UNIT 1. States of matter
Matter is any substance that occupies physical space. The kinetic theory of matter states that matter is
made of tiny particles (i.e. atoms and molecules) and that they are always in constant motion.
There are three states of matter that you need to be aware of: Solids, liquids, and gases.
Solids have particles that are packed closely together. The atoms are arranged in a regular pattern due
to the strong attractive forces that exists between the particles. These particles therefore cannot move.
Instead, they vibrate constantly on the spot.
By giving a solid energy (heating), the particles will begin to vibrate more vigorously as they gain more
energy. Eventually, the particles begin to separate as they start to overcome attractive bonds, the
particles will have enough separation to move past one another, and become a liquid.
Liquids therefore have particles that are randomly arranged (i.e. can move past one another). They
therefore take up the shape of the container. The particles themselves are constantly moving in constant,
random motion.
By further heating the liquid, the particles gain even more energy. This will separate the particles even
more as they overcome most of the remaining forces of attraction that exists.
Gases therefore have particles that are very far apart. Again, the particles will be moving in constant,
random motion and also take up the shape of its container. Unlike solids and liquids, gases can be
compressed.
If energy is taken away from a gas (cooling) then individual particles will have less energy to overcome
attractive forces, and will eventually turn back into a liquid. Similarly, by cooling a liquid, it will turn
into a solid.
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Transitioning from one state of matter to another
There is a name for each of the respective changes from one state to another.
Diffusion - Diffusion is a mixing process. Particles move in a random motion, collide with air molecule
and fill all the available space.
The constant random movement of particles (and their kinetic energy) allows diffusion to occur.
Ultimately this means that particles will spread out from one place to another.
There are many things that can affect the rate of diffusion. Molecular mass is one of these things.
Heavier molecules will travel slower than lighter molecules. (the higher the molecular mass, the slower
the rate of diffusion).
To demonstrate diffusion in gases, a long glass tube is set up with cotton wool soaked with hydrochloric
acid at one end, and cotton wool soaked with ammonia at the other end. The hydrogen chloride and the
ammonia gases diffuse along the tube from either end, because the particles are constantly, randomly
moving. Also the white ring of NH4Cl is formed closer to HCl due to its greater Mr.
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Describe the effects of temperature and pressure on the volume of a gas
Effect of temperature on the volume of gas:
When the gas is heated, its volume increases. The particles gain kinetic energy and move faster. They
collide more often with greater force on the container walls. Temperature is directly proportional to the
volume of the gas.
Effect of pressure on the volume of gas:
the volume of gas decreases with an increase in pressure, due to which gas converts into liquid and
eventually into a solid state. The number of molecules in the gas is directly proportional to the pressure
of gas and volume is inversely proportional to the gas.
According to the Kinetic-Molecular Theory, changes of the state of matter occur when energy is added
or removed from a substance. The addition of energy makes particles move more, pushing them apart
and weakening the bonds between them. Removing energy makes particle movement slow down, reforming the bonds.
Increase in temperature increases kinetic energy of the particles. Particles moves faster and collide with
each other more frequently and effectively. This leads to more particles to attain activation energy and
the rate increases.
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UNIT 2. Atoms, Elements and Compounds
Atomic structure
Atoms are the smallest particle of a chemical element that can exist. Elements are substances that are
composed of just a single type of atom.
For example, the element carbon is made of only carbon atoms. Likewise, the element oxygen is made
of only oxygen atoms.
The structure of an atom is made up of three sub-atomic particles: Protons, neutrons, and electrons.
The above diagram is an example of a helium atom.
1. Location
• Protons & neutrons are always found in the nucleus
• Electrons are found in shells, and they orbit the nucleus
2. Charge
• Protons have positive charge (+)
• Neutrons have zero charge (0)
• Electrons have negative charge (-)
3. Mass
• Protons have a relative mass of 1
• Neutrons have a relative mass of 1
• Electrons have a negligible relative mass of 1/1840, which is essentially zero.
The table below is a summary:
Isotopes are atoms of same element with same atomic number (no. of protons) but with different mass
number (no. of neutrons and protons). Neutron number will differ in the isotopes. Example carbon has
three isotopes C612, C613 and C614.
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Periodic table
There are many chemical elements on earth, and the periodic table summarizes all of them into a single
table:
Not only does it tell us the full names of all existing elements (and their respective shortened symbols),
but it also gives us important information regarding the structure of a single atom of that particular
element.
Take helium for example: He24
Firstly, it tells you the full name and shortened chemical symbol for the element. In this case – Helium
(He). This is fairly straight forward.
Secondly, it tells you the proton number and the mass number of a helium atom. Here are a couple
of extremely important things to remember:
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The proton number (aka atomic number) is the number of protons in the atom
The mass number is the number total number of protons and neutrons (recall that neutrons/protons
have mass but electrons do not)
The number of electrons will always equal the number of protons
From the information provided by the periodic table, we can calculate the number of protons,
neutrons and electrons of the atoms of any particular element.
example: O816
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An oxygen atom has 8 protons
This means that it has 8 electrons
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Since its mass number (proton + neutron) is 16 it must mean that it has 8 neutrons (16 – 8)
Electron arrangement
Recall that electrons are held in shells. Shells are represented as rings around the nucleus.
It is really important to understand that the maximum number of electrons that a single shell can hold
can vary.
Take a look at this diagram below:
As described above, the first (inner) shell can hold up to 2 electrons. The second and third shell can
hold up to 8 electrons. It will follow 2n2 rule where n – is the number of shell.
Example: Oxygen
Above, we established that a single oxygen atom holds 8 protons, 8 neutrons and 8 electrons.
Remember, electrons always fill from the inner shell first. Since we know that the first electron shell
can only hold up to 2 electrons, it must mean that the rest of the electrons (6 of them) are held in the
2nd shell. This is what an oxygen atom therefore would look like diagrammatically:
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Knowing the electron arrangement of atoms are extremely important because it defines the entire
foundation of chemistry. Why do you think atoms react with one another? It’s because all atoms have
a goal. Do you know what that goal is? It’s simple: To achieve a full outer shell of electrons.
Reactivity of elements
As mentioned above, all atoms have a simple goal of wanting to achieve a full outer shell of electrons.
If you look at the diagram of oxygen above, you will see that an oxygen atom has 6 electrons in its most
outer shell. So how could oxygen achieve its goal? There are possible two main ways:
1. Gain 2 electrons
• If an oxygen atom just added two extra electrons into its outer shell, then it would have 8, and
therefore a full outer shell!
2. Lose 6 electrons
• If an oxygen atom lost all six of its outer electrons, then that shell would simply disappear. That
means the inner shell (with two electrons) will become the most “outer shell”. This would also
mean that indeed the atom would now have a full outer shell since two electrons is the maximum
(for that shell)!
Would option 1 be easier or option 2? Indeed, gaining 2 is a lot easier than losing 6 and therefore this
is what happens in reality. Oxygen either gains 2 extra electrons by sharing with other atoms or by
a transfer process.
So really, the reason why chemicals react with each other to begin with is because these reactions allow
atoms to obtain full outer shells.
Now you may notice that some elements in the periodic table already have full outer shells. These are
called noble gases and they are placed in the most right hand side of the table (i.e. Helium, Neon, etc.).
As you would expect, these noble gases are inert (do not react) because they simply do not need to.
They already have a full outer shell electrons.
Chemical bonding
There are several types of chemical bonds. We will be looking at ionic bonds, covalent bonds, and
metallic bonds.
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Elements are substances made of just one type of atom.
Compounds on the other hand, are substances that are made from chemical bonds between two
or more different elements.
A mixture is a combination of two or more different substances in the absence of chemical
bonds.
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Ionic bonding [Metals & non-metals]
When atoms loss or gain electrons to obtain a full outer shell, the neutral charge of the atom will be
disrupted since proton number will now be unequal to the electron number. If this happens, the atom is
now called an ion. The ion can have a positive charge (cation) or a negative charge (anion).
Metal atoms lose electrons to form cations and non-metal atoms gain electrons to form anions. Since
cations and anions have opposite charge, they are attracted to each other via strong electrostatic forces.
This is called ionic bonding: The bonding between anions and cations via strong electrostatic forces of
attraction.
In ionic bonding, metallic elements will donate their outer electrons to non-metal elements that need it.
Both elements will therefore achieve full outer shells and turn into cations & anions that get bonded by
electrostatic forces.
Example #1
Elements of group 1 (metals) and group 7 (non-metals) in the periodic table form ionic bonds. This is
because group 1 elements need to lose 1 electron to be happy, whilst group 7 elements need to gain 1
electron. This is a win-win situation! The group 1 metal will simply donate an electron to the group 7
non-metal and this will result in the formation of cations and anions that are bonded via ionic bonding.
