1. Acids & Bases

2. A. Properties ACIDS BASES electrolytes  electrolytes sour taste
bitter taste
turn litmus red
turn litmus blue
react with metals to form H2 gas
slippery feel
vinegar, milk, soda, apples, citrus fruits
ammonia, lye, antacid, baking soda
ChemASAP

4. Arrhenius acids & bases
Arrhenius - In aqueous solution…
Acids form hydronium ions (H3O+)
Hydrogen ions (protons) plus water
HCl + H2O  H3O+ + Cl– H Cl O – +
acid

5. Strong vs. Weak

6. Increasing polarity = increasing strength

7. Organic acids Acids that contain the carboxyl group -COOH Weak acids
Acetic acid – CH3COOH
See your notes for other examples

8. pouvoir hydrogène (Fr.)
pH Scale 14 7
INCREASING
ACIDITY
INCREASING
BASICITY
NEUTRAL
pH = -log[H3O+]
pouvoir hydrogène (Fr.)
“hydrogen power”

9. pH of Common Substances
B. pH Scale
pH of Common Substances

10. NH3 + H2O  NH4+ + OH- B. Definitions – +
Arrhenius - In aqueous solution…
Bases form hydroxide ions (OH-)
NH3 + H2O  NH4+ + OH- H N O – +
base

11. Alkalinity
Alkaline: solution in which a base has completely dissociated in water to yield aqueous OH- ions
Alkalinity depends on concentration of OH- in solution

12. Strong vs. Weak Bases Strong base Weak base
Strong electrolytes – complete dissociation
Weak base
Removes H+ from water leaving OH-behind

13. -Substances that already contain ions (i.e. ionic compounds)
merely release their ions as the crystal dissociates.
-Substances that are neutral molecules
(i.e. molecular/covalent compounds) must react with water to produce ions (called ionization)
-CONCENTRATED/DILUTE – tell how much acid or base is dissolved in solution (molarity, molality, %conc.)
-STRONG/WEAK – refer to the extent of ionization or dissociation of an acid or base

14. Arrhenius Theory Had some inadequacies:
Requires that substances be aqueous
NH3, Na2CO3, and C6H5NH2 (aniline) are all basic but they are NOT hydroxides
So his theory had to be expanded (added on to)

15. Bronsted Lowry Includes all Arrhenius acids and bases
Plus substances that are not in aqueous solutions and do not have hydroxide

16. HCl + H2O  Cl– + H3O+ Bronsted Lowry acid conjugate base base
Brønsted-Lowry
Acids are proton (H+) donors.
Bases are proton (H+) acceptors.
HCl + H2O  Cl– + H3O+
acid
conjugate base
base
conjugate acid

17. Proton Transfer Reaction
H2O + HNO3  H3O+ + NO3– B A CA CB

18. Acids – turn to page 465-466 for examples
Monoprotic
Polyprotic
Diprotic
Triprotic

19. Just a note The terms ACID & BASE are used for reactants
The terms CONJUGATE ACID & CONJUGATE BASE are used for products
Reactants must be stronger than the products in order for the reaction to proceed/occur with any significant extent

20. Conjugate acid/base
Conjugate base – remains after B-L acid has given up proton
Conjugate acid – remains after B-L base has accepted proton
Conjugate pair – differ from each other by the loss or gain of 1 H+ ion

21. Strength Strong acids have weak conjugate bases and vice versa
The larger the Ka/Kb value – the stronger the acid or base (see chart)
Reactions always proceed from the strong to the weak.
Strong acid/base ionize readily, the conjugates must be too weak to compete

22. Strength

23. C. Strength Strong Acid/Base 100% ionized in water strong electrolyte- +
HCl
HNO3
H2SO4
HBr HI
HClO4
NaOH
KOH
Ca(OH)2
Ba(OH)2

24. C. Strength Weak Acid/Base does not ionize completely weak electrolyte- + HF
CH3COOH
H3PO4
H2CO3
HCN
NH3

25. Amphoteric - can be an acid or a base.
B. Definitions
NH3 + H2O  NH4+ + OH- B A CA CB
Amphoteric - can be an acid or a base.

26. B. Definitions HF H3PO4 H3O+ F - H2PO4- H2O
Give the conjugate base for each of the following: HF
H3PO4
H3O+
F -
H2PO4-
H2O

27. B. Definitions Br - HSO4- CO32- HBr H2SO4 HCO3-
Give the conjugate acid for each of the following:
Br -
HSO4-
CO32-
HBr
H2SO4
HCO3-

28. Lewis Acid Base Theory
Arrhenius theory & Bronsted-Lowry theory still did not cover all the substances that were acidic & basic SO there must be more
Lewis based his theory on Bonding Structure

29. Lewis Acids & Bases Lewis base Lewis acid
Acids are electron pair acceptors.
Bases are electron pair donors.
Draw the example
Lewis base
Lewis acid

30. Lewis Acids & Bases Lewis acids = e- pair acceptor
Lewis base = e- pair donor
Lewis reaction = the formation of one or more covalent bonds between an electron pair donor and an electron pair acceptor

31. Neutralization
Reaction of hydronium and hydroxide ions to form water molecules and another compound
Salt – an ionic compound composed of cation from the base and anion from an acid
ACID + BASE>SALT + WATER

32. Acid Rain
Acid Rain – any precipitation that has an unusually high concentration of sulfuric or nitric acids, resulting from chemical pollution of the air
Unpolluted rain has a pH of 5.6
Acid rain = pH of 4.6 (10 times more acidic)

33. Acid Rain Can dissolve marble buildings and statues
Kills plankton and other aquatic animals
Dissolves metals and nutrients in the soil
Usually caused by sulfur trioxide and nitrogen oxide when they combine with water

34. Mole Ratios
Use the coefficients in the balanced equations as ratios to do your calculations