HSC CHEMISTRY CORE TOPIC 2

HSC CHEMISTRY CORE TOPIC 2

THE ACIDIC THE ACIDIC ENVIRONMENTENVIRONMENT

THE ACIDIC ENVIRONMENT

VOLUMETRIC ANALYSIS

TITRATION

BRONSTEDLOWRYTHEORY

ESTERIFICATION

HISTORYLavoisier

DavyArrhenius ACIDIC

OXIDES

LE CHATELIER’S

PRINCIPLE

CARBOXYLICACIDS

ACID RAIN

pH

CHEMICAL EQUILIBRIUM

INDICATORS

NEUTRALISATION

Subsection 1Indicators were identified with the observation that the colour of some flowers depends on soil composition

Indicators

• substances that have distinctive colours in different types of chemical environments

• natural acid-base indicators are vegetable dyes that provide the colour of flowers and vegetables

• LITMUS is a pink mixture of compounds extracted from lichens grown mainly in the Netherlands

• today many indicators used are manufactured dyes

INDICATORS

INDICATOR pH RANGECOLOUR RANGE(low pH – high pH)

Methyl orange 3.1 – 4.4 red (orange) yellow

Bromothymol blue 6.0 – 7.6 yellow (green) blue

Phenolphthalein 8.2 – 10.0 colourless - crimson

Litmus 5.5 – 8.0 red - blue

INDICATORS

INDICATORS

Universal Indicator is a mixture of • thymol blue• methyl red• bromothymol blue• phenolphthalein• dissolved in methanol, propan-1-ol and water

INDICATORS

A solution gives the following colours in each indicator. Deduce the approximate pH

Indicator Colour

methyl orange yellow

bromothymol blue yellow

phenolphthalein colourless

>4.4

<6.0

<8.0

pH of solution is 4.4 – 6.0

INDICATORS

A solution gives the following colours in each indicator. Deduce the approximate pH

Indicator Colour

bromothymol blue blue

thymolphthalein colourless

phenol red red

> 7.6

< 9.5

> 8.4

pH of solution is 8.4 – 9.5

INDICATORS

USES• Universal indicator used to test soil

acidity/alkalinity (pH)• plants have preference for alkaline/neutral/acid

soils – choice of crop• diseases that affect plants thrive in soils with a

particular pH range• pH affects availability of nutrients

• Phenol red used to test acidity/alkalinity of a swimming pool – level of disinfection

Clay Minerals in Soil

Metal ions bind to negative surface

H+ are able to displace the surface cations from the clay. If aluminium ions are displaced by H+ the soil becomes toxic to crops growing in the soil.

Subsection 2

While we usually think of the air around us as neutral, the atmosphere naturally contains acidic oxides of carbon, nitrogen and sulfur. The concentrations of these acidic oxides have been increasing since the Industrial Revolution

Oxides of Period 3

NaOH(s)

MgO(s) + H2O(aq)

S(s) + O2(g) SO2(g) + H2O(aq)

P4(s) + O2(g) P2O5(s)

P2O5(s) + H2O(aq)

Oxides of Period 3 NaOH(s) Na+

(aq) + OH-(aq)

MgO + H2O Mg(OH)2

S + O2 SO2

SO2 + H2O H2SO3 (sulfurous acid)

2H2SO3 + O2 2H2SO4 (sulfuric acid)

P4 + 5O2 2P2O5

P2O5 + 3H2O 2H3PO4 (phosphoric acid)

Cl2O7 + H2O 2HClO4 (perchloric acid)

Oxides of Period 3

Oxide/

HydroxideNa2O

(NaOH)

MgO Al2O3 SiO2 P2O5 SO2 Cl2O7

Bonding

Acid/base property

amphoteric acidic

Acid/Base Properties of Oxides

BASIC OXIDES• metal oxides and hydroxides (ionic

compounds)• soluble oxides react with water to form

alkaline/basic solutions Na2O(s) + H2O(l) 2NaOH(aq)

