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8.2 The Chemical Earth

1. Mixtures in the Earth

1.2 Elements, mixtures, compounds and the particle theory

• The particle theory states that all matter is made up of tiny particles which are continuously moving
• A mixture is a substance which can be separated into two or more substances by physical or mechanical means such as filtering.
  • The difference between a mixture and a pure substance is as follows:

Mixture	Pure Substance
Can be separated into two or more substances by physical or mechanical means	Cannot be separated into two or more substances by physical or mechanical means
May be homogenous (uniform composition throughout) or heterogeneous (non uniform composition)	Can only be homogenous
Displays the properties of the pure substances it is composed up	Has properties which are constant throughout the whole sample
Properties can change as the relative amount of substance changes	Properties cannot change regardless of how it is prepared or how many times it is purified
Has a variable composition	Has a fixed composition
E.g. Tea, tap water, milk, petrol, coins	E.g. Salt, copper, gold, sugar, water, alcohol

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• An element is a pure substance which cannot be decomposed into simpler substances
• A compound is a pure substance which can be decomposed into simpler substances such as elements
  • It has properties that are quite different from those elements which make is up

1.3 The spheres

• The atmosphere is the layer if gases around the earth
  • An example of a mixture which exists in this ‘sphere’ is air
  • Examples of elements in this ‘sphere’ are nitrogen, oxygen, and argon
  • Examples of compounds in this sphere are H₂O, CO₂, NO₂, SO_2, and CO
• The biosphere is any part of the on earth where life exists
  • An example of a mixture in this sphere is an animal cell
  • An example of an element in this sphere is oxygen
  • Examples of compounds in this sphere are H₂O, CO₂, proteins, lipids, and carbohydrates
• The hydrosphere is the zone containing water and all its states
  • An example of a mixture in this sphere is seawater
  • An example of an element in this sphere is oxygen
  • Examples of compounds in this sphere are H₂O, NaCl, CaCl₂ and MgCl₂
• The lithosphere is rigid outer rock layer of the Earth including all of the crust and the upper mantle
  • Examples of mixtures in this sphere are granite rock and clay
  • Examples of elements in this sphere are iron, gold, silver and copper
  • Examples of compounds in this sphere are Fe₂O₃ and SiO₂

1.4 Separation of naturally occurring mixtures

• Solids of different sizes
  • Separated by sieving
    • Sieving is a separation process which involves allowing only particles of a certain size or smaller to pass through a sieve, in order to separate different substance.
  • Eg. Pebbles and sand
• Solids and liquids (Solids are insoluble)
  • Separated by sedimentation
    • Sedimentation is the process of allowing dense solids to settle at the bottom of the container, which separates it from the liquid
  • Separated by decanting
    • Decanting is the process of carefully pouring off liquid and leaving solid undisturbed at the bottom of the container
  • Separated by filtering
    • Filtration involves passing the mixture through filter paper, where the solid particles are trapped and unable to pass.
    • The filtrate is the solution which passes through the solution.
  • Separated by suspension
    • A suspension is a dispersion of particles through a liquid with the particles being sufficiently large, that they settle out on standing.
  • Eg. Sand and water
• Dissolved solids in liquids
  • The solid in the liquid is known as the solution, whereas the liquid by itself is known as the solvent, and the solid is known as the solute
  • Separated by evaporation
    • Evaporation involves heating the solution in an evaporating basin to drive off the solvent and leave the solute behind
  • Separated by distillation
    • Distillation is the method of separating liquids from solutions. It involves boiling the substance and condensing the resulting vapour back to liquid in a different part of the apparatus.
    • The liquid collected from distillation is known as distillate

• Separated by crystallization
  • Crystallization involves forming solid crystals from a homogenous solution.
• Eg. Sodium chloride and water

• Liquids
  • Separated by fraction distillation
    • Fractional distillation is a type of distillation where a substance is heated to certain boiling points where the most volatile of the liquids will evaporate, before being condensed back to liquid in a different part of the apparatus
    • Volatile means ‘able to be converted to vapour’
      • The more volatile a liquid, the lower its boiling point
  • Separated by a separating funnel
    • A separating funnel is a method which involves the densest of the immiscible liquids to be allowed to run out of the pear shaped device without contaminating any other liquids.
    • Two liquids are said to be immiscible if when mixed, do not form a homogenous liquid but instead, stay as drops of one liquid dispersed over the other liquid.
    • Eg. Ethanol and water
• Gases
  • Separation by liquefaction followed by fractional distillation
    • This involves converting the mixture of gases into liquids, before using fractional distillation

1.5 Properties which enable separation

• Each separation method is only useful due to a physical property of the mixture. They are as follows:
  • Sieving – Particle size
  • Filtering – Solubility and particle size
  • Sedimentation, decanting and suspension – Density
  • Distillation – Boiling point
  • Separating funnel – Density of immiscible liquids
  • Evaporation – Liquid has much lower boiling point than solid
  • Crystallization – Differences in solubility

1.6 Gravimetric analysis

• Gravimetric analysis is used to determine the quantities of substance present in a sample
  • It involve weighing, separating a component from the mixture, filtering, drying, and weighing again
• Gravimetric analysis is used to find:
  • To find if a mineral deposit contains sufficiently high percentage of the required compound to make its extraction from that deposit economically viable
  • To determine percentage composition of the soil and see if it is suitable for growing a particular crop
  • To decide how polluted a particular sample of air or water is
  • To decide whether a particular commercial mixture being sold has the same percentage composition as a similar mixture marketed by a rival company.

1.7 Inorganic and organic

• Any compound containing carbon is organic except
  • CO₂
  • CO
  • CO₃^2-
  • Or any other carbon oxide

1.8 Names of carbon compounds

• Hydrocarbons are organic compounds containing only the atoms of hydrogen and carbon
• There are 3 families, and each family is known as a homologous series
  • A Homologous series is a grouping of atoms which have something common in all members of the series. It is also known as a functional group
• The 3 families are:
  • Alkanes, which have single bonds
    • General formula of C_nH_2n+2
  • Alkenes , which have double bonds
    • General formula of C_nH_2n
  • Alkynes, which have triple bonds
    • General formula of C_nH_2n-2
• Other homologous series include
  • Alkanols, which are in the functional group OH (hydroxyl group)
    • General formula of C_nH_2n+1OH
  • Alkanoic Acid, which are in the functional group COOH (carboxyl group)
    • General formula of C_nH_2nO₂
• Isomers are compounds that have the same molecular formula but different structural formula
  • Eg. 1-butene and 2-butene

1.9 Case Study – Froth Flotation

• This industrial separation process is used to separate mineral ores (mainly zinc, lead and copper sulfides) from the lithosphere
• The process involves
  • Crushing ore into small particles and mixing with water in a large tank
  • Air is blown into the tank and chemicals added to mixture to make it froth and help metal minerals cling to the froth
  • The metal sulfides cling to the bubbles and float to the top of the tank, whilst gangue minerals (unwanted minerals) do not cling since they have different surface tension properties.
  • The metal sulfides are then skimmed from the bubbles, and are heated in a smelter to extract the metals
• The products of the separation and their uses are:
  • Lead – Lead-acid batteries and crystal glass
  • Zinc – Alloy, to galvanize steel
  • Copper – To make electrical wiring and water pipes
• Issues associated with wastes produced
  • Disposing of waste safely due to toxic residue
  • Preventing wastes from entering food chains

2. Elements
2.1 Reactivity of an element and natural existence

• The less reactive an element, the higher the chance of finding it in the Earth as an uncombined element
  • Eg. Inert gases such as helium and metals with low reactivity such as gold are found as uncombined elements
  • Extremely reactive elements are never found as uncombined elements naturally

2.2 Physical properties of metals, non-metals, and semi metals

• Metals
  • Solid at room temperature (except mercury)
  • Shiny or lustrous appearance
  • Good conductors of heat and electricity
  • Malleable (Can be hammered into shapes) and ductile (Can be drawn into wires)
  • High melting and boiling points
  • High densities (except some such as sodium and aluminium)
• Non-metals
  • Poor conductors of heat and electricity
  • Dull appearance
  • Brittle
• Semi-metals
  • Dull
  • Solid at room temperature with high melting and boiling points
  • Poor conductors of electricity and heat (except graphite)

2.3 Uses of metals and non-metals

• Metals
  • Gold
    • Used for jewelry since it is shiny and malleable
  • Aluminium
    • Aircrafts since it is light weight and malleable
  • Iron
    • Used for steel since it is high in strength
• Semi-metals
  • Carbon
    • Used for diamonds since it is shiny and very hard
• Non-metals
  • Helium
    • Used for filling balloons since it has a low density
  • Chlorine
    • Used for disinfectant and bleach since it has a low melting and boiling point

