Transcript C4C5C6

In 1897 JJ Thomson “Plum Pudding” model for
the atom.
small negative charges (“electrons”) arranged
in a sphere of positive charge.
Rutherford “gold foil experiment”.
few positively charged particles were
deflected backwards.
Bohr - Each shell represents an
energy level. All electron in the same
shell have fixed energy
-1 charge, 0.0005 mass
+1 charge, 1 mass
0 charge, 1 mass
Relative Atomic
Mass – protons and
neutrons
Atomic Number –
protons (also number
of electrons)
Always the
biggest number!
Neutrons = atomic
mass – atomic number
Isotopes – a form of an atom with the same atomic
number (same number of protons) but a different
number of neutrons so it has a different relative
atomic mass.
Going across the periodic table the
atoms are arranged into periods.
Going down the periodic table the
Elements in the same period have the
atoms are organised into groups.
same number of shells
Elements in the same group have
the same number of electrons in
the outer shell.
They all have similar properties
because they have the same
number of electrons in the outer
shell.
In 1828, Dobereiner grouped the known
elements into triads based on their chemical
properties.
Middle element has the atomic mass which is the
mean of the total atomic mass of the triad
Newlands every 8th element had similar
properties so he organised them into
groups of 7 called “octaves”.
His work was not accepted because:
1) Not all elements had similar properties.
2) mixed up metals / non metals
3) no space for undiscovered elements.
Mendeleev - elements in order of atomic mass.
groups according to similar properties.
left gaps - predicted properties of undiscovered
elements.
Rules for drawing electron
configuration
• Atomic number tells the
number of electrons
around the nucleus.
• Electrons always occupy
shells (also known as
energy levels)
• Lowest energy levels are
always filled first.
• First shell – max 2
electrons
• Second shell onwards –
max 8 electrons
• Ca – 2.8.8.2
Ionic Bonding
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ions strongly attracted to each other. (+ to -)
Group 1/2 are keen to lose electrons.
Groups 6/7 are keen to gain electrons
Metal – non metal
Na loses e- =
positively charged.
Cl gains e- =
negatively charged.
Ionic Compounds
• giant ionic lattices
• ions strongly attracted to
each other / unable to
move.
• high melting / boiling
points
• don’t conduct electricity
when solid
• When melted or dissolved
- conduct electricity – ions
free to move.
• Draw the ionic bonding for:
– MgO
– CaCl2
Covalent Bonding (strong)
Methane (CH4)
• non-metal atoms - share pairs of
electrons.
• The forces between molecules, are
weak
• very low melting / boiling points.
• don’t conduct electricity.
Hydrogen Gas (H2)
Carbon Dioxide (CO2)
Water (H2O)
Chlorine Gas (Cl2)
Group 1 – The Alkali Metals
– Group 1 Metals - 1e• Reacting with water:
in outer shell
– Move around the
– go down group 1 surface, fizzing
metals more reactive.
violently, produce
– outer electron further
away from nucleus so
hydrogen.
less energy is needed
2Na + 2H2O  2NaOH + H2
to remove it.
– all have low melting
points, low density Sodium + Water  Sodium Hydroxide + Water
and are very soft.
Alkali metals burn with characteristic
colours:
Lithium: Red
Sodium: Yellow
Potassium: Lilac
• 7 e- in outer shell.
Group
7 – The Halogens
• go down group 7 - less reactive
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because there is less inclination to fill the outer shell as its
further from the nucleus.
Melting / boiling points increase.
At room temp:
– Cl2 is a poisonous green gas
– Br2 is a poisonous orange liquid.
– I2 is a grey solid
Halogens react with Alkali metals to form Metal Salts
2Na + Cl2  2NaCl
Sodium + Chlorine  Sodium Chloride
Displacement Reactions
A more reactive halogen can displace a less reactive halogen from its salt
chlorine + sodium bromide
Cl2 + 2NaBr
→
→
sodium chloride + bromine
2NaCl + Br2
Superconductors
• At low temperatures
• have little or no electrical
resistance.
