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LATTICE ENERGY
Lattice energy is the enthalpy change when 1 mole of an ionic
compound is formed from its gaseous ions under standard conditions.
Lattice energy is always exothermic, the more exothermic, the
stronger the lattices ionic bond. The standard enthalpy change of
atomisation is the enthalpy change when 1 mole of gaseous atoms is
formed from its element under standard conditions. The first
electron affinity is the enthalpy change when 1 mole of electrons
is added to 1 mole of gaseous atoms to form 1 mole of gaseous 1-
ions under standard conditions. (exothermic) The second electron
affinity is the enthalpy change when 1 mole of electrons is added
to one mole of gaseous 1- ions to form 1 mole of gaseous 2- ions
under standard conditions. (2nd and 3rd electron affinities are
always endothermic because there must be an input of energy to
overcome the repulsive forces between the electron and the negative
ion) Born-Haber Cycle
1. Write its elements in standard conditions on the left-hand
side in the amount as needed for the lattice (e.g. O2 (g) is
allowed).
2. Point downwards to the bottom for Hf of the lattice.
3. Continue the steps as needed to form the ions in the gaseous
state.
Point the arrows accordingly whether the reaction is endothermic
or exothermic.
Pay attention to the ionisation energy and electron affinity.
(e.g. to make Mg2+, the 1st and 2nd ionisation energies must be
taken into account; to make 2Cl(g) from Cl2 (g), calculate 2Hat;
2Cl(g) + 2e- 2Cl-(g) requires 2Hea1)
4. Once the elements are in its ionic gaseous state and with the
right amount needed to form the lattice, point downwards to the
bottom to form Hlatt.
Factors affecting the value of lattice energy: Lattice energy
arises from the electrostatic force of attraction of oppositely
charged ions when the crystalline is formed.
More exothermic as the size of the ion decreases.
Ions with the same charge have a lower charge density if their
radius is larger. A higher charge density results in stronger
electrostatic forces of attraction in the ionic lattice, so more
exothermic.
More exothermic as the charge on the ion increases.
Ion polarisation The positive charge on the cation in an ionic
lattice may attract the electrons in the anion towards it resulting
in ion polarisation. Polarising power: The ability of a cation to
attract electrons and distort an anion. Factors affecting ion
polarisation: The degree of polarisation of an anion depends
on:
The charge density of the cation. The ease with which the anion
can be
polarised its polarisibility. An anion is more likely to be
polarised if:
The cation is small The anion is large The cation has a charge
of 2+ or 3+ The anion has a charge of 2- or 3-
Many ionic compounds have some covalent character due to ion
polarisation. Many covalent compounds have some degree of charge
separation, i.e. they are polar, due to bond polarisation. Soluble
and insoluble are only relative terms. No metallic salts are
absolutely insoluble in water, but dissolves to a very small
extent. If salts were completely insoluble they could not have a
value for Hsol.
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The thermal stability of Group II carbonates and nitrates The
Group II carbonates decompose to their oxides and CO2 on heating.
The more positive the enthalpy change, the more stable is the
carbonate relative to its oxide and CO2. The relative stabilities
of these carbonates increases down the group:
MgCO3 < CaCO3 < SrCO3 < BaCO3
(least positive Hr) (most positive Hr)
The carbonate ion has a relatively large ionic radius so it is
easily polarised by a small highly charged cation.
The ionic radius of the Group II cations increase down the group
(the smaller the better it is at polarising)
The greater the polarisation, the easier it is to weaken a C-O
bond in the carbonate and form CO2 and the oxide on heating.
The same goes for Group II nitrates: these decompose to form the
oxide, NO2 and O2. The order of stability is:
Mg(NO3)2 < Ca(NO3)2 < Sr(NO3)2 < Ba(NO3)2 (most
polarised) (least polarised)
Enthalpy changes in solution The enthalpy change of solution is
the energy absorbed or released when 1 mole of an ionic solid
dissolves in sufficient water to form a very dilute solution. (can
be exothermic or endothermic) A compound is likely to be soluble in
water only if Hsol is negative or has a small positive value.
Substances with large positive values of Hsol are relatively
insoluble. The enthalpy change of hydration is the enthalpy change
when 1 mole of a specified gaseous ion dissolves in sufficient
water to form a very dilute solution. (always exothermic because it
forms bonds with water) ions (g) ions (aq) When an ionic solid
dissolves in water, bonds are formed between water molecules and
the ions (ion-dipole bonds).
The - oxygen atoms in water molecules are attracted to the
positive ions in the ionic compound. The + hydrogen atoms in water
molecules are attracted to the negative ions in the ionic compound.
