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Doc
Brown's Chemistry Clinic
My unofficial
support for Salters
A2
Advanced Chemistry
Salters A2 Chemistry - 'exam
bashing' thoughts for
Unit O "The
Oceans" - part of module
2854
My revision
index * extra O stuff *
O unit
map-learning
objectives *
My Salters AS homepage *
My Salters A2 homepage *
Email
PLEASE REMEMBER, THESE ARE NOT 'STAND
ALONE' NOTES, and were designed for my classes for use alongside the Salters
resources - Chemical Ideas, Chemical Storylines, Practical
Activities-Investigations and the AS-A2 Revision guides all published by
Heinemann Secondary Series, to reduce the reading workload and offer a study
strategy. From your teacher (not
me!), its handy to have the answers to the Chemical Ideas, Storylines
Assignments and Activities Questions side by side with the texts and these
strategy pages. You haven't time to redo the Q's but a quick read of the Q's and
connecting with the official answers is valuable revision - there is too much
hit and miss revision from doing past papers in my opinion.
Chemical Storylines
O1 The Edge of
the land
-
Products from the sea, examples of food, dye materials,
salts and elements such as bromine, chlorine and magnesium.
-
Reminder of bromine from
seawater and chlorine from NaCl (unit M),
-
99% of dissolved materials are ionic,
origin of
the ‘saltiness’ of sea-water (Fig 5, part of the Earth’s complex
geochemical cycle): rainwater leachings from rocks (but not source of Br, Cl, S based
ions).
-
Volcanic gases from mid-ocean ridges, leaching from shattered cooled lava in
sea-water, hydrothermal vents in hot cracked crust.
-
Salt content fairly constant around
the seas/oceans but higher in hot areas where rate of evaporation high.
-
Salt content lower
where there is a source of fresh water eg estuaries, high rainfall or icebergs.
-
NaCl
important component of our diet, NaCl from evaporated sea-water, often via fractional
crystallisation (eg CaSO4, crystallise first due to lower solubility, this can
be ‘skimmed’ off, also allows NaCl to be removed before bitter-tasting Mg salts
crystallise.
-
Sea organisms are potential sources of organic chemicals eg lead
(NOT Pb!) molecules
for pharmaceutical industry.
-
Pollution problems eg heavy metal ions from industry, sewage
and excess fertiliser from agriculture.
Chemical Ideas 5.1 Ions in solids and
solutions (revision)
-
Cations and anions, NaCl lattice, double salts and modified lattice,
hydrated salts, solutions of ions, ionic equations (with or without
precipitation) eg AgCl(s) formation or neutralisation,
recognising which are spectator ions, and writing ionic equations for
salt crystallisation.
-
Describing the process of
hydration and the influence of size and charge on the number of weakly
'attached' water molecules surrounding the central metal ion.
-
Also, thinking synoptically,
don't forget CI 1.5 "Concentration of
solutions" - mol and mass conversions, g/litre, g/dm3, mol/dm3,
calculation of concentrations or amounts in solution, volumetric calculations (eg
acid-alkali, iodine- thiosulphate).
Chemical
Ideas 3.2 The size of ions
-
The idea of charge density
(function of charge/radius) and the higher it is the greater the
-
Know the reasons for decreased
ion radius trend across a period and increasing down group.
-
The change of effective ion
radius on hydration - the smaller and more highly charged the central Mn+
ion the more strongly it attracts water producing an effectively larger ion
radius,
-
and so the larger the effective
radius the less strong the attraction (eg to ion-exchange resins or clay
minerals in soil)
Chemical Ideas 4.5 Energy changes in
solution
-
BE ABLE TO
WRITE OUT any definition clearly, quote standard conditions eg 298K,1
atm., 1 mole of
.., and any associated equations.
-
Lattice enthalpy ( HLE)
definition and related equations.
-
The factors affecting HLE
eg ionic charge and
radius.