• benefits:
– Power transmission without
loss
– Super-fast electronic
circuits
– Powerful electromagnets
– Superconducting
electromagnets used in
hospital MRI scanners
Drawbacks:
At the moment, superconductors have to be REALLY COLD. This is expensive to
achieve and takes a lot of energy.
Identifying transition metals
• Add sodium hydroxide =
Thermal Decomposition
displacement reaction.
• Na more reactive metal, displaces
transition metal
Transition
• substance breaks down into •two
or moremetal hydroxide is
insoluble in water = precipitate.
substances,when heated.
• Transition metal carbonates Copper hydroxide: blue precipitate
Iron (II) Hydroxide: Grey/green precipitate
• CuCO3  CuO + CO2
Iron (III) Hydroxide: Orange/Brown
• Test for carbon dioxide-limewater-cloudy.
precipitate
CuSO4 + 2NaOH  Cu(OH)2 + Na2SO4
Copper (II) Sulphate + Sodium Hydroxide  Copper (II) Hydroxide +
Sodium Sulphate
FeSO4 + 2NaOH  Fe(OH)2 + Na2SO4
Iron (II) Sulphate + Sodium Hydroxide  Iron (II) Hydroxide + Sodium
Sulphate
Fe2(SO4)3 + 2NaOH  2Fe(OH)3 + 3Na2SO4
Iron (III) Sulphate + Sodium Hydroxide  Iron (III) Hydroxide + Sodium
Sulphate
Water
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reservoirs, lakes, rivers, bore holes, aquifers.
Pollutants get into water... Factory output, Leaks in pipes
Natural disasters
Bad sanitation
Waterborne disease
Lead pipes dissolving into the water
Pesticides
Nitrates from fertilisers
STEPS
1. Sedimentation – particles drop to the
bottom
2. Filtration – of particles using sand
3. Chlorination – to kill microbes
PRECIPITATION reactions are used to test for the presence of IONS
in water.
IONS to test for:
Chemical used
Sulphate SO42- ......... barium chloride – white precipitate
Barium chloride(aq) + sodium sulfate(aq)→sodium chloride(aq) + barium
sulfate(s)
BaCl2(aq) + Na2SO4(aq) → 2NaCl(aq) + BaSO4(s)
Chloride ClBromide BrIodide I-
……… silver nitrate
White
Cream
Pale Yellow
Silver nitrate(aq)+sodium bromide(aq)→sodium nitrate(aq)+silver bromide(s)
AgNO3(aq) + NaCl(aq) → NaNO3(aq) + AgCl(s)
Write BALANCED symbol equations for these :
1. Silver nitrate + Sodium chloride
2. Silver nitrate + Sodium bromide
3. Silver nitrate + Sodium iodide
MOLES
In 1 mole there are 6.02x10^23 particles
Avagadro’s number = 6.02x10^23
• The mass of one mole is its molar
mass (RFM in grams)
For example ... Hydrogen's RFM is
1 ... Its molar mass is 1g
• RFM is relative to 1/12 carbon
Mass
(g)
Amount
of moles
Concent
ration
Volume
(dm3)
Converting concentration from mol/dm3 to g/dm3
Amount
Molar
of
mass
moles
1. How many moles are there in
66g Carbon Dioxide?
2. What mass of carbon is there
in 4 moles of carbon dioxide?
You have a solution of sulphuric acid of
0.04mol/dm3. What is the concentration
in g?
STEP 1: Work out RFM
H2SO4 = 98
STEP 2: Convert the conc in moles to
conc in grams. So in 1dm3
Mass = moles x RFM
0.04 x 98 = 3.92 g
• Calculate the mass of aluminium oxide when 135g of aluminium is
burned in air.
• Step 1: Write the balanced equation for the reaction
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(4Al + 3O2

2Al203)
• Step 2: Calculate the moles for the part you have the information for.
(moles of aluminium = 135/ 27 = 5)
• Step 3: Look at the ratio to give you the moles for the part that you
want.
(4 moles of Al react to form 2 molesAl2O3 so 5 moles would give
2.5 moles of aluminium oxide)
• Step 4: Use the equation mass = moles x Mr
(mass of aluminium oxide = 2.5 x 102 =255g)
• empirical formula shows you the simplest ratio of atoms
in a compound (C2H6 would become CH3.).