2.4 Physical properties of common elements

Element	Appearance	Malleable	Conductive	Density (g/cm3)	Melting.P (ºC)	Boiling.P (ºC)	Hardness (mohs)	Uses
Al	Shiny, silver	Yes	Yes	2.9	660	2950	2.5	To make aircraft
Pb	Dull, grey, lustrous	Yes	Yes	11.4	327	17400	1.5	Batteries
Fe	Shiny silver	Yes	Yes	7.6	1535	2750	6.5	Railings
Mg	Shiny silver	Yes	Yes	1.7	650	1110	2.5	Flares, fireworks
Zn	Shiny silver	Yes	Yes	7.1	419	907	2	To galvanize iron
Sn	Silver	Yes	Yes	7.3	232	2602	1.5	Jewelry Mirrors
Cu	Brass shiny	Yes	Yes	8.96	1085	2572	3	Electrical wiring Water pipes
I	Black grey	No	No	4.94	114	184	N/A	Sterilization
S	Yellow powder	No	No	2.07	113	445	1.5	Sulfuric acid
C	Dull black-grey	Yes	Yes	2.26	3727	3642	10	Jewelry and pencils

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2.5 Periodic Table

• Metals, semi-metals, and non-metals

• Solids, liquids and gases (at 25°C)

3. Compounds

3.1 Particle nature of matter

• All matter is made up of small particles that are continuously in random motion and interacting with each other
  • Solids have particles which are close together and vibrate in fixed positions. Due to this, solids:
    • Have a definite shape
    • Have a definite volume
    • Cannot be compressed
    • Cannot diffuse
    • Have particles with the least kinetic energy
  • Liquids have particles which are close together but able to move more freely. Due to this, liquids:
    • Do not have a definite shape, but take up the shape of their container
    • Have a definite volume
    • Can be compressed
    • Can diffuse
    • Have particles with some kinetic energy. Due to this, liquids have weaker forces between neighboring particles
  • Gases have particles which are much further apart and moving very freely in rapid motion. Due to this, gases:
    • Does not have a definite shape. It’s shape depends on the container
    • Spread out quickly to fill the whole volume
    • Can be compressed
    • Can diffuse
    • Have particles with great amounts of kinetic energy

3.2 Energy levels of electrons

• Atoms do not collapse when electrons are attracted to the nucleus, since the electrons have posses enough energy to resist moving straight towards the nucleus
• Electrons exist in energy levels, also known as layers of energy shells
  • Electrons fill the shells closer to the nucleus first.
  • The closer the to the nucleus a shell is, the lower its energy level is
  • Each energy level can only hold a certain maximum number of electrons. The amount of electrons held in each shell can be figured from the formula
    • which denotes the maximum amount of electrons held in the n^th energy level
    • An energy level may hold less electrons than its maximum capacity, but no more
  • The electrons fill as follows
    • Fills 1^st 2 (2)
    • Then 8 (2.8)
    • Then next 8 (2.8.8)
    • Now instead of continuing to 9, its fills till 2 are complete in next shell since its more stable that way (2.8.8.2)
    • Then it goes back and starts filling the 3^rd shell till its full (2.8.18.2)
      • Exceptions at Cr and Cu, see back for Aufbau principle
    • Then the 4^th shell begins to get full till it reaches Kr (2.8.18.8)
    • Then like before, fills till 2 in the next shell for stability (2.8.18.8.2)
    • Then begins filling 4^th shell till 2.8.18.10.2
    • Then random exceptions, see Aufbau principle
• The number of electrons in the outermost, highest energy shell is known as the valence shell
  • It determines the interactions between atom

3.3 Mass number and atomic number

• An atom contains an extremely small dense nucleus surrounded by an electron cloud
  • The bulk of the volume is given by the nucleus
  • The nucleus contains protons and neutrons, whereas the electrons are in the electron cloud
  • Atoms or electrically neutral
• The atomic number of an element is the number of protons in the nucleus of an atom (also equal to the number of electrons in an atom)
• The mass number of an element is the number of protons + neutrons in the nucleus of an atom

3.4 Formation of ions

• An ion forms when an atom loses or gains electrons
  • Atoms which lose electrons are known as cations
  • Atoms which gain electrons are known as anions
• Ions which have the same electronic configuration are known as isoelectric
  • Eg. Na⁺, Ca²⁺, Al³⁺, and F^- all have the configuration of 2.8, and are thus isoelectric with the atom Ne

• Valency is the combining power, the tendency to lose, gain, or share one or more electrons
• Ions are positively or negatively charged particles which are usually bonded with other ions by strong forces of electrostatic attraction.

3.5 Periodic table for ions

• Elements in Group 1 have an ion of 1+ charge
• Elements in Group 2 have an ion of 2+ charge
• Elements in Group 3 have an ion of 3+ charge
• Elements in Group 4 have an ion of 4+ charge
• Elements in Group 5 have an ion of 3- charge
• Elements in Group 6 have an ion of 2- charge
• Elements in Group 7 have an ion of 1- charge
• Elements in Group 8 do not form ions]

Lewis electron dot diagram

• Metals
  • Sodium
  • Calcium
• Non-metals
  • Chlorine
  • Oxygen
• Ions
  • Potassium
  • Nitrogen
• Ionic bonding
  • Sodium chloride
  • Iron (II) oxide
• Covalent bonding
  • O₂
  • H₂O
  • CO₃
  • C₂H₄

3.6 Ionic bonding

• Ionic bonding is the type of chemical bonding which involves outright transfer of electrons from one atom to another
  • This bonding consists of electrostatic between the positive and negative ions formed by this transfer of electrons
• In ionic compounds there are no discrete molecules, only an infinite array of positive and negative ions.
  • An empirical formulae are formulae which give the ration by atoms of elements in a compound rather than the actual numbers of atoms in a molecule.

3.7 Molecules

• A molecule is the smallest particle of a substance that is capable of separate existence
  • They are particles which can move independently of each other
  • Molecules are usually formed by covalent bonding of non-metals
• A special type of molecule is a monatomic molecule
  • This is a molecule with only one atom in it
  • These molecules are made of noble gases, which already have a saturated outer shell
  • Similarly, diatomic atoms are made of 2 atoms, which share electrons by covalent bonding

3.8 Covalent molecules

• A covalent molecule is a molecule where its atoms bond covalently
  • Covalent bonds are formed between pairs of atoms by sharing electrons
  • Shared electrons occupy a volume of space that surrounds both atoms.
  • The bonding resulting for the sharing of electrons is known as covalent bonding
    • - represents a single pair of electrons being shared
    • = represents a double pair of electrons being shared
    • represents a triple pair of electrons being shared
  • E.g. two oxygen both have valences of ‘2-‘. Thus they both share two electrons resulting in a diatomic molecule known as O₂ with the structure as O=O
  • The number of electrons needed by an atom to attain a noble gas configuration tells us how many covalent bonds that atom can form

4. Chemical Attraction

4.1 Physical and chemical change

• A physical change is a change in which no new substance is formed
  • Examples are:
    • Changing state
    • Changing physical appearance
    • Dissolving solid into liquid
  • Physical changes are easily reversible
  • Physical changes have relatively small energy changes
  • Physical properties are those which relate to physical changes, such as MP, BP, density, conductivity, and hardness
• A chemical change is a change in which at least one new substance is formed. (also known as chemical reaction)
  • Examples are:
    • Precipitate formed
    • Change in colour
    • Odour is produced
    • Significant change in temperature
    • Disappearance of solid which is not merely dissolution
  • Chemical changes are difficult to reverse
  • Chemical reactions involve a large input or output of energy
  • Mass is conserved in a chemical reaction
  • In a chemical reaction, the starting substances are known as reactants, whereas the resulting substances are known as products
  • Chemical properties are those which relate to chemical changes, such as decomposition, effect of light, and reactivity with other substances

1. Boiling and Electrolysis of Water

• These two processes clearly illustrate the difference between physical and chemical changes
  • Electrolysis of water involves passing a current through water to decompose it into hydrogen and oxygen
  • Boiling water involves heating water causing the particles to vibrate with more intensity due to an increase in kinetic energy until the dipole-dipole attractions break and the liquid changes state and becomes vapour.
• Electrolysis is a chemical change whereas boiling is a physical change since
  • Electrolysis produced two new substances where as boiling does not produce any new substance
  • Electrolysis is difficult to reverse (products need to be mixed together and ignited at high temp) whereas boiling is easily reversed (cooling vapour changes its state to liquid)
  • Electrolysis requires much more energy (20 – 30 kj/g) compared to boiling (2.3kj/g)
  • Electrolysis actually breaks up the water molecules into hydrogen and oxygen, whereas boiling just separates the WHOLE molecules from one another, and doesn’t actually affect the molecule itself.

4.3 Energy during chemical reactions

• Light, heat, and electricity are common forms of energy that may be released during synthesis or absorbed during decomposition
  • Decomposing a compound into elements required a large input of energy because it is necessary to overcome the strong chemical bonds holding the atoms together
  • The stronger the chemical bond, the more energy is required to break the compound into elements. Alternatively, the stronger the chemical bonding in a compound, the higher the energy is released
• Everyday applications of decomposing include
  • Airbags, where sodium azide is decomposed into sodium and nitrogen gas
  • Aluminium is decomposed by electrolyzing molten aluminium oxide
  • Calcium carbonate is decomposed by heating it to make lime, cement, and glass
• Every day applications of synthesis (direct combination reactions) include
  • Rusting of iron and steel to form iron(III) oxide
  • Burning of coke (carbon) which releases much heat energy for smelting
  • Lightning which causes nitrogen and oxygen to form nitric oxide (NO)

4.4 Energy and bonding

• As seen in 4.3, the higher the amount of energy needed to separate atoms in a compound, the greater the strength of attraction or bond between those atoms

4.5 Effects of light on silver salt

• When light hits a silver salt, a redox reaction occurs causing the Ag⁺ to become Ag_(s)m while the Cl^- becomes oxidized to Cl_(aq)
  • This reaction appears as if dark spots are appearing in the white powdery AgCl, as metal silver clumps are being formed.
• This reaction is used in photography, to make negatives.