-
Definition of Enthalpy of
hydration, Hhydration
(cation or anion), [ion(g) ==> ion(aq)],
energy change when gaseous ion becomes hydrated to form an aqueous solution.
-
The factors affecting enthalpy of hydration of ions
(solvation) eg size of charge and radius of ion.
-
Definition of Enthalpy of
solution, Hsolution
, energy changes when eg a salt dissolves in water (solid ionic solid ==>
aqueous solution of free ions).
-
Using Hess’s Law cycles
and enthalpy level diagrams to solve problems (be familiar with both cycle
and graph styles). WATCH
THE SIGNS and ARROW DIRECTION.
-
Reason why ionic substances do not readily dissolve in
non-polar solvents.
-
Read comments on Group 2
solubility trends.
-
Need CI 4.4 Entropy
etc. to use entropy factor promoting
dissolving process, and so understand why some solids dissolve despite the
fact that the Hsolution
may be endothermic.
-
No need for
section on the Solubility of Group 2 Compounds on pages 81-82 or
problem 6 on page 83
Chemical Ideas 4.6 The Born-Haber cycle
-
BE ABLE TO
WRITE OUT any definition clearly, quote standard conditions eg 298K,1
atm., 1 mole of
.., and any associated equations.
-
Reminder of lattice enthalpy,
HLE,
definition and related equations, factors affecting LE eg ionic charge and
radius.
-
The idea of breaking down the
formation of an ionic compound from its constituent elements into several
'theoretical' stages as means of understanding the energy changes for
the overall process.
-
Reminder of definitions of 1st
ionisation enthalpy, H1stIE,
and learning two new ones, 1st electron affinity, H1stEF,
and enthalpy of atomisation, Hatom.
-
Solving Born-Haber Cycle
problems in the style of Fig 26 or Fig 27.
-
WATCH THE
SIGNS and ARROW DIRECTION.
Activity
O1.1 What is the relationship between a
solvent and the substance that dissolves in it?
Activity
O1.2 What change occurs when an ionic solid
dissolves
Activity
O1.3 What factors affect the enthalpy
change of formation of an ionic compound
Chemical Storylines
O2 Wider Still and Deeper
-
Significance of the 70% water surface of the Earth, 97% of
all water is in the oceans.
-
This water transports enormous quantities of materials and energy
.
-
Its role as a
storehouse of food and chemicals is dwarfed by the impact it has on climate.
-
The oceans
are the central feature of what controls global conditions (conditions in which we live
and life has evolved).
-
Oceans are continuously surveyed to monitor its state and develop
‘global’ computer models.
-
The acid rain problem illustrates the complexity of
global models and problems so, in particular, study p244-245 and Fig 15 sulphur cycle and
appreciate the ‘missing’ sulphur
from dimethyl sulphide produced by marine algae.
-
Assignment 5e(ii) and
Assignment 6 are typically synoptic.
Chemical Storylines
O3 Oceans of Energy
-
The global central heating system,
apart from energy
released by fossil and nuclear fuels etc., the bulk of energy input to Earth comes from
the Sun’s radiation.
-
The heat absorbed drives the wind and waves.
-
Some of Sun’s rays reflected,
some absorbed by
atmosphere or Earth’s surface, and some re-radiated (see Figs 17-19).
-
If it wasn’t
for ‘global circulation’ of air and water the poles would be colder and the
tropics warmer.
-
The tropics are cooled by endothermic evaporation of water
and higher latitudes are
warmed by exothermic condensation of water.
-
In the Atlantic wind/water currents move
energy from the tropics SW to NE warming Northern Europe (Figs 24-25).
-
Explaining state changes (gas 
liquid
solid) in terms of intermolecular forces, energy involved and entropy
changes.
-
Fig 26 global water cycle, ideas of how energy is transferred from low to high
latitudes and warming the land.
-
The high enthalpy of vaporisation makes water a good energy
carrier (g/l) AND its high specific heat capacity (l).