The energy released in forming ion-dipole bonds is sufficient to
compensate for the energy that must be put in to separate the ionic
bonds in a lattice. The value of Hhyd is more exothermic for ions
with higher charge density. Hlatt + Hsol = Hhyd (add the Hhyd
values for both cations and anions) Energy level diagram to
calculate for the enthalpy change in solution
1. Write the gaseous ions. 2. Point downwards for Hlatt to form
the ionic
solid, then downwards again for Hsol to form the ions (aq).
3. Point downwards from the gaseous ions to the bottom for Hhyd
to form the ions (aq).
The solubility of Group II sulfates The solubility decreases as
the radius of the metal ion decreases. The order of the solubility
is:
MgSO4 > CaSO4 > SrSO4 > BaSO4
Change is hydration enthalpy down the group Smaller ions with
the same charge have
greater enthalpy changes of hydration. So the enthalpy change of
hydration gets
less exothermic down Group II. This decrease is relatively large
down the
group and it depends entirely on the increase in the size of the
cation since the anion is unchanged.
Change in lattice energy down the group
Lattice energy is greater if the ions with the same charge
forming the lattice are small.
The lattice energy decreases down Group II. The lattice energy
is inversely proportional
to the sum of radii of the anion and cation. The decrease in
lattice energy is relatively
smaller down the group and it is determined more by the size of
the much larger sulfate ion than the size of the cations.
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Difference in enthalpy change of solution of Group II
sulfates
The lattice energy of the sulfates decreases by relatively
smaller values down the group.
The enthalpy change of hydration decreases by relatively larger
values down the group.
So applying Hesss law, the value of Hsol gets more endothermic
down the group.
So the solubility of the Group II sulfates decreases down the
group.
ELECTRODE POTENTIALS
The standard hydrogen electrode consists of: H2 gas at 101kPa
pressure (1 atm) H+ ions of concentration 1.00 mol dm-3, in
equilibrium with H2 gas A platinum electrode
The more positive the electrode potential, the easier it is to
reduce the ions. So the metal on the right is relatively unreactive
and is a relatively poor reducing agent. The more negative the
electrode potential, the more difficult it is to reduce the ions.
So the metal on the right is relatively reactive and is a
relatively good reducing agent. Combining half-cells Standard
electrode potential (E):
Concentration of ions at 1.00 mol dm-3 A temperature of 25C (298
K) Pressure of 1 atm (101 kPa)
The standard electrode potential for a half-cell is the voltage
measured under standard conditions with a standard hydrogen
electrode as the other half-cell. Half-cells are connected together
using:
Wires connecting the metal rods in each half-cell to a
high-resistance voltmeter; the electrons flow round this external
circuit not through the electrolyte solution from the more negative
electrode potential to the more positive electrode.
A salt bridge to complete the chemical circuit allowing the
movement of ions between the two half-cells so that ionic balance
is maintained; a salt bridge does not allow the movement of
electrons.
A salt bridge can be made from a strip of filter paper (or other
inert porous material) soaked in a saturated solution of KNO3.
Metal/metal ion half-cell
Reduction takes place at the positive terminal of the cell (the
half-cell with the more positive E value)
Oxidation takes place at the negative terminal of the cell (the
half-cell with the more negative E value)
Non-metal/non-metal ion half-cell Electrical contact with the
solution is made by using platinum wire or platinum foil as an
electrode. The redox equilibrium is established at the surface of
the platinum. The platinum electrode is inert so plays no part in
the reaction. The platinum must be in contact with both the element
and the aqueous solution of its ions. Ion-ion of the same element
in different oxidation states half-cell The concentration of each
ion present is 1 mol dm-3. Using E values Standard cell potential:
the difference in standard electrode potential between two
half-cells. In order to calculate the cell voltage, always subtract
the less positive E value from the more positive E- value. The more
positive the value of E, the greater is the tendency for the
half-equation to proceed in the forward direction. The more
positive the value of E, the easier it is to reduce the species on
the left of the half-equation.
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A reaction is said to be feasible if it is likely to occur. Will
____ oxidise __ ions to __ ions?
1. Compare the E values for both half-cell reactions.
2. The reaction with the more positive E values will go to the
forward reaction, while the other reaction will go in the reverse
direction.
3. Check whether the combined equation satisfies the
question.
4. If so, then the reaction is said to be feasible. A reaction
will occur if the E value is positive. How does E vary with ion
concentration?
If the E values of the two half reactions involved under
non-standard conditions differ by more than 0.30 V, then the
reaction predicted by the E values is highly likely to occur. The
feasibility of a reaction based on E values is no guarantee that a
reaction will proceed quickly. It only tells us that a reaction is
possible, and that the reverse reaction does not occur. Some
reactions are feasible, but they proceed so slowly that they do not
seem to be taking place. It is the rate of reaction rather than
then value of E which is determining the lack of reactivity. Cells
and batteries Primary cells use redox reactions until the reactants
reach a low concentration and the voltage of the cell declines.
