MELTING CURVE
Note! The horizontal points on the boiling curve indicate where heat energy is absorbed
to overcome forces of attraction between the particles to change state from solid to
liquid then liquid to gas, hence, there is no temperature increase at these points. Similarly,
for the melting curve, the horizontal portions indicate where heat energy is released as the
particles move closer together.
© Dhakshenya A. D, 2021
KINETIC PARTICLE THEORY
BOILING CURVE
Heat Energy
Heat Energy
ABSORBED
RELEASED
WHAT'S THE DIFFERENCE BETWEEN
EVAPORATION AND BOILING?
ARRANGEMENT OF ELECTRONS
ISOTOPES
- atoms of the same element with the same
number of protons and electrons but
different number of neutrons
eg. Carbon-12, Carbon-13
- Electrons are arranged around the
nucleus in shells
- shell 1 = up to 2e- shell 2 = up to 8e- shell 3 = up to 8e-
- isotopes have the same chemical
properties as each other (because they have
the same number and arrangement of
electrons) but different physical properties
(as these are determined by an element's
mass - which is different for isotopes.)
- This allocation and numbering should
be followed for elements up to atomic
number 20.
- The outermost shell = valence shell
- Electrons in the valence shell = valence
electrons
Atomic Structure
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ATOMIC STRUCTURE
- the relative atomic mass is the weighted
average of all the isotopes of the element
Protons = electrons
The total number of protons equals the
total number of electrons in an atom of
an element. The negative and positive
charges cancel each other out, leaving
the atom with no overall charge.
Periodic Table
Depending on the periodic table, sometimes
the orientation of mass number and atomic
number may be reversed. It can be helpful to
remember that the bigger number, tends to
be the mass number.
GASES
Measuring gas volume: Gas syringe
To collect the gas evolved in a reaction, the
following methods can be employed:
for gases that are insoluble in
water. e.g. CO2, O2, H2, (Soluble
gases will dissolve in the water.)
Downward Delivery
3.
Upward Delivery
for gases soluble in water and
for gases soluble in water and
denser than air. e.g. Cl2, HCl, SO2
less dense ("lighter") than air.
e.g. NH3
2.
VOLUMES
Beaker - different sizes for different volumes e.g. 10ml, 50ml, 100ml
Measuring Cylinder - accurate to the nearest 0.5. e.g. 80.0 cm3, 55.5 cm3
© Dhakshenya A. D, 2021
EXPERIMENTAL TECHNIQUES
1. Displacement of water
Pipette - for accurate measurement of fixed volumes e.g. 20.0 cm 3/ 25.0 cm 3
Burette - accurate to the nearest 0.05. e.g. 37.00 cm3, 58.65 cm3
Note! Always read from the bottom of the meniscus
Increasing Accuracy
G E N E R AL AP P AR AT U S AN D M E AS U R E M E N T S
Mass - beam balance or electronic balance (g/kg)
Time - stopwatch (secs/mins)
Temperature - thermometer or temperature sensor and data logger (degrees
Celsius oC / Kelvin)
1.Evaporation to dryness
eg. Salt solution
- Heat the solution until all the liquid evaporates, leaving the soluble solid
behind
*If the solute in the solution will thermally decompose upon heating to
dryness, you can use another method:
2.Crystallisation
eg. Sugar solution
- Heat the solution till saturated (dip a cool glass rod into the solution
and if crystals form on the rod when you remove it from the solution, it is
saturated)
- Leave to cool and crystals of the soluble solid start to form
- Filter then pat dry the crystals with filter paper
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SEPARATION & PURIFICATION
SOLUBLE SOLID FROM SOLUTION
INSOLUBLE SOLID
FROM SOLUTION
IMMISCIBLE
LIQUIDS
Filtration
Separating funnel
MISCIBLE LIQUIDS
Fractional Distillation
eg. Ethanol (bp 78 degrees celsius) and water
(bp 100 degrees celsius)
-The liquid with the lower boiling point (bp) will
boil off first, be condensed and collected in the
conical flask at the end. Collect all the distillate.
- As you keep heating, the temperature will
continue to rise. At the boiling point of the next
liquid, replace the conical flask to collect the next
distillate.
LIQUID FROM SOLUTION
Simple Distillation
eg. Salt solution, where you wish to collect the
water (ie. you do not want to evaporate the water
away)
-The setup is identical to the diagram shown, less
the fractionating column and glass beads. Since
there is only one liquid of interest, as soon as the
temperature reaches its boiling point, the liquid
will evaporate, condense and be collected in the
receiving flask at the end.
CHROMATOGRAPHY
- To separate dissolved components in a solution (eg. dyes in ink)
- To identify banned substances (eg. in food)
- To separate amino acids (which come from proteins)
- To determine the purity of a sample
- Substances which travel further are more soluble in the solvent
- If it doesn't travel at all, it may be insoluble in the solvent or
colourless in which case you'd need to apply a locating agent.
ELEMENTS
Monoatomic eg.
