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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

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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

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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.