Secondary cells can be recharged by passing an electric current
through them. The products are then changed back to reactants so
the cell can function again.
Advantages of fuel cells: Water is the only product made. They
produce more energy. They are very efficient. There are no
moving
parts where energy is wasted as heat. Limitations to
hydrogen-oxygen fuel cells:
High cost: the materials used to make the electrodes and
membrane are expensive.
Manufacturing of fuel cells involves the production of toxic
by-products
Storage of hydrogen: high-pressure tanks are needed in order to
store a sufficient amount of fuel. At present, refueling has to be
done more often compared to petrol.
Manufacturing hydrogen: the hydrogen needed for fuel cells can
only be produced cheaply by using fossil fuels.
Fuel cells do not work well at low temperatures.
Electrolysis of molten electrolytes When pure molten ionic
compounds containing 2 simple ions are electrolysed, a metal is
formed at the cathode and a non-metal at the anode. Electrolysis of
aqueous solution Aqueous solution of electrolytes contain more than
one cation and more than one anion. The cation in the half-equation
with the most positive E value will be discharged. At the anode,
the ease of discharge of anions follows the order: SO42- (aq) NO3-
(aq) Cl- (aq) OH- (aq) Br- (aq) I- (aq)
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increasing ease of discharge/oxidation
Electrolysis products and solution concentration An ion, Z,
higher in the discharge series (more difficult to discharge) may be
discharges in preference to one below it if Z is present at a
relatively high concentration than normal. For this to be possible,
the E values of the competing ions are usually less than 0.30V
different from each other.
Electrode potential
Concen
tration Reactant Product
+ +
-
Quantitative electrolysis Faradays first law: the mass of
substance produced at an electrode is proportional to the quantity
of electricity passed Faradays second law: the number of Faradays
required to discharge 1 mole of an ion at an electrode is equal to
the charge on the ion. (the number of moles of the electrons needed
to reduce/oxidise 1 mole of a substance) 1 F = 96 500 C mol-1
Finding for Avogadros constant: L = charge on 1 mole of
electronscharge on 1 electron (1.60 10!!" C)
How to find the charge on 1 mole of e- 1. Weigh the pure copper
anode and pure
copper cathode separately. 2. Do an electrolysis experiment. 3.
Remove the cathode and anode and wash
and dry them with distilled water and propanone.
4. Reweigh the cathode and anode. 5. Measure the decrease in
mass of the anode. 6. Use proportions to find the charge
required
for 1 mole of the element. 7. Calculate for the charge on 1 mole
of
electrons by using proportions again.
IONIC EQUILIBRIA
Strong acids ionise completely in water. Weak acids only ionise
to a small extent in water. The ionic product of water, Kw
H2O (l) H+ (aq) + OH- (aq) Kw at 298 K = 1.00 10-14 mol2 dm-6 pH
= log10[H+] [H+] = 10-pH log10 [H+] log10 [OH-] = 14 The pH of
strong acids The concentration of hydrogen ions in solution is
approximately the same as the concentration of the acid. Diluting
the acid 10 times reduces the value of the H+ ion concentration by
one-tenth and increases the pH by a value of one. The pH of strong
bases The concentration of hydroxide ions in a solution is
approximately the same as the concentration of the base. H! =
K![OH!] Weak acids Ka is the acid dissociation constant.
HA (aq) H+(aq) + A- (aq) K! = H! [A!][HA] = H! ![HA]
A high value of Ka indicates that the position of equilibrium
lies to the right. The acid is almost completely ionised. A low
value for Ka indicates that the position of equilibrium lies to the
left. The acid is only slightly ionised and exists mainly as HA
molecules and comparatively few H+ and A- ions.
pKa = log10 Ka
The less positive the value of pKa, the more acidic is the acid.
In order to calculate the value of Ka, we make 2 assumptions:
We ignore the concentration of [H+] produced by the ionisation
of the water molecules present in the solution. This is reasonable
because the ionic product of water (Kw) is negligible compared with
the values for most weak acids.
We assume that the ionisation of the weak acid is so small that
the concentration of undissociated HA molecules present at
equilibrium is approximately the same as that of the original
acid.
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Acid-base indicators The acid and conjugate base have different
colours. HIn H+ + In-
un-ionised indicator conjugate base
colour A colour B
Adding an acid to this indicator solution shifts the position of
equilibrium to the left.
Adding an alkali shifts the position of equilibrium to the
right. There are now more ions of colour B.
The colour of the indicator depends on the relative
concentrations of HIn and In-. The colour of the indicator during a
titration depends on the concentration of H+ ions present.