Neon atoms Ne
Definition: The simplest form of a pure substance, that cannot be broken
down into 2 or more simpler substances by chemical processes
Classifications: Elements can be classified as metals, metalloids, nonmetals, or by physical state at room temperature and pressure (solids,
liquids, gases)
Elements can also exist as atoms or molecules
Atoms are the smallest particles of an element
Molecules are made up of 2 or more atoms of an element chemically
combined (2 atoms = diatomic, more than 2 = polyatomic)
Ionic compound
e.g. Sodium Chloride
NaCl
Definition: a pure substance containing two or more
elements chemically combined in a fixed ratio
Compounds can exist as:
molecules (eg. covalent compounds like CO2) or ionic
compounds (NaCl)
Polyatomic
eg. Ozone
molecules O3
Covalent compound
e.g. Carbon dioxide CO2
M IXT U R E S
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ELEMENTS, COMPOUNDS &
MIXTURES
COMPOUNDS
Diatomic eg.
Nitrogen
molecules N2
Definition: An impure substance made up of two or more
substances not chemically combined (ie. physically combined) in
no fixed ratio. eg. Sand and metal fillings
eg. Air - Oxygen and
nitrogen (amongst other
gases)
eg. Alloy: Brass mixture of copper
and zinc
Definition: electrostatic forces of attraction between ions of oppositely
charged ions ie. positive (cations) and negative (anions) ions.
eg. the force of attraction tightly holding each sodium ion to chloride ion
is an ionic bond.
>> Ionic bonds are formed by the transfer of electrons from a metal to a
non-metal.
Positive ions are formed by atoms (mostly metals) losing electrons
eg. A sodium atom loses one electron to form a sodium ion.
Negative ions are formed when atoms (mostly non-metals) gain electrons
eg. A chlorine atom gains one electron to form a chloride ion.
>> Ions have a full outer shell of electrons and noble gas electronic
configuration. (see below)
COVALENT BONDS
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CHEMICAL BONDING
IONIC BONDS
Definition: a bond formed by
the sharing of electrons
between two non-metals.
Covalent bonds can form
between atoms of the same
element (eg. Oxygen molecule)
or atoms of different elements
(eg. Carbon Dioxide)
Ionic compound
eg. Sodium Chloride
NaCl
METALLIC BONDS*
Definition: electrostatic
forces of attraction
between positive metal
ions and a sea of
delocalised electrons.
PROPERTIES OF IONIC & COVALENT COMPOUNDS
WH AT IS T H E M O L E ?
Think of it as a unit of measurement. For example, 100cm in 1m, 1000g in 1kg, 6x1023
particles in 1 mole.
Here, particles can be atoms, molecules or ions. 6x1023 is known as Avogadro's constant.
Relative atomic mass (Ar): This doesn't refer to the absolute mass of an element because the actual
mass is extremely small. Instead, this refers to the average mass of an atom of an element compared
to 1/12 the mass of a carbon-12 atom.
The relative atomic mass of an element can be found in the periodic table - mass number of the
element. eg. Ar of Magnesium = 24. Ar of Zinc = 65. No units.
Relative molecular mass (M r): This refers to the total of the relative masses of all the elements in a
molecule. In other words, the average mass of one molecule of that element or compound compared to
1/12 the mass of a carbon-12 atom. No units.
The way to calculate this is to total up the individual relative atomic masses of the elements in the
compound or molecular element. eg. M r of O2 = (16x2) = 32. Mr of NH3 = 14 + (1x3) = 17.
Relative formula mass: Has the same definition as relative molecular mass but is used for ionic
compounds - which are not molecules. eg. Mr of NaOH = 23+16+1 = 40.
Molar mass: This refers to the mass of one mole of something (atom/molecule/compound) and is
calculated in the same way as the relative molecular mass but it has the units g/mol. eg. Molar mass
of carbon = 12g/mol. Molar mass of NaCl = 23 + 35.5 = 58.5g/mol.
© Dhakshenya A. D, 2021
THE MOLE
S O M E D E F IN IT IO N S
eg. How many atoms in 2 moles of Argon?
2 moles x 6x1023 = 1.2 x 1024 atoms
*Note! The calculation is the same even if you were
looking for the number of particles in 2 moles of Zinc,
carbon, hydrogen etc. The number of particles in 1
mole of anything is 6x1023.
eg. How many moles in 25g of Magnesium?
Refer to the periodic table to find the molar mass of
magnesium = 24g/mol.
No. of moles = 25 (mass)/24 (molar mass) = 1.04 mol
*Note! Be careful with the units when it comes to these
questions. Sometimes the mass may be in Kg. You must
remember to change it to grams to use this formula
triangle correctly.
eg. How many moles in 250cm3 of 2 mol/dm3 of HCl?
Note the difference in units here. Convert 250 cm 3 to
0.25dm3.
No. of moles = 0.25 dm3 x 2 mol/dm3 = 0.5 moles
M O L AR VO L U M E O F G AS E S
1 mol of any gas occupies a volume of 24 dm3. 1dm3 is equivalent to 1000 cm3
eg. What is the volume (in dm3) of 8g of oxygen gas (O2)?
You would first need to use the mass and molar mass equation above to calculate the
number of moles of 8g of Oxygen which is 8/(16x2 for O2) = 8/32 = 0.25 mol. 0.25 x 24 dm3 =
6 dm3.