This is an example of sodium chloride:
Example #2
Magnesium is a group 2 element and needs to remove 2 electrons to achieve a full outer shell. Similar
to the situation above, it can also form ionic bonds with fluorine (a group 7 element) by donating its
electrons. The only difference is that magnesium will be donating to two chlorine atoms (giving 1
electron each).
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Final structure of an ionic compound
Whilst the above diagrams are used to demonstrate ionic bonding diagrammatically, it
does NOT represent the final structure of an ionic compound. In fact, all ionic compounds have a 3D
lattice structure.
In sodium chloride for example, many sodium cations and chloride anions will join each other in regular
arrangements (called a lattice) forming a 3-dimensional structure full of cations and anions joined by
ionic bonds. This is what the final structure would look like:
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Covalent bonds [Non-metals & non-metals]
Atoms can achieve a full outer electron shell via sharing electrons. A pair of electrons (one from each
atom) can be shared. This is a single covalent bond and it holds the two atoms together.
Please note that atoms can be bonded via a single bond (sharing a single pair of electrons), double bond
(sharing two pairs) or a triple bond (sharing three pairs). Moreover, covalent bonds will ONLY exist
between two non-metals.
The examples below show that by sharing electrons, all atoms in the bond successfully achieve a noble
gas configuration.
Complex examples
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Intermolecular vs intramolecular forces
Knowing the difference between inter-molecular forces and intra-molecular forces is extremely
important.
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When you are melting or boiling a substance, it is the inter-molecular forces that you are breaking,
NOT the intra-molecular attraction.
Inter-molecular forces are attractive forces that exist between one molecules to another. These are
usually quite weak.
Intra-molecular forces are attractive forces that exists between atoms within the molecule. These are
usually extremely strong.
Differences between ionic and covalent compounds
Macromolecules
All of the examples of covalent molecules that we have looked at above are simple molecules. This
means that atoms are bonded to one or few other atoms to make a molecule or a compound that are
attracted to one another via inter-molecular forces (as described above).
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Macromolecules on the other hand are giant structures made of millions of atoms all joined by
covalent bonding. In other words, a huge number of atoms are joined via intra-molecular forces which
are extremely powerful (also mentioned above).
Diamond
This is a crystalline form of the element carbon. It has a three-dimensional structure in which every
carbon atom is covalently bonded to 4 other carbon atoms.
This is a small part of the structure of a diamond:
Graphite
This is another form of the element carbon. The atoms are covalently bonded in layers, with each atom
is strongly bonded to 3 other atoms in the same layer.
An important thing to note is that when carbon forms 3 covalent bonds with other carbon atoms, each
carbon atom will actually have a spare electron left over (you do not need to know the specifics of this).
These free electrons are called the ‘sea of electrons’ and they are free to move within the layers of
graphite. It is also because of these electrons that the layers are held together (weakly).
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Silicon (IV) Oxide
One silicon atom is bonded to four oxygen atoms and each oxygen atom is bonded to two silicon atoms
in a tetrahedral arrangement. This structure is very similar to the structure of a diamond and
consequently, the properties are also very similar. Silicon (IV) Oxide in the form of quartz exists as
colourless crystals. They are very hard, have a high melting point, and they do not conduct electricity.
RED = Oxygen [2 bonds per atom]
BLACK = Silicon [4 bonds per atom]
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Metallic bonding
In metals, the atoms shed their outer electrons to become cations. The cations are arranged in a regular
lattice structure whereas the removed electrons are delocalized and free to move throughout the
structure (this is called the sea of electrons).
The lattice arrangement of cations are therefore surrounded by free electrons and since cations and
electrons have opposite charge, they attract each other which bonds the structure together.
Metallic bonding is therefore defined as the electrostatic forces of attraction between cations and their
surrounding sea of electrons.
Positive metal ions are arranged in lattice between the pool of negative delocalised electrons.
The attraction is between positive metals ions and negative electrons.
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UNIT 3. Stoichiometry
Formula of simple compounds
To figure out the chemical formula of a given compound all you need to know are the valencies of the
atoms that make the compound.
The valency is the amount of electrons an atom has to gain or lose in order to achieve a full outer shell.
You can figure out the valency of an atom via the periodic table.
The group of the element (on the periodic table) tells you how many electrons there are in the atom’s
outer shell. From this information, you can figure out how many electrons it needs to gain or lose to
achieve a full outer shell. Groups 1-3 (metals) will LOSE electrons. Groups 5-7 (non-metals) will GAIN
electrons. Group 4 can do either.
Also remember, many atoms will become ions due to the loss or gain of electrons in ionic bonding. This
means the valency of an atom will also tell you the charge of its respective ion (i.e. Sodium ion = +1
charge. Chloride ion = -1 charge etc.)
So once you know the valencies of the atoms, all you need to do is swap the numbers around and cancel
them out if they are equal. Take a look at the examples below, you will understand what I mean:
Examples
1. What is the formula for Magnesium Chloride?
• Mg (valency 2) + Chlorine (valency 1)
• Swap the two numbers around
• Formula is therefore
2. What is the formula for Aluminium Oxide?
• Al (valency 3) + Oxygen (valency 2)
• Swap the two numbers around
• Formula is therefore
3. What is the formula for Calcium Oxide?
• Ca (valency 2) + Oxygen (valency 2)
• In this case, because the valencies are equal, you must cancel them out
• Formula is therefore just CaO
Writing equations
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This skill will be briefly introduced here and developed further throughout the course. All chemical
reactions can be represented by equations. You need to know how to write both word equations and
symbol equations.
Word equation
These are very simple. You simply write the equation in words. For example:
Magnesium + Oxygen -> Magnesium Oxide
Symbol equations
This is the most common type. Atoms and compounds are represented by their symbols. This is a bit
more complex because the equation needs to balance.
A balanced equation means that there are the same number of each type of atom on both sides of the
chemical equation. For example:
As we learnt above, we know that the chemical formula for magnesium oxide is MgO because
magnesium and oxygen both have a valency of 2 which cancels out.
Now if you look closely, the above equation is not balanced. Why? Because the left hand side has two
oxygen atoms, but the right hand side only has one.
So how the hell do we balance this thing? Well, we do so by adding numbers in front of the reactant or
products like so:
Now if you look at it, the equation is balanced!
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LHS = 2 Mg and 2 O
RHS = 2 Mg and 2 O
State symbols
State symbols represent the physical state in which the reactions or products are in a chemical reaction.
For example:
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Definitions
Mole equations
Many of the calculations that you’ll need to do will involve the concept of moles. There are three
important equations that you need to learn:
Moles and masses
Example 1 – Calculate the relative formula mass of the following: (You can use the periodic table)
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Remember, the relative formula mass (Mr) is just a sum of all the different relative atomic masses
(Ar). The Ar is really just a fancier term for “mass number” and this can easily be found in the periodic
table!
Example 2 – Calculate the mass of one mole of the following:
Remember, one mole is equivalent to 6X10^23 atoms, molecules or ions of the substance. The mass of
one mole of a substance is equal to the relative formula mass (Mr)!
Example 3 – Calculate the mass of each of the following:
Now that you know the mass of one mole of any compound is equal to the relative formula mass, you
can calculate the mass of a given compound as long as you know how many moles there are!
Example 4 – Calculate the mass of magnesium oxide formed when 3.0g of magnesium reacts with
excess oxygen
* Step 1: Write down a balanced chemical equation
* Step 2: Calculate the amount of moles of the reactant
* Step 3: Calculate the amount of moles of the product
* Step 4: Calculate the mass of product
Firstly, write down a balanced equation:
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3.0g of Mg = 3/24 = 0.125 moles (of Mg)
The mole ratio between Mg and MgO is 2:2 (or 1:1). Therefore, 0.125 moles of Mg will form 0.125
moles of MgO.
*The mole ratio is the ratio of ‘big numbers’ in front of the reactants and products inside the equation.
In this case, Mg and MgO both have a number 2 at the front. Therefore the ratio is 2:2 (and thus 1:1).
In this scenario, one mole of Mg will form one mole of MgO. Theoretically, if the ratio was 1:2 that
would mean one mole of Mg would make 2 moles of MgO.
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Since 0.125 moles of MgO is formed, the mass can be calculated via mole equation 1
0.125 X (24+16) = 5.0g of MgO has been produced from 3g of Mg in excess oxygen.
Moles and volumes
Example 1 – Calculate the amount of moles of oxygen molecules in the following volumes of oxygen
at rtp:
Example 2 – Calculate the amount of volume of oxygen at rtp for each of the following masses of gas:
Do not be confused. This is simple! First figure out the number of moles of oxygen (mass/Mr). Now
simply apply mole equation 2 and you’re done!
Example 3 – Calculate the volume of oxygen at rtp required to burn 1.4g of butene:
The balanced chemical equation will most likely be provided:
*The way you approach this question is very similar to example 4 from ‘reacting masses’ above. To
calculate the volume of oxygen required, all you need to know is the moles of oxygen that is required.