• react with acids/acidic oxides to form salts CuO(s) + H2SO4(aq) CuSO4(aq) + H2O(l)

CaO(s) + SO2(g) CaSO3(s)


ACIDIC OXIDES• generally oxides of non-metals (covalent

compounds• called acid anhydrides• react with water to produce an acidic solution

CO2(g) + H2O(l) H2CO3(aq)

• react with bases to form saltsCO2(g) + CaO(s) CaCO3(s)

2NaOH(l) + SiO2(s) Na2SiO3(s) + H2O(l)


AMPHOTERIC OXIDES react with acids and bases Al2O3, BeO, ZnO, PbO, SnO

Al2O3(s) + 6HCl(aq) 2AlCl3(aq) + 3H2O(l)

Al2O3(s) + 2NaOH(aq) 2NaAlO2(aq) + H2O(l)

NEUTRAL OXIDES do not react with acids or bases CO, N2O, NO

Group A Oxides

105

Db107

Bh

basic amphoteric acidic3A 4A 5A 6A 7A

1A 8A

Group B

Salts

in general acids and bases form salts the type of salt is determined by the acid

HCl chloride salts HNO3 nitrate salts

H2SO4 sulfate salts

H3PO4 phosphate salts

A Saturated Solution of NaCl

Dissolution

NaCl(s) Na+(aq) + Cl-

(aq)

Solid NaCl

Solution with Na+ and Cl- ions

Na+

Cl-

NaCl

System at 25oC


Solid NaCl


Radioactive 24NaCl

Na+

Cl-

Radioactive 24NaCl is introduced into the beaker


Dissolution

NaCl(s) Na+(aq) + Cl_

(aq)

Recrystallisation

NaCl(s) Na+(aq) + Cl_

(aq)


Solid NaCl


Radioactive 24NaCl

Na+

Cl-

24NaCl

24Na+

24NaCl

Cl-

Chemical Equilibrium


• only occurs in a closed system• no interchange of matter between system and

surroundings• occurs in physical and chemical systems• temperature of the system remains constant• the rate of the forward reaction equals the

rate of the reverse reaction• refers to reversible reactions where the

forward reaction occurs simultaneously with reverse reaction


Le CHATELIER’S PRINCIPLE

• French chemist Henri Le Chatelier (1850-1936) studied changes in systems that were in a state of equilibrium

• If a stress is applied to a system in a state of chemical equilibrium, the system changes to relieve the stress

• may be changes in • concentration• volume and pressure• temperature

CONCENTRATION• increasing the concentration of a reactant or

product will cause the system to favour the direction which will decrease the concentration of that substance

• decreasing the concentration of a reactant or product means the rate of the reaction using up that substance will decrease in rate• the rate of the other reaction to produce that

substance will now have the faster rate

CONCENTRATION

equilibria with solids and pure liquids

e.g. C(s) + H2O(g) CO(g) + H2(g)

PRESSURE• changes in pressure have little effect on solids

and liquids as they are only very slightly compressible

• changes in pressure have significant effects on the concentration of GASES

• gas pressure is proportional to the number of molecules

• changing the partial pressure of a gas changes its concentration

Partial Pressure of a Gas

PRESSURE

Adding an inert gasdoes not change the partial pressures of any of the other gasesno effect on equilibrium

VOLUME CHANGES

reducing the volume of the gas by half doubles the pressure of the gas

VOLUME CHANGES

in the reaction above there will be a shift to the right as the forward reaction increases in rate to minimise the pressure changes

e.g. 2SO2(g) + O2(g) 2SO3(g)

• decreasing volume - reaction rate increases in the direction that produces the smaller number of molecules

• producing less molecules reduces the pressure• increasing the volume decreases the pressure

causes a shift to the left in the above reaction as this produces the larger number of molecules


analyse the changes made to the equilibrium system shown in the diagram at the left

2NO2 N2O4

TEMPERATURE• increasing the temperature increases the reaction rate• this is due to an increase in the fraction of collisions

in which the total kinetic energy of reacting particles is at least equal to the activation energy

• increasing the temperature favours the endothermic reaction

• decreasing the temperature favours the exothermic reaction

TEMPERATURE

Mass-Gas Volume

1. Calculate the mass of 18.25L of ammonia gas at 25.0oC and 100.0kPa.

2. At 100.0 kPa and 25.0oC, how many litres of carbon dioxide gas will be produced when 75.0g of calcium carbonate is decomposed into calcium oxide?

