1. Aqueous Stuff Aqueous Stuff

2. Reactions Between Ions Ionic compounds, also called salts, consist of both positive and negative ions When an ionic compound dissolves in water, it dissociates to aqueous ions What happens when we mix aqueous solutions of two different ionic compounds? if two of the ions combine to form a water-insoluble compound, a precipitate will form otherwise no physical change will be observed

3. Reactions Between Ions Example: suppose we prepare these two aqueous solutions if we then mix the two solutions, we have four ions present; of these, Ag + and Cl - react to form AgCl(s) which precipitates

4. Reactions Between Ions we can simplify the equation for the formation of AgCl by omitting all ions that do not participate in the reaction net ionic equation the simplified equation is called a net ionic equation; it shows only the ions that react spectator ions ions that do not participate in a reaction are called spectator ions

5. Reactions Between Ions In general, ions in solution react with each other when one of the following can happen two of them form a compound that is insoluble in water two of them react to form a gas that escapes from the reaction mixture as bubbles, as for example when we mix aqueous solutions of sodium bicarbonate and hydrochloric acid an acid neutralizes a base one of the materials can oxidize another

6. Reactions Between Ions Following are some generalizations about which ionic solids are soluble in water and which are insoluble all compounds containing Na +, K +, and NH 4 + are soluble in water all nitrates (NO 3 - ) and acetates (CH 3 COO - ) are soluble in water most chlorides (Cl - ) and sulfates (SO 4 2- ) are soluble; exceptions are AgCl, BaSO 4, and PbSO 4 most carbonates (CO 3 2- ), phosphates (PO 4 3- ), sulfides (S 2- ), and hydroxides (OH - ) are insoluble in water; exceptions are LiOH, NaOH, KOH, and NH 4 OH which are soluble in water

7. Oxidation-Reduction Oxidation: Oxidation: the loss of electrons Reduction: Reduction: the gain of electrons Oxidation-reduction (redox) reaction: Oxidation-reduction (redox) reaction: any reaction in which electrons are transferred from one species to another

8. Oxidation-Reduction Example: if we put a piece of zinc metal in a beaker containing a solution of copper(II) sulfate some of the zinc metal dissolves some of the copper ions deposit on the zinc metal the blue color of Cu 2+ ions gradually disappears In this oxidation-reduction reaction zinc metal loses electrons to copper ions copper ions gain electrons from the zinc

9. Oxidation-Reduction we summarize these oxidation-reduction relationships in this way

10. Oxidation-Reduction Although the definitions of oxidation (loss of electrons) and reduction (gain of electrons) are easy to apply to many redox reactions, they are not easy to apply to others for example, the combustion of methane An alternative definition of oxidation-reduction is oxidation: oxidation: the gain of oxygen or loss of hydrogen reduction: reduction: the loss of oxygen or gain of hydrogen

11. Oxidation-Reduction using these alternative definitions for the combustion of methane

12. Oxidation-Reduction Five important types of redox reactions combustion: combustion: burning in air. The products of complete combustion of carbon compounds are CO 2 and H 2 O. respiration: respiration: the process by which living organisms use O 2 to oxidize carbon-containing compounds to produce CO 2 and H 2 O. The importance of these reaction is not the CO 2 produced, but the energy released. rusting: rusting: the oxidation of iron to a mixture of iron oxides bleaching: bleaching: the oxidation of colored compounds to products which are colorless batteries: batteries: in most cases, the reaction taking place in a battery is a redox-reaction

13. Heat of Reaction In almost all chemical reactions, heat is either given off or absorbed example: the combustion (oxidation) of carbon liberates 94.0 kcal per mole of carbon oxidized Heat of reaction: Heat of reaction: the heat given off or absorbed in a chemical reaction exothermic reaction: exothermic reaction: one that gives off heat endothermic reaction: endothermic reaction: one that absorbs heat heat of combustion: heat of combustion: the heat given off in a combustion reaction; all combustion reactions are exothermic

14. Properties of Acids & Bases Reaction with metal oxides strong acids react with metal oxides to give water plus a salt

15. Properties of Acids & Bases Reaction with carbonates and bicarbonates strong acids react with carbonates to give carbonic acid, which rapidly decomposes to CO 2 and H 2 O strong acids react similarly with bicarbonates

16. Properties of Acids & Bases Reaction with ammonia and amines any acid stronger than NH 4 + is strong enough to react with NH 3 to give a salt

17. Self-Ionization of Water pure water contains a very small number of H 3 O + ions and OH - ions formed by proton transfer from one water molecule to another the equilibrium expression for this reaction is we can treat [H 2 O] as a constant = 55.5 mol/L

18. Self-Ionization of Water ion product of water, K w combining these constants gives a new constant called the ion product of water, K w in pure water, the value of K w is 1.0 x 10 -14 this means that in pure water

19. Self-Ionization of Water the product of [H 3 O + ] and [OH - ] in any aqueous solution is equal to 1.0 x 10 -14 for solutions as well. for example, if we add 0.010 mole of HCl to 1 liter of pure water, it reacts completely with water to give 0.010 mole of H 3 O + in this solution, [H 3 O + ] is 0.010 or 1.0 x 10 -2 this means that the concentration of hydroxide ion is