We know that the mole ratio of butene to oxygen is 1:6 so that means for every mole of butene, 6 moles
of oxygen is required. All we need to know, then, is the amount of moles in 1.4g of butene and that is
easy!
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1.4g of butene = 1.4 / (4X12 + 1X8) = 0.0259 moles
0.0259 moles of butene requires (0.025 X 6) moles of Oxygen due to 1:6 mole ratio. This equates to
0.15 moles (of oxygen)
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0.15 X 24 = 3.6
is the volume of oxygen required
Moles and concentrations
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Example 1 – Calculate the volume of sodium hydroxide, concentration 0.16 moldm-3, needed to
neutralize 20cm3 of sulphuric acid, concentration 0.2 moldm-3
The balanced chemical equation is as follows:
*Using the mole equation above, start by calculating the amount of moles in sulphuric acid. The mole
ratio here is 1:2, meaning for every mole of sulphuric acid, double the amount of sodium hydroxide will
be required. Once you’ve figured out the required amount of moles of sodium hydroxide, the volume
can easily be obtained by rearranging the formula.
*Also remember, cm-3 needs to be converted into dm-3 by dividing by 1000
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Moles (sulfuric acid) = 0.2 X (20/1000) = 0.004
Moles (NaOH) = 0.004 X 2 = 0.008
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Volume (NaOH) = 0.008/0.16 = 0.05
% Yield
In a chemical reaction, the expected/calculated amount of product is the theoretical yield. Unfortunately
however, the product amount actually produced (actual yield) is often lower than this theoretical
amount. Taking a ratio of these two values gives us the % yield.
Example – Excess magnesium carbonate was added to 25cm-3 sulfuric acid, concentration of 2.0moldm-3. The
unreacted magnesium carbonate was removed by filtration. The solution of magnesium sulfate was evaporated to
give 6.7g of hydrated magnesium sulfate crystals. Calculate the percentage yield
Balanced chemical equation is as follows:
The question tells you that 6.7g of crystals were formed, so therefore this is the actual yield.
The theoretical yield can be calculated as follows:
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Moles (H2SO4) = 25/1000 X 2 = 0.05
Moles (MgSO4.7H2O) = 0.05 (Because mole ratio is 1:1)
Relative formula mass (MgSO4.7H2O) = 246
Mass = 0.05 X 246 = 12.3g (theoretical yield)
% yield = 6.7/12.3 = 0.545 (54.5%)
*The original answer is in decimals. To convert decimals to percentage, multiply by 100
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% Purity
Example – 7.0g of impure calcium carbonate was heated and 2.42g of carbon dioxide was collected.
Calculate the percentage purity of the calcium carbonate
The key thing here is that carbon dioxide can only be made from pure calcium carbonate. The impurities
in the original sample (of 7.0g) will not contribute to the production of carbon dioxide.
The percentage purity can therefore be calculated as follows:
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Moles (CO2) = 2.42 / (12 + 32) = 0.055
Moles (CaCO3) = 0.055 (Mole ratio 1:1)
Relative formula mass (CaCO3) = 100
Mass (Pure CaCO3) = 0.055 X 100 = 5.5g
Percentage purity = 5.5/7 = 0.786 (78.6%)
Empirical and molecular formula
Example – A hydrocarbon contains 92.3% carbon and 7.7% hydrogen. It’s relative molecular mass is
78. Calculate it’s empirical and molecular formulae
The ratio of carbon to hydrogen is therefore 1:1
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Empirical formula (simplest ratio) is therefore CH
The molecular formula is a multiple of the empirical formula. The Mr of the hydrocarbon is 78
and the Mr of our empirical formula (CH) is 13, giving us a multiple of 6. The molecular formula
is thus:
ee Exam Academy
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UNIT 4. ELECTRO CHEMISTRY
Definitions
Electrolysis
Electrolysis is the breakdown of an ionic compound (molten or aqueous solution) by the passage of
electricity.
Fundamentals
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Reactions at the cathode or anode
Electrons flow from the battery to the cathode. Cations (usually metal and hydrogen ions) in the
electrolyte are attracted to the cathode (negative electrode). Cations accepted electrons from the
cathode, and therefore metals and hydrogen are formed at the cathode. For example:
Electrons flow from the anode to the battery. Negative ions (non-metals except hydrogen) are attracted
to the anode (positive electrode).
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If the anode is inert (i.e. carbon or platinum) the negative ions lose electrons to the anode:
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If the anode is not inert (i.e. silver, copper, or other reactive metals) the metal atoms of the anode
lose electrons and form positive ions. The anode will therefore dissolve and become smaller:
Ions of an electrolyte the electrolyte can either be molten or aqueous.
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A molten substance means that the substance has been melted down. The ions therefore come
only from the substance itself.
An aqueous solution means that the substance is dissolved in water. The water molecules
themselves can ionize so you will always find hydrogen and hydroxide ions in addition to the
ions from the solute.
The discharge of ions
As we looked at above, ions are discharged at the anode or cathode.
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In many cases, you will find that there are more than one cations or anions inside the electrolyte. For
example:
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At the cathode you will find that the H+ will get discharged rather than Na+.
At the anode you will find that OH- will get discharged rather than Cl-.
The electrochemical series tells us which ions discharge easier than others. The lower ion of each series
will be the one to get discharged.
The electrochemical series
Series of positive ions
Series of negative ions
Potassium k+
Sulfate ion SO42- no product
Sodium Na+
Nitrate ion NO3- noproduct
Calcium Ca2+
Oxide O2-/ Hydroxide OH-
Magnesium Mg2+
Chloride ion Cl-
Aluminium Al3+
Bromide ion Br-
Zinc Zn2+
Iodide ion I-
Iron Fe2+
Tin Sn2+
Lead Pb2+
Hydrogen H+
Copper Cu2+
Silver Ag+
Follow the basic principles for each example. Firstly, figure out the ions inside the electrolyte. Secondly,
figure out which ions will be discharged (from the electrochemical series). Write down the reactions at
the electrodes and also figure out what remains inside the final electrolyte.
Molten sodium chloride (inert electrodes)
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Ions present:
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Reactions in electrodes:
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Sodium chloride is therefore decomposed
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Concentrated aqueous sodium chloride (inert electrodes)
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Ions present:
Reactions in electrodes:
Na+ and OH- remain in the electrolyte (which is sodium hydroxide)
If the solution is very dilute, then OH- would get discharged instead of the Cl-. This means Na+ and
Cl- would remain in the electrolyte and the solution will become more and more concentrated (as
water is used up).
Concentrated hydrochloric acid (inert electrodes)
Ions present:
Reactions in electrodes:
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Acid therefore gets used up in the electrolyte
Dilute sulfuric acid (inert electrodes)
Ions present:
Reactions in electrodes:
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Acid gets more concentrated as water gets used up
Aqueous copper (II) sulphate (Inert electrodes)
Ions present:
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Reactions in electrodes:
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H+ and SO42- ions remain in the solution (which is sulfuric acid)
Aqueous copper (II) sulphate (copper electrodes)
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Ions present:
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The only difference is that the anode is not inert. This means that the metal anode itself will
react by losing electrons to form ions.
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Copper deposited at the cathode becomes thicker. Copper is removed at the anode and it gets
thinner. The electrolyte remains the same since one electrode produces copper ions whereas the
other removes them. This process is used to electroplate other metals with copper.
Commercial use of electrolysis
Electroplating
This is used to plate one metal with another. The general arrangement for electroplating is shown
here:
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The metals commonly used to electroplate are copper, chromium, nickel, and silver. The two main
reasons for electroplating are appearance and protection from corrosion.
Refining metals
Metals can be refined or purified by electrolysis. The impure metal forms the anode, the cathode is a
small piece of pure metal and electrolyte is an aqueous metal salt. In the refining of copper, the
following reactions occur
Cathode:
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Copper ions from solution lose their charge and copper is deposited
Anode:
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Copper atoms lose their valency electrons and go into solution as ions
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Overall pure copper is transferred from the anode to the cathode. The impurities from the copper are
left as ‘anode slime’ and the cathode becomes a large piece of pure copper.
Aluminium extraction
Critical information:
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Main ore of aluminium is called bauxite
It is changed to pure aluminium oxide (alumina)
Graphite cathode and anode (therefore made of carbon)
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Electrolyte is molten mixture of pure aluminium oxide dissolved in cryolite
The point in cryolite is to lower the temperature from approximately 2000 to 900 degrees.
Reactions at electrodes:
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The carbon anodes burn away in oxygen and is replaced periodically
Industrial use of sodium chloride
As we looked at above, the use of concentrated sodium chloride can be used in electrolysis to make
hydrogen gas, chlorine gas, and sodium hydroxide.
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Chlorine can be then used in making solvents, treating drinking water, bleach, etc.
