THE LIVING AND NON LIVING COMPONENTS OF THE EARTH CONTAIN MIXTURES

Particle TheoryAll matter consists of many very small particles constantly moving in a continual state of motion and the degree of movement depends on the energy and the relationship to other particles

Elements (pure) Atoms (pure) Compounds Mixtures (impure) Only one

kind of atom Cannot be

simplified Can be

metals (mostly), non metals or semi metals

Smallest particle elements are made up of them

John Dalton Theory = atoms are indestructible and indivisible

All atoms of an element are identical

Atoms of different elements combine to form compounds

Connt be created/destroyed

Elements combined chemically (constant composition)

Metals react with non metals = ionic

Metals react with each other = covalent

E.g. H2O is a compounds

Pure substances

Two or more substances mixed but not chemically combined

Can be physically separated Homogeneous

o Substances spread evenly throughout

o Cant see different partso E.g. sea water

Heterogeneouso Not spread evenly

throughouto Includes suspensionso E.g. dirt and water and oil

and water, salad dressing

SPHERES

Lithosphere All cold hard solid rock of planets i.e. crust, semi solid under crust, hot liquid near

centre, solid iron centre Different minerals combined in different proportions from rocks e.g. oxides,

carbonates, sulfides which are mine to manufacture metals E.g. quartz, calcite, olivine (all minerals) MOST COMMON ELEMENTS = oxygen (46%), silicon (28%) and aluminum (8%)

Hydrosphere Refers to whole planet’s water in any form Most abundant compound in this sphere is water = most abundant elements are

hydrogen and oxygen Most found in oceans and polar ice caps

MOST COMMON ELEMENTS = oxygen (86%), hydrogen (11%) and chlorine (1.9%)

Atmosphere There are different layers we live in the troposphere Refers to all of Earth’s air Major gases = nitrogen and oxygen MOST COMMON ELEMENTS = nitrogen (79%), oxygen (21%) and argon (nearly

1%)

Biosphere Refers to planet’s living organisms Ecosystems occupy all other spheres Composed of cells – therefore large amount of oxygen and hydrogen Cells = complex carbon compounds carbs, fats, proteins etc.

SEPARATION TECHNIQUES

Method Mixture used for Property used to achieve separation ExampleFiltration Solids (which are

not dissolved in liquids)

Used for heterogeneous mixtures

Particle size

Residue – substance left in filter paperFiltrate – soluble substance dissolved in solvent (through filter paper)

Sand and water

Evaporation Dissolved solids in liquids

Differing boiling points of substances

Good for when you want to collect the solid rather than the liquid

Salt water

Crystallization Used to purify impure salts

A concentrated solution of the impure salt is allowed to cool and crystallize resulting in a higher purity of crystal

Used to purify impure salts – copper sulfate

Sieving Solids of different sizes

Particle size Separate sand from gravel in a rock quarry

Sedimentation Undissolved solids from liquids

Particle size and density

Solids are left to settle and the liquid is then carefully poured off the top (decanting)

Soil from water

Chromatography Dissolved solids in liquids esp. if coloured

Different adsorption to a stationary phase

Chromatography paper contains pores draw line in pencil and place dots of dye on the line. Then place it into a solvent, solvent travels up the paper through capillary action

Analytical testing e.g. water contaminants

Separate food colouring or inks

Distillation Liquids and dissolved solids

This is good for when you want the liquid (cant use evaporation)

Use of different boiling points

Solution is heated, liquid evaporates and condense back into liquid in a separate container

Salt water when you want the water

Fractional distillation

Separating liquids that are miscible or gases

Uses boiling points (better for when substances have boiling points closer together

Same apparatus but slightly different process

Cordial and water

Separating funnal Separating immiscible liquids

Different densities

Valve is closed after the more dense liquid is let through

Oil and water

GRAVIMETRIC ANALYSIS Determining masses of substances present in a sample

Reasons for wanting to knowo Decide whether newly discovered minerals contain a high enough % of

required compound for its extraction to be efficiento Determine the composition of soil suitable for certain cropso View the amount of particles in the air to determine pollution growtho Compare the make up of competing products

SYSTEMATIC NAMING

First elements name is unchanged Second elements name is modified (add ide) First element is usually the one which is further to the left of the periodic table (or if in

the same group, further down)

Naming compounds that have multiples of same element

1 – mono2 – di3 – tri4 – tetra 5 – penta6 – hexa

EXAMPLESCO – carbon monoxideCO2 – carbon dioxidePCl3 – phosphorus trichlorideCCl4 – carbon tetrachlorideH2O – dihydrogen monoxideH2O2 – dihydrogen dioxide

ALTHOUGH MOST ELEMENTS ARE FOUND IN COMBINATIONS, SOME ARE FOUND AS UNCOMBINED ELEMENTS

REACTIVITY

Most elements are chemically reactive (react to form compounds) General rule: more reactive = less chance of finding it uncombined in nature Examples: sodium, potassium, calcium, fluorine are all v. reactive therefore never

found as free elements Noble gases and metals are found pure = stable = full valence shells

PROPERTIES OF METALS NON METALS AND SEMI METALS

Metals Non metals Semi metals Solids at room temp Shiny (lustrous) Good conductors of

heat and electricity Malleable Ductile

Solid, liquid (only bromine) and gases at room temp

Dull appearance Brittle Non conductors

Low sheen Moderately malleable Semi conductors of electricity B, Si, Ge, As, Sb in pure state

wont conduct, but will when impurity is added

TESTS USED

Property TestLustre Clean with abrasive paperConductivity of electricity

Connect power supply, globe or ammeter and material in simple circuit and if the globe lights/ammeter reads it conducts

Hardness Scratch with steel knife to see if harder than steelMalleability Use hammer to see if it will flatten out, shatter or form a powderSolubility in water Dip sample in boiling water for 1 minute to observe

PROPERTIES RELATIONSHIP TO USES Physical properties determine which metal/element will be used for particular

products Properties include melting point, density, conductivity, hardness, strength etc.

METALS

Main uses for common metals: building materials, cars, planes, machinery, electrical wiring, appliances and household goods

EXAMPLESo Aluminum aircraft because of low density and high strengtho Copper electrical wiring, high conductivity and ductilityo Iron cars and trains because of high tensile strength

Other factors determine the uses such as cost and chemical reactivity Few are used in uncombined form (risk of being reactive)

NON METALS Uses are based on their properties EXAMPLES

o Helium lighter than air and unreactive therefore used in balloonso Carbon (graphite) conducts electricity and is soft therefore used as an

electrode in dry cells

SEMI METALS o Mixtures of silicon and germanium conduct electricity and are used as

semi conductors in computer chipso

ELEMENTS ON EARTH ARE MOSTLY FOUND AS COMPOUNDS BECAUSE OF ATOMIC INTERACTIONS

SOLIDS LIQUIDS AND GASES

Solids Liquids Gases Particles packed

tightly together Strong forces

holding particles to one another

Definite shapes therefore cannot be compressed

Relatively hard Particles not

stationary, but vibrate in their places

Arranged in much less orderly fashion

Move about much more freely (more energy)

Forces between particles are relatively weaker

Takes shape of container Motion of particles =

actual free movement therefore cannot be compressed much

Solids have greater density (except water)

Particles are much further apart

Very rapid motion No significant inter

particle forces Expand to fill

whole volume (rapid motion)

Big spaces between particles = easily compressible

ELECTRON CONFIGURATION

1st shell 2 electrons2nd shell 8 electrons3rd shell (holds up to 18, but for first 20 elements, only holds 8)4th shell 32 electrons5th shell 50 electrons

E.g. Ca (20 atomic number) electron configuration 2, 8, 8, 2

E.g. Sc (21 atomic number) electron configuration 2, 8, 9, 2

Each electron in the first shell has a constant amount of energy as does each in the 2nd where those in outermost shells have higher energy amounts

Clear relation between reactivity and electron configurationo Noble gases: no chemical reactions = completely stable and full outer shello Alkali metals: similar chemical properties (all lose one electron)o Halogens: similar chemical properties (all gain one electron)

Chemical reactions:o All elements undergo chemical reactions to form compounds to achieve

noble gas configuration

Valence electrons:o Electrons in incompletely filled highest shell are called valenceo These are involved in chemical bonding

FORMATION OF IONS

By gaining/losing/sharing electrons, atoms achieve full shells = stable Metal atoms = positive ions (lose electrons) Non metals = negative ions (gain electrons) This is because electrons have a negative charge

e.g. Na (2, 8, 1) wants to lose one electron and Cl (2, 8, 7) wants to gain one electronand therefore they form NaCl

PERIODIC TABLE AND IONS

Atomic number number of protons = number of electrons Mass number mass of nucleus (protons + neutrons) Group 1 metals lose 1 electron = singly charged positive ions

Group 2 metals lose 2 electrons = doubly charged positive ions Group 6 non metals gain 2 electrons = doubly charged negative ions Group 7 halogens gain one electron = single charged negative ions Transition metals all lose electrons to form positive ions Elements 3 places away from noble gas conficguration may form ions e.g. Al+

General rule: metals generally form positive ions and non metals (when they form ions) form negative ions

LEWIS DOT DIAGRAMS

Concerned with valence electrons

e.g.

