HSC CHEMISTRY CORE TOPIC 2

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HSC CHEMISTRY CORE TOPIC 2

• THE ACIDIC ENVIRONMENT

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INDICATORS
LE CHATELIERS PRINCIPLE
TITRATION
pH
VOLUMETRIC ANALYSIS
NEUTRALISATION
CHEMICAL EQUILIBRIUM
THE ACIDIC ENVIRONMENT
HISTORY Lavoisier Davy Arrhenius
CARBOXYLIC ACIDS
ACIDIC OXIDES
BRONSTED LOWRY THEORY
ESTERIFICATION
ACID RAIN
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Subsection 1 Indicators were identified with the
observation that the colour of some flowers
depends on soil composition
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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

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INDICATORS
INDICATOR pH RANGE COLOUR 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
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INDICATORS
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INDICATORS

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

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INDICATORS

• A solution gives the following colours in each
  indicator. Deduce the approximate pH
• Indicator Colour
• methyl orange yellow
• bromothymol blue yellow
• phenolphthalein colourless

gt4.4 lt6.0 lt8.0
pH of solution is 4.4 6.0
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INDICATORS

• A solution gives the following colours in each
  indicator. Deduce the approximate pH
• Indicator Colour
• bromothymol blue blue
• thymolphthalein colourless
• phenol red red

gt 7.6 lt 9.5 gt 8.4
pH of solution is 8.4 9.5
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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

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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.
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• 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

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Oxides of Period 3

• NaOH(s) ?
• MgO(s) H2O(aq) g
• S(s) O2(g) g
• SO2(g) H2O(aq) g
• P4(s) O2(g) g P2O5(s)
• P2O5(s) H2O(aq) g

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Oxides of Period 3

• NaOH(s) g Na(aq) OH-(aq)
• MgO H2O g Mg(OH)2
• S O2 g SO2
• SO2 H2O g H2SO3 (sulfurous acid)
• 2H2SO3 O2 g 2H2SO4 (sulfuric acid)
• P4 5O2 g 2P2O5
• P2O5 3H2O g 2H3PO4 (phosphoric acid)
• Cl2O7 H2O g 2HClO4 (perchloric acid)

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Oxides of Period 3
Oxide/ Hydroxide Na2O (NaOH) MgO Al2O3 SiO2 P2O5 SO2 Cl2O7
Bonding
Acid/base property amphoteric acidic
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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)

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Acid/Base Properties of Oxides

• 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 salts
• CO2(g) CaO(s) ? CaCO3(s)
• 2NaOH(l) SiO2(s) ? Na2SiO3(s) H2O(l)

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Acid/Base Properties of Oxides

• AMPHOTERIC OXIDES
• react with acids and bases
• Al2O3, BeO, ZnO, PbO, SnO
• Al2O3(s) 6HCl(aq) g 2AlCl3(aq) 3H2O(l)
• Al2O3(s) 2NaOH(aq) g2NaAlO2(aq) H2O(l)
• NEUTRAL OXIDES
• do not react with acids or bases
• CO, N2O, NO

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Group A Oxides
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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

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A Saturated Solution of NaCl
Dissolution NaCl(s)g Na(aq) Cl-(aq)
NaCl
System at 25oC
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A Saturated Solution of NaCl
Radioactive 24NaCl is introduced into the beaker
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A Saturated Solution of NaCl
CHEMICAL EQUILIBRIUM
Dissolution NaCl(s) g Na(aq)
Cl_(aq) Recrystallisation NaCl(s) f Na(aq)
Cl_(aq)
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Chemical Equilibrium
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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

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CHEMICAL EQUILIBRIUM
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Le CHATELIERS 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

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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

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CONCENTRATION

• equilibria with solids and pure liquids
• e.g. C(s) H2O(g) n CO(g)
  H2(g)

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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

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Partial Pressure of a Gas
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PRESSURE

• Adding an inert gas
• does not change the partial pressures of any of
  the other gases
• no effect on equilibrium

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VOLUME CHANGES

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

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VOLUME CHANGES

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

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

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

• 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

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CHEMICAL EQUILIBRIUM
2NO2 n N2O4

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

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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

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TEMPERATURE
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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?

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Mass-Gas Volume

• 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?

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Acid Rain

• Sulfur 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 environment
• corrosion metals, limestone buildings (CaCO3)

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Acid Rain

• Nitrogen oxides NOx (NO, NO2)
• natural and man-made sources
• reactions to produce it
• high temperature engines
• N2 O2 g 2NO (neutral oxide)
• 2NO O2 g 2NO2 (acidic oxide)
• reactions to produce acidic solutions
• 2NO2 H2O g HNO2 HNO3
• effects
• living things and environment

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Production of Ozone
photodecomposition NO2 g NO O ozone formation O
O2 g O3
regeneration of nitrogen dioxide O3 NO g O2
NO2
Ozone is a secondary pollutant in the troposphere
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Photochemical Smog
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What is acid rain?
More appropriate term is acidic
deposition -snow, fog, sleet, haze, dry
deposition
What is Acid Rain? Pure water pH 7 Natural rain
pH 5-6 Acid rain pH lt 5
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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

