preliminary chemistry notes
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Preliminary Chemistry NotesTRANSCRIPT
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Chemical Earth
1. The living and non-living components of the Earth contain mixtures
Construct word equations from observations and written descriptions of a range of
chemical equations
Acid + Base -> Salt + Water
Acid + Metal -> Hydrogen + Salt
Acid + Carbonate -> Salt + Water + Carbon Dioxide
Identify the difference between elements, compounds and mixtures in terms of
particle theory
The particle theory states that all matter consists of particles which are constantly
moving.
Element: An element is a pure substance made up of one type of atom. It cannot be
decomposed into simpler substances
Compounds: Compounds are pure substances made up of two or more elements. It can be
decomposed into simpler substances.
Mixtures: Mixtures are substances made of parts in which the parts keep their own
properties. They are either made of compounds mixed together, elements mixed together
or both. They do not have a definite composition
Homogenous (Uniform composition throughout) e.g. salt water, petrol
Heterogenous (Non-uniform composition throughout) e.g. granite, sand
Identify that the biosphere, lithosphere, hydrosphere and atmosphere contain
examples of mixtures of elements and compounds
The spheres of the Earth Description Examples
Biosphere All living things Blood, Cell sap in plants
Lithosphere The rocks and crusts of the
Earth
Soil
Hydrosphere The waters of the Earth Salt Water
Atmosphere The gases of the Earth Air
Identify and describe procedures that can be used to separate naturally occurring
mixtures of:
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Solids of different sizes – Sieving
Solids and liquids – Filtration (Substance remaining – residue) (Substance passing
through the filter paper – filtrate)
Dissolved solids in liquids – Evaporation or crystallisation
Liquids – Distillation, fractional distillation or decantation
Gases – Fractional Distillation
The process of distillation: The separation technique of distillation can be used to
separate liquids of different boiling points by heating the mixture with a Bunsen burner.
The liquid with the lowest boiling point evaporates first, and then is condensed (water
vapour becoming a liquid) into a condenser to form a liquid which can be collected. The liquid
with the higher boiling point is left behind in a round bottom flask.
The process of fractional distillation:
Assess separation techniques for their suitability in separating examples of Earth
materials, identifying the differences in properties which enable these separations
Mixture separated Method of separation Property used in separation
Solids of different sizes Sieving Particles of different sizes
Solids and liquids Filtration Particles of different sizes
Dissolved solids in liquids Crystallisation Liquid has a lower boiling
point than solid
Liquids Distillation/Fractional Large/Small difference in
boiling points
Gases Fractional Distillation Small difference in boiling
points
Describe situations in which gravimetric analysis supplies useful data for chemists
and other scientists
Gravimetric analysis is quantitative analysis by weight or mass. Examples of gravimetric
analysis include:
1. Mining company – wants to know the composition of an ore sample to see if it is
financially viable to mine the ore
2. Health authority – wants to know the composition of air near an industrial area to
see if the air is polluted
3. Pharmacies – can use the process in mixing various chemicals to manufacture the
drugs.
Apply systematic naming of inorganic compounds as they are introduced in the
laboratory
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Naming binary compounds –
Binary compounds are those that consists of two elements only (Note: Metallic is not a
binary compound)
Ionic Covalent
Metal and non-metal form a
composition
Non-metal and non-metal form a
composition
Metal is written first The non-metal furthest to the left in
the periodic table is written first in
the name
Non-metal modified with ‘ide’
E.g. NaCl – Sodium Chloride
MgO – Magnesium Chloride
Because non-metals have differing
valencies, the prefixes mono, di, tri,
tetra etc are used to show the
numbers of atoms of the non-metal in
the molecule
e.g. CO Carbon Monoxide
CO2 Carbon Dioxide
Covalent Prefixes:
Di-, tri-, tetra, penta, hexa-, hepta-, oct-, non- and dec-
They are used to indicate HOW MANY atoms there are for that element
2. Although most elements are found in combinations on Earth, some are found
uncombined.
Explain the relationship between the reactivity of an element and the likelihood of
its existing as an uncombined element.
The higher the reactivity, the less likely it is to exist as an uncombined element. Less
reactive elements include the noble gases (stable outer shell – no reaction).
Highly reactive elements include the alkali metals (Group 1)
Classify elements as metals, non-metals and semi-metals according to their physical
properties
Metals Non-Metals Semi-Metals
Are solids at room
temperature except
mercury (liquid)
High electrical
conductivity
Shiny lustre
Malleable
(hammered)
Ductile (drawn into
Poor electrical
conductivity
Dull lustre
Brittle (break when
trying to break them)
Non-ductile
Low density, melting
and boiling point
Poor conductors
Properties of metals
and non-metals (also
known as metalloids
e.g. Boron, Silicon)
Low density and
malleability
High lustre e.g.
Silicon is shiny
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wires)
High range melting
and boiling point
Medium range melting
and boiling point
Account for the uses of metals and non-metals in terms of their physical properties
Metals Non-metals
1. Mercury - Used in thermometers –
liquid state expands when heated
2. Iron – used in building construction
and car-making –high tensile strength
and hardness
3. Copper – Good conductors for
domestic appliances
1. Carbon – Good conductors used in
batteries
2. Oxygen - for medical purposes to
help with breathing
3. Argon - gas for filling light bulbs
3. Elements in Earth materials are present mostly as compounds because of
interactions at the atomic level
Identify that matter is made of particles that are continuously moving and
interacting.
Matter is made of particles which are constantly moving and interacting. The movement of
the particles occurs in:
Solids – Particles are compactly packed, they vibrate within the space available (vibrational
movement)
Liquids – In liquids, the particles are all over each other. They move about freely.
Gases – Particles are well separated in space and more freely
Describe atoms in terms of mass number and atomic number
Mass number – the number of protons plus neutrons in the atom
Atomic number – the number of protons in the atom
Describe qualitatively the energy levels of electrons in atoms
The electron is an extremely small negatively charged particle 1/1836 the mass of a proton.
