CHAPTER 3

CONCEPT OF ACID-BASE NEUTRALIZATION

CONCEPT OF ACID-BASE NEUTRALIZATION

CHM 138

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ACID-BASE THEORY

ARRHENIUS BRONSTED-LOWRY LEWIS

ARRHENIUS BRONSTED-LOWRY LEWIS

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Arrhenius theoryArrhenius theory

Acid : a substance that, when dissolved in water, produce H+ ions

Base : a substance that, when dissolved in water, increases the concentration of OH- ions.

Acid : a substance that can donate a proton to a base

Base : a substance that can accept a proton from an acid.

An acid is a proton donor and a base is a proton acceptor.

Bronsted Lowry theory

When an acid gives up its proton, what remains is called the conjugate base of that acid.

When a base accepts a proton, the resulting chemical is called the conjugate acid of that original base.

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Arrhenius theory

acid base

Conjugate acid

Conjugate base

Learning Check!

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

HCl + OHHCl + OH--   Cl   Cl-- + H + H22OO HCl + OHHCl + OH--   Cl   Cl-- + H + H22OO

HH22O + HO + H22SOSO44   HSO   HSO44-- + H + H33OO

++ HH22O + HO + H22SOSO44   HSO   HSO44-- + H + H33OO

++

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LEWIS ACID-BASE

Lewis acid - a substance that accepts an electron pair

Lewis base - a substance that donates an electron pair

Example : Formation of hydronium ion H

H

H

BASE

••••••

O—HO—H

H+

ACID

ACID

E.g : HCl, HNO3, H2SO4 – strong acids CH3COOH – weak acid (due to its incomplete ionisation) CH3COOH (aq) CH3COO- (aq) + H+ (aq)

Properties of acids:- Sour taste- Change litmus colour from blue to red- React with certain metals- React with carbonates and bicarbonates to produce CO2 gas.- Conduct electricity- React with base forming salts and water.

CONT….

ACID Acids can be described as monobasic, dibasic or tribasic etc.

• depending on the maximum number of protons that are available for transfer in an acid-base reaction.

monobasic acids e.g. • hydrochloric HCl, nitric HNO3, ethanoic CH3COOH (the alkyl H's are not

acidic), dibasic acids e.g.

• sulphuric H2SO4, ethanedioic (COOH)2, and the three isomeric

• benzene-x,y-dicarboxylic acids (x,y = 1,1 1,2 and 1,3) C6H4(COOH)2, tribasic acids e.g.

• boric acid H3BO3, phosphoric(V) H3PO4,

• citric acid , the middle-left hydrogen of the HO-C (alcohol) is not acidic in water.

BASE

E.g. NaOH, Ba(OH2), LiOH, KOH – strong bases

NH3 – weak base

Properties of bases Bitter taste Slippery Change litmus colour from red to blues Conduct electricity React with acids to form salts and water

ACID-BASE NEUTRALISATION

Acids and bases neutralise one another forming salt and water.

Acid + base salt + water Salt – made up of a cation other than H+ and an

anion other than OH- or O2-.

E.g: HNO3 (aq) + NaOH (aq) NaNO3 (aq) +H20(l)

ACIDS & BASES STRENGTH

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ACIDS AND BASES STRENGTH

% Ionisation of acid = [H+] X 100 initial molarity of acid

% Ionisation of base = [OH-] X 100 initial molarity of base

Acid

Strong acids are completely dissociated in water.

Weak acids only dissociate partially in water.

Strong Acids

the seven strong acids are HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4.

These are strong electrolytes and exist totally as ions in aqueous solution

Strong Bases

Strong bases are the soluble hydroxides, which are the alkali metal (NaOH, KOH)and heavier alkaline earth metal hydroxides (Ca(OH)2 and Ba(OH)2).

Again, these substances dissociate completely in aqueous solution.