• To calculate this, all you need is the experimental
masses and the relative atomic mass (Ar), which is
found on the periodic table.
• Example: Find the empirical formula of iron oxide when
44.8g of iron reacts with 19.2g of oxygen
• There are 5 steps:
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1) List the elements 2) Write down the experimental masses 3) Divide each experimental mass by the Ar of each 4) Divide by smallest
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What is the empirical formula of
a) C7H14?
b) C6H12O6?
c) Al2O6?
• Find the empirical formula when:
• a) 2.4g of carbon react with 0.8g of hydrogen
• b) 21.9g of magnesium react with 29.3g of sulfur
and 58.3g of oxygen
Titrations use single indicators so it
makes it easy to see the end point of
the titration.
E.G phenolphthalein
Universal indicator is made from a
mixture of different indicators so
each colour indicates a range of pH
values.
Titrations
• Titrations are used to find out exactly
how much acid is needed to
neutralise an alkali or vice versa.
• It can then be used to calculate
unknown concentrations.
• Method:
– Fill a conical flask with 25cm3 alkali
of unknown concentration
– Add 2-3 drops indicator
– Fill a burette with acid
– Using the burette add the acid a bit at
a time (say 5cm3)
– When indicator changes colour you
have reached the end point. You now
have a rough estimate of how much
(to the nearest 5cm)
– Now repeat adding a smaller amount
of acid each time.
– To increase the accuracy you need to
get several consistent readings!
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Concentration = moles x volume
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You start off with 25cm3 of sodium hydroxide that has a concentration of
0.100 moles per dm3. It takes 49cm3 of hydrochloric acid to neutralise the
sodium hydroxide. What is the concentration of the hydrochloric acid used?
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Step 1: Work out how many moles of the “known” substance you have.
Number of moles = conc x volume
= 0.1 x (25/1000)
= 0.0025 moles of sodium hydroxide
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Step 2: Write the balanced symbol equation for the reaction. Work out how
many moles of the “unknown” stuff you had.
NaOH + HCl  NaCl + H2O
For every mole of NaOH, you need one mole of HCl
So you must need 0.0025 moles of HCl
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Step 3: Work out the concentration of the “unknown” stuff!
Concentration = no of moles / volume
= 0.0025 / (49/100)
= 0.0510 mol/dm3
When you do a titration there is a gradual change in
pH. At the end point there is a sudden change in pH.
The rate of a reaction can be measured
by the amount of gas produced.
Gas Syringe
Upturned measuring
cylinder/burette
Mass Loss
Method
Gas Syringe
Pros
Can be used to collect
pretty much any gas.
Accurate volumes – to
the nearest cm3
Upturned
measuring
cylinder
Mass Loss
Accurate volumes – to
the nearest cm3
Greater accuracy.
Cons
If the reaction is too
vigorous it can blow the
plunger out of the
syringe.
Cannot collect gases
such as Hydrogen
chloride or ammonia as
these dissolve in water.
Gas is released straight
into the room so not
suitable for reactions
that produce poisonous
gases.
Interpreting Rates of Reaction
Amount of Product
End of Reaction
Steeper
gradient –
faster rate of
reaction
More
reactant
Time
Reaching an Equilibrium – the Haber
Process
Equilibrium
1) As nitrogen and hydrogen react
• Haber Process
together their concentrations fall.
• N2 + 3H2
2NH3
The initial rate of reaction will begin
• equilibrium – rate of the forward reaction is equal to the
to slow down.
rate of the backward reaction
2) As more and more of the product
• closed system
ammonia is made, its concentration
rises and it begins to turn back into
nitrogen and hydrogen.
3) As more is made, the rate of the
reverse reaction speeds up.
4) After a while the forward reaction
will be going at the same rate as the
backward reaction.
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The position of the equilibrium can be in the middle, to the left or to the
right.
This tells us about the amounts of the products and reactants.
If the equilibrium is in the middle then there are the same amounts of
reactants as products.
If the equilibrium is to the right then there is more product and not so much
reactant.
If the equilibrium is to the left then there is more reactant and not so much
product.