The midpoint of the sharp fall in the titration graph
corresponds to the point at which the H+ ions in the acid have
exactly reacted with the OH- ions in the alkali; this is the
end-point of the titration. For weak acids with weak bases: There
is no sharp fall in the graph line. No acid-base indicator is
suitable to determine the end-point of this reaction. Buffer
solution A buffer solution is a solution in which the pH does not
change significantly when small amounts of acids or alkalis are
added. A buffer solution is used to keep pH (almost) constant. One
type of buffer solution is a mixture of a weak acid and one of its
salts. The weak acid stays mostly in the unionised form and only
gives rise to a low concentration of ions in solution. The pH of a
buffer solution depends on the ratio of the concentration of the
acid and the concentration of its conjugate base. If this does not
change very much, the pH changes very little. CH3COOH (aq) H+ (aq)
+ CH3COO- (aq) (high conc. of ethanoic acid and ethanoate ion) When
H+ ions are added to the buffer solution:
P.O.E. shifts to the left The large reserve supply of ethanoate
ion
and ethanoic acid ensures that the conc. in solution does not
change significantly.
So the pH does not change significantly.
When OH- ions are added to the buffer solution: The added OH-
combine with H+ to form H2O P.O.E. shifts to the right The large
reserve supply of ethanoate ion
and ethanoic acid ensures that the conc. in solution does not
change significantly.
So the pH does not change significantly. If very large amounts
of acid or alkali are added, the pH will change significantly.
Buffer solutions which resist changes in pH in alkaline regions are
usually a mixture of a weak base and its conjugate acid. pH of
buffer solution: pH = pK! + log!" [salt][acid] Uses of buffer
solutions:
Industrial processes including electroplating Manufacture of
dyes Treatment of leather To calibrate pH metres
In humans, blood pH is kept constant between 7.35 7.45 by a
number of buffers:
HCO3- Haemoglobin and plasma proteins H2PO4- and HPO42-
CO2 + H2O H+ + HCO3- (catalysed by carbonic anhydrase)
If H+ ion concentration increases: P.O.E. shifts to the left.
Reduces the concentration of H+ in the
blood and keeps the pH constant. If H+ ion concentration
decreases:
P.O.E. shifts to the right. Increases the concentration of H+ in
the
blood and keeps the pH constant. Equilibrium and solubility A
solution is saturated when no more solute dissolves in it.
Solubility product Ksp is the product of the concentrations of each
ion in a saturated solution of a sparingly soluble salt at 298 K,
raised to the power of their relative concentrations.
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The idea of solubility product only applies to ionic compounds
which are only slightly soluble. The smaller the value of Ksp, the
lower is the solubility of the salt. If Ion product > Ksp,
precipitation occurs. The common ion effect The common ion effect
is the reduction in the solubility of a dissolved salt achieved by
adding a solution of a compound which has an ion in common with the
dissolved salt. This often results in precipitation.
For example, AgCl (s) Ag+ (aq) + Cl- (aq)
Then add NaCl:
The Cl- is common to both NaCl and AgCl. P.O.E. shifts to the
left AgCl is precipitated because [Ag+][Cl-] > Ksp
The solubility of an ionic compound in aqueous solution
containing a common ion is less than its solubility in water.
REACTION KINETICS
Methods for following the course of a reaction Sampling This
method involves taking small samples of the reaction mixture at
various times and then carrying out a chemical analysis on each
sample. Samples are removed at various times and quenched to stop
or slow down the reaction. E.g. by cooling the sample in ice.
Continuous A physical property of the reaction mixture is monitored
over a period of time. The rate of reaction Rate of reaction = k
[A]m[B]n
k = rate constant m and n = orders of the reaction (usually 0,
1, 2, 3) The order of reaction with respect to a particular
reactant is the power to which the concentration of reactant is
raised in the rate equation. How to identify the order of a
reaction:
Plot a graph of reaction rate against concentration of
reactant
Plot a graph of concentration of reactant against time
Deduce successive half lives from graphs concentration against
time
Graphs of reaction rate against concentration It is very rare to
obtain an order with respect to a particular reagent higher than
second order.
Zero-order The graph is a horizontal straight line. The rate
does not change with concentration. First-order The graph is an
inclined straight line going through the origin. The rate is
directly proportional to the concentration. Second-order The graph
is an upwardly curved line. The rate is directly proportional to
the square of the concentration. Graphs of concentration of
reactant against time Zero-order The graph is a descending straight
line. The rate of reaction is the slope of the graph. First-order
Declines in a shallow curve. Second-order Declines in a deeper
curve which then levels out.
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Half-life and reaction rates Half-life, t1/2, is the time taken
for the concentration of a reactant to fall to half of its original
value. Zero-order Has successive half-lives which decrease with
time. First-order Has a half-life which is constant. Second-order
Has successive half-lives which increase with time. Calculating k
from half-life For a first-order reaction, half-life is related to
the rate constant by the expression k = 0.693t!/! Kinetics and
reaction mechanisms
A reactant that appears in the chemical equation may have no
effect on the reaction rate.
