1The Chemistry of Acids and Bases

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Some Properties of Acids

þ React with certain metals to produce hydrogen gas.

þ React with carbonates and bicarbonates to produce carbon dioxide gas

þ Taste sour

þ Corrode metals

þ Electrolytes

þ React with bases to form a salt and water

þ pH is less than 7

þ Turns blue litmus paper to red “Blue to Red”

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Some Properties of Bases

Taste bitter, chalky

Are electrolytes

Feel soapy, slippery

React with acids to form salts and water

pH greater than 7

Corrosive

Turns red litmus paper to blue “Basic Blue”

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Acid and Bases

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Acid and Bases

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Acid and Bases

8Indicators

• Indicators are dyes that can be added in small amounts that will change color in the presence of an acid or base.

• Some indicators only work in a specific range of pH

9Examples: Indicators

• Litmus paper• Phenolpthalein• Bromothymol blue• Methyl Orange

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Universal Indicator

• Universal indicator is a pH indicator composed of a blend of several compounds that changes colour over a wide range of pH values from 0-14.

Cabbage juice

pH paper

12• A universal indicator is typically composed of water, methanol, proan-1-ol, phenolpthalein, sodium salt, methyl red, bromothylmol blue, monosodim salt and thymol blue monosodium salt.

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pH• Measure of H+ ions in

solution. • High [H+] ions = more

acidity and low pH.• pH meter tests the

voltage of the electrolyte

• Converts the voltage to pH

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The pH scale is a way of expressing the strength of acids and bases. Under 7 = acid

7 = neutral Over 7 = alkaline (base)

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pH of Common Substances

16pH is a logarithmic function

• pH 2 ____ more acidic than a pH of 3• pH 2 ____ more acidic than a pH of 4• pH 2 ____ more acidic than a pH of 5

pH = - log [H+](Remember that the [ ] mean

Molarity)

17Strong and Weak Acids/Bases

The strength of an acid (or base) is determined by the amount of IONIZATION (dissociation).

18Strong Acids

• A strong acid is one that completely 100% ionizes in water.

Example:

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

HNO3, HCl, H2SO4 and HClO4 are among the only known strong acids. Most acids are weak.

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• Strong Base: 100% dissociated into its respective ions water.

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

Strong Bases

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• Weak acids ionize less than 100% in water.

Weak Acids

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HClO (aq) H+(aq) + ClO-

(aq)

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Weak Acid Dissociation(only partially

ionizes)

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Weak base: less than 100% ionized in waterOne of the best known weak bases is ammonia

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

Weak Bases

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Weak Bases

25Conductivity

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More About Water

In pure water there can be AUTOIONIZATION

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Equilibrium constant for water = Kw

It can be seen from the above equation that H3O+ and

OH¯ concentrations are in the molar ratio of one-

to-one. This means that [H3O+] = [OH¯].

Kw = [H3O+] [OH-] = 1.00 x 10-14 at 25 oC

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Acid/Base definitions

Arrhenius (Swedish chemist)

29Arrhenius acid is a substance that produces H+ in water

Arrhenius base is a substance that produces OH- in water

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POLYPROTIC ACIDS

Monoprotic

Diprotic

Triprotic

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Limitations to Arrhenius’ Theory

An acid is expected to be an acid in any solvent.

But that’s not the case nowadays. For example HCL acts as an Arrhenius acid when dissolved in water. However when HCL is dissolved in benzene there is no dissociation. This is against Arrhenius theory; Arrhenius states that dissociation occurs in any aqueous solution. The properties of acid and bases play a critical role.  Water does not have to be the only solvent.

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Limitations to Arrhenius’ Theory

2- In Arrhenius theory all salts produced in the neutralization reaction should produce solutions that are neither acidic nor basic. But there are some exceptions against this theory. For example if equal amounts of HCl and ammonia react, the solution is slightly acidic. If equal amounts of acetic acid and sodium hydroxide react, the resulting solution is basic. Arrhenius theory does not include any explanation for this.  

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Limitations to Arrhenius’ Theory

3- The proton H+ liberated from the acid actually combines with water to form hydronium ions in a water solution. Therefore this reaction: H2O + H+ H3O+ Occurs most of the time. 

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Limitations to Arrhenius’ Theory

4- The need for hydroxide as the base led Arrhenius to propose the formula NH4OH as the formula for ammonia in water. This led to the misunderstanding that NH4OH is the actual base. But the actual base is NH3. 

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Acid/Base DefinitionsBrønsted – Lowry

Acids – any species that can donate a proton (H+ ions) in solution ‘proton donor’

Bases – any species that accepts a proton (H+ ions) in solution proton acceptor

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A Brønsted-Lowry acid is a proton donorA Brønsted-Lowry base is a proton acceptor

acidconjugate

basebase conjugate

acid

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ACID-BASE THEORIES

The Brønsted definition means NH3 is a BASE in water — and water is itself an ACID

BaseAcidAcidBaseNH4

+ + OH-NH3 + H2O

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Conjugate Pairs

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Learning Check!

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

HCl + OH-   Cl- + H2O

H2O + H2SO4   HSO4- + H3O

+

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Amphoteric

• A substance that can act as an acid in one situation and a base in another.

• Ex• H20 + HCl H3O+ + Cl-• H20 + NH3 NH4 + OH-

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ACID-BASE REACTIONSTitrations

H2C2O4(aq) + 2 NaOH(aq) --->

acid base

Na2C2O4(aq) + 2 H2O(liq)

Carry out this reaction using a TITRATION.

Oxalic acid,

H2C2O4

42Setup for titrating an acid with a base

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Titration1. Add solution from the buret.2. Reagent (base) reacts with

compound (acid) in solution in the flask.

3. Indicator shows when exact stoichiometric reaction has occurred. (Acid = Base)

This is called NEUTRALIZATION.