1. Acids and Bases

2. 1 14 H2O H+ OH- Ex. HCl HSO3 Ex. NaOH KOH Base Acid Strong Weak Weak
Acetic Acid
NH3
NaOH 7
H2O H+
OH-
Base
Acid
Ex. HCl
HSO3
Ex. NaOH
KOH

3. Arrhenius Acids
Substances that produce H+ (hydrogen ion, proton) in soln.
Properties of acids are related to the H+ ions they produce
Increase H+ concentration when added to H2O
Decrease OH- concentration when added to H2O
Decreases pH, x<7
Turn blue litmus paper red
Sour
No effect on phenolphthalein, stays colorless
React with certain metals to produce salt and H2
The stronger the acid, the greater the hydrogen ion concentration in comparison to the OH- ion concentration

4. Arrhenius Acids Some definitions:
Not all compounds that contain hydrogens are acids. Ex. CH3COOH : acetic acid, i.e. Vinegar
Only hydrogens in very polar bonds are ionizable.
Hydrogen ions are stabilized by solvation.
Some definitions:
Monoprotic Acids – acids that contain one ionizable hydrogen. Ex. HNO3
Diprotic Acids – acids that contain two ionizable hydrogens. Ex. H2SO4
Triprotic Acids – acids that contain three ionizable hydrogens. Ex. H3PO4

5. Arrhenius Base Substances that produce OH- (hydroxide ion) in soln.
Properties of bases are related to the OH- ions they produce
Increase OH- concentration when added to H2O
Decrease H+ concentration when added to H2O
Increases pH, x>7
Turn red litmus paper blue
Bitter and slippery
Turns colorless phenolphthalein, pink.
Most bases are ionic compounds.
The stronger the base, the greater the OH- ion concentration in comparison to H+ ion concentration

6. Bronsted-Lowry Acid Bronsted-Lowry Base
Substances that DONATE H+ (hydrogen ion, proton) in rx.
H2O is an acid b/c it gives up an H+ and becomes OH- (a base)
Substances that ACCEPT H+ in rx.
NH3 is a base b/c it accepts the H+ and becomes NH4+ (an acid)

7. Conjugate Acids & Bases
Bronsted- Lowry Theory
Conjugate Acids & Bases
Conjugate - coupled, connected, or related, in particular.
In the case of the reverse reaction
the conjugate acid is the acid; H+ donor on the “product” side
the conjugate base is base; H+ acceptor on the “product” side.
In essence, the reversible reaction of ammonia and water has two acids and two bases.

8. Conjugate Acids & Bases
Bronsted- Lowry Theory
Conjugate Acids & Bases
Which substance is the B-L acid (H+ donor)?
Which substance is the B-L base (H+ acceptor)?
What substance is the conjugate base?
What substance is the conjugate acid?

9. Conjugate Acids & Bases
Bronsted- Lowry Theory
Conjugate Acids & Bases
Conjugate acids are always paired with a base, and conjugate bases are always paired with an acid.
A conjugate acid-base pair consists of two ions or molecules related by the loss or gain of one hydrogen ion.

10. Conjugate Acids & Bases
Bronsted- Lowry Theory
Conjugate Acids & Bases
Which substance is the B-L acid (H+ donor)?
Which substance is the B-L base (H+ acceptor)?
What substance is the conjugate base?
What substance is the conjugate acid?
The chloride ion is the conjugate base of the acid HCl.
The hydronium ion is the conjugate acid of the water base.

11. Conjugate Acids & Bases
Bronsted- Lowry Theory
Conjugate Acids & Bases
Which substance is the B-L acid (H+ donor)?
Which substance is the B-L base (H+ acceptor)?
What substance is the conjugate base?
What substance is the conjugate acid?
The chloride ion is the conjugate base of the acid HCl.
The hydronium ion is the conjugate acid of the water base.

12. Amphoteric Substances
Bronsted- Lowry Theory
Amphoteric Substances
A substance that can act as either an acid or a base is said to be amphoteric.
Water is amphoteric.
In the reaction with hydrochloric acid, water accepts a proton and is therefore a base.
In the reaction with ammonia, water donates a proton and is therefore an acid.