A substance which is not a reactant in the chemical equation can
affect reaction rate.
Rate-determining step: the slowest step in a reaction mechanism.
If a substance does not appear in the overall rate equation, it
does not take part in the rate-determining step. (and vice versa)
The mechanism is not deduced from the kinetic data. The kinetic
data simply shows us that a proposed reaction mechanism is
possible. The slow step may not involve molecules A and B but the
intermediate which is involved in the slow step which is derived
from A and B, then A and B will appear in the rate reaction. If a
molecule is not involved until after the rate-determining step,
then the rate reaction does not depend on the concentration of that
molecule.
Catalysis Homogeneous catalysis occurs when the catalyst is in
the same phase as the reaction mixture. Homogeneous catalysis often
involves changes in oxidation number of the ions involved in the
catalysis. Ions of transition elements are often good catalysts
because of their ability to change oxidation number. Heterogeneous
catalysis occurs when the catalyst is in a different phase to the
reaction mixture. A redox reaction occurs in heterogeneous
catalysis. Energy must be supplied to overcome the negative charges
of the two reactants by using a catalyst with a positive charge.
Heterogeneous catalysis often involves gaseous molecules bonding to
the surface of a solid catalyst in a process called adsorption.
Transition elements such as nickels are good at chemisorbing
hydrogen gas. The weak van der Waals forces link the hydrogen
molecule to the nickel, causes the weakening of the H-H covalent
bond. Heterogeneous catalysis in the Haber process The iron
catalyst works by allowing H2 and N2 molecules to come close
together on the surface of the iron. They are then more likely to
react.
1. Diffusion of H2 and N2 to the surface of Fe. 2. Adsorption.
The bonds formed between the
reactants and Fe are strong enough to weaken the covalent bonds
within H2 and N2 so they can react with each other but weak enough
to break and allow the products to leave the surface.
3. Reaction of H2 with N2 4. Desorption. The bonds between
ammonia
and the surface of Fe weaken and broken. 5. Diffusion. Ammonia
diffuses away from the
surface of Fe.
Transition elements in catalytic convertersThe honeycomb
structure inside the catalytic converter contains small beads
coated with platinum, palladium or rhodium which act as
heterogeneous catalysts.
The steps are: Adsorption of NO and CO onto the catalyst surface
Weakening of the covalent bonds within NO & CO Formation of new
bonds to form N2 and CO2 Desorption of N2 and CO2 from the catalyst
surface.
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GROUP IV
The metallic character of the elements increases down the group.
Metalloids differ from metals in two ways:
Metalloids have very small electrical conductivities at room
temperature; metals conduct electricity very well.
The electrical conductivity of metalloids increases with
increase in temperature; the conductivity of metals decreases with
increase in temperature.
Tin and lead are less reactive than metals from Group I and II.
Melting points C, Si, Ge have a similar structure to diamond (giant
molecular), the difference in melting points can be related to
their bond energies. Sn and Pb have a metallic structure at room
temperature. The ions are held together by the electrostatic
attraction between their positive charges and the delocalised
electrons. Sn and Pb have relatively large ions, so the metallic
bonding is relatively weak and the melting points are relatively
low. Electrical conductivity Carbon (diamond) does not conduct
electricity because of the strong covalent bonds so no delocalised
electrons free to move around. Carbon (graphite) does conduct
electricity. Si and Ge conduct to a small extent. The weaker
bonding allows some electrons to move out of position especially if
there are traces of other contaminating atoms in the lattice, but
the electrons are not delocalised. Sn and Pb are conductors because
of the free delocalised electrons.
The tetrachlorides The tetrachlorides all have:
The general formula XCl4 Simple covalent molecules A tetrahedral
structure They are all volatile liquids at room temp. Low boiling
points
*arrow head always pointing to the largest/the most ___
Thermal stability
CCl4, SiCl4, GeCl4 are stable at high temp. They do not
decompose on heating.
SnCl4 decomposes readily on heating. SnCl! l SnCl! s + Cl!(g)
PbCl4 decomposes explosively.
Reaction with water
All the tetrachlorides except CCl4 are hydrolysed by water.
SiCl! l + 2H!O l SiO! s + 4HCl (g)
The ease of hydrolysis increases from SiCl4 to PbCl4 as the
metallic nature increases.
The oxides
Oxidation state +2 Oxidation state +4 CO CO2 SiO SiO2 GeO GeO2
SnO SnO2 PbO PbO2
Most of the Group IV oxides do not decompose on heating because
of the strong bonding in their structure.
