the Bronsted and Lowry Bronsted and Lowry described the reactions of acids as involving the donation of a hydrogen ion (H + ). A hydrogen ion is a hydrogen that has lost its only electron. In most cases, a hydrogen ion is a proton.

5. Bronsted-Lowry Acids and Bases According to the Bronsted-Lowry theory, a substance behaves as an acid when it donates a proton, ie H + to a base. A substance behaves as a base when it accepts a proton from an acid. Hence: – Acids are proton donors and – Bases are protons acceptors.

6. Bronsted-Lowry Acids and Bases As protons are exchanged from an acid to a base, this definition explains why acids and bases react together. In an aqueous solution of hydrogen chloride, nearly all the hydrogen chloride is present as ions – virtually no molecules of hydrogen chloride remain. This solution is known as hydrochloric acid. In this reaction, each hydrogen chloride molecule has donated a proton to a water molecule. According to the Bronsted-Lowry theory, the hydrogen chloride has acted as an acid. The water molecule has accepted a proton from the hydrogen from the hydrogen chloride, so has acted as a base. HCl(g) + H 2 O(l)  H 3 O + (aq) + Cl - (aq)

7. Acid-base Conjugate Pairs Because HCl and Cl- can be formed from each other by the loss or gain of a single proton, they are called a conjugate acid/base pair. Similarly, H 3 O + and H 2 O are also a conjugate pair. A conjugate pair is two species which differ by a proton. For the reaction between HCl and H 2 O, the conjugate pairs are shown as: HCl(g) + H 2 O(l)  H 3 O + (aq) + Cl - (aq) Blue = bases Red = acids

8. The H + ion in Water A hydrogen ion (or proton) in solution is represented as H 3 O + (aq) or more simply H + (aq) and is called the hydronium ion. The hydronium ion itself attracts more water molecules and is further hydrated. However, these water molecules are not as strongly attracted and their number is not constant.

9. Some Common Acids & Bases

10. Amphiprotic Substances Some substances can behave as either acids or bases, depending on what they are reacting with. These substances are given the name amphiprotic substances. In equation 1 below, water readily accepts a proton from sulfuric acid and acts as a base. In equation 2, water donates a proton to the oxide ion and acts as an acid. Eqn 1:H 2 SO 4 (aq) + H 2 O(l)  HSO 4 - (aq) + H 3 O + (aq) Eqn 2: O 2- (aq) + H 2 O(l)  OH - (aq) + OH - (aq)

11. Amphiprotic Substances cont… If the solute is a stronger acid than water, then water will act as a base. If the solute is a stronger base than water, then the water will act as an acid.

12. Amphiprotic Substances cont… When an amphiprotic substance is placed in water, it reacts as both an acid and a base. For example, the hydrogen carbonate (HCO 3 - ) ion reacts according to the equations: HCO 3 - (aq) + H 2 O(l)  H 2 CO 3 (aq) + OH - (aq) HCO 3 - (aq) + H 2 O(l)  CO 3 2- (aq) + H 3 O + (aq) Since HCO 3 - can act as both acid and base, it is amphiprotic. Although both reactions are possible for all amphiprotic substances in water, generally one of these reactions occurs to a greater extent. The dominant reaction can be identified by measuring the pH of the solution.

13. Acid & Base Strength Experiments show that different acid solutions of the same concentration do not have the same pH. Some acids donate a proton more readily than others. The strength of an acid is based on its ability to donate hydrogen ions. The strength of a base is based on its ability to accept hydrogen ions. Since aqueous solutions of acids and bases are most commonly used, it is convenient to use an acid’s tendency to donate a proton to water, or a base’s tendency to accept a proton, as a measure of its strength.

14. Strong Acids Acids that ionise completely in solution are called strong acids. Strong acids donate protons easily. Solutions of strong acids would contain ions and virtually no unreacted acid molecules. The most common strong acids are hydrochloric acid, sulfuric acid and nitric acid.

15. Weak Acids An acid that does not fully ionise is called a weak acid. An example of a weak acid is ethanoic acid. Only a small proportion of ethanoic acid molecules are ionised. A weak acid can be shown be the presence of reversible arrows. CH 3 COOH(l) + H 2 O(l) CH 3 COO - (aq) + H 3 O + (aq)

16. Strong Bases The ionic compound sodium oxide (Na 2 O) dissociates in water, releasing sodium ions (Na + ) and oxide ions (O 2- ). The oxide ions react completely with the water, accepting a proton to form hydroxide ions (OH - ). The oxide ion is an example of a strong base. Strong bases accept protons easily.

17. Weak Bases Ammonia is a covalent molecular compound that ionises in water by accepting a proton. This ionisation process can be represented by the equation: NH3(aq) + H2O(l)NH4 + (aq) + OH - (aq) Only a small proportion of ammonia molecules ionise. This is shown in the equation by the presence of reversible arrows. Ammonia is a weak base in water.

