2. Acids and Just the Bases Mrs. Herrmann

3. Concentrations of Solutions Concentration—quantitive; a measure of the amount of solute in a given amount of solvent or solution. Molarity (M)—number of moles solute per liter of solution M = moles of solute M = moles of solute liters of solution liters of solution

4. Aqueous Solutions Two Confusing and Subtle Definitions Dissociation: The separation of ions that occurs when an ionic compound dissolves. (Ions that are already present separate from one another.) Dissociation: The separation of ions that occurs when an ionic compound dissolves. (Ions that are already present separate from one another.) H 2 O H 2 O Example: NaCl(s)  Na + (aq) + Cl - (aq) Example: NaCl(s)  Na + (aq) + Cl - (aq) Ionization: The process by which ions are formed from solute molecules (polar) by the action of the solvent. (Ions form where none were present.) Ionization: The process by which ions are formed from solute molecules (polar) by the action of the solvent. (Ions form where none were present.) H 2 O H 2 O Example: HCl  H + (aq) + Cl - (aq) Example: HCl  H + (aq) + Cl - (aq)

5. Two Easy Definitions Electrolyte : Substances that dissolve in water to form ions Electrolyte : Substances that dissolve in water to form ions Strong electrolyte—almost ALL the dissolved compound exists as ions in aqueous solution Weak electrolyte—relatively small amount of the dissolved compound exists as ions in aqueous solution Note: degree of ionization or dissociation indicates whether an electrolyte is strong or weak NOT the concentration of the solution! Nonelectrolyte : substances that dissolve in water but do not form any ions Nonelectrolyte : substances that dissolve in water but do not form any ions

6. The Hydronium Ion H + ion attracts other molecules or ions so strongly that it normally does not exist alone. Ionization of monoprotic or polyprotic acid in water best described as a transfer of a proton from the acid to a water molecule to form the Hydronium ion. H 3 O + Highly exothermic which generally provides enough energy to ionize a molecular solute Example: H 2 O(l) + HCl(g)  H 3 O + (aq) + Cl - (aq)

7. Acids and Bases Properties Acids Bases Acids Bases Sour taste Bitter taste Sour taste Bitter taste Change the color of acid- Change the color of acid- Change the color of acid- Change the color of acid- base indicators base indicators base indicators base indicators Blue litmus turns red Red litmus turns blue Blue litmus turns red Red litmus turns blue React with active metals Dilute aqueous solutions React with active metals Dilute aqueous solutions to release hydrogen gas feel slippery to release hydrogen gas feel slippery React with bases to form React with acids to form React with bases to form React with acids to form salts and water salts and water salts and water salts and water Some conduct electric Conduct electric current Some conduct electric Conduct electric current current current

8. Arrhenius Acids and Bases Arrhenius Acid: a chemical compound that increases the concentration of hydrogen ions, H+, in aqueous solution. Arrhenius Base: a substance that increases the concentration of hydroxide ions, OH-, in aqueous solution.

9. Brønsted-Lowry Acids and Bases Brønsted-Lowry Acid—is a molecule or ion that is a proton donor. Brønsted-Lowry Acid—is a molecule or ion that is a proton donor. Brønsted-Lowry Base—is a molecule or ion that is a proton acceptor. Brønsted-Lowry Base—is a molecule or ion that is a proton acceptor. Example: HCl + NH 3  NH 4 + + Cl - Example: HCl + NH 3  NH 4 + + Cl - where HCl is the Brønsted-Lowry Acid and where HCl is the Brønsted-Lowry Acid and NH 3 is the Brønsted-Lowry Base NH 3 is the Brønsted-Lowry Base

10. Conjugate Acids and Bases Conjugate Acid—is the species that is formed when a Brønsted-Lowry base gains a proton. Conjugate Acid—is the species that is formed when a Brønsted-Lowry base gains a proton. Conjugate Base—is the species that that remains after a Brønsted-Lowry acid has given up a proton. Conjugate Base—is the species that that remains after a Brønsted-Lowry acid has given up a proton. HF(aq) + H 2 0(l)  F - (aq) + H 3 O + (aq) HF(aq) + H 2 0(l)  F - (aq) + H 3 O + (aq) acid base conjugate base conjugate acid acid base conjugate base conjugate acid

11. Lewis Acids and Bases Lewis Acid is an atom, ion, or molecule that accepts an electron pair to form a covalent bond. Lewis Base is an atom, ion, or molecule that donates an electron pair to form a covalent bond. (an anion) Example: BF3(aq) + F-(aq) BF4-(aq) Where BF3 is the Lewis Acid and F- i is the Lewis Base

12. Ionization of Water In the self-ionization of water: In the self-ionization of water: H 2 O(l) + H 2 O(l)  H 3 O + (aq) + OH - (aq) H 2 O(l) + H 2 O(l)  H 3 O + (aq) + OH - (aq) at constant temperature the concentrations of hydronium ion and hydroxide ions remain constant in water and dilute aqueous solutions, so that the ionization constant of water is: K w = [H 3 O + ][OH - ] K w = [H 3 O + ][OH - ] K w = (1.0 x 10 -7 ) (1.0 x 10 -7 ) K w = (1.0 x 10 -7 ) (1.0 x 10 -7 ) K w = 1.0 x 10 -14 K w = 1.0 x 10 -14

13. pH and [H 3 O + ] pH—French word “pouvoir hydrogène” which means “hydrogen power” pH—French word “pouvoir hydrogène” which means “hydrogen power” pH of a solution is defined as the negative of the common logarithm of the hydronium ion concentration pH of a solution is defined as the negative of the common logarithm of the hydronium ion concentration pH = -log[H 3 O + ] pH = -log[H 3 O + ] [H 3 O + ] ranges from 1 to 10 -14 so [H 3 O + ] ranges from 1 to 10 -14 so pH ranges from 0 to 14 pH ranges from 0 to 14

14. pOH and [OH - ]  pOH of a solution is defined as the negative of the common logarithm of the hydroxide ion concentration, [OH - ]  pOH = -log[OH - ]  [OH - ] ranges from 1 to 10 -14 so  pOH ranges from 0 to 14

15. pH and pOH Ranges So for a neutral solution of water K w = [H 3 O + ][OH - ] K w = [H 3 O + ][OH - ] 1.0 x 10 -14 = (1.0 x 10 -7 ) (1.0 x 10 -7 ) so -log K w = -log[H 3 O + ] + -log[OH - ] so -log K w = -log[H 3 O + ] + -log[OH - ] pK w = pH + pOH pK w = pH + pOH 14 = 7 + 7 14 = 7 + 7 But if the [H 3 O + ] > [OH - ] then pH < pOH pH < pOH and pH 7 and is defined as acidic and pH 7 and is defined as acidic if the [H 3 O + ] < [OH - ] then if the [H 3 O + ] < [OH - ] then pH > pOH pH > pOH and pH > 7 and pOH 7 and pOH < 7 and is defined as basic

16. Calculating [H 3 O + ] and [OH - ] from pH Remember pH = -log [H3O+] so log [H3O+] = -pH [H3O+] = antilog(-pH) [H3O+] = 10-pH

17. That’s all !