Unit 4 Acids,Bases and Salts

1. Unit 4 Acids,Bases and Salts

2. Quiz 1 topics • Definitions • Theories • Properties • Hebden IV.1-IV.4 (pages 109-118)

3. Properties of an Acid • An acid is any substance that turns blue litmas paper red • Reacts with metals to produce hydrogen • Conduct electricity • Sour taste • React Exothermically with Bases • Functional Definition of an Acid: An acid is any substance that turns blue litmas paper red (Grade 10 definition)

4. Properties of A Base • Any substance that turns red litmus paper Blue • Are electrolytes • Are Slippery (Corrosive to skin) • Taste Bitter • React Exothermically with Acids • Functional Definition: Any substance that turns red litmus paper Blue (Grade 10)

5. Problems with the functional Definitions of Acids &Bases • The functional definitions of acids and bases allowed chemists to classify compounds, however, they are not very useful because they provide no theory as to how they work.

6. Common Acids & bases • Acids • Sulphuric acid • Hydrochloric Acid (Muriatic) • Nitric Acid • Acetic Acid • Bases • Sodium Hydroxide • Potassium Hydroxide (potash) • Ammonia (Ammonium Hydroxide)

7. Arrhenius Theory/Definition • An Acid is any substance which releases H+ in water • Base is any substance which releases OH- in water • A Salt is the neutralization product which results when an acid and a base react. Acid + Base => Water + Salt (A Salt is any ionic compound which is neither an acid nor a base)

8. Problems with the Arrhenius theory • Arrhenius’s theory and definitions arte better than the functional definitions because they explain how acids and bases interact. • The main problem with the Arrhenius’s theory is that there are examples of compounds that follow the functional definitions of a base but not the Arrhenius definition (Example NH3)

9. Classification of Acids • Monoprotic Acid, is an acid which can supply only one proton. • Diprotic Acid, is an acid that can supply two protons • Triprotic Acid, is an acid that can supply three protons • Polyprotic acid, is a general term for an acid that can supply mote than 1 proton.

10. Bronsted-Lowery Theory(Grade 12 Definitions) • An Acid is any substance which Donates a Proton to another substance • A Base is any substance which Accepts a Proton from another Substance • Examples

11. Amphiprotic compounds • Amphiprotic compounds are compounds that can act either as an acid or a base. I.e. as a proton acceptor or a proton donor. Most amphiprotic compounds contain hydrogen ions that can be donated and have a negative charge (I.e are anions) (exception is water) • Examples

12. Conjugate Acid/base pairs • A conjugate acid-base pair (conjugate pair) is a pair of chemical species that differ by only one proton. • Conjugate Acid: is the member of the pair that has the extra proton (H+ ion) • Conjugate base: is the member of the pair that lacks the extra proton.

13. Quiz 2 • Conjugate acid/base pairs • Strengths of acids and bases • Ka, Kb, and Kw’s • Hebden IV.5- IV.10 (120-133)

14. Strength of Acids and Bases • Acid strength is a measure of how willing a compound is to donate a proton (H+ ion) • The strength of a base is a measure of how willing a compound is to accept a proton • ** In water, the strongest acid is the H3O+ ion. In water the strongest base is the OH- ion. This is called the Leveling effect • A Strong Acid is a compound that donates a proton and dissociates 100% in water • A Strong base is a compound is 100% ionized in water and will accept a proton

15. Dissociation Expression for H2O • Keq =[products]/[reactants]= [H3O+][OH-]=1.00x10-14 =Kw (at 25C) • [H3O+]=[OH-]=1.00x10-7 • We say H2O is neutral because the [H3O+]=[OH-] • We say a solution is acidic if [H3O+]>[OH-] • We say a solution is basic if [H3O+]<[OH-] • Note: KW does depend on temperature as can be seen by the dissociation eqs.