Product
Reactant
Reactant
Equilibrium
Product
C5
• Three factors affect the
position
of the equilibrium:
Changing
Equilibrium
Temperature
For all reversible reactions, one direction is an exothermic
reaction and the reverse direction is endothermic.
• decrease temperature - rate of endothermic reaction decreases
•equilibrium will shift towards the exothermic reaction so that more
heat is produced.
Pressure (only for gases)
increase pressure-equilibrium tries to reduce it
equilibrium moves in the direction where there are fewer moles of
gas
Concentration
Increase concentration-equilibrium tries to reduce this (so more
product is made) = shifts to reduce this so more turns back into
the reactants.
The contact process
o
• 450 C -compromise, forward
1.)Burn sulphur in air to make reaction is exothermic so high
temps reduce yield and shift
sulphur dioxide
equilibrium to left. But, at high
S + 02
SO2
temps rate of reaction is quicker
so chemical is produced faster
2.)React sulphur dioxide with
• Atmospheric pressuremore oxygen in air to
compromise, 3 gas molecules on
create sulphur trioxide
the left and 2 on the right so high
2SO2 +02
2SO3
pressure increases yield by forcing
equilibrium to the right. However,
3. React SO3 with water to
equilibrium already lies to right so
makes sulphuric acid
the cost of thicker walls etc. to
SO3 + H2O
H2SO4
withstand higher pressure is not
economical
• vanadium pentoxide (V2O5)does
not affect the position of the
equilibrium but makes the
reaction go faster
Strong and Weak Acids
• Strong Acids ionise completely in water. This means that
the compound dissociates (e.g HCl  H+ + Cl-). There
is a higher concentration of H+ ions ready to react.
• Weak Acids only partially ionise in water. It is a
reversible reaction which sets up an equilibrium mixture.
• (e.g CH3COOH
H+ + CH3OO-)
• Only a few H+ ions are released at once so the
equilibrium is off to the left.
• Once these H+ ions have been used up a few more are
released.
• Strong acids are better electrical conductors because
they have a higher concentration of hydrogen ions to
carry the charge.
Preventing Rusting
Redox Reactions
1) Making alloys e.g steel
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2) Painting and oiling/greasing
Oxidation is Loss,
3) Galvanising – coating with a tin plate
Reduction is Gain
4) Sacrificial Protection – place a more
OIL RIG
reactive metal with the iron. The water and
Oxidising Agent – accepts oxygen
ewill react with this instead.
Reducing Agent – donates e-
Iron + Water + Oxygen Hydrated iron(III)oxide
• Electrolysis of aqueous sulfuric
acid.
• Ions: H+, OH-, SO42• Hydrogen ions accept electrons
from the cathode to make
hydrogen gas.
• At the anode, hydroxide ions
lose electrons to make oxygen
gas
• Products:
Cathode: Hydrogen
2H+ +2e-  H2
Anode: Oxygen
4OH- -4e-  O2 + H20
• Electrolysis of copper sulfate
using carbon electrodes.
which anion/cation is
easier to discharge?
The ion discharged first is
the one which is lower in
the reactivity series.
• Ions: Cu2+, H+, OH-, SO42• Copper ions accept electrons
from the cathode to make
copper.
• At the anode, hydroxide ions
lose electrons to make oxygen
gas
• Products:
• Cathode: Copper
Cu2+ +2e-  Cu
Anode: Oxygen
4OH- -4e-  O2 + H20
What’s the link between current and
charge when talking about electrolysis?
• The amount of product is proportional to
time and current
• Q = It
• t Time (seconds)
• I Current (amps)
• Q Charge (coulomb)
• Example question:
– A current of 0.1A for 2 hours increased the
mass of an anode by 0.24g. How much
charge was transferred?
Pollution
•Contain poisonous catalyst
•Burn fossil fuels to produce
hydrogen and oxygen
http://www.youtube.com/watch?v=6UwSazq8GTU
Making Ethanol
Fermentation
•Renewable – raw material from plant material
•Yeast is used to ferment glucose solution
•glucose→Ethanol+Carbon Dioxide
•C6H12O6 → 2C2H5OH + 2CO2
•(enzyme inactive) 25 - 50 degrees (denatured)
•Absence of air from fermentation prevents the formation of ethanoic acid by
oxidation of the ethanol.