For covalent bonding molecules and molecular ions1. Write formula of compound2. Count all valence electrons of atoms in formula (valency of an element is also

the number of bonds it wants to make)3. Draw skeletal diagram and place 1 atom in the middle4. Place pair of electrons between each atoms in skeletal5. Check to see if surrounding atoms have their required number to be stable

(filled shell)6. Place remaining pairs around central atoms

e.g. NH3

ENERGY IS REQUIRED TO EXTRACT ELEMENTS FROM THEIR NATURALLY OCCURRING SOURCES

PHYSICAL AND CHEMICAL CHANGE

Chemical change (reaction) Physical change At least one new substance is

formed Difficult to reverse (because atoms

have been rearranged or bonded and therefore a lot of energy is needed to break these bonds)

Generally a large input or output of energy

Mass is conserved (atoms are rearranged but no new energy is created)

No new substance is formed Easily reversed e.g. changing state Relatively small energy changes

involved e.g. evaporate alcohol Mass is conserved (no new matter is

created)

Chemical Properties Physical properties Properties that relate to the chemical

reactions that substance undergo E.g. ease of decomposition by heat Effect of light Reactivity with other substances

such as oxygen, chlorine and sulfur

Properties that relate to physical change

E.g. boiling and melting points Appearance Density Conductivity Hardness

DIFFERENCE BETWEEN BOILING AND ELECTROLYSIS OF WATER

These two processes show the differences between physical and chemical change Electrolysis produces two new substances (hydrogen and oxygen) whereas boiling

doesn’t produce any new substances it just changes the state of the liquid to a gas Electrolysis is difficult to reverse as you need to mix gases together at very high

temperatures whereas boiling is easily reversed Electrolysis requires much more energy than boiling (needs 20-30 kj per gram)

whereas boiling only needs approximately 2.3 kj per gram – because you are separating the bonds between hydrogen and oxygen whereas when boiling you are merely weakening the weaker intermolecular forces

Energy release during reactions Decomposing a compound requires a large input of energy need to break strong

chemical bonds holding atoms together in the molecule (strong electrostatic forces in ionic bonds and strong covalent bonds holding atoms together in molecules or lattices)

As a rule: the stronger the chemical bonding in a compound is, the more energy is needed to decompose the compound into atoms

The stronger the chemical bonding in a compound is, the more energy is released when the compound is formed from its atoms in the first place

Everyday Exampleso Decomposition: air bags in cars, sodium azide is decomposed to sodium and

nitrogen by igniting it with a detonating capo Rusting of iron to form iron oxide

THE PROPERTIES OF ELEMENTS AND COMPOUNDS ARE DETERMINED BY THEIR BONDING AND STRUCTURE

BONDING

Ionic Covalent Molecular Metallic Giant Covalent Always compounds Solids at room temperature High melting points

(usually over 400˚C) High boiling points (usually

over 1000˚C) Very high melting and

boiling points because when you melt an ionic solid you break up the electrostatic forces between ions and they are very strong so much energy is needed to break them

Hard because of the strong electrostatic forces between ions

Brittle if enough force is applied the ions of the same charge can come close together, they repel, causing the crystal to shatter

As solids, they DO NOT conduct electricity the ions are tightly bound into an orderly array and so there are no free moving electrons allowing the movement of a current through the substance

When molten or in aqueous solution they DO conduct electricity when they melt this arrangement of ions is broken up, allowing electrons and therefore current to travel freely through

Ionic compound formulas are empirical as they show the ratio of the atoms as they appear in the compound

Can be either compounds or elements e.g. H2 or O2

At room temperature can be in any stateGases = e.g. nitrogenLiquids = water, methanolSolids = CBr4

Low melting points (usually below 200˚C) involves disrupting the orderly arrangement of molecules and is therefore easier as these forces are weaker

Low boiling points (usually below 400˚C) covalent molecules involve strong bonds within the molecule but the intermolecular forces are weak, therefore it needs not much energy to break the forces between each molecule, making it easy to boil

When solid they are soft intermolecular forces are easy to overcome therefore it is easy to distort this solid

Pure covalent substances DO NOT conduct electricity with as solids or liquids (only when impure) there are no charged particles free to move as the electrons are localized in pairs and there are no ions

In aqueous solution they DO NOT conduct electricity unless they react with water to form ions

Solid at room temperature (except mercury)

Hard and have high melting points very strong electrostatic bonds between sea of electrons and positive ions

Conduct electricity delocalized electrons allow current to flow through

Consist of an orderly giant 3D structure (lattice) which is an array of positive ions held together by a sea of delocalized electrons

Metallic bond = force of attraction between free electrons and metal positive ions

Metals are malleable and ductile allows the layers to slide over each other meaning that their shape can be distorted and they do not break

Solids in which the covalent bonding extends throughout the whole crystal

Have the same properties as molecular covalent bonded substances

Oppositely charged ions held in a 3D lattice by electrostatic attraction

1:1 ratio (each pos. surrounded by one neg.)

EXAMPLESo Diamond

Each carbon atom is covalently bonded to four other carbon atoms billions of atoms are bonded in this way to form a covalent network solid

o SilicaSiO2 – silicon wants to form 4 covalent bonds and oxygen wants to form two = each silicon atom is bonded to four oxygen’s formula for a covalent lattice compound shows ratio in which the atoms are present = empirical formula

Diamond

HOW IT WORKS

Ionic Covalent

EXAMPLESElements as covalent molecules:

H2, F2, Cl2, O2, N2 all diatomic gases Phosphorus and sulfur exist as covalent molecules P4 and S8

Carbon exists as diamond which is a 3D lattice and as graphite (2D lattice) Noble gases (He, Ne, Ar, Kr, Xe, Rn) all exists as monoatomic molecules Graphite

o Form of carbono Covalently bonded however it conducts electricity layered structure where

each carbon only uses three electrons for bonding and the fourth is delocalized, allowed to move freely throughout the layer – free moving electrons = can conduct

DEFINITIONSSolution: homogeneous mixture in which one substance (solute) is dissolved in another subatnce (solvent)

Suspensions: heterogeneous mixture where particles are large enough to be seen

Colloid: mixture where size of particels are betwee suspension and solution

Diatomic molecules: more than one atom e.g. O2 N2 H2 etc, = all gases

Monoatomic molecules: exist as independent atoms e.g. all noble gases, do not need to bond to be stable

Covalent bonds: pairs of atoms that are sharing electrons form covalent bonds

METALS

1. METALS HAVE BEEN USED FOR MANY THOUSANDS OF YEARS

Uses of metals through history and today

Metal Historical Use Modern Use PropertiesCopper Ornaments, domestic

utensilsElectrical wiring, pipes and plumbing fittings, electroplating, jewellery and household decorations

High conductivity Ductility High melting point

(hard to melt), soft and doesn’t corrode

Iron and steel Tools and weapons Railways and bridges, car bodies, ships and trains, reinforcing in concrete, pipes, nails, nuts and bolts, containers and heavy machinery

Higher melting point than copper

Hard and tough (high tensile strength)

Iron is malleable – easily rolled into sheets

However it corrodes

Bronze Tools and weapons Statues, door and window frames, rails, chutes

Hard, resists corrosion, easily cast

Aluminium Building Planes, car parts, high voltage electrical transmission lines, drink cans

Low density Good tensile

strength High resistance to

corrosionZinc Galvanising iron, protective paints, casing for dry

cells (batteries) Low melting point

Lead Car batteries, plumbing, in crystal glass as glaze for pottery

High resistance to corrosion

Use of alloys in relation to properties

Alloy: homogeneous mixture of a metal with one or more other elements. The proportions of the elements in the alloy determine the properties and use of the alloy

Alloy Use PropertiesSolder (30-60% tin with lead) Ships’ propellers,

casting statues Hard, easily cast,

resists corrosionBrass (50-60% copper with zinc)

Plumbing fittings, musical instruments, decorations

Lustrous gold appearance, hard but easily machined

Steel MILD: Car bodies, MILD: soft, malleable

pipes, nuts and bolts STRUCTURAL:

beams, railways, reinforcing

HIGH-CARBON: knives and tools e.g. hammers

ALLOY STEELS: electromagnets, safes, ball bearings

STAINLESS: food processing things, kitchen sinks, surgical and dental instruments, razor blades

STRUCTURAL: hard, high tensile strength

HIGH-CARBON: very hard

ALLOY STEELS: hard and shock resistant hard at high temps, easily de/magnetised

STAINLESS: hard, resist corrosion

Energy and oresIn order to extract metals from their ores, energy is required to break the existing bonds in the minerals. In ancient times, ores were heated with carbon (charcoal).