16. Ka & Kb for determining the relative strengths of Acids and Bases • Using the fact that water is neutral I.e., [H3O+]=[OH-] and amphiprotic, we can establish the relative strengths of acids and bases by comparing them to H2O • In order to be able to use Ka as a measure of acid strength, the same base must be used in both cases. Water is chosen because it neutral

17. Quiz 3 • pH, pOH, [H3O+], [OH-] • For strong acids and bases

18. Ka Calculations • Since Strong Acids/Bases dissociate 100% so pH, pOH, etc calculations are straight forward. • This is not the case for Weak Acids or bases. • In order to calculate H3O+ or OH- concentrations Ka/Kb data must be used

19. Types of Ka problems • Given: [HA], Ka Calculate pH or [H3O+] • Given: [HA], pH or [H3O+] Calculate Ka • Given: Ka, pH or [H3O+] Calculate [HA] • HA is a weak acid

20. Kb Problems • Kb problems are very similar to Ka problems except you will be writing Kb expressions and you will have to calculate Kb from Ka • Given: [A-], Kb Calculate pOH or [OH-] • Given: [A-], pOH or [OH-] Calculate KB • Given: Kb, pOH or [OH-] Calculate [A-] • A- is a weak base

21. Making Acids from Anhydrides • You can prepare an acid from an anhydride by hydrating it (i.e adding water) • Ex1: Covalently bonded molecules containg oxygen • Ex adding water to an Acidic anhydride

22. Hydrolysis & Anhydrides • An anhydride is a chemical compound formed by removing water. Salts formed in acid/base reactions are anhydrides because: • Hydrolysis, is when an Anhydride reacts with water Anhydride rx. Hydrolysis rx.

23. Anhydrides • Anhydrides can be prepared by driving off water, or, any acid/base rx will produce an anhydride salt. • Examples • ** not all anhydride salts can under go hydrolysis • The conjugates of strong acids & strong bases can not under go hydrolysis and are called spectator ions in solution

24. Spectator Ions(will not under go Hydrolysis) • Cations: all group 1 and group 2 metals • Anions: 5 anions found at top of Acid Strength table ( ClO4-, I-, Br-, Cl-, NO3-)

25. Titrations Quiz • Titration Curves • How to choose and indicator • How an indicator works • Primary standard vs. Secondary standard • Titration problems

26. Titration Problems • Titration: is an experimental process of determining when a balanced reaction has reached its correct stoichiometric ratio • Grade 11- finding the concentration of an unknown acid/base (N,C,V problem) • Grade 12 problems • Percent purity problem • 34.786 g of NaHSO4(impure) is diluted to 250.0 ml of solution, 25.0 ml of this was titrated with 26.77 ml of NaOH with a Molality of 0.9974 what is the % purity of the NaHSO4

27. Titration Problems • Determining molar mass • 3.2357 g of a monoprotic acid is diluted to 250.00 ml 25.00 is titrated with 16.94 mL of a 0.1208 M solution NaOH what is the molar mass of the compound? • Partial Neutralizations involving multi-protic acids • An equivalence point is reached by reacting 25.00 mL of a 0.11255 M NaOH solution with 38.74 ml of a 0.02700 M H4P2O7 How many protons have been removed? What is the balanced reaction?

28. Primary vs. Secondary Standards • Primary standard: a substance which can be obtained in a pure and stable form (which does not absorb water or CO2 from the air) and from which a solution of exactly known concentration can be prepared • Secondary standard: are derived from primary standards by preforming titrations to determine the exact concentration (example NaOH solution)

29. How an Indicator works • An indicator is a week acid in which its conjugate base is a different color • NOTE: The pH of the end point color of an Indicator is = Pka or the Indicator(week acid)

30. Buffers • Definition: A buffer is a solution containing appreciable amounts of a weak acid and its conjugate pair • Example: HF +H2O <=> F- + H3O+ ( 1 M) (1 M) To prepare a Buffer you make a solution containing equal concentration of a weak acid and its conjugate • The pH of a buffered solution will equal the Pka of the weak acid. This is because [HA] = [A-] so that in the Ka expression they cancel out in this manner you can prepare buffer systems with specific pH’s.

31. Buffers in Biological systems • Blood/Oxygen equilibrium • CO2/HCO3- Buffer System CO2 + 2H2O <=> HCO3- +H3O+