*Be very careful with the units. If they ask for the answer in cm3, convert the 6 dm3 to 6000 cm3.
SOME DEFINITIONS
ACIDS: substances which ionise (form ions) in
solution to form hydrogen (H+) ions.
BASES: metal oxides or hydroxides which
react with acids to form salt and water only
ALKALIS: soluble bases which ionise/dissolve
in solution to form hydroxide (OH-) ions.
PROPERTIES
© Dhakshenya A. D, 2021
ACIDS & BASES
Acids
- sour taste
- conducts electricity when dissolved in
solution (enabled by the presence of ions)
- turns blue litmus paper red
Reactions:
Acid + reactive metal --> salt + hydrogen
*Note: Acids do not react with unreactive
metals like copper, silver and gold
Acid + carbonate --> salt + water + CO2
*Note: recall the limewater test to test for
carbon dioxide in this reaction
Acid + metal oxide/hydroxide --> salt
and water
*Note: this is a neutralisation reaction, reaction
of acid with a base to produce salt (neutral, pH
7) and water.
Bases
- bitter taste
- soapy texture
- turns red litmus paper blue
- also conduct electricity when dissolved
Reactions:
Alkali + acid --> salt + water
Alkali + ammonium salt --> salt + water +
ammonia gas
*Note: recall the test for ammonia gas, the gas
turns red litmus blue
*The acidity of soil can be controlled by
treating with bases - quicklime (calcium oxide)
or slaked lime (calcium hydroxide). This
process is called "Liming" the soil. pH of soil:
between 5-9.
Importance of pH
- What happens when the pH deviates
too far from the norm?
Blood has a slightly alkaline pH of 7.4.
When medication or vaccinations are
injected into the body, the pH needs to be
accounted for. Deviating too far from this
pH of 7.4 will affect enzyme activity and
other chemical reactions, causing harm to
the body.
N
O
N
M
E
T
A
L
L
I
C
ACIDIC OXIDES - react with alkalis to
produce salt + water
- dissolve in water to produce acids
eg. Carbon dioxide --> carbonic acid
Sulfur trioxide --> sulfuric acid
NEUTRAL OXIDES - neither acidic nor
basic, so don't react with acids or bases
- examples: Water, NO, CO
M
E
T
A
L
L
I
C
BASIC OXIDES - react with acids to
produce salt + water
- dissolve in water to produce alkalis
eg. CaO--dissolve--> Calcium Hydroxide
AMPHOTERIC OXIDES - react with
both acids and bases
- zinc, lead and aluminium oxide
WHAT ARE SALTS?
Ionic compounds made up of a cation (+ve ion) and anion (-ve ion).
A salt is produced when...
an acid reacts with: (1) metal, (2) carbonate, (3) base or (4) alkali.
(1) Acid + metal --> salt + hydrogen gas
(2) Acid + carbonate --> salt + water + carbon dioxide
(3) Acid + base --> salt + water
(4) Acid + alkali --> salt + water
Hydrated salts contain water of crytallisation. They exist as salt crystals.
Anhydrous salts do not contain any water of crystallisation. They exist as powders.
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SALTS
Why does solubility matter?
It is important to know whether a salt is
soluble or insoluble when you want to
prepare a salt from scratch.
All SPAN salts are soluble. SPAN? Sodium,
Potassium, Ammonium and Nitrates.
Bearing this in mind, all carbonates are
INSOLUBLE, except for the SPA salts
(sodium, potassium, ammonium)
All sulfates are soluble except for CLuB
(Calcium, Lead and Barium)
All chlorides are soluble except for two
(lead and silver)
METHODS OF SALT PREPARATION
The solubility will determine which method of preparation you can use.
There are three methods of preparation:
(1) Precipitation
(2) Titration
(3) Acid + excess insoluble solid
(-- this solid can be a metal, metal carbonate or metal oxide/hydroxide)
Titration
Precipitation
- to prepare soluble salts
- to prepare insoluble salts
- starting materials must both - starting materials must be both
soluble solutions
be soluble solutions
- used for the preparation of SPA
salts
(1) Mix two solutions one
containing the cation, and one (1) Perform the titration with
acid in the burette and base in
containing the anion for the
the conical flask (with indicator).
salt you wish to prepare.
(2) Record the volume of
(2) Filter the insoluble salt
solution needed for
out of the solution
neutralisation.
(3) Dry crystals with filter
(3) Mix the now known volumes
paper
of the two solutions together
(4) Crystallise the solution
eg. Barium sulfate.
Mix barium nitrate (soluble - (5) Filter to collect crystals.
supplies the cation) with
sodium sulfate (soluble supplies the anion), to
prepare Barium Sulfate.
eg. Nitric acid + Sodium
hydroxide to prepare Sodium
Nitrate
Acid + excess insoluble solid
- to prepare soluble salts which
are NOT SPA salts.
- starting materials must be
one acid and one insoluble
solid (eg. metal, carbonate or
metal oxide/hydroxide)
(1) Mix acid with excess solid
(2) Filter to remove excess
solid
(3) Crystallise to prepare salt
crystals
(4) Filter the mixture and dry
the crystals.
eg. Copper (II) nitrate.