More metals now than 200 years ago – why? Before 1800 only ten metals were in use: copper, iron, tin, lead, gold, silver, zinc,

mercury, bismuth and platinum. They either occurred in the Earth as uncombined elements or could be extracted by heating with carbon.

Developments of technology more metals can be extracted from their ores and can be detected in the ground and separated from their combined forms

Abundance in the Earth’s crust was a major factor in using metals such as rhodium Availability of metals for commercial and industrial use depends upon:

o Abundance of metal in the Earth’s crusto Ease of extracting the metal from its naturally occurring ore

Chronology of ages

BRONZE AGE (2300-700 BC)

The Bronze Age in the Middle East (known at this time as the Near East) is divided into three main periods (the dates are very approximate):

EBA - Early Bronze Age (c.3500-2000 BC) Turkey = copper and tin MBA - Middle Bronze Age (c.2000-1600 BC) Copper mined in Egypt LBA - Late Bronze Age (c.1600-1200 BC) Dominating battles using iron weapons

IRON AGE (700BC – AD1)

Iron used predominantly for tools and weapons. Extracted through heating it with charcoal then softening – alloys included castiron

MODERN AGE (AD1 – PRESENT)

Aluminium and iron used Extraction: Blast furnace for iron, electrolysis from aluminium oxide for aluminium Alloys: carbon, stainless and alloy steel

FUTURE AGE

Important substances: ceramics, plastics, composite metals Ceramics – from sand (fired in kiln) used for windows, glasses, cooing utensils, food

storage etc. because of its high melting temp, low density, high strength, corrosion resistance

Plastics – used for electrical wiring, insulation, bottles, ropes, toys because of its hard/soft, no corrosion, flexible, tough and lightweight

Composites – used for sport equipment, race cars, structural and building materials because of its superior properties compared to that of its singular components

2. METALS DIFFER IN THEIR REACTIVITY WITH OTHER CHEMICALS AND THIS INFLUENCES THEIR USES

Reactions of metals with water, oxygen and acids

Metal and oxygen React to form oxide

Metal and water Some react to produce hydroxides/oxides and hydrogen

When water in form of steam = forms oxides

When water in liquid form = hydroxide

Metal and acid Many react with acids (hydrochloric and sulphuric) to produce a salt and hydrogen

To test for hydrogen you do the pop test (with a lit splint)

METAL + OXYGEN METAL OXIDE

METAL + ACID SALT + HYDROGEN

METAL + WATER (steam) METAL OXIDE + HYDROGEN

METAL + WATER (liquid) METAL HYDROXIDE + HYDROGEN

Activity series

Uses the reactivity of metals with oxygen, water and dilute acids to draw up a list of metals in order of decreasing reactivity (activity series). This tells us that any metal above another in the activity series can displace another in a reaction.

K (potassium)Na (sodium)Li (lithium)

Ba (barium)Ca (calcium)

Mg (Magnesium)Al (Aluminium)

Zn (Zinc)Fe (Iron)Sn (Tin)

Pb (Lead)Cu (Copper)Ag (Silver)

Pt (platinum)Au (gold)

A more reactive metal (top of activity series) when reacted with a less reactive metal will displace that metal in the equation

Therefore the less reactive metal will be left in its pure form and the more reactive metal will take its place in the compound

Element Reaction with oxygen

Reaction with water Reaction with acid

PotassiumTarnishes Rapidly

ViolentViolentlySodium

Calcium Moderately

MagnesiumBurns brightly when heated

Slowly in cold water VigorouslyAluminum

Reacts only in steamModeratelyCarbon

Zinc

Reacts slowly when heated

IronTin SlowlyLead Extremely slowly in

steamVery slowly

Hydrogen

No reactionCopper

No reaction in most acids

SilverNo reaction when heated

GoldPlatinum

NOTE: do not worry about carbon or hydrogen regarding their reactions – simply included to see how reactive they are in relation to the other metals

Uses of metals based on properties

Some situations where the choice of metal is based heavily upon chemical reactivity (particularly with oxygen, water and dilute acids) are:

Roof guttering non-reactive but expensive aluminium Water pipes expensive but non-reactive copper Body implants expensive but extremely inert titanium

Relationship between reactivity of metals and their position of periodic table

The reactivity of the metals decreases as you move from left to right across the table because the atoms have more electrons that they need to lose and energy is needed to strip those electrons away from the increasingly more positive nucleus.

The reactivity of metals increases as you go down the group because the outer electrons (which are lost when the metals reacts) are increasingly further from the nucleus. This means that the nucleus exerts less of an electrostatic pull on the outer electron.

Ionisation energy

What is it: measure of the energy that is required to remove an electron from an atom. The atom is in its gaseous state. The electron that is removed comes from the outermost shell E.g. Na (g) – e- Na+ (g)

The higher the ionisation energy, the more difficult it is to remove an electronGenerally the relationship between the periodic table and ionisation energy is:

LHS RHS ionisation energy increases TOP BOTTOM ionisation energy decreases

The lower the ionisation energy, the more reactive is the element

The first ionisation energy measures the energy required to remove the first electron from an atom

The second ionisation energy measures the energy required to remove the 2nd electron from the gaseous +1 ion

Oxidation and reduction

Whenever a metal undergoes a reaction especially with water and acids, they lose electrons to form positive ions

1 st Definition

When an atom loses electrons = oxidised

When an atom gains electrons = reduction

In any equation there is no overall loss of electrons therefore oxidation and reduction occur simultaneously and are known as redox reactions

Agents

Oxidising Agent Reducing Agent Accepts electrons Causes the oxidation of another

substance Is always itself reduced Has its oxidation state decreased

Donates electrons Causes reduction of another

substance Is always itself oxidised Has its oxidation state increased

2 nd Definition

Gain of oxygen = oxidation

Loss of oxygen = reduction

Oxidation States

Idea of oxidation state (number) is a way to track oxidation and reduction changes that occur in chemical reactions

A set of rules have been assigned to elements, ions and compounds where they have an oxidation state at the start and end of reaction

Rules1. All elements in their standard states have oxidation numbers of 02. The charges on the ion is its oxidation number

E.g. Na+ = +1 oxidation number O2- = -2 oxidation number

3. Certain atoms have been assigned fixed oxidation numbers e.g. H in all of its compounds will have an oxidation number of +1 O in all of its compounds will be -2 F will always be -1 in its compounds Cl will always be -1 in its compounds provided that oxygen isn’t also present

Using these rules we can determine the other elements present. The sum of the oxidation numbers in a compound or ion, is equal to the charge or O if the charge is neutral

Third Definition

Oxidation: increase in oxidation number

Reduction: reduction in oxidation number

Example

1) Write equation

Mg (s) + 2HCl (aq) MgCl2 (aq) + H2 (g)

2) Rewrite equation, separating the compounds into their lone ions

HCl = H+ and Cl-

MgCl2 = Mg2+ and 2Cl-

Therefore new equation: Mg + 2H+ + 2Cl- Mg2+ + 2Cl- + H2

3) Take note of the spectator ions (the ones that do not change after the reaction)

In this case the 2Cl- is the spectator ion, the rest have changed after the reaction

4) Write net ionic equation (equations where spectator ions are left out)

Mg (s) +2H+ (aq) Mg2+ (aq) +H2

5) Write half equations, separating the separate reactions between elements

Mg (s) – 2e- Mg2+ (aq)2H+ + 2e- H2 (g)Therefore it is obvious that 2 electrons have been donated from magnesium to hydrogen

