ACIDS and BASES pH indicators pH indicators are valuable tool for determining if a substance is an acid or a base. The indicator will change colors in.

ACIDS and BASES

pH indicators pH indicators are valuable tool for

determining if a substance is an acid or a base.

The indicator will change colors in solution.

Things to use… pH meter will indicate the numeric

value of acid or base based on the pH range

Chemical indicators: phenolphthalein, universal indicator…

Natural indicators: poinsettia, red cabbage juice…

Properties of Acids and Bases ACIDS

Have a sour taste Change the color

of many indicators Are corrosive

(react with metals)

Neutralize bases Conduct an

electric current

BASES Have a bitter taste Change the color

of many indicators Have a slippery

feeling Neutralize acids Conduct an

electric current

The Arrhenius Theory of Acids and Bases

Arrhenius Theory of Acids and Bases:

an acid contains hydrogen and ionizes in solutions to produce H+ ions:

HCl H+(aq) + Cl-(aq)

Arrhenius Theory of Acids and Bases:

a base contains an OH- group and ionizes in solutions to produce OH- ions:

NaOH Na+(aq) + OH-(aq)

Neutralization Neutralization: the combination

of H+ with OH- to form water.

H+(aq) + OH-(aq) H2O (l)

Hydrogen ions (H+) in solution form hydronium ions (H3O+)

In Reality…

H+ + H2O H3O+

Hydronium Ion

(Can be used interchangeably with H+)

Commentary on Arrhenius Theory…

One problem with the Arrhenius theory is that it’s not comprehensive enough. Some compounds act like acids and bases that don’t fit the standard definition. So others came up with their own theories!!!!

Bronsted-Lowry Theory of Acids & Bases

Bronsted-Lowry Theory of Acids & Bases:

An acid is a proton (H+) donor

A base is a proton (H+) acceptor

for example…

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

Proton transfer

Acid

Base

NH3(aq) + H2O(l) NH4+ (aq) + OH-

(aq)

BASE

ACID

CONJUGATE ACID

CONJUGATE BASE

Ammonia is a proton acceptor, and thus a base

another example…

Water is a proton donor, and thus an

acid.

Amphoteric Substances

A substance that can act as both an acid and a base (depending on what it is reacting with) is termed amphoteric.

Water is a prime example.

Conjugate acid-base pairs

Conjugate acid-base pairs differ by one proton (H+)

A conjugate acid is the particle formed when a base gains a proton.

A conjugate base is the particle that remains when an acid gives off a proton.

Examples: In the following reactions, label the conjugate acid-base pairs:

H3PO4 + NO2- HNO2 + H2PO4

-

CN- + HCO3- HCN + CO3

2-

HCN + SO32- HSO3

- + CN-

H2O + HF F- + H3O+

acid base c. acid c. base

acidbase c. acid c. base

acid base c. acid c. base

acidbase c. acidc. base

SUMMARY OF ACID-BASE THEORIES

Theory Acid Definition Base Definition

Arrhenius Theory

Any substance which releases H+ ions in water solution.

Any substance which releases OH- ions in water solution

Brǿnsted-Lowry Theory

Any substance which donates a proton.

Any substance which accepts a proton.

Strength of Acids and Bases

A strong acid dissociates completely in sol’n: HCl H+(aq) + Cl-(aq)

A weak acid dissociates only partly in sol’n: HNO2 H+(aq) + NO2

-(aq)

A strong base dissociates completely in sol’n: NaOH Na+(aq) + OH-(aq)

A weak base dissociates only partly in sol’n: NH3(aq) + H2O(l) NH4

+(aq) + OH-(aq)

Acid-Base Reactions Neutralization reactions: reactions

between acids and metal hydroxide bases which produce a salt and water.

A salt is a metal and a non metal!! H+ ions and OH- ions combine to form

water molecules:

H+(aq) + OH-(aq) H2O(l)

Buffered Solutions

A solution of a weak acid and a common ion is called a buffered solution.

Thus, the solution maintains it’s pH in spite of added acid or base.

Universal Indicator Color ChartThis is in Chapter 19

pH scale 0 7 14

Acid Neutral Base

pH and pOH

Ionization of water Experiments have shown that pure

water ionizes very slightly: 2H2O H3O+ + OH-

Measurements show that: [H3O+] = [OH-]=1 x 10-7 M

Pure water contains equal concentrations of H3O+ + OH-, so it

is neutral.

So what is pH?

pH is a measure of the concentration of hydronium ions in a solution.

pH = -log [H3O+] or

pH = -log [H+]

Example: What is the pH of a solution where [H3O+] = 1 x 10-7 M?

pH = -log [H3O+] pH = -log(1 x 10-7)pH = 7

Example: What is the pH of a solution where [H3O+] = 1 x 10-5 M?

pH = -log [H3O+] pH = -log(1 x 10-5)pH = 5

When acid is added to water, the [H3O+] increases, and the pH decreases.