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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
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Acid Rain
Acid rain damage Blue Ridge Mountains North
Carolina
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Acid Rain
Acid rain damage on monument CaCO3 (s) H2SO4
(aq) ? CaSO4 (aq) CO2 (g) H2O (l)
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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.
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Acid Rain

• 1984 reported almost half of Germanys 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

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Acid Rain

• Sulfur dioxide
• natural and man-made sources
• reactions to produce it
• sulfur compounds in coal
• smelting of sulfide ores
• oxidation of H2S decay and industrial
• reactions to produce acidic solutions
• effects
• living things and environment
• corrosion metals, limestone buildings (CaCO3)

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Acid Rain
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SO2 Pollution
Killer smogs of London 1952, 1956, 1957, 1962
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Acid Rain
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Acid Rain

• Nitrogen oxides NOx (NO, NO2)
• natural and man-made sources
• reactions to produce it
• high temperature engines
• N2 O2 g 2NO (neutral oxide)
• 2NO O2 g 2NO2 (acidic oxide)
• reactions to produce acidic solutions
• 2NO2 H2O g HNO2 HNO3
• effects
• living things and environment

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http//www.csiro.au/promos/ozadvances/series14acid
rainmovb.htm
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Changes in Sulfate across the USA
http//nadp.sws.uiuc.edu/data/amaps/so4/amaps.html
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• NADP Annual Maps

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Production of Ozone
photodecomposition NO2 g NO O ozone formation O
O2 g O3
O3 NO g O2 NO2
Ozone is a secondary pollutant in the troposphere
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Photochemical Smog
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Table 4.4 Text p. 124
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Common Acids
Acetic acid
Phosphoric acid
Sulfuric acid
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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)

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

Aspirin acetylsalicylic acid
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Amino acids
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Acids

• Manufactured/Synthetic
• sulfuric acid
• car batteries, fertiliser (NH3)2SO4, detergents,
  catalyst production ethanol and esters
• nitric acid
• fertilisers , explosives

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Bases

• Naturally 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

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Bases
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Acids Bases
Text p. 131-133
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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) n 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

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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?

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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?

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The pH Scale

• proposed in 1909 by Danish scientist Soren
  Sorensen
• pH means power of the Hydrogen ion
• pH -logH the negative logarithm of the
  hydrogen ion concentration
• neutral, acidic, basic solutions
• to obtain the H given the pH
• H 10-pH

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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 107 M. What is the pH of the blood?

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pH

• Show that a change in pH from 4.75 to 3.75
• corresponds to a tenfold increase in hydrogen ion
• concentration.

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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.

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4.2 Bronsted-Lowry theory
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3.2.2 plan and perform a first-hand
investigation to measure the pH of identical
concentrations of strong and weak acids
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Strong Acid
unionised acid molecule
hydrogen ion
Would the solution conduct (be an electrolyte)?
anion from acid
100 ionisation of HA
HA g H A-
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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)

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Weak Acid
At any one time, only a fraction of the molecules
are ionised
HA
H
A-
Would the solution be conductive?
Partial ionisation of HA
HA ? H A-
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Weak Acid

• Note the use of the double arrow
• The unionised acid molecules are in EQUILIBRIUM
  with the ionised hydrogen ion and anion from the
  acid
• CH3COOH(aq) n H(aq) CH3COO-(aq)
• OR
• CH3COOH(l) H2O(l) nH3O(aq) CH3COO-(aq)

HA ? H A-
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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-

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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

87
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 n 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

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Acids and Bases

• strength
• weak limited ionisation forming an equilibrium
  system
• strong complete (100) ionisation
• concentration
• dilute
• concentrated

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Dilution
water (solvent)
solute
moles of solute remain constant
diluted, Mfinal
Vfinal
molesinitial molesfinal
Vinitial
concentrated, Minitial
adding water lowers the solute concentration
Mfinal x Vfinal Minitial x Vinitial
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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

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Monoprotic Acid

• contains only one ionisable hydrogen
• HCl, HNO3, CH3COOH

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Diprotic Acid

• contains 2 ionisable hydrogens
• 2-step ionisation
• First ionisation
• H2SO4 ? H HSO4- (complete)
• Second ionisation
• HSO4- n H SO42- (partial)

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Triprotic Acid

• contains 3 ionisable hydrogens
• phosphoric acid H3PO4
• First ionisation
• H3PO4 n H H2PO4- (partial)
• Second ionisation
• H2PO4- n H HPO42-
• Third ionisation
• HPO42- n H PO43-

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Strong Acid
unionised acid molecule
hydrogen ion
anion from acid
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Weak Acid
HA
H
A-
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Gas-neutralisation Problems

• 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

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Gas-neutralisation Problems

• 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?
• 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?

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Gas-neutralisation Problems

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

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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.