6) Decide which have been oxidised and which have been reduced

In this case the magnesium has been oxidised and the hydrogen has been reduced

Oxidation ReductionIs IsLoss Gain

THE RELATIVE ABUNDANCE AND EASE OF EXTRACTION OF METALS INFLUENCES THEIR VALUE AND BREADTH OF USE IN THE COMMUNITY

History of periodic table

Dalton Recognised that diferent atoms have different weights and used triads to arrange the elements in order by weight

Dobereiner (1829) Wanted to order elements by chemical patterns and properties. Idea of triads elements reacting in similar ways with therefore similar chemical properties, ultimately leading to groups

Newlands (1864) Pattern of octaves: every 8 elements patters would repeat led to idea of periods

At this time there were approximately 60 elements

Mendeleev (1869) Arranged table according to atomic weights and properties

Left gaps for elements that he believed were left to be discovered

Ramsay (1894) Discovered noble gasesMolesley (1914) Determined every element’s atomic

numberSeaborg (1940) Synthesised transuranic elements

(those after uranium in the Periodic Table)

Relationship between position on periodic table and properties

Electrical Conductivity Left Right = conductivity decreases the further across a period you go as the metallic properties do

Top bottom = further down on the table has high conductivity

Ionisation Energy Left Right = across a period, ionisation energy increases

Top bottom = ionisation energy decreases as you go down a group

Atomic radius Left Right = decrease in atomic radius because the positively charged nucleus that you get as move right gets stronger therefore bringing electrons closer to the nucleus

Top bottom = increases the further down you go as you get more shells = larger radius

Melting/boiling point Peak in group 4 (they tend to form covalent network solids)

Noble gases = v. low = almost never bond

Groups 1 and 2 = metallic bonding = moderate and as you go down it decreases as the metallic bonding decreases

Group 7 = weak intermolecular forces = low m/bp but increases as you go down a group (same as group 8)

Valency Peak in group 4 Refers to the number of elements it

can combine withElectronegativity Relative power of an element to

attract electrons to itself in a covalent bond

Left right = increases (more protons therefore more attraction and pull power)

Top bottom = decreases Therefore F is the most

electronegative elementReactivity Increases going down a group for

group 1 Group 7 – decreases going down No trend going across, just in

separate groups

FOR EFFICIENT RESOURCE USE, INDUSTRIAL CHEMICAL REACTIONS MUST USE MEASURED AMOUNTS OF EACH REACTANT

Relative atomic mass (atomic weight) Most of an atom’s mass is from the nucleus All atomic weights are relative to the mass of Carbon-12 which is exactly 12 Carbon 12 = 6 protons and 6 neutrons Mass is very small and therefore difficult to isolate the mass of one atom For this reason, chemists determine the relative mass of atoms. E.g. a silver atom

has 4 times the mass of a carbon atom. Since they are relative they have no units Examples of relative masses found on the periodic table

o Carbon = 12o Aluminium = 27o Chlorine = 35.5o Gold = 197o Lead = 207.2o Silver = 107.9

Isotopes

Many elements don’t have whole number atomic weights This is because most elements have more than one isotope (atoms with different

numbers of neutrons) Therefore the relative atomic weights are averages of these isotopes

E.g. 99% of carbon is carbon 120.9% is carbon 130.1% is carbon 14 (this is all determine by a mass spectrometer)Average mass = (99x12) + (0.9x13) + (0.1x14) 100 = 12.011

Molecular mass

Molecular mass/weight: sum of the atomic weights of the atoms in a molecular formula

E.g.Molecular weight of C12H22011

M.W. = (12 x Ac) + (22 x AH) + (11 x Ao) = (12 x 12.0) + (22 x 1.01) + (11 x 16.0) = 342.2 (but no units because mass is only relative)

Formula mass

Formula mass/weight: sum of atomic weights atoms in a compound with no discreet molecules (e.g. ionic compound). They describe the ratio of atoms present, but are calculated in the same way as molecular weights

E.g.Ca3(PO4)2 is:Ca3, P2, O8

F.W. = (3 x ACa) + (2 x AP) + (8 x A0) = (3 x 40.1) + (2 x 31) + (8 x 16) = 310.3

The mole

1 mole of any substance is equal to 6.022 x 1023 particles including atoms, ions, electrons, molecules etc.

1 mole = number of particles in 12g of Carbon-12 = 6.022 x 1023

For example: 1 mole of carbon = 12g 1 mole of oxygen = 16g 1 mole of O2 has mass = 32g 1 mole of H2O molecules = 18g

12g of carbon-12 contain 6.022 x 1023 atoms. This is Avogadro’s number and is equivalent to one mole of carbon atoms

Avogadro’s Number = 6.022 x 1023 = 1 mole

n = number of molesm = mass (g)mr = molar mass

n = mass mr

Empirical formula

Empirical formula: simplest whole number ratio of all the elements in a compoundFor example:C4H10 is the molecular formula BUT C2H5 is the empirical formula

Examples1. Find the empirical formula of a compound made up of 20g of Ca, 6g of Carbon and

24g of OxygenCa = 40C = 12O = 16

20 : 6 : 24 (divide by the mass of each element)40 12 16

1 : 1 : 3 (simplify ratios)2 2 2

1 : 1 : 3 (get rid of common denominator)

Therefore the empirical formula is CaCO3

2. Empirical formula from composition A compund is found to be %C = 55.5%, %H = 9.1% and %O = 36.4%

Step 1 assume there is 100g and therefore the percentages are grams

55.5g of Carbon9.1g of Hydrogen36.4g of Oxygen

Step 2 turn g into moles

55.5g/12 = 4.538 moles C9.1g/1 = 9.09 moles H36.4g/16 = 2.275 moles O

Step 3 write in ratio form1.538 : 9.1 : 2.275

Step 4 divide by smallest (to get the ratio to 1: something)4.538 : 9.1 : 2.2752.275 : 2.275: 2.265

Therefore empirical formula is C2H4O

Gay-lussac Law states that gases at equal temperatures and pressures react in whole number

ratios to one another Notice above that the volume ratio is equal to the coefficients in the reaction and

there if no conservation of volumes precursor to idea of moels

Avogadro Law states that equal volumes of gases at the same temperature and pressure

contain the same numbers of molecules Now we have a convenient way of determining gas volumes in a reaction since we

can replace gas volumes for moles E.g. 2 moles of Hydrogen gas and 1 mole of oxygen gas 2 moles of water gas is

same as 2H2 + O2 2H2O In the above equation there are 2 volumes of hydrogen, 1 volume of oxygen and 2

volumes of water

Volume calculations from these laws

All problems are under conditions of 25˚C and 100kPa pressure Mole of a gas 24.79 litres

EXAMPLEWhat volume of oxygen is required for the complete combustion of 11.5g of sodium?4Na (s) + O2 (g) 2Na2O (s)

Ratio is 4:1Number of moles of sodium = 11.5/23 = 0.5Divide this by 4 to get number of moles of oxygen = 0.125Therefore the volume of oxygen = n x 24.79

= 3.09875 litres

Mole calculations

1. Calculating number of moles when given atoms

E.g. 3.01 x 1024 carbon atoms – find the number of molesDivide number of atoms by Avogadro’s number to get number of moles

2. Finding number of moles when given grams

Number of moles = mass (g) usually given in question Molar mass of the compound/element

E.g. How many moles are contained in 1.01g of neon gas?N = 1.01/20.2 = 0.5

3. Calculating mass from number of moles

Mass = no. of moles x atomic weight (found on periodic table)

E.g. calculate the mass of 2.05 moles of rubidium atomsMass = 2.05 x 85.5 = 175g

4. Calculating mass from ratio (when given grams of wrong element)

First – determine ratio between elements by looking at the coefficients of the substances in balanced equationSecond – work out number of moles in the substance that you are given the grams for

Third – use the ratio to work out the moles of the other substanceFourth – use number of moles to calculate mass in g

EXAMPLE: Calculate the mass of sodium needed to react with 10g of titanium chlorideTiCl4 + 4Na Ti + 4NaCl

First – It is obvious that the ratio is 1:4Second – using the 10g of titanium we can work out the number of moles of titanium chloride

Number of moles = 10/19 = 1/19

Third – but its in ratio 1:4 so multiply by four to get the number of moles of the sodium1/19 x 4 = 4/19 moles

Fourth – work out mass of 4/19 moles of sodium

Mass = 4/19 x 23 = 4.8 (1dp) grams

Yield

Theoretical yield what you expect to gain from reactionActual yield what you actually get in reality% yield = actual yield/theoretical yield

EXAMPLE

Part a) Calculate mass of ammonia that can be formed from 12g of Hydrogen3H2 + N2 2NH3

Calculate mass normally by finding no. of moles and multiplying by molar massFound 68g to be the theoretical yield

Part b) 20g was formed in this reaction. Calculate percentage yield20/68 x 100 = 29.4%

THE RELATIVE ABUNDANCE AND EASE OF EXTRACTION OF METALS INFLUENCES THEIR VALUE AND BREADTH OF USE IN THE COMMUNITY

Minerals and ores

MineralsNaturally occurring compounds found in the Earth – where most metals (except Au and Ag) are found. Most common minerals that contain metals in Aus = oxides and sulphides. Minerals are pure compounds and usually crystalline

OresNon-renewable mineral deposits that are economically feasible to extract metals from. Ores are impure and may contain a number of compounds including minerals. They are non-renewable substances because there is a finite number of them that occur naturally on Earth and their abundance can be extinguished.

