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Acids and bases, salts and solutions

Acids and bases, salts and solutions

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Acids and bases, salts and solutions

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  1. Acids and bases, salts and solutions Chapter 10-1 – 10-9, 11-1 – 11-4

  2. Key concepts • Compare and contrast the Arrhenius and Brønsted-Lowry theories of acids and bases • Describe hydrated protons • Properties of acid and base solutions • Arrange acids according to acid strength • Balance acid-base equations • Amphoterism • Lewis acid-base theory • Molarity calculations in titrations • Equivalents

  3. Various properties of acids and bases

  4. Arrhenius theory of acids/bases • Developed in 1884. • An acid is a substance containing hydrogen that produces _______ in aqueous solution. • A base is a substance containing the OH group that produces _______ in aqueous solution.

  5. Protons are not alone…. • Protons combine with water molecules to form __________ ___________. • We commonly represent this as the hydronium ion, H3O+(aq), but writing H+(aq) means the same thing.

  6. Brønsted-Lowry theory (1923) • An acid is a _______ _________. • A base is a ________ __________. • Bases are no longer restricted to compounds that release OH- in solution. For instance, NH3 is a base. What does it look like after reacting with a proton?

  7. Ionization of weak acids/bases • While strong acids dissociate completely, not all reactions are complete and irreversible (in fact, most are not). • Rxns with weak acids/bases are reversible. Example: HF + H2O • What is the acid? What is the base? (depends on which side of rxn you look at)

  8. Conjugate acid-base pairs • Conjugate acid-base pairs differ in structure by ___ _________. • Some examples of conjugate pairs:

  9. Acid/base strength • The strength of an acid is _______ proportional to the strength of its base. • Strong acids have ________________. • Weaker acids have ________________. As the acid gets weaker and weaker, what happens to the conjugate base? What does this tell you about the amount of ionization taking place?

  10. Amphoterism • Some substances can both give and accept protons. This process is called amphoterism. • Water is the prime example of amphiprotic behavior. 2H2O → H3O+ + OH-

  11. Acid strength • The hydrohalic acids: HF, HCl, HBr, HI • What are the sizes of the halogens? How will this affect the H-X bond? • HF bond is very strong vs. the other halogens. • F- causes ordering of the H2O molecules (how does that happen?)

  12. Leveling solvents • In aqueous solution, no acid is stronger than H3O+(aq). All other acids completely dissolve in water to form H3O+. • Because of this, all strong acids are of equal strength in water. • A similar effect is observed for strong bases, which completely dissolve to form OH-.

  13. What is a ternary acid? Ternary acids are hydroxyl compounds of a ______________. Ionize to produce H+. Compare to other hydroxyl compounds… Metal hydroxides—ionize to produce ________________ and are ________ in aqueous solution. Ternary acids and bases

  14. Ternary acid strength • H2SO4 vs H2SO3. What’s the difference in acid strength? • Compare oxidation number of sulfur in each. • Acid strength increases with _______ oxidation number of the central atom. • Order the following acids from weakest to strongest: • HBrO3, HBrO, HBrO4, HBrO2

  15. important! • When comparing ternary acid strength, make sure the compounds have similar structure. • Where are the hydrogens located? • (H3PO3 vs H3PO4)

  16. Neutralization of Brønsted-Lowry acids/bases HA + MB → HB + MA • In many cases, HB ends up being ______. • Classic example: strong acid + strong base. • What happens in the reaction of hydrochloric acid and sodium hydroxide? • What is the net ionic equation?

  17. Weak acid + strong base • General reaction: HxA (aq) + x OH- (aq) → A- (aq) + x H2O (l) • When does x vary? • Examples:

  18. Acid salts • Acid salts are salts of ______ acids that still contain __________ ___________.

  19. Lewis theory • The most general of all acid-base theories • Discards the proton acceptor/donator all together. • A Lewis acid _______ a share in an electron pair. • A Lewis base _________ a share in an electron pair. • Lewis acids and bases are neutralized when a ________ _______ forms.

  20. When is Lewis theory used? Arrhenius or B-L theory a better description for most aqueous sol’ns LEWIS THEORY Bronsted- Lowry Arrhenius Lewis theory a good descriptor for nonaqueous solvents or transition metals

  21. Acid-Base calculations • Molarity calculations play an important part in acid-base reaction stoichiometry • Much of what we will learned in Chapter 3 will be used here.

  22. Molarity • M = mol/L or • M = mmol/mL • we can use moles and liters, or millimoles and milliliters, and the molarity is still the same.

  23. Similarities between acid-base and other reaction calculations • We still compare moles to moles, not volumes to volumes or molarities to molarities. • Additionally, knowing the limiting reactant is very important (i.e, what will run out first—acid or base?)

  24. Some examples • what volume of 0.800 M NH3 is required to neutralize 22.0 mL of 12.0 M HCl? • 25.0 mL of 0.0500 M Ca(OH)2 added to 10.0 mL of HNO3. • Is the solution now acidic or basic? • how many moles excess acid or base are in the solution? • how much additional Ca(OH)2 or HNO3 sol’n required to neutralize solution?

  25. TITRATIONS • Combining a known concentration with an unknown concentration solution. • Titrant: The solution of one reactant (usually of unknown concentration) that is carefully added to the solution of the other reactant until the resulting solution is just neutralized (no excess acid or base). • How do we know when to stop?

  26. Titrations (con’t) • indicators: • How to measure the volume of titrant? • Buret: • equivalence point: The point where _____________ _______________ amounts of acid and base have reacted. • end point: The point where the indicator ____________ ___________. • For accurate work, one wants the end point and equivalence point to coincide with each other.

  27. Primary and secondary standards • Reading on standardization: your text goes over the requirements of a primary standard. You should be familiar with these requirements. • Primary standards are used to determine the concentration of solutions, which become secondary standards. • Example: KHP and NaOH.

  28. EQUIVALENT WEIGHTS AND NORMALITY • One mole of acid is _____________________________ • But, one equivalent of acid contains ______________________________. • The equivalent weight, then, corresponds to molar mass/(# of equ./mol)

  29. normality • number of equivalents per liter, or N = eq/L = meq/mL • N = M  eq/mol • Let’s do a couple of examples…

  30. EQUIVALENTS in acid/base reactions • 1 eq acid always reacts with 1 eq base. • Va Na = Vb Nb