Element Symbol Metal or non-metal? MP /C Electrical Conductivity
Bond
Bond energy/ kJ mol-1
Carbon C Non-metal 3550 Non-conductor Strong covalent 350
Silicon Si Metalloid 1410 Semi-conductor Strong covalent 222
Germanium Ge Metalloid 937 Semi-conductor Strong covalent 188 Tin
Sn Metal 232 Conductor Metallic bonding Lead Pb Metal 327 Conductor
Metallic bonding
Tetrachloride BP /C Thermal Stability Can be
hydrolysed? Ease
CCl4 76
SiCl4 57
GeCl4 87 SnCl4 114 PbCl4 105
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The +2 oxides Oxidation state +2 Structure, bond
Thermal stability Acid-base character
Strength as a reducing agent
CO SM, Strong Covalent
Very weakly acidic
SiO GM, Weak Covalent GeO GM, Weak Covalent Amphoteric SnO Ionic
Amphoteric PbO Ionic Amphoteric (B) Bonding and thermal
stability
CO has a simple molecular structure. The strong triple covalent
bond makes the molecules thermally stable.
SiO and GeO have giant molecular structure. They have weak
covalent bonds and they decompose on heating in a
disproportionation reaction. 2GeO s GeO! s + Ge(s)
SnO and PbO do not decompose on heating. When PbO is heated, it
forms a bright red
solid called triplumbic tetroxide Pb3O4, also known as red lead
oxide. This compound behaves chemically as PbO2.2PbO
Acid-base properties
CO is very slightly soluble in water. It forms a neutral
solution. It reacts with hot conc. NaOH solution. So it is an
acidic oxide.
GeO, SnO, PbO all react with HCl. SnO s + 2HCl (aq) SnCl! aq +
H!O(l) GeO, SnO, PbO all react with NaOH. PbO s + 2OH! (aq) PbO!!!
aq + H!O(l)
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The +4 oxides Oxidation state +4 Structure, bond
Thermal stability Acid-base character
Strength as an oxidising agent
Ease of reaction with alkali
CO2 SM, Strong Covalent
Acidic
SiO2 GM Very weakly acidic GeO2 GM Amphoteric (A) SnO2 GC
(~ionic) Amphoteric PbO2 GC (~ionic) Amphoteric Bonding and thermal
stability
PbO2 is the only one that decomposes readily on heating. 2PbO! s
2PbO s + O!(g)
Acid-base properties
CO2 is slightly soluble in water. CO! aq + H!O l HCO!! aq + H!
aq CO!!! aq + 2H! (aq) The P.O.E. lies well over to the left. Most
of the dissolved CO2 is in the form of CO2 (aq), so carbon dioxide
solution is only weakly acidic.
CO2 also reacts with alkalis CO! g + 2OH! (aq) CO!!! aq + H!O(l)
SiO2 does not react with acids (except
hydrofluoric acid). It reacts with hot concentrated alkalis to
form silicate (IV) ions and water, so it is an acidic oxide.
SiO2 + 2OH- SiO32- + H2O GeO2, SnO2, PbO2 are all amphoteric.
They
react with concentrated HCl to form the tetrachlorides. SnO2 (s)
+ 4HCl (aq) SnCl4 (l) + 2H2O (l)
GeO2, SnO2, PbO2 all react with hot concentrated alkalis.
SnO2 (s) + 2OH- SnO32- (aq) + H2O (l) GeO2 and PbO2 react in a
similar manner to
form germinate (IV) and plumbate (IV) ions (needs molten NaOH to
make plumbate (IV))
The reason why +4 oxidation state is less stable at the bottom
of the group can only be explained using complex Born-Haber
cycles:
Compounds low in the group having the +2 ox. state have greater
ionic character than those in the +4 ox. state.
+4 compounds low in the group have more weak covalent character.
The energy released on forming these bonds is low.
It takes less energy to oxidise an element to a low oxidation
state than to a high ox. state
At the bottom of the group, the energy released on forming
covalent bonds or ions in the +4 ox. state is not enough to
compensate for the extra energy required to form the +4 ox.
state.
At the bottom of the group, the +4 ox. state becomes less
energetically stable with respect to the +2 ox. state.
Towards the bottom, the oxidized form (e.g. Ge4+ or Pb4+) is
more readily reduced to the +2 state.
The value of E gets more positive down the group.
Ceramics from silicon (IV) oxide Also known as silicon dioxide.
It is used to make a variety of ceramics. It is either used on its
own or mixed with clay.
Ceramics containing silicon (IV) oxide are used:
For furnace linings As abrasives In the manufacture of glass and
porcelain
The properties of silicon (IV) oxide which make it an ideal
ceramic include:
Very high melting and boiling point. It needs a high temperature
to break the strong covalent bonds in the giant molecular
structure.
