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Chemistry Year 12

By nicdontdoit Mar 05, 2013 4004 Words
Core Module 1: The Chemical Earth
Contextual Outline
The Earth includes a clearly identifiable biosphere, lithosphere, hydrosphere and atmosphere. All of these are mixtures of thousands of substances and the use of this pool of resources requires the separation of useful substances. The processes of separation will be determined by the physical and chemical properties of the substances.

In order to use the Earth’s resources effectively and efficiently, it is necessary to understand the properties of the elements and compounds found in mixtures that make up earth materials. Applying appropriate models, theories and laws of chemistry to the range of earth materials allows a useful classification of the materials and a better understanding of the properties of substances.

This module increases students’ understanding of the nature, practice, applications and uses of chemistry.

8.2.1 The living and non-living components of the Earth contain mixtures

.2.1 a) Construct word and balanced formulae equations of chemical reactions as they are encountered Method:
1. Identify reactants, products and reaction conditions
2. Write reactants and products with formulae (note valency numbers and balance electric charges) 3. Balance equation using laws of conservation of mass

.2.1 b) Identify the difference between elements, compounds and mixtures in terms of particle theory Atom: The smallest particle that cannot be divided by any physical or chemical means Molecule: Two or more atoms (the same or different) that are chemically bonded together Lattice: 3D array of oppositely charged particles (ions) held together by an electrostatic attraction Element: Consists of only one type of atom

Compound: Composed for two or more elements that are chemically bonded together. Contains a fixed number of atoms of each component element. Pure Substance: Have a fixed composition and fixed properties (for that specific substance). Cannot be decomposed by physical separation methods. Mixtures: Have variable composition and variable properties. Can be separated into components by physical separation methods.

* Elements are composed of atoms or molecules
* Compounds are composed of a fixed number of atoms of different elements * Mixtures have various particle types and compositions
* Homogenous mixtures Particles are distributed uniformly (E.g. glass, salt water – appears the same by naked eye) * Heterogeneous mixtures Not distributed uniformly

.2.2 Identify that the biosphere, lithosphere, hydrosphere and atmosphere contain examples of mixtures of elements and compounds Lithosphere
Rocks and weathering products (soil). Mixtures of minerals. * Minerals naturally occurring crystalline solid that has a fixed chemical composition (element or compounds) * Combined in different proportions to form sedimentary, igneous and metamorphic rocks * Examples:

* Compounds: Sand – silicon dioxide, Extraction of metals from minerals (aluminium from compounds) * Elements: Gold
* Mixtures: Mud
* Most abundant: oxygen and silicon at crust, iron because of core. Hydrosphere
Water with varying quantities of particles of chlorine, sodium, magnesium etc. * Most abundant: water (Hydrogen and oxygen)
* Examples:
- Mixtures: Salt water
Mixture of uncombined light elemnts
* Most abundant: Nitrogen (78%) and oxygen (21%)
* Examples
* Compounds: Carbon dioxide
* Elements: Oxygen molecules
Living things: composed for many common chemical characteristics and Carbon as a basis * Most abundant: Oxygen and hydrogen (because of water in cells of living organisms) * Examples:
* Compounds: Carbon compounded basis for life (CHO, fats, proteins, nucleic acids)

.2.3 Identify and describe procedures that can be used to separate naturally occurring mixtures of: * Solids of different sizes
A) Sieving:
* Where: Metal mining industries
* Process: Mined ore is crushed/ground to separate useful minerals from gangue. Sieved through wire mesh * Property utilized: Particle size
B) Sedimentation:
* Where: Heavy minerals in alluvial deposits (E.g. Gold)
* Process:
1. Panned in water with lighter material washed out
2. Finely ground ore added to water, gravity separators mean heavy particles sediment faster, finely ground material/heavy material are separated * Property utilized: density
C) Froth Flotation:
* Where: Silver, lead and zinc mines
* Process: Ground materials mixed with oily chemicals and water, air blown through mixture to make froth (minerals grains adhere), froth layer (concentrate – higher amount of metallic minerals) scraped off, gangue sediments to bottom * Property utilized: Solubility in oil