Examples: Bauxite Al2O3 and Sphalerite ZnS

Economic considerations of mining

Factors affecting metal prices: Worldwide demand for metal Cost of purification cost of transporting metal, relative cost of production including

amount of energy required Abundance of metals in ores

Extraction of copper

1) Copper ore: consists of:Chalcolite

Chalcopyrite

2) Froth flotationCopper mixed with water and frothing agent

This sticks to copper, bring it to the top3) Roasting

Ore is heated with oxygen to remove sulphur components (impurities)

4) Iron RemovalProduced iron silicate

5) ReductionCopper sulphide copper metal (98%) – blister copper

6) Electrolysis = pure copper

Recycling aluminium1. Collection of used products from homes and businesses2. Transport to recycling facility3. Separate Al from impurities (Al in labels and cans etc.)4. Re-smelt the Al into ingots for transport to product manufacturers

Why recycle Less energy required to recycle a metal than is required to extract it from its ore Metal ores are non-renewable natural resources that need to be conserved Less waste to dispose of

WATER

WATER IS DISTRIBUTED ON EARTH AS A SOLID, LIQUID AND GAS

Solute: a substance that is dissolved by a solvent in a solution

Solvent: a substance that dissolves a solute in a solution

Solution: homogeneous mixtures that contain a solvent and solute

Examples: Sea water solvent is water, solute is salt Blood solvent is water, solutes are oxygen and nutrients

Significance on Earth

For living things Habitat for life Natural resource Alters landforms Transport for

nutrients in cells Raw material for

plants in photosynthesis

Solvent for nutrients and oxygen in blood

Solvent for wastes (e.g. CO2 and sweat)

Place where aquatic plants and animals live

Much less fluctuation in temperature – high heat capacity

Drinking and food prep

Washing clothes, dishes etc.

Irrigation Recreation Transportation Hydroelectricity

Moving water in rivers = canyons

Can erode rocks by dissolving minerals

Glaciers change the landscapes from mountains to sea

Spheres

Biosphere Lithosphere Hydrosphere Atmosphere Liquid (as a

solvent for nutrients)

60-90% in most living things (50-75 in humans)

Solid, liquid or chemically bound as waters of hydration

Variable %’s in groundwater, aquifers and rocks

Liquid and solid Approx. 95% in

oceans, lakes/rivers and icecaps

Gas and liquid Up to 5% in air,

variable depending upon environment

Density

DENSITY = MASS/VOLUME

Most substances contract as they cool = increase in density = decrease in kinetic energy of particles

Temperature of water decreases (to 5˚C) the H-bonds arrange themselves so that there is more space between water molecules density drops until solid ice is formed

Low density = ice floats on water

EXPERIMENT

Density of water1. Add an amount of water to beaker and weigh2. Measure temp of water3. Repeat in 10mL increments

What happens: as volume increases, density increases Density of ice

1. Weigh ice cube 2. Measure 100mL of water into beaker3. Submerge ice cube and calculate difference in volume4. Use d = m/v to find density

THE WIDE DISTRIBUTION AND IMPORTANCE OF WATER ON EARTH IS A CONSEQUENCE OF ITS MOLECULAR STRUCTURE AND HYDROGEN BONDING

Lewis Dot Diagrams

CH4 (methane) NH3 (ammonia) H2S (Hydrogen sulfide)

H2O (water) CO2 (carbon dioxide)

Molecular Structure Comparison

Water Ammonia Hydrogen SulfideShape of molecule Bent Pyramidal Bent

M.p (˚C) 0 -78 -86B.p. (˚C) 100 -33 -60

Valence Shell Repulsion Theory

CH4 NH3 H2O CO2

Hydrogen Bonding

Type of intermolecular force that involves a hydrogen atom bonded to an O, N or F atom in one molecule becoming attached to an O, N or F from a different molecule

This results in an unequal sharing of electrons leading to a partial positive charge on the H atom

These bonds are stronger than dipole-dipole and dispersion forces, but about one tenth of that of a normal covalent bond

Page 191 picture

Ammonia Water

Polarity

Bond can be polar or non-polar Depends on the difference in electronegativity of the atoms in that bond Electronegativity: measure of the ease at which an atom attracts electrons (recall that

F is the most electronegative element)

Example 1

Hydrogen fluoride

Polar bond = polar molecule The F has a slightly negative polarity because it is more electronegative and therefore

has more pull on the electrons For something to be a polar molecule with a polar bond one end must be positive

and one end must be negatively charged It is permanently like this

Example 2

NH3 (ammonia)

Ammonia contains 3 separate polar bonds Therefore ammonia is a polar molecule

Example 3

CH4 (methane)

Both carbon and Hydrogen share the same electronegativity Therefore the charges on both are the same Therefore the bond is non-polar and methane is a non-polar molecule

Dipole-Dipole Forces

What are they? Because polar molecules have positive and negative ends, they are able to line up so

that the positive end of one molecule attracts the negative end of another molecule Therefore electrostatic attraction holds the molecules to one another more strongly

than in non-polar molecules These electrostatic attractions are called dipole-dipole froces

Polar Covalent Bonding 2 atoms have difference in electronegativity – the one with HIGHER value tends to

have a partially NEGATIVE charge (there is a stronger attraction to electrons in higher electronegative atoms)

For the same reason usually – the LOWER electronegativity tends to have a POSITIVE charge

Molecular dipoles occur due to the unequal sharing of electrons between atoms in a molecule. 

Net Dipole Not all molecules with polar covalent bonds are polar overall or have a net dipole Some molecules contain polar bonds that cancel out the effect of any single bond

Dipole-Dipole Interactions

When there is an overall net dipole – interactions between them is electrostatic Positive end is attracted to negative end of another molecule Dipole-dipole interactions are not as strong as hydrogen bonds

Properties – Intermolecular Forces

Surface Tension Viscosity Melting/boiling point Cohesive forces = water

molecules are attracted to each other in all

Reistance of a liquid to flow

Two factors affect

The larger the molecule, the higher the melting and boiling

This molecule contains two polar covalent bonds, but no overall net dipole.

directions At surface – molecules

are only attracted either side or down (obviously cant be attracted upwards)

This results in an overall force in towards the liquid

Unbalanced force results in surface tension – also explains why water forms droplets as all of the force is directed inwards

viscosity: size and complexity of molecules AND intermolecular forces

The stronger the forces the more resistance to flow and therefore the higher the viscosity

Water has a relatively high viscosity due to strong hydrogen bonding

point The stronger the

intermolecular forces are, the higher the melting and boiling points are (as melting and boiling involve the breaking of intermolecular forces)

Water has a high melting/boiling point due to the strength of hydrogen bonding

Expansion on Freezing Ice – nearly all molecules are hydrogen bonded to each other Each oxygen has 4 H atoms around it – two covalently bonded and two hydrogen

bonded This arrangement = open structure When melts – individual water molecules are able to move – breaking up hydrogen

bonding Less hydrogen bonding = molecules move closer water higher density than ice

(same mass occupies smaller volumes)

WATER IS AN IMPORTANT SOLVENT

Type of substance Particle interaction with water

Soluble Ionic Compound

E.g. NaCl – made up of pos Na ions and neg Cl ions Lattice – electrostatic attraction Water is polar – pos and neg ends as well Water is attracted to charged ions in salt lattice Dissociation: water molecules surround ions, overcoming

forces between salt particles

Ions in lattice are dispersed in solvent However not all ionic compounds are soluble in water