Mix nitric acid with excess
copper (II) oxide (which is an
insoluble solid) and continue
from step 2 above to prepare
Copper(II) Nitrate crystals.
GAS TESTS
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SALTS
CHEMICAL
TESTS FOR
CATIONS
- POSITIVE
IONS
CHEMICAL
TESTS
FOR ANIONS
- NEGATIVE
IONS
Oxidation
- gain of oxygen
- loss of hydrogen
- loss of electrons
(positive ion forms)
Reduction
- loss of oxygen
- gain of hydrogen
- gain of electrons
(negative ion forms)
- increase in
oxidation state
(see below)
- decrease in oxidation
state
(see below)
OXIDATION STATES
Oxidation state refers to the charge an atom of an element would have if it existed in a compound as
an ion (even covalently bonded atoms)
Rule 1: the oxidation state of any element is 0
eg. Mg, C, He
Rule 2: oxidsation state of an ion is the same as the charge of the ion
© Dhakshenya A. D, 2021
OXIDATION & REDUCTION
OXIDATION VERSUS REDUCTION
Rule 3: Some elements have fixed oxidation states
eg. all group 1 elements have oxidation state of +1, group 2 elements (+2), hydrogen is almost always
(+1), oxygen is almost always (2-)
Rule 4: Oxidation states of elements in a compound add up to zero
eg. CaO (2+, 2-)
Rule 5: Oxidation states of the atoms in a polyatomic ion add up to the charge of the ion
eg.
(oxygen has an oxidation state of -2. In manganate there are 4 O, 4x(-2) = -8. Overall
charge of the ion is -1, hence the oxidation state of manganese must be +7. -8 + 7 = -1.)
eg. OH- (O = 2-, H = 1+, 1-2 = -1 which is the charge of the ion).
An INCREASE in oxidation state indicates OXIDATION
A DECREASE in oxidation state indicates REDUCTION (see below for examples)
REDUCING & OXIDISING AGENTS
Test for oxidising agents
*oxidising agents are themselves REDUCED
Test: add aqueous potassium iodide (KI) to the
test solution. If a brown solution forms, the
test solution contains an oxidising agent.
The iodide ions in KI are oxidised (lose
electrons) to form iodine - which is where the
brown colour comes from
Test for reducing agents
*reducing agents are themselves OXIDISED
Test for gases: place filter paper soaked in
acidified potassium manganate (VII) at the
mouth of the test tube where the gas is
Test for solutions: add acidified potassium
manganate (VII) to the test solution
Positive result: the potassium manganate
changes colour from purple (manganate
ion) to colourless (manganese 2+ ion).
The oxidation state of iodine increases from -1
in iodide to 0 in iodine = OXIDATION.
The oxidation state of manganese decreases
from +7 to +2 in = REDUCTION.
PHYSICAL PROPERTIES
(1) Good conductors of heat and electricity
- the free electrons in the metallic structure are able to move through
the metal to carry charge/heat
(2) High densities, melting and boiling points
- (with a few exceptions eg. Alkali metals have low mp and bp.
Mercury is a liquid at room temperature)
- metallic atoms are tightly packed and held together by strong forces
of attraction = metallic bonds which require plenty of heat energy to
overcome. This contributes to the high mp, bp and density.
(3) Malleable (flexible) and ductile (can be drawn into wires)
- regular arrangement of atoms means the layers can slide past each
other, hence, the metal can be bent into different shapes
© Dhakshenya A. D, 2021
METALS
ALLOYS
What are they?
A mixture of a metal with one or more elements
Why do we need them?
sometimes metals have unfavourable properties we would like
to change for some usage application.
Pure metals are soft (see malleable + ductile property above). -- Alloys enable us to have a stronger & harder alternative
Pure metals may react with air and water and corrode easily. --- Alloys allow us to have alternatives which are less corrosive +
reactive
Examples:
- Brass - alloy of copper and zinc mixed together. It is stronger
than the individual metals
- Pewter - alloy of tin, antimony and copper. The mixture means
the alloy looks prettier than pure tin and can be used to make
ornaments
- Bronze - alloy of copper and tin.
Dark blue circles =
element added to the
metal (colourless circles).
No more regular layers of
atoms (like in metals).
The layers of atoms now
can't slide past each
other = no longer
malleable and ductile.
The alloy is harder than
the pure metal
REACTIVITY SERIES
- ARRANGEMENT OF METALS FROM MOST TO LEAST
Please
REACTIVE
Potassium
Sodium
Calcium
Magnesium
Aluminium
Zinc
Iron
Lead
Copper
Silver
Gold
Here's how you can remember it!
Stop
Calling
My
Adorable
Zebra
In
Language
Class
She
Grunts
METHODS OF EXTRACTION
More reactive metals (In Red above) are extracted by Electrolysis - using electricity to
extract the pure metal from the metal ore (metal + impurities).