D) Magnetic Separation:
* Where: Iron mineral grains
* Process: Mixture passes along conveyer belt with magnetic roller, non magnetic falls off first thus separating * Property utilized: Magnetic property

* Insoluble Solids and liquids

A) Filtration:
* Where: Our water filtration system
* Process: Adding mixture to water, stirred, poured through filter paper, insoluble solid remains in filter paper * Property utilized: Particle Size
B) Centrifugation:
* Where: Cream separated from milk, blood cells separated from plasma * Process: Spun at high speed so sediment collects at the base of tubes (layers according to size and weight) * Property utilized: Particle size, density

* Soluble solids in liquids
A) Evaporation:
* Where: Salt from salt water
* Process: Heating and in turn vaporising one component leaving the other * Property utilized: Boiling point
B) Crystallisation:
* Where: Purify impure salts (E.g. Sugar crystals extracted from sugarcane syrup) * Process: Impure salt dissolved in water at high temperature to make a concentrated solution, cooled so salt crystallises, leaving impurity in solution * Property utilized: Solubility

* Liquids
A) Separating Funnel:
* Where: Immiscible liquids - Separate oil and water
* Process: Pouring mixture into separating funnel, opening tap to let lower layer through Property utilized: * Property utilized: Density
B) Distillation:
* Where: Miscible liquids - salt water to get fresh water, alcohol from water, crude oil * Process: Liquid mixture boiled in distillation apparatus, lowest boiling point vaporises first, then condenses. Simple (very different boiling points) or Fractional (for similar boiling points) * Property utilized: Boiling point/volatility

* Gases
A) Zeolite Sieves:
* Where: Oxygen for hospitals
* Process: Molecular sieve called zeolite (with surfaces and channels) selectively absorb certain gases, so oxygen remaining is almost entirely pure. * Property utilized: ???
B) Cryogenic Air Separation:
* Where: Oxygen, nitrogen (liquid nitrogen for snap freezing) and argon (neon signs) * Process: Gradually decreases air temperature to boiling point of gases (in order) to liquefy and be drained, which is then warmed and reconverted to isolate each. * Property utilized: Boiling point

.2.4 Assess separation techniques for their suitability in separating examples of earth materials, identifying the differences in properties that enable these separations Earth materials are mixtures of elements and compounds, each with its own physical properties. The difference in the physical properties allows for separation of mixtures. Therefore, different techniques are suitable for different earth mixtures as they separate different elements and compounds with different physical properties.

.2.5 Describe situations in which gravimetric analysis supplies useful data for chemists and other scientists Gravimetric analysis Separating components of a material to accurately determine their mass so percentage composition can be calculated

* Can determine
* Composition of a mixture using physical separation techniques * Percentage composition of a compound using chemical and physical separation techniques

* Examples of when used:
* Percentage by weight of ingredients (sugar, fat, fibre) in food * Purity and composition of alloys for building construction * Extent of heavy metal pollution in river water
* In a newly discovered mineral deposit, engineers need to know the concentration of the required mineral to decide if the extraction of the deposit is commercially valuable.

.2.6 Apply systematic naming of inorganic compounds as they are introduced in the laboratory Binary Compounds:
Ionic compounds:
1. Element closer to the bottom left corner of periodic table is named first 2. Name of first element stays the same
3. Name of second element changes to anion name – suffix ‘ide’ 4. If a metal has several different valencies, a Roman numeral is used to indicate the valency in the name of the compound

Covalent compounds:
1. Element closer to the bottom left corner of periodic table is named first 2. Name of first element stays the same
3. Name of second element changes to anion name – suffix ‘ide’ 4. Prefixes are used to distinguish combinations - carbon dioxide, carbon monoxide 1| 2| 3| 4| 5| 6| 7| 8| 9| 10|
Mono| Di| Tri| Tetra| Pent| Hex| Hept| Oct| Non| Dec|

