1. Chapter 15 Acids and Bases, A Molecular Look

2. Brønsted–Lowry Acid/Base Definition
Johannes Brønsted 1879 – 1947
Thomas Lowry 1874 – 1936

3. Some Definitions HCl H+ (aq) + Cl- (aq) HCl + H2O H3O+ (aq) + Cl- (aq)
Brønsted–Lowry
Acid – chemical substance that acts a proton donor (species that produces H+)
Base – chemical subtance that acts as a proton acceptor (species that takes in H+)
HCl H+ (aq) Cl- (aq)
H2O
HCl H2O H3O+ (aq) Cl- (aq)
H+ = H3O+
hydronium ion

4. A Brønsted–Lowry acid… …must have a removable (acidic) proton.
A Brønsted–Lowry base…
…must have a pair of nonbonding electrons.
At their heart, acid-base reactions are proton transfer reactions.
The proton from the acid must be transferred to something.

5. If it can be either… …it is amphoteric. H2CO3 HCO3 CO32-
H2SO4 HSO4 SO42
H3O H2O OH-
Each of these substances can act as a base, accepting a proton, to become a neutral species or in the case of water, a hydronium ion.
They can also donate a proton to become a dianion in the case of bicarbonate and bisulfate or a hydroxide anion in the case of water.
Phosphoric acid H3PO4 is another example of an acid that can form an amphoteric species after it has donated a proton.

6. What Happens When an Acid Dissolves in H2O?
Water acts as a Brønsted–Lowry base and abstracts a proton (H+) from the acid.
As a result, the conjugate base of the acid and a hydronium ion are formed.
In the top equation HCl serves as a Brønsted–Lowry acid.
Water serves as a Brønsted–Lowry base.
Using its lone pairs it abstracts a proton (H+) from the acid.

7. H+ = H3O+ proton = hydronium ion
In aqueous solutions a “naked” proton does not exist.
The hydronium ion is the acid in aqueous solutions.
H+ = H3O+
proton = hydronium ion

8. Conjugates – an acid/base pair that differs only by a single H+
HNO2 / NO2− are conjugate acid/base pair
H3O+ / H2O are conjugate acid/base pair

9. What is the conjugate base of HCl ?
HCl H+ (aq) Cl− (aq)
H2O
HCl / Cl− are conjugates
(acid/base)
What is the conjugate acid of NH3 ?
NH3 (aq) + H2O
NH4+ (aq) + OH− (aq)
NH3 / NH4+ are conjugates
(base/acid)

10. What is the conjugate acid of HS−?
HS− / H2S are conjugates
(base/acid)
What is the conjugate base of HS−?
HS− / S2− are conjugates
(acid/base)
Is SO42− the conjugate base of H2SO4 ?
No way……
HSO4− is the conjugate base

11. Strong Electrolyte
– substance that completely dissociates into ions in aqueous solution
NaCl (aq) Na+ (aq) + Cl- (aq)
H2O

12. HCl (g) + H2O H3O+ (aq) + Cl- (aq)
Strong Acid – an acid that completely, 100% dissociates into ions in aqueous solution
HCl (g) + H2O H3O+ (aq) + Cl- (aq)
Note the single arrow.
6 common strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4
Completely ionized in water.
The conjugate bases of strong acids are exceptionally weak.
Note the single arrow when we write the reaction of a strong acid.

13. NaOH (s) Na+ (aq) + OH- (aq)
Strong Base – a base that completely, 100% dissociates into ions in aqueous solution
NaOH (s) Na+ (aq) OH- (aq)
H2O
6 common strong bases: LiOH, NaOH, KOH, RbOH, CsOH, Ba(OH)2
Completely ionized in water.
The strong bases listed here are the Group I and Ground II hydroxides.
Note the single arrow when we write the reaction of a strong base.

14. Weak Acid – an acid that only slightly dissociates into ions in aqueous solution
CH3COOH + H2O CH3COO- (aq) + H3O+ (aq)
Less than completely ionized in water.
The conjugate bases of weak acids are weak bases.
Note the double arrow when we write the reaction of a weak acid.
Note the double arrow.

15. – reacts with water to produce hydroxide ion
Weak Base – a base that only slightly dissociates into ions in aqueous solution
– reacts with water to produce hydroxide ion
Less than completely ionized in water.
The conjugate acids of weak bases are weak acids.
Note the double arrow when we write the reaction of a weak base.

16. Acid and Base Strength Strong acids – very weak conjugate bases.
Weak acids – weak conjugate bases.
Very weak acids – strong bases.

17. Acid and Base Strength
The more easily a substance gives up a proton, the less easily its conjugate base accepts a proton.
The more easily a base accepts a proton, the less easily its conjugate acid gives up a proton.
Strong acids – very weak conjugate bases.
Weak acids – weak conjugate bases.
Very weak acids – strong bases.
The stronger an acid, the weaker its conjugate base, and the stronger a base, the weaker its conjugate acid.