13. Lewis Acids and Bases
A Lewis acid is a substance that can accept a pair of electrons to form a covalent bond.
A Lewis base is a substance that can donate a pair of electrons to form a covalent bond.
The Lewis definitions include all the Brønsted-Lowry acids and bases.
This definition is more general than those offered by Arrhenius or by Brønsted and Lowry.

14. Acid Base Theories Summary
Acid-Base Definitions
Type
Acid
Base
Arrhenius
H+ producer
OH– producer
Brønsted-Lowry
H+ donor
H+ acceptor
Lewis
electron-pair acceptor
electron-pair donor
The Lewis definition is the broadest.
It extends to compounds that the Brønsted-Lowry theory does not classify as acids and bases.

15. Practice Problem
Identify the Lewis acid and the Lewis base in this reaction between ammonia and boron trifluoride.
NH3 + BF3 → NH3BF3
When a Lewis acid reacts with a Lewis base, the base donates a pair of electrons and the acid accepts the donated pair.

16. Additional Practice Problems
Identify the Lewis acid and Lewis base in each of the reactions.
H+ + H2O  H3O+
AlCl3 + Cl-  AlCl4-
Would you predict PCl3 to be a Lewis acid or a Lewis base in a typical reaction? Explain your prediction.
KOH + HBr  KBr + H2O
HCl + H2O  H3O+ + Cl-
H+ is the Lewis acid, H2O is the Lewis base
AlCl3 is the Lewis acid; Cl- is the Lewis base
PCl3 would be a Lewis base, it has a nonbonding pair of electrons that it can donate
KOH is the base, HBr is the acid
HCl is the acid, H2O is the base

17. Question #7 on Page 593
7. Write equations for the ionization of HNO3,in water and the reaction of CO32- with water. For each equation, identify the hydrogen-ion donor and the hydrogen-ion acceptor. Then label the conjugate acid-base pairs in each equation.
HNO3 + H2O  H3O+ + NO3-
Hydrogen-ion donor (Acid): HNO3
Conjugate Base: NO3-
Hydrogen- ion acceptor (Base): H2O
Conjugate Acid: H3O+

18. Acid Base Indicators
Substances that change colors in the presence of an acid or base.

19. Hydrogen Ions and Acidity
Section 2
Hydrogen Ions and Acidity

20. Ion Product Constant for Water
The ionization of water is a reversible reaction, so Le Châtelier’s principle applies.
Adding either hydrogen ions or hydroxide ions to an aqueous solution is a stress on the system.
In response, the equilibrium will shift toward the formation of water.
The concentration of the other ion will decrease.

21. Ion Product Constant for Water
The ionization of water is a reversible reaction, so Le Châtelier’s principle applies.
In any aqueous solution, when [H+] increases, [OH−] decreases. Likewise, when [H+] decreases, [OH−] increases.

22. Ion Product Constant for Water
Acidic
HCl(aq) → H+(aq) + Cl−(aq)
A solution in which [H+] is greater than [OH−] is an acidic solution.
The [H+] is greater than 1 × 10−7M.
Basic
NaOH(aq) → Na+(aq) + OH−(aq)
A basic solution is one in which [H+] is less than [OH−].
The [H+] is less than 1 × 10−7M.
AKA alkaline solutions
Acids Notes: When some substances dissolve in water, they release hydrogen ions.
Bases Notes: When sodium hydroxide dissolves in water, it forms hydroxide ions in solution.

23. Ion Product Constant for Water
For aqueous solutions, the product of the hydrogen-ion concentration and the hydroxide-ion concentration equals × 10−14.
The product of the concentrations of the hydrogen ions and the hydroxide ions in water is called the ion-product constant for water (Kw).
Kw = [H+] × [OH−] = 1.0 × 10−14

24. Using the Ion Product Constant for Water
Sample Problem 19.2
Using the Ion Product Constant for Water
If the [H+] in a solution is 1.0 × 10−5M, is the solution acidic, basic, or neutral? What is the [OH−] of this solution?
[H+] is 1.0 × 10−5M, which is greater than 1.0 × 10−7M.
Rearrange Kw = [H+] × [OH−] 
[OH−] = Kw
[H+]
[OH −] = 1.0 × 10−9M