Mass-Gas Volume

3. Solid lithium hydroxide has been used in space craft to remove carbon dioxide from air. Lithium carbonate and water are formed.

What mass of lithium hydroxide would be needed to remove 250.0 L of carbon dioxide at 100.0 kPa and 25.0oC?

Acid RainSulfur dioxide

natural and man-made sources reactions to produce it

sulfur compounds in coal S(s) + O2(g) SO2(g)

smelting of sulfide ores ZnS(s) + O2(g) Zn(s) + SO2(g)

oxidation of H2S – decay and industrial

2H2S + 3O2(g) 2H2O(l) + 2SO2(g)

reactions to produce acidic solutions effects

living things and environmentcorrosion metals, limestone buildings (CaCO3)

Acid RainNitrogen oxides NOx (NO, NO2)


high temperature engines

N2 + O2 2NO (neutral oxide)

2NO + O2 2NO2 (acidic oxide)

reactions to produce acidic solutions 2NO2 + H2O HNO2 + HNO3

effectsliving things and environment

Production of Ozone

photodecomposition

NO2 NO + O

ozone formation

O + O2 O3

regeneration of nitrogen dioxide

O3 + NO O2 + NO2

Ozone is a secondary pollutant in the troposphere

Photochemical Smog

What is acid rain?

More appropriate term is“acidic deposition”-snow, fog, sleet, haze, dry deposition

What is Acid Rain?Pure water: pH 7Natural rain: pH 5-6Acid rain pH < 5

Acid Rain

1730s – originated at height of Industrial Revolution

1872 – Robert Smith, an English chemist, coined the phrase “acid rain”

1950s – lake acidification first described

1960s – became more noticeable and subsequently became worse in rural areas

tall chimneys on factories allow wind to transport pollutants far away from sources of production

Acid Rain

Out west, in the Rocky Mountains scientists are finding that power plant emissions are saturating high-elevation watersheds in Colorado with acid-causing nitrogen. Evergreen forests are losing their needles and tree health is declining throughout the forest range

Acid Rain

Acid rain damage Blue Ridge Mountains North Carolina

Acid Rain

Acid rain damage on monumentCaCO3 (s) + H2SO4 (aq) CaSO4 (aq) + CO2 (g) + H2O (l)

Acid Rain

Tasmania - Queenstown emerged as a boomtown of the 1890s when gold and minerals were discovered at Mount Lyell. The strange but arresting 'moonscape' that surrounds the town was caused by acid-rain during the mining era.

Acid Rain1984 – reported almost half of Germany’s Black

Forest damaged by acid rain

Other areas

- acidification of lakes in Scandinavia

- Taj Mahal and many statues in Europe increased deterioration due to acid rain

- substantial problem in Europe, China and Russia as burn higher S-containing coal to generate electricity

- aluminium

Acid RainSulfur dioxide


sulfur compounds in coal smelting of sulfide ores oxidation of H2S – decay and industrial

reactions to produce acidic solutions effects

living things and environmentcorrosion metals, limestone buildings (CaCO3)

Acid Rain

SO2 Pollution

Killer smogs of London – 1952, 1956, 1957, 1962

Acid Rain

Acid RainNitrogen oxides NOx (NO, NO2)


high temperature engines

N2 + O2 2NO (neutral oxide)