Less reactive metals (in blue above) are extracted by reduction of their oxides with carbon
(eg. Iron via the reactions in the blast furnace)
REACTIONS WITH COLD WATER,
STEAM AND ACID
The nature of the reaction indicates how
reactive the metal is. Metals (in the table) are
arranged from most (top) to least (bottom)
reactive.
-- = no reaction
E = reacts explosively
V = reacts violently
R = reacts readily
M = moderately
S = reacts slowly
Iron ore = Haematite = made up of iron oxide, sand (silicon dioxide) and clay.
The aim here is to extract pure iron from the Haematite. A series of reactions occur in the
blast furnace to do this.
(1) Carbon (coke) reacts with oxygen
to produce --> Carbon dioxide
(2) Carbon dioxide reacts with more
carbon to produce --> Carbon
Monoxide
© Dhakshenya A. D, 2021
METALS
BLAST FURNACE - EXTRACTION OF IRON FROM ITS ORE
(1) Haematite
(iron ore)
(2) Coke
(mostly carbon)
(3) Limestone
(Calcium
carbonate)
(3) Carbon monoxide reduces iron
(III) oxide in haematite to produce -->
(molten) iron + carbon dioxide
To remove impurities:
(4) Limestone is decomposed
under the heat to produce -->
calcium oxide & carbon dioxide
(5) Calcium oxide reacts with the
silicon dioxide (sand) --> Molten slag
RUSTING
What is rust?
It is the corrosion of iron. In other words, the
reaction of iron with oxygen in the presence of
water.
IRON + OXYGEN + WATER ==>
HYDRATED IRON (III) OXIDE
Rust prevention methods provide a protective
layer over the iron, preventing it from coming
into contact with air and water.
Some methods:
- Painting - layer of paint over the iron (eg. cars)
- Oiling & Greasing - protective layer of
oil/grease
- Plastic coating - protective plastic layer
- Electroplating - coating iron with another
metal (eg. chrome, tin, copper) using electricity
- Galvanising - coating iron with zinc (a more
reactive metal). Sacrificial method (sacrificing
the more reactive metal, in this case - Zinc).
(1) Molten iron
(pure iron)
(2) Waste Gases
(Carbon
monoxide,
carbon dioxide &
nitrogen)
(3) Molten slag
(Calcium silicate)
RECYCLING
Metals are a finite resource - they will run out
eventually, hence, recycling is necessary
Conserve natural resources - extracting
metals from ores requires burning of
fossil fuels. Recycling means we save this
fuel
Reduce environmental problems extracting from ores requires large plots
of land and generates large amounts of
waste
Recycling saves cost associated with
extracting metals from their ores
Some issues/disadvantages of recycling
-- Recycling can be expensive (sorting,
cleaning and transporting the scrap metal)
-- Social issues (recycling is only effective if
large communities practice it)
-- Environmental issues (recycling sometimes
creates pollution problems)
ARRANGEMENT OF ELEMENTS IN THE PERIODIC TABLE
- Elements are arranged according to increasing proton number in the periodic table
- Vertical Columns are GROUPS - Check the numbering in the periodic table. There are 8
groups. The block of elements in the middle don't belong to any group. These are called
TRANSITION METALS.
- Horizontal Rows are PERIODS - Hydrogen is in Period 1
Elements in the same group:
- have the same number of valance electrons (group number indicates no. of valence electrons)
- have similar chemical properties
GROUP 1 ELEMENTS
- ALKALI METALS
© Dhakshenya A. D, 2021
PERIODIC TABLE
Moving from left to right across the periodic table: ------------->>>>
- change in properties from metals to metalloids to non-metals
- metalloids are the elements touching the zig-zag line on the periodic table (except Al)
- elements in the same period have the same number of electron shells (period number =
number of electron shells).
- change in type of oxide from basic (to amphoteric) to acidic
Physical properties
- soft & can be cut with a knife
(tarnishes when freshly cut)
- low melting and boiling points
(compared to general metals)
- low densities
(lithium, sodium & potassium float on water)
Chemical properties
- alkali metals react with water:
alkali metal + water --> alkali + hydrogen
- they are powerful reducing agents
(very readily lose their valence electron to
form ions with charge +1)
- react with non-metals to form ionic
compounds
Trends down the group
_ reactivity increases
(lithium< sodium< potassium etc.)
- melting and boiling points decrease
- densities increase
GROUP 7 ELEMENTS
- HALOGENS
Physical properties
- diatomic covalent molecules
(they exist in pairs F2, Cl2)
- low melting and boiling points
- they are coloured
Chemical properties
- undergo displacement reactions
more reactive halogen will displace a less
reactive halogen from a halide solution:
chlorine+ sodium bromide --> bromine +
sodium chloride
- they are powerful oxidising agents
(very readily gain an electron to form ions
with a full outer shell and charge of -1)
Trends down the group
_ reactivity decreases (fluorine is the most
reactive)
- melting and boiling points increase
- densities increase
(fluorine gas, bromine liquid, iodine solid)
- colour intensity increases (gets darker)
GROUP 0 ELEMENTS - NOBLE GASES
Properties
- monoatomic (exist as lone atoms)
- unreactive (because they have a full outer
shell of electrons. Hence, do not need to
lose, gain or share electrons)
- low melting and boiling points
- gases at room temperature
- colourless
- insoluble in water
Uses
- Helium: to fill up balloons
- Neon: for making bright lights and
advertisement signs
- Argon: for welding stainless steel & to fill
tungsten bulbs
ENERGY CHANGES
Energy can neither be created nor destroyed. It can only be converted from one form to
another.
