Transcript C2-Revision-Mats-(condensed)

T1: Reading the Periodic
Table
T1: Mendeleev
•
•
•
•
•
•
•
•
•
•
Arranged elements by increasing atomic
mass but….
He broke this rule and left some gaps if an
element’s properties weren’t similar to the
one above it.
He thought the gaps were for elements that
hadn’t been discovered yet and predicted
their properties.
When they were discovered, the properties
matched the predictions
20
90.5
21
0.3
22
9.2
•
• Note: on some periodic tables,
they are the wrong way up, just
remember that the smaller
number is the proton number.
•
•
T2: Common Ions
•
For example, Neon has three isotopes
Relative Abundance
(%)
•
Atomic number
(aka proton number):
The number of protons or
electrons.
Versions of an element with same atomic number
but different atomic mass.
Number of protons is the same, but number of
neutrons is different.
Relative Atomic Mass is average of the masses of
the isotopes, weighted by their relative abundance
Neon Isotope
Mass
Simple Covalent Molecules
Relative Atomic Mass
(aka nucleon number):
The total number of protons and
neutrons added together.
T1: Isotopes (HT)
•
Relative atomic mass of Neon =
20 × 90.5 + 21 × 0.3 + 22 × 9.2
= 20.2
90.5 + 0.3 + 9.2
This is why some atoms have a relative atomic mass
with a decimal point.
T1: What’s in my
atom?
T3: Covalent Structures
Molecule = A particle made of a
small group of atoms, covalently
bonded together.
Low melting and boiling point,
due to weak attractive forces
between molecules..
Electrical insulator as no electrons
free to move.
Examples: water, ammonia,
oxygen
Protons = atomic number
Electrons = atomic number
Neutrons = relative atomic mass
.
– atomic number
Electrons in
outer shell
Ion
formed
Examples
1
1
+
Li+, Na+, K+
2
2
2+
Be2+, Mg2+, Ca2+
3
6
2-
O2-, S2-
4
7
-
F-, Cl-, Br-, I-
T1: Electron Configuration
Electrons orbit the nucleus in shells.
First shell holds two electrons
Second and third shell hold 8 electrons
Note: the third shell can actually hold more,
but we won’t worry about this until A-level.
•When an insoluble salt is formed from the reaction of two
soluble salts.
•Goes cloudy as small particles of solid are made.
•Predicting precipitates: simply choose a combination of
soluble salts where you tell that if the ions swapped over you
would get an insoluble salt: use the solubility table for help.
•Example:
Lead nitrate + potassium iodide  lead iodide + potassium nitrate
Pb(NO3)2(aq) +
2KI(aq)
 PbI2(s) + 2KNO3(aq)
•
Example: Silicon
Atomic number is 14, so it has 14
electrons.
You build up electrons from the
first shell outwards, so in this case:
- First shell has 2
- Second shell has 8
- Third shell has 4
This can be written as: 2.8.4; or drawn as:
Note: Si is in period three and
group four of the periodic
table; it also has three
electron shells and four
electrons in the outer shell –
this is no coincidence!
T1: Relative Atomic Mass
T1: Atoms and Elements
•Element = substance containing only one type
of atom.
•Protons and electrons: same for every atom
of an element…it is the number of protons
that decides the element.
•Neutrons: can differ…atoms with the same
number of protons but different numbers of
neutrons are called isotopes
• Immiscible = when liquids do not dissolve in
each other….like oil and water, one floats on
top of the other.
• Can be separated with a separating funnel;
the denser layer is tapped-off at the bottom.
•
•
•
Example 1: Magnesium reacting with
chlorine.
• Anion: Cl forms Cl- ions
• Cation: Mg forms Mg2+ ions
• Formula = MgCl2
• Why: two Cl- gives a 2- charge to
balance 2+ from Mg2+.
• Name: magnesium chloride
Example 2: aluminium reacting with oxygen.
• Anion: O forms O2- ions
• Cation: Al forms Al3+ ions
• Formula = Al2O3
• Why: Two Al3+ gives a 6+ charge, three
O2- gives a 6- charge.
• Name: aluminium oxide
T3: Separating Miscible Liquids
•Miscible = when liquids
dissolve in each other…like
alcohol and water.
•Separate with fractional
distillation using a
fractionating column.
•The components of the mixture
have different boiling points,
so if you heat it, each
component will boil at a
different time, allowing you to
collect and condense the pure
vapour.