Soluble molecular compound

E.g. sucrose – Contains many polar OH groups that attract water molecules

This attraction breaks the crystal structure and distributes individual sugar molecules through water

The only soluble molecular compounds are highly polar or can form hydrogen bonds with water – this allows them to break the crystal structure

E.g. ethanol

Soluble/partially Covalent molecular held together by dipole-dipole or

soluble element/compound

dispersion (weak compared to hydrogen bonds) – therefore this type is usually insoluble

E.g. iodine: non-polar substance, held together by weak dispersion forces – therefore insoluble in water – more so in non-polar solvents

E.g. oxygen gas: many gases like O2, N2, CO2 are non-polar and therefore insoluble in water

E.g. Hydrogen Chloride: Molecules go through ionization – HCl reacts with water to produce H and Cl ions in solution, meaning they are highly soluble in water

The solubility of a substance in water depends on whether it reacts with water

Covalent network structure

E.g. Silicon dioxide: v. strong covalent bonds – in rigid lattice Strong bonds cannot be broken by attraction to water molecul Therefore insoluble in water

Substance with large molecules

E.g. cellulose: There are polar OH groups, but log chains lie beside each other and are involved in hydrogen bonding – meaning the OH groups are not accessible to the water molecules. Therefore cellulose is insoluble in water

E.g. polyethene: non-polar and therefore is insoluble in water Generally: the larger the molecule, the less likely it will be

soluble in water unless there are a large number of accessible polar groups

Solubility Relationships Summary

Water is polar molecule with slight negative on oxygen end and slight positive on hydrogen end

LIKE DISSOLVES LIKE – polar solutes dissolve up in polar solvents and non-polar solutes dissolve up in non-polar solvents

Dissolution: process where ions in a solid are dispersed in a solvent (happens with NaCl as they break down into ions)

Ionization: process where covalent molecule reacts with water to form ions in solution (happens between HCl and water)

Covalent network: highly insoluble due to strong lattice structures Large molecules: tend to be insoluble unless they have lots of accessible polar sites Most non-polar molecules: insoluble or slightly soluble in water

Solubility Rules

Soluble Compounds Insoluble Compounds All sodium, Potassium and

Lithium All ammonium Chlorides, Bromides and

Iodides except those of Silver, Mercury and Lead

All Nitrates and Acetates and Sulfates, except Calcium, Strontium, Barium, Lead, Mercury and Silver

Carbonates, Chromates and Phosphates except Sodium, Potassium, Lithium, Caesium and Ammonium

Sulfides except Sodium, Potassium, Lithium, Caesium, Beryllium, Magnesium, Calcium, Strontium, Barium and Ammonium

Hydroxides and Oxides except Sodium, Potassium, Lithium, Caesium, Beryllium, Magnesium, Calcium, Strontium and Barium

Hydrogen Bonding

Attraction between a hydrogen on one molecule and an oxygen particle on another molecule

Intermolecular attraction OR interaction between an F atom and an H atom OR interaction between N atom and an H atom F, N, O, H: all have high electronegativity and therefore always have a slightly

neg/pos charge – causes attractions E.g. ammonia can dissolve up in water due to its polarity

DRAWING FROM EXERCISE BOOK

THE CONCENTRATION OF SALTS IN WATER WILL VARY ACCORDING TO THEIR SOLUBILITY, AND PRECIPITATION CAN OCCUR WHEN THE IONS OF AN INSOLUBLE SALT ARE IN SOLUTION TOGETHER

Formation of Precipitate

If ionic bonds are stronger than interaction with water, compound will be insoluble If two solutions of different ionic compounds are mixed, an insoluble compound

(precipitate) may form Precipitate Examples:

Barium nitrate(aq) + sodium sulphate(aq) barium sulphate(s) + sodium nitrate(aq)

Al3+(aq) + 3NH3(aq) + 3H2O(aq) Al(OH)3(aq) +3NH4

+(aq)

Equilibrium Products react to form reactants as well as reactants forming products

Known as reversible reactions Both take place simultaneously Forward and reverse reactions happening at same rate = dynamic equilibrium

where we can see no observable changes in the system This only happens in ionic compounds Therefore, it looks like nothing is happening because ions are moving in and out of

solution at the same rate

Equilibrium in Saturated Solution Saturated solution: maximum amount of solute is dissolved in a given amount of

solvent In a saturated solution – dissolved solute is in equilibrium with any undissolved solid Any solid that is added to the system doesn’t affect the equilibrium Rate of crystallization and dissolution will be the same

Concentration of Solutions Defined as amount of solute in a specified amount of solvent – there are many ways

of expressing concentration depending on the purpose and circumstances

C = n AND n = mass v mr

Mass per unit volume (g/L)

Used when describing solubility or making solutions involving dissolving a solid in solvent

Example: 2g of fertilizer is dissolved in enough water to make a 50mL solution. Calculate the concentration of the fertilizer in g/L

Solution: m = 2g and V = 0.050L Concentration

Percent by mass (w/w)

Compares weight of solute with weight of solution and is then converted to a percentage

E.g. concentration = 35 (mass of solute) /135 (mass of solution) x 100 = 26%

Percent by volume

Used when liquids are dissolved in liquids Ratio is calculated and converted to percentage E.g. concentration = 2.5 (volume of solute) /30 (volume of solution) x

100 = 5.0% (v/v)Parts per million

Parts of particular solute per million parts of mixture E.g. is the maximum allowable concentration of DDT in beef is 1.5ppm,

would a 500g sample of beef containing 0.075g of DDT be acceptable Solution: M of DDT = 0.075g and M of beef = 500g (make sure both are

in same units) Concentration = 0/075/500 x 1 000 000 = 150ppm (this is 100 times the

allowable limit therefore not permitted)

Dilution Calculations

C1 x V1 = C2 x V2

C1 = initial concentrationV1 = volume to be taken from the original solution and transferred into the second solution with new concentrationC2 = new concentrationV2 = volume of solution (usually stays constant)

WATER HAS A HIGHER HEAT CAPACITY THAN MANY OTHER LIQUIDS

Specific Heat Capacity

Measure of the amount of energy (in joules J) that is needed to raise 1g of a substance by 1˚C

The higher the heat capacity the more resistant a substance is to a change in temperature

Water has a heat capacity of 4.18 JK/g

∆H = -mc∆T(only add in the negative if you need to see whether the reaction is endo or exo thermic)

∆H energy released/absorbed (J)m mass of substance (g)C specific heat capacity (e.g. 4.18Jg-1K-1)∆T change in temperature (final minus initial temperature)

Example:What quantity of energy is required to raise the temperature of 0.5L of water from 20˚C to 100˚C?

∆H = mc∆T∆H = 500g x 4.18 x 80˚C = 167200J = 167kJ

In Chemical Reactions:

∆Hrxn = H(products) – H(reactants)

It is possible to measure the changes in heat content in chemical reactions Change in enthalpy is known as heat of reaction There are two possibilities:

1. Energy of reactants is positive 2. Energy or products will be positive

Energy Changes Involving Dissolution

When an ionic substance dissolves, it may be endothermic (absorbs heat/gets cold) or it may be exothermic (releases heat/warms up). The change in heat (enthalpy) when one mole of a solute dissolves in a solvent is called the molar heat of solution,

∆Hsolutione give Breaking the bonds that hold ions together requires energy (endothermic) Forming bonds with water (hydration) releases energy (exothermic) So, ∆ Hsoln = heat absorbed when breaking bonds – heat released when forming

bonds (in kJ/mol.)