-
As well as precipitation of rain,
the Atlantic ‘conveyer belts’ of water also warms Western Europe (Fig
25-26), this 'Gulf Stream' meets currents of cooler less dense water from melted ice/snow from Greenland,
this cools some of Gulf Stream, which is already more dense due to salts, so cooled more
salty water sinks, this produces a deep water cold current flowing in the opposite
direction (Fig 25).
-
Another deep water current is generated in the Antarctic (S Pole!),
this is produced by increased saltiness in the Antarctic water, as the water freezes the
salt remains in solution, residual water is more concentrated and dense, this sinks and
forms a conveyer belt of water heading for the Pacific.
-
The two deep water currents meet
in the South Atlantic producing a total global circulation system (Fig 26).
-
All of this
is ok BUT if the ‘conveyers’ are inhibited, the climate consequences are
dramatic! eg an ice age in Northern Europe. If the northern polar ice melts eg with
global warming ('Greenhouse Effect'), the water becomes less salty and less dense and the more salty/dense deep
ocean current slows down/switches off and the warm northern flowing Gulf stream then
suffers the same fate. So we in NW Europe get colder, but most of rest of world gets warmer
grrrr!!!! or should we say brrrr!!!!!
Chemical Ideas 4.4 Energy, entropy and
equilibrium
-
1st Revise CI 4.3 "Entropy and the direction of
change"?
-
Diffusion via random particle movement,
the mixed situation is the
most probable outcome.
-
Similar arguments for miscible liquids BUT if intermolecular forces
are strong in one of the liquids then two layers form eg oil/water (H bonding!).
-
Entropy
is a measure of the number of ways a system can be arranged, from this we can tell how or
direction the system will change.
-
In general the more spread out/mixed up/disordered, the
higher the entropy, this means effectively 'more ways to arrange the system'
-
In general entropy (S) for: gases > liquids > solids - but
entropy is not just about arrangement, we must also consider the ways that energy is
‘arranged’ and distributed in particles.
-
CI 4.4
-
CI 4.4: Relate g/l/s particle picture to properties and state changes.
-
The definition and units
of specific heat capacity (SHC), relate SHC to the KE of the particles in terms of
translational/rotational/vibrational KE AND most important of all now ...
-
Energy distributed not just in translational
KE, but also in rotation, vibration and also distributed in
electronic energy
levels (if input great enough, bond breaks).
-
All 4 forms quantised and note order of
quanta ‘gap’ differences between the four types of energy/quantum levels (Fig
15).
-
Entropy (S) and energy distribution,
the energy is distributed among
the energy levels in the particles to maximise entropy.
-
Entropy is a measure of both: the
way the particles are arranged AND the ways the quanta of energy can be arranged.
-
Appreciate why S(g) >
S(l) > S(s) and S is greater for larger atoms and larger molecules - excellent summary
p72.
-
Consider the enthalpy and entropy changes for H2O(l)
=> H2O(s) ...
-
Apply ideas to freezing seawater:
-
be able
to explain why it freezes at a lower temperature.
-
Its due to bigger decrease in
Sø(sys),
because of forming ordered ice out of a more complex mixture than pure
water.
-
Hence Sø(surr)
must be
greater, and the only way it can here is to have a smaller temperature
(in K), so that Sø(tot)
is still positive
-
Apply Sø(sys/surr/tot)
ideas to chemical changes to test feasibility of a
reaction:
-
ie is Sø(tot)
must be >0 ie positive
for a chemical change to be feasible
-
In the example on page 75, Sø(sys)
is
given
-
Sø(surr)
is
-Hø/T(K)
and delta H is very endothermic,
-
so at low temperature the Sø(surr)
term is too negative for Sø(tot)
to be plus overall,
-
but as the temperature is
raised the Sø(surr)
term becomes less negative and eventually Sø(tot)
becomes plus overall, ie the point of feasibility.
-
SCRIBBLE
ON ENTROPY related to these two example
-
For
equilibrium S(tot) =
0 and using the idea that at
equilibrium Sø(tot)
is 0 you can calculate the temperature when a reaction is
likely to become spontaneous or the temperature at which a change of state
occurs.