•We can do this to separate the
gases in air by first cooling the
air to turn the gases to liquid.
T1: Sub-atomic particles
•
•
Atoms are made from smaller particles called
subatomic particles.
There are three types we need to know about,
summarised below.
Particle
Proton
Neutron
Electron
Relative
Relative
charge
mass
1
Positive, +1
1
Neutral, 0
Neglible Negative, -1
1
(
)
1840
Found?
In nucleus
In nucleus
In shells orbiting
nucleus
T2: Properties of Ionic Compounds
•
•
•
•
•
Melting point: High due to strong bonds between ions.
Boiling point: Higher, due to strong bond between ions.
Solid: do not conduct electricity
Molten (liquid): do conduct electricity
Dissolved (aqueous): do conduct electricity
Why? (HT)
Electrical Conductivity
• Electricity is conducted when there are charged
particles that are free to move.
• Solid: there are charged particles (the ions), but they are
not free to move, so they do not conduct.
• Liquid/Aqueous: the ions are now free to move, so they
do conduct
High Melting/Boiling Points
• Ionic bonds (attraction between positive and negative
ions) are very strong.
• Melting and boiling require these bonds to be broken.
• This takes lots of (heat) energy.
•
•
•
Example 1: Water
Each hydrogen needs one more electron to
complete it’s outer shell and the oxygen
needs two more. Oxygen forms two single
bonds: one to each hydrogen.
H
H
O
Example 2: Carbon dioxide (HT only)
Carbon needs two more electrons to
complete it’s outer shell and each oxygen
needs two more. Carbon forms two double
bonds: one to each oxygen.
O
C
•A repeating 3D
lattice of
positive and
negative ions.
•Strong
electrostatic
bonds
between ions.
T2: Barium
Meals
• A patient is given a
drink containing
barium sulfate.
• This can show up on
a x-ray, helping
doctors to investigate
the digestive system.
Clean a metal loop in acid
Did loop in a metal salt.
Heat in roaring Bunsen flame.
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Sodium, Na+  Yellow
Potassium, K+  Lilac
Calcium, Ca2+  Red
Copper, Cu2+  Green-blue
Precipitation Tests
Chloride: add acidified silver nitrate to
get a white precipitate if chloride is
present.
Sulfate: add acidified barium chloride
to get a white precipitate if sulfate is
present.
Carbonate Test
1.
2.
Add acid to the sample
Pass any gas produced through
limewater: will go cloudy if the
sample contained carbonate
Soluble: a compound dissolves in a given liquid.
Insoluble: a compound does not dissolve.
Soluble in water
In soluble in water
All sodium, potassium,
ammonium salts
All nitrates
Most chlorides
Except: silver and lead
chlorides
Most sulfates
Except: lead, barium
and calcium sulfates.
Except: sodium, potassium Most carbonates
and ammonium carbonates
Except: sodium, potassium Most hydroxides
and ammonium hydroxides
T3: Diamond vs Graphite (HT)
Diamond:
•
•
•
Very hard, as all carbon atoms joined
with strong covalent bonds.
Used to make cutting tools
Insulator as all electrons locked-tight
in bonds, so can’t move.
Graphite:
•
O
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•
T2: Flame tests
1.
2.
3.
T2: Solubility
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T2: Ionic
Structures (HT)
Layers of hexagonal carbon mesh
that rub away from each other, as
there are only weak forces between
the layers.
Used as a lubricant.
Conductor as the electrons between
the layers are free to move. This is
very rare for a giant covalent
structure.
PERIODS….increasing atomic mass, differing properties
GROUPS……similar properties
•This is the mass of an element relative
to 1/12th the mass of 12C.
•Element: substance containing only one
type of atom.
•Protons and electrons: same for every
atom of an element…it is the number of
protons that decides the element.
•Neutrons: can differ…atoms with the
same number of protons but different
numbers of neutrons are called
isotopes.
T3: Separating Immiscible Liquids
An ionic bond is the attraction between a
positive and a negative ion.
The overall number of positive and
negative charges must cancel out.
Form between a metal and a non-metal
Ionic compounds do not form molecules
T3: Covalent Bonds
Form when non-metals share electrons
between them.
Attraction between each atom and the
shared electron pair.
Atoms share electrons to complete their
outer shells
One bond is formed for each ‘gap’ in the
outer shell
Bonding represented with dot-and-cross
diagrams showing only the outer-shell
electrons.
T2: Forming Ions
Cations are positive (cat…pussitive!) ions
They are formed when atoms lose electrons.
