Topic 8 Acids and bases 6 hours

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Topic 8 Acids and bases 6 hours

Hydronium ion (H

3

O

⁺

) = more stable form of hydrogen ion (H

⁺

) H

⁺

+ H

2

O → H

3

O

⁺

8.1 Theories of acids and bases 2 hours

1. Arrhenius H-X / M-OH

ACID – a substance that dissociates in water to produce one or more hydrogen ions (H

⁺

)

H-X, where X is any anion ex. hydrochloric acid HCl

(aq)

→ H

⁺(aq)

+ Cl

–(aq)

BASE – a substance that dissociates in water to produce one or more hydroxide ions (OH

^–

)

M-OH metal hydroxides ex. sodium hydroxide NaOH

(aq)

→ Na

+(aq)

+ OH

–(aq) 8.1.1 Define acids and bases according to the Brønsted–Lowry and Lewis theories.

Discuss the value of using different theories to explain the same phenomenon. What is the relationship between depth and simplicity?

8.1.2 Deduce whether or not a species could act as a Brønsted–Lowry and/or a Lewis acid or base.

8.1.3 Deduce the formula of the conjugate acid (or base) of any Brønsted–Lowry base (or acid).

Students should make clear the location of the proton transferred, for example, CH3COOH/CH3COO^– rather than C2H4O2/C2H3O2–.

2. Brønsted-Lowry ACID = PROTON DONOR BASE = PROTON ACCEPTOR Proton p

⁺

= H

⁺

→ hydrogen nucleus

ACID – a substance that donates a proton to another substance (PROTON DONOR) ex. acetic acid (vinegar) CH

3

COOH É CH

3

COO

–

+ H

⁺

carboxylic acids have a carboxyl group (COOH) that is stable with a H

⁺

removed

BASE – a substance that accepts a proton from another substance (PROTON ACCEPTOR)

ex. ammonia NH

3

+ H

⁺

É NH

4+

ammonia is a Brønsted-Lowry base 3. Lewis ACID = ELECTRON ACCEPTOR BASE = ELECTRON DONOR

Gilbert Lewis based his definition on chemical bonding theory where pairs of valence electrons form bonds ACID – a substance that can accept a pair of non-bonding electrons from a Lewis base (lone pair).

BASE – a substance that can donate a pair of non-bonding electrons to a Lewis acid

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8.2 Properties of Acids and Bases 1 hour

8.2.1 Outline the characteristic properties of acids and bases in aqueous solution.

Bases that are not hydroxides, such as ammonia, soluble carbonates and hydrogencarbonates, should be included.

Alkalis are bases that dissolve in water.

Students should consider the effects on indicators and the reactions of acids with bases, metals and carbonates.

Property ACID BASE

Arrhenius form H

⁺

(H

3

O

⁺

) ions form OH

^–

ions

Brønsted-Lowry proton (H

⁺

) donor proton (H

⁺

) acceptor

D ef init ion

Lewis Lone pair acceptor Lone pair donor

Feel  soapy

Se ns es

Taste  sour bitter

M ea sur e

pH LOW pH (pH < 7) HIGH pH (pH > 7)

Litmus red blue

Phenolphthalein clear pink

Indic at or s

Bromothymol Blue yellow blue

Carbonates

(baking soda,limestone) carbon dioxide gas evolved (bubbles)

MCO3(s) + HX(aq) → MX^(aq) + H2O(l) + CO2(g) ↑

no reaction

R ea ct ions

Active Metals

(Mg, Al, Zn, Fe) hydrogen gas evolved (bubbles)

M(s) + HX (aq) → MX^(aq) + H2(g)↑

no reaction Strong

complete (100%) ionization

HCl, hydrochloric acid H

2

SO

4

, sulfuric acid

HNO

3

, nitric acid

NaOH, sodium hydroxide (lye) KOH, potassium hydroxide (lye)

Weak

partial ionization –

equilibrium

H

2

CO

3

, carbonic acid All carboxylic acids:

CH3COOH, acetic acid; HCOOH, formic acid

Some amino acids: Asp, Glu

NH

3

, ammonia (NH

4

OH) Amines (1º, 2º, 3º) Some amino acids: Arg, Lys carbonates

(CaCO3 -chalk, limestone)

hydrogencarbonates

(NaHCO3 - baking soda)

Oxides Non-metal oxides (CO

2

, SO

3

) Metal oxides (Na

2

O, CaO) Household citric acid, lactic acid, vinegar detergents, cleaners

 = IF AN ACID OR BASE TOUCHES YOUR SKIN, RINSE IMMEDIATELY WITH WATER!

 = NEVER TASTE CHEMICALS IN THE LAB!

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8.3 Strong and weak acids and bases 2 hours

8.3.1 Distinguish between strong and weak acids and bases in terms of the extent of dissociation, reaction with water and electrical conductivity.

Although weakly acidic solutions are relatively safe, they still cause damage over long periods of time.

Students could consider the effects of acid deposition on limestone buildings and living things.

8.3.2 State whether a given acid or base is strong or weak.

Students should consider hydrochloric acid, nitric acid and sulfuric acid as examples of strong acids, and carboxylic acids and carbonic acid (aqueous carbon dioxide) as weak acids.