2NO + O2 2NO2 (acidic oxide)

reactions to produce acidic solutions 2NO2 + H2O HNO2 + HNO3

effectsliving things and environment

http://www.csiro.au/promos/ozadvances/series14acidrainmovb.htm

Changes in Sulfate across the USA

http://nadp.sws.uiuc.edu/data/amaps/so4/amaps.html

NADP Annual Maps

Production of Ozone

photodecomposition

NO2 NO + O

ozone formation

O + O2 O3

O3 + NO O2 + NO2

Ozone is a secondary pollutant in the troposphere

Photochemical Smog

TABLE 4.4 TEXT P. 124

Common Acids

Acetic acid

Phosphoric acid

Sulfuric acid

3. Acids occur in many foods, drinks and even within our stomachs

Naturally occurring acetic/ethanoic (vinegar) citric/2-hydroxypropane-

1,2,3-tricarboxylic acid (citrus fruit)

hydrochloric (stomach)

3. Acids occur in many foods, drinks and even within our stomachs

Acids

Aspirinacetylsalicylic acid

Amino acids

Acids

Manufactured/Synthetic sulfuric acid

car batteries, fertiliser (NH3)2SO4, detergents, catalyst production ethanol and esters

nitric acid fertilisers , explosives

BasesNaturally Occurring ammonia NH3

also manufactured to produce fertilisers (Haber process)

metal oxides – Fe2O3, CuO

carbonates CO32- (Na2CO3, CaCO3)

Manufactured/Synthetic sodium hydroxide – soap, Draino (NaOH) calcium oxide, calcium hydroxide

Bases

Acids & Bases

Text p. 131-133

Self-Ionisaton/Autolysis of H2O in a sample of pure water a very small amount

of the molecules react with each other this is called the self-ionisation of water.

H2O(l) + H2O(l) H3O+(aq) + OH–

(aq)

at 25oC [OH-] = [H3O+] = 1.0x10-7 mol/L

KW = [OH-] x [H3O+] = 1.0 x 10-14

in any aqueous solution the [OH-] and [H3O+] are interdependent but KW is constant

aqueous solutions are neutral, acidic, basic

Using Kw

If the hydroxide ion concentration of a sodium

hydroxide solution is 1.5 x 10-3mol/L at 25oC,

what is the hydrogen ion concentration?

Using Kw

At 25oC an aqueous solution has a hydrogen ion concentration of 2.4 x 10-3mol/L.

What is the hydroxide ion concentration in this solution?

The pH Scale

proposed in 1909 by Danish scientist Soren Sorensen

pH means: power of the Hydrogen ion pH = -log[H+] the negative logarithm of the

hydrogen ion concentration neutral, acidic, basic solutions to obtain the [H+] given the pH

[H+] = 10-pH

pH – A measure of acidity Nitric acid (HNO3) is used in the production of

fertilizer, dyes, drugs, and explosives. Calculate the pH of a HNO3 solution having a hydrogen ion concentration of 0.76 M.

The pH of a brand of orange juice is 3.33. Calculate the H+ ion concentration.

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

pH

Show that a change in pH from 4.75 to 3.75

corresponds to a tenfold increase in hydrogen ion

concentration.

Problems with B-L Theory

The theory works very nicely in all protic solvents

but fails to explain acid-base behavior in aprotic solvents and some non-solvent situations.

A more general concept of acids and bases was proposed by G.N. Lewis at about the same time Bronsted-Lowry theory was proposed.

4.2 Bronsted-Lowry theory

3.2.2 plan and perform a first-hand investigation to measure the pH of identical

concentrations of strong and weak acids

HA

H+

A-

100% ionisation of HA

Would the solution conduct (be an electrolyte)?

Strong Acid

HA H+ + A-

unionised acid molecule

hydrogen ion

anion from acid

Strong Acid

For example

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

(aq)

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

(aq)

HNO3(aq) H+(aq) + NO3

-(aq)

OR

HNO3(l) + H2O(l) H3O+(aq) + NO3

-(aq)

HA

H+

A-

Partial ionisation of HA

Would the solution be conductive?