18. The conjugate base of a strong acid exhibits NO basic properties whatsoever
The conjugate acid of a strong base exhibits NO acidic properties whatsoever
The conjugate base of a weak acid exhibits weak basic properties
The conjugate acid of a weak base exhibits weak acidic properties

19. Salt Solutions: Ions as Weak Acids and Bases
Cations as Acids
Conjugate acids of weak molecular bases are weak acids and can affect pH of solution (make pH < 7.00).
e.g. NH4+ , HC17H19O3N+, etc.
The cation of a strong base is too weak to influence the pH of solution
e.g. NaOH ionizes 100%

20. Salt Solutions: Ions as Weak Acids and Bases
Anions as Bases
Conjugate bases of weak acids are weak bases and can influence the pH of solution. It will tend to make the solution basic (pH > 7.00)
The anion of a strong acid is too weak a base to influence the pH of solution
e.g. HCl ionizes 100%, conjugate base extremely weak
Little or no tendency to attract H+ to it

21. Predicting Acid-Base Properties of Salt
If neither cation nor anion can affect pH, solution will be neutral.
e.g. NaCl
If only cation is acidic, solution is acidic.
e.g. NH4I
If only anion is basic, solution is basic.
e.g. NaCHO2 (sodium formate)
If both cation is acidic and anion is basic, pH depends on relative strengths of acid and base.
e.g. NH4CHO2 (ammonium formate)

22. Learning Check
Predict whether a 0.10 M solution of NH4Br will be acidic, basic or neutral.
NH4+ is conjugate acid of weak base ammonia
It’s a weak acid
Will tend to make solution acidic
Br – is conjugate base of strong acid HBr
Extremely weak base
Will not affect the pH
Conclusion: Solution acidic
pH <7.00

23. Is a Soln of N2H5OCl Acidic, Basic or Neutral?
For N2H5+ :
For OCl– :
Comparing equilibrium constants we see that:
The base OCl– is stronger than the acid N2H5+
Solution will be basic

25. Trends in Binary Acid Strength
Binary Acids = HnX (two elements)
X = Cl, Br, P, As, S, Se, etc.
Acid strength increases from left to right within same period (across row)
Acid strength increases as electronegativity of X increases
e.g. HCl is stronger acid than H2S which is stronger acid than PH3
or PH3 < H2S < HCl
Binary acids only have two components.

26. Trends in Binary Acid Strength
Binary Acids = HnX
X = Cl, Br, P, As, S, Se, etc.
2. Acid strength increase from top to bottom within group
Acid strength increases as size of X and bond length increases
e.g. HCl is weaker acid than HBr which is weaker acid than HI
or HCl < HBr < HI

27. Synopsis: Factors Affecting Acid Strength
The more polar the H–X bond and/or the weaker the H–X bond, the more acidic the compound.
Acidity increases from left to right across a row and from top to bottom down a group.
Binary Acids

28. Learning Check Which is stronger? H2S H2S or H2O NH3 CH4 or NH3 HI
HF or HI

29. Oxoacids Oxoacids (HnX Om) Acids of H, O, and one other element
HClO, HIO4, H2SO3, H2SO4, etc.
As electronegativity of element X , electron density is drawn away from O, which draws electron density away from the O—H bond. This makes the bond more polar and makes the molecule a better proton donor.

30. Factors Affecting Oxoacid Strength
In oxoacids, in which an –OH is bonded to another atom, Y, the more electronegative Y is, the more acidic the acid.

31. Factors Affecting Oxoacid Strength
For a series of oxyacids, acidity increases as the number of oxygens increases.

32. Learning Check Which is the stronger acid in each pair? H2SO4 or H3PO4
HNO3 or H3PO3
H2SO4 or H2SO3
HNO3 or HNO2
H2SO4
HNO3

33. Learning Check
Which corresponds to the correct order of acidity from weakest to strongest acid ? A. HBrO3, HBrO, HBrO2 B. HBrO, HBrO2, HBrO3 C. HBrO, HBrO3, HBrO2 D. HBrO3, HBrO2, HBrO

34. Learning Check
Arrange the following in order of increasing acid strength:
HBr, AsH3, H2Se
AsH3 < H2Se < HBr
H2SeO4, H2SO4, H2TeO4
H2TeO4 < H2SeO4 < H2SO4
HBrO3, HBrO, HBrO4, HBrO2
HBrO < HBrO2 < HBrO3 < HBrO4
Note that in the second example, the electronegativity of the central element in the oxoacids is the criterion for determining acidity.

35. Strength of Organic Acids
Organic acid —COOH
Presence of electronegative atoms (halide, nitrogen or other oxygen) near —COOH group
Withdraws electron density from O—H bond
Makes organic acid, stronger acids
e.g. CH3CO2H < CH2ClCO2H < CHCl2CO2H < CCl3CO2H

36. Which of the following is the strongest organic acid?B C D E

37. Lewis Acids Lewis acids are defined as electron-pair acceptors.
A Coordinate Covalent Bond
Lewis acids are defined as electron-pair acceptors.
Atoms with an empty valence shells can be Lewis acids.

38. Lewis Bases Lewis bases are defined as electron-pair donors.
Anything that could be a Brønsted–Lowry base is a Lewis base.
Lewis bases can interact with things other than protons, however.

39. Learning Check Identify the Lewis acid and base in the following:
NH3 + H NH4+
Base Acid
F– BF BF4–
SeO3 + O2– SeO42–
Acid Base