Compounds containing polyatomic ions:
* Ionic compound of 3 elements, where third element is Oxygen, has the suffix ‘ate’. * The one with less oxygen has the suffix ‘ite’
* The corresponding acids with hydrogen have suffixes ‘ic’ and ‘ous’ Example:
Na2So3 Sodium Sulfite H2So3 Sulfurous acid
Na2So4 Sodium SulfateH2So4 Sulfuric acid

.2.7 Identify IUPAC names for carbon compounds as they are encountered

8.2.2 Although most elements are found in combinations on Earth, some elements are found uncombined

.2.1 Explain the relationship between the reactivity of an element and the likelihood of its existing as an uncombined element Reactivity Chemical property that is related to the electronic structure/valency of the element. That is, the more easily an atom can lose or gain electrons in its valence shell to achieve a full outer shell makes it more reactive. * Unreactive elements can exist as free elements in nature * Reactive elements combine with other substances in the environment to form compounds Native metals those metals found as free elements in nature (gold) Noble gases those gases found as free elements in nature (helium, neon)

To Exist Uncombined:
* Highly Unreactive (native metal/noble gas)
* Other elements it is around are not reactive (oxygen in the air is a free element because argon and nitrogen aren’t highly reactive) – also continuous supply of free oxygen from photosynthesis * Environment protects from reaction (sulfur is insoluble and not reactive with oxygen at low temperature so can exist in underground ore bodies)

.2.2 Classify elements as metals, non-metals and semi-metals according to their physical properties Property| Metals| Semi-Metals| Non-Metals|
Appearance| Lustrous | Low sheen| Dull |
State (at room temp)| Solid (except mercury)| Solid| Gas (except bromine (l), carbon(s), sulfur(s))| Thermal Conductivity| High | High| Low (insulator)|
Electrical Conductivity| High| Semi conductor| Low (insulator - except graphite)| Density| High | Moderate| Low|
Ductility| High| Moderate| Nil (brittle)|
Malleability| High| Moderate| Nil (brittle)|
Boiling Point| High| Very high| Low|
Melting Point | High| Very high| Low|
Strength| High| Variable| Low|

Melting point lowest temperature at which a solid changes to a liquid Boiling point lowest temperature at which a liquid changes to a gas Density mass of the substance per unit of volume
Electrical Conductivity quantity of electric current transmitted though a unit cube of material when there is a potential difference of 1 volt across the cube Thermal conductivity the rate at which heat energy is transferred through a unit cube when there is a 1 degree temperature difference across the cube Malleable Ability to be shaped

Ductile Ability to be drawn into wires
Density The amount of atoms in a unit size of the substance

.2.3 Account for the uses of metals and non-metals in terms of their physical properties Uses of metals are directly related to their chemical and physical properties.

8.2.3 Elements in Earth materials are present mostly as compounds because of interactions at the atomic level

.2.1 Identify that matter is made of particles that are continuously moving and interacting Kinetic Theory of Matter:
* Everything is made of particles; particles are moving at all times. * Change in state is a result of change in temperature (average kinetic energy) “



* Solids: Particles vibrate at a fixed position
* Liquids: Particles slide over each other and vibrate
* Gases: Particles move around very freely

* Solid/liquid refers to the particles themselves
* Gas refers to the collision of the particles within the walls of the container * (<0.5% is taken by particles themselves)
* Volume of gas directly proportional to temperature x number of particles / pressure In process of change state
In the state that it is coming from
State (S, L, G)
In process of change state
In the state that it is coming from
State (S, L, G)

.2.2 Describe qualitatively the energy levels of electrons in atoms Electron Configuration Rules:
1. Orbiting shells fill low energy inner shells before high energy outer shells * Energy is inversely proportional to attraction
2. Any atom can have up to 7 shells
* For the nth shell it can fit up to 2n2 electrons
3. The outer most shell (valence shell) can only fit up to 8 electrons * Transition metals:
* Most have 2 in valence shell
* The new electrons fit into 2nd and 3rd outer most shell