Weak Acid

At any one time, only a fraction of the molecules are ionised

HA H+ + A-

Weak Acid Note the use of the double arrowdouble arrow The unionised acid molecules are in

EQUILIBRIUM with the ionised hydrogen ion and anion from the acid

CH3COOH(aq) H+(aq) + CH3COO-(aq)

OR

CH3COOH(l) + H2O(l) H3O+(aq) + CH3COO-

(aq)

HA H+ + A-

Acids and Bases

STRONG ACIDS HCl, HBr, HI, H2SO4, HNO3, HClO3, HClO4

WEAK ACIDS organic acids, and H2SO3, HNO2, H3PO4, H2CO3

STRONG BASES ionise completely in water to produce OH- ions

LiOH, Na2O, KOH, Ba(OH)2 ALKALIS – strong soluble bases

WEAK BASES NH3, CO32-, HCO3

-

Acids and Bases

Weak bases like NH3 react with water to produce hydroxide ions

This also forms an EQUILIBRIUM

NH3(g) + H2O(l) NH4+

(aq) + OH-(aq)

ammonium

ion

Acids

If the degree of ionisation of a weak acid is known then the pH of the solution can be determined.

e.g. If a solution of 0.037M hydrofluoric acid, HF, is 12.9% ionised what is the pH of the solution? HF H+ + F-

[H+] = 12.9/100 x 0.037 M = 0.00477 M

pH = -log(0.00477) = 2.32

Finish worksheet on p.100 in SSB

Acids and Bases

strength weak – limited

ionisation forming an equilibrium system

strong – complete (100%) ionisation

concentration dilute concentrated

water (solvent) solute

concentrated, Minitial

diluted, Mfinal

adding water lowers the solute concentration

moles of solute remain constant

Vinitial

Vfinal

molesinitial = molesfinal

Mfinal x Vfinal = Minitial x Vinitial

Dilution

Acid Concentration

dilute solution of a strong acid low number of moles of acid molecules per L of

solution all acid molecules completely ionised

concentrated solution of weak acid higher number of moles of acid molecules per L

of solution acid molecules only partially ionised

Monoprotic Acid

contains only one ionisable hydrogen HCl, HNO3, CH3COOH

Diprotic Acid

contains 2 ionisable hydrogens 2-step ionisation

First ionisation

H2SO4 H+ + HSO4- (complete)

Second ionisation

HSO4- H+ + SO4

2- (partial)

Triprotic Acid

contains 3 ionisable hydrogens phosphoric acid H3PO4

First ionisation

H3PO4 H+ + H2PO4- (partial)

Second ionisation

H2PO4- H+ + HPO4

2-

Third ionisation

HPO42- H+ + PO4

3-

HA

H+

A-

Strong Acid

unionised acid molecule

hydrogen ion

anion from acid

HA

H+

A-

Weak Acid

Gas-neutralisation Problems

1. At 25oC and 100 kPa, 2.5 litres of hydrogen chloride gas is bubbled through a sodium hydroxide solution. If the solution is 0.50M what volume would be needed to completely neutralise the gas? balanced equation moles of HCl moles of sodium hydroxide needed volume of solution


2. 3.0 litres of carbon dioxide is bubbled through 200.0 mL of 0.15 M calcium hydroxide solution at 25oC and 100 kPa. What mass of calcium carbonate precipitate will form?

3. If 350.0 mL of a solution of potassium hydroxide completely neutralises 5.0 L of sulfur dioxide gas at 25oC and 100 kPa, what is the concentration of the solution?


4. What volume of 0.25M barium hydroxide solution would completely neutralise 10.0 L of hydrogen chloride gas at 25oC and 100 kPa?

5. 500.0mL of hydrogen chloride gas at 25oC and 100kPa is bubbled through 800.0mL of distilled water. Assuming all the hydrogen chloride reacts, what is the pH of the solution?

pH of Solutions

1. 175.0mL of a 0.085M solution of sodium hydroxide is mixed with 150.0mL of a 0.15M solution of hydrochloric acid. Determine the pH of the final solution.