25. Relating pH to H+ Ion Concentration
Mathematically, pH is defined as the –log of H+ ion concentration of a soln.
pH = -log[H+] or [H+] = 10-pH
pH is a measure of how much H+ are in a solution.
If H+ concen. Is not given as
1 x 10-x :
Ex. [H+] = 5.4 x 10-8M
pH = -log(5.4 x 10-8) = 7.3
Expressing [H+] in scientific notation can make it easier to calculate pH. If H+ is written in scientific notation and has a coefficient of 1, then the pH of the solution equals the exponent, with the sign changed from negative to positive.

26. Calculating pH
To calculate the pH of an aqueous solution you need to know the concentration of the hydronium ion in moles per liter (molarity). The pH is then calculated using the expression:
pH = - log [H3O+]
Example:
Find the pH of a M HCl solution. The HCl is a strong acid and is 100% ionized in water. The hydronium ion concentration is M. Thus:
pH = - log (0.0025) = - ( ) = 2.60

27. Calculating pH more problems
Example:
What is the pH of a solution with a hydrogen-ion concentration of × 10−10M?
pH = - log [H3O+]
pH = - log [4.2 x 10-10M]
pH = - ( )
pH = (Base)

28. BELL RINGER Activity 1. 2.

29. Calculating the Hydronium Ion Concentration from pH
pH = - log [H3O+]
Calculating the Hydronium Ion Concentration from pH
The hydronium ion concentration can be found from the pH by the reverse of the mathematical operation employed to find the pH.
[H3O+] = 10-pH
Example: What is the hydronium ion concentration in a solution that has a pH of 8.34?
8.34 = - log [H3O+]
= log [H3O+]
[H3O+]= = 4.57 x 10-9 M

30. Calculating the Hydronium Ion Concentration from pH
The pH of an unknown solution is What is the hydrogen-ion concentration?
6.35 = - log [H3O+]
= log [H3O+]
[H3O+]= = 4.5 x 10-7 M

31. Calculate the [H+] for each solution. pH = 5.00 pH = 12.38
Practice Problems
Calculate the [H+] for each solution.
pH = 5.00
pH = 12.38
When the pH is a whole number, you can assume the coefficient to be 1
When the pH is a decimal, use the antilog.
What are the hydrogen ion concentrations for solutions with the following pH values?
pH = 4.00
pH = 11.55

32. Relationship between pH and pOH.
pH scale ranges from
For aqueous solutions, the product of the hydrogen-ion concentration and the hydroxide-ion concentration equals 1.0 × 10−14.
[H+] + [OH−] = 1.0 × 10−14

33. Calculating pOH
To calculate the pOH of a solution you need to know the concentration of the hydroxide ion in moles per liter (molarity). The pOH is then calculated using the expression:
pOH = - log [OH-]
Example:
What is the pOH of a solution that has a hydroxide ion concentration of 4.82 x 10-5 M?
pOH = - log [4.82 x 10-5]
= - ( ) = 4.32

34. Calculating pH from [OH−]
If you know the [OH−] of a solution, you can find its pH.
You can use the ion-product constant to determine [H+] for a known [OH−].
Then you use [H+] to calculate the pH.

35. Practice Problem
What is the pH of a solution if the [OH-] = 4.0 x 10-11M?
Start with the ion-product constant to find [H+].
Rearrange the equation to solve for [H+].
[H+] = Kw
[OH−]
Kw = [OH−] × [H+]
[H+] =
1.0 × 10−14
4.0 × 10−11
= 0.25 × 10−3M 3.
[H+] = 2.5 × 10−4M 4.
pH = −log[H+]
= −log(2.5 × 10−4)
pH = 3.60
= −(− )

36. Practice Problems Calculate the pH of each solution
[OH-] = 4.3 x 10-5M
[OH-] = 4.3 x 10-11M
[H+] = 5.0 x 10-5M
[H+] = 8.3 x 10-10M
A. pH = 9.63
B. pH = 3.65
A. pH = 4.30
B. pH = 9.08