18. Polyprotic Acids Some acids are capable of donating more than one proton from each molecule and are said to be polyprotic. The number of hydrogen ions an acid can donate depends on the structure of the acid. Monoprotic acids: can donate only one proton and include HCl, HF, HNO 3, CH 3 COOH. Diprotic acids: can donate two protons and include H 2 SO 4, H 2 CO 3, Triprotic acids: can donate three protons and include H 3 PO 4, H 3 BO 3.

19. Polyprotic Acids cont… Polyprotic acids do not donate all protons at once, but do so in steps when reacting with a base. Sulfuric acid (H 2 SO 4 ) is diprotic, meaning it has two protons that it can donate to a base. A diprotic acid ionises in two stages, for example: STAGE 1: H 2 SO 4 (l) + H 2 O(l)  HSO 4 - (aq) + H 3 O + (aq) STAGE 2: HSO 4 - (aq) + H 2 O(l)  SO 4 2- (aq) + H 3 O + (aq)

20. Polyprotic Acids cont… When added to a base stronger than water, a weak acid will ionise to a greater extent. For example, a strong base such as OH - will accept a second proton from H 2 SO 4 and the second and third proton from H 3 PO 4. Similarly a weak base will ions to a greater extent if added to a strong acid. Sometimes there are more hydrogens in a molecule than can actually be donated. For example CH 3 COOH contains four hydrogen and yet will only donate one. Only the hydrogen involved in the polar OH - bond is donated. In general each hydrogen ion that is donated by an acid molecule is involved in a polar bond.

21. Relative Strengths of Acid Base Pairs

22. Strength vs. Concentration It is important that the terms strong and weak are not confused with the terms concentrated and dilute. Concentrated and dilute describe the amount of acid or base dissolved in a given volume of solution. The terms strong and weak describe how readily an acids donates, or base accepts a proton.

23. Strength vs. Concentration cont…

24. Qualitative vs. Quantitative Terms such as concentrated and dilute, or weak and strong are qualitative, or descriptive terms. Solutions can be more accurately described by stating concentration in mol/L or g/L. This is a quantitative description.

25. Acidic, Basic and Neutral Solutions The acidity of a solution is a measure of the concentration of hydrogen ions present. The higher the concentration of hydrogen ions, the more acidic the solution. Water has the ability to act as either an acid or a base. Pure water undergoes self ionisation to a small extent with allows it to conduct electricity slightly. This can be represented by the equation: H 2 O(l) + H 2 O(l) H 3 O + (aq) + OH - (aq)

26. Acidic, Basic and Neutral Solutions cont… Acidic solutions contain a greater concentration of H 3 O + than OH -. Neutral solutions contain equal concentrations of H 3 O + and OH -. Basic solutions contain a lower concentration of H 3 O + than OH -.

27. Measuring Acidity [H 3 O + ] x [OH - ] = 10 -14 M 2 Pure water is neutral so [H 3 O + ] = [OH - ] If either the [H 3 O + ] or [OH - ] in an aqueous solution is increased, the other must decrease proportionally. At 25°C, a solution is: Acidic if [H 3 O + ]>10 -7 M and [OH - ]<10 -7 M Neutral if [H 3 O + ] = 10 -7 M = [OH - ] Basic if [H 3 O + ] 10 -7 M

28. Acidity Example In a 5.6x10 -6 M HNO 3, solution at 25°C, calculate the concentration of: a.H 3 O + ions HNO 3 is a strong acid and ionises completely to produce 5.6x10 -6 M of H + ions. b. OH - ions [H 3 O + ] x [OH - ] = 10 -14 5.6x10 -6 x [OH - ] = 10 [OH - ] = 10 -14 /5.6x10 -6 [OH - ] = 1.79 x 10 -9 M

29. The pH Scale This scale is a useful way of indicating the acidity of a solution. pH = -log 10 [H 3 O + ] The pH of a solution decreases as the concentration of hydrogen ions increases. Acidic solutions have a pH<7 Basic solutions have a pH>7 Neutral solutions have a pH=7

30. Calculating pH Example 1… What is the pH of a solution in which [H + ] = 0.0135M pH = -log[H + ] pH = -log(0.0135) pH = -(-1.87) pH = 1.87

31. Calculating pH Example 2… What is the pH of a 0.0050M of Ba(OH) 2 ? Step 1: Find concentration of H + Ba(OH) 2 (aq)  Ba 2+ (aq) + 2OH - (aq) Ba(OH)2 is completely dissociated in water and each mole of Ba(OH) 2 dissociates to release 2 moles of OH - ions So, [OH-] = 2 x [Ba(OH) 2 ] = 2 x 0.0050 = 0.010M Since [H + ] x [OH - ] = 10 -14 [H + ] x 0.010 = 10 -14 [H + ] = 10 -14 / 0.010 [H + ] = 10 -12 Step 2: Calculate the pH pH = -log[H + ] = -log(10 -12 ) = 12

32. Calculating the Concentration of H + in a solution of a given pH [H + ] = 10 -pH If the pH is 5.00, what is the [H + ]? [H + ] = 10 -5 = 0.0001M