In reactions where heat energy is released to the surroundings, this is called an EXOTHERMIC
change (exo- like exit)
Bonds are always being broken down in the reactants and re-joined to form the products.
Bond breaking is endothermic (absorbs energy) and bond formation is exothermic (releases
energy). So, these processes are happening all the time. When MORE energy is absorbed than
released, the reaction is considered endothermic. When MORE energy is released than
absorbed, the reaction is exothermic.
© Dhakshenya A. D, 2021
ENERGY CHANGES
In reactions where heat energy is absorbed into the reaction mixture, it is called an
ENDOTHERMIC change.
*Neutralisation - is an exothermic reaction
eg. Sodium hydroxide + Hydrochloric acid --> Sodium chloride + Water
The ionic equation for the reaction is:
When the hydrogen (H+) and hydroxide (OH-) ions combine to form water, heat energy is
released. There is a rise in temperature and the reaction mixture feels warm.
*Combustion of fuels - fuels burn easily in air to release energy. This energy can be used
to power cars, for cooking, insulating or lighting homes. The most commonly used fossil
fuels are wood, coal, petroleum and natural gas (mostly methane).
The combustion of fuels can be complete:
methane + oxygen --> carbon dioxide + water vapour + heat energy
or incomplete (in limited air supply):
methane + oxygen --> carbon monoxide + water vapour + carbon + heat energy
Either way, all combustion reactions release heat energy hence, it is an exothermic change.
*Thermal decomposition - requires heating of the substance to be thermally
decomposed. It absorbs heat energy and decomposes. ENDOTHERMIC. Eg. Calcium
carbonate decomposes to calcium oxide and carbon dioxide.
*Sodium hydrogen carbonate - reacts with acids in an ENDOthermic reaction
Sodium carbonate - reacts with acids in an EXOthermic reaction
DO ALL REACTIONS HAPPEN AT THE SAME SPEED?
Fast Reactions - explosion of petrol in air
Moderately fast reactions - browning of a
cut apple
Slow reactions - rusting of iron
HOW CAN WE DETERMINE THE SPEED OF A REACTION?
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SPEED OF REACTION
1. Measure the amount of product produced in a given amount of time
2. Measure the amount of reactant used up in a given amount of time
Measure volume of gas produced
If a gas is produced in a reaction, (1) collect the
gas using a delivery tube and gas syringe, (2)
note down the volume produced at regular time
intervals until the reaction is complete.
Plot a graph of time against volume of gas
produced to calculate the speed of reaction at
any point in the reaction.
Total volume of gas produced/total time taken =
gives the average speed of reaction.
To find the speed of reaction at a specific time
point, draw a tangent to the graph and calculate
the gradient. This will reveal the speed of
reaction.
Measure mass of reactant used up
Alternatively, the loss in mass of the reactant can
be measured using an electronic balance, at
regular time intervals.
Plot a graph as described above.
Total loss in mass/total time = gives the average
speed of the reaction.
*A quick indicator of reaction speed is
the steepness of a curve. The steeper
the curve, the faster the speed of the
reaction.
FACTORS AFFECTING THE SPEED OF REACTION
What exactly are we talking about when we mention reaction speed? What makes a
successful reaction?
Effective collisions between reactant particles. The more effective collisions occur, the higher the
reaction speed.
What is an effective collision?
One where the reactant particles collide with sufficient energy - this is called the activation energy.
ACTIVATION ENERGY - the minimum amount of energy needed by reactant particles for a reaction
to happen when they collide.
- concentration: a higher concentration of reactants causes a
rise in reaction speed. Why? Higher concentration means more
reactant particles in a given volume hence, a higher likelihood
of effective collisions.
- pressure: similarly, a higher pressure means more reactant
particles in a smaller volume, hence, a higher likelihood of
effective collisions.
- particle size: a smaller particle size means the surface area is
greater which results in more exposed surfaces for the reactant
particles to collide with hence, a higher speed of reaction.
- temperature: increasing temperature causes particles to
move faster (higher kinetic energy) so they collide with each
other more frequently and with sufficient/more energy resulting
in a higher speed of reaction.
COMPOSITION OF AIR
*Liquid air can be
separated via
fractional distillation
into its constituent
gases. Liquid nitrogen,
for example, can be
obtained and used to
store items are very low
temperatures.
Air pollution
- refers to a state where the air contains high concentrations of chemicals which can harm
living things (humans/animals/plants) or damage non-living things (like buildings).
COMMON AIR POLLUTANTS
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AIR & ATMOSPHERE
*The noble
gases
predominantly
comprise
ARGON
Carbon monoxide
Sulfur dioxide
Oxides of nitrogen
Methane - formed when plant and animal matter decay. It is the main constituent of
natural gas - a fossil fuel. Also released by sheep and cows due to digestion of food.