•slow
•Ethanol purified by distillation, lots of energy, expensive
Hydration
•Ethene
Ethanol
•Ethene passed over heated phosphoric acid catalyst with steam.
•Ethene + water → Ethanol
•C2H4 + H2O → C2H5OH
•non-renewable as the ethene will be made by cracking components of crude oil.
•quicker
Depletion of Ozone
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In stratosphere
Ozone filters out and stops harmful ultraviolet light from reaching the surface of the earth
CFCs were used as refrigerants and in aerosols because they have a low boiling point, are
insoluble in water and are very unreactive.
Use of CFCs in the UK is now banned to stop any more damage to the ozone layer.
Instead hydrocarbons (alkanes) or hydrofluorocarbons (HFCs) are now used as safer
alternatives to CFCs.
The depletion of ozone in the atmosphere allows increased levels of harmful ultraviolet light to
reach the earth and this can cause……
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An increased risk of sunburn
Increased ageing of the skin
Skin cancers
Increased risk of cataracts
Free Radicals
• CFC molecule is hit by UV light a chlorine atom is produced. A chlorine atom is called a free
radical.
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The chlorine free radical reacts with an ozone molecule to form a chlorine oxide molecule and an oxygen
molecule
• Cl + O3 → ClO + O2
The chlorine oxide molecule then reacts with an oxygen atom to produce a chlorine free radical and an
oxygen molecule
• ClO + O → Cl + O2
The chlorine free radical is regenerated by this chain reaction and can go on to destroy many more
ozone molecules in a a chain reaction. Therefore a few chlorine atoms can destroy large amounts of
ozone.
Water Hardness
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– Hard = does not lather with soap
– Soft = lathers well with soap
Calcium and magnesium ions form dissolved salts which cause hardness in
water
There are two types of hardness in water
– Permanent = caused by dissolved calcium sulphate. Cannot be
removed by boiling
– Temporary
Calcium carbonate + water + carbon dioxide
calcium hydrogencarbonate
Can be removed by boiling
Decomposition of Ca(HCO3)2
Ca(HCO3)2
CaCO3 + H2O + CO2 (insoluble limescale)
Removes temp. and permanent
1. Ion Exchange resin –- sodium ions come off the resin and go into the
water, while calcium ions come out of the water and stick to the resin
2. Washing Soda – Na2CO3
Ca2+(aq) + CO32–(aq)→ CaCO3(s) Insoluble Limescale
(solids) Fats and Oils (liquids)
Oils and fats are ESTERS that can be obtained from animals or vegetables.
– Saturated = carbon-carbon single bonds
– Unsaturated = at least one carbon-carbon double bond
Test for Unsaturated fats
• Shake with bromine water : Orange to colourless
Natural Fats and Oils
• Animal oils/fats often saturated - Vegetable oils/fats often unsaturated
• More unsaturated = reduce build up of cholesterol.
Mixing Fats and Oils
• Oil and water are immiscible - do not mix.
• vegetable oil added to water + shaken well = emulsion.
• An emulsion is one liquid finely dispersed in another
• The shaking breaks up the oil into small droplets that disperse (spread out)
in the water.
– Milk is an oil-in-water emulsion that is mostly water with tiny droplets of
oil dispersed in it.
– Butter is a water-in-oil emulsion that is mostly oil with droplets of water
dispersed in it.
Saponification - vegetable oil + hot sodium hydroxide
glycerol + soap
• hydrolysis - breaking up ester groups in the oil molecule using an alkali
• Margarine : vegetable oil + hydrogen : nickel catalyst
solid saturated fat
Detergents
• hydrophilic head forms strong intermolecular forces with water molecules
• hydrophobic tail forms strong intermolecular forces with fat and oil molecules
Dry cleaning
• A greasy stain may be removed using dry cleaning solvent.
• There are weak intermolecular forces between the grease molecules, and
there are weak intermolecular forces between the solvent molecules.
• The solvent molecules can also form intermolecular forces with the grease
molecules. This lets the solvent molecules surround the grease molecules.