Hydrogen is used in the Haber process, making fuels in cells, making margarine etc.
Sodium hydroxide is used in soap manufacture
Electric cables: Conductors and insulators
Copper and aluminium are commonly used as conductors in electric cables. You need to know why
they are good for this purpose.
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Copper
• Good conductor of electricity
• Ductile
• Easily purified
Aluminium
• Good conductor
• Resists corrosion
• Low density, allowing high diameter cables to be used. This reduces resistance and sagging.
Plastics and ceramics are often used as insulators in electric cables.
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Plastics
• Do not conduct electricity
• Flexible & easily molded
• Non-biodegradable
Ceramics
• Do not conduct electricity
• High melting points allowing use at high temperatures
• Not affected by water or oxygen
• Can be molded into complex shapes
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UNIT 5. Chemical energetics
Energetics of a reaction
Definitions
All chemical reactions fall into two categories: Exothermic or endothermic.
An exothermic reaction gives out heat. The total chemical energy of the reactants is larger than the
products. This difference in energy is transferred to surroundings as heat.
An endothermic reaction takes in heat. The chemical energy of the reactants is smaller than the products
so this difference in energy is transferred from the surroundings to the chemicals.
The activation energy is the minimum amount of energy which the reacting species must possess in
order to undergo a specified reaction.
Bond breaking and bond making
The process of bond breaking is an endothermic process. Energy must be “taken in” to break bonds
apart.
The process of bond making is an exothermic process. Energy (heat) is released when this happens.
Quite often they will give you an equation such as:
30
They will also give you the relevant bond energies:
•
•
•
•
C-H = 413 kJ/mol
F-F = 158 kJ/mol
H-F = 565 kJ/mol
C-F = 495 kJ/mol
An endothermic reaction has a (+) sign because energy is taken in. An exothermic reaction has a (-)
sign because energy is lost. For instance, if you break one mole of C-H bonds it is denoted as +413. If
you form one mole of C-H bonds, then it is denoted as -413.
From the information above, can you figure out whether the overall reaction is exothermic or
endothermic?
Energy transfer
The most common way of producing heat energy is by burning fossil fuels – natural gas, coal, petroleum
products
Hydrogen as fuel - The combustion of hydrogen is highly exothermic. It is only used as a rocket fuel,
in experimental vehicles, and fuel cells. There are advantages and disadvantages of hydrogen fuel.
•
Advantages
• Very energy rich
• No pollutants are formed
• Nitrogen oxides not formed (these are environmentally harmful i.e. acid rain)
•
Disadvantages
• Expensive to produce
• Difficult (and expensive) to store
31
Hydrogen fuel cells
Much like the set-up for electrolysis, you have an anode (+) and cathode (-). Reactants are supplied to
each electrode. The fuel (containing hydrogen) is supplied to the anode and the air (containing oxygen)
is supplied to the cathode.
At the anode, hydrogen molecules lose electrons to form ions in the electrolyte:
At the cathode, oxygen gains electrons:
The ions formed react to product water:
The overall equation is:
The overall equation above demonstrating the formation of water is essentially an exothermic reaction.
When hydrogen reacts with oxygen to form water in a fuel cell, electrical energy is produced.
32
UNIT 6. Chemical Reactions
Physical vs chemical changes
It is important for you to understand the difference between chemical and physical changes. Some
changes are obvious, but there are some basic ideas you should know.
Physical changes are usually about physical states of matter. Chemical changes happen on
a molecular level when you have two or more molecules that interact. Chemical changes happen
when atomic bonds are broken or created during chemical reactions.
Collision theory
At an atomic level, a chemical reaction will occur when two conditions are met. Two particles need to
collide and they must have enough energy to react. This means the reaction rate will depend on these
two factors: Collision rate & particle energy.
If two atoms collide but they don’t have enough energy, then the reaction will not occur. If the particles
have enough energy but they don’t collide, then again, the reaction will not occur.
It is very important to realize that there are certain things such as concentration, pressure, temperature,
catalysts, and particle size that affect the collision rate & particle energies, and therefore directly affect
the chemical reaction rate.
Factors affecting rate of reaction
We briefly mentioned above the several factors that affect reaction rates. We will be going through each
of these in a bit more detail.
Firstly however, it is important to note that the best definition for the rate of reaction is:
This means that the rate is dependent on concentration (i.e. amount of substance in a specified
volume) rather than just the amount of substance.
Concentration
When the concentration is increased, the rate of reaction is also increased due to higher collision rates
(since there are more particles per unit volume).
33
Pressure
The pressure only affects reactions with gases. An increased pressure means gas molecules are closer
together. This increases the collision rate and thus the reaction rate.
Temperature
When the temperature is increased, the rate of chemical reaction will increase due to larger amounts of
energies of individual particles and a higher collision rate (since particles are moving quicker).
Particle size
This only affects reactions involving solids. Smaller particle sizes mean that there is a larger surface
area for collisions to occur, which in turn, increases the reaction rate.
This diagram below demonstrates this concept well:
Catalysts
A catalyst increases the rate of a reaction but remains chemically unchanged. Enzymes are biological
catalysts.
Experimental methods
CIE expects you to understand how to devise suitable experiments to investigate each of the above
variables on reaction rates.
Most experimental techniques should be learnt in the lab as a part of your curriculum. We will touch
on the details of experimental techniques in a separate section.
Photochemical reactions
Reduction of silver (I) Halide
This is the basis of photography. A photographic film i coated with a layer of silver (I) bromide. When
exposed to light, silver ions accept electrons from bromide ions and form silver atoms. This is called
reduction (more details down the page)
34
Parts of the film that have been exposed to light turn black, while unexposed portions remain white,
The rate of reaction depends on intensity (i.e. brightness) of the light.
Photosynthesis
Green plants make carbohydrates via this reaction.:
The reaction is catalysed by chlorophyll (the green pigment in plants) and occurs only in sunlight.
Again, the rate of reaction is dependent on light intensity.
Reversible reactions
A reversible reaction is a chemical reaction where the reactants form products that, in
turn, react together to give the reactants back.
When hydrated copper (II) sulphate is heated, it decomposes. This is the forward reaction which is
endothermic.
When the products are cooled and mixed, the reverse reaction occurs. This reaction is exothermic.
Therefore the overall reversible reaction can be written into one equation:
Equilibrium
As we saw above, in a reversible reaction, the reactants make the products and the products make the
reactants.
Eventually, the reaction will reach an equilibrium whereby the rate of the forward reaction and the
reverse reaction are equal. This means that the concentrations of reactants and products will stay exactly
the same (unless the conditions are changed).
The concentrations of reactants and products in an equilibrium is called the position of equilibrium.
35
•
•
If the position of equilibrium moves to the right, it means that in the “new” equilibrium, the
concentration of products has increased whereas the concentration of reactants has decreased.
If the position of equilibrium moves to the left, it means that in the new equilibrium, the concentration
of reactants has increased and the concentration of products have decreased.
For example:
If the equilibrium shifts to the RIGHT that means more C is being made from A and B. Therefore, the
concentration of C increases whilst the concentration of A and B decreases.
If the equilibrium shifts to the LEFT that means more A and B is being made from C. Therefore, the
concentration of C decreases whilst the concentration of A and B increases.
There
are
certain
conditions
that
affect
the
position
of
equilibrium.
The position of equilibrium will shift in the direction that OPPOSES the change in condition.
For example, increasing the concentration of product C will shift the equilibrium to the left (to try and
reduce the concentration). Bear this in mind when reading through the separate conditions below.
1. Concentration
• Increasing the concentration of the product will shift the equilibrium to the left (to reduce the
product concentration). Decreasing the product concentration will shift the equilibrium to the
right (to produce more).
• Increasing the concentration of the reactant will shift the equilibrium to the right. Decreasing the
reactant will shift it to the left.
2. Temperature
• In a reversible reaction, one reaction is exothermic and the other is equally endothermic.
• Increasing the temperature will therefore shift the equilibrium towards the endothermic reaction
(to reduce heat)
• Decreasing the temperature will shift the equilibrium towards the exothermic reaction (to
increase heat)
3. Pressure
• This factor is only relevant for reacts that involve gases. Please refer to this example
• Note that the total number of gas molecules in the LHS = 4 and the RHS = 2
Increasing the pressure will move the equilibrium to the side with a smaller number of gas
molecules (i.e. Right in this example)
• Decreasing the pressure will move the equilibrium to the side with a larger number of gas
molecules (i.e. Left in this example).
4. Catalysts
• CIE loves to trick you with this so be careful
• Catalysts do NOT affect the position of equilibrium. It affects the rate of reaction only.
•
Redox
Oxygen gain/loss
Redox is shortened for ‘reduction’ and ‘oxidation’. This can be explained via the gain or loss of oxygen.