2. 250.0mL of a 0.15M solution of potassium hydroxide is mixed with 275.0mL of a 0.085M solution of nitric acid. Determine the pH of the final solution.

pH of Solutions

3. 50.0mL of a 0.050M solution of barium hydroxide is mixed with 75.0mL of a 0.100M solution of hydrochloric acid. Determine the pH of the final solution.

pH Meter tests the voltage of the

electrolyte converts the voltage to

pH very cheap, accurate must be calibrated with

buffer solutions non-destructive testing –

does not change solution being tested

pH of Acid Solutions

Acid: Molarity: pH (±0.1): [H+]:

HCl 0.1 1.0 0.10

C6H8O7 0.1 1.5 0.032

CH3COOH 0.1 2.9 0.00013

3.3.2 plan and perform a first-hand investigation to measure the pH of identical concentrations of strong

and weak acids

Molecular Structure and Acid Strength

the strength of an acid depends on its tendency to ionize.

for general acids of the type H–X:1. The stronger the bond, the weaker the acid.

2. The more polar the bond, the stronger the acid.for the hydrohalic acids, bond strength plays the key role giving: HF < HCl < HBr < HI

Molecular Structure and Acid Strength

The electrostatic potential maps show all the hydrohalic acids are polar. The variation in polarity is less significant than the bond strength which decreases from 567 kJ/mol for HF to 299 kJ/mol for HI.

Acids

1. Write equations to show the 2-step ionisation in water of the weak sulfurous acid, H2SO3

Acids as Food Additives

acidulant – gives a sharp/tart taste to food antimicrobials – lowers pH to inhibit growth

of bacteria, yeasts or molds antioxidants – slows oxidation which causes

spoilage e.g. fats and oils inhibit/block enzymes that continue natural

ripening after harvest – causes browning

3.3.6 Identify data, gather and process information from secondary sources to identify examples of naturally

occurring acids and bases and their chemical composition

Name: Formula: pH in natural form: Naturally found in:

Acetic CH3COOH 3-5 Vinegar, grapes, wine

Ascorbic C6H8O6 2-3 Fruit (esp. citrus), vegetables

Carbonic H2CO3 2-3 Acid rain

Citric C6H8O7 2-3 Citrus fruits

Formic CHOOH 3-5 Poison of stinging ants/insects

Hydrochloric HCl 0.1-2 Gastric juice in stomach

Ammonia NH3 9-11 Volcanic gases, decomposed plant/animal matter

Caffeine C8H10N4O2 8-10 Coffee beans, cola nuts

Nicotine C8H14N2 8-10 Tobacco leaves

Limestone CaCO3 8-10 Limestone

Acids and Bases

OPERATIONAL

DEFINITION based on observed

properties what do they do?

Acids • taste sour• change the colour of indicators e.g. blue

litmus to red• neutralise bases and basic oxides• some are corrosive• react with active metals such as zinc,

magnesium giving off hydrogen gas• aqueous solutions of acids conduct electricity –

they are ELECTROLYTES

Bases

• taste bitter• change the colour of indicators e.g turn red

litmus blue• neutralise acids and acidic oxides• some are corrosive• solutions of soluble bases in water are

electrolytes

Acids and Bases

CONCEPTUAL

DEFINITIONS a theoretical framework to

explain observed properties

more likely to change as our knowledge increases

4.1 Outline the historical development of ideas about acids

1778 – Antoine Lavoisier• oxides of P and S combined with water to

produce acidic solutions

S + O2 SO2 + H2O H2SO3

• oxygen is responsible for acidity• named oxygen from Greek “oxys” =sharp/sour

and “genes” = born/form (acid former)