4.6 Electrolysis of water

• When a current is passed through water, the negative electrode (anode) attracts hydrogen, while the positive electrode (cathode) attracts oxygen, thus separating the two.
• The hydrogen and oxygen are created in a ratio of 1:2
  • This reaction appears as bubbles forming around the electrodes, and water levels dropping in the test tubes. The test tube with the cathode will have half the amount of water left compared to the test tube with the anode
  • The hydrogen and oxygen go to the top of the test tubes since they are less dense than water.
• This reaction is used to split water into hydrogen and oxygen
  • The hydrogen can be used in the production of ammonia
  • The oxygen can be used in medicine to supply patients with oxygen

5. Bonding and Structure

5.1 Physical and Chemical properties

Physical Properties	Chemical Properties
Properties of a substance by itself	Property of a substance reacting with another chemical
No new substance	New substances (at least one)
Easy to reverse	Difficult to reverse
Eg. Malleability, density, electrical conductivity	Eg. Reaction with oxygen, acids, alkalids

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5.2 Physical properties for ionic, covalent molecular or covalent network compounds

Ionic compound	Covalent molecular compound	Covalent network compound
Eg. NaCl, Na2SO4	Eg. CO2, H2O, CH4	Eg. SiO2, SiC, C
Conduct electricity when molten or when aqueous if soluble due to having mobile ions in these states	Never conducts electricity since in never has mobile ions	Never conducts electricity since in never has mobile ions (except graphite)
High MP, BP	Low MP, BP	Extremely high BP, MP
Hard	Hard	Hard
Solid at room temperature	Gas/ liquid at room temperature	Solid at room temperature

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5.3 Metallic, ionic and covalent bonds

• Metallic bonds
  • Occur in metals where they hold the positive metal ions together
  • Metallic bond caused by the random motion of the delocalized outer-shell electrons of the metals atoms, and their attraction to the positive ions
• Ionic bonds
  • Occur in ionic compounds where they hold positively charged metal ions together with negatively charged non-metal ions (or polyatomic ions)
  • Ionic bonds are caused by the electrostatic attraction between oppositely charged ions
• Covalent bonds
  • Occur in molecular compounds and in elements from Groups V, VI, VII and VIII, of the Periodic Table
  • Covalent bonds occur in both small covalent molecules and large covalent networks
  • Covalent bonds are caused by atoms sharing one or more pairs of electrons

5.4 Model of metals and ionic compounds

• Metals are 3D lattices of ions in a sea of electrons
• Ionic compounds are 3D lattices of ions

Metallic bonding model Ionic bonding model


5.5 The empirical and molecular formula

• The empirical formula shows the simplest whole number ration of atoms or ions present in a compound
  • Eg. For ionic compounds, NaCl does not mean that there is one Na ion and one Cl ion in the compound. It is an empirical formula which shows the ratio of Na ions to Cl ions, and not the number of ions in the compound
  • Eg. Similarly for metals, Fe_(s) does not mean there is only one Fe ion in the metallic structure.
• The molecular formula shows the number of atoms of the elements present in a molecule of the compound
  • The molecular formula is used to describe compounds with covalent bonds.

5.6 Common elements as molecules or covalent lattices

• Monatomic molecules
  • He, Ne, Ar, Kr, Xe
• Covalent molecules
  • H₂, O₂, N₂, F₂, Cl₂, Br₂, I₂,
• Covalent lattices
  • SiC, SiO₂, C

5.7 Properties of various structures

• Metallic
  • Held together by metallic bonding caused by delocalized electrons
  • Electrically conductive since they have free moving delocalized electrons
  • Hard since the metallic bonds are difficult to break, yet they are malleable
  • High MP and BP since metallic bonds are difficult to break
  • NOT soluble since water does not break the metallic bonds
• Ionic compounds
  • Held together by electrostatic forces
  • Electrically conductive ONLY when molten or dissolved in water (if soluble) since it is only those states when ionic compounds have delocalized electrons
  • Hard since ionic bonds are difficult to break, yet they are brittle
  • High MP and BP since metallic bonds are difficult to break
  • May be soluble (see solubility table)
• Covalent molecular
  • Held together by intermolecular dispersion forces
  • NEVER electrically conductive since it doesn’t have delocalized electrons
  • Soft since intermolecular forces are weak
  • Low MP and BP since intermolecular forces are weak
  • Soluble
• Covalent network
  • Held together by many strong covalent bonds
  • NEVER electrically conductive since it doesn’t have delocalized electrons (except graphite which has delocalized electrons, and thus can conduct electricity)
  • Hard due to strong covalent bonds, however still brittle
  • V. high MP and BP due to strong covalent bonds which are difficult to break
  • Never soluble

5.8 Limitations of models

• Advantages
  • They summarise what we know
  • They are based on practical experiences
  • Help us to visualize and understand ideas
  • Help us understand the mechanisms of chemical reactions
  • Can be used to make predictions and design further experiments in order to test the model
• Limitations
  • Models are not facts. They are ideas and so they depend on how people interpret observations.
  • Models may be based on incomplete or incorrect information. Models are developed at a certain time in history based on what is known at that time.
    • New experimental discoveries often lead the model being modified
  • Models may be simplifications designed to convey the main idea across, and not focus on the specific details.
  • There are assumptions behind all models
  • Models cannot describe various properties, such as strength of bonds, densities, etc.

8.3 Metals

1. Metals and Alloys

1.1 Metals through history

• Human history is closely linked to the use of metals
• The different ‘Ages’ in history after the Stone Age are named after main metals or alloys being extracted and used by people at that time. These ‘Ages’ were:
  • Stone Age: Up to 3000 BC
  • Copper Age: 3000 BC – 2500BC
  • Bronze Age: 2500BC – 1000 BC
  • Iron Age: 1000BC – 1800 AD
  • Modern Metal Age: 1800AD to present
• Gold was the first metal used by ancient people for ornaments and decorations.
  • It was mainly used because it looked good, was found uncombined in rocks, was malleable, and did not rust
• Copper was seen as valuable because it could be used for tools and weapons
  • Copper compounds in rocks decomposed when the rocks were heated, to extract pure copper. However sometimes it was found uncombined naturally.
• As extraction methods improved, Iron was discovered, followed by Aluminium and Sodium.
• The use of different metals at various stages in history illustrates how technology and chemistry improve due to one-another.

1.2 Alloys

• Alloys are mixtures of metals, which can be mixed in any proportion so they do not have a constant composition or chemical formula
  • Properties of alloys vary with the composition
• Alloys are giant lattice structures of metal ions surrounded by a sea of delocalized electrons
  • Alloys are held by strong metallic bonds
• Alloys are difficult to recycle since the metals are very hard to separate
• The composition of some common alloys are given in the following table

Name of alloy	Composition	Uses	Special Properties
Stainless Steel	80% iron, 18% chromium 2% nickel	Cutlery, sinks, machinery	Strong and resists corrosion
Carbon Steel	99% iron, 1% carbon	Bridges, scissors, car bodies	Extremely hard
Brass	65% copper, 35% zinc	Musical instruments	Hard, resists corrosion, attractive appearance
Solder	63% tin, 37% lead	Joining metal joints	Low melting point, easily worked
Bronze	85% copper, 15% tin	Statues, medals, weapons	Hard, resists corrosion, attractive appearance

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Tungsten steel	70-80% iron, 10-18% chromium, 1-8% nickel	Tools for cutting and grinding	Hard, even when heated
Alnico	62% iron, 21% nickel, 12% Al, 5% cobalt	Permanent magnets	Retains magnetism
Zinc aluminium	45% zinc, 55% Al	Coating to use for roofing and walls	Strong, resists corrosion

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1.3 Energy input for extraction of metals

• Energy is needed to physically break up the ore and separate the mineral from the gangue
  • It is also needed to chemically decompose the mineral to extract the metal
  • The most active metals form the most stable bonds, thus they need the most energy to decompose them
    • The easiest metal compounds to decompose (eg. Copper) can be decomposed by roasting with sand
    • The next easiest (eg. Iron) can be decomposed by roasting with coke (carbon) in a furnace
    • The most difficult (eg. Na and Al) are decomposed by electrolysis of molten compounds or aqueous solutions.
• Only a few metals such as gold, copper and silver occur in nature. Most metals occur as compounds in rocks
  • Gangue is the unwanted waste of an ore after the mineral has been extracted
    • Pollution can results from inappropriate disposal of this waste
    • Rehabilitation is the process of returning mined areas to their original state
• The steps in extracting metal are:
  • Mining the Ore
  • Separating the mineral from the gangue using physical processes such as froth flotation, or chemical processes such as dissolving
  • Extracting the metal by chemical processes such as roasting or electrolysis

