Transcript Chemistry revision booklets!

Chemistry revision!
1.
2.
The early periodic table
The modern periodic
table
3. Group 1 – alkali metals
4. Group 7 – halogens
5. Transition elements
6. Strong and weak acids
and alkalis
7. Titrations
8. Titration calculations
9. How ideas about acids
and alkalis developed
10. Water and solubility
11. Solubility curves
12. Hard water
13. Removing hardness
14. Water treatment
15. Comparing the energy
produced by fuels
16. Energy changes in
reactions
17. Calculations using bond
energies
18. Test for positive ions
19. Tests for negative ions
20. Testing for organic
substances
21. Instrumental analysis 1
22. Instrumental analysis 2
In the 19th Century some chemist
dudes had a fight…
• In the 1800s people were on a roll discovering
elements. Scientists didn’t like how disorganised
they were so they tried to organise them. The
problem was they didn’t know a whole lot about
them!
• Not much was known about atoms
• Each element had several names
• A load of elements hadn’t been discovered yet!
John Dalton
• Dalton put the elements in
order of mass, measured by
doing different chemical
reactions.
• Problem was, that didn’t tell
anyone a lot about the
elements and when more
elements were discovered
the list had to be changed!
John Newlands
• Newlands introduced a bit more order; the law of octaves
(8)
• He based this on the fact that every eighth element seemed
to have similar properties.
• Problem is, he got a bit full of himself and tried to make all
the elements fit in the octaves even if they didn’t fit the
pattern, assuming all had been discovered in spite of the
fact new ones were being discovered all the time. He even
put 2 elements in one spot to make them all fit.
• All the other scientists laughed at him .
Alexandre-Emile Beguyer
de Chancourtois
• This French dude
copied Newlands idea
of octaves, but put
them in a much clearer
diagram.
• Unfortunately when his
work was published the
diagram was missed
out!
• Without the diagram
everyone got confused
so they ignored him 
•
Dmitri
Mendeleev
This Russian dude cracked it.
• By the time he ordered the elements 50
had been discovered. He put them in
order of atomic masses and then
arranged them so a periodic pattern in
their physical and chemical properties
could be seen.
• The clever part was he left GAPS for
elements that had not been discovered
yet when the pattern wasn’t followed.
He could predict the properties of
these elements thanks to his table.
• A few years later elements were
discovered that fit Mendeleev’s
predictions.
• All the other scientists thought he was
brilliant 
1. The Early Periodic Table
Dalton – elements in order of mass
Newlands – law of octaves, organised into
similar properties
Mendeleev – periodic table used today. In
order of atomic masses, and structured
according to chemical and physical
properties. He left gaps which were
filled when new elements were
discovered
2. The modern Periodic Table
• Elements are arranged in order of their
atomic (proton) number, and in groups of
similar properties.
• Groups have the same number of
electrons in their outer shell.
3. Alkali Metals
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VERY reactive, stored in oil (fizz and burn H2 with water)
Reactivity increases down the group
Low density (1st 3 float on water!)
Soft (can cut with a knife)
Low melting and boiling points for metals
Melting points decrease down the group
React with non-metals, losing their outer electron
Lithium + water → lithium hydroxide + hydrogen
2Li (s) + 2H2O (l) → 2LiOH (aq) + H2 (g)
Sodium + chlorine → sodium chloride
2Na (s) + Cl2 (g) → 2NaCl (s)
4. Halogens
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Poisonous non-metals
Coloured vapours
Low melting and boiling points
Poor conductors of heat and electricity
Fluorine – yellow gas
Chlorine – green gas
Bromine – orange/brown liquid
Iodine – gray solid – violet vapour
Go around in pairs (F2, Cl2…)
Less reactive as you go down the group
H2(g) + F2(g) → 2HF(g)
Displacement: Cl2 + 2KBr → 2KCl + Br2
5. Transition Elements
• Held together by metallic bonding
• Good conductors of heat and electricity
because of delocalised electrons
• Hard, tough and strong
• Malleable
• High melting points (apart from mercury)
• Much less reactive than alkali metals
(corrode slowly)
• Combining them makes useful alloys
• Make coloured compounds
Task:
• Page 222-3
• Summary questions 1 - 6
Strong and weak acids/alkalis
• Acids form H+ ions when we add them to water
HCl(g) → H+(aq) + Cl-(aq)
• The H+ (hydrogen ion) is the acidic part.