•
•
Oxidation is the gain of oxygen
Reduction is the loss of oxygen
36
For example:
In the equation above, CuO has been reduced because it has “lost” an oxygen to become Cu.
Meanwhile, the hydrogen has been oxidized because it has “gained” an electron to become H2O.
Electron transfer
The concept of reduction and oxygen can also be explained in terms of electron gain or loss.
•
•
Oxidation is the loss of electrons
Reduction is the gain of electrons
For example:
In the example above, the magnesium atom loses electrons to become an ion so therefore it has
been oxidized.
The chlorine molecule on the other hand, gains two electrons to become chloride ions and therefore it
has been reduced.
UNIT 7. Acids, bases, and salts
Definitions
The pH scale
The pH scale is a measure of acidity or alkalinity of water soluble substances.
As you can see above, the scale range from 0 to 14. Water has a pH of 7 and it is therefore neutral. The
lower the pH, the more acidic it is and the higher the pH the more basic it is.
•
The lower the pH, the more acidic the solution is with higher the concentration of hydrogen ions
(H+) in aqueous
37
•
The higher the pH, the more basic the solution is with a higher concentration of hydroxide ions (OH) in aqueous
pH indicators
Litmus paper, & methyl orange are used as pH indictors which allow us to determine whether a solution
is acidic, basic, or neutral.
•
Litmus paper
• Turns red in acid
• Turns blue in alkali (soluble base)
• No change in water
•
Methyl orange
• Turns red in acid
• Turns yellow in alkali (soluble base)
• Orange in water (original colour)
Simple definition of acid & base
So what exactly are acids and bases?
The most basic definition of an acid is a substance that generates hydrogen H+ ions (or protons) in
aqueous.
The most basic definition of a base is a substance that generates hydroxide ions (OH-) in aqueous.
•
•
Bases are generally metal oxides or hydroxides
Bases that are soluble in water are called alkalis
A more complex definition of acids and bases is associated with proton (H+) transfer. An acid is defined
as a proton donor whereas a base is defined as a proton acceptor.
Neutralization
Neutralization is the process by which an acid and base react to form water. Using the definitions above,
it is easy to imagine this process. The H+ ions from the acid gets “neutralized” by the OH- ions from
the base.
Properties of acids
Now that you understand that acids are chemicals that release H+ ions in water, lets take a look at some
reactions that CIE wants you to learn.
38
Acid + Metal
In chemistry, a salt is just another word for an ionic compound. An acid-metal reaction will always
form hydrogen and the respective salt.
Acids + bases
Acids + carbonates
Properties of bases Bases + acids (again)
Bases + ammonium salts
Strong acids and bases
Remember, acids release H+ ions in aqueous solution and bases release OH- ions. The concentration of
these respective ions dictate the strength of the acid or base.
•
A strong acid will have a high concentration of H+ ions (i.e. low pH)
•
A strong base will have a high concentration of OH- ions (i.e. high pH)
39
A strong acid or base will therefore release a lot of hydrogen ions or hydroxide ions respectively. This
is because the molecules are completely ionized in aqueous solution (denoted by single arrow).
Hydrochloric acid is an example of a strong acid:
Sodium hydroxide is an example of a strong base:
Weak acids and bases
•
A weak acid will have a low concentration of H+ ions
•
A weak base will have a low concentration of OHA weak acid or base will therefore release small amounts of hydrogen ions or hydroxide ions
respectively. This is because the molecules are partially ionized in aqueous solution (denoted
by double arrow).
Ethanoic acid is an example of a weak acid:
Ammonia is an example of a weak base:
Types of oxides
From a given oxide, CIE want you to derive whether the oxide is an acid, base, neutral, or amphoteric.
Most metal oxides are basic. Soluble bases are called alkalis and they turn litmus paper blue. Insoluble
bases will not affect litmus paper. As discussed above, these basic oxides will undergo a neutralization
reaction with acids.
Non-metal oxides are usually neutral or acidic
•
•
Neutral oxides such as water, nitrogen (II) oxide, and carbon monoxide do not react with acids or
bases
Acidic oxides such as sulphur dioxide, sulphur trioxide, carbon dioxide, and oxides of phosphorus
will turn litmus paper red and neutralize bases
40
Amphoteric oxides react with either a base or acid to form salt and water. This means that these oxides
have the properties of a base and acid. Examples are zinc oxide/hydroxide and aluminium
oxide/hydroxide. You need to learn these two examples:
Behaving as a base
Behaving as an acid
Salt preparation
Titration
•
This is used to prepare a soluble salt from a soluble base (i.e. alkali) and an acid. This method is
used to make salts of group 1 metals and ammonium salts.
Neutralization of insoluble base by acid
•
An excess of the base is added to an acid, and the excess is removed via filtration. The filtrate is
partially evaporated to obtain crystals of the salt. Soluble salts of most metals that are not group
1 are made by this method.
Metal reacting with acid
•
This method is basically the same as above. It can make magnesium, zinc, aluminium, and iron
(II) salts. However it cannot be used to prepare salts of reactive metals (such as sodium and
potassium) due to the violent reaction.
Preparing insoluble salts via precipitation
•
Here is a list of the insoluble compounds that you need to know. Salts not on this list are
considered soluble.
• All carbonates except aluminium carbonate and all group 1 carbonates
• All hydroxides except calcium, strontium, barium, and group 1 hydroxides
• Barium, calcium, and lead sulphates
• All chlorides, bromides, and iodides of silver and lead
41
To make barium sulphate for example (an insoluble salt), two solutions must be mixed. One solution
must contain a soluble barium salt (i.e. barium chloride) and the other containing a soluble sulphate (i.e.
sodium sulphate). The precipitate of barium sulphate is filtered off, washed, and dried.
Ion & gas identification
42
UNIT 8. The periodic table
The periodic table is a method of classifying elements and its use to predict properties of elements.
Periodic trends
Groups - Elements are arranged in order of proton number. Elements with similar chemical properties
are placed in the same vertical column called groups. Elements in a group have similar chemical
properties, same outer electron numbers, and usually the same valency. Going down the group, the
elements become more metallic in character.
Periods -The horizontal rows are called periods. Moving across a period, the elements change from
metallic to non-metallic. The number of valency electrons increases across the period but the number
of occupied energy levels (i.e. shells) stays the same.
Metals vs non-metals - The main physical differences are summarized in this table:
The main chemical differences are summarized in this table:
43
Group properties
Group 1 – Alkali metals (i.e. Li, Na, K)
These are extremely reactive metals. The alkali metals (despite being metallic) are rather soft and have
low m.p/b.p compared to other metals. They are good conductors of heat and electricity and demonstrate
shiny surfaces when freshly cut.
Due to their low valency, these metals are extremely reactive. Each element will react with cold water
to form a hydroxide and hydrogen.
Physical trends down the group
•
•
•
Increasing softness
Decreasing melting/boiling points
Increasing densities
Chemical trends down the group
•
Increasing reactivity
• Lithium reacts steadily with water whereas potassium may cause an explosion
Group 7 – Halogens (i.e. Cl, Br, I)
Halogens are a collection of diatomic non-metals showing both physical and chemical trends down the
group.
Physical trends down the group
•
•
Colour gets darker down the group
• Chlorine is yellow/green
• Bromine is brown
• Iodine is a black solid with purple vapour
M.p and b.p increase down the group
44
•
•
•
Chlorine is a gas (rtp)
Bromine is a liquid (rtp)
Iodine is a solid (rtp)
Chemical trends down the group
•
Decreasing reactivity
• Chlorine can displace
both
bromine
and
iodine
from
their
compounds:
* chlorine is more reactive than both bromine and iodine. Therefore, in reaction (1), chlorine has the
ability to ‘kick out’ the bromide ion from the compound and take its place. Chlorine will therefore
become an ion to form an ionic bond with potassium (KCl), whereas bromine becomes a molecule
Transition elements
These are metallic elements placed in the middle of the periodic table.
•
•
Physics properties
• Compared to groups I and II, they have higher densities and melting points. They are also harder
and stronger.
Chemical properties
• Compared to groups I and II, they are a lot less reactive
• They do not react with cold water but many react when heated in steam:
•
•
They have more than one valency. Iron (Fe), for example, forms two different types of ions:
• This means that they can either decide to lose 2 or 3 electrons to become an ion
They form coloured compounds
• Iron (II) salts are pale green whereas iron (III) salts are yellow/brown
• Copper (II) salts are blue
• Nickel salts are bright green
Group 0 – Noble gases
Noble gases are unreactive; they have a valency of 0. Their outer electron shell is already complete.