1.4 Why are there more metals available than 200 years ago?

• Al, Ti, Mg, and NA are examples of metals that were not available 200 years
  • This is due to the lack of technology for separating such reactive metals
  • Development of technological methods of separation have enable compounds such as Al₂O₃ to be separated, which cannot be decomposed by heat

1.5 Formulae for metal extractions and the future of metals

• Cu₂S + O₂ → 2Cu + SO*₂
• Fe₂O₃ + 3CO → 2Fe + 3CO₂
• The future of metals will probably be replaced with plastics and ceramics

2. Metals and reactions

2.1 Reactions with water, dilute acid, and oxygen

• Reaction with oxygen
  • K, Na, Li, Ba, Ca react rapidly at room temperature
    • Burn to form white powder
  • Mg, Al, Zn, Fe react slowly at room temperature but burn vigorously of heated, since the reactions reach their activation energies much faster
  • Sn, Pb, Cu react extremely slowly and only when heated.
  • The resulting metals oxides have none of the physical properties of the original metal
  • Eg. 4K_(s) + O_{2 (g)} → 2K₂O_(s)
• Reaction with water
  • K, Na, Li, Ba, Ca react with water at room temperature
    • Exothermic reaction producing bubbles
  • Mg, Al, Zn, Fe react with steam at elevated temperatures, in order to reach their activation energies
  • Sn, Pb, Cu, Ag, Pt, Au do not react at all
  • These reactions displace the hydrogen gas to create metal oxides
  • Eg. 2K_(s) + 2H₂O_(l) → 2KOH_(aq) + H_{2 (g)}
• Reaction with dilute acids
  • Acids are substances which can donate H⁺ ions
  • All metals except Cu and under react with dilute acids to produce hydrogen gas.
  • The reactive metals produce vigorous bubbling, with the metal dissolving almost instantly
  • The medium reactive metals have a much slower reaction, but still produce bubbles.
  • Eg. 2K_(s) + 2HCl_(aq) → 2KCl_(aq) + H_{2 (g)}

2.2 The activity series of metals

• By observing the reactions of metals in water, oxygen, and dilute acid, a list can be drawn up depicting the metals in order of decreasing reactivity.
  • This list is known as the activity series
• The activity series is:
  • K, Na, Li, Ba, Ca, Mg, Al, Zn, Fe, Sn, Pb, Cu, Ag, Pt, Au
• This can be easily remembered by
  • Knalibacam Galzn Fesnpb, Cuagaupt
• The lower the metal in the activity series, the more likely it will be found as an uncombined element
• The order of metals discovered is approximately the same order as least reactive metal to the most reactive metal.

2.3 Electron transfer in reactions

• Neutral species equations show the actual chemical substances used in the reaction
• Complete net ionic equations show all the ions involved in the reaction
• Net ionic equations show the actual ionic species that undergo change in the reaction, ignoring spectator ions (ions which do not undergo any change during a reaction

• Eg. Zn_(s) + 2HCl_(aq) → ZnCl_{2 (aq)} + H_{2 (g)} ……… Neutral Species Equation

Zn_(s) + 2H⁺_(aq) + 2Cl^-_(aq) → Zn²⁺_(aq) + 2Cl^-_(aq) + H_{2 (g)*} …….. Complete ionic equation




Zn_(s) + 2H⁺_(aq) → Zn²⁺_(aq) + H_{2 (g)} ……… Net ionic equation

• Oxidation and Reduction
  • When an atom loses electrons, it is oxidized.
    • The oxidized atom is the reducing agent
  • When an atom gains electrons, it is reduced.
    • The reducing atom is the oxidizing agent
  • OILRIG – Oxidation Is Loss Reduction Is Gain
  • Since oxidation and reduction occur at the same time, the reaction is called a redox reaction.
    • Redox reactions are net movement of electrons from one reactant to another. The movement of electrons occurs from the reactant with less electronegativity to the reactant with more electronegativity
  • Half reactions (half equations) are reactions which describe the oxidation and reduction terms separately in terms of electrons lost or gained.
  • Example Reactions
    • Mg burning in O
      • Oxidation: 2Mg → 2Mg²⁺ + 4e^-
      • Reduction: O₂ + 4e^- → 2O^2-
      • Main reaction: 2Mg + O*₂ → 2MgO
    • Zn reacting with HCl
      • Oxidation: Zn → Zn²⁺ + 2e^-
      • Reduction: 2H⁺ + 2e^- → H₂
      • Main Reaction: Zn + 2HCl → ZnCl₂ + H₂
    • Aluminium reacting with HCl
      • Oxidation: Al → Al³⁺ + 3e^-
        • 2Al → 2Al³⁺ + 6e^-

• Reduction: 2H⁺ + 2e^- → H₂
  • 6H⁺ + 6e^- → 3H₂
• Main reaction: 2Al + 6HCl → 2AlCl₃ + 3H₂

• During the reaction between metals and dilute acids, electrons are transferred from the metal to the acid
  • This reaction is an electron transfer reaction
  • When a metal gives up electrons, it becomes a positive ion
  • When acid gains an electron, its hydrogen ions become hydrogen atoms which then form molecules of hydrogen gas

2.4 Uses of metals based on their activity

Name of metal	Use	Reason for choice
Silver	Silver salts are used in photography	Silver is unreactive, thus its compounds are relatively unstable. Thus they decompose in light releasing silver which darkens the film
Lead	Lead is used in batteries and in solder	Lead is a relatively unreactive metal so it has a low corrosion rate
Magnesium	Used in sacrificial anodes to slow the rate of corrosion of the metal in boats.	It’s a reactive metal, thus it reacts in preference to iron or aluminium used in boat building

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2.5 Relative activity and the Periodic Table

• Metal reactivity decreases across a period
  • Eg. An element in Group I is more reactive than a metal in Group III of the same period
  • This is because as you move across a period, the number of electrons in the outer shell increase, causing the atomic radius to become smaller due to the increased attraction between the nucleus and the electrons in the outer shell. Because of this greater attraction force, it is harder for the metal to lose its electrons, thus making it less reactive.
• Metal reactivity increases down a group
  • Eg. An element in period 4 be more reactive an element in a period above it in the same group
  • This is because as you move down a group, the radius the atom becomes larger, thus the attraction force between the metal’s nucleus and its outer electrons is reduced. Because of the lesser reaction force, electrons are lost more easily, thus making the metal more reactive.

2.6 The first ionization energy

• The ionization energy is the energy required to remove an electron from the atom of an element in the gas state.
  • The first ionization energy is the energy required to move one electron from an atom
• Trend of ionization energy
  • As you move across a period, ionization energy generally increases. This is because the number of electrons in the outer shell increases, thus the atomic radii reduces and the attraction force between the electrons and the nucleus is greater. Due to this greater attraction force, a larger amount of energy is needed to remove an electron. However the exceptions in this general trend are caused by electrons entering the p-subshell at Groups III and VI
  • As you move down a group, ionization energy generally decreases. This is because the number of electrons in the outer shell stays the same, yet the atomic radius becomes much larger due to another shell being added on. Thus the attraction force between the nucleus and the outer electrons reduces due the increased distance. Due to this reduced attraction force; a smaller amount of energy would be required to remove an electron.

2.7 Metal Reactions

• Metal + acid → salt + hydrogen
• Metal + water → hydroxide + hydrogen
• Metal + oxygen → metal oxide

3. The Periodic Table

 

[powlmao]

3.1 Models of Atomic structure

About 2000 years ago, Democritus suggested the idea that all matter was made of tiny particles
This was dropped since Aristotle did not agree with the idea.
In 1808, Dalton put forward the atomic theory, stating that:
Elements are made of very small particles called atoms
Atoms cannot be broken into anything simpler
Each element has its own kind of atoms
Atoms cannot be created or destroyed in a chemical reaction
Joseph Thomson showed that atoms were not solve spheres, but made of tinier particles themselves.
He suggested the plum pudding model, where negative particles were scattered throughout the nucleus
Ernest Rutherford showed atoms were largely empty space with most of the mass concentrated in a central positively charged nucleus
He states that each negative particle was an electron which orbited around the nucleus
Niels Bohr altered this, showing the electrons orbit in energy levels
This model of an atom is referred to as the Bohr model of an atom.