Metals form cations by losing the electrons in their
outer shells
In the example, aluminium loses its three outer-shell
electrons to become Al3+…each lost electrons cause 1 ‘+’
charge.
3+
Anions are negative ions
They are formed when atoms gain electrons.
Non-metals form anions by filling their outer shells.
Name ends with ‘-ide’ to show it is a negative ion,
In the example, oxygen gains two outer-shell electrons
to become O2-, giving it 8 electrons in its outer shell.
T2: Precipitates and Precipitation
• Repeating pattern of
many millions of atoms
covalently bonded.
• High melting/boiling
point because much
heat energy needed to
break strong covalent
bonds.
• Electrical insulator as
no electrons free to
move.
• Examples: silicon
dioxide, diamond,
graphite
T2: Making Ionic Compounds
•
2.
Atomic number = 9
Relative Atomic mass = 19
Protons = 9
Electrons = 9
Neutrons = 19-9 = 10
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React solutions of (the right)
two soluble salts together.
Filter the mixture to collect
the precipitate.
Rinse the filter residue with
distilled water to remove
impurities.
Allow the residue to dry.
4.
Giant Covalent
There are also some ‘compound’ ions made of
more than one atom with an overall charge:
•
Hydroxide: OH•
Nitrate: NO3•
Sulphate, SO42•
Carbonate, CO32•
Ammonium, NH4+
1.
3.
You should try to memorise the ions formed
by various species:
Group
T2: Making Insoluble
Salts
Element Type
= non-metal
= metal
• Group 1: Lithium (Li), Sodium (Na), Potassium (K)…
• Properties: low melting point, soft (can be cut with
a knife).
• React with water as follows:
General equation: metal + water  metal
hydroxide + hydrogen
For example: 2K(s) + 2H2O(l) 
2KOH(aq)
+ H2(g)
T5: Rates of Reaction (Intro)
T4: Transition
Metals
T4: Alkali Metals
• High melting points
• Form brightly coloured
compounds
Reactivity
Explaining Reactivity (HT only)
• All reactions require you to remove the outer-shell
electron/
• Atoms get bigger going down the group  outershell electrons further from nucleus  easier to
remove the outer shell electron.
T4: Halogens and Their Reactions
• Group 7: Fluorine (F) – pale yellow gas, Chlorine (Cl) – pale
green gas, Bromine (Br) – orangey-brown liquid, Iodine (I) –
grey solid.
• Most reactive at top of group, and get less reactive as you go
down.
• Form halide ions with a charge of ‘-1’
Reaction with metals
• React with metals to form metal halides
• General equation: metal + halogen  metal halide
• For example: magnesium + iodine  magnesium iodide
Mg(s) + I2(s) 
MgI2(s)
Note: Mg forms a 2+ ion, so two I- ions are
needed.
Reaction with hydrogen
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•
•
•
React with hydrogen to form hydrogen halides.
Hydrogen halides dissolve in water to form acids.
General equation: metal + halogen  hydrogen halide
For example: hydrogen + fluorine  hydrogen fluoride
H2(g) + F2(g) 
2HF(g)
Note: hydrogen fluoride dissolves to make
hydrofluoric acid.
Displacement Reactions
• More reactive halogens can react with the ions of less
reactive halogens and displace them from compounds.
• For example: 2KI(aq) + Br2(aq)  2KBr(aq) + I2(aq)
• This reaction works because bromine is more reactive
than iodine.
• The orange colour of bromine would change to the
brown colour of aqueous iodine.
• The reverse reaction would not work.
T5: Endothermic and Exothermic
Exothermic Reactions
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•
•
Chemical energy is converted to heat energy.
The surroundings get hotter.
For example: combustion reactions:
Methane + oxygen  carbon dioxide + water
CH4 + 2O2 
CO2
+ 2H2O
•
Explosions are just very fast exothermic
reactions.
Endothermic Reactions
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Energy Diagrams (HT only)
Halogen
•
Displacement reactions can be used to determine the
order of reactivity of the halogens.
Try reacting each halogen with solutions of each halide
salt, the halogen that does most reactions is most reactive.
Halide Salt
Potassium Potassium Potassium Potassium
fluoride
chloride
bromide
iodide
Fluorine
x
Reaction
Reaction
Reaction
Chlorine
No
x
Reaction
Reaction
reaction
Bromine
No
No reaction
x
Reaction
reaction
Iodine
No
No reaction
No
x
reaction
reaction
ENDOTHERMIC
Reactants
Products
Energy
released
so gets
hotter
Products
•
•
Reactivity Series of Halogens
•
In reactions, old chemical bonds are broken,
and then new ones are made.
