Chapters 14 & 15

Acid and Base

Equilibrium

History of Acids and Bases

In the early days of chemistry chemists were organizing physical and chemical properties of substances. They discovered that many substances could be placed in two different property categories:

Substance A

1. Sour taste
2. Reacts with carbonates to make CO₂
3. Reacts with metals to produce H₂
4. Turns blue litmus pink
5. Reacts with B substances to make salt and water

Substance B

1. Bitter taste
2. Reacts with fats to make soaps
3. Do not react with metals
4. Turns red litmus blue
5. Reacts with A substances make salt and water

​

​

​

Arrhenius Theory

The Swedish chemist Svante Arrhenius

proposed the first definition of acids and bases.

(Substances A and B became known as

acids and bases)

​

​

​

“acids are substances that dissociate in water to produce H⁺ ions and bases are substances that dissociate in water to produce OH^- ions”

​

NaOH (aq) Na⁺ (aq) + OH^- (aq) Base

HCl (aq) H⁺ (aq) + Cl^- (aq) Acid

BUT WHAT IF THE ACID/BASE IS NOT �DISSOLVED IN WATER?�

• The Arrhenius definition for acids and bases only refers to compounds dissolved in water.
• Does this mean that acids and bases cannot exist out of water?
• Not quite, that’s where the Bronsted-Lowry definition comes in.

Bronsted

Lowry

Johannes Brønsted and Thomas Lowry revised Arrhenius’s acid-base theory to include other solvents besides water.They defined acids and bases as follows:

​

An acid is a hydrogen containing species that

donates a proton.

​

A base is any substance that

accepts a proton”

​

​

HCl (aq) + H₂O (l) Cl^- (aq) + H₃O⁺ (aq)

​

In the above example,

what is the Brønsted acid?

what is the Brønsted base?

Bronsted Lowry Theory

HCl (aq) + H₂O (l) Cl ^- ( aq) + H₃O⁺ (aq)

In reality, the reaction of HCl with H₂O is an equilibrium and occurs in both directions, although in this case the equilibrium lies far to the right.

For the reverse reaction Cl ^- behaves as a Brønsted base and H₃O⁺ behaves as a Brønsted acid. The Cl^- is called the conjugate base of HCl.

​

Brønsted acids and bases always exist as conjugate acid-base pairs. Their formulas differ by only one proton.

Conjugate Pairs

Give it a Try !

Label the acid, base, conjugate acid, and conjugate base in each reaction:

HCl + OH^- →   Cl^- + H₂O

H₂O + H₂SO₄ →   HSO₄^- + H₃O⁺

Strong acids ionize 100% and weak ones do not!

​

A single arrow is used to represent the ionization of strong acids. HCl _(g) H⁺ _(aq) + Cl ^- _(aq)

​

double arrows are used to represent ionization of weak acids because an equilibrium is created.

HF _(g) H⁺ _(aq) + F ^– _(aq)

​

​

​

​

​

Common Strong Acids and Bases�Easy, memorize them!

6 Common Strong Acids

• These are the strong acids. What makes them 'strong' is that they completely dissociate into their ions when they are mixed with water. Any other acid is a weak acid!

​

• As the strong acids become more concentrated, they may be unable to fully dissociate.

​

• The rule of thumb is that a strong acid is 100% dissociated in solutions of 1.0 M or less.

​

6 COMMON STRONG BASES

• Strong bases are bases which completely dissociate in water into the cation and OH^- (hydroxide ion).

​

• The hydroxides of the Group I and Group II metals usually are considered to be strong bases.

In pure water (no solute) water molecules behave as both an acid and base!!

​

It is called amphoteric meaning it will act as either an acid or a base depending on the situation.

​

H₂O (l) + H₂O (l) H₃O⁺ (aq) + OH^- (aq)

​

This is called the autoionizationof water. Although the equilibrium lies far to the left it is very important to take into consideration, especially for living systems.

​

For pure water [OH-] = [H+]

​

​

​

​

Autoionization of Water

K_w is called ionization constant of water and is very small. As with all Kw values, it is temperature dependent.

K_w = 1.0 x 10 ^{-14 @} 25^oC

​

​

K_w = [H⁺][OH^-]

K_w =(1 x 10^-7 )(1 x10^-7 )

​

​

​

​

, This only means that the neutral value for pH is getting lower, it does not mean that the solution is becoming more acidic as the temperature increase. it

We define an aqueous solution as being

• neutral when the [H⁺] = [OH^-]
• acidic when [H⁺] > [OH^-]
• basic when [H⁺] < [OH^-]

​

​

CHECK THIS OUT !

[H⁺] = 0.0000001 = 10^-7

​

How can this be abbreviated further?