Noble gases have various uses:
•
•
Helium is used in balloons. The balloon will float because helium is less dense than air and also safer
(because they cannot catch fire due to their unreactiveness)
Argon is used to fill electric light bulbs because it is inert
45
UNIT 9. Metals
Metallic properties
Physical properties
We have touched on the physical properties of metals in the previous topics. Here is a brief summary:
•
•
•
•
•
•
Shiny
Good conductors of heat/electricity
High density
Malleable and ductile
Usually solid at room temperature
Sonorous (makes bell-like sounds when struck)
Chemical properties
Metal + Acid
Metal + Oxygen
Metal + Cold Water
Metal + Steam
Note: Reactive metals such as sodium and potassium will react fine in cold water to produce hydroxide
salts. Less reactive metals like copper will not react in cold water. They will only react in steam and
produce oxide salts instead of hydroxide salts.
46
Alloys
An alloy is a mixture of two or more metals, or a mixture of one or more metals with a non-metal.
Alloys are used in preference to pure metals because they can be designed to have properties for
whatever usage purpose. For example, they may be made to be harder and more resistant to corrosion.
Alloys are harder than pure metals because the presence of different sized atoms will make the layers
less mobile and prevent them from slipping.
This diagram below represents a simple alloy: The mixture of metallic atoms (red) with other different
atoms (blue):
Some examples of alloys are:
•
•
•
Brass: mixture of zinc and copper
Mild steel: Iron and up to 0.3% carbon
Stainless steel: Iron, nickel, chromium
Reactivity series
Different metals have varying reactivities. This all dependent on the tendency of a metal to form its
positive ion. The greater the tendency to form the ion, the greater the reactivity of the metal.
The reactivity series orders metals from most reactive to least reactive:
47
NOTE: The main ones you need to know are potassium, sodium, calcium, magnesium, zinc, iron,
hydrogen, and copper
The reactivity series tells us that potassium (highest in series) atoms have a much higher tendency to
become cations than, say, platinum (lowest), and therefore much more reactive.
We can demonstrate the difference in reactivities by observing the reactions of each of these metals
with steam, dilute acid, and also the reduction of the oxides with carbon.
48
As you can see, the reactions become less vigorous down the table, suggesting the reduction of
relativities of the metals.
Reduction of metal oxides via carbon
Remember, reduction is the loss of oxygen. If a metal oxide gets “reduced” by carbon, it means that
carbon “steals” an oxygen from the metal oxide. For example:
The rule is, only a more reactive element can “steal” an oxygen from an oxide.
In this scenario, zinc oxide is reduced by carbon because carbon is more reactive than zinc (refer to the
reactivity series).
This is why the metal oxides that have metals above carbon on the reactivity series can not be reduced
but those that are below it can.
Displacement reactions
As mentioned above, the reactivity series is based upon the metal’s tendency to become cations. The
greater the tendency, the greater its reactivity.
Displacement reactions involve one ion replacing another. Whether or not a metal can displace another
metal in a compound is strictly dependent on their relative reactivities.
49
Examine the reaction above. What has happened to each of the metals in the equation?
The zinc has changed from the zinc metal to the ion form (in zinc sulfate). Meanwhile, the copper ions
(in copper sulfate) has become copper metal:
This is a redox reaction. Zinc loses 2 electrons (oxidation) and donates them to the copper ions that
ultimately gain the 2 electrons (reduction).
Quite simply, the more reactive metal zinc has displaced the copper in copper sulfate. This displacement
occurs because of the transfer of electrons in the redox equation above.
In this example, zinc can NOT displace the magnesium in magnesium chloride because it’s tendency to
form ions is lower (i.e. less reactive).
Decomposition reactions (via heat)
A decomposition reaction occurs when one reactant breaks down into two or more products. The more
reactive the metal, the more stable its compounds are (and thus harder to decompose).
Metal hydroxides
NOTE: Sodium and potassium hydroxides are exceptions and do not decompose when heated.
Metal carbonates
Extraction of metals
50
Reactive metals such as potassium, sodium, calcium, magnesium, and aluminium are extracted from
their ores via electrolysis of a molten compound.
CIE requires you to understand the extraction of zinc, iron, and aluminium.
Zinc extraction
•
•
The ore is zinc blende (ZnS)
This
is
roasted
in
the
•
The
carbon
in
•
Zinc distills out of the furnace
oxide
is
heated
with
a
air
to
form
the
furnace,
where
it
reduced
is
oxide
to
zinc
Iron extraction
•
•
•
The ore is called haematite
Haematite, coke (carbon), and limestone are added to a furnace
Carbon dioxide is formed
• From reaction between coke and oxygen
•
From decomposition of limestone
•
Carbon dioxide gets reduced to carbon monoxide
•
Carbon monoxide reduces the iron (III) oxide to iron
•
The impurity in the ore is sand (silivon IV oxide)
• This reacts with calcium oxide to form slag
•
•
Molten slag floats on molten iron
This can be run off separately and used as building material
Aluminium extraction
The details of aluminium extraction has been covered in the topic of electrolysis. Please click here and
scroll down the page to find the relevant information.
NOTE: Aluminium is a reactive metal, and quite often it reacts with oxygen in the air to form a
‘aluminium oxide coating’. This oxide coating makes the metal seem unreactive.
51
Uses of metals
Here we go through some brief uses of several different metals
Aluminium
•
•
Aircraft manufacture due to strength and low density
Food contains due to corrosion resistance
Zinc
•
Galvanizing and brass making
Copper
•
Electrical wiring and utensils
Steel
•
•
Car bodies and machinery
Stainless steel
52
UNIT 10. Chemistry of the environment
Water
Chemical Tests for the Presence of Water
Test
Type of test
Positive result
Anhydrous Cobalt (II) Chloride
Chemical
It turns from blue to pink
Anhydrous Copper (II) Sulfate
Chemical
It turns from white to blue
Test Melting and Boiling Point
Physical
M.P at 0℃ and B.P at 100℃
•
Distilled Water is used in practical chemistry rather than tap Water because it has fewer
chemical impurities.
Water from Natural Sources
Water is an essential source in the natural world. However, with the changing world and massive
urbanisation. Our water may contain substances such as:
1. Dissolved Oxygen (this is important for aquatic life)
2. Metal Compounds (Provide essential minerals for life - however, some are toxic)
3. Plastics (harm aquatic life)
4. Sewage (contains harmful microbes which cause diseases)
5. Harmful microbes
6. Nitrates from fertilisers
7. Phosphate from fertilisers and detergents (leads to deoxygenation of water and damage to
aquatic life (Eutrophication)
Treatment of Domestic Water Supply
1. Water is pumped into screens to remove solid, insoluble impurities.
53
2. A sedimentation process makes small clay pieces stick together and are then removed.
3. The water then undergoes filtration through layers of sand and gravel to remove larger,
insoluble debris.
4. Carbon is also added into filtered water to remove taste and odour.
5. The chlorination process adds chlorine gas bubbled into the water to kill bacteria and other
microbes; the acidic effect on the water is reversed by adding an alkali, sodium hydroxide.
Step one and five, chlorine is added before going to domestic areas.
Fertilisers
Fertilisers: Substances added to the soil and taken up by plants to increase crop yield.
Fertilisers
contain
substances
such
as nitrates and ammonium
salts.
N.P.K - Nitrogen, Phosphorus, and Potassium Fertilisers are found inside fertilisers, essential
to improve plant growth.
Functions of Elements
1. Nitrogen - Makes chlorophyll and protein. Promotes healthy leaves
2. Phosphorus - Promotes healthy roots
3. Potassium - Promotes growth and healthy fruits and flowers
54
Reaction with any alkali substance (except ammonia) displaces ammonia from its compound, for
example:
Calcium hydroxide + Ammonium chloride → Calcium chloride + Ammonia + Water
Air Quality and Climate
The pie chart below presents the components present in clean air:
•
Primary: Nitrogen (78%), Oxygen (21%)
•
Secondary: Noble
gases
(mainly
Argon)
and
Carbon
Dioxide
(1%)
Air Pollutants and their Adverse Effects
Pollutant
Source
Negative impact
Carbon
monoxide (CO)
Incomplete combustion of carboncontaining fuels (ex. Internal combustion
engines)
Binds with haemoglobin, constricting
oxygen supply in cells; leads to fatigue/
death
Carbon Dioxide
(CO2)
Complete Combustion
Containing Fuels
Carbon
Increased global warming leads to
climate change.
Methane (CH4)
Decomposition of vegetation and waste
gases from digestion in animals
Increased global warming leads to
climate change.
of
55
Pollutant
Source
Negative impact
Sulfur
(SO2)
Dioxide
Combustion of fossil fuels which contain
sulfur compounds
It causes acid rain.