Electrons in an atom move very rapidly and randomly, but not in fixed positions.
Electron configuration: the arrangement of electrons around the nucleus
They are arranged in different shells or energy levels. The more valence shells an atom has,
the more PROTONS, NEUTRONS AND ELECTRONS it contains. Therefore, there is more
energy.
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Describe the formation of ions in terms of atoms gaining or losing electrons
Ions are charged atoms or charged groups of atoms (review chemistry formulae sheets!)
Ions
Monatomic Ions Polyatomic Ions - Radicals
E.g. Na+, Ca2+, K+, Cl- NH4+, CO32-, SO42
Apply the periodic table to predict the ions formed by atoms of metals and non-
metals
1. Cations (Positively charged ions) lose electrons
2. Anions (Negatively charged ions) gain electrons for an octet (stable shell) of
electrons
Apply Lewis electron dot structures to:
- the formation of ions
- the electron sharing in some simple molecules
Lewis symbols are also known as electron dot symbols. It consists of the symbols of the
element and the dots representing the valence electrons.
Note: Practice Lewis dot diagrams
Describe the formation of ionic compounds in terms of the attraction of ions of
opposite charge
Ionic compounds form from a metal (+ve cation) with a non metal (-ve anion) e.g. NaCl. The
ionic bonds are the strong electrostatic attraction between positive and negative ions. The
formula of an ionic compound gives the ratio in which ions are present (empirical formula)
Note: Practice drawing models
Describe molecules as particles which can move independently of each other
Molecules form when two or more atoms bond by sharing electrons (covalent bonds).
Molecules can exist as elements e.g. 02, N2, Cl2 and compounds e.g. H2O, CO2, CH4, SO2
etc.
Within molecule forces – Intramolecular forces (electrolysis)
Forces between the molecules – Intermolecular forces (boiling)
Distinguish between molecules containing one atom (the noble gases) and molecules
with more than one atom
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Molecules
Monatomic Gases
E.g. Noble Gases
It contains full outer shell and they don’t
combine with other atoms. Exists as single
atoms.
Diatomic Gases
E.g. O2, H2, Cl2
Two atoms bonded covalently
Describe the formation of covalent molecules in terms of sharing of electrons
Bonding: The forces that hold an atom together is known as bonding. Bonding is of
different types:
1. Ionic Bonding
2. Metallic Bonding
3. Covalent Bonding.
Ionic bonding – the attractions between the positive and negative ions due to electrostatic
attraction
Metallic bonding – attractions between the positive ions and the sea of delocalised
electrons
Covalent Bonding – The sharing of electrons between non-metal atoms (negative ions,
cations)
4. Energy is required to extract elements from their naturally occurring sources
Identify the difference between physical and chemical change in terms of the
rearrangement of particles
- Physical change is a change of state, with no new product made. It can be easily
reversed, and less energy changes are involved. The particles in a physical
change are not rearranged.
Indications of physical changes include
1. Melting lead
2. Boiling water
3. Solid dissolving in a liquid.
- Chemical change is the formation of a new substance, with a new product made.
It is difficult to reverse, and high energy changes are involved. The particles in
a chemical change are rearranged to form new substances.
Indications of chemical changes include:
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1. New gas formed
2. Precipitate (solid formed)
3. Colour change
4. Significant change in temperature
5. Disappearance of a precipitate
6. Odour is produced
Summarise the differences between the boiling and electrolysis of water as an
example of the difference between physical and chemical change
Boiling and electrolysis of water
Boiling Electrolysis
1. Does not produce any new
substances, just a conversion of a
liquid to gas
2. Easily reversed by cooling the vapour
3. Requires less energy
4. Does not alter the actual particles, it
just separates them from one
another
1. Electrolysis is the process by which
an electric current produces a
chemical change
2. Produces two new substances,
hydrogen and oxygen gases i.e. H2
and O2
3. Difficult to reverse
4. Requires much more energy for the
two gases to be decomposed
5. Breaks the particles up into H2 and
O2
Identify light, heat and electricity as the common forms of energy that may be
released or absorbed during the decomposition or synthesis of substances and
identify examples of these changes occurring in everyday life
1. Light given off often in oxidization
2. Heat given off in all exothermic reactions
3. Electricity given off in reactions occurring within a battery
Decomposition (breaking of a chemical substance into simpler substances) occurring in
everyday life:
1. Sodium Azide decomposes to sodium and N gas, air bag
2. Limestone, calcium carbonate, is decomposed to CO2 and CaO to make glass, cement
etc
Synthesis (formation of a compound from simpler compounds or elements) in everyday life:
1. Rust Fe FeO
2. Burning coal, C CO2
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Explain that the amount of energy needed to separate atoms in a compound is an
indication of the strength of the attraction, or bond, between them
The stronger the chemical bonding in a compound, the more energy that is required to
break the compound into atoms. Alternatively, the stronger the chemical bonding in a
compound, the more energy that is released when the compound is formed from its
atoms.
5. The properties of elements and compounds are determined by their bonding
structure
Identify the differences between physical and chemical properties of elements,
compounds and mixtures
- Physical properties are those related to changes of state and physical changes,
including lustre, hardness, ductility, conductivity etc.
- Chemical properties include relating to chemical changes, including reactivity and
valency, which bonds will work and which won’t etc.
Describe the physical properties used to classify compounds as ionic or covalent
molecular of covalent network
Molecular solids Lattice Solids
Covalent
molecular
Metallic Ionic Covalent
network
Melting and
boiling points
Low Variable High Very high
Conduct
electricity?