* Noble Gas Electron Configuration Theory:
* All noble gases (Group 8 elements) don’t react
* An atom with a full outer shell wont react
* Atoms react in order to achieve a full outer shell

.2.3 Describe atoms in terms of mass number and atomic number

Subatomic Particle| Location| Mass| Charge|
Proton| Nucleus| 1 amu| +1 |
Neutron| Nucleus| 1 amu| 0|
Electron| Orbiting shells| 1/1800 amu| -1|
Isotopes Different number of neutrons

.2.4 Describe the formation of ions in terms of atoms gaining or losing electrons Ions Charged atoms or charged molecules, both positively and negatively

* Lose their valence electrons to achieve noble gas configuration * Become positively charged – cation
* Gain valence electrons to achieve noble gas configuration * Become negatively charged – anion

.2.5 Apply the Periodic Table to predict the ions formed by atoms of metals and non-metals Period 7 horizontal rows - all elements in the same period have the same number of electron shells Group 8 vertical columns - all elements in the same group have the same number of valence electrons

* Valence number = group number
* Transition metals: valence number = 2
* Valence number = group number minus 8

.2.6 Apply Lewis electron dot structures to:
* The formation of ions
* The electron sharing in some simple molecules
1. Only draw valence electrons
2. Electrons are expressed as: x or • or °
3. For covalent bonds, circles around shared electrons are used. • x • x
2 circles = 2 pairs shared
4. For ionic bonds brackets are used [ +] [ -]

.2.7 Describe the formation of ionic compounds in terms of the attraction of ions of opposite charge * A metal atom loses its electrons, to form a cation
* A non-metal atom gains these electrons, to form an anion * These two charged atoms attract each other, an electrostatic attraction, and therefore form an ionic bond. * The compound formed is an ionic lattice.

* Examples: Sodium Chloride, Magnesium Oxide

* The ratio of atoms in an ionic compound depends on the valency of each element * Valency the number of electrons gained or lost by each element in the reaction * The number of electrons lost must equal the number of electrons gained * The formula for an ionic compound gives the simplest ration Empirical formula

.2.8 Describe molecules as particles that can move independently of each other Extension of previous molecule definition: Molecules are the smallest part of a pure substance that can exist separately, that is; they are particles that can move independently of one another.

.2.9 Distinguish between molecules containing one atom (the noble gases) and molecules with more than one atom Monatomic Molecules
* Comprise of just one atom
* Noble gases
* Examples: He, Ar
Diatomic Molecules
* Comprise of two atoms
* Examples: O2, N2, HI, CO
Triatomic Molecules
* Comprise of three atoms
* Examples: O3, H2O, SO2
Tetra-atomic Molecules
* Comprise of four atoms
* Examples: P4, NH3

.2.10 Describe the formation of covalent molecules in terms of sharing of electrons Covalent bonds Non metal atoms can achieve electron shell stability by sharing electron pairs with other non metals (forming molecules) * Single bond – one pair shared

* Double bond – two pairs are shared
* Triple bond – three pairs are shared

* The ratio of atoms/number of bonds in a covalent compound depends on the valency of each element * Valency the number of electron pairs that are shared by an element in the molecular compound * The formula for a molecular compound gives the molecular formula may not be the simplest, but cannot exist in smaller terms

.2.11 Construct formulae for compounds formed from:
* Ions
* The sum of the positive and negative valencies of ionic compounds is zero (swap valence numbers) * The metal (closer to bottom left corner goes first)
* For polyatomic remember to bracket
* Example: Ag has valency of +1, S has valency of -2. Therefore becomes: Ag2S

* Atoms sharing electrons
* Simple binary sum of valencies of one element should match the sum of valencies for other * Three elements sum of the valencies of the first 2 should equal the total valency of the third * The element closest to bottom left hand corner goes first * Example: Br has valency of 5, F has valency of 1. Therefore becomes: BrF5