Nitrogen Oxides
(NO2)
High temperatures that trigger a reaction
between N2 and O2 (from air)
Causes respiratory problems and
photochemical smog; contributes to
acid rain
Lead
Compounds
Combustion of leaded fuels
Damages brain and nerve cells in young
children
Impact of Greenhouse Gases
1. Short wavelength radiation from the Sun reaches the Earth's surface
2. Some thermal energy is absorbed and heats oceans/lands
3. Earth radiates some thermal energy as more prolonged wavelength radiation
4. Greenhouse gases absorb some of the infrared radiation and re-emit in all directions
5. Some infrared radiation comes back to Earth's surface, and this reduces the heat loss to space
and
leads
to
global
warming
Strategies to Reduce the Effect of Environmental Issues
56
1. Climate Change: planting trees, reducing livestock farming, decreasing the use of fossil fuels,
increasing the use of hydrogen and renewable energy, e.g. wind, solar
2. Acid Rain: use of catalytic converters in vehicles, reducing emissions of sulfur dioxide by using
low-sulfur fuels and flue gas desulfurisation with calcium oxide
Photosynthesis
Photosynthesis: the reaction between carbon dioxide and water to produce glucose and oxygen in
the presence of chlorophyll and using energy from light.
Word Equation: Carbon Dioxide + Water → Glucose + Oxygen
Balanced Chemical Equation:
6CO2 + 6H2O → C6H12O6 + 6O2
Catalytic Converters
1. Present in car exhausts; contains transition metal catalysts of platinum and rhodium
2. Aids redox reactions to neutralize toxic pollutants formed as a result of incomplete fuel
combustion: (a) Carbon Monoxide, (b) Nitrogen Oxides, (c) Unburned hydrocarbons
3. Reaction equations:
(a) 2CO+ O2 → 2CO2
(b) 2NO+ 2CO→ N2+ 2CO2
(c) C8H18 + 12½O2 → 8CO2 + 9H2O
57
UNIT 11. Organic Chemistry
Names of compounds
Organic chemistry is the branch of chemistry that deals with carbon compounds (other than simple salts
such as carbonates, oxides, and carbides).
A homologous series is a family of similar compounds with similar chemical properties due to the
presence of the same functional group.
A functional group is a set of atoms which dictate the chemical properties of that compound. Therefore,
compounds that share that same functional group will inherently share similar chemical properties.
The table below demonstrates the names of all the different functional groups that exist in organic
chemistry.
Do not be overwhelmed! Each of the relevant functional groups will be explained in detail in the next
few topics.
For now it is important to be aware that compounds with the same functional groups will always end
with the same suffix.
For example
•
•
•
•
•
All alkanes will end with ‘___ane’
All alkenes will end with ‘___ene’
All alcohols will end with ‘___anol‘
All carboxylic acids will end with ‘___anoic acid’
All esters will end with ‘___yl ___anoate”
58
Naming an organic compound can be done in 3 easy steps:
•
•
•
Find the suffix (dictated by its functional group)
Find the prefix (dictated by number of carbon atoms)
Find the position of the functional group
We will use this compound below as an example. Let’s determine it’s name step by step:
Step 1: Find the suffix
The molecule above is an example of an alcohol because it contains the functional group -O-H.
The suffix of this molecule is therefore ‘___ol’
Step 2: Find the prefix
The prefix of organic molecules are strictly dependent on the number of carbon atoms present in the
carbon chain of the molecule.
The example molecule contains 4 carbons in its chain so therefore the prefix is ‘but___‘.
Step 3: Find the position of the functional group
The position of the functional group must be specified in the name of the compound. This is simply the
position of the carbon atom which is bonded to the functional group (-OH in this case).
In our example above, the OH functional group is attached to the second carbon on the carbon chain.
This position number goes between the prefix and suffix in the name of the compound.
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It is important that you make sure that the functional group is closest to the beginning of the carbon
chain. For example, I want you to imagine a chain of 5 carbons: C-C-C-C-C. Now counting from the
LHS, imagine the OH functional group being attached to the fourth carbon on the chain. The position
of the carbon is NOT four. It is in fact one since that would make the OH functional group closest to
the beginning of the chain (counting from the LHS).
ANSWER:
Butan-2-ol
Fuels
There are three main fuels that you need to be familiar with:
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•
•
Coal has been formed by the anaerobic decay of vegetation over millions of years. Anaerobic means
in the absence of oxygen
Natural gas is mainly methane
Petroleum (or crude oil) is a complex mixture of hydrocarbons
Fractional distillation separates petroleum into more useful mixtures of hydrocarbons called fractions.
You need to know the different uses of each of the fractions.
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Alkanes
Properties of alkanes
Alkanes have the general formula:
They are hydrocarbons (compounds that contain only carbon and hydrogen) that only have single bonds.
Since no more atoms can add onto the molecule they are described as saturated hydrocarbons. They
are generally unreactive except for combustion and chlorination (substitution reaction).
Note: The combustion and substitution reactions of alkanes will be covered down below
Alkane homologous series
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Alkane reactions - Combustion The complete combustion of an alkane gives carbon dioxide and
water
The incomplete combustion of an alkane gives carbon monoxide and water
Substitution
Alkanes react with chlorine in bright light to give a mixture of chloroalkanes. One hydrogen atom is
substituted by one chlorine atom.
Provided that there is enough chlorine present, the reaction can continue until all hydrogen atoms in the
compound have been substituted.
Lets take methane as an example:
Alkenes - Properties of alkenes
Alkenes have a general formula:
Alkenes contain carbon double bonds (C=C). Unlike alkanes, alkenes can undergo addition reactions
due to the C=C double bond, and are therefore called unsaturated hydrocarbons.
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Alkene homologous series
Alkene manufacture
Alkenes are made by cracking alkanes. Large alkane molecules obtained by fractional distillation of
petroleum, are passed over a heated catalyst (silicon IV oxide & aluminium oxide).
The idea is that larger alkanes can be broken down into simpler alkanes, alkenes, and possibly hydrogen
But there is more than just one possibility of products. For example:
Hydrogen can also be made as well:
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Alkene reactions
Addition reactions
The C=C double bond in alkenes can be broken to “add” molecules onto the compound. There will
therefore two reactants but only one product formed.
Alkenes can undergo addition reactions with bromine, hydrogen, and steam.
Bromine (test for alkenes)
Aqueous bromine undergoes addition reactions with alkenes. As a result, the original brown colour of
aqueous bromine will turn colourless in the presence of alkenes.
Notice how the C=C double bond in the ethene/ethylene molecule becomes broken and bromine atoms
add onto each of the respective carbons.
Hydrogen
Hydrogen reacts with alkenes to produce alkanes.
The conditions required for this reaction are:
•
•
Temperature 150 degrees
Nickel catalyst
Water (steam)
Water can react with alkenes to make alcohols. This type of reaction is called hydration.
The conditions required for this reaction are:
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•
•
•
Temperature 300 degrees
Pressure 60 atmospheres
Phosphoric acid catalyst
Addition polymerization
Polymerization is the formation of long chain molecules called polymers froma large number of
monomer molecules. down below shows addition polymerization because there is only one product –
The polymer.
Alcohols
ByFree Exam AcademyDecember 26, 2018
Properties of alcohols
Alcohols have the general formula:
All alcohols have the -OH functional group and their names end with -ol.
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Alcohol homologous series
Ethanol manufacture
Ethanol can be manufactured by two methods:
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Catalytic hydration of ethene
Fermentation
Hydration of ethene (ethylene)
This has been covered before in the topic of ‘alkenes‘.
Water can react with alkenes to make alcohols. This type of reaction is called hydration.
The conditions required for this reaction are:
•
•
•
Temperature 300 degrees
Pressure 60 atmospheres
Phosphoric acid catalyst
The advantages of hydration (for ethanol manufacture) are
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•
•
No waste products
Produces ethanol continuously
The disadvantages are
•
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Use of crude oil which is non-renewable
Requires a lot of energy for high heat and pressures
Fermentation
Fermentation is the chemical breakdown of glucose by yeast (or other microorganisms). Carbon dioxide
and ethanol is produced in the process and the reaction is catalysed by yeast enzymes.
The advantages of fermentation
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Renewable (uses plants)
Uses less energy (lower temperature & pressure)
The disadvantages are
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Slow, batch process i.e. needing new batch once yeast die
Produces carbon dioxide as waste
Properties & uses of ethanol
Ethanol burns with blue flame. Combustion of ethanol will produce carbon dioxide and water.
Ethanol can be used:
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As a fuel (i.e. combusted)
as a solvent in perfume and food industries
In some cultures in alcoholic drinks
To make other organic chemicals such as esters
Carboxylic acids
All organic acids (carboxylic acids) have the general formula:
All of them also have the functional group -COOH and their names end in -oic acid.
Carboxylic acid homologous series
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Ethanoic acid manufacture
Ethanoic acid is the second carboxylic acid in the homologous series (look above). Ethanoic acid can
be formed in two ways:
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Oxidation of ethanol via fermentation
Oxidation of ethanol via acidified potassium manganate (VII)
Oxidation via fermentation
Acetobacter bacteria can oxidise (ferment) ethanol into ethanoic acid.