Example: What is the pH of a solution where [H3O+] = 1 x 10-10 M?

pH = -log [H3O+] pH = -log(1 x 10-10)pH = 10

When base is added to water, the [H3O+] decreases, and the pH increases.

The pH Scale

Acid Neutral Base

0 7 14

What is pOH?

pOH is a measure of the concentration of hydroxide ions in a solution.

pOH = -log [OH-]

Example: What is the pOH of a solution where [OH-] = 1 x 10-5 M?

pOH = -log [OH-] pOH = -log(1 x 10-5)pOH = 5

How are pH and pOH related?

At every pH, the following relationships hold true:

[H3O+] • [OH-] = 1 x 10-14 M

pH + pOH = 14

Example 1: What is the pH of a solution where [H+] = 3.4 x 10-5 M?

pH = -log [H+] pH = -log(3.4 x 10-5 M)pH = 4.5

Example 2: What is the pH of a solution where [H+] = 5.4 x 10-6 M?

pH = -log [H+] pH = -log(5.4 x 10-6)pH = 5.3

Example 3: What is the [OH-] and pOH for the solution in example #2?

[H3O+][OH-]= 1 x 10-14

(5.4 x 10-6)[OH-] = 1 x 10-14

[OH-] = 1.9 x 10-9 M

pH + pOH = 14 pOH = 14 – 5.3 = 8.7

Example #4 Classify each solution as acidic, basic, or

neutral *MUST SOLVE FOR pH and use the pH

scale

a. [H+] = 6.0 x 10-10 Mb. [OH-] = 3.0 x 10-2 Mc. [H+] = 2.0 x 10-7 Md. [OH-] = 1.0 x 10-7 M

basicbasicacidicneutral

Example #5 Which is the MOST basic from

question #4?

B.

Acids and bases: Titrations The amount of acid or base in a

solution is determined by carrying out a neutralization reaction;

an appropriate acid-base indicator (changes color in specific pH range) must be used to show when the neutralization is completed.

Buret

Solution with Indicator

This process is called a titration: the addition of a known amount of solution to determine the volume or concentration of another solution.

Read a buret volume to 2 decimal places

Textbook page 619

End point: the point at which the indicator changes color

3 Steps to do a titration (pg. 618):

1. Add a measured amount of an acid of unknown concentration to a flask.

2. Add an appropriate indicator to the flask

3. Add measured amounts of a base of known concentration using a buret. Continue until the indicator shows that neutralization has occurred. This is called the end point of the titration

(show lab in demo form…)

4 steps to a titration CALC: 1) balance the equation 2) calculate the number of moles of acid

or base in known solution 3) calculate the number of moles in

unknown solution used during the titration

4) determine molarity of unknown solution and the pH (M=mol solute/ L solvent)

Example: In a titration, 27.4 mL of 0.0154 M

Ba(OH)2 is added to a 20.0 mL sample of HCl solution of unknown concentration. What is the molarity and pH of the acid solution?

Equation: (Step 1)Ba(OH)2 + 2 HCl BaCl2 + 2 H2O

(steps)

Step 2 Calculate the number of moles of

known solution (Ba(OH)2)

Calculate moles of known solution:

Mol Ba(OH)2 = 0.0154 127.4

1 1000

mol LmLx x

L mL

4.22 x 10-4 mol Ba(OH)2

Step 3 Calculate moles of unknown

solution

Calculate moles of unknown solution:

Mol HCl = 4.22 x 10-4 mol Ba(OH)2 x 2 mol HCl= 8.44 x 10-4 mol HCl

1 mol Ba(OH)2

Step 4 Determine molarity and pH

Molarity = 8.44 x 10-4 mol HCl = 4.22 x 10-2 M HCl

0.0200 L

pH: HCl dissociates into H+ and Cl- ions

4.22 x 10-2 M HCl = 4.22 x 10-2 M H+ = 4.22 x 10-2 M Cl-

pH = -log[4.22 x 10-2 M H+] = 1.375 = 1.38

Extra Calculation (same steps as the one we just did)

A 25 mL solution of H2SO4 is completely neutralized by 18 mL of 1.0 M NaOH. What is the concentration of the H2SO4 solution?

Titration Curve (pg. 619) A graph showing how the pH

changes as a function of the amount of added titrant in a titration.

Data for the graph is obtained by titrating a solution and measuring the pH after EVERY drop of added titrant.

Equivalence point = The point on the curve where the

moles of acid equal the moles of base

the midpoint of the steepest part of the curve is a good approximation of the equivalence point.

Knowledge of the equivalence point can then be used to choose a suitable indicator for a given titration;

the indicator must change color (end point) at a pH that corresponds to the equivalence point.

HANDOUT: titration curve WS