1. Topic 8 Acids and Bases
What are the current working definitions of an acid and a base? How have the definitions changed over time? Are there varying but equally accepted definitions?
How can we identify an acid or a base? What are the patterns of acid/base behavior?

2. Topic 8.1 A Brønsted-Lowry definition: Amphiprotic species can act as
acid is a proton (H+) donor
base is a proton (H+) acceptor
Amphiprotic species can act as
A pair of species differing by a single proton is called a conjugate acid-base pair.

3. What is an Acid?
In an acidic solution the hydronium ions outnumber the hydroxide ions (more H3O+).
Acids taste tart or sour- lemons, citrus fruits, vinegar, etc.
pH < 7.0

4. Acids & H+ identity crisis
Arrhenius definition – H+ in solution
Hydronium = hydrogen ion = proton
“explained by the extreme reluctance of a hydrogen nucleus to lead an isolated existence”
Lowry 1923

6. -ΔH Properties of Acids Aqueous solutions of acids have a sour taste
Shouldn’t have to say this but: don’t taste test in lab...
Acids change the color of acid-base indicators
Some acids react with active metals and release hydrogen gas, H2.
Acids react with bases to produce salts, water, and heat
Neutralization reactions
-ΔH

7. HF HI Acid Nomenclature
Binary acid- acid that contains only two different elements (hydrogen + another element that is not oxygen)
Name starts with the prefix “hydro-”
The root name of the second element follows the prefix
Root is followed by the suffix “-ic”
Followed by the word “acid”
Ex. HCl
“hydro-” + “chlor” + “-ic” + “acid” = hydrochloric acid HF
Hydrofluoric acid HI
Hydroiodic acid

8. Acid Nomenclature
Oxyacid- acid that is a compound of hydrogen, oxygen, + another element.
Base name of ion + “-ic” + acid
Ex. H2SO4
SO4-2 = sulfate
H2SO4 = sulfur- + “-ic” + acid = sulfuric acid
What if you have multiple base names? (nitrite, nitrate, etc.)

9. More on naming acids
Number of Oxygen Atoms (compared to starting compound)
Ion Name
Acid Name
+1 Oxygen
ClO4- (perchlorate)
HClO4 (perchloric acid)
Starting Compound
ClO3- (chlorate)
HClO3 (chloric acid)
-1 Oxygen
ClO2- (chlorite)
HClO2 (chlorous acid)
-2 Oxygen
ClO- (hypochlorite)
HClO (hypochlorous acid)
No Oxygen
Cl- (chloride)
HCl (hydrochloric acid)

10. Some Common Acids used in Industry
Sulfuric Acid
Petroleum refining, metallurgy, fertilizers, dehydrating agent + many industrial processes
Nitric Acid
Explosives, rubber, plastics, dyes (yellow), pharmaceuticals, + industry
Phosphoric Acid
Fertilizers, animal feed, flavoring agent, cleaning agent, detergents, ceramics
Hydrochloric Acid [muriatic acid]
Your stomach, “pickling” iron and steel, cleaning, food processing, activating oil wells, etc.
Acetic Acid
Vinegar, synthesizing chemicals in plastic manufacturing, food supplements, fungicide.

11. What is a Base?
In a basic solution the hydroxide ions outnumber the hydronium ions (more OH-)
Bitter taste, feel slippery- soapy water, milk of magnesia
pH > 7

12. Indicator colors vary – only in litmus is blue/pink
Base Properties
Aqueous solutions of bases taste bitter
Don’t try tasting bases, either
Bases change the color of acid-base indicators
Dilute aqueous solutions of bases feel slippery
Bases react with acids to produce salts, water, and heat
Neutralization reactions
Indicator colors vary – only in litmus is blue/pink

13. Some Definitions Arrhenius Brønsted–Lowry
An acid is a substance that, when dissolved in water, increases the concentration of hydrogen ions.
A base is a substance that, when dissolved in water, increases the concentration of hydroxide ions.
Brønsted–Lowry
An acid is a proton donor.
A base is a proton acceptor.

14. Brønsted–Lowry Acid and Base
A Brønsted–Lowry acid must have at least one removable (acidic) proton (H+) to donate. A Brønsted–Lowry base must have at least one nonbonding pair of electrons to accept a proton (H+).

15. What Is Different about Water?
Water = amphiprotic.
What’s going on here?

16. Conjugate Acids and Bases
The term conjugate means “joined together as a pair.”
Reactions between acids and bases always yield their conjugate bases and acids.

17. acid
base
base
acid

19. What does ‘alkali’ mean?
Alkalis = substances which form OH- ions in solution
Bases = substances which accept H+ ions in solution

20. Acids react with metals, bases, and carbonates to form salts
Salt- ionic compound formed when the hydrogen of an acid is replaced by a metal or other positive ion
Parent Acid/Parent Base- starting acid/base in a reaction

21. 1. Acid + metal -> salt + hydrogen
(write an example):

22. 2. Acid + base -> salt + water
(write an example):
What type of reaction is this?
What does enthalpy of neutralization mean?

