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Chapter 14 – Acids and Bases

Explore the history of acids and bases, from ancient times to modern definitions. Learn about the old definitions and the more useful Bronsted-Lowry theory. Understand the concept of acid-base conjugate pairs and the strength of different acids.

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Chapter 14 – Acids and Bases

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  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 12th 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 17th 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 18th 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- H2O Arrhenius 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) + H2O(l)  H3O+(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, H3O+ and H2O are also a conjugate pair. • A conjugate pair is two species which differ by a proton. • For the reaction between HCl and H2O, the conjugate pairs are shown as: HCl(g) + H2O(l) H3O+(aq) + Cl-(aq) Blue = bases Red = acids

  8. The H+ ion in Water • A hydrogen ion (or proton) in solution is represented as H3O+(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:H2SO4(aq) + H2O(l)  HSO4-(aq) + H3O+(aq) Eqn 2: O2-(aq) + H2O(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 (HCO3-) ion reacts according to the equations: HCO3-(aq) + H2O(l)  H2CO3(aq) + OH-(aq) HCO3-(aq) + H2O(l)  CO32-(aq) + H3O+(aq) • Since HCO3- 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. CH3COOH(l) + H2O(l) CH3COO-(aq) + H3O+(aq)

  16. Strong Bases • The ionic compound sodium oxide (Na2O) dissociates in water, releasing sodium ions (Na+) and oxide ions (O2-). • 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, HNO3, CH3COOH. • Diprotic acids: can donate two protons and include H2SO4, H2CO3, • Triprotic acids: can donate three protons and include H3PO4, H3BO3.

  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 (H2SO4) 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: H2SO4(l) + H2O(l)  HSO4-(aq) + H3O+(aq) STAGE 2: HSO4-(aq) + H2O(l)  SO42-(aq) + H3O+(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 H2SO4 and the second and third proton from H3PO4. • 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 CH3COOH 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: H2O(l) + H2O(l) H3O+(aq) + OH-(aq)

  26. Acidic, Basic and Neutral Solutions cont… • Acidic solutions contain a greater concentration of H3O+ than OH-. • Neutral solutions contain equal concentrations of H3O+ and OH-. • Basic solutions contain a lower concentration of H3O+ than OH-.

  27. Measuring Acidity • [H3O+] x [OH-] = 10-14M2 • Pure water is neutral so [H3O+] = [OH-] • If either the [H3O+] or [OH-] in an aqueous solution is increased, the other must decrease proportionally. • At 25°C, a solution is: Acidic if [H3O+]>10-7M and [OH-]<10-7M Neutral if [H3O+] = 10-7M = [OH-] Basic if [H3O+]<10-7M and [OH-]>10-7M

  28. Acidity Example • In a 5.6x10-6M HNO3, solution at 25°C, calculate the concentration of: • H3O+ ions HNO3 is a strong acid and ionises completely to produce 5.6x10-6M of H+ ions. b. OH- ions [H3O+] x [OH-] = 10-14 5.6x10-6 x [OH-] = 10 [OH-] = 10-14/5.6x10-6 [OH-] = 1.79 x 10-9M

  29. The pH Scale • This scale is a useful way of indicating the acidity of a solution. • pH = -log10[H3O+] • 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)  Ba2+(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

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