Breaking bonds takes in energy; making
bonds gives out energy.
Stronger bonds take more energy to break,
and give out more when made.
In exothermic reactions, weaker bonds are
broken and stronger bonds are made.
In endothermic reactions, stronger bonds are
broken and weaker bonds are made.
EXOTHERMIC
Simple molecular
Giant
Molecular
Swapping electrons
to form ions
Sharing electrons
Sharing
electrons
Examples
Sodium chloride,
magnesium oxide
Water, methane,
nitrogen
Quartz (silicon
dioxide)
Bond
strength
Strong
How the
bonds form
Melting and
boiling point
Energy
absorbed
so gets
colder
Reactants
T5: Collision Theory (HT)
To react: particles must collide with enough
energy.
To increase rate: increase the amount of
collisions or the energy of the collisions.
Conduct
electricity?
•
Increasing concentration increases the number
of reacting particles.
This increases the number of collisions.
Effect of Surface Area:
•
•
Increasing the surface area increases the
proportion of (solid) particles available to
react.
This increases the number of collisions.
Effect of Temperature:
•
•
Increasing the temperature increases the
speed that particles are moving
This means there are more collisions, and those
collisions have more energy.
• Toxic carbon monoxide and unburned
hydrocarbons (from petrol) are converted into
carbon monoxide and water.
• The catalytic converter has a fine honeycomb
structure coated with the catalyst.
• The catalyst contains a mixture of platinum,
rhodium and palladium.
• The metals are expensive, so only a very thin
coating is used.
• The catalysts work best at high temperatures, so
car exhaust is more damaging when the car has
only just started and hasn’t warmed up.
Strong bonds, weak Strong bonds
intermolecular forces
High
Low
High
Most in water
Some in water
Insoluble in
water
Only when molten or
dissolved
No
No (except
graphite)
• Lord Rayleigh noticed the density of nitrogen made in
reactions was less than nitrogen made from air.
• Sir William Ramsey hypothesised that the nitrogen in
the air must also contain a denser gas that had not yet
been discovered.
• Through careful experiments, Rayleigh and Ramsey
discovered a gas that they named ‘argon’.
• They also discovered helium, and then later Ne, Kr
and Xe.
Uses:
T4: Metallic Bonding
• He and Ar were used to stop in filament in old bulbs
burning.
• Ar and He used in welding to stop hot metal oxidising.
• Ar used in fire extinguishing systems in server rooms.
• He used in airships/blimps due to low density.
• Neon lights due to red colour of light produce by
neon.
T6: Percentage by Mass
• This is the percentage of the mass of a compound due to a particular
element.
𝑃𝑒𝑟𝑐𝑒𝑛𝑡𝑎𝑔𝑒 𝑏𝑦 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑎𝑛 𝑒𝑙𝑒𝑚𝑒𝑛𝑡
𝑁𝑜. 𝑜𝑓 𝑎𝑡𝑜𝑚𝑠 𝑜𝑓 𝑒𝑙𝑒𝑚𝑒𝑛𝑡 × 𝐴𝑟 𝑜𝑓 𝑒𝑙𝑒𝑚𝑒𝑛𝑡
=
× 100
𝑀𝑟 𝑜𝑓 𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑
For example: what is the carbon in ethanol, C2H6O?      
Calculate Mr of C2H6O
• Electrons are delocalised, moving
freely between all the atoms
creating a ‘sea of electrons’
• All atoms have a positive charge
as their outer-shell electrons
have left them.
• The bond is the attraction
between the positive ions and
the sea of electrons.
• Conduct electricity as electrons
are free to move.
• Malleable (change shape but
don’t shatter when hit) because
rows of atoms slide past each
other when hit
Mr = (2 x 12) + (6 x 1) +
16 = 46
Number of C in C2H6O
2
Relative atomic mass of
C
12
Percentage by mass of C
% 𝑜𝑓 𝐶 𝑖𝑛 𝐶2𝐻6𝑂 =
2 ×12
× 100 = 52.1%
46
T6: Yield
T6: Relative Masses
• Theoretical yield: the amount of
product you would expect according
to the calculation in the ‘Reacting
Quantities’ box.
• Actual yield: the amount of product
you actually get in practice.
• Percentage yield: the proportion of
the theoretical yield that you actually
achieve.