Chemical Ideas 5.4 Forces between
molecules: hydrogen bonding (revision)
-
1st Revision of CI 5.3 "Forces between molecules:
temporary and permanent dipoles"? - relating boiling points to intermolecular forces,
polarisation. Understand origin/examples of permanent, temporary/instantaneous and
induced dipoles d+ and
d-.
The three kinds of dipole interaction (1)permanent
dipole-permanent dipole, (2)permanent dipole-induced dipole, (3)instantaneous
dipole-induced dipole. Shape can influence the strength of intermolecular attraction
-
CI5.4: Looking in more detail at permanent
dipoles, bond polarity and dipole d+/d- diagrams, dipole originates from
electronegativity differences.
-
Permanent dipoles and origin of large dipole
effects ...
-
one or more very electronegative atoms in
molecule bonded to a less electronegative atom.
-
where dipoles can approach closely as in
the 'special' case of H-bonding, very electronegative O
and small H atom.
-
The elements O,N and F are all very electronegative,
(carrying the d-),
-
and lone pair of electrons to
line up with H, (d+)
-
H-bonding diagrams (remember it isn’t an actual covalent bond).
-
Examples and effects of H-bonding in eg HF (wrt to other HX),
H2O, nylon polymers,
proteins etc. already fully encountered.
-
Water has an unusually
high Hø(vap), mpt, bpt and
specific heat capacity (SHC in eg Jg-1K-1), it also decreases in density on freezing (compare water
properties with the other hydrides of Group 6).
-
All of these anomalies are accounted for
via hydrogen bonding. The H-bonding increases intermolecular forces (permanent
dipole-permanent dipole) so more energy needed, at a higher temperature. To effect state
changes more energy stored in water to weaken the H-bonds, so this raises the SHC.
-
Ice has an
‘open’ crystalline structure held by H-bonding (Fig 28) and as this breaks down on
melting, molecules can get closer, density increases.
-
Eventually with increased translational
KE with increasing temperature, you get normal thermal expansion and decreasing density with
increase in temperature [be able to explain Fig 27 using Fig 28!].
-
Its not
just global consequences eg burst pipes
but good news for fish and other aquatic life, as water freezes downwards!
Activity
O3.1 The enthalpy change of vaporisation of
water
Activity
O3.2 What crystals form when a solution is
cooled?
Chemical Storylines
O4 A safe Place to Grow
-
Consider solubility of CO2(g) in water,
increases with inc. P, dec. with inc. T (xref opened fizzy drinks going ‘flat’
fast in warm room!)
-
CO2(g) more soluble than other gases in air - low T and
higher P make more CO2(g) dissolve in oceans and this helps maintain a stable
environment
-
CO2(g) exchange between atmosphere and ocean is fast,
uptake of
CO2 is speeded up by the action of marine life ie photosynthesis in
phytoplankton.
-
CO2 has polar d+C=Od- bonds, is made more soluble via hydrogen
bonding with water,
-
(1) CO2(g)
CO2(aq) - the equilibrium is
moved further in the direction of increased solubility by two chemical equilibria
producing mainly hydrogen ions and hydrogencarbonate ions with a little carbonate ions:
-
(2) CO2(aq) + H2O(l)
H+(aq)
+ HCO3-(aq)
-
(3) HCO3-(aq)
H+(aq)
+ CO32-(aq)
-
If reactions 1 to 3 are ‘added’ you get (4)
CO2(aq) + H2O(l)
2H+(aq) + CO32-(aq)
(= H2CO3, see later)
-
35-50% of the excess CO2 from fossil fuel
combustion is dissolved in the oceans, more so in colder water - Le Chatelier’s
Principle says removing H+(aq), making it more alkaline (not easy or
normal!), will cause more CO2 to dissolve
-
Some marine organisms use the CO2
as CaCO3 in their shells - the Earth’s primeval atmosphere contained much
more CO2 but this has been drastically reduced by the evolution of marine
‘plant’ organisms and further CO2 reduction by shell production of
other marine life has also flourished, ultimately producing limestone /chalk (CaCO3)
-
CaCO3 sparingly soluble solid (refer
to Le Chatelier and Fig 36 for this section) ...