No Yes As solid: no
Molten: yes
Dissolved:
Yes
No
Hardness and/or
workability
Soft Variable
hardness;
malleable and
ductile
Hard and
brittle
Hard and
brittle
Forces holding
particles
together in the
solid
Intermolecular
(between
molecules)
Delocalized
electrons
(metallic
bonding)
Electrostatic Covalent
bonding
throughout the
crystal
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Distinguish between metallic, ionic and covalent bonds
Metallic bonds are held together by a sea of delocalized electrons, with positive ions
(cations)
Ionic bonds are held together by the electrostatic attraction of positive and negative ions
Covalent bonds are held together between anions (negatively charged ions) with weak
intermolecular forces
Describe metals as three-dimensional lattices of ions in a sea of electrons
Metallic bonding results in an orderly 3D array of positive ions held together by a ‘sea’ of
delocalised electrons (valence electrons only). These electrons move freely throughout the
lattice, holding it together and causing the metal to be a conductor of electricity. These
delocalised electrons also hold the metal together when distorted – thus making metals
malleable and ductile. The lattice formation makes metals hard.
Describe ionic compounds in terms of repeating three-dimensional lattices of ions
Ionic bonding forms crystals; the electrostatic attraction between the opposite charges
extends throughout the entire lattice. This strong attraction makes ionic substances hard,
but also brittle; distorting the crystal bring opposite charges together – they repel each
other – causing the crystal to shatter. This orderly array means that in solid form, ionic
substances do not conduct electricity; the ions are not free to move towards a charged
electrode. However, when melted or dissolved in water, the arrangement of ions is broken
up, allowing the ions to move towards an electrode, hence conducting electricity.
Explain why the formula for an ionic compound is an empirical formula
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Empirical Formula is the lowest ratio formula. It does not tell you the number of atoms
present
For example, NaCl means that for every 1 Na there is 1 Cl (not the exact amount of each
element)
¬ The empirical formula is used for ionic compounds because the size of the lattice is unknown therefore a ratio must be used
Another example is glucose
C6H12O6 is its molecular formula, whereas its empirical formula (lowest ratio) is CH2O
Molecular Formula does tell you the number of atoms present.
Identify common elements that exist as molecules or as covalent lattices
Covalent Molecules:
- Noble gases (monatomic)
- H2, O2, F2, N2 and Cl2 (diatomic gases)
- Br2 (liquid) and I2 (solid)
Covalent Lattices:
- Carbon; diamond (3D) and graphite (2D)
Explain the relationship between the properties of conductivity and hardness and
the structure of ionic, covalent molecular and covalent network structures
Ionic Lattices: the strong electrostatic attraction between the ions makes ionic
substances hard, and they are brittle because any distortion of the crystal brings opposite
charges into contact, and they repel each other. As solids, they do not conduct electricity
as their ‘orderly array’ does not allow the ions to migrate towards a charged electrode.
When melted or dissolved in an aqueous solution, the orderly arrangement is broken up,
allowing for the movement of ions towards the oppositely charged electrode.
Covalent molecules: covalent molecular substances are soft, because their intermolecular
forces (forces between pairs of molecules) are weak. These weak intermolecular forces also
mean that covalent molecular have low melting and boiling points. Because covalent
molecules are a neutral species, they do not conduct electricity as a pure substance or in a
solution – however, some covalent substances react with water when mixed, and form ions
capable of conducting electricity. Eg hydrogen peroxide hydrochloric acid.
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Covalent network solids: covalent lattices are very hard and brittle (e.g. diamond, quartz),
due to the covalent bonds that extend throughout the entire lattice. Apart from graphite,
covalent lattices do not conduct electricity, as they do not contain any ions, and all the
electrons are tied up in covalent bonds.
Metals
1. Metals have been extracted and used for many thousands of years
Outline and examine some uses of different metals through history, including
contemporary uses, as uncombined metals or as alloys
The Copper Age was 3200BC to 2300BC. It is the period that archaeological records
indicate that copper was the first metal to be extracted from its ore. Copper was heated
with charcoal and globules of copper formed. Molten copper was used to make ornaments
and domestic utensils.
The Bronze Age was 2300BC to 1200BC. It was later discovered that heating copper with
tin produces an alloy, bronze. Bronze was harder than copper and more easily melted to be
molded due to its low melting point. Bronze was used for tools and weapons.
The Iron Age was 1200BC to 1AD. Iron is more reactive than copper, so it need a higher
temperature to melt. Hematite was mixed with charcoal in primitive furnaces by blowing air
and obtaining a sufficiently high temperature. By 1000BC, iron had replaced bronze for
tools and weapons because it was harder and had hard tensile strength.
The Modern Age is 1Ad to present. There had been more extraction and uses of other
metals such as aluminium, chromium and metal alloys. Iron is the most widely used metal
today. Many other metals have come into common use due to the advancement in extraction
technology.
Describe the use of common alloys including steel, brass and solder and explain how
these relate to their properties
Alloys Properties Uses
Brass (50-60% copper with
zinc)
Lustrous gold appearance
Hard but easily machined
Plumbing fittings
Musical instruments
Decorations
Bronze (80-90% copper with
tin)
Hard
Resists corrosion
Easily cast
Ships’ propellers
Casting statues
Solder (30-60% tin with
lead)
Low melting point
Adheres firmly to other
metals when molten
Joining metals together in
plumbing and electronics
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Steels
Mild Steel
Structural Steel
High – carbon Steel
Stainless Steel
Soft, malleable
Hard, high tensile strength
Very hard
Hard, resists corrosion,
lustrous appearance
Car bodies, pipes, nuts and
bolts, roofing
Beams and girders, railways,
concrete reinforcement
Knives and tools such as drill
bits, chisels, hammers
Food processing machinery,
kitchen sinks and appliances,
cutlery, surgical instruments
Explain why energy input is necessary to extract a metal from its ore
Energy such as electricity and heat is required to extract a metal from its ore in order to
break the chemical bonds within the compounds. The higher the chemical, the more energy
is required to break the chemical bonds.
Identify why there are more metals available for people to use now than there were
200 years ago
Many metals have been available for use to due lower cost of generating electricity and
more advanced in commercial extraction techniques. Two hundred years ago, there was a
lack of extraction technology and scarcity of metals and resulted in only a limited amount
of metals being able to be extracted and used. Some metal ores have very high melting
points and it would have been difficult to reach a very high melting point two hundred years
ago with the lack of technology.