8.2.4 Energy is required to extract elements from their naturally occurring sources

.2.1 Identify the differences between physical and chemical change in terms of rearrangement of particles Chemical Change| Physical Change|
A new chemical substance is formed * Bonds are broken in reactants, bonds are formed in products * There is a rearrangement of valency electrons Requires lots of energy Difficult to reverse Examples: * Extracting a metal from a mineral * Electrolysis of water * Dissolving a metal in an acid * | Does not lead to the formation of new chemical substanceExamples: * Changing state – atoms gain/lose kinetic energy * Changing shape * Separating mixtures Can be reversed|

.2.2 Summarise the differences between the boiling and electrolysis of water as an example of the difference between physical and chemical change Electrolysis of Water Chemical Change
* Process: Dilute sulphuric acid is added to water to increase the conductivity of the water. The water is electrolysed, breaking it down to form hydrogen and oxygen gases * The bonds between the hydrogen and oxygen are broken -- which rearrange into diatomic molecules * The ratio formed is 2H:1O (with each hydrogen taking one electron, and each oxygen taking 6 electrons) * Requires large amounts of energy from electrolysis

* Difficult to reverse, cannot put gases together, must have a flame/spark

Boiling of Water Physical Change
* Process: A flame heats liquid water. Water evaporates, changing state to water vapour (gas) * No bonds are broken or formed in this process, however the atoms have gained kinetic energy (from the flame) making it into a gas separate individual water molecules * Can easily be reversed by cooling the gas (condensation)

.2.3 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 Decomposition reaction the process of breaking a compound down into its component elements or simpler compounds Synthesis reaction the formation of a compound from its elements or a more complex compound from simpler compounds

Decomposition Reactions:
* Heat energy in industrialised society to decompose minerals to produce metals in smelters * UV Light energy in nature decomposes ozone molecules into oxygen gas and oxygen radicals (prevents most UV rays reaching Earth’s surface) * Electrical energy in lightning provides energy to gas molecules for decomposition reactions in atmosphere * Chemical energy in airbags sodium azide decomposes by detonation to produce a large volume of nitrogen gas to inflate bag for protection

Synthesis Reactions:
* Electrical energy in lightning provides energy to nitrogen and oxygen molecules for synthesis of nitric oxide. Electrical energy in car engine spark causes nitrogen and oxygen molecules to combine * Heat energy in ammonia industry synthesises ammonia directly by combining nitrogen and hydrogen gases at high temperatures and pressures with a catalyst * Light energy in photosynthesis needed for the synthesis of carbon dioxide and water molecules to form glucose

.2.4 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 amount of energy (heat, light, electrical) needed to separate atoms in a compound is a measure of strength of bond (attraction between ions or attraction between bonds)

Ionic Bonds
The stronger the bonds, the more difficult it is to decompose

Covalent Bonds
Multiple bonds are stronger than single bonds

8.2.5 The properties of elements and compounds are determined by their bonding and structure

.2.1 Identify differences between physical and chemical properties of elements, compounds and mixtures Chemical and physical properties of elements, compounds and mixtures are very different. Example Iron oxide, iron and oxygen, iron ore

| Iron| Oxygen| Iron Oxide|
Appearance| Grey-silvery lustrous metal| Colourless gas| Red/brown solid| Density (g/cm3)| 7.9| 0.0013| 5.2|
Melting point (°C)| 1535| -219 | 1565|
Boiling point (°C)| 2750| -183| |
Solubility in water (g/100g)| Insoluble| 0.004| Insoluble | Reaction with hydrochloric acid| Forms iron chloride and hydrogen gas| No reaction| Forms iron chloride and water|
.2.2 Describe the physical properties used to classify compounds as ionic or covalent molecular or covalent network Compounds can be classified into three groups depending on their physical properties: | Ionic| Covalent Molecular| Covalent Network|