3.2 History of the Periodic Table

In 1787, Lavoisier classified the known elements into metals and non-metals, based on their physical and chemical properties
Dalton in 1808 put forward his atomic theory.
From this, he suggested that the atomic mass of hydrogen, should be given as 1 since it was the lightest element.
These masses came to be known as relative atomic masses
In 1829 Dobereiner, noticed some similarities in some sets of three elements, which he called triads.
Eg. He found Li, Na, and K to have similar properties
In 1863, Newlands put forward his law of octaves.
He pointed out that if elements were placed in order of increasing mass, there was a repetition every eighth element.
Newlands work however broke down when some elements did not fit his ideas.
In 1869, Mendeleev proposed the periodic law
This states that if the elements were arranged in order of increasing atomic masses, then elements with similar physical and chemical properties (including valency) occurred at regular intervals.
His table was successful because he placed elements where he though they should go and left gaps for elements he predicted would be discovered later.
He could even predict the properties of these unknown elements, due to their position on his table.
In 1914, Henry Moseley stated the modern period law.
This states that when atoms are arranged in order of increasing atomic number (number of protons) they show a repeating pattern of properties.
Using atomic numbers rather than atomic mass solved some problems, such as K and Ar, where K after Ar but has a lighter atomic mass.
The modern periodic table placed the elements in the vertical Groups and horizontal Periods, and includes the nobles gases discovered between 1893 and 1898

3.3 Periodic Trends

Atomic radius
The radius of an atom decreases across a period due to the increasing nuclear charge causing the valence shell to be attracted closer to the nucleus
The radius of an atom increases down a group due to the extra shell being added onto the atom whilst the number of electrons in the outer shell remain the same.
Boiling and melting point
The BP and MP of an atom across a period relates to bonding:
From Group I-IV, BP and MP increases since bonding becomes stronger
From Group V-VIII, BP and MP decreases since intermolecular forces become weaker
The BP and MP of an atom down a group varies, dependant on the bonding and type of crystal structure
Going down Group I-IV, BP and MP decreases
Going down Group V-VIII, BP and MP increases
Electrical conductivity
Conductivity generally decreases across a period as elements become less metallic (except transition metals where it increases again since they are very metallic)
Conductivity is variable
Conductivity of Group I-III decreases as you go down the group
Conductivity in Group IV-VII increases as you go up, since the elements become more metallic.
Electronegativity
It is the measure of the ability of a atom to attract electrons towards itself
Across a period, electronegativity increases since atomic radii decreases, thus electrons can approach closer to the nucleus to feel a stronger attraction force
Down a group, electronegativity decreases, since atomic radii increases yet the electrons in the valence shell remain the same, causing the electrons to have a weaker attraction force to the nucleus.
Valency
Maximum valency relates to the number of valence electrons an element has
Maximum valency increases across a period, since the number valence electrons increase across a period
Maximum valency is constant down a group, since the number of valence electrons remain constant down a group.




Ionization energy
See 2.6
Reactivity
See 2.5

3.4 Graphing period trends











4. Molar Calculations

4.1 Avogadro’s Number

The mole is a chemical counting unit
1 mole is defined as the number of atoms in exactly 12 grams of the carbon-12 isotope
This number of atoms is: 6.022 x 10²³ (Avogadro’s number)
Number of moles can be calculated by:

Where:
n is number of moles (mol)
m is mass of substance (g)
M is molar mass of substance

4.2 The mole and chemical equations

The coefficients in a chemical equation are the mole ratios of reactant to products
Eg. 2H₂ + O₂ → 2H₂O
Moles of hydrogen : moles of oxygen : moles of water is:
2 : 1 : 2
The volume of 1 more of gas is called the molar volume
One mole of ANY gas will always occupy 24.79L at 25°C and 22.71L at 0°C

4.3 Gay Lussac and his contribution

Gay Lussac released the Law of combining gases which states that:
When two gaseous elements combine, the volumes of the gases that react are in simple whole number ratio to each other and also to the volume of the product of it is a gas. (All volumes measure under same conditions)
It is a summary of experimental results

4.4 Avogadro’s Law

Avogadro summarized his ideas into what is now called Avogadro’s Law which states:
Under the same conditions of temperature and pressure, equal volumes of all gases contain the same number of molecules
ie. 100ml of H₂ has same number of moles as 100ml of O₂
He proposed that the particles in gaseous elements were not single atoms as Dalton believed by were made of more than one atom
He coined these particles “molecules”

4.5 Empirical and molecular formulae

The empirical formula is defined as:
The simplest ratio of atoms or ions in a compound
In general, ionic compounds and giant covalent structures only have empirical formula
The molecular formula is defined as: The number of atoms in a molecule
Covalent compounds such as ethene can have both empirical and molecular formulas
The molecular formula is a multiple of the empirical formula
Some covalent compounds like water are both empirical and molecular formulas

4.6 Volumes of gases

The molar volume of a gas has a mass equal to the molar mass.
When temperature changes, the molar volume changes but the molar mass and number of molecules stay the same.

4.7 More formula






























5. Extraction and Recycling of Metals

5.1 Minerals and Ores

Minerals are any useful naturally occurring elements (eg gold) or compounds (eg bauxite) from the earth
Ores are metal bearing substances (mixtures) from the earth with commercial value

Metal Ore Name Ore composition
Iron Haematite Fe₂O₃
Copper Cuprite Cu₂O
Aluminium Bauxite Al₂O₃
Zinc Sphalerite ZnS

The economic concerns in mining are:
Percentage of metal in the ore
Location of site
Transportation costs
Facilities and employment costs (infrastructure, maintenance, etc)
Processing costs (extraction of metal from ore)
Predicted yield
Predicated yield is considered together with the costs of mining in that location, processing costs, and the money that can be obtained from the extracted metal.
These factors influence the mining company’s decision on whether or not it is economically viable to mine an ore body.

5.2 Prices of metals

Usually, the more abundant a metal, the less the commercial price.
This is not always true however, due to extraction costs.
The harder the extraction process for an element, the more expensive it it.
Eg. Al is the most abundant metal, yet extraction costs are extremely high due to its reactivity. Thus Al is more expensive then say Fe, which is less abundant, but much easier and inexpensive to extract.
The demand of metal also affects their price
Metals in order or price per tonne (decreasing)
Au, Ni, Ti, Hg, Sn, Cu, Mg, Al, Zn, Pb, Fe

5.3 Non-renewable Ores

Ores are non-renewable resources
This is because the rate at which they are used by humans is much greater than the rate at which ores are deposited as sediments.
A resource is any physical or virtual entity which is useful


5.4 Extraction of Copper

There are 4 types of copper ores
Malachite (CuCO₃)
Chalcopyrite (CuFeS₂)
Copper Sulfide (Cu₂S)
Cuprite (Cu₂O)
Stages of copper extraction
First, the ore is crushed into small particles
The crushed ore is then subjected to a physical process such as froth flotation (See 8.2 Chemical Earth - 1.9)
This produces a copper concentrate which is about 25-30% pure copper
Next the copper concentrate is heated with sand (SiO₂). The sand combines with iron oxide to form iron silicate, a liquid slag which is discarded
FeO + SiO₂ → FeSiO₃
The resulting copper sulfide is then heated on its own while air bubbles through it, reducing the copper sulfide to copper metal, and producing sulfur dioxide
Cu*₂S_(l) + O_2*(g) → 2Cu_(l) + SO_2(g)
This remaining copper is blister copper, which is about 98% pure.
This process is known as smelting.
Smelting is the extraction of metals by heating substances to high temperatures to produce a molten material from which metal can be obtained
The blister copper is then subjected to electrolysis to purify it.
The impure copper is used as an anode (positive since it is an electrolytic cell) while the cathode is a thin sheet of pure copper in a sulfate solution
The anode copper gives up electrons to the circuit and goes into the solution
Cu → Cu²⁺ + 2e^-
The cathode copper ions take electrons from the circuit, which deposit as copper metal
Cu²⁺ + 2e^- + Cu
The cathode is then removed containing 99.95% pure copper
Throughout the process, there is a considerable input of energy, mainly in the form of heat.

5.5 Recycling aluminium

The steps involved in recycling aluminium is as follows
The collection of used cans
The delivery of these cans to the recycling plant
The preparation of aluminium from other materials such as labels
The melting reprocessing of aluminium
The making of a new product such as new cans


5.6 Why increase recycling?

Recycling of aluminium should be increased, as there are many advantages which are as follows:
Uses of less fossil fuels
Reduces carbon emissions which are emitted during electrolysis
Less energy is used (saving of about 90% of energy)
Thus there is less pollution from power stations
Bauxite ore is preserved for future generations
Less landfills needed to bury waste from aluminium extraction
Less land needed for new mining sites
Aluminium can be continuously recycled with no loss of quality in formed products
Aluminum currents accounts for about 30% of aluminium production in Australia

5.7 Cost and energy for extraction and recycling

The cost for extraction of aluminum from bauxite is much higher than the cost of recycling, since the processes involved in extracting are much more expensive than those of recycling
The energy for extracting aluminium is 200 Kj/kg (14.7 KW/kg)
The energy for recycling aluminium is 7 Kj/kg (0.7 KW/kg)


8.4 Water

1. Water in Nature

1.1 Define the terms solute, solvent and solution

Solute is the substance (usually a solid) that is dissolved in a solution
Solvent is the substance (usually a liquid) used to dissolve a solute
A solution is a homogeneous mixture in which the dispersed dissolved particles (ions or molecules) are so small that they do not settle out. It is composed of solute dissolved in solvent.