2. Metals differ in their reactivity with other chemicals and this influences their uses
Describe observable changes when metals react with dilute acid, water and oxygen
- Reactions of metals with oxygen – the majority of metals will react with oxygen in
the air at room temperature to form metallic oxides. When metals react with gases
in the atmosphere to form new substances, they go through corrosion. This
corrosion can cause the metal to lose some of its strength.
- Reactions of metals with water – most metals when placed in cold water undergo no
observable changes and show no sign of chemical reaction. Most metals combine with
water to form hydrogen and a metal hydroxide.
- Reactions of metals with acids – during a reaction between a metal and an acid, the
metal dissolves as it loses electrons and forms cations. Hydrogen ions from the acid
gain electrons to form hydrogen gas. This reaction involves the transfer of
electrons and it is a redox reaction.
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Describe and justify the criteria used to place metals into an order of activity
based on their ease of reaction with oxygen, water and dilute acids
Graphical representations of ionization energy can be used to represent the periodic trends
in ionization energy. Emphasise that ionization trends across the period can be related to
the effective nuclear charge; whilst trends inionization energy down a group of
representative elements can be related to the size of the atom.
Identify the reaction of metals with acids as requiring the transfer of electrons
During the reaction between a metal and an acid the metal dissolved as it loses electrons
and forms positively charged ions. Hydrogen ions from the acid gain electrons to form
hydrogen gas. As this reaction involves a transfer of electrons it is an redox reaction.
Outline example of the selection of metals for different purposes based on the
reactivity, with a particular emphasis on current developments in the use of metals
- Magnesium is a highly reactive metal and some of its used is a result of reactivity.
Magnesium is used in the cathodic protection of less reactive metals to protect
them from corrosion. Magnesium is called a sacrificial anode.
- Calcium is highly reactive and is restricted to situations where its reactivity can be
used as an advantage. Calcium is added to steels to remove any remaining traces of
oxygen, sulfur and phosphorus.
- The reactivity of zinc makes it suitable for use in batteries such as dry cells and
button batteries. In these cells the zinc is oxidized and the electrons it loses travel
through an external circuit producing as electric current.
Outline the relationship between the relative activities of metals and their position
on the Periodic Table
The activity series for metals from least reactive to most reactive are Au, Pt, Hg, Ag, Cu,
Pb, Sn, Ni, Co, Cd, Fe, Cr, Zn, Al, Mg, Ca, Na and K. a comparison of the activity series for
metals with the position of these metals on the periodic table show some trends. The most
reactive metals are generally found on the left side of the periodic table whereas the least
reactive metals tend to be found in the middle of the periodic table.
Identify the importance of first ionization energy in determining the relative
reactivity of metals
Ionization energy is a measure of the energy needed to remove an electron from the
electro – static attractive force of the positively charged nucleus. The ionization of an
atom or ion is defined as the amount of energy required to remove the most loosely bound
electron from the atom of ion in gaseous state. The energy required to remove the first
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electron from an atom is called the first ionization energy. Reactive metals tend to have low
ionization energies and less reactive metals have higher ionization energies.
3. As metals and other elements were discovered, scientists recognized that patterns
in their physical and chemical properties could be used to organize the elements in
the periodic table
Identify an appropriate model that has been developed to describe atomic structure
Bohr’s model is an appropriate model to describe the atomic structure. The nucleus is the
central part of the atoms which contains the protons and neutrons. It has a positive equal
charge equal to the number of protons. The electrons move through a relatively large space
outside the nucleus. The electrons are kept moving around the nucleus by attractive electro
– static forces between the positively charged nucleus and negatively charged electrons.
Outline the history of the development of the Periodic Table including its origins,
the original data used to construct it and the predictions made after its
construction
In the 1800s, 30 naturally occurring chemical elements were known.
French chemist, Antoine Lavoiser classified the elements into two groups, metals and non –
metals based on their physical properties.
In 1829, a German chemist, Dobereiner recognized the similarities of several groups of
three elements in which he called the triads.
In 1864, an Englishman, John Newlands, proposed the law of octaves where the elements
were ordered according to their atomic weight.
In 1869, Mendeleev proposed the periodic law where the properties of the elements vary
periodically with their atomic weight. He arranged the elements with increasing atomic
weight and grouped them with elements with similar properties. Mendeleev knew that there
were still more elements to be discovered and left spaces in his periodic table.
In 1914, a British chemist, Henry Moseley, proposed a modified periodic law where the
properties of the elements vary periodically with their atomic numbers.
Explain the relationship between the position of elements in the Periodic Table and:
o Electrical conductivity
Across a period, the electrical conductivity of elements decreases because elements are
less metallic. Non metals do not have free mobile electrons in their crystal lattice. Down a
group, the electrical conductivity of elements increases because they are more metallic.
Down a group, the valence shell is further away from the nucleus and can more easily escape
into the lattice.
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o Ionization energy
Ionization energy in the energy required to remove an electron from an atom of the
element in the gaseous state. Across a period, the ionization energy increases because the
atomic radius decreases across a period. The valence electrons closer to the nucleus
experience a stronger nuclear pull. Down a group, the ionization energy decreases because
the atomic radius is bigger and outer electrons are not as attracted to the nucleus of
atoms.
o Atomic radius
The atomic radius is the average distance from the nucleus to the valence shell. Across a
period, the atomic radius decreases as the valence shells are closer to the nucleus. Down a
group, the atomic radius increases because the number of electron shells increases.
o Melting point and boiling point
Across a period, the melting point increases from group I to group IV the decrease from
group IV to group VIII. The lattice changes from metallic bonding to covalent network and
then covalent molecular. Down a group, it decreases from groups I to IV and increases from
groups V to VIII
o Combining power (valency)
The combining power of a group increases down the periodic table. Across the periodic
table, the combining power decreases.
o Electronegativity
Electronegativity is the tendency of an atom of an element to attract electrons. Across a
period, the Electronegativity increases as the metallic character decreases. Down a group,
the Electronegativity decreases as the metallic character increases.
o Reactivity
The reactivity of elements down a group increases and it decreases as it goes across a
period.
