1. The Arrhenius Theory of acids and bases The theory Acids are substances which produce hydrogen ions in solution. Acids are substances which produce hydrogen ions in solution. Bases are substances which produce hydroxide ions in solution. Bases are substances which produce hydroxide ions in solution. Neutralisation happens because hydrogen ions and hydroxide ions react to produce water. In the sodium hydroxide case, hydrogen ions from the acid are reacting with hydroxide ions from the sodium hydroxide - in line with the Arrhenius theory. However, in the ammonia case, there don't appear to be any hydroxide ions! THEORIES OF ACIDS AND BASES

2. The Bronsted-Lowry Theory of acids and bases An acid is a proton (hydrogen ion) donor. An acid is a proton (hydrogen ion) donor. A base is a proton (hydrogen ion) acceptor A base is a proton (hydrogen ion) acceptor The relationship between the Bronsted-Lowry and Arrhenius theory. The Bronsted-Lowry theory doesn't go against the Arrhenius theory in any way - it just adds to it. Hydroxide ions are still bases because they accept hydrogen ions from acids and form water. An acid produces hydrogen ions in solution because it reacts with the water molecules by giving a proton to them.When hydrogen chloride gas dissolves in water to produce hydrochloric acid, the hydrogen chloride molecule gives a proton (a hydrogen ion) to a water molecule. A co- ordinate (dative covalent) bond is formed between one of the lone pairs on the oxygen and the hydrogen from the HCl. Hydroxonium ions, H 3 O +, are produced. THEORIES OF ACIDS AND BASES

3. When the acid, HA, loses a proton it forms a base, A -. When the base, A -, accepts a proton back again, it obviously refoms the acid, HA. These two are a conjugate pair. Ammonia is a base because it is accepting hydrogen ions from the water. The ammonium ion is its conjugate acid - it can release that hydrogen ion again to reform the ammonia.The water is acting as an acid, and its conjugate base is the hydroxide ion. The hydroxide ion can accept a hydrogen ion to reform the water.Looking at it from the other side, the ammonium ion is an acid, and ammonia is its conjugate base. The hydroxide ion is a base and water is its conjugate acid. Conjugate pairs

4. A substance which can act as either an acid or a base is described as being amphoteric. Amphoteric substances

5. The theory An acid is an electron pair acceptor. An acid is an electron pair acceptor. A base is an electron pair donor. A base is an electron pair donor. The relationship between the Lewis theory and the Bronsted-Lowry theory Lewis bases It is easiest to see the relationship by looking at exactly what Bronsted- Lowry bases do when they accept hydrogen ions. Three Bronsted-Lowry bases we've looked at are hydroxide ions, ammonia and water, and they are typical of all the rest. The Lewis Theory of acids and bases

6. Lewis acids Lewis acids are electron pair acceptors. In the above example, the BF 3 is acting as the Lewis acid by accepting the nitrogen's lone pair. On the Bronsted-Lowry theory, the BF 3 has nothing remotely acidic about it. This is an extension of the term acid well beyond any common use. A final comment on Lewis acids and bases: A Lewis acid is an electron pair acceptor. A Lewis acid is an electron pair acceptor. A Lewis base is an electron pair donor. A Lewis base is an electron pair donor. Conjugate pairs

7. Fundamental Role in synthesis, analytical behavior, reactivity including phsyiological behavior. General Equation: Strengths of Acids: 1. Inverse relationship rule Acid/Base Chemistry

8. 2. Periodic effects: a. Acidity increases from left to right: CH < NH < OH < FH b. Acidity increases from top to bottom: HI > HBr > HCl > HF; H 2 S > H 2 O Acid Strength depends on the ability of the conjugate base to stabilize a negative charge. a) Presence of electronegative elements: Acid/Base Chemistry

9. 2. Resonance stabilization: Acid/Base Chemistry

10. Measure of Acid/Base Strength Aqueous systems - pKa, pKb Scale: 0 to 14, neutral = 7 Acid/Base Chemistry

11. Base strength is often reported in terms of pKa which is the strength of the conjugate acid. Example: Pyridine, pKa = 5.19 (Merck Index) Rank the following compounds in order of relative basicity: Resonance stabilization Cyclohexylamine vs. aniline Resonance stabilization Cyclohexylamine vs. aniline Acid/Base Chemistry

12. Basicity trends for amines: Amides are very weakly basic (pKa = -1) Solvation effects: Explanation: (CH 3 ) 3 NH + is less solvated in H 2 O Acid/Base Chemistry

13. Amino acids as zwitterions Zwitterions in simple amino acid solutions. An amino acid has both a basic amine group and an acidic carboxylic acid group. THE ACID-BASE BEHAVIOUR OF AMINO ACIDS

14. Amino acids as zwitterions Zwitterions in simple amino acid solutions. An amino acid has both a basic amine group and an acidic carboxylic acid group. There is an internal transfer of a hydrogen ion from the -COOH group to the -NH 2 group to leave an ion with both a negative charge and a positive charge. This is called a zwitterion. THE ACID-BASE BEHAVIOUR OF AMINO ACIDS

15. Adding an alkali to an amino acid solution If you increase the pH of a solution of an amino acid by adding hydroxide ions, the hydrogen ion is removed from the -NH 3 + group. You could show that the amino acid now existed as a negative ion using electrophoresis. THE ACID-BASE BEHAVIOUR OF AMINO ACIDS

16. Adding an acid to an amino acid solution If you decrease the pH by adding an acid to a solution of an amino acid, the -COO - part of the zwitterion picks up a hydrogen ion. This time, during electrophoresis, the amino acid would move towards the cathode (the negative electrode). THE ACID-BASE BEHAVIOUR OF AMINO ACIDS

17. Adding an acid to an amino acid solution If you decrease the pH by adding an acid to a solution of an amino acid, the -COO - part of the zwitterion picks up a hydrogen ion. This time, during electrophoresis, the amino acid would move towards the cathode (the negative electrode). The pH at which this lack of movement during electrophoresis happens is known as the isoelectric point of the amino acid. This pH varies from amino acid to amino acid. THE ACID-BASE BEHAVIOUR OF AMINO ACIDS

18. Why is phenol acidic? Unlike alcohols (which also contain an -OH group) phenol is a weak acid. A hydrogen ion can break away from the - OH group and transfer to a base. For example, in solution in water: Phenol is a very weak acid and the position of equilibrium lies well to the left. Phenol can lose a hydrogen ion because the phenoxide ion formed is stabilised to some extent. The negative charge on the oxygen atom is delocalised around the ring. The more stable the ion is, the more likely it is to form. THE ACIDITY OF PHENOL

19. One of the lone pairs on the oxygen atom overlaps with the delocalised electrons on the benzene ring. THE ACIDITY OF PHENOL

20. This overlap leads to a delocalisation which extends from the ring out over the oxygen atom. As a result, the negative charge is no longer entirely localised on the oxygen, but is spread out around the whole ion. THE ACIDITY OF PHENOL