• Acids are PROTON DONORS. A proton is an H+
• Form OH- ions when we add them to water
NaOH(s) → Na+(aq) + OH-(aq)
• OH- (hydroxide ion) is the alkaline part.
• Alkalis are PROTON ACCEPTORS.
• An acid or base is STRONG if it completely ionises in
water.
• An acid or base is WEAK if it only partially ionises in
water.
7. Titrations
Indicators are used to show the endpoints of neutralisation
reactions:
• Strong acid + strong alkali - Universal indicator or any
other
• Weak acid + strong alkali - Phenolphthalein
• Strong acid + weak alkali - Methyl orange
In a titration we need to use the correct indicator
We also use pipettes to measure out fixed volume of
solution and burettes to measure the volume of solution
added.
8. Titration Calculations
• On the bottles you find
the name of the acid and
alkali and a number
followed by a M.
• M means the number of
moles in 1000cm3 (or 1dm3)
of solution. ‘Molarity’ or
‘molar concentration’ mean
the same thing.
• n = M x V (dm3) for liquids
9. How ideas about acids and
alkalis developed
• Liebig defined an acid as a compound that
contained hydrogen which could react with a
metal to produce hydrogen gas.
• Arrhenius defined an acid as a substance that
produces hydrogen ions (H+) in water, and
bases as a substance that produces hydroxide
ions (OH-) when dissolved in water.
• As Arrhenius’ definition only worked in
aqueous solution another definition had to be
developed.
• Brønsted and Lowry defined an acid as a
proton donor and a base as a proton acceptor
Task
Page 232-3
Summary Questions 1 – 3
Exam style questions 1 - 2
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10. Water and solubility
Soluble - substances that can dissolve.
Insoluble – substances that cannot dissolve.
Solution - the mixture formed when a substance dissolves.
Solute - the substance that dissolves.
Solvent - the liquid in the solution.
A saturated solution is one
in which no more solute will
dissolve at that temperature.
When a hot saturated solution
cools some of the solute will
separate from the solution.
11. Solubility Curves
• The solubility of most
solid solutes
increases as the
temperature
increases
• The solubility of
gases decreases as
temperature rises.
• The solubility of
gases increases as
pressure increases
12. Hard water
• Soft water readily forms lather with soap. Hard water reacts with soap
to form scum and so more soap is needed to form a lather.
• Hard water contains dissolved compounds, usually of calcium or
magnesium. The compounds are dissolved when water comes into contact
with rocks.
• Using hard water can increase costs because more soap is needed.
• Salts in the water react with the soap to produce stearates (scum)
• When hard water is heated it can produce scale (calcium carbonate)
that reduces the efficiency of heating systems and kettles.
• Hard water has some health benefits because calcium compounds are
good for health.
• Calcium is good for strong teeth and bones.
• There is evidence that drinking hard water reduces the chances of
heart disease.
13. Removing hardness
• Hardness is from dissolved calcium and
magnesium ions. The ions come from rocks
which the water has filtered through.
• The ions can be removed using washing soda
(sodium carbonate), which precipitates the
ions.
• An ion exchange column can be used. They
contain sodium ions which can be exchanged
with the calcium or magnesium ions.
14. Water Treatment
• Water from boreholes is usually pretty clean, it’s
been filtered by the surrounding rocks. Usually you
just use chlorine to kill of germs.
• Water from rivers or reservoirs needs more
treatment. Treatment involves chemical processes,
like adding aluminium sulphate and lime, and physical
processes, like filtration.
Water that has been treated is not pure. It still
contains substances dissolved in it.
Pure water is produced by distilling it, boiling and
condensing the steam produced, or deionising it by
using an ion exchange column.