4. For efficient resource use, industrial chemical reactions must use measured
amounts of each reactant
Define the mole as the number of atoms in exactly 12g of carbon-12 (Avagadro’s
number)
Atoms and molecules are too small to weigh out individually. Chemists measure the amount
of any substance in terms of moles. A mole is defined as the amount of a substance that
contains the same number of particles as there are atoms in exactly 12 g of carbon of
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Carbon – 12. Chemists have determined that the number of atoms in 12 g of carbon – 12 is
6.02 x
Compare mass changes in samples of metals when they combine with oxygen
Metals exhibit very varied reactivities in their reactions with oxygen. An example, lithium
and sodium and potassium tarnish rapidly when exposed to air and must therefore be stored
in liquid paraffin old. Also other metals react with oxygen and explode. Due to the
different reactivities of metals, the less reactive a metal is, then the more the metal
weighs. If a more reactive metal reacts with oxygen, then it can result in an explosion and
it is less weight.
Describe the contribution of Gay – Lussac to the understanding of gaseous reactions
and apply this to an understanding of the mole concept
Gay – Lussac found that gases always combine in simple whole number ratios. French
chemist, Gay – Lussac was conducting an experiment with gases and determines the volume
in which they combined.
Gay – Lussac’s law of combining gas volume states:
When measured at constant temperature and pressure, the volumes of gases taking part in
the chemical reaction who simple whole number ratio to another.
Recount Avagadro’s law and describe its importance in developing the mole concept
Amadeo Avogadro proposed in 1811 that elements could exist as atomic aggregates called
molecules.
Avagadro’s law states that under the same condition as temperature and pressure, equal
values of all gases contain the same number of molecules.
Due to Avagadro’s contribution, the number of atoms or molecules in 1 mole is called
Avagadro’s number – 6.02 x 1023
Distinguish between empirical formulae and molecular formulae
The empirical formula of a compound is the simplest whole number ratio of the numbers of
atoms of each element in the compound. The molecular formula specifies the actual number
of atom of each element in a molecule. E.g. the compound, hydrogen peroxide has the
molecular formula of . The molecule contains two hydrogen atoms and two oxygen
atoms bonded together. The empirical formula of hydrogen peroxide would be HO.
5. The relative abundance and ease of extraction of metals influences their value and
breadth of use in the community
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Define the terms mineral and ore with reference to economic and non-economic
deposits of natural resources
Minerals are naturally occurring inorganic substances, usually compounds with a particular
chemical composition and a definite crystal structure. Examples of minerals include
hematite, magnetite, gibbsite, boehmite, malachite and chalcopyrite.
Ores are naturally occurring deposits that are mixtures of minerals from which a substance,
usually a metal can be economically extracted. Examples of ores include bauxite and iron
ore.
Describe the relationship between the commercial prices of common metals, their
abundances and relative costs of production
The commercial price of metals depends on a few factors including their relative
abundances and the cost of production.
The greater the abundance of a metal the lower the commercial price of the metal would be.
The cost of production of the metals depends on where it is located and the amount of
energy input. If the location of the ore is located in a high population zone, the mining
procedure would be difficult because there would be damages done to the environment and
increase the cost of production.
If an ore is located in remote places, then the cost of production would increase because it
would cost money to transport the raw materials to refinery plants.
The more reactive the metal is, then the higher the energy input is needed for extraction
and it would increase the cost of extraction.
Explain why ores are non-renewable resources
Ores are deposits of naturally occurring minerals which were formed during the evolution
of the universe and the planets; therefore they are non – renewable resources.
Describe the separation processes, chemical reactions and energy considerations
involved in the extraction of copper from one of its ores
- Mining, crushing and grinding
The mined ore (containing ammonium of 6.5% of copper by weight) us placed in a crusher
and converted to pebbles. The pebbles are then grounded in a grinding mill to liberate
the mineral crystals from the rock
- Concentration
Using froth flotation, 30% of the copper is obtained by weight.
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- Roasting and smelting
It is roasting in the air.
The mixture is then heated to a sufficiently high temperature to produce material from
which the required metal can be obtained.
The mixture of Copper (I) sulfide and Iron oxide with sand is heated to a sufficient
high temperature where it produces two immiscible liquids.
The liquid is removed. The copper (I) sulfide is then heated on its own to a higher
temperature while air is bubbled through it. This reduces sulfide to copper metal and
sulfur dioxide is produced.
The liquid copper is left to cool and solidify.
Recount the steps taken to recycle aluminium
- Collect the used products from homes, shopping centres and factories
- Transport the collected material to a central processing plant
- Separate the required metal from the impurities
- Re – melt the metal into stock ingots and transport them to product manufacturers
Water
1. Water is distributed on Earth as a solid, liquid and gas
Point 1.1 – Solute is the substance dissolved in a solution. Solvent is the substance
which does the dissolving. A solution is a homogenous mixture in which the dispersed
particles are so small (molecules or ions) that they never settle out.
Point 1.2 – Water is important as a solvent because:
Most of the chemical reactions responsible for ‘life’ occur in water solutions
(aqueous solutions).
Water carries waste products away from cells.
Used in many household products.
Aqueous solutions are involved in the production of chlorine, common fertilisers
and zinc.
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Point 1.3 -
Point 1.4 - Earth is the only planet in our Solar System with abundant water. Water is
important on Earth because:
It is necessary for all forms of life (major constituent of cells – used as a
solvent in which reactions take place and for metabolism)
It is a habitat with few temperature extremes for many forms of life
It is a major force in the shaping of the planet, as a liquid and a solid (i.e.
weathering and erosion)
It is a natural resource for humans and other organisms.
2. The wide distribution and importance of water on Earth is a consequence of its
molecular structure and hydrogen bonding.
