1. Acids and Bases

2. Properties of Acids They are corrosive
pH range is 0 to 7, lower the pH, more acidic the acid
Taste sour
Turns litmus paper red
Often reacts with metals to produce H2 gas
Often produces H+ in water. Example; HCl  H+ + Cl-

3. Symbol used for acids

4. Properties of Bases
pH range is from 7 to 14, higher the pH more alkaline the base
They are caustic
Taste bitter
Turns litmus paper blue
Feels slippery
Group 1 and 2 metals make common bases. Example: Na + H2O  NaOH + 1/2H2

5. The pH of some common acid and bases

6. Common Acids Binary or Halide Acids
HCl, hydrochloric acid: is stomach and pool acid
HF, hydrofluoric acid: used to etch glass

7. Common acids continued
Oxyacids: acids with oxygen (more oxygen- stronger the acid)
HNO3, Nitric acid, used for making explosives, when it reacts with metals it produces NO2gas instead of H2
H2SO4, sulfuric acid: used in car batteries, used in many industrial process as a dehydrator.

8. Chemical burn from H2SO4
AHHHHHH SATAN!!!!!!! =]

9. Reaction of Cu with HNO3

10. Oxyacids continued
H2CO3, carbonic acid, found in carbonated sodas, and in rainwater. It is responsible for the sour taste, or sting in sodas. It is also found in blood to a very small extent.
H3PO4, phosphoric acid, used as lime scale remover, and as flavoring in the colorless sodas like Sprite.

11. Acids continued Organic acids Carbon based acids.
Structure ends in COOH
COOH is called a carboxyl group.
CH3COOH, ethanoic acid or vinegar. Used as a preservative and flavoring O
C-OH

12. Bases Group 1 bases Group 2 bases Strongest bases
NaOH, sodium hydroxide, and KOH, potassium hydroxide, both of these bases are used to make commercial cleansers and drain cleaners.
Group 2 bases
Weak bases,
Ca(OH)2 calcium hydroxide, make limewater which is used to test for CO2,,used to make antacids.

13. Bases Amines. This group has a -NH2 at the end of a molecule
These are formed from decomposing proteins
They have a “fishy” smell
NH3, ammonia, this is used as glass cleaner and to make fertilizers.
Examples CH3NH2, methyl amine, CH3CH2NH2 ethyl amine. Etc.

15. Bases Carbonates (they are antacids) Have CO3 -2 in them
This group produces CO2 when they react with acids.
NaHCO3 sodium bicarbonate, or baking soda
CaCO3, calcium carbonate, or baking powder.
MgCO3 and Al2(CO3)3 are used as antacids, for example: MgCO3 + 2HCl MgCl2(aq) + H2O(l) + CO2(g)

16. Arrhenius definition of acids and bases
Acids produce a hydrogen ion(H+) or hydronium ion (H3O+) when they dissolve in water. (An H+ ion is considered too reactive to exist so an H3O+ ion is used)
Acids dissolving in water producing an H+
HCl  H+ + Cl-
HNO3  H+ + NO3-
H2SO4  2H+ + SO4-2
CH3COOH  CH3COO- + H+

17. Carbonates are in antacids
Turner's stash

18. Arrhenius definition continued
Acids dissolving in water form a hydronium ion:
HCl + H2O  H3O+ + Cl-
HNO3 + H2O  H3O+ + NO3-
H2SO4 + 2H2O  2H3O+ + SO4-2
CH3COOH + H2O  CH3COO- + H3O+
Note: H2O is written in the reaction when H3O+ is used, but not when H+ is used.

20. Arrhenius definition of bases
Bases dissolve in water to produce an OH- called a hydroxide ion.
NaOH  Na+ + OH-
Ca(OH)2  Ca+2 + 2OH-
NH3 + H2O  NH4+ + OH-
Note: H2O is written in the reaction only with amines

21. Bronsted-Lowry def. of Acids/Bases
Acids are proton donors
Bases are proton acceptors
Ex: HNO2 + H2O H3O+ +NO2-
Ex: NH3+ HCO3-   NH4+ + CO3-2
Ex: HCO HSO4-   SO4-2 + H2CO3

22. Determining the conj. Acid of a base
--Add a H+
Base Conj. Acid
HCO3- H2CO3
H2O H3O+
OH- H2O

23. Determining the conjugate base of an acid.
--remove an H+
Acid Conj. Base
H2O OH-
HSO4- SO4-2
NH4+ NH3

24. Continued Bronsted-Lowry
Stronger the acid/ base, the weaker its conjugate base/ acid.
Amphoteric or Amphiprotic: A substance which can be either an acid or a base.
Example is H2O.
Water as an acid: H2O + Cl-  OH- + HCl
Water as a base: H2O + HCl  H3O+ + Cl-