Particles forming| Cations and anions| Molecules| Atoms| Forces holding together| Ionic bonds (relatively strong)| Intermolecular forces (weak)| Covalent bonds (strong)| Electrical conductivity | S: ZeroMolten: GoodAq: Good| Zero| Zero | Melting Point| High| Low| Very high|

Hardness| Hard| Soft| Very hard|
Malleability| Brittle| Brittle| Brittle |

.2.3 Distinguish between metallic, ionic and covalent bonds
Ionic Atoms lose or gain electrons to become cations or anions, which are then bonded by electrostatic attraction Covalent Atoms share electrons
Metallic Metals lose their valence electron/s, which form a negative sea holding the now negative metal ions together by electrostatic attraction

.2.4 Describe metals as threedimensional lattices of ions in a sea of electrons Model of metals:
* Positive metal ions are arranged in a three-dimensional lattice * Delocalised electrons move throughout the lattice
* Have been lost from the valence shell of each metal atom * Attraction between positive metal ions and the delocalised electrons stabilise and hold the lattice together * This attraction is known as a metallic bond

* Metallic bonds are strong

.2.5 Describe ionic compounds in terms of repeating three-dimensional lattices of ions Continuous 3D arrangement of cations and anions, held together by electrostatic attraction…

.2.6 Explain why the formula for an ionic compound is an empirical formula Empirical formula represents the atomic composition as the simplest, whole number ratio Molecular formula indicates the atomic composition in a particular molecule

* Covalent molecular compounds are discreet, therefore the ratio is very clear. For example, the formula for ethane is C2H6. This is its molecular formula. However, because this isn’t the simplest formula, the empirical formula is CH3 (this compound doesn’t exist). * Ionic compounds are continuous in 3D dimensions, without discrete molecules, therefore the simplest repeating unit of the crystal is identifies, known as a unit cell. This determines the simplest ratio of atoms, its empirical formula.

.2.7 Identify common elements that exist as molecules or as covalent lattices Covalent Network:
* Diamond (allotrope of Carbon) stronger bonds than graphite * Graphite (allotrope of Carbon)
* Silicon dioxide (compound)
Covalent Molecular
* Ice (solid water… compound)
* Dry ice (solid carbon dioxide… compound)
* Sulfur
* Iodine
* Phosphorus

.2.8 Explain the relationship between the properties of conductivity and hardness and the structure of ionic, covalent molecular and covalent network structures Ionic Lattice
* Strong ionic bonds:
* High melting point
* Hard
* Brittle shear force pushes similarly charged ions together = strong repulsion * Don’t have free electrons:
* Don’t conduct electricity
* Conduct in molten/aqueus state ions are no longer bound together, so there are available electrons

Covalent Molecular
* Discrete molecules arranged in patterns (held in place by WEAK intermolecular forces little energy required to break) * Low melting points
* Very soft
* No mobile electrons (each molecule is neutral)
* Don’t conduct electricity (not even when dissolved or melted)

Covalent Network
* Very strong covalent bonds:
* Very high melting point
* Very hard
* Brittle shear force pushes similarly charged ions together = VERY strong repulsion * No free electrons or ions
* Doesn’t conduct electricity (except graphite because carbon structure has free electrons) * Highly insoluble

Exothermic reaction: releases energy
* Therefore less energy needed to break bonds in reactants * Bonds are stronger in products, so more energy released when bonds are formed * Occur on their own
Opposite in endothermic reactions

Type of Reaction| Description | Example|
Combustion| Oxidation reaction – burning/fast| |
Corrosion| Oxidation reaction – deterioration of a metal/slow | 2Fe + 3O2 --> 2Fe2O3| Precipitation| | |
Acids on Metals| | |
Acids on Carbonates| | |
Neutralisation | | |
Decomposition | | |
Reactive Metal and Water| Reactive metal + water metallic hydroxide and hydrogen| Calcium + Water Calcium hydroxide + hydrogen |

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