Point 2.1 –
Point 2.2 – Water is bent shape. Contains hydrogen bonding and therefore has higher
melting and boiling points than ammonia and hydrogen sulphide.
Ammonia is pyramidal shape. It also contains hydrogen bonding. It also has a high
melting point and boiling point than expected.
Hydrogen sulfide is bent shape. Instead of hydrogen bonding, it contains dipole-dipole
forces and therefore has lower boiling and melting points than ammonia and water.
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Point 2.3 – Hydrogen bonding only occurs when H atoms of one substance is attracted
to F, O or N atoms of another substance. This is the cause of high melting and boiling
points of a substance.
Point 2.4 – Water molecule is a polar molecule because in the actual bond, there is
uneven sharing of electrons between H and O atoms. O atom has a higher level of
electronegativity (wanting of electrons) than H. Therefore water has polar bonds. The
fact that water is bent in shape, it confirms water a polar molecule. In the molecule
there is a slightly negative end and a slightly positive end.
Point 2.5 – Because polar molecules have slightly positive and negative regions, they line
up so that opposite poles line up and so electrostatic attractions hold the molecules
more strongly than dispersion forces. These forces are called dipole-dipole
interactions.
Point 2.6 – Water has a high surface tension due to strong intermolecular forces
(hydrogen bonding). Water also has a high level of viscosity because it has hydrogen
bonding. The H atoms are attracted to O atoms in the glass and therefore water has a
high resistance of flow than most other liquids. Water has high melting and boiling
points because of hydrogen bonding. Lots of energy is needed to break the hydrogen
bonds and this is why water has a higher mp and bp than most other similar substances.
3. Water is an important solvent in biological systems, transporting materials into and
out of cells.
Point 3.1 – When a soluble ionic compound (such as sodium chloride) interacts with
water, they break up into positive and negative ions. These ions move freely and
independently to each other.
When a soluble molecular compound (such as sucrose) interacts with water, the
crystals of the solid break up and disperse throughout the solvent (water) and they
break down to the molecular level.
When a soluble or partially soluble molecular element or compound (such as iodine,
oxygen or hydrogen chloride) interacts with water, the solvent-solute interactions are
weak dispersion forces and this is why the solubilities of such substances are quite low.
When a covalent network structure substance (such as silicon dioxide) interacts with
water, nothing happens because water is not able to break the strong covalent bonds
between the particles (atoms) in these lattice solids.
When a substance with large molecules (such as cellulose or polyethylene) interacts
with water, nothing happens because water is not able to break the strong covalent
21
bonds between the particles (molecules) in these solids. However some large molecules
such as amylose and glycogen are soluble in water as it contains F, O or N atoms which
form hydrogen bonding with the water.
Point 3.2 – Remember like dissolves like. Substances which are polar molecules are
soluble in water as water is a polar solvent. Forces between the substance and the
water may be hydrogen bonding, dipole-dipole force.
4. The concentration of salts in water will vary according to their solubility and
precipitation can occur when the ions of an insoluble salts are in solution together.
Point 4.1 -
Point 4.2 –
Point 4.3 – When an ionic substance dissolves in water, it breaks up into ions which move
independently through the solution. When a solution becomes saturated, ions still continue
to break away from the crystals of solid and go into solution but in addition an equal number
of ion pairs from the solution precipitate out on to the solid. When this has happened,
dynamic equilibrium has been achieved. This is when a solid is in contact with its saturated
22
solution there is a dynamic balance between dissolution and precipitation: both are
occurring but at equal rates so that there is no overall change in concentration in the
solution.
Point 4.4 - The molarity of a solution is the number of moles of solute per litre of
solution.
Eg. 17.54g of barium hydroxide was dissolved in water and made up to 500mL.
Calculate the molarity of the solution.
First step: Find the number of moles of Ba(OH)2
Use formula n = m/M
n = 17.54/171.316
n = 0.102 moles
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Second step: Calculate molarity
Use formula c = n/v
c = (0.102)/ (0.5)
= 0.20 mol/L
Point 4.5 – A variety of ways of expressing concentration is used because each method
has advantages for particular situations. In commerce and industry and in shopping
where the main concern is with how much solute is present, then mass per unit volume is
very convenient. In environmental contexts concentrations are usually very low. Masses
per unit volume or percent compositions generally lead to very small numbers so parts
per million (ppm) gives more manageable numbers.
Point 4.6 –
Pb2+ (aq) + 2NO3-
(aq) ⇌ Pb(NO3)2 (s)
Forward Reaction (precipitation – ions in solution
precipitate out onto solid).
Reverse Reaction (dissolution – ions in solid break away
into solution).
Point 4.7 – Make sure for neutral, full ionic and net equations, the state of each
substance is written down. NOTE: The solid remains unchanged in the full ionic and net
equations.
Point 4.9 –
Eg. 17.54g of barium hydroxide was dissolved in water and made up to 500mL.
Calculate the molarity of the solution.
First step: Find the number of moles of Ba(OH)2
Use formula n = m/M
n = 17.54/171.316
n = 0.102 moles
Second step: Calculate molarity
Use formula c = n/v
c = (0.102)/ (0.5)
= 0.20 mol/L
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5. Water has a higher heat capacity than many other liquids
Point 5.1 – The specific heat capacity of a substance (C) is the heat needed to raise the
temperature of 1 gram of the substance by 1 Kelvin (or ºC). It is measured in Joules per
Kelvin per Gram.
Point 5.2 – The specific heat capacity of water is much higher than other common
solvents. The specific heat capacity of water is 4.18 Joules per Kelvin per Gram.
Point 5.3 -
E.g.
How much energy will be required to raise the temperature of 1.00 L of water from 17 C to
100 C?
Change in temperature = 83ºC
Specific Heat Capacity of Water = 4.18
Mass = 1000g
Energy = (83) x (4.18) x (1000)
= 346 940 J
= 346.94 kJ
Point 5.4 – Water is widely used to measure energy changes in chemical reactions and as
a result of its ability to absorb heat.