38. More Practice Problems
Find the pH and the pOH of each solution.
[H+] = 1 x 10-4M
[H+] = M
What are the pH and pOH values of the following solutions, based on their hydrogen-ion concentrations?
[H+] = 1 x 10-12M
[H+] = 0.045M
Use either pH Calculations worksheet #1 or #2 here

39. Complete Topic 8 Acids, Bases, and Salts note packet page 146, practice problems # 26-29.

40. Relative H+ Concentrations of Solutions
The lower the pH, the more H+ ions are in a soln.
As [H+] in a soln. increases, pH of the soln. decreases.
Ex. A soln. with a pH of 3 has more H+ ions then a soln. of pH 4
Difference in H+ of two soln. = 10 (difference in pH)
That Means:

42. Relative H+ Concentrations of Soln.
Complete Topic 8 Acids, Bases, and Salts note packet page 146, practice problems #

43. Strong Vs Weak Acids
Determined by degree in which the acid dissociates in water.
Strong Acid – completely ionized in aqueous soln.
High concentration of H3O+
Ex.
Weak Acid – partially ionized in aqueous soln.
Low concentration of H3O+
HCl(g) + H2O(l) → H3O+(aq) + Cl–(aq)
CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO–(aq)

45. Acid Dissociation Constant;Ka
Acid dissociation constant (Ka) is the ratio of the concentration of the dissociated form of an acid to the concentration of the undissociated form.
Ka reflects the fraction of an acid that is ionized
Ka = [products] [reactants]
Weak acids have small Ka values.
Strong acids have large Ka values.
Dissociation/ ionization is more complete
Example: Nitrous acid (HNO2) has a Ka of 4.4 × 10−4 : strong
Ethanoic acid (CH3COOH) has a Ka of 1.8 × 10−5 : weak
dissociation constants are sometimes called ionization constants.

46. Acid Dissociation Constant;Ka
Example
CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO–(aq)
Keq=
[H3O+] × [CH3COO–]
[CH3COOH] × [H2O]

47. Calculating Dissociation Constants
You can find the Ka of an acid in water by substituting the equilibrium concentrations of the acid, [HA], the anion from the dissociation of the acid, [A−], and the hydrogen ion, [H+], into the equation below.
Ka =
[H+][A−]
[HA]

48. Calculating Dissociation Constants
In a M solution of ethanoic acid, [H+] = 1.34 × 10−3M. Calculate the Ka of this acid. Refer to the table for the ionization equation for ethanoic acid.
Start by determining the equilibrium concentration of the ions.
[H+] = [CH3COO−] = 1.34 × 10−3M
( – )M = M
Concentration
[CH3COOH]
[H+]
[CH3COO−]
Initial
0.1000
Change
−1.34 × 10−3
1.34 × 10−3
Equilibrium
0.0987

49. Strong Vs Weak Bases
Determined by degree in which the metal ions and hydroxide ions dissociate in water.
Strong Base – completely ionized in aqueous soln.
Ex.
Weak Acid – partially ionized in aqueous soln.
NaOH(aq) + H2O(l) → OH-(aq) + Na+(aq)
NH3(aq) + H2O(l) NH4+(aq) + OH–(aq)
Base
Acid
Conjugate Acid
Conjugate Base

50. Neutralization
A rx between an acid and a base to produce water and a salt.
Generally a double replacement reaction

51. Titration
A lab process used for determining the concen. Of an unknown soln. by reacting it with a soln. of known concen.
Involves an acid and a base
When moles of H+ and OH- are equal, neutralization has occurred, and the endpoint of the titrations is reached.
Indicators are used to identify the endpoint. Change in indicator color = endpoint reached.

52. Salts
Ionic compounds composed of a positive ion (other than H+) and a negative ion (other than OH-)
Salt is one of the products of an acid-base neutralization rx
Salts are electrolytes (conduct electricity when dissolved in H2O)
Soluble salts are better electrolytes than insoluble salts
Table F can be used to determine soluble and insoluble salts

53. Electrolytes
Substances that can conduct electricity when dissolved in H2O
Dissolve in H2O to produce a soln. with “+” and “–” ions.
Conduct electricity b/c of mobile ions in the soln.
Acids (Table K), bases (Table L), and salts are electrolytes