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pH of Solutions

• 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.

101
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

102
pH of Acid Solutions
3.3.2 plan and perform a first-hand investigation
to measure the pH of identical concentrations of
strong and weak acids
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
103
Molecular Structure and Acid Strength

• the strength of an acid depends on its tendency
  to ionize.
• for general acids of the type HX
• The stronger the bond, the weaker the acid.
• The more polar the bond, the stronger the acid.
• for the hydrohalic acids, bond strength plays the
  key role giving HF lt HCl lt HBr lt HI

104
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.

105
Acids

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

106
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

107
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
108
Acids and Bases

• OPERATIONAL
• DEFINITION
• based on observed properties
• what do they do?

109
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

110
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

111
Acids and Bases

• CONCEPTUAL
• DEFINITIONS
• a theoretical framework to explain observed
  properties
• more likely to change as our knowledge increases

112
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)

113
4.1 Outline the historical development of ideas
about acids

• 1811 Sir Humphrey Davy
• acids contain the element hydrogen - so hydrogen
  is responsible for acidity

114
4.1 Outline the historical development of ideas
about acids

• 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

115
4.1 Outline the historical development of ideas
about acids

• 1887 Svante Arrhenius
• acids are substances that release H
• in aqueous solution
• e.g. HCl(aq) g H(aq) Cl-(aq)
• H2SO4(aq) g 2H(aq) SO42-(aq)
• bases are substances that release OH- ions in
  aqueous solution
• e.g. NaOH(aq) g Na(aq) OH-(aq)
• Ba(OH)2(aq) g Ba2(aq) 2OH-(aq)

116
4.1 Outline the historical development of ideas
about acids

• Neutralisation
• HCl(aq) NaOH(aq) ? NaCl(aq) H2O(l)
• H(aq) OH-(aq) ? H2O(l)

117
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

118
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

119
Subsection 4 Because 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
120
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)

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4.2 Bronsted-Lowry theory
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4.2 Bronsted-Lowry Theory

• CH3COOH(l) H2O(l) nH3O(aq) CH3COO-(aq)
• NH3(g) H2O(l) n 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

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4.2 Bronsted-Lowry theory
The role of the solvent Hydrogen chloride in
liquid ammonia
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4.2 Bronsted-Lowry theory

• 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

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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 joined

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4 Bronsted-Lowry theory
conjugate acid
base
acid
conjugate base
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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?

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4 AMPHIPROTIC SPECIES

• Molecules or ions that can accept OR donate a
  proton
• Act as acids or act as bases
• e.g. H2O(l) H2O(l) n H3O(aq) OH(aq)

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4 AMPHIPROTIC SPECIES

• 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) D CO2(g) H2O(l)

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4 AMPHIPROTIC SPECIES

• hydrogen carbonate ion and a weak acid/base
• HCO3- H2O D CO32- H3O
• acid base
• HCO3- H2O D H2CO3 OH-
• base acid

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4 AMPHIPROTIC SPECIES

• The hydrogen sulfate ion is amphiprotic.
• Write balanced equations to show this behaviour.
  (use H3O and OH-)
• A solution of sodium hydrogen sulfate in water
  turns blue Litmus red. Use an equation to
  explain this behaviour.

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4.2.8 4.3.3 TITRATIONS
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Validity Reliability

• TITRATION - 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

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

• TEXT p. 154 TABLE 5.4
• Summary of salts formed from different types of
  acids bases

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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.

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Strong Acid with Strong Base
HCl NaOH g NaCl H2O
139
Strong Acid with Strong Base
pH finishes high
8.3-10
Equivalence point pH 7
pH starts low
3.1-4.4
HCl NaOH g NaCl H2O
140
Weak Acid with Strong Base
phenolphth
methyl orange
CH3COOH NaOH g NaCH3COO H2O
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Weak Acid with Strong Base
pH finishes high
Equivalence point
phenolphth
pH starts higher
methyl orange
CH3COOH NaOH g NaCH3COO H2O
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Weak Base with Strong Acid
phenolphth
methyl orange
NH3 HCl g NH4Cl H2O
143
Weak Base with Strong Acid
pH starts moderately high
Equivalence point
phenolphth
methyl orange
pH finishes low
NH3 HCl g NH4Cl H2O
144
Weak Acid with Weak Base
pH
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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,
• give the formula
• state the acid and base that produced the salt
• state whether you would expect 0.1M aqueous
  solutions to be neutral, acidic or basic
• 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

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4.2.7 Neutralisation

• a proton transfer reaction
• exothermic reaction
• for example
• HCl(aq) NaOH(aq) g NaCl(aq) H2O(aq)
• DH -56.1 kJmol-1

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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

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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- ? CO32- H2O
• HCO3- H ? H2CO3 ? CO2 H2O
• Booklet p. 142-144

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Sources of H in the Body
Ketone bodies Acetone Betahydroxybutyric
acid Acetoacetate (CH3COCH2COOH)
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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