It is an electrical and thermal insulator (no free
electrons)
It is hard. It is difficult to break the 3D network of strong
covalent bonds.
It is generally chemically unreactive. It can be moulded at a
very high
temperature into a variety of shapes without affecting its
strength.
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TRANSITION ELEMENTS
A transition element is a d-block element which forms one or
more stable ions with an incomplete d sub-shell. Sc and Zn are not
transition elements.
Sc forms only one ion Sc3+ and this has no electrons in its 3d
sub-shell. (Ar) 3d0 4s0
Zn forms only one ion Zn2+ and this has a complete 3d sub-shell.
(Ar) 3d10 4s0
SPECIAL EXCEPTIONS Cr: 3d5 4s1 Cu: 3d10 4s1
The transition elements are all metals. Their atoms tend to lose
electrons so they
form positively charged ion. Transition metals have variable
oxidation
states and the resulting ions often have different colours.
The existence of variable oxidation states means that the names
of compounds containing transition elements must have their
oxidation number included. E.g. manganese (IV) oxide.
When transition elements form ions, their atoms lose electrons
from the 4s sub-shell first, followed by 3d electrons.
The most common ox. state is +2, formed when the two 4s
electrons are lost.
The max. ox. number of the transition elements at the start of
the row involves all the 4s and 3d electrons in the atoms.
At the end of the row, (from iron onwards) the +2 ox. state
dominates as 3d electrons become increasingly harder to remove as
the nuclear charge increases across the period.
The higher ox. states are found in complex ions or in compounds
formed with O2 or F. e.g. chromate (VI) ion & manganate (VII)
ion.
PHYSICAL PROPERTIES OF TRANSITION ELEMENTS
High melting point High density Hard and rigid, useful for
construction Good conductors of electricity & heat The 1st
ionisation energy, the atomic radius
& the ionic radius do not vary much as we go across the
first row.
Comparing the transition elements with an s-block element
(Ca)
The melting point of Ca < TE The density of Ca < TE The
atomic radius of Ca > TE The ionic radius of Ca > TE The 1st
ionisation energy of Ca > TE The electrical conductivity of Ca
> TE
(except copper) Redox Reactions When a compound of a transition
element is treated with a suitable reagent, the oxidation state of
the transition element can change. Balance a redox reaction by
making the number of electrons equal to both half-equation (does
not effect the value of E) The redox reaction equation can be used
to calculate the amount of a reactant by carrying out a titration.
Use Cr2O72- rather than MnO4- to obtain a more accurate result by
titration because compounds such as K2Cr2O7 can be prepared to a
higher degree of purity than KMnO4. Ligands and complex formation A
ligand is a species that contains one or more lone pairs of
electrons that forms a dative bond to a metal ion. A complex is a
compound formed by a central metal atom surrounded by one or more
ligands. Co-ordination number is the number of dative bonds formed
by ligands to the central transition metal ion in a complex.
Monodendate are ligands (e.g. H2O & NH3) which can form only
one dative bond from each ion/molecule to the central transition
metal ion. Bidentate are ligands which can form 2 dative bonds from
each ion/molecule to the central transition metal ion. The shape of
a complex with 6 ligands is octahedral. charge of a complex =
charge on (CMI + ligands)
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Tetrahedral complex e.g. [CuCl4]2- and [CoCl4]2-. The copper(II)
and cobalt(II) ions have four chloride ions bonded to them rather
than six, because the chloride ions are too big to fit any more
around the central metal ion. Square planar complex e.g.
[Ni(CN)4]2- and Pt(NH3)2Cl2 Pt(NH3)2Cl2 is neutral because the 2+
charge of the original platinum(II) ion is exactly cancelled by the
two negative charges supplied by the chloride ions. Substitution of
ligands The ligands in a complex can be exchanged, wholly or
partially, for other ligands. This is a type of substitution
reaction. It happens if the new complex formed is more stable than
the original complex. The colour of complexes The colour of
complexes containing transition metal ions arises because part of
the visible spectrum is absorbed by transition metal ions. The 5 d
orbitals in an isolated transition metal atom/ion are described as
degenerate (the orbitals are at the same energy level) In the
presence of ligands, a transition metal ion is not isolated. The
co-ordinate bonding from the ligands causes the 5 d orbitals in the
transition metal ion to split into 2 states of non-degenerate
orbitals at slightly different energy levels.
The lone pairs donated by the ligands into the transition metal
ion repel electrons in the two d!!!!! and d!! orbitals more than
those in the other three d orbitals. This is because these d
orbitals line up with the co-ordinate bonds in the complexs
octahedral shape and so they are closer to the bonding electrons in
the octahedral arrangement, increasing repulsion between electrons.