1.2 Identify the importance of water as a solvent

Water is an essential solvent which is required for all living organisms.
Eg. In fluid mediums of both plants and animals
ie. Inter and intra cellular fluid, blood and phloem.
Seawater also has water as a solvent with many salts
Water is a polar solvent
This is because the oxygen atom is more electronegative that then hydrogen atom, and is asymmetrical
It is due to being a polar solvent that water can dissolve a wide range of substances like ionic compounds, polar covalent compounds, and some molecular elements

1.3 Compare the state, percentage and distribution of water in the biosphere, lithosphere, hydrosphere and atmosphere

Sphere State % water Distribution
BIOSPHERE Liquid

60 – 90 % overall
50 – 75% in humans depending on age

In all living things such as fluid mediums
LITHOSPEHE Liquid

1.5 – 20%

Ground water and crystallized minerals
HYDROSPHERE Liquid and solid

95 – 99%

Rivers, seas, oceans and glaciers
ATMOSPEHRE Liquid, solid and gas

0.5 – 5%

Rain, clouds, hail and snow

1.4 Outline the significance of the different states of water on Earth in terms of water

Water is a constituent of cells and has a role at both being a solvent and a raw material in metabolism
Chemical reactions occur in aqueous solutions, essential for metabolism
Liquid water comprises 70% of cells, as this allows ions and molecules to be transported around and react
Essential reactant of photosynthesis
Helps maintain turgidity in plant cells
Major component of blood as a transport medium
Coolant in evaporation of sweat
Water is a habitat in which temperature extremes are less than nearby terrestrial habitats
Liquid water has a higher heat specific capacity (4.18 J/g K), meaning it can absorb and release large amounts of heat energy without a large change in temperature
This is important to living things which are not adapted to large temperature changes
An insulating layer of ice over bodies of water enables aquatic life to survive in winter
Water is an agent of weathering rocks both liquid and solid
Waves weather away cliffs and beaches
River cut paths in valleys
Rain wears away rocks over a long time
Glaciers moving down valleys cause erosion and grind rocks
Water which enters a fissure and freezes over will expand, causing the crack to become bigger
Water is a natural resource for humans and other organisms
Used as drinking water and for other domestic purposes (eg washing clothes)
Used for generating electricity (hydroelectric power)
Used for irrigation in agriculture
Allows transportation of goods
Allows recreational activities such as swimming and ice skating

1.5 Perform an investigation relating to water and its density

Density is
The density of water is highest at 4°C
This is because the water molecules are packed closest together at this temperature, due to hydrogen bonding
This means that water expands as it freezes, which is very unusual
This is why ice floats on water, since it is less dense

1.6 Analyse information by using models for accounting different densities of ice and liquid water

Ice has a crystal structure which is an open framework
Each molecule of water in ice is tetrahedrally surrounded by 4 other water molecules and attracted to them by hydrogen bonds.







Liquid water has particles packed together in a random fashion
In liquid water, some of the hydrogen bonds are broken, allowing the particle to pack closer together








1.7 Plan and perform an investigation to identify and describe the effect of anti freeze or salt on the boiling point of water

See prac book
































2. Structure and Bonding of Water

2.1 Construct Lewis electron structures of water, ammonia, and hydrogen sulfide to identify the distribution of electrons

Water can be described as having strong intra-molecular forces (covalent bonds) and weaker intermolecular forces (hydrogen bonds and dispersion forces)
The Lewis Dot diagrams of H₂O, NH₃ and H₂S are drawn below


Water Ammonia Hydrogen Sulfide

In the Lewis Dot structure of water:
There are two bonding pairs of electrons
There are two non-bonding pairs of electrons (lone pairs)
HYDROGEN SULFIDE has a similar Lewis Dot structure
In the Lew Dot structure of AMMONIA:
There are three bonding pairs of electrons
There is only one non-bonding pair

2.2 Compare the molecular structure of water, ammonia and hydrogen sulfide, the differences in their molecular shapes and in the melting and boiling points.

Molecule Molecular Structure Electron pair geometry Molecular Shape Melting Point Boiling Point
WATER

Tetrahedral Bent Shape 0°C 100°C
AMMONIA

Tetrahedral Pyramidal Shape -78°C -33°C
HYDROGEN SULFIDE

Tetrahedral Bent Shape -83°C -62°C


2.3 Describe hydrogen bonding between molecules

Hydrogen bonds are electrostatic attractions between slightly positive H atoms of one atom and a highly electronegative atom from another molecule.
They result when a hydrogen atom bound to one of the 3 most electronegative elements (N, O, F) is attracted to one of these atoms in another molecule
The strength of a hydrogen bond is typically one-tenth of a normal covalent bond
The term ‘hydrogen bonding’ is incorrect, as these forces are not really bonds, but forces

2.4 Identify the water molecule as a polar molecule

A polar bond is a covalent bond in which the shared electrons are not shared equally between the atoms forming the bond
These occur in heteroatomic molecules (eg HCl), where the electron pairs are shared unevenly and spend more time near one nucleus compared to another, due to that atom being more electronegative.
A dipole is a pair of equal and opposite charges separated in space
A polar molecule is a molecule with an overall unbalanced charge distribution (net dipole). It either has only one polar bond or multiple polar bonds which DO NOT cancel each other out.
HOWEVER, the presence of polar covalent bonds in polyatomic molecules does not guarantee a polar molecule, as dipoles may cancel each other out due to the shape of the molecule
Due to the bent molecular structure of water, there is a net dipole causing the molecule to be POLAR
Due to the molecular structure of hydrocarbons, the dipoles cancels each other out causing them all to be NON-POLAR


Water Methane

Electronegativity imbalances between atoms decide whether a molecule contains any polar bonds
Then the shape of the molecule determines whether the possible dipoles cancel out (non-polar molecule) or produce a net dipole (polar molecule)

2.5 Describe the attractive forces between polar molecules as dipole-dipole forces

Intermolecular forces are the forces between molecules
They are not given the title ‘bond’, as they are much weaker than covalent bonds



Intermolecular forces fall into two categories
Dispersion Forces are the forces between polar and non-polar molecules
They are also known as London Forces or Van der Walls Forces
They arise from the temporary changes in electron distribution in molecules resulting in one end being slightly positive and the other slightly negative
This results in temporary attractive forces between the molecules
Dipole-dipole forces are the forces between polar molecules ONLY.
They arise from the permanent dipoles of molecules which result in electrostatic attraction between molecules
The relative strength of various bonds and forces are seen below

Force Typical Energy
(kJ mol^-1)
O-H covalent bond 463
Hydrogen bond 20
Dipole-dipole force 10
Dispersion force 2


2.6 Explain the following properties of water in terms of its intermolecular forces

Surface tension is a measure of the resistance of a liquid to increase its surface area
It is caused due to the unbalances downward forces from other molecules acting on molecules at the surface
Since the forces acting on water molecules are strong hydrogen bonds, the molecules at the surface experience stronger forces in comparison to other liquid, resulting in higher surface tension of water comparatively
Mercury has higher surface tension than water, ethanol has less

Viscosity is a measure of a liquid’s resistance to flow or to be poured
It is dependent on how strong the intermolecular forces are between the liquid’s molecules
The forces between water molecules are strong hydrogen bonds, which do not readily allow water molecules to move relative to each other
This results in water having a higher viscosity in comparison to other liquids
Glycerol has higher viscosity, ethanol has less
MP and BP of water is comparatively higher than other substances
This is because water molecules are held together by dispersion forces, dipole-dipole forces, AND hydrogen bonds.
Thus it requires more energy to break all these bonds, which is why is has a higher MP and BP compared to other similar sized molecules

3. Solutions and Solubility

3.1 Explain changes to various particles when the interact with water and account for those changes

Type of Substance Example Explanation
Soluble ionic NaCl

Ionic bonds are broken since water molecules reduce the electrostatic forces of attraction between oppositely charged ion
Water molecules then surround each ion, with the negative ion attracting positive dipoles of water and positive ions attracting negative dipoles of water
This hydration of an ionic compounds results in the formation of ion-dipole attractions

Soluble molecular compound C₁₂H*₂₂O₁₁

Intermolecular forces between sucrose molecules are broken since the total strength of these hydrogen bonds are comparatively weaker to the total strength of the increased number of hydrogen bonds with water molecules.

Partially soluble element I₂, O₂

Although these molecules have extremely weak intermolecular forces known as dispersion forces, these substances are only partially soluble as they are non-polar whilst water is a polar solvent.

Soluble molecular compound HCl

HCl molecules are polar and thus they dissolve in the polar solvent water.
The covalent bonds are broken since the total strength of these bonds is comparatively less than the total strength in all the ion-dipole attractions formed with water molecules
HCl is ionized to form hydronium ions (H₃O⁺) and chloride ions (Cl^-) and is named hydrochloric acid

Covalent network SiO₂

Since a covalent network has a large number of strong multi-directional covalent bonds, the total strength of these bonds is much higher than the strength of any total possible number of intermolecular forces formed with water molecules.
Thus these covalent bonds cannot be broken by water, meaning that that the covalent network cannot dissolve in water

Large molecules Polyethylene

The total strength of the covalent bonds in such large molecule sis comparatively greater to the strength of any possible number of intermolecular forces formed with water molecules
Thus these covalent bonds cannot by broken by water, meaning such large molecules cannot dissolve in water


3.2 Analyse the relationship between the solubility of substances in water and the polar nature of the water molecule

Ionic substances often dissolve since the positive ion is attracted to the negative dipole of the water molecule, whilst the negative ion is attracted to the positive dipole of the water molecule, forming ion-dipole attractions.
The number of water molecules that can cluster around the dissolved ions depends on the size of the ion
Polar substances usually dissolve in polar solvents such as water, because there are dipole-dipole attraction forces or hydrogen bonding between the polar substances and the water molecules.
HCl_(g) is a polar molecules which ionizes in water, as covalent bonds break to form ions.