Point 5.5 – When sodium hydroxide dissolves in water, the solution heats up. The
dissolution process releases heat which then warms up the solution. It is exothermic.
Also when dissolving lithium bromide or sulfuric acid in water is exothermic.
Point 5.6 - When potassium nitrate dissolves in water, the solution cools. Dissolving
potassium nitrate in water requires an input of energy: this energy is taken from the
normal thermal energy of the water and solid substance so the mixture (solution)
becomes colder. It is endothermic. Also dissolving ammonium chloride or silver nitrate
is endothermic.
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Point 5.7 - For aquatic organisms, water’s high heat capacity (ability to absorb heat)
means that their environment (such as lakes, rivers, oceans and ponds) maintains a much
more stable temperature than the surrounding atmosphere or land. For example on a
hot day, the temperature on land can reach 50ºC making it difficult for organisms on
land to survive. However the temperature of the water may only be 25ºC. This allows
aquatic organisms to thrive.
Not only is water’s ability to absorb heat important to aquatic organisms, it is also
important to life on earth generally. Water is such a large component of the biosphere.
As a result it has a moderating influence on global temperatures, stabilising the day-to-
night and summer-to-winter temperature fluctuations. This in turn produces a more
hospitable and friendly environment for all life forms.
Point 5.8 - Thermal pollution is the harm to the environment (lakes and rivers) resulting
from the release of excessive waste heat. Increase in temperature as a result of
thermal pollution, can kill aquatic life as they cannot cope with the temperature rise.
Energy
1. Living organisms make compounds which are important sources of energy.
Point 1.1 - Photosynthesis is the process in which plants use energy from the sun (light
energy) to convert carbon dioxide from the air and water from the ground into
carbohydrates such as glucose and starch. It is an endothermic reaction.
Equation of photosynthesis:
Carbon Dioxide + Water Glucose + Oxygen
6CO2 (g) + 6H2O (l) C6H12O6 (aq) + 6O2 (g)
The solar energy collected by the plants is converted into chemical energy in the process of
photosynthesis.
Point 1.2 - Carbohydrates are mainly sugars and starches constituting one of the three
principal types of nutrients used as energy sources (calories) by the body. They
originate from the photosynthesis of plants. It is a product in this process. They are
converted for all life to use (animals, humans).
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Point 1.3 – Coal is formed from plant material buried in swampy conditions. There is
chemical energy stored in the plant material in the coal. This energy had been converted
from light energy in photosynthesis. Natural gas and petroleum are formed from
phytoplankton, zooplankton and algae buried on the sea floor under pressure.
Point 1.4 – Compounds found in coal include:
Hydrocarbons
Carbon dioxide
Organoarsenic compounds
2. There is a wide variety of carbon compounds
Point 2.1 – Carbon is located in Group 4 in the Periodic Table and its electronic
configuration is 2, 4; which means it has 4 valence electrons (outer shell).
Point 2.2 - The allotrope diamond consists of carbon atoms each covalently bonded to
four other carbon atoms. In diamond, the shape around each carbon atom is tetrahedral.
The carbon atoms are arranged in six-membered rings; the rings are buckled and not
flat. The three dimensional structure means that diamond is very hard. With its valence
electrons tied up in strong covalent bonds, diamond has no mobile electrons and so it
does not conduct electricity. The orderly arrangement of the atoms throughout the
whole crystal gives its transparency and brilliance.
The allotrope graphite is also a covalent lattice but unlike diamond each carbon atom is
bonded to only three other carbon atoms to form a planar structure. The structure of
graphite consists of flat six-membered rings. The three bonds per carbon atom mean
that it leaves each carbon with a free valence electron and as a result graphite unlike
diamond can conduct electricity. Graphite is packed in layers. Because they are only
weak intermolecular forces between these layers, they can easily slide across one
another and this explains the slipperiness of graphite and its good lubricating
properties.
Point 2.3 – Carbon can form single, double and triple bonds with other carbon atoms.
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Point 2.4 – Carbon has a valency power of 4. As shown above carbon can form single,
double and triple bonds and as a result carbon forms a huge array of compounds.
Point 2.6 –
Uses - Diamond Related physical properties
Diamond is widely used for jewellery.
(rings, necklaces)
This is because of their brilliance sparkle
and hardness (diamond cannot be dulled by
scratches).
Diamond is used for drills and cutting
implements. (tips of drills)
This is because diamond is extremely hard
and resists corrosion.
Diamond is also used in long-lasting dies
for drawing fine wire (e.g. for a light globe
filaments)
This is because diamond is extremely hard,
has a very high melting point and it has a
high resistance to any chemical attack.
Uses - Graphite Related physical properties
Graphite is used as electrodes in ordinary
and alkaline dry-cell batteries.
This is because graphite conducts
electricity.
Graphite is used as a dry lubricant (often
on door catches in motor cars) and in the
‘lead’ of lead pencils.
This is because graphite is very slippery
(as the planar layers slide over one
another).
Graphite is used for making kitchen This is because graphite resists corrosion.
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benches and other household surfaces.
Graphite is also used as a chemical plant
structural material (e.g. absorption towers
in the manufacture of hydrochloric acid
are built from graphite blocks).
This is because graphite is a good
conductor of heat, low coefficient of
thermal expansion, easily machined, has a
very high melting point, durable at high
temperatures and because it is very
resistant to chemical attack.
3. A variety of carbon compounds are extracted from organic sources
Point 3.1 - Fractional distillation involves vaporising the petroleum by heating it to 350oC
in a fractionating column. Using fractional distillation, components of petroleum are
separated by according to their boiling points. The components with low b.p vaporise to
the top of the fractioning tower and condense back.