25. Conjugate acid-base pairs

26. Definition of a Lewis acid and base
Acids are electron pair acceptors in a dative covalent bond.
Bases are electron pair donors in a dative covalent bond.
These acids are reserved for those molecules/ions which make dative covalent bonds but are NOT ALREADY A BRONSTED-LOWRY ACID.
Examples: NH3 is a base because the nitrogen has a pair of electrons which can be used in a dative bond. AlF3 can be a Lewis acid since Al has a vacant pair of orbitals that can accept e-.

27. Lewis acids will include substances that are not considered typical acids.
Examples:
Some metal ions. Ex Group 3, Trans. metals

28. Strong acids and bases completely dissociate in water

29. Strong and weak acids and bases
Strong acid/base when it dissolves in H2O, nearly 100% dissociates into ions.
Ex: HCl H+ + Cl- nearly 100% of the molecules break up into ions
Ex: NaOH Na+ + OH- nearly 100% of the molecules break up
(Remember: the H+ and OH- gives the properties associated with Acids and Bases.)
▪ For weak acid and bases, when they dissolve in water only a small % will dissociate to form ions.
CH3COOH(aq) CH3COO- (aq)+ H+(aq) Very little of the acid forms ions. Same condition for weak bases.
▪ Both the strength and the concentration of the acid or base determines its pH and its harmfulness. So when working with acids and bases one needs to be concerned with
1. Is the acid or base strong or weak?
2. Is the acid or base concentrated or dilute?

31. WWCND
How burgers are made

32. The pH equation The pH equation is: pH = -log [H+]
Where [H+] is the molar concentration.

33. The pH scale Development of the pH scale
Based upon the dissociation of water: H2O  H+ + OH-
Based upon the equilibrium of water: Kw = [H+] [OH-]
at equilibrium the concentrations of H+ and OH- are each 1x10-7M (this was experimentally determined)
Substitute these concentrations into the Kw expression: Kw = [1x10-7][1x10-7]; this = 1x10-14

34. Development of the pH scale cont.
According to the rules of equilibrium changing conc does not change the Kw constant. So if acid is added to water, the H+ conc goes up, and the OH- goes down, but K equals 1x The same thing happens if a base is added to water. Knowing that Kw is always 1x10-14, then
The largest conc for either an H+ or OH – is 1M
The smallest conc H+ or OH – is 1x10-14 that we work with
the –log of 1M = 0 the low end of the pH scale
the –log of 1x10-14 = 14 the high end of the pH scale
Also, pure H2O has a conc of H+ = 1x10-7, the –log = 7, pure H2O has a pH of 7 or it is neutral

35. Acid and Base Neutralizations
General equation: acid + base  salt+ H2O + energy
Ex: HCl+ NaOH  H2O + NaCl
CH3COOH + NaOH HOH + NaCH3COO
H2 SO4 + 2NaOH  2 HOH + Na2SO4
HNO3 + NaOH  NaNO3 + H2O
2HCl+ Ca(OH)2  CaCl2 + 2H2O

36. Titration Curves for strong acid and bases

37. Titration curve for strong acid and bases

38. Explanation of a titration curve
A titration involves an acid base reaction:
Example NaOH + HCl  NaCl + NaOH
All titration curves have the same shape but they do not all start at the same pH
Why this shape?
Start with 1M HCl titrated with 1M NaOH
Initial pH=O, titrate 10% of 1MHCl, 90% of HCl remains. [H+]= .9M or pH= .05
Titrate 50% of 1M HCl, 50% remains [H+]= .5M or pH= .3 (not a great change in pH yet!)
Titrate 90% of 1MHCl , 10% remains, [H+]=.1M or pH=1
Titrate 100% of HCl, acid and bases are now equal, pH=7 the cure shoots up until excess NaOH is added. Now curve gradually increases again with the addition of more NaOH

39. Tiration curves with weak acids and bases.

42. IB optional material: Equilibrium and weak acids and bases
Since weak acids and bases disassociate partially:
They are reversible and have a measurable equilibrium.
Their pH cannot be based upon the initial molarity of the acid or base. An equilibrium equation must be used to measure it.

43. The Ka or Kb (equilibrium constant) of weak acids or bases