Ozone - formed when nitrogen dioxide in the air reacts with unburnt hydrocarbons in
the presence of sunlight
Unburnt hydrocarbons - from car exhaust fumes
Acid Rain
- when pollutants like sulfur dioxide and nitrogen dioxide dissolve in the water in the
atmosphere, acid rain forms (sulfuric acid/nitric acid).
*sulfur dioxide + water --> sulfurous acid --> oxidises in the air to sulfuric acid
*nitrogen dioxide + water + oxygen --> nitric acid
- acid rain roughly has a pH of 4
- normal unpolluted rain water naturally has a slightly acidic pH (slightly below 7) because
carbon dioxide in the air dissolves in water vapour to form carbonic acid which forms rainwater
that is weakly acidic.
Harmful effects of acid rain:
- reacts with and corrodes limestone buildings or metals structures like bridges
- the acid rain which falls on the ground gets absorbed by the soil and causes plants to wither
and die
- acid rain lowers the pH of lakes and streams, killing fish and aquatic life
Organic compounds - mostly refer to compounds containing carbon. Most organic
compounds also contain hydrogen
Hydrocarbons - compounds that ONLY contain carbon and hydrogen
Functional group
- an atom or a group of atoms which give a molecule a defining factor.
eg. For Alkenes the defining factor is a carbon-carbon (C=C) double bond
Alcohols = oxygen-hydrogen (O-H) component = called a hydroxyl group
Carboxylic acids = COOH component = carboxyl group
NAMING ORGANIC
COMPOUNDS
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ORGANIC CHEMISTRY
Homologous series
- a group of organic compounds with the same functional group and similar chemical
properties, gradual change in physical properties when moving through the group
Petroleum (or crude oil) is a mixture of hydrocarbons and like natural gas (mostly
methane) is a fossil fuel.
How did these fossil fuels form?
1. Dead sea creatures and plants sank to the seabed millions of years ago.
2. They got covered by mud and sand
3. Heat and pressure from the earth was exerted on these dead organisms
4. Over millions of years, petroleum and natural gas formed
5. Suppressed by heat and pressure, the fossil fuels are hundreds of metres below
the surface of the earth and deep wells need to be drilled to reach and extract them.
Petroleum contains a mixture of hydrocarbons which can be separated into more useful fractions
via Fractional Distillation:
1.Petroleum is heated in a furnace till it
becomes a vapour which is passed
through a fractionating column (image)
2.The hot vapour first rises and then
begins to cool and condense at
different levels in the column
3.Those with shorter carbon chains are
lighter, have lower boiling points and
are collected closer to the top as
GASES
4.Those hydrocarbons with longer
carbon chains, are heavier, have higher
boiling points and are collected at the
lower part of the column
C1 to C4
Petroleum gas - for cooking and heating
C5 to C10
Petrol - fuel for cars
C7 to C14
Naphtha - raw materials for detergents
C9 to C16
Kerosene - fuel for airplanes
C15 to C25
Diesel oil - for diesel engines (bus/lorry)
C20 to C35
Lubricating oil - lubricating machines;
making wax/polish
>C70
Bitumen - making road & roof surfaces
*Homologous series - Alkanes
*Functional group = C-C single bonds*
Definition - alkanes are a group of
hydrocarbons that only contain carboncarbon single bonds.
Alkanes are SATURATED hydrocarbons each carbon atom has four valence
electrons, so it can bond with up to 4 other
atoms. Since alkanes are made up solely of
single bonds, each carbon atom forms the
maximum number of bonds, with four other
atoms - hence the molecule is saturated. It
cannot bond with any further atoms.
Physical properties
low melting and boiling points
insoluble in water
soluble in organic solvents
Trends going down the series
(ie. as carbon chain length increases
meth, eth, prop etc.)
melting and boiling points increase
they become more viscous (thick)
they become less flammable
Chemical Properties
Alkanes are generally UNreactive because their C-C and C-H are so strong they are hard to break.
But, alkanes will engage in these 2 types of reactions:
(1) Combustion (exothermic reaction, makes alkanes good fuels)
alkane + oxygen --> carbon dioxide + water vapour
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ALKALNES & ALKENES
ALKANES
(2) Substitution reactions (alkanes react with halogens - group 7 elements - in the
presence of UV (ultraviolet) light)
methane + chlorine --> chloromethane + hydrogen chloride
*A hydrogen in methane is replaced or 'substituted' with a chlorine = chloromethane.
ALKENES
*Homologous series - Alkenes
*Functional group = C=C double bond
Definition - a group of hydrocarbons that
contain one or more carbon-carbon
DOUBLE bonds
Alkenes are UNSATURATED
hydrocarbons - the double bond is formed
by the sharing of TWO pairs of electrons. It
can be broken to form more single bonds.
Physical properties
ethene, propene and butene at gases at
room temperature
Trends going down the series
(ie. as carbon chain length increases
meth, eth, prop etc.)
melting and boiling points increase
Chemical Properties
(1) Combustion
alkene + oxygen --> carbon dioxide + water vapour
(2) Addition reactions - the double bonds in an unsaturated hydrocarbon are broken when it reacts
with another substance to form a new compound with single bonds. Hence, an unsaturated
hydrocarbon becomes a saturated hydrocarbon.