3.3 Perform a first-hand investigation to test the solubilities in water of a range of substances.

See Prac book

3.4 Process information from secondary sources to visualize the dissolution of water of various types of substances


Dissolving of NaCl in H₂O Dissolving of ethanol in H₂O Dissolving of HCl_(g) **in H₂O








4. Concentration of solutions

4.1 Identify some combinations of solutions which will produce precipitates using solubility data

The Solubility Table

Ion (s) Solubility
Group I, NH₄⁺, NO₃^-, CH₃COO^- Always soluble
Br ^-, Cl^-, I^- Always soluble EXCEPT Ag⁺, Pb²⁺
SO₄^- Always soluble EXCEPT Ca²⁺, Ba²⁺, Pb ²⁺, Ag⁺
CO₃^2-, PO₄^3-, SO₃^2- Always insoluble EXCEPT G1, NH₄⁺
OH^- Always insoluble EXCEPT G1, NH₄⁺
S^2- Always insoluble EXCEPT G1, G2, NH₄⁺

4.2 Describe a model that traces the movement of ions when solution and precipitation occur

Dissolution is the process by which a solid or liquid forms a homogenous mixture with a solvent.
Dissociation is the general process in which ionic compounds split into smaller ions or radicals, usually in a reversible manner (Hydration in water)
Ionization is the process in which an atom or molecule is converted into an ion
Precipitation is the opposite of dissolution, where a substance “falls out” of a solution as an insoluble sold.

4.3 Identify the dynamic nature of ion movement in a saturated solution

A saturated solution is a solution in which more solute will dissolve at a particular temperature.
The limit of dissolving of the solute is reached much earlier with an INSOLUBLE compound than with a SOLUBLE compound
In these solutions, the dissolution and precipitation reactions occur simultaneously rate at the same time.
Neither reaction goes to completion
Equilibrium is a dynamic situation in a closed system where there is continual interchange between reactants and products in a reaction at the atomic level but with no observable change
In an equilibrium situation
The container must be sealed so no solvent can escape
The temperature must be kept constant
The concentration of ions and the mass of undissolved solid will not change
An equilibrium equation is written with double arrows
Eg. PbI_2(s) Û Pb²⁺_(aq) + 2I^-_(aq)






4.4 Describe the molarity of a solution as the number of moles of solute per liter of solution.


Where:
C is Concentration (mol L^-1)
N is the Number of Moles (mol)
V is the Volume (L)

From this we derive:



4.5 Explain why different measurements of concentration are important

It is useful to have different measure of concentration because:
People other than chemists also use concentration and they may not be familiar with molarity
In commerce and industry, the measure of concentration focuses on the mass or volume of solute, rather than the number of moles
In areas such as environmental studies or drug testing, the concentrations of solutes are so small that molarity becomes inconvenient and so parts per million (ppm) are preferred
These different methods of expressing concentration are seen below:

Method Units Definition Where used
Percentage by volume % (v/v) Volume of solute expressed as a % of total volume of solution General labeling in medicines and foods for liquid solutes
Percentage by mass % (w/w) Mass of solute expressed as a % of total mass of the solution General commercial labeling for solid solutes
Grams in a volume g L^-1 Grams of solute per litre of solution General commercial labeling for solid solutes
Parts per million by volume ppm (v/v) Volume in millilitres per kilolitre of solution Environmental studies such as air pollution
Parts per million by mass ppm (w/w) Mass in milligrams per kilogram of solution Environmental studies such as water pollution

mg/L = ppm







5. Water and Energy

5.1 Explain what is meant by the specific heat capacity of a substance

Specific heat capacity (C) of a substance is the measure of the number of joules of heat energy required to raise the temperature of 1g of the substance by 1°C or 1K.
It is measured in J/gK

5.2 Compare the specific heat capacity of water with a range of other solvents

Substance Specific Heat Capacity
Water 4.18
Ethanol 2.44
Glycerol 2.38
Antifreeze 2.39

5.3 Explain the equation regarding change of heat


Where
ΔH = Change in enthalpy ie. the amount of heat energy released or absorbed (J)
m = The mass of water in grams (g)
C = The specific heat capacity of the substance (J/gK or J/g°C)
ΔT = Change in temperature (K or °C)
ΔH is positive when:
Heat is absorbed from endothermic reactions
The temp. of surroundings gets cooler
ΔH is negative when:
Heat is absorbed from exothermic reactions
The temp. of surroundings gets hotter

5.4 Explain how water’s ability to absorb heat is used to measure energy changes in chemical reactions

Many reactions take place in water since water has a higher heat capacity, allowing it to absorb heat better than most other common liquids
A Calorimeter is a thermally insulated container containing a known mass of water.
Heat released or absorbed in a chemical reaction is then released or absorbed by the water, resulting in a temperature change of the water (measured by a thermometer or probe)
Using the ‘change of heat’ equation, the amount of heat released or absorbed can be calculated.





5.5 Describe dissolutions which released heat as exothermic reactions and give examples

An exothermic reaction is when heat is released
The ΔH change is negative
As a result, the water gets warmer since it absorbs the released energy
In exothermic reactions
The energy required to break the ionic bonds in the crystal lattice of the solute and break the intermolecular hydrogen bonds between water molecules is LESS than the energy released when the separated ions are hydrated (forming ion-dipole forces)
Energy required for bond breaking is LESS than the energy released from bond formation
Examples of exothermic reactions are:
Dissolution of NaOH, CaCl₂, and concentrated H₂SO₄

5.6 Describe dissolutions which absorb heat as endothermic reactions and give examples

An endothermic reaction is when heat is released
The ΔH change is positive
As a result, the water gets cooler since it releases the absorbed energy
In endothermic reactions
The energy required to break the ionic bonds in the crystal lattice of the solute and break the intermolecular hydrogen bonds between water molecules is MORE than the energy released when the separated ions are hydrated (forming ion-dipole forces)
Energy required for bond breaking is MORE than the energy released from bond formation
Examples of endothermic reactions are:
Dissolution of NH₄NO₃, NaCl, and KNO₃

5.7 Explain why water’s ability to absorb heat is important to aquatic organisms and to life on earth generally

Water’s high heat specific heat capacity minimizes the fluctuations in the surrounding temperature.
This is because it can absorb a large amount of heat without having a large rise in temperature.
This is very important to many aquatic organisms, as most are ectothermic and thus depend on the surroundings to determine body temperature.
Since their enzymes and other functions of the body only operate within a narrow temperature range, aquatic ectotherms must rely on the water to maintain a constant body temperature and therefore sustain life.






5.8 Explain what is meant by thermal pollution and discuss the implications for life if a body of water is affected by thermal pollution

Thermal pollution is the discharge of quantities of hot water from power stations (previously used as a coolant) into rivers or lakes that are large enough to increase significantly the temperature of the water body
A 2 – 5°C increase can be significant
The detrimental effects of increased water temperature are:
Less dissolved oxygen, causing stress to organisms and even suffocation
Increased metabolism rates which increases demand for oxygen and thus aggravates the low dissolved oxygen problem
Fish eggs cannot develop or hatch if water temp. is too high
False temperature cues are given to aquatic life, thus setting off migrations and spawning at wrong times of the year.
Lethal temperature limits may be exceeded causing the death of aquatic life and affecting food chains
Sudden temperature changes can kill fish eggs even when temp. changes are within the survival range of the eggs

5.9 Choose resources and perform a first-hand investigation to measure the change in temperature when substances dissolve in water and calculate the molar heat of solution

Molar heat of solution is the heat released or absorbed when 1 mole of a solid is dissolved in a large quantity of water
See prac book

5.10 Process and present information from secondary sources to assess the limitations of calorimetric experiments and design modifications to equipment used

Limitations are:
The insulation may not have been adequate, and thus heat may have entered of escaped from the calorimeter
Some heat from the dissolving substance may have been lost to the container and to the thermometer itself
Solute might not be completely dissolved
Reading on thermometer might fall between measurement markings on thermometer.
Modifications
Uses a thermos flask to increased insulation
Use a thermometer with a smaller range of temperatures to improve accuracy of the measurement or use a data logger and a temperature probe
Use a burette rather than a measuring cylinder to measure out a more accurate volume of water
Repeat the experiment with several calorimeters and take an average
Add solute quickly and stir well


8.5 – Energy

4. Combustion

4.1 Describe the indicators of a chemical reaction

A gas is produced
A change in colour
A solid is produced by the reaction [ie. precipitate]
There is a change in temperature indicating changes in energy
Light is released