The difference between the non-degenerate orbitals is E. When light
shines on the solution, an electron absorbs E and jumps into the
higher energy level among the 2 non-degenerate orbitals. E is
affected by many factors: The identity of the ligand that surrounds
the transition metal ion. If E is different, then the amount of
energy being absorbed by electrons will be different. Therefore, a
different colour is seen.
APPLICATIONS OF ANALYTICAL CHEMISTRY
Electrophoresis: The separation of charged particles by their
different rates of movement in an electric field. The sample is
placed on absorbent paper or on a gel supported on a solid base
such as a glass plate. A buffer solution carries the solution
along. The rate at which the ions move towards the oppositely
charged electrode depends on the size and the charge on the ions.
(smaller and higher charge is faster) You get a series of lines or
bands on the paper/gel one a chemical is applied. Sometimes UV
light is
used to show the bands up. The series of bands is called an
electropherogram. Electrolysis is used in biochemical analysis. It
can be used to separate, identify and purify proteins & nucleic
acids. Buffer solution is needed to maintain the pH because pH will
affect the movement of ions during electrophoresis. When separating
a mixture of proteins, they are usually first treated with a
chemical that makes them negatively charged. A dye can also be
added.
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All the proteins move towards the positive electrode but larger
proteins move more slowly. Uses of DNA fingerprinting/profiling
Forensic Science Genetic fingerprinting: a technique based on
matching the minisatellite regions of a persons DNA to a database
of reference samples. We inherit DNA half from our mother and half
from our father. The process:
1. Restriction enzymes are used to cut the DNA molecule at
specific places where the same sequences occur, making smaller
fragments for analysis.
2. DNA fragments are all negatively charged because of the P
groups present, so they all move towards the + electrode. The
larger ones find it harder to get through the matrix formed by the
gel so they do not travel as far in a given time.
3. The bands are made visible by radioactive labeling of the
bands with the P-32 isotope, which causes the photographic film to
fog. Alternatively, use a probe that makes the bands fluoresce in
UV light.
Short-tandem repeat (STR) analysis
1. Short sequences of bases that make up genes are multiplied
using DNA polymerase to copy selected sequences.
2. People differ in the number of these short sequences that the
DNA contains. This affects the distances that the bands travel.
How an NMR works The nucleus of each hydrogen atom (proton) in
an organic molecule behaves like a tiny magnet. This proton can
spin. This movement of the (+) charged proton causes a very small
magnetic field to be set up. In NMR, we put the sample to be
analysed in the magnetic field. The proton, by spinning, either
line up with or against the field. There is a tiny difference in
energy between the oppositely spinning nuclei. This difference
corresponds to the energy carried by waves in the radio wave range
of the E.M. radiation spectrum.
The nuclei flip between the 2 energy levels. They absorb energy
in the range of frequencies that are analysed. The size of the gap
between nuclear energy levels varies depending on the other atoms
in the molecule (the molecular environment) In NMR spectroscopy, we
vary the magnetic field rather than the wavelength of the
radiowaves. As the magnetic field is varied, the H nuclei in
different molecular environment flip at different field strengths.
The different field strengths are measured relative to
tetramethylsilane (TMS SiCl4 an inert, volatile liquid which mixes
well with most organic compounds and its H atoms are equivalent)
which is zero. Spin-spin coupling: Peaks are made up of a series of
closely grouped peaks because the magnetic fields generated by
spinning nuclei interfere slightly with those of neighbouring
nuclei. The number of signals a peak splits into equals n+1 where n
is the number of H atoms on the adjacent C atom. Steps to interpret
high-resolution NMR spectrum
1. Use values to identify the environment of the protons present
at each peak.
2. Look at the relative areas under each peak to determine how
many of each type of non-equivalent protons are present.
3. Apply the n + 1 rule to see which protons are on adjacent C
atoms in the unknown molecule.
4. Put all this information together to identify the unknown
molecule.
For ethanol: The peak of the OH group is not split by the H on
the neighbouring CH2 group because the OH proton exchanges very
rapidly with protons in any traces of water (or acid) present.
(labile proton). This also occurs in amines and amides which
contain the NH group.
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Identifying the OH or NH signal in an NMR spectrum Add a small
amount of D2O (deuterium oxide) to the sample then their peaks
disappear from the spectra. The deuterium (2H) atoms exchange
reversibly with the protons in the OH or NH groups. The deuterium
atoms do not absorb in the same region of the E.M. spectrum as
protons, so the OH or NH signal disappears. X-ray crystallography
To get the best results, we need to have a very pure crystal of the
sample. Many large biological molecules can be crystallised from
solution. This technique relies on the diffraction of X-rays as
they pass into a crystal which is caused by the electrons in the
atoms present. The larger the atom, the more electrons it contains
and the more intense the spot it produces in its diffraction
pattern.