Students should consider all group 1 hydroxides and barium hydroxide as strong bases, and ammonia and amines as weak bases.

8.3.3 Distinguish between strong and weak acids and bases, and determine the relative strengths of acids and bases, using experimental data.

The following is a comparison using solutions of the same concentration (such as 1.0 mol dm

^–³

)

STRONG weak

Ionisation

^molecule

→

^cation⁺^{+ anion}^–

 complete ionization (Kc > 10¹⁰)

 IONS ONLY - no molecular form

molecule

É

^cation⁺^{+ anion}^–

 partial ionisation (equilibrium)

 mostly molecular form, some ions

ACID

reactions with water

HA + H

2

O → A

^–

+ H

3

O

⁺ [H⁺]= [HA]

H2SO4 + H2O → HSO^4– + H3O+

HNO3, H2SO4, HCl

HA + H

2

O É A

^–

+ H

3

O

⁺ [H⁺]<< [HA]

CH3COOH + H2O É CH3COO^– + H3O⁺ carboxylic/organic acids, carbonic acid

BASE

reactions in water

MOH → M

+

+ OH

^– [OH^–]= [MOH]

KOH → K⁺+ OH^– NaOH, KOH, Ba(OH)2

MOH É M

+

+ OH

^– [OH^–]<< [MOH]

NH4OH É NH4+ + OH^– NH3, amines

Conductivity

based on total ion concentration

HIGHER conductivity higher [ion] =[HA] or [MOH]

lower conductivity small [ion] – little ionisation

Reaction Rate

acids with reactive metals or carbonates based on concentration of reactive ions

faster initial rate

 [ions] →collision frequency

2 HCl + Mg → MgCl² + H2(g)

HNO3 + NaHCO3 → NaNO³ + H2O + CO2(g)

bubbles faster

slower initial rate

 [ions] →collision frequency

2 CH3COOH + Mg → Mg(CH³COO)2 + H2(g)

CH3COOH + NaHCO3

→ CH³COO^–Na⁺ + H2O + CO2(g) bubbles slowly

Conjugate Acid – Base Pairs (p. 554 – 556)

Conjugate Acid – Base Pair two molecules or ions that are related by the transfer of one proton

• Conjugate Acid = Base + 1 H

⁺

NH

3

+ H

⁺

⇔ NH

4+

Base: NH

3

Conjugate Acid: NH

4+

• Conjugate Base = Acid – 1 H

⁺

H

2

SO

4

⇔ HSO

4-

+ H

⁺

Acid: H

2

SO

4

Conjugate Base: HSO

4-

HSO

4-

⇔ SO

4-2

+ H

+

Acid: HSO

4-

Conjugate Base: SO

42-

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Monoprotic and Polyprotic Acids

Monoprotic Acids  Molecules that dissociate in water to produce only one H

⁺

(proton) ex. HCl, CH

3

COOH, HNO

3

Polyprotic Acids  molecules that dissociate in water to produce two or more H

⁺

(protons) ex. H

2

SO

4

, H

2

CO

3

, H

3

PO

4

The first proton dissociates more (high K

a

value) H

2

SO

4

⇔ HSO

4-

+ H

⁺

Subsequent protons do not dissociate as easily (low K

a

) HSO

4-

⇔ SO

4-2

+ H

+

Functional Definitions

Acid – any substance that neutralizes a base

– reduces the amount of hydroxide ions in solution OH

–(aq)

Base – any substance that neutralizes an acid

– reduces the amount of hydronium ions in solution H

+(aq)

Amphiprotic Substance any substance that neutralizes both acids and bases can act as a proton donor (acid) or a proton acceptor (base)

ex. Hydrogen sulfate ion (HSO

4–

) HCl + HSO

4–

⇔ H

2

SO

4

+ Cl

^–

neutralizes acid (base)

NaOH + HSO

4–

⇔ NaSO

4–

+ HOH neutralizes base (acid)

Hydrogen carbonate ion (HCO

3–

) HCO

3–

⇔ H

+

+ CO

32–

proton donor (acid)

HCO

3–

+ H

⁺

⇔ H

2

CO

3

proton acceptor (base)

Neutral Solutions pH = 7

1. Non-ionic substances elements, most molecules, all ionic compounds

all organic compounds (except carboxylic acids and amines) 2. Water and Amino Acids equal amounts of H

+

and OH

^-

produced

Amino acids with equal numbers of COOH and NH

2

groups

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8.4 The pH scale 1 hour

8.4.1 Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale.

8.4.2 Identify which of two or more aqueous solutions is more acidic or alkaline using pH values.

Students should be familiar with the use of a pH meter and universal indicator.

8.4.3 State that each change of one pH unit represents a 10-fold change in the hydrogen ion concentration [H⁺(aq)].

Relate integral values of pH to [H+(aq)] expressed as powers of 10. Calculation of pH from [H+(aq)] is not required.

TOK: The distinction between artificial and natural scales could be discussed.

8.4.4 Deduce changes in [H⁺(aq)] when the pH of a solution changes by more than one pH unit.

Aim 8: A study of the effects of small pH changes in natural environments could be included.