Chemistry Unit 2 Acids and Bases

Transcription

1 Chemistry 3202 Unit 2 Acids and Bases

2 Definitions of Acids and Bases An operational definition is one that is based on the observable properties, behaviours or uses of an entity. The earliest definitions of acids and bases were operational. Acids were defined as sour substances that dissolved metals. Bases were defined as bitter extracts of wood ash or seaweed that were good cleaners or good for making soaps. As you can see, these definitions were based on sensory perception and the uses of the substances.

3 Operational Definitions ACIDS: Taste sour. React with certain metals (Zn, Fe, etc.) to produce hydrogen gas. Turn litmus paper red. React with bases to form salts and water. Will conduct electricity.

4 BASES: Operational Definitions feel slippery or soapy Taste bitter Turn litmus paper blue Will conduct electricity. React with acids to form salts and water

5 Theoretical Definitions In 1884, Svante Arrhenius presented his Ph.D research on the concept of dissociation. He pointed out that certain species, when dissolved in water, are present as ions only. He extended this idea to acids and bases in Acids dissociate to produce hydrogen ions (H + ) in water. Bases dissociate to produce hydroxide ions (OH - ) in water. Arrhenius's definitions of acids and bases are theoretical because they explain acidity and alkalinity in terms of the species responsible for those conditions. They are not operational because actually observing these ions being formed is not possible.

6 Arrhenius Theory: Acids The key to Arrhenius' acid-base theory is the ionization equation (which is often called the dissociation equation). An acid molecule like HCl, releases a hydrogen ion and an anion when dissolved in water. In other words, when it is dissolved in water, HCl doesn't exist as HCl molecules, it separates into hydrogen ions (H + ) and chloride ions (Cl - ). Such equations can be written for other acids. The pattern is represented by this general form of the Arrhenius acid equation: where HA is any acid, H + is the hydrogen ion and A - is an anion.

7 Arrhenius Theory: Bases Arrhenius defined bases as compounds that dissociate in water to form cations and hydroxide ions. For example:the general form of the Arrhenius base equation is: Dissociation of Barium Hydroxide Dissociation of Sodium Hydroxide

8 Limitations of Arrhenius Theory Although the Arrhenius theory successfully explains the behaviour of many acids and bases, it does not explain the properties of all acids and bases. Ammonia (NH 3 ) is an example. The formula for ammonia obviously lacks the hydroxide ion, yet whenever an aqueous solution of ammonia is tested, it exhibits the properties of a base. Arrhenius proposed the formula NH 4 OH (aq) or aqueous ammonium hydroxide to accommodate the problem. He suggested that this compound dissociates in water to produce ammonium ions (NH 4+ ) and hydroxide. Note that NH 4 OH (aq) is a theoretical compound.

9 Limitations of Arrhenius Theory Another problem involves the ion that is responsible for acidity: H +. Look again at the equation for the dissociation of hydrochloric acid. HCl (aq) H + (aq) + Cl (aq) This dissociation occurs in aqueous solution, but chemists often leave out H 2 O as a component of the reaction. They simply assume that it is there. What happens if you put H 2 O into the equation? HCl (aq) + H2O (l) H + (aq) + Cl (aq) + H 2 O (l) Notice that the water is unchanged when the reaction is represented this way. However, water is a polar molecule to which H + ions will attach. The result is a hydrated molecule called hydronium. (H 3 O + (aq))

10 Limitations of Arrhenius Theory Yet another limitation of the Arrhenius theory is its inability to predict successfully whether a given aqueous compound is acidic or basic. For example, the hydrogen phosphate ion, HPO 4 2 (aq), appears to contain H + ions. You would expect that solutions containing this ion would be acidic. And yet, when tested with litmus paper, solutions that contain this ion are found to be basic. Similarly, salts that contain carbonate ions (CO 3(aq) ) have basic properties when dissolved in water. The Arrhenius theory cannot predict or explain these facts.

11 Modernizing the Arrhenius Theory A modernized Arrhenius theory acknowledges the role of water and the production of hydronium ions. HCl (aq) + H 2 O (l) H 3 O + (aq) + Cl (aq) Similarly, the dissociation of sulfuric acid may be reinterpreted in light of a modernized Arrhenius theory. H 2 SO 4 (aq) + H 2 O (l) H 3 O + (aq) + HSO 4 (aq) A modernized Arrhenius theory also acknowledges the dissociation of ammonia in water by way of a the two-step process: NH 3 (g) + H 2 O (l) NH 4 OH (aq) NH 4 OH (aq) NH 4 + (aq) + OH (aq)

12 The Brønsted-Lowry Theory of Acids and Bases In 1923, two chemists working independently of each other proposed a new theory of acids and bases. Johannes Brønsted in Copenhagen, Denmark, and Thomas Lowry in London, England, proposed what is called the Brønsted-Lowry theory of acids and bases. This theory overcame the limitations of the Arrhenius theory. Brønsted Lowry

13 The Brønsted-Lowry Theory of Acids and Bases The Brønsted-Lowry Theory of Acids and Bases states: An acid is a substance from which a proton (H + ion) can be removed. A base is a substance that can remove a proton (H + ion) from an acid.

14 The Brønsted-Lowry Theory of Acids and Bases Proton removal explains the behaviour of ammonia in water, which Arrhenius failed to do.

15 The Brønsted-Lowry Theory of Acids and Bases Proton removal also explains the behaviour of the hydrogen carbonate (bicarbonate) ion in water. Recall that Arrhenius theory predicts hydrogen carbonate to be an acid but that hydrogen carbonate has the properties of a base in water. Brønsted-Lowry theory shows how it can behave as a base in an acidbase reaction with water: H 2 O loses/donates a proton to become OH -, so it is a Brønsted-Lowry acid. HCO 3 - removes/accepts a proton from water to become H 2 CO 3, so it is a Brønsted-Lowry base.

16 Conjugate Acid-Base Pairs Two molecules or ions that are related by the transfer of a proton are called a conjugate acid-base pair. (Conjugate means linked together. ) The conjugate base of an acid is the particle that remains when a proton is removed from the acid. The conjugate acid of a base is the particle that results when the base receives the proton from the acid. It is common practice to label conjugate acid-base pairs. Ammonium is the conjugate acid of ammonia. Hydroxide is the conjugate base of water because water lost a proton to become.

17 Conjugate Acid-Base Pairs More examples.

18 Complete #1-9 p Practice