Task:
• Page 246-7
• Summary questions 1 - 3
15. Comparing the energy produced by fuels
• The relative amounts of energy released when
substances burn can be measured by simple
calorimetry, e.g. by heating water in a glass or metal
container. This method can be used to compare the
amount of energy produced by fuels and foods.
• Energy is normally measured in joules (J). Some
dietary information is given in calories, which are equal
to 4.2 joules.
Different foods produce
different amounts of energy.
Foods with higher proportions
of carbohydrates, fats and oils
produce relatively large amounts
of energy.
16. Energy changes in reactions
• Energy (heat) is being put in to break bonds in the reactants.
• At the top of the curve, the bonds in the reactants have been broken.
The amount of energy put in to break these bonds is called the
activation energy.
• The activation energy is the minimum amount of energy needed for the
reaction to occur. A catalyst may work by lowering the activation energy
for a reaction.
• Energy (heat) is given out as bonds form in the products.
The difference in energy levels
between the reactants and the
products is given the symbol DH
(pronounced 'delta H').
This is the amount of heat given out
(or taken in) during the reaction.
For an exothermic reaction,
DH is negative.
For an endothermic reaction,
DH is positive.
17. Calculations using bond energies
H2 + Br2 → 2HBr
H–H
Br – Br
H – Br
H - Br
• H-H = 436 kJ/mol, Br-Br = 193jK/mol, H-Br =
366kJ/mol
• Reactants: 436 + 193 = 629 kJ/mol
• Product: 366 x 2 = 732 kJ/mol
• 629 – 732 = -103 kJ/mol
-ve means exothermic
Task:
• Page 256-7
• Summary questions 1 - 2
18. Tests for positive ions
Element Flame
Colour
Li
Bright
red
Na
Golden
yellow
K
Lilac
Ca
Ba
Brick
red
Green
Add sodium hydroxide (NaOH)
• Cu (II) – light blue ppt
• Iron(II) – dirty green ppt
• Iron(III) – red/brown ppt
• Al, Ca, Mg – white ppt
Add more NaOH to Al/Ca/Mg:
• Al – white ppt dissolves again
• Ca, Mg – white ppt does not
dissolve
• Add NaOH to ammonium ions
(NH4+)
• Makes ammonia!
• Warm soln and use damp red
litmus on gas – turns blue
19. Tests for negative ions
Carbonates (CO32-):
- Add dilute acid – makes
CO2.
- CO2 turns limewater
cloudy/milky
• Copper carbonate turns
green to black when heated
• Zinc carbonate turns white
to yellow when heated (and
kept hot!)
Nitrates:
• Add NaOH and warm it. If
no ammonia detected add
Al and test for ammonia
again (damp red litmus
blue)
Sulphates:
• Add hydrochloric acid
and barium chloride.
• Makes white ppt (barium
sulphate)
Halides:
• Add dilute nitric acid and
silver nitrate.
• Cl ions – white ppt
• Br ions – cream ppt
• I ions – yellow ppt
20. Testing for organic substances
• Contain carbon
• Burn or char on heating
• You can detect C=C double bond with
bromine water (orange to colourless)
• COMBUSTION ANALYSIS allows you
to determine the empirical formula of
an organic substance.
21. Instrumental Analysis
Technology is awesome:
1. Highly accurate
2. Quicker
3. Enable small quantities to be analysed
Technology has not so awesome:
1. Usually expensive
2. Takes special training to use
3. Gives results that can often be interpreted only by
comparison with already available known specimens
22. Instrumental analysis 2
• AAS – atomic absorption spectroscopy, measures
concentrations of metals in liquids
• Mass spectrometry – compares masses of different atoms
using magnets!
• UV – visible spectroscopy
• NMR (nuclear magnetic resonance) spectroscopy
Chromatography:
• Gas-liquid – separates compounds easily vaporised
• Gel permeation – separates according to size of molecules
• Ion-exchange – separates according to charge
• High performance liquid – separates compounds in solution