Endo/Exothermic

Endothermic Exothermic Enthalpy/heat content of reactant < products Enthalpy of reactants > products

Heat is absorbed ∆H is positive Heat from surroundings is absorbed - temperature

of surroundings will go down Final temp – initial temp = NEGATIVE This is because the temp of surroundings go

down Final temp < initial temp When solute/solvent relationships are weaker Therefore:

∆T is always negative ∆H is always positive

Heat is released ∆H is negative Heat is given out to the surroundings Final temp – initial temp = POSITIVE This is because the temp of surroundings

(water) goes up Final temp > initial temp When solute/solvent relationship is stronger Therefore:

∆T is always positive ∆H is always negative

Examples from exercise book

Possible Causes of Error

Draft lost from room – heat goes into air and surroundings go up rather than water Depends on how well you read the thermometer Insulation of cup/container Whole experiment is dependent on weighing and measuring the temperature –

therefore they are sources of error

Dissolving Solutes

Ionic Compounds Molecular Compounds Completely dissolve and

dissociate to give ions in solution

E.g: NaCl Na+ + Cl-

Therefore, NaCl is an electrolyte and will conduct electricity

Those that dissolve include: glucose, sucrose They dissolve, but don’t dissociate to give

ions C6H12O6(s) C6H12O6(aq)

Therefore no ions are released and it is a NON-electrolyte

Thus, has no conductivity

Ionic compounds with high solubility = strong electrolyte Ionic compounds with low solubility = weak electrolyte Molecular compounds which don’t dissolve = non-electrolyte

Thermal Properties of Water for Organisms

Need narrow temperature ranges to survive Water in cells provides necessary temperature regulation It can do this due to:

High heat capacity: large additions of hear = small change in temperature High thermal conductivity relative to other liquids (quickly removes heat from

a hot place to a cold one) Water is such a large proportion of most living organisms

Aquatic organisms – high heat capacity of water means their environment maintains a much more stable temperature than land

Water’s prominence in the biosphere means that it has a moderating influence on global temperatures

Thermal Pollution

Refers to human acitivity dramatically changing the temperature of waterwyas

Most commonly use of natural water bodies for industrial cooling e.g. power plants Thermal pollution can have detrimental eggects on aquatic life:

Solubility of oxygen decreases as temperature increases Sharp increase in temperature may directly kill fish and other organisms High temperatures can prevent development of fish eggs

ENERGY

LIVING ORGANISMS MAKE COMPOUNDS WHICH ARE IMPORTANT SOURCES OF ENERGY

Photosynthesis

6CO2(g) + 6H2O(l) (light and chlorophyll) C6H12O6(aq) + 6O2

Endothermic Autotrophs make their own food by photosynethsis Using chlorophyll, they convert light energy to chemical energy Release

Respiration

C6H12O6 CO2 + H2O + energy

Cellular respiration releases the stored chemical energy The amount of energy released during respiration per mole of glucose is equal to the

amount of energy absorbed during photosynthesisUses for respiration

Used directly for activities Converted to protein for growth/repair of tissues Stored as fat for energy reserves Most is released as heat back to the environment

Carbohydrates

Photosynthesis is the production of glucose However, photosynthesis actually makes many other carbohydrates as well Carbohydrates contain the chemical energy that is released by living things

The general equation of a carbohydrate:

Cx(H2O)y insert x and y for different carbs

Common carbs: Glucose and fructose Sucrose and maltose Starch, cellulose and glycogen

Fossil Fuels

Normally as plants die, decomposers (insects, worms, bacteria) break down decaying material water CO2, minerals

Decaying plants/animals (energy released)(decomposers) CO2, H2O, nutrients

Formed when dead and decaying material was buried before complete decomposition

Formed over millions of years due to the heat/pressure beneath the Earth’s surface Energy-rich compounds known as hydrocarbons Chemical potential energy is released when burning in oxygen

Sequence of Production

Coal Often formed in swamps and mangroves

Plant material is anaerobically decomposed (i.e. without oxygen) by anaerobic bacteria

As more and more layers of material are deposited, carbon content increases

Temperature and pressure conditions reduce the amount of oxygen (as CO2) and hydrogen (as CH4)

Some impurities are sulphur and other inorganicsPetroleum Mostly formed from the remains of buried aquatic organisms (e.g.

plankton) broken down by anaerobic bacteria They contain a mixture of hydrocarbons commonly known as ‘crude

oil’ Oil is deposited in porous sedimentary rocks and the less dense oil

moves upwards unless blocked by impermeable rock Most of Australia’s oil deposits are found offshore (e.g. the Gippsland

Basin (Bass Strait) and the North-West Shelf (WA))Natural Gas Mostly the remains of buried aquatic organisms

Often found in a trapped layer just above petroleum deposits Often contains up to 90% methane Also contains propane (C

3H

8) and butane (C

4H

10) which are

liquefied to produce LPG (liquefied petroleum gas)

THERE IS A WIDE VARIETY OF CARBON COMPOUNDS

Carbon Group 4 of periodic table 4 valence electrons Can form 4 covalent bonds Carbon can form 4 single bonds, double bonds and triple bonds with a wide variety of

elements forming nearly 10 000 000 known compounds

Single Bonds Double Bonds Triple Bonds

Tetravalent Carbon

Carbon has 4 valence electrons – therefore forms 4 bonds with other elements to make a full valence shell

All valence electrons are involved in bonding This bonding leads to tetrahedral shapes when all bonds are single Hydrocarbons are made up of carbon bonded to H, but many elements can take its

place such as halogens (Cl, F) and N, O, S

Forms of Carbon

Allotropes: forms of an element with different properties. Carbon has 4:

Arnorphous Soot from incomplete burning of hydrocarbons consisting of shapeless particles

Graphite Thin sheets of six-sided carbon rings in layers held together by weak intermolecular forces

Diamond Crystalline covalent network substanceBuckminsterfullerene ‘Bucky balls’

Contains 60 carbons with 5 and 6 carbon rings arranged in a structure similar to a soccer ball

A VARIETY OF CARBON COMPOUNDS ARE EXTRACTED FROM ORGANIC SOURCES

Fractional Distillation

Crude oil contains a mixture of hydrocarbons ranging from carbon (C1) up to more than C24.

Fractional distillation allows for these components to be separating through a fractionating column

Heat is applied to bottom of the column and lighter compounds with lower boiling points rise to the top

Heavier compounds remain towards the bottom Chains of carbon bonding get longer as you go down The gases at the very top are used for BBQ, camping gas and cooking (they don’t

condense) In the past, the most important organic chemicals were derived from coal, but now

natural gas and crude petroleum provide an alternative source Composition of crude petroleum varies according to its source Rough fractions can then be distilled further to obtain narrower boiling ranges

IUPAC Nonclemature

Homologous series: a family of compounds with the same general formulaStraight-chain alkanes: carbons joined together to form a single chain with no branching.Isomers: compounds that have the same molecular formula, but different structure.

ALKANES

General Formula:

CnH2n+2

All have similar chemical properties Homologous series No functional group Different physical properties because: as chain gets longer, intermolecular forces

between chains gets stronger therefore bouling and melting points start to increase i.e. first four members are gases – as chain gets longer it becomes less volatile and more viscose

No. of C Molecular Empirical Structure1 CH4

MethaneCH4

2 C2H6

EthaneCH3

3 C3H8

PropaneC3H8

4 C4H10 C2H5

Butane

5 C5H12

Pentane(from here the prefixes are normal – hex, hept, oct etc.)

C2H12

ALKENES

General Formula:

CnH2n (where Cn > 2)

Related to the alkanes All contain one C = C double bond There is a functional group (unlike the alkanes) as there is a double bond and this is

where they undergo reactions) The longer the chain the higher the forces between molecules therefore higher

melting/boiling points for longer chains Alkenes have isomers because the double bond can be in a different location above

C4. The location of the double bond is indicated by a numerical prefix counting from the shortest end.

Prefix for 7 carbons is hept

No. of Carbon Molecular Empirical Structure2 C2H4

EtheneCH2

(this is always the same)

3 C3H3

PropeneCH2

4 C4H8

ButeneCH2

5 C5H10

PenteneCH2

Comparison between Anes and Enes

Alkanes Alkenes Boiling and melting points increase

with number of carbon atoms C1 to C4 are gases at room temp C5 to C17 are liquids Greater than C18 are solid

Boiling and melting points increase with increased number of carbons

C2 to C4 are gases at room temp C5 to C15 are liquids All above C16 are solid

Weak dispersion forces that increase in strength with increasing molecular mass

More electrons = stronger forces

Similar to alkanes

Non-polar due to symmetrical shape and similarity of electronegativity of H and C

Similar to alkanes

Densities increase with increased molecular mass

All are less dense than water (float on top)

Similar to alkanes

Volatility Volatility refers to how easily a substance evaporates or changes state Liquid hydrocarbons such as petrol readily evaporate In a closed container, a dynamic equilibrium will be achieved between evaporation

and condensation (i.e. same rate). Once dynamic equilibrium is established, a constant pressure is exerted on the

container known as vapour pressure. The weaker the intermolecular forces, the lower the boiling point and the greater the

volatility.