Fraction Boiling Point (ºC) Carbon atoms per
molecule
Major uses
Gases Less than 30 1 to 4 Liquefied
petroleum gas
(LPG)
Petroleum ether 30 – 80 5 to 6 Industrial solvents
Gasoline 70 – 200 6 to 12 Motor fuel
Kerosene 175 – 250 12 to 16 Jet fuel, domestic
heating
Gas oil 250 – 350 15 to 18 Diesel fuel,
industrial and
domestic heating
Lubricating oil Greater than 350 18 to 25 Motor oils
Greases Greater than 350 Greater than 20 Lubrication
Asphalt and tar Residue Greater than 25 Road-making,
roofing
Point 3.2 –
C1 meth- C6 hex-
C2 eth- C7 hept-
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C3 prop- C8 oct-
C4 but- C9 non-
C5 pent- C10 dec-
Single bonds = alkanes
Double bonds = alkenes
Triple bonds = alkynes
Points 3.3 and 3.4 –
In terms of molecular structure…
Non-polar covalent bonds (C – C is non-polar, C – H is slightly polar but geometry
(shape) tends to cancel it out).
Only intermolecular forces are dispersion forces.
Dispersion forces increase as molecular weight (no. of C atoms) increases.
These bonds explain the following properties…
Melting and Boiling Point
Melting & boiling point increases as number of C atoms increases (as molecular
weight increases, dispersion forces increase)
Melting & boiling points of the alkanes are higher than the corresponding alkene
(lower molecular weight of alkene lower dispersion forces)
Solubility
Insoluble in water (non-polar), soluble in non-polar solvents
Volatility
Volatility is the ease at which a substance can be converted to a vapour.
Volatility decreases as molecular weight increases. (Molecular weight increases
dispersion forces increases)
Other
Density < 1 g/cm3 (ie. floats on water)
Do not conduct electricity (no free electrons)
Generally unreactive (except combustion)
Homologous series:
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Family of compounds that can be represented by one general formula
Common functional group
Similar structures & chemical properties
Gradation in physical properties based on order of molecular weight (eg. boiling point)
Functional Group:
Atom or group of atoms that determines the chemical behaviour of the compound.
Eg.
Alkanes – single bonds
Alkenes – double bonds
Alkynes – triple bonds
Point 3.5 – Because alkanes and alkenes contain weak dispersion forces, they are very
volatile. (i.e. vaporise very easily – have low boiling points). As a result safety
precautions must be taken:
Well-maintained cylinders and fittings for gaseous hydrocarbons
Added odours for early detection of leaks
Sturdy containers for liquids
Minimise the quantity in everyday use
Do not handle these liquids in confined places
Keep hydrocarbons away from naked flames or sparks
Use fume hoods
4. Combustion provides another opportunity to examine the conditions under which
chemical reactions occur.
Point 4.1 – Indicators of chemical reactions:
Gas evolved
Precipitate formed
Significant temperature rise
Disappearance of a solid
Odour produced
Light produced
Colour change
Point 4.2 – Combustion is an exothermic chemical reaction because it releases heat and
new products are formed.
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Point 4.3 – In chemical reactions, bonds are broken and rearranged to form new bonds.
Reactants are altered to produce new products. Therefore the molecular configuration
is changed.
Point 4.4 - Energy is absorbed from the surroundings to break the bonds (endothermic).
Energy is released when bonds are formed (exothermic).
Point 4.5 – Activation energy is the minimum amount of energy needed to start a
reaction. On an energy profile diagram, activation energy is the space between peak and
reactants.
Point 4.6 – Energy profile for exothermic reaction: reactants higher than products.
Energy profile for endothermic reaction: reactants less than products.
Point 4.7 – The ignition temperature of a substance is the temperature above which
spontaneous combustion will start. (I explained activation energy earlier). There are
both minimum temperatures required for molecules to start reacting.
Point 4.8 – Partial combustion occurs in ‘fossil’ fuels when they are burnt in limited
oxygen for example when the sleeve of a Bunsen burner is partly or fully closed. In the
case of petrol in cars in cities when a temperature inversion occurs ozone and
peroxyacetyl nitrate (PAN) as well as other oxides of nitrogen are produced causing
photochemical smog. This can be avoided by allowing more oxygen and in the case of the
Bunsen burner opening the hole.
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Point 4.9 –
5. The rate of energy release is affected by factors such as types of reactants.
Point 5.1 –
Slow:
Slow combustion occurs when we use big lumps of fuel and limit the supply of air
(oxygen gas). This means that burning occurs only on the surface of the big lumps
and its speed is controlled by the limited supply of air.
Spontaneous:
Spontaneous combustion occurs when a substance catches on fire without the
application of heat from the outside. The oxidation between the two substances in
contact starts the fire.
An explosive reaction reacts much more rapidly than a normal combustion reaction.
The conditions in which it reacts under are: an excess amount of heated air and high
pressure. For example in petrol engines a spark is used to ignite a heated mixture of
petrol and air. In this case, the conditions used to promote a very rapid reaction.
Point 5.2 – Collision theory: molecules must collide (energy) and at orientation (must
collide at right position). The more successful collisions there are between the
molecules, the faster the reaction rates.
Point 5.3 – The higher the temperature, the higher the kinetic energy of the
particles. If you lower the temperature, the lower the kinetic energy of the
particles.
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Point 5.4 – A catalyst is a substance which increases the rate of a reaction without
undergoing permanent chemical change in the reaction. The role of catalysts in
chemical reactions is to increase the rate of reaction by helping break chemical
bonds in reactant molecules and provide a 'different pathway' for the reaction. The
catalyst used in the Haber process (production of ammonia) is iron (iron substrate).
Point 5.5 – The role of catalysts is to lower the activation energy of a reaction.
Lowering the activation energy increases the rate of reaction. For endothermic
reactions, more activation energy is needed than in exothermic reactions.
Point 5.7 - Explosions occur when reactions become extremely rapid. This usually
occurs when there is ‘good’ contact between reactant particles and when the
reaction is highly exothermic with high activation energy. It is related to the
collision theory because the molecules of each substance are basically colliding into
each other (as a result of the heated conditions). The substances may or may not
react – it depends if they collide the right way around and if they collide with
enough energy for the bonds to break.