1811 – Sir Humphrey Davy• acids contain the element hydrogen - so

hydrogen is responsible for acidity


1887 Svante Arrhenius acidic and basic solutions conduct

electricity so electrolytes (ions) acids react with metals to produce

hydrogen so ions involved developed ionic theory of electrolytes for

which he received a Nobel Prize in 1903


1887 Svante Arrhenius acids are substances that release H+

in aqueous solutione.g. HCl(aq) H+

(aq) + Cl-(aq)

H2SO4(aq) 2H+(aq) + SO4

2-(aq)

bases are substances that release OH- ions in aqueous solution e.g. NaOH(aq) Na+

(aq) + OH-(aq)

Ba(OH)2(aq) Ba2+(aq) + 2OH-

(aq)


Neutralisation

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

H+(aq) + OH-

(aq) H2O(l)

Problems with Arrhenius Theory

the role of the solvent? – is an acid an acid in any solvent

all salts should produce neutral solutions – neither acidic nor basic

the need for hydroxide as the base

e.g. NH4OH as the base and not NH3

The Hydrogen Ion

a proton with a 1+ charge and extremely small mass/volume high charge density and intense electric field

too reactive to exist independently in a very polar solvent like water

the hydronium ion, H3O+

Subsection 4Because of the prevalence and importance of

acids, they have been used and studied for hundreds of years. Over time, the definitions

of acid and base have been refined

4.2 Outline the Bronsted-Lowry theory of acids and bases

in 1923 a more general theory of acid-base behaviour was independently proposed by Danish chemist J Bronsted and English chemist T Lowry

Bronsted-Lowry theory defines: an acid as a species from which a proton

can be removed (acids are proton donors) a base as a species that can remove a

proton from an acid (bases are proton acceptors)


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

(aq)

+ -

hydronium ion

4.2 Bronsted-Lowry Theory

CH3COOH(l) + H2O(l) H3O+(aq) + CH3COO-

(aq)

NH3(g) + H2O(l) NH4+

(aq) + OH-(aq)

an acid-base reaction is one in which a proton is transferred from an acid to a base

a proton-transfer reaction


The role of the solvent:

Hydrogen chloride in liquid ammonia


a broader definition which shows the complementary nature of acids and bases

shows the role of the solvent which can be a proton acceptor or proton donor

includes more species that Arrhenius Theory - molecules and ions

acid must contain hydrogen to have a proton removed

4 Bronsted-Lowry theory

each B-L reaction involves two acid-base pairs called CONJUGATE PAIRS - two species that differ by a proton

conjugate means “coupled or joinedcoupled or joined”

4 Bronsted-Lowry theory

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

(aq)

+ -

acid conjugate base

base conjugate acid

Acids and Bases

1. What is the pH of a solution made by diluting 2.50mL of 6.0M HCl to 500.0mL?

2. What is the pH of a 0.035M solution of Ba(OH)2 ?

3. The pH of a HCl solution is 1.25. If 200.0mL of this solution is diluted to 500.0mL, what is the pH of this new solution?

4 AMPHIPROTIC SPECIES

Molecules or ions that can accept OROR donate a proton

Act as acids or act as bases

e.g. H2O(l) + H2O(l) H3O+(aq) + OH–

(aq)


hydrogen carbonate ion and a strong acid and base

HCO3-(aq) + OH-

(aq) CO32-

(aq) + H2O(l)

acid base

HCO3-(aq) + H3O+

(aq) H2CO3(aq) + H2O(l)

base acid

H2CO3(aq) CO2(g) + H2O(l)


hydrogen carbonate ion and a weak acid/base

HCO3- + H2O CO3

2- + H3O+

acid base

HCO3- + H2O H2CO3 + OH-

base acid


The hydrogen sulfate ion is amphiprotic.

a) Write balanced equations to show this behaviour. (use H3O+ and OH-)

b) A solution of sodium hydrogen sulfate in water turns blue Litmus red. Use an equation to explain this behaviour.