Safe Storage of Alkanes and Alkenes

Weak intermolecular forces and low molecular weight of alkanes = extreme flammability. Some alkanes are carcinogenic (potential to cause cancer), so safe storage of alkanes (and many other hydrocarbons) needs to consider the following:

Minimise the quantity of material to be stored Store in cool place with good ventilation (flammable cabinet) Avoid inhaling the vapours Use in a fume hood Keep away from sparks or naked flames Store in approved containers (sturdy with narrow neck) Gas cylinders should be regularly checked and stored outside, under cover in well-

ventilated area. They should also be strapped to a permanent structure.

Branching and Naming

COMBUSTION PROVIDES ANOTHER OPPORTUNTY TO EXAMINE THE CONDITIONS UNDER WHICH CHEMICAL REACTIONS OCCUR

Indicators of chemical reactions At least one new substance is formed A change in colour (not always a chemical reaction) A gas is given off (acid + carbonate produces CO2)

Heat is produced or absorbed (combustion) Light is given off (chemiluminescence) A precipitate forms (insoluble ionic compounds)

Combustion is Exothermic Combustion of a material involves its combination with oxygen gas Because combustion reactions release energy in the form of heat and light, they are

exothermic.

Examples of exothermic (many more) Examples of endothermic Combustion reactions Reactions between metals and water Reactions between metals and acid Reactions of acids and bases

Decomposition reactions Photosynthesis Some precipitation reactions Formation of synthesis gas

Why is combustion exothermic?

Energy is required to break bonds therefore the breaking of bonds is endothermic Energy is released when bonds form therefore the forming of bonds is exothermic E.g. when methane burns in oxygen to form carbon dioxde and water C-H bonds in

methane and O-O bonds in oxygen gas are broken and C-O and H-O bonds are formed to make carbon dioxed and water

This is a combustion reaction: 2H2 + O2 2H2O (it is obvious that bonds have been formed and therefore combustion is exothermic as it releases energy)

∆H = (energy needed to break reactant bonds) – (energy needed to make bonds for products)Where ∆H = enthalpy change for a reaction

In other terms, where ∆H is enthalpy change:

∆ Hrxn = ∑ E (bonds broken) - ∑ E (bonds formed)

Enthalpy

Enthalpy it is the measure of the total energy possessed by a substance or group of substances. We can’t measure the enthalpy itself, but we can measure changes in itEnthalpy change defined as the heat absorbed when the reaction occurs at a constant pressure (basically the heat change)

∆H = enthalpy change Therefore the enthalpy change: ∆H = enthalpy of products – enthalpy of reactants For ENDOTHERMIC reactions, ∆H is POSITIVE For EXOTHERMIC reactions, ∆H is NEGATIVE

Measuring Enthalpy Changes

Calculating molar enthalpy changes from data:1. Calculate the amount of heat released or absorbed by using ∆H = mC∆T2. Calculate the number of moles that reacted from no. of moles = mass/mr3. Calculate the heat relased/absorbed per mole from: total heat rel or abs/no. of moles4. ∆H is plus the heat absorbed or minus the heat released

Energy Profile Diagrams

Activation Energy Ignition Energy

The activation energy is the minimum amount of energy that is required by the reactants for the reaction to proceed to

the products.

The minimum temperature required by

a substance to ignite.

Relationship: The higher the activation energy, the higher the ignition temperature In order for a fuel to ignite, a certain amount of heat energy must be applied to the

substance Not all of the particles must be heated to this temperature to start the reaction. Once the combustion reaction has started, the resulting release of energy provides

sufficient temperatures for neighbouring particles to ignite.

Sources of Pollution

Incomplete combustion can produce carbon monoxide and carbon (CO and C) Sulfur Dioxide/Trioxide

Sulfur Dioxide : Sulfur is an impurity in fossil fuels, mainly coal (up to 5%). When sulfur burns in air sulfur dioxide is produced: S(s) + O2(g) SO2(g) (or sulfur trioxide

through further oxidation of SO2) These two gases can combine with water in the atmosphere to produce acid rain Oxides of Nitrogen : Nitrogen and oxygen in the air can react at high temperatures

that are present inside car engines: N2(g) + O2(g) 2NO (g) causing brown haze

known as photochemical smog and respiratory problems in many people

Controlling Pollution

Use low sulfur coal (Australian coal is relatively low in sulfur) Remove sulfur dioxide from effluent gas at power stations (very expensive) Keep car engines tuned properly so that incomplete combustion is minimised. Catalytic converters mounted to motor vehicle exhaust systems remove unburned

hydrocarbons and oxides of nitrogen.

Complete and Incomplete Combustion

Complete Incomplete All hydrocarbons produce carbon

dioxide (a greenhouse gas) and water. These are the only products when there is sufficient oxygen (i.e. complete combustion).

Complete combustion of methane:

CH4(g) + 2O2(g) CO2(g) + 2H2O(g) + 832kJ

When the oxygen available is insufficient, incomplete combustion results in other pollutants.

As the quantity of oxygen decreases, combustion of hydrocarbons becomes more and more incomplete, leading to the production of the pollutants CO and C

Incomplete Combustion (the red are the products due to the lack of oxygen)

C5H12(l) + 4O2(g) 2CO(g) + 3C(s) + 6H2O (g)

THE RATE OF ENERGY RELEASE IS AFFECTED BY FACTORS SUCH AS TYPES OF REACTANTS

Combustion Rates

Slow Combustion Fast Combustion Explosive Combustion Slow combustion:

like in slow combustion stoves where large lumps of fuel take many hours to burn

Occurs when we use big lumps of fuel and limit the supply of air

This means that burning only occurs on the surface of the big lumps and its speed is controlled by the limited supply of air

Fast combustion: such as burning methane, kerosene or heating oil in stoves or heating appliances

Often in power stations Coal is ground into very

small particles that are sprayed into plentiful air supply – meaning there is a larger surface area and a larger amount of oxygen, therefore speeding up the process as more collisions take place

In other fast combustion processes the gaseous nature of the mixture and high temps ensure that fuel is always in contact with oxygen

Explosive combustion: as in the cylinders of petrol and diesel engines in vehicles

In petrol and diesel engines, the conditions (often involving large amounts of heated air) promote very rapid reaction

An explosion is just an extremely rapid reaction, which completes within a few microseconds

Spontaneous Reactions: Once ignited they proceed without further assistance and continue to go until all of

the fuel is used up All combustion reactions are spontaneous

Collision Theory and Reaction Rates

For a reaction to proceed to products, the reactants must collide with one another. There must also be enough energy present (known as activation energy) and the

particles must collide in a certain relative angle or direction (referred to as steric effect)

The rate of a reaction is the rate that reactants disappear and products form. As the reaction proceeds and reactants are used up, the rate decreases. Since the reacting particles must collide to react, increasing the rate of collisions

increases the rate of reaction. Ways to increase the number of collisions and therefore the rate of the reaction:

Increasing surface area Increasing temperature Increasing concentration Use of catalysts

Concentration Surface area Measures no. of particles of a

substance per unit volume Increasing concentration puts more

particles per unit volume Therefore increases the chance of a

collision and speeds up the reaction

The greater the area of the solid in a solute/solvent reaction, the more collisions that can occur in a given time, so reaction rate increases

Temperature As temperature increases, the average kinetic energy (and therefore the speed) of

particles increases This means that the rate of collision will increase, which will cause an increase in

reaction rate However, this effect alone does not explain the very rapid increase in the rate In order for a reaction to take place, it is necessary for the reactant molecules to

collide AND for the colliding molecules to possess a minimum amount of kinetic energy so that they can reach the top of the energy barrier (activation energy)

If colliding molecules have insufficient energy, they just bounce apart and stay as reactants

If sample is heated, the kinetic energy of molecules is increased, meaning that the molecules have more than enough energy to scale the energy barrier quickly, increasing the rate of reaction

Catalysts

Useful when reactions have high activation energies (and is therefore very slow) Catalyst provides a pathway of lower activation energy The addition of a catalyst does not ultimately change the ∆H (enthalpy change) of the

reaction Industrial Examples:

Catalytic cracking of hydrocarbons in oil refining to convert higher boiling point fractions into petrol – cracking breaks long chains of hydrocarbons into shorter ones and therefore speeds up the reaction

The Haber Process: The production of ammonia from nitrogen and hydrogen gases is catalysed by an iron catalyst.

Catalytic Converter: Rhodium and platinum catalysts are coated on a ceramic honeycomb block to remove unburnt hydrocarbons, nitric oxide and carbon monoxide from motor vehicles.