23. 3. Acid + carbonate -> salt + water + carbon dioxide
(write an example):
What does effervescence mean?
Stalagtites/Stalagmites

25. Indicators
Change color based on [H+] or [OH-] concentration, section 22 of data booklet
Universal Indicator- can be used across a range of different acids and bases

27. Titrations
Neutralization reactions in titrations are used to determine the exact concentration of an acid or a base when the other concentration is known. [solution of known concentration = standard solution]
Gradual addition of the standard solution from a buret to a volume of the unknown using an indicator

29. 8.3 The pH scale pH = -log[H+(aq)] and [H+] = 10-pH
A change of one pH unit represents a 10-fold change in the hydrogen ion concentration [H+]
pH values distinguish between acidic, neutral, and alkaline solutions

30. Aqueous Solutions Can Be Acidic, Basic, or Neutral
If a solution is neutral, [H+] = [OH–].
If a solution is acidic, [H+] > [OH–].
If a solution is basic, [H+] < [OH–].

31. pH pH is a method of reporting hydrogen ion concentration.
pH = –log[H+]
Neutral pH is 7.00.
Acidic pH is below 7.00.
Basic pH is above 7.00.

32. How Do We Measure pH?
Indicators, including litmus paper, are used for less accurate measurements; an indicator is one color in its acid form and another color in its basic form.
pH meters are used for accurate measurement of pH; electrodes indicate small changes in voltage to detect pH.

33. More about pH pH numbers are usually positive and have no units
The pH number is inversely related to the [H+]
A change of one pH unit represents a ten-fold change in [H+]

34. pH = -log[H+(aq)] and [H+] = 10-pH

36. Autoionization of Water
Water is amphoteric.
In pure water, a few molecules act as bases and a few act as acids.
This is referred to as autoionization.
Remember H+ identity crisis…
What do we do with pure liquids in Kc?

37. Ion Product Constant The equilibrium expression for this process is
Kc = [H3O+][OH]
This special equilibrium constant is referred to as the ion product constant for water, Kw.
At 25 °C, Kw = 1.0  1014

38. Other “p” Scales
The “p” in pH tells us to take the –log of a quantity (in this case, hydrogen ions).
Some other “p” systems are
pOH: –log[OH]
pKw: –log Kw

39. –log[H3O+] + –log[OH] = –log Kw = 14.00
Relating pH and pOH
Because
[H3O+][OH] = Kw = 1.0  1014
we can take the –log of the equation
–log[H3O+] + –log[OH] = –log Kw = 14.00
which results in
pH + pOH = pKw = 14.00

41. 8.4 Strong and weak acids and bases
Strong and weak acids and bases differ in the extent of ionization/dissociation
Strong acids and bases have higher conductivities
A strong acid is a good proton donor and has a weak conjugate base
A strong base is a good proton acceptor and has a weak conjugate acid

42. Strong and Weak Acids Strong acid- ionizes completely in solution
Strong acid = strong electrolyte
Weak acid- releases few hydrogen ions in aqueous solution
Reverse reaction also occurring

43. Strong Acids
Seven strong acids are HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4.
These are, by definition, strong electrolytes and exist totally as ions in aqueous solution; e.g.,
HA + H2O → H3O+ + A–
So, for the monoprotic strong acids,
[H3O+] = [acid]

44. Strong Bases
Strong bases are the soluble hydroxides, which are the alkali metal and heavier alkaline earth metal hydroxides (Ca2+, Sr2+, and Ba2+).
Again, these substances dissociate completely in aqueous solution; e.g.,
MOH(aq) → M+(aq) + OH–(aq) or
M(OH)2(aq) → M2+(aq) OH–(aq)

45. We will revisit Ka Weak Acids
For a weak acid, the equation for its dissociation is
HA(aq) + H2O(l) ⇌ H3O+(aq) + A–(aq)
Since it is an equilibrium, there is an equilibrium constant related to it, called the acid-dissociation constant, Ka:Ka = [H3O+][A–] / [HA]
The greater the value of Ka, the stronger is the acid.
We will revisit Ka

46. We will revisit Kb Weak Bases Ammonia, NH3, is a weak base.
Like weak acids, weak bases have an equilibrium constant called the base dissociation constant.
Equilibrium calculations work the same as for acids, using the base dissociation constant instead.
We will revisit Kb

47. Comparing Strong and Weak Acids
What is present in solution for a strong acid versus a weak acid?
Strong acids completely dissociate to ions.
Weak acids only partially dissociate to ions.

48. Strong vs. Weak Acids— Another Comparison
Strong Acid: [H+]eq = [HA]i
Weak Acid: [H+]eq < [HA]i
This creates a difference in conductivity and in rates of chemical reactions.

49. IB table p.362

50. Comparing Acids and Bases
Electrical conductivity- more mobile ions = better conductors
2. Rate of Reaction- higher H+ concentration = faster reaction rate
3. pH- higher H+ concentration = lower pH