20. pH and pOH we commonly express these concentrations as pH, where pH = -log [H 3 O + ] we can now state the definitions of acidic and basic solutions in terms of pH acidic solution: acidic solution: one whose pH is less than 7.0 basic solution: basic solution: one whose pH is greater than 7.0 neutral solution: neutral solution: one whose pH is equal to 7.0

21. pH and pOH just as pH is a convenient way to designate the concentration of H 3 O +, pOH is a convenient way to designate the concentration of OH - pOH = -log[OH - ] the ion product of water, K w, is 1.0 x 10 -14 taking the logarithm of this equation gives pH + pOH = 14 thus, if we know the pH of an aqueous solution, we can easily calculate its pOH

22. pH of Salt Solutions When some salts dissolve in pure water, there is no change in pH from that of pure water Many salts, however, are acidic or basic and cause a change the pH when they dissolve We are concerned in this section with basic salts and acidic salts

23. pH of Salt Solutions Basic salt: Basic salt: raises the pH as an example of a basic salt is sodium acetate when this salt dissolves in water, it ionizes; Na + ions do not react with water, but CH 3 COO - ions do the position of equilibrium lies to the left nevertheless, there are enough OH - ions present in 0.10 M sodium acetate to raise the pH to 8.88

24. pH of Salt Solutions Acidic salt: Acidic salt: lowers the pH an example of an acidic salt is ammonium chloride chloride ion does not react with water, but the ammonium ion does although the position of this equilibrium lies to the left, there are enough H 3 O + ions present to make the solution acidic

25. Acid-Base Titrations Titration: Titration: an analytical procedure in which a solute in a solution of known concentration reacts with a known stoichiometry with a substance whose concentration is to be determined

26. Acid-Base Titrations An acid-base titration must meet these requirement 1. we must know the equation for the reaction so that we can determine the stoichiometric ratio of reactants to use in our calculations 2. the reaction must be rapid and complete end point 3. there must be a clear-cut change in a measurable property at the end point (when the reagents have combined exactly) 4. we must have precise measurements of the amount of each reactant

27. Acid-Base Titrations As an example, let us use 0.108 M H 2 SO 4 to determine the concentration of a NaOH solution requirement 1: requirement 1: we know the balanced equation requirement 2: requirement 2: the reaction between H 3 O + and OH - is rapid and complete requirement 3: requirement 3: we can use either an acid-base indicator or a pH meter to observe the sudden change in pH that occurs at the end point of the titration requirement 4: requirement 4: we use volumetric glassware

28. Acid-Base Titrations experimental measurements doing the calculations

29. pH Buffers pH buffer: pH buffer: a solution that resists change in pH when limited amounts of acid or base are added to it a pH buffer as an acid or base “shock absorber” a pH buffer is common called simply a buffer the most common buffers consist of approximately equal molar amounts of a weak acid and a salt of the conjugate base of the weak acid for example, if we dissolve 1.0 mole of acetic acid and 1.0 mole of its conjugate base (in the form of sodium acetate) in water, we have an acetate buffer

30. pH Buffers How an acetate buffer resists changes in pH if we add a strong acid, such as HCl, added H 3 O + ions react with acetate ions and are removed from solution if we add a strong base, such as NaOH, added OH - ions react with acetic acid and are removed from solution

31. pH Buffers The effect of a buffer can be quite dramatic consider a phosphate buffer prepared by dissolving 0.10 mole of NaH 2 PO 4 (a weak acid) and 0.10 mole of Na 2 HPO 4 (the salt of its conjugate base) in enough water to make 1 liter of solution

32. pH Buffers Buffer pH Buffer pH if we mix equal molar amounts of a weak acid and a salt of its conjugate base, the pH of the solution will be equal to the pK a of the weak acid if we want a buffer of pH 9.14, for example, we can mix equal molar amounts of boric acid (H 3 BO 3 ), pK a 9.14, and sodium dihydrogen borate (NaH 2 BO 3 ), the salt of its conjugate base

33. pH Buffers Buffer capacity depends both its pH and its concentration

34. Henderson-Hasselbalch Eg. Henderson-Hasselbalch equation: Henderson-Hasselbalch equation: a mathematical relationship between pH, pK a of the weak acid, HA concentrations HA, and its conjugate base, A - It is derived in the following way taking the logarithm of this equation gives

35. Henderson-Hasselbalch Eg. multiplying through by -1 gives -log K a is by definition pK a, and -log [H 3 O + ] is by definition pH; making these substitutions gives rearranging terms gives

36. Henderson-Hasselbalch Eg. Example: Example: what is the pH of a phosphate buffer solution containing 1.0 mole of NaH 2 PO 4 and 0.50 mole of Na 2 HPO 4 dissolved in enough water to make 1.0 liter of solution

37. Henderson-Hasselbalch Eg. Example: Example: what is the pH of a phosphate buffer solution containing 1.0 mole of NaH 2 PO 4 and 0.50 mole of Na 2 HPO 4 in enough water to make one liter of solution Solution Solution the equilibrium we are dealing with and its pK a are substituting these values in the H-H equation gives