By just describing the power called the POWER OF H

pH = 7

A pH Number line

The pH scale is a way of expressing the strength of acids and bases.

pH = 14

pH = 12

pH = 7

pH = 2

[H⁺] = 10^-2

[OH ^-] = 10^-12

[H⁺] = 10^-7

[OH ^-] = 10^-7

[H⁺] =10^-12

[OH ^-] = 10^-2

[H⁺] =10^-14

[OH ^-] = 10⁰

[H⁺] < [OH ^-]

basic

[H⁺] < [OH ^-]

basic

[H⁺] = [OH ^-]

neutral

[H⁺] > [OH ^-]

acidic

acidic

basic

To calculate pH or pOH

​

pH = -log [H⁺], or pOH = -log [OH^-]

​

pH + pOH = 14 for water solutions

​

Find the pH of these:

1. 0.15 M solution of Hydrochloric acid

2) 3.00 X 10^-7 M solution of Nitric acid

​

pH Calculations

​

M= 2^nd log = 10 ^–pH M= 2^nd log = 10 ^–pOH

​

​

The number of decimal places in the log answer is equal to the number of sig figs in the original Molar [ ]

​

Find the Molarity:

A solution has a pH of 8.5. What is the Molarity of hydrogen ions in the solution?

​

Calculations for

Molarity Concentrations:

This is the calculator button

You only need one piece of information to be able to determine all the values associated with acids and base.

Acid and Base Equilibrium

• Strong acids and strong bases completely dissociate in water. Therefore they DO NOT create an equilibrium system.

​

• Weak acids and weak bases disassociate to a far lesser degree in an aqueous solution. They DO create an equilibrium system .

​

• They have their own K_c constants.

​

​

​

​

They have their own Kc Constants

K is designated K_a ACID Dissociation constant

​

K gives the ratio of ions to molecules

​

Pure substances are not shown in the K_a constant

​

The greater the dissociation, the larger the K_a will be.

Equilibrium Constants �for Weak Acids

Weak acid has K_a < 1

Leads to small [H₃O⁺] and a pH of 2 - 7

Equilibrium Constants �for Weak Bases

Weak base has K_b < 1

Leads to small [OH^-] and a pH of 12 - 7

• In general, if the percent dissociation is less than or equal to 5%, then it is completely acceptable to ignore the X value in the ice table.
• The justification for this is that K_a and K_b values are typically known only to an accuracy of ±5%. 

​

Copyright © Cengage Learning. All rights reserved

24

% dissociation = [H⁺] X 100%

[HA]

Section 14.5

Calculating the pH of Weak Acid Solutions

Calculate the pH of a 0.50 M aqueous solution of the weak acid HF. Calculate the % dissociation

(K_a = 7.2 × 10^–4)

Copyright © Cengage Learning. All rights reserved

25

EXERCISE!

Section 14.5

Calculating the pH of Weak Acid Solutions

CHAPTER 15

• Neutralization Reactions
• Buffers
• Titration Curves

​

NEUTRALIZATION REACTIONS

​

• When an acid and a base mix, the acid will donate protons to the base in what is called a neutralization reaction.

​

• There are four different ways that this can occur and it depends on the relative strengths of the acids and bases involved.

​

# 1) STRONG ACID + STRONG BASE

• Both substances dissociate completely
• The only important ions are the H+ and OH-, the rest are spectator ions.
• Water and a salt will always be created.

​

• HCl _(aq) + NaOH(aq) H₂O(l) + NaCl(salt)

​

• H⁺ (aq) + OH^- (aq) H₂O(l) (net ionic)

​

• The net ionic equation is identical for ALL strong acid/strong base reactions.

​

• pH = 7 at 25⁰C, only if there were no excess ions in the solution. This assumes that all ions have been neutralized.

​

​

​

​

​

​

​

# 2) STRONG ACID + weak base

​

​

• The strong acid dissociates completely and will donate a proton to the weak base.

​

• The product will be the conjugate acid of the weak base.

​

• HCl _(aq) + NaC₂H₃O_{2 (aq)} HC₂H₃O_{2 (aq)} + Na+ + Cl^-_(aq)

​

​

• Net ionic equation involves the proton transfer from the strong acid to the weak base.

​

• H⁺ _(aq) + C₂H₃O₂^- _(aq) HC₂H₃O_{2 (aq)} (Net Ionic)

​

#3) weak acid + STRONG BASE

​

• The strong base, will accept protons from the weak acid

​

• The products are the conjugate base of the weak acid and water.

HC₂H₃O₂ + NaOH H₂O + C₂H₃O₂^- + Na⁺

​

​

Net Ionic

HC₂H₃O_{2 (aq)} + OH^- _(aq) H₂O _(l) + C₂H₃O₂^- _(aq)

​

​

# 4) weak acid + weak base

​

• The solution contains both cations and anions and the pH is dependent on the relative numbers of each produced.