ELECTROLYTES

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electrolyte any substance containing free ions that make the substance electrically

conductive. STRONG ELECTROLYTE• Completely ionised in aqueous solution• Conducts a strong electric current Good conductor WEAK ELECTROLYTE• Only slightly ionised in aqueous solution• Conducts a weak electric current Weak conductor NON ELECTROLYTE• Does not ionise in aqueous solution• Remains as molecules• Does not conduct an electric current Non conductor 19

pH, pOH, [ H+] , [ OH-]

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pH The pH of a solution is given by the negative logarithm of

the hydrogen ion concentration, [H+] (in mol/dm3 or M)

pH = - log [H+] pH for acidic solutions < 7, basic solutions > 7 and neutral solutions = 7

pOH = - log [OH-] pH + pOH = 14.00

[OH-]

[H+] pOH

pH

10 -pOH

10 -pH-Log[H+]

-Log[OH

-Log[OH

--]]

14 -

pOH

14 -

pH

1.0

x 10

1.0

x 10-1

4-14

[OH[O

H-- ]]

1.0

x 10-1

4

[H

+ ]

Calculating the pH

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

Example: If [H+] = 1 X 10-10

pH = - log 1 X 10-10

pH = - (- 10)pH = 10

Example: If [H+] = 1.8 X 10-5

pH = - log 1.8 X 10-5

pH = - (- 4.74)pH = 4.74

Try These!

Find the pH of these:1) A 0.15 M solution of Hydrochloric acid2) A 3.00 X 10-7 M solution of Nitric acid

pH calculations – Solving for H+pH calculations – Solving for H+

If the pH of Coke is 3.12, [H+] = ???

Because pH = - log [H+] then

- pH = log [H+]

Take antilog (10x) of both sides and get

Antilog -pHAntilog -pH == [H[H++]][H[H++] = antilog -3.12 = 7.6 x 10] = antilog -3.12 = 7.6 x 10-4-4 M M *** to find antilog on your calculator, look for “Shift” or “2*** to find antilog on your calculator, look for “Shift” or “2nd nd

function” and then the log buttonfunction” and then the log button

pH calculations – Solving for H+pH calculations – Solving for H+

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

pOH Since acids and bases are opposites, pH

and pOH are opposites!

pOH does not really exist, but it is useful for changing bases to pH.

pOH looks at the perspective of a basepOH = - log [OH-]

Since pH and pOH are on opposite ends,

pH + pOH = 14

[H[H33OO++], [OH], [OH--] and pH] and pH

What is the pH of the 0.0010 M NaOH solution?

The pH of rainwater collected in a certain region of the northeastern United States on a particular day was 4.82. What is the H+ ion concentration of the rainwater?

The OH- ion concentration of a blood sample is 2.5 x 10-7 M. What is the pH of the blood?

Practice

ACID-BASE TITRATION

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ACID-BASE TITRATIONS Acid-base titration is a procedure used in quantitative analysis to

determine the molarity of acid or alkali.

Point at which the acid-base indicator changed colour is the end-point.

Equivalence point is the point where equal mole of H3O+ ions and OH- ions in the titration flask. At this point, all acid has been neutralized by base and contain only salt and water.

Strong acid-strong base titrations (end-point pH 7)

Setup for titrating an acid with a baseSetup for titrating an acid with a base

Strong acid-strong base titrations

base

acid

Strong acid-strong base titrations

base

acid

acid

base

INDICATORS Weak organic acid/base of distinctly different colours in its ionised

and unionised forms, Used to indicate the equivalence point in an acid-base titration by a

change in colour.

VOLUMETRIC ANALYSIS

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MOLARITY• The number of moles of solute per

liter of solution

MOLARITY, M = NO OF MOLE OF SOLUTE, g

VOLUME OF SOLUTION, L

CONCENTRATION = molarity x molar mass

Begin with a balanced equation for the reaction: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

na = 1 nb = 1 (mole ratios of acid and base)

Mole = Molarity X volumeFor the acid: na = MaVa

For the base: nb = MbVb

na : nb (stoichiometry mole ratio)

MaVa : MbVb

Volumetric analysis

Ma Va / Mb Vb = na / nb

Chapter 15 39

Example

Consider the titration of acetic acid with sodium hydroxide. A 10.0 mL sample of acetic acid requires 37.55 mL of 0.223 M NaOH. What is the molarity of the acetic acid?

A 10.0 mL sample of 0.555 M H2SO4 is titrated with 0.233 M NaOH. Calculate volume of NaOH required for the titration?

THE END ….THE END ….

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