4.2.8 & 4.3.3 TITRATIONS

Validity & ReliabilityTITRATION - VALIDITY appropriate reaction – acid and base primary std or standardised secondary std appropriate indicator for type of titration accurate measuring instruments – volumetric

glassware – volumetric pipette, burette correct washing procedures and use e.g. method of

operating the pipette and burette

RELIABILITY 3 or more trials – reproducible average titre

4.2.4 Identify a range of salts which form acidic, basic and neutral solutions and explain their

acidic, neutral or basic nature

0.1M Salt SolutionpH Universal

Indicator pH Probe

NaCl 6-7 6.5

NH4Cl 4-5 4.5

NaCH3COO 8-9 9.5

NaNO3 6-7 7.2

Na2CO3 10-11 9.9

TEXT p. 154 TABLE 5.4

Summary of salts formed from different types of acids & bases

4.2.4 Identify a range of salts which form acidic, basic and neutral solutions and explain their

acidic, neutral or basic nature

Indicators

Phenolphthalein is a commonly used indicator for titrations, and is a weak acid.

the weak acid is colourless and its ion is bright pink. Adding extra hydrogen ions shifts the position of

equilibrium to the left, and turns the indicator colourless.

Adding hydroxide ions removes the hydrogen ions from the equilibrium which shifts to the right to replace them - turning the indicator pink.

Strong Acid with Strong Base

HCl + NaOH NaCl + H2O

Strong Acid with Strong Base

pH starts low

Equivalence

point pH = 7

pH finishes high

HCl + NaOH NaCl + H2O

8.3-10

3.1-4.4

Weak Acid with Strong Base

Volume of base added (mL)

CH3COOH + NaOH NaCH3COO + H2O

phenolphth

methyl orange

Weak Acid with Strong Base

Volume of base added (mL)

pH starts higher

Equivalence point

pH finishes high

CH3COOH + NaOH NaCH3COO + H2O

phenolphth

methyl orange

Weak Base with Strong Acid

Volume of acid added (mL)

NH3 + HCl NH4Cl + H2O

phenolphth

methyl orange

Weak Base with Strong Acid

Volume of acid added (mL)

pH starts moderately high

Equivalence point

pH finishes low

NH3 + HCl NH4Cl + H2O

phenolphth

methyl orange

Weak Acid with Weak Base

pH

4.2.4 Identify a range of salts which form acidic, basic and neutral solutions and explain their acidic, neutral or basic

nature

For each of the salts below,

i) give the formula

ii) state the acid and base that produced the salt

iii) state whether you would expect 0.1M aqueous solutions to be neutral, acidic or basic

iv) explain why, giving appropriate equations where necessary

1. barium nitrate 2. sodium methanoate

3. sodium carbonate 4. ammonium nitrate

5. sodium sulfite 6. potassium bromide

4.2.7 Neutralisation

a proton transfer reaction exothermic reaction for example

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(aq)

H = -56.1 kJmol-1

4.3.5 Neutralisation – safety measure and minimise damage in chemical spills

Factors to consider type of acid or base – weak or strong,

concentrated or dilute volume – few mL on laboratory bench or

much larger volume in more public place

4.3.5 Neutralisation – safety measure and minimise damage in chemical spills

weak acids and bases are safer to use neutralise acids

Na2CO3 – solid, cheap, easy to use, excess does not present problems of disposal

neutralise acids and alkalis NaHCO3 – amphiprotic

HCO3- + OH- CO3

2- + H2O

HCO3- + H+ H2CO3 CO2 + H2O

Booklet p. 142-144

Sources of H+ in the Body

Ketone bodies

Acetone

Betahydroxybutyric acid

Acetoacetate

(CH3COCH2COOH)

4.2.9 BUFFERS

A buffer is a solution that resists a change in its pH when acid (H3O+) or base (OH-) is added to it. based on chemical equilibrium

A solution of a weak acid and its conjugate base OR a weak base and its conjugate acid nearly all biochemical reactions are influenced

by the pH of their fluid environment maintaining the pH of blood

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HSC CHEMISTRY CORE TOPIC 2