- solubility
product is (5) Ksp(CaCO3) = [Ca2+(aq)] [CO32-(aq)]
= 5.0 x 10-9 mol2 dm-6
- If Ksp exceeded,
ie [Ca2+(aq)] x [CO32-(aq)] > Ksp,
precipitation occurs, if not, ions remain in solution.
- Usually [Ca2+(aq)] and
[CO32-(aq)] concentrations are high enough near the surface so that
shells do not dissolve (but there is an equilibrium between the ions and the solid and
constant ion exchange).
- but things are different deep down where
the remains of dead
organisms/waste products of live creatures fall because the organic decomposed gives CO2
and the higher pressure and lower temperature increase solubility of CO2,
so although at first the
shells fall intact, at greater depths they dissolve because of reaction ...
- (6) CaCO3(s) + CO2(aq) + H2O(l)
Ca2+(aq) + 2HCO3-(aq)
moves right,
- and also (5) CaCO3(s)
Ca2+(aq)
+ CO32-(aq), this is more to the right because it is
exothermic.
- * These two reactions,
(6) and (5)*, explain why limestone
must have been laid down in shallower and warmer seas and there are no shells at the
bottom of the ocean. (* don't confuse with
endothermic thermal decomposition, got hydration enthalpies here!)
The formation of stalactites:
-
Rain
water dissolves more CO2 from ‘soil air’ which is 10-40x higher in CO2,
-
the ‘carbonated’ rainwater dissolves limestone via reaction (6) left to
right,
-
lower down the CO2 air concentration is normal and reaction (6) reverses to
precipitate CaCO3 as the calcium hydrogencarbonate decomposes (summary Fig
38)
The origins of life on Earth:
- Hydrothermal vents (p259) release CH4 , H2S and black
sulphide mineral specks.
- Colonies of tube worms rely on energy from bacteria
(probably similar to earliest forms of light and can live without light and oxygen).
- These bacteria gain energy by using sulphate ions to oxidise CH4 and H2S.
- Photosynthesis became possible with the evolution of cyanobacteria
and these produce
oxygen which was used up by soluble reducing agents (note cyanobacteria cannot tolerate
oxygen), so sulphate(VI) and nitrate(V) ions were the oxidising agents in
‘respiration’.
- Later marine organisms could use oxygen dissolved or in
atmosphere and formation of ozone layer is relatively recent.
- Cyanobacteria live on today in oxygen free environment and are very tolerant of uv.
- The
main changes in the Earth’s atmosphere reducing gases like methane and acidic gases
like carbon dioxide have been replaced by a neutral oxidising mix of oxygen and nitrogen
Need to understand that evolution in the oceans
requires pH stability (see Fig 42 p261)
- (pH 8 for millions of years, despite earlier high concentrations
of upto 35% CO2 - one reason is CO2(aq) is weak acid:
- reaction
(7) HA(aq) + H2O(l)
H3O +(aq) + A-(aq)
HA
is the weak acid ie H2CO3
The equilibrium is well over to the left, so only few
oxonium ions are formed, the dissolving of CO2 is an equilibrium itself (see
reactions 1-3) so amount dissolved is proportional to the amount of CO2 in the
atmosphere.