4.2 Identify combustion as an exothermic chemical reaction

Combustion is an exothermic chemical reaction which involves the reaction of a substance with oxygen
One or more compounds are produced and energy is released
Examples of complete combustion reactions are:



4.3 Outline the changes in molecules during chemical reactions in terms of bond-breaking and bond-making

During a chemical reaction, there are 3 stages of change the molecules undergo.
Ionic or covalent bonds are broken within the main reactant
Covalent bonds of oxygen are broken
Ionic or covalent bonds are made between the separated ions or atoms, causing the formation of new molecules

4.4 Explain that energy is required to break bonds and energy is released when bonds are formed

In any chemical reaction there are always energy changes.
Energy is used to break bonds so that atoms or ions can become available to react and be rearranged.
When new bonds form, energy is released
In an exothermic reaction; energy released > energy absorbed
\Exothermic reactions release energy in to the environment
In an exothermic reaction; energy released < energy absorbed
\Endothermic reactions absorb energy from the environment

4.5 Describe the energy needed to begin a chemical reaction as activation energy

Activation energy is the energy needed to break existing bonds so that atoms and ions are available to react
It is generally expressed in kJ/mol

4.6 Describe the energy profile diagram for both endothermic and exothermic reactions

Exothermic







Where:

E_R = energy of reactants
E_P = energy if products
E_A = activation energy
ΔH = enthalpy change Note

In exothermic reactions, the energy in the products is lower than the energy in the reactants.

Thus the ΔH is negative
Endothermic






Where:

E_R = energy of reactants
E_P = energy if products
E_A = activation energy
ΔH = enthalpy change Note

In endothermic reactions, the energy in the products is higher than the energy in the reactants.

Thus ΔH is positive

4.7 Explain the relationship between ignition temperature and activation energy

Ignition temperature is the temperature to which a substance must be heated in order for it to catch alight, without apply a spark or flame
This indicates the amount of energy needed to break the bonds of the reactants so that they are ready to react.
This is the activation energy
Thus the ignition temp. is the temperature at which the particles reach the activation energy for the reaction to occur.

4.8 Identify the sources of pollution which accompany combustion of organic compounds and explain how these can be avoided

Organic compounds are carbon based compounds such as fossil fuels, alkanes, carbohydrates, proteins, etc
Most of the pollutants are formed as a result of burning fossil fuels, and are usually in the form of oxides
Carbon dioxide
This is colourless and odourless gas formed when plenty of O₂ is available for combustion
Major cause of the enhanced greenhouse effect and contributes to acid rain
Can be used to make fire extinguishers and fizzy drink

Carbon monoxide
This is a colourless and odourless gas produced by incomplete combustion when the oxygen supply is restricted
Toxic as it combined with the haemoglobin in red blood cells in preference to oxygen, reducing the ability of blood to transport O₂
A catalyst can be used to convert CO to CO₂
Sulfur dioxide
Formed by the combustion of sulfur in fossil fuels
Causes respiratory problems such as asthma, as well as acid rain
Can be used to make sulfuric acid
Nitrogen oxides
Formed by combustion of nitrogen in fuels at high temperature
Causes respiratory problems, as well as contributes to the formation of acid rain. Also causes photochemical smog
Catalysts can be used to remove nitric oxides by increased that rate of its reaction with CO
Particles are small particles of solids (such as ash) and tiny droplets of liquids, that stay suspended in air
Formed when mineral matter is combusted
Reduce visibility, causes respiratory problems, damage machinery and contribute to photochemical smog.
Can be filtered out in power stations using electrostatic precipitators
Additives in petrol (Case: Lead)
Lead , Pb(C₂H₅)₄ was added to petrol to help produce even combustion
Emitted through car exhaust and entered food chains, accumulating in animal tissue as a heavy metal
Has been phased out completely with unleaded petrol
POLLUTION can be avoided by:
Ensuring complete combustion of fuels, by ensuring there is enough oxygen
Remove particle pollutants using filters and electrostatic precipitators
Find uses for wastes eg. SO₂ for H₂SO₄
Lower combustion temperature to avoid production of nitrogen oxides

4.9 Solve problems and perform a first-hand investigation to measure the change in mass when a mixture such as wood is burned in an open container

See Prac Book

4.10 Identify the changes of state involved in combustion of a burning candle

Candle wax is a solid fuel, comprised of large chains of hydrocarbons (C₃₀H₆₂)
First the wax melts, to change into a liquid state
Then the molten wax moves up the wick and vaporises, turning into a gas state
It is the wick vapour which actually burns when a candle is lit.

4.11 Perform a first-hand investigation to observe and describe examples of endothermic and exothermic chemical reactions: SEE PRAC BOOK
5. Reaction Kinetics

5.1 Describe combustion in terms of slow, spontaneous, and explosive reactions and explain the conditions under which these occur

Combustion can occur at varying rates
These are:
Slow combustion
This is combustion at a very slow rate, such as burning lumps of wood or rusting oxygen
The fuel takes a long time to burn in this combustion
This usually occurs when big lumps of fuel are used and air supply is limited
Burning only occurs on the surface of the fuel and its speed is controlled by the limited supply of air
Fast combustion
This is combustion which occurs at a much faster rate, such as burning methane or kerosene
This fuel burns very quickly in this combustion
This usually occurs when gaseous or crushed fuels having large surface areas are exposed to excess oxygen
Explosive combustion
This is combustion that occurs extremely rapidly, such as burning hydrogen.
The fuel burns almost instantaneously
These are usually initiated with a spark, with the fuel being exposed to an excess amount of HEATED air.
All forms of combustion are spontaneous reactions
This means that once started, they proceed without further assistance and continue to go on until all the fuel is used up

5.2 Explain the importance of collisions between reacting particles as a criterion for determining reaction rates

Rate of reaction is the rate at which reactants disappear in a reaction
In a Concentration vs time graph, the gradient is the rate of reaction
Typically, reaction rate decreases as the reaction proceeds
In a reaction, reacting particles must collide SUCCESSFULLY in order for reaction to begin.
The rate of reaction can be increased by increasing the number of successful collisions which occur.
It can be increased in following ways:
High concentrations
The more reactant particles present per unit volume, the greater the chance of successful collisions occurring between reactant particles as the particles are packed closer together.

Temperature increase
An increase in temperature gives particles more kinetic energy, allowing them to move faster.
This results in the average energy of the particles to increase, thus being a higher change for particles to overcome the activation energy and collide successfully.
After successfully colliding, the kinetic energy may be converted to heat, light and chemical cnergy
Gas particles
By changing the state of the reactants to gas, the resulting gas particles move more freely and rapidly
This results in a greater chance of successful collisions occurring compared to solids and liquids
Reactions with gases are known as homogenous reactions
Heterogeneous reactions are those where the reactants differ in state
Extra factors influencing heterogeneous reactions are:
Surface Area
The greater the surface area of a solid reactant, the more its particles are exposed to other colliding particles, thus increasing the chance of successful collisions
Stirring
By stirring the mixture, solid reactant particles are constantly suspended and thus expose maximum surface area to other colliding particles, rather than settling on the bottom and being covered
Thus by stirring, the rate of reaction is increased.

5.3 Describe the role of catalysts in chemical reactions, using a named industrial catalyst as an example

A catalyst is a chemical which speeds up the rate of a chemical reaction without being used up by the reaction.
It can only change the speed at which the products form, not the nature or concentration of the products.
Eg. Vanadium pentoxide is a catalyst involved in the manufacturing of sulfuric acid

5.4 Explain the role of catalysts in changing the activation energy and hence the rate of chemical reaction

A catalyst can increase the rate of a chemical reaction by providing an alternative pathway for the reaction with a lower activation energy
The lower activation energy means that less energy will be needed for the particles to start reacting
Catalysts are particularly useful when the uncatalysed reaction has a very high activation energy







5.5 Investigate the condition under which explosions occur and relate these to the importance of collisions between reacting particles

Explosions occur when reactions become rapid
This usually happens when there is a good contact between reactant particles and when the reaction is highly exothermic with a high activation energy
Once the reaction is initiated, it liberates energy which heats up the reaction mixture
This makes the reaction go faster, in a self propelling cycle leading to an extreme escalation in reaction rate, causing an explosion
In order for the rate to increase this way, there must be a good supply of oxygen available to the fuel.
Otherwise lack of oxygen slows down the reaction

5.6 Relate the conditions under which explosions occur to the need for safety in environments where fine particles mix with air

Working conditions should ensure there can be no build-up of concentrations of flammable particles
Formation of flammable dust should be minimized, and dust which is there should be removed
Not only is dust from obvious fuels such a coal dangerous, but also dust from other fuels such as wheat or fabrics can also be dangerous.

5.7 Develop a model to simulate the role of catalysts in chancing the rate of chemical reactions

 

[Gujju]

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[powlmao]

To annoy people that they have to copy it in word if they want the notes

 

[Gujju]

To annoy people that they have to copy it in word if they want the notes

LOL you funny fella btw do u have notes for english (adv)???

 

[Gujju]

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[powlmao]

i hope u do know abt ctrl + a ??????

obviously,

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[ditzy bub]

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