𝑎𝑐𝑡𝑢𝑎𝑙 𝑦𝑖𝑒𝑙𝑑
% 𝑌𝑖𝑒𝑙𝑑 =
× 100
𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑦𝑖𝑒𝑙𝑑
Relative Atomic Mass, Ar
% yield is always less than 100
because:
Relative Atomic Mass, Ar
• The reaction may be incomplete
• Some product may be lost during the
steps to prepare it.
• Some reactants may also produce
products other than the desired one.
•
•
The mass of atom relative to
the mass of 12C (carbon-12).
For example…    

Element
Relative Mass
Hydrogen, H
1
Carbon, C
12
16
Sodium, Na
23
.
Chlorine, Cl
35.5
Molecular Formula
Empirical Formula
Water, H2O
H2O
Ethane, C2H6
CH3
CH2O
•
•
Relative atomic mass
Divide by relative
atomic mass
Divide both sides by
smallest answer
Empirical formula
• Combining relative masses with balanced equations lets us
work out the masses of chemicals involved in reactions.
• We can use this mathematical relationship:
• m = mass of substance present
• Mr = relative formula mass of
𝑚1
𝑚2
substance
=
𝑀𝑟1 𝑛1 𝑀𝑟2 𝑛2 • n = number of substance in balanced
•
10.0
133.3
12
80
What mass of carbon dioxide can be produced by burning
15g ethene (C2H4) in excess oxygen (O2)?
C2H4 + 3O2  2CO2 + 2H2O
𝑟1 1
𝑟2 2
𝑚1
15
=
44 × 2 28 × 1
1.67/0.83 = 2
𝑚1 =
MgBr2
equation
• 1 refers to the first substance
• 2 refers to the second substance
• Substance 2 will be ethene, substance 1 will be carbon dioxide.
• Calculate relative masses:
• Mr(ethene) = 2 x 12 + 4 x 1 = 28
• Mr(carbon dioxide) = 12 + 2 x 16 = 44
• Then:
𝑚1
𝑚
Write out the equation.
= 𝑀 2𝑛
𝑀 𝑛
10 / 12 = 0.83 133.3 / 80 = 1.67
0.83 / 0.83 = 1
Relative formula mass of sodium chlorate, NaClO3
Mr = Ar(Na) + Ar(Cl) + 3 x Ar(O)
= 23 + 35 + (3 x 16)
Example:
The empirical formula can be calculated from the
masses of substances that react with each other as
below.
For example: 10.0g of magnesium reacts with 133.3 g of
bromine.
Mg
Br
Mass in g
This is the sum of all the relative masses in a
formula.
Relative formula mass of carbon dioxide, CO2:
Mr = Ar(C) + 2 x Ar(O)
= 12 + (2 x 16) = 44
T6: Reacting Quantities (HT)
The lowest whole number ratio of atoms in a molecule.
For example:
Glucose, C6H12O6
•
•
T6: Empirical Formulae
•
•
Relative Formula Mass, Mr
•
Oxygen, O
Effect of Concentration:
•
• Part of exhaust pipe that helps make car exhaust
less environmentally damaging.
Discovery:
Ionic
Solubility
Heat energy is converted to chemical energy.
The surroundings get colder.
Examples: ammonium nitrate dissolving in
water, photosynthesis
Making and Breaking Chemical Bonds
•
Type of
Bonding
T5: Catalytic Converters
• Group 0 in the periodic table.
• Helium ((He, Neon (Ne), Argon (Ar), Krypton (Kr)
Xenon (Xe), Radon (Rn)
• Full outer shells so extremely unreactive: inert.
Note: you increase the surface area by breaking a large piece into many smaller pieces, with
powder being the best.
Chemical Energy
• Reactivity increases down the group:
•
Lithium just fizzes before disappearing
•
Sodium fizzes and gets hot enough to melt
into a ball, occasionally catching fire
•
Potassium fizzes very vigorously, getting hot
enough to burn with a lilac flame
T4: Noble Gases
• The rate of a reaction is its speed, how quickly
products are made.
• Reactions happen when particles collide with each
other.
• Concentration: increasing concentration (the amount
of solute (dissolved stuff) in a given volume) will
increase the rate.
• Temperature: increasing temperature will increase
the rate.
• Surface area: increasing surface area will increase the
rate.
15
× 44 × 2
28 × 1
= 47.1 𝑔
Sub in the numbers
Rearrange to make m1 the subject.