But the proportion which reacts is greatest when the CO2 is low,
so the equilibrium of reaction (7) lies to the rhs because of the large excess of water,
so the weak acid nature of CO2(aq) regulates its acidity ie there is never a
high concentration of H3O +(aq) ions
But a much more effective buffer
system is operating in the oceans [eg mix of a weak acid (carbonic acid H2CO3)
and salt of a weak acid and a
strong base (limestone CaCO3)] - the buffering reactions are (3, 5, 8 and 9)
- regard dissolved CO2 as a solution of
‘carbonic acid’
- (10) CO2(aq) +
H2O(l)
H2CO3(aq)
- (3) HCO3-(aq)
H+(aq) +
CO32-(aq) [ note H+(aq) used for simplicity instead of H3O
+(aq) ]
- (5) CaCO3(s)
Ca2+(aq) + CO32-(aq)
- (8) H2CO3(aq)
H+(aq) +
HCO3-(aq)
- (9) CO2(g) + H2O(aq)
H2CO3(aq)
Reactions (3) and (8) can provide a ‘weak’
source of hydrogen ions to neutralise any alkali eg
- (11) H+(aq) + OH-(aq)
=> H2O(l),
and because H2CO3(aq) or HCO3-(aq)
are weak acids there is always plenty of them left, as well as replacement from the
atmosphere by reactions (8) and (9)
If there is a rise in acidity, reaction (3) right to
left, will mop up hydrogen ions and reaction
(5), via limestone/chalk dissolving will
replenish the carbonate ions
Rather a neat system, instead of the weak acid and its
salt being permanently present (as in laboratory buffers), the ‘stock’ buffering
chemicals are stored in the atmosphere (CO2) and sedimentary rocks (CaCO3)!
If the CO2 concentration rose to what
it was 2 billion years ago, solid CaCO3 would dissolve to produce HCO3-
and CO32- which are needed to remove H+ ions.
The equilibrium constants are such that few CO32- ions would
remain, most of the carbon would be in the form of HCO3- and
dissolved CO2 and the sea would be like ‘Perrier water and
bicarbonate of soda’ and the white limestone cliffs and shells of sea
creatures would dissolve! [ by reaction (6)]
For this reason, the first rocks were
silicaceous,
calcium carbonate deposits could only be formed later, when phytoplankton and their
descendents began to significant reduce the carbon dioxide concentration in the atmosphere
(this would reduce the ‘greenhouse effect’ and allow ‘ice ages’ to
occur)
The carbon dioxide/calcium carbonate system is fast
acting but later a 2nd powerful buffering system via ion exchange between H+
and Na+ or K+ ions in clay sediments (see AA unit).
This process can only take place at the bottom of the ocean where seawater and sediment
are in contact. Deep ocean water circulates slowly so the 1st buffering
mechanism described, is important on a shorter time scale for keeping the pH of the
surface water and oceans stable.
Assignments 9 to 14 are all good 'applied'
chemistry questions with a strong synoptic flavour.
Chemical Ideas 7.1 Chemical equilibrium
(revision)
-
meaning of dynamic equilibrium and sign - conditions for
equilibrium eg rate forward = rate backward, no net change in concentrations - reversible
reactions - introduction to equilibrium expressions and quantitative interpretation -
position of equilibrium - constancy of K at constant temperature - apply ideas to the
dissolving carbon dioxide in water.
-
revise Le Chatelier’s Principle - direction of
change if pressure is changed (inc P, shift to less gas molecules ie to try to reduce
enforced inc P) - direction of change if temperature changed (inc T favours endothermic
direction, dec T favours exothermic direction) - only temperature affects the value of K,
the equilibrium constant - a catalyst does not affect the position of an
equilibrium, only alters rate at which equilibrium position is reached
-
problems: predicting
direction of change, Kc expressions, affect of changes on Kc and
equilibrium position, deduction from given data for a reaction - whether increase or
decrease in gas molecules or exothermic or endothermic reaction, suggest highest yield
conditions
-
Problems on p131 – position of equilibrium -
concentration effects - prediction of equilibrium position * these section should have
been done in earlier unit A, EP or SS
Chemical Ideas 7.2 Equilibria and
concentrations (revision)
-
writing out equilibrium
expressions with/without powers, numerical analysis of data to show constancy of
Kc, units of Kc, how position of equilibrium changes if particular concentrations of
components are changed, Le Chatelier’s principle, writing
out equilibrium expressions and solving problems via Kc expressions.