​

• Less tendency to proceed to completion than neutralizations involving strong acids and strong bases.

​

• HA (aq) + B (aq) BH⁺ (aq) + A^- (aq)

​

Summary of steps for determining� the ph of Strong acid and bases neutralizations

1. Write out the equation
2. Determine the moles of each reactant
3. Create an ICE table and determine the limiting reactant and the excess reactant.

4. Determine the [H₃O⁺] of the excess reactant

5. Use the formula:

pH = -log[H₃O⁺] to determine pH

​

​

Sample problem :

​

Calculate the pH of a solution if you mix 200ml of 0.15 M Hydrobromic acid with 100 ml of .2M potassium hydroxide.

​

OH NO! HOW DO I CALCULATE THE PH OF A SOLUTION WHENEVER WEAK BASES OR WEAK ACIDS ARE INVOLVED?

The Henderson-Hasselbalch equation is widely used by many chemists, biologists and pharmacists to accomplish this.

Use the Henderson-Hasselbalch Equation

​

​

​

​

• pK_a is found by taking the - log of K_a
• A^- is the conjugate base of the corresponding weak acid HA

​

​

​

• pK_b is found by taking the - log of K_b
• BH⁺ is the conjugate acid of the corresponding weak base B

SUMMARY OF STEPS FOR WEAK ACID AND �WEAK BASE NEUTRALIZATIONS

​

1. Write out the equation

​

2. Determine the moles of each reactant

​

3.Create an ICE table and determine the limiting reactant

and the excess reactant.

​

4. Determine the [ ] of any excess reactants and the [ ] of any conjugate acid or conjugate base products.

​

5. Calculate either the pK_a or the pK_b

​

6. Use the appropriate HH Equation

THE HENDERSON-HASSELBALCH EQUATION

Find the pH of a mixture of :

​

100.0 ml of 0.30 NH₃ and 180.0 ml of 0.10 M HCl

​

K_{b NH3} = 1.8 x 10 ^-5

​

BUFFERS

• What is it?

A buffer is a solution with a very stable pH

You can add acid or base to the buffer without greatly affecting the pH of the solution.

​

• How do buffer solutions work?

A buffer solution contains compounds that can remove any hydrogen ions or hydroxide ions that you might add to it, otherwise the pH will change.

​

Simply put, a buffer is a solution of a weak acid or base and its salt [which is its conjugate]. Therefore, any titration involving a weak A/B is buffer problem.

AP TEST TIP!

• If you are asked to construct a buffer of a specific pH and given a table of Ka’s, choose a Ka with an exponent close to the desired pH and use equal concentrations of the acid and base.

​

• The above works because pH = pK_a at the ½ equivalence point, buffer region.

​

WHAT IS A TITRATION?

• IN THIS TECHNIQUE, AN ACID (OR BASE) SOLUTION OF KNOWN CONCENTRATION (TITRANT) IS SLOWLY ADDED TO A BASE (OR ACID) SOLUTION OF UNKNOWN CONCENTRATION(ANALYTE).

​

• A PH METER OR INDICATORS�ARE USED TO DETERMINE WHEN THE SOLUTION HAS REACHED THE EQ.

�EQUIVALENCE POINT(EQ): �THE POINT AT WHICH EQUIVALENT QUANTITIES OF ACID AND BASE HAVE BEEN BROUGHT TOGETHER.

​

Titrant

Analyte

© 2015 Pearson Education, Inc.

COMPARISON OF TITRATION CURVES

This is the type you need to be most familiar with for the AP Exam

© 2015 Pearson Education, Inc.

K_a and initial pH for weak acid and strong base titrations

The pH will be always be above 7 at the equivalence point

The initial pH will be usually be at 2 or above for the weak acids.

Excess OH-

pH= 14 – (-logOH^-)

Equivalence point

Weak Acid and Strong Base Titration

Calculations forWEAK ACID/ STRONG BASE Titration Curve

​

• The pH before the titration begins. Treat as usual, the acid or base in the flask determines the pH. If weak, use an ICE table.

​

• The pH on the way to the equivalence point. You are in the “land of buffer” as soon as the first drop from the burette makes a splash and reacts to form the salt. Use your Henderson-Hasselbach

​

• The pH at the midpoint of the titration (½ eq point): pH = pKa.

​

• The pH at the equivalence point.—you are calculating the pH of the salt, all the acid or base is now neutralized [to salt + water!]

​

• The pH beyond the equivalence point—it’s stoichiometry again with a limiting reactant. Calculate the molarity of the EXCESS and solve for either pH directly (excess H+ ) or pOH (excess OH− ) and subtract it from 14 to arrive at pH.

​