Oxidation via acidified potassium manganate (VII)
When ethanol is heated with the oxidising agent acidified potassium manganate (VII), ethanoic acid is
formed.
Properties of aqueous ethanoic acid
All carboxylic acids (including ethanoic acid) is a weak acid. This means that they demonstrate typical
acid properties (see chapter 8) and they only partially ionize in aqueous solution.
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Formation of esters - Carboxylic acids react with alcohols to form an ester and water.
The conditions required for this reaction are:
•
•
Heat
Concentrated sulfuric acid catalyst
This is a condensation reaction (removal of water). The ‘OH’ from the -COOH functional group reacts
with the ‘H’ from the -OH functional group to form water. The carbon (from -COOH) and oxygen (from
-OH) join together to form the ester bond.
When naming the ester, the alcoholic portion always comes first (i.e. ethyl) and the carboxylic portion
comes second (i.e. propanoate)
Polymers
Polymers are large molecules built from small units (monomers). Different polymers are built from
different monomers and have varying linkages between the monomers.
Synthetic polymers
Synthetic polymers are man-made polymers such as nylon (polyamide) and terylene (polyester).
There are two main methods of polymerization:
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•
Addition polymerization
Condensation polymerization
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Both nylon and terylene are made via condensation polymerization.
Addition polymerization
By breaking apart the double bond of alkenes, repeating units (monomers) can be joined together into
a larger, longer molecule (polymer).
The contents inside the square brackets indicate the repeating unit (monomer) with the letter ‘n’
indicating the total number of the monomers throughout the structure. Take notice of the fact that
the double bond has become a single bond in the polymer.
Condensation polymerization
In a condensation reaction, two monomers react together and join. During the reaction a water molecule
is lost in the process (therefore condensation).
Polyamides (nylon)
In polyamides (such as nylon) the two monomers are always a dicarboxylic acid (or diacid) and a
diamine. These two monomers join together via an amide link as shown below:
Here is the official structure of nylon as drawn in the CIE syllabus:
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Polyesters (tetrylene)
In polyesters (such as tetrylene) the two monomers are a diacid and a diol. These monomers are held
via an ester link as shown below:
Structure of terylene as drawn in the CIE syllabus:
Plastics and pollution
Polymers are non-biodegradable, meaning that they do not decay. They are a major source of visual
pollution and fill up available waste sites. Burning them isn’t good because poisonous gases form as a
result. The best long-term solution to disposal would be to recycle polymer waste.
Natural polymers
Proteins and carbohydrates are important constituents of food. These molecules are natural polymers
that you find in the body.
Proteins
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Proteins are built from amino acid monomers that are joined via amide links (through condensation
polymerization) – Exactly like nylon but different units!
Structure of protein in the CIE syllabus:
There are many different types of amino acids (up to 20) that can be found. Each block above
demonstrates a different amino acid.
As above, amino acids can be joined together via condensation polymerization to form proteins.
Meanwhile, proteins can then be hydrolysed back to amino acids by boiling with hydrochloric acid.
Paper chromatography can be useful in determining what amino acids are present after hydrolysing the
protein.
UNIT 12. Experimental Techniques
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Measurement
1. Time - Digital stop watch - Measures up to 0.01s
2. Temperature - Mercury-in-glass thermometer, alcohol in glass thermometer -Measures up to nearest °C
3. Mass - Electric top-pan balance - Measures up to 0.01g
4. Volume - Beaker -Used to estimate liquid volume
5. Measuring cylinder- Measures up to 0.1cm3 (More accurate than beaker)
6. Pipette - Measures fixed volumes of liquids accurately (i.e. 20cm3)- Measures up to 0.1cm3
7. Burette - Used for measuring variable volumes of liquids accurately - Measures up to 0.1cm3
Criteria of purity
The purity of a substance is defined as the degree to which a substance is undiluted or unmixed with
other material. A pure substance therefore would be made of a single substance. Pure substance will
have sharp melting point and will produce only one spot in chromatography.
Purity assessment from melting point/boiling points
The melting point of a substance is the temperature in which the substance melts. Similarly, the boiling
point of a substance is the temperature at which it boils.
Interestingly, the boiling point and melting point of a substance can give us an indication of how pure
it is. The table below summarizes this quite well (Hodder IGCSE Chemistry Revision Guide):
Paper chromatography
Paper chromatography is a separation technique that is used to separate and identify the components in
a mixture.
How it works is fairly easy. Let’s imagine you have an unknown liquid (Liquid A). You want to find
out whether or not this liquid is impure (i.e. a mixture) and if so, how many substances are in this
mixture and what exactly are they?
Firstly, you simply get a drop of liquid A and place it onto the chromatography paper. You then draw a
horizontal line marking that drop (you’ll see why this is important later).
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You then set up the chromatography paper inside a beaker so that the bottom of the paper
is just immersed inside the solvent (propanone or water). An example of this set up may look like this:
As time passes, the solvent will travel up the chromatography paper. As the solvent moves up, the
sample spot of liquid A will dissolve in the solvent. If liquid A was a mixture, the various substances
inside the mixture will begin to separate because they have different solubilities. Some substances will
travel up the paper slower than others and reach a different end point. The end result may look like this:
In this particular example, it is clear that the ink spot (liquid A) is a mixture. Why? Because you can
see that it has separated out into 3 different components (green, purple, and yellow. If liquid A was pure
then you would only see one component.
*If liquid A was colourless, then the process can be carried out exactly as before but a locating agent
like ninhydrin is required to “locate” all the separated spots later in order to measure the Rf values
Finally, since we know that liquid A is a mixture, we can actually determine what each of the substances
are exactly. To do so, we need to calculate the Rf value of each of the separated components on the
chromatogram.
Rf Value = Distance travelled by spot (from the base line) / Distance travelled by solvent (from the base line)
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•
The Rf value for: spot 3: 5 / 6 = 0.83 spot 2 = 3/6 = 0.50 1 spot = 1/6 = 0.17
All substances have a unique Rf value, and therefore you will be able to find out what exactly the
substance is if you have a reference table. In an examination, they will always provide you with this.
Filtration, centrifuging, decanting
All of these methods are used to separate an insoluble solid from a liquid.
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Filtration is carried out by pouring the mixture into a funnel covered by a filter paper. While the
liquid will pass through the filter, solids will get caught, thereby separating them.
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Centrifuging is a technique which uses a spinning tube. The spinning generates a strong centripetal
force which causes denser materials (i.e. solids) to travel towards the bottom of the centrifuge tube
at a faster rate than normal gravity.
Decanting is the simple process of letting insoluble solids settle in the liquid before gently pouring
it out later.
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Evaporation
This is a simple process of separating the crystals of a solute from a solution. Simply let the solvent
evaporate off and it will leave the solids behind.
Crystallization
This technique is used to separate two soluble solids from a solution (given that they have different
solubility). It works by dissolving the two solids in minimal water, and then slowly cooling it. The less
soluble salt will crystallize first.
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Simple distillation
This method is used to separate a volatile liquid (easy to evaporate) from a solution with a nonvolatile solid. For example, salt water can be purified using this method. The equipment set-up is as
follows:
The heat causes the water to vaporize, leaving the salt behind. The water vapour then turns back into
(pure) water as it passes through the condenser (which cools the vapour down).
Fractional distillation
This method is used to separate the different liquids from a liquid mixture. For example, a mixture of
water and ethanol can be separated using this method.
Fractional distillation works by using the fact that different liquids have different boiling points. It is a
bit more complicated than simple distillation but here is how it works:
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Example: Let’s take the round flask contains a mixture of liquid A (boiling point 50°C) and liquid B
(boiling point 100°C).
As you heat the flask with the Bunsen burner, the temperature of the flask and the fractionating column
will begin to rise. It is really important to understand here that the bottom of the long column is always
going to be hotter than the top. This is simply because we are heating from the bottom and it takes time
for the heat to rise up.
So lets imagine the round flask hits 50°C. What do you think will happen? Well, liquid A will start to
boil in the flask but the vapour won’t get far in the column before cooling back down into liquid (back
into the flask). This is because of the temperature gradient in the column.
The very bottom of the column might be 50°C but the top will be cooler than that, meaning the gas will
just condense back to a liquid before reaching the top. Therefore the main point to take away here is
that the gas can only go up as far as the temperature in the column allows it to.
Eventually however, the column will heat up sufficiently. When the top of the column reaches 50°C,
the vapour of liquid A will reach all the way to the top and get condensed into liquid by the condenser.
Once all of liquid A has entered the flask, we simply need to replace it with an empty one.
The column will then continue to heat up. Eventually, the top of the column will hit 100°C whereby
liquid B will reach the top of the column and become liquefied by the condenser and into our empty
flask.
By utilizing the differing boiling points of liquid A and B, we have successfully separated it from the
original mixture.
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Purification summary
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