Chemical Ideas 8.1 Acid-base reactions
(revision)
-
definitions and properties of acids in terms of H+
transfer, alkalis as soluble bases, oxonium ion, acid-base pairs, strengths of acids and
bases, function of indicators.
Chemical Ideas 7.7 Solubility
equilibria
-
idea of a sparingly soluble ionic solid, equilibrium expression of dissolving
process .
-
Kc expression reformulated as the Ksp solubility
product expression
-
examples of Ksp calculations to show whether or
not precipitation can occur ie when the product of the actual (*) ion
concentrations exceeds the Ksp expression.
-
Writing out Ksp
expressions for more complex formula (watch the powers, especially in calculations)
-
The
effect on adding a ‘common’ ion on the solubility both qualitatively and
quantitatively - check Ksp problems on p190 and (*) watch out for
concentration adjustments on mixing solutions! If you mix 50:50 the
'effective' molarity for any calculation is halved, conversely, if you have
to work an original concentration to ..., then you double the concentration
that comes out of the Ksp
expression.
Chemical Ideas 11.2 The s block: Groups 1
and 2 (revision)
-
physical properties, chemical similarity, reaction of metals with water or
acid, reaction of oxides with water, hydroxides and their pH in water, reaction of oxides,
hydroxides and carbonates with acids, effect of heating the carbonate, solubility trends
of the hydroxides and carbonates.
Chemical Ideas 8.2 Weak acids and pH
-
Acids are hydrogen ion donors,
but different donating capacities ie 'strength'.
-
Difference between weak and
strong acids is about extent of ionisation (a few % compared to approximately 100%).
-
Equilibrium equations and Ka expressions for weak acids.
-
The pH scale
(Fig 5) and the definition of pH = -lg[H+(aq)],
meaning minus log to the base 10 of the hydrogen concentration in mol dm-3.
(note: p means –lg)
-
An aqueous solution pH is measured accurately
with a pH meter calibrated with buffers of known pH.
-
Examples of calculating the
pH of ...
-
strong acid, quite
straightforward
-
weak acid solutions more
tricky using the acidity or dissociation constant Ka equilibrium
expression, be able to quote the assumptions and justification of them.
-
Be able to 'reverse' the
calculation for a weak acid: measuring pH of
a weak acid of known molarity, hence calculate Ka. Its
a matter of working through carefully and logically.
-
Ka values
can be very wide ranging so Ka values are sometimes expressed as pKa = -lg Ka
(so make sure you can convert back to Ka for the purpose of calculations!)
-
The self-ionisation
of water and the Kw = [H+(aq)] [OH-(aq)]
and using the
Kw expression to calculate the pH of a strong base solution.
-
Examples of comparing weak
and strong acid phenomena using solutions of similar molarity eg (i) pH
of solutions, (ii) their electrical conduction (function of ion
concentration) and (iii) rate of reaction with a metal or a carbonate (rate
of fizz!).
Chemical Ideas 8.3 Buffer solutions
-
Definition of a buffer solution,
solution that minimises pH change when small amounts of acid or alkali are
added to a solution.
-
How buffers work in mopping
up H+(aq) or OH-(aq) ions - Bronsted Lowry acid-base
reactions that remove H+ or OH+ ions and check notes on Activity
O4.2 below.
-
The buffer calculations
follow on from the weak acid calculations in CI 8.2, again, be able to quote
and justify the assumptions. Its a matter of
working through carefully and logically.
Activity O4.1 Finding out more about weak
acids
Activity O4.2 Investigating some buffer
solutions
Activity
O5 Check your notes on The Oceans, Storylines O5 Summary, are incorporated in
the O
learning objective list and don't forget your O UNIT TEST etc. and
note what you got wrong! Q4c(iii)/(iv) not needed now.
GENERAL
REVISION
NOTES
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Sø(products) -