- (a) hydrogenation - addition of hydrogen to an alkene
alkene + hydrogen --in the presence of nickel, 200 degrees celsius ----> alkane
- (b) bromination - addition of bromine to an alkene. The reddish-brown bromine solution is
decolourised in this reaction. It is a good test to distinguish between an alkane and alkene.
alkene + bromine --> bromoalkane
- (c) polymerisation - joining of a repeat unit molecule - alkene - together in a long chain
CRACKING
a process which involves breaking down long chain hydrocarbons from petroleum
into smaller molecules
eg. Hexane can be broken down into butane and ethene:
Long chain alkanes ---> [mixture of short chain ALKENES] +
[mixture of short chain ALKANES or hydrogen gas]
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ALKALNES & ALKENES
The process can be sped up with a catalyst.
In industries, cracking is carried out at high temperatures (600 degrees celsius), with a
petroleum fraction containing long chain hydrocarbons (eg. naphtha) and is passed over a
catalyst (usually Aluminium Oxide and Silicon dioxide ). Here is what happens:
Why it this important?
1. This is a process which can generate short chain alkenes eg. ethene which is a
necessary starting material to make ethanol.
Long chain alkane --> shorter chain alkane + short chain ALKENE
2. Cracking can help to produce hydrogen. Hydrogen is an integral component when
manufacturing ammonia which is required to make fertilisers
C19 chain alkane ---> C9 chain alkane + C10 chain alkene + HYDROGEN gas
3. Cracking is useful in the production of petrol. More than half the fractions refined from
petroleum are heavy and less useful (eg. lubricating oil). These can be converted by
catalytic cracking to shorter chain fractions such as Petrol, which are in much higher
demand.
Lubricating oil (long chain alkane) ----cracking----> PETROL (short chain alkane)
ALKANES VS ALKENES
Definition: a homologous series of organic compounds with an -OH (hydroxyl) functional
group
Physical properties
liquids at room temperature and pressure
methanol, ethanol, propanol - soluble in water (butanol is only slightly soluble)
relatively low boiling points (ethanol bp 78 degrees celsius)
Trends going down the series (ie. as carbon chain length increases)
Solubility decreases
Boiling point increases
Chemical Properties
(1) Combustion - (like alkanes and alkenes)
Alcohol + oxygen --> carbon dioxide + water vapour
(2) Oxidation - when an alcohol is heated with an oxidising agent such as acidified
potassium manganate (VII)
Alcohol + oxidising agent ---heat---> carboxylic acid + water
eg. ethanol + oxygen from oxidising agent ---heat---> ethanoic acid + water
Production of ethanol by Fermentation
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ALCOHOLS &
CARBOXYLIC ACIDS
ALCOHOLS
- fermentation refers the action of micro-organisms (like yeast) on carbohydrates (like
glucose), in the absence of oxygen to produce an alcohol and carbon dioxide gas.
glucose solution ---yeast, no oxygen----> ethanol + carbon dioxide
1.
Glucose solution is mixed with yeast and kept at 37 degrees Celsius - this is the
optimum temperature for enzyme activity in this fermentation reaction
2. As carbon dioxide is produced, frothing can be observed during fermentation. (The
gas also produces a white precipitate when passed through limewater solution.)
3. A dilute solution of about 15% ethanol solution is produced. (if the alcohol content
exceeds this level, the yeast will be unable to survive and fermentation halts.)
4. Ethanol is extracted from the solution via Fractional Distillation (check separation
techniques for more details)
Carboxylic acids
- a homologous series of organic acids which have a -COOH (carboxyl) group
Properties - same typical properties as acids
- react with reactive metals to produce a salt and hydrogen gas
- react with metal carbonates to form a salt, water and carbon dioxide
-react with bases to form a salt and water
Production of ethanoic acid
- it is produced by the oxidation of ethanol. There are two ways this can be done:
1.Oxidation by acidified potassium manganate
When ethanol is heated with acidified potassium manganate (VII), potassium manganate
acts as the oxidising agent and is reduced. (There is a colour change from purple to
colourless). Ethanol is the reducing agent and is itself oxidised - to form Ethanoic acid.
Ethanol + Oxygen (from the oxidising agent) ---heat---> Ethanoic acid + water
2.Oxidation by atmospheric oxygen
Bacteria in the air (acetobacter) uses atmospheric oxygen around us to oxidise ethanol to
form a weak solution of ethanoic acid.
Ethanol + Oxygen (from the air) ---bacteria----> Ethanoic acid + water
Summary of organic compounds
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ALCOHOLS &
CARBOXYLIC ACIDS
Molecular formula - tells you the number of atoms for each element
Structural formula - tells you how the atoms are structurally grouped together in the
compound
Displayed formula - visual representation of the structural formula
Methane
Propanol
Ethane
Ethene
Propanoic Acid
*There's no
Methene because
the functional group
for alkenes is the
C=C double bond.
Hence, the first
member of the
series must have 2
carbon atoms
hence, ethene. The
"meth" prefix is
associated with a 1
carbon chain