Chemistry Preliminary Exam Notes

Atchaya senthilkumar

The Chemical Earth The particle theory states that all matter consists of particles, which are constantly moving. Element: pure substances cannot be decomposed into simpler substances Compounds: pure substances that can be decomposed into simpler substances, example: elements Mixtures: Variable in composition, classified as homogenous (uniform in composition) or heterogeneous (non-uniform in composition) Biosphere All living things Mixture: blood

Compounds: sugars Elements: present in living things

Atmosphere Gases of the Earth Mixture: Air Compounds: Water Elements: Nitrogen

Lithosphere Rocks, land and crust of Earth Mixture: Metal Ores Compounds: Quartz (sand) Element: O2

Hydrosphere Waters of the Earth Mixture: Sea water Compounds: CO2 Elements: H2

Mixture separated Method of separation Property used in separation

Solids of different sizes Sieving Particles of different sizes Solids and liquids Filtration Particles of different sizes Dissolved solids in liquids Crystallisation Liquid has a lower boiling point

than solid

Solids and Liquid Sedimentation and Decantation Solid and liquid, liquid slowly poured off leaving solid

Liquids Distillation/Fractional Large/Small difference in boiling points

Liquids Decantation (separating funnel) Immiscible liquids

Gases Fractional Distillation Small difference in boiling points

Gravimetric Analysis Gravimetric analysis is an analytical method used by chemists. It involves separating the components of the material and accurately determining their mass. The percentage composition of the material can then be calculated.

It can be used to determine the:

• Percentage by weight of ingredients (sugar, fat, fibre) in food. • Purity and composition of alloys used for building construction • Extent of heavy metal pollution in river water and human food • Percentage composition of new compounds produced by chemical and medical research.

Combined or Uncombined The higher the reactivity, the less likely it is to exist as an Uncombined element. Less reactive elements include the noble gases (stable outer shell – no reaction).

Ø Highly reactive elements include the alkali metals (Group 1) Ø The less electrons it has in it's outer shell, the more reactive it is

Classification Physical Properties Metals Semi-metals Non-metals Lustre Lustrous Lustrous Dull Malleability Malleable Malleable Non-malleable Electrical and Heat Conductivity

Conductive Non-conductive Non-conductive

Ductility Ductile Ductile Non-ductile Hardness Hard Hard Soft State (25°C) Solid (except Mercury) Solid Gas

Melting and boiling points High High Low

Uses

Metals Non-metals 1. Mercury - Used in thermometers – liquid state

expands when heated 2. Iron – used in building construction and car-

making –high tensile strength and hardness 3. Copper – Good conductors for domestic

appliances

1. Carbon – Good conductors used in batteries 2. Oxygen - for medical purposes to help with

breathing 3. Argon - gas for filling light bulbs

Energy levels for Electrons

Ø Exists in discrete energy levels Ø Energy increases as levels increase Ø Electron configuration: the arrangement of electrons around the nucleus

Formation of Ions Anions: Cation:

Ø Group 1metals tend to lose one electron and therefore positive ions: Li+, ���Na+, K+, Rb+, Cs+.

Ø Group 6 (non-metals) gain two electrons and thus negative ions: ���O2–, S2–, Se2–, Te2–.

Ø Transition metals all lose electrons to form positive ions (Fe2+, Cu2+, Ag+, Zn2+)

Lewis Dot Diagram Ø For molecules/atoms such as HCL, NH3, H2 – all normal dot diagram Ø For Ions such as K+, O2-, F- these need brackets and the dots

Molecules as particles move independently Ø Molecules are particles that may move independently of each other. The forces that hold the molecules

together are known as intermolecular forces. If these forces are weak the molecules may move independently of each other. However if these forces are extremely strong the molecules maybe tightly bound and their movement maybe restricted to just vibrations.

Physical Properties Ionic Bonding (metal with non-m)

Covalent Molecular (b/w non metals) Covalent Network Metallic

Melting Point High because of strong bonds.

Low because of weak bonds.

Very high melting points because of very strong bonds.

High because of strong bonds between atom and delocalised electrons.

Boiling Point High because of strong bonds.

Low because of weak bonds.

Very high boiling points because of very strong bonds.

High because of strong bonds between atom and delocalised electrons.

Electrical Conductivity

Free moving ions in aqueos form allows electrical conductivity. Solid form is non-conductive.

Poor conductors because there are no moving electrons/ions.

Graphite is the only conductor of electricity because of it extra electrons.

Good conductors because of the sea of delocalised electrons moving about.

Malleability

Brittle because slight movement causes the anions to be in line with each other and they repel.

Non-malleable.

Depends on the substance. Eg. Diamond is brittle; others are not.

Malleable due to the ability of ions to slide over each other even when stress is applied.

Lustre Some are lustrous; others are not. Quite dull.

Depends on network. Graphite = dull. Diamond = shiny.

Lustrous because of the delocalised electron's ability to reflect light.

Other Can be dissolved in water

Mainly insoluble and easy to decompose.

Insoluble in water. Conducts heat.

Physical Changes: not lead to the formation of new substances

Ø Filtration, evaporation and distillation Ø do not cause the formation of a new

substance(s) Ø cutting, hammering and rolling Ø change of state Ø are easily reversible ��� Ø require relatively small energy changes

Chemical Changes: leads to the formation of at least one new substance.

Ø Effervescence Ø Odour Ø Heat Ø Colour change Ø Precipitate Ø Large input of energy or output of energy Ø Difficult to reverse

Boiling and electrolysis of water

Boiling Electrolysis 1. Does not produce any new substances,

just a conversion of a liquid to gas 2. Easily reversed by cooling the vapour 3. Requires less energy 4. Does not alter the actual particles, it just

separates them from one another

1. Electrolysis is the process by which an electric current produces a chemical change

2. Produces two new substances, hydrogen and oxygen gases i.e. H2 and O2

3. Difficult to reverse 4. Requires much more energy for the two

gases to be decomposed 5. Breaks the particles up into H2 and O2

Decomposition reactions

Ø Reactions which involve the breakdown of substances into simpler constituent elements Ø Thermal, light and electrolytic

Thermal decomposition of gold oxide Ø When a sample of brown gold (III) oxide is heated over a Bunsen burner flame in a test tube, it readily

decomposes to produce a sample of lustrous gold.

Light decomposition (photolysis) of silver bromide ��� Ø Silver bromide is decomposed by light in the ultraviolet part of the spectrum. The white crystals darken as

black grains of silver metal form. Bromine vapour is released in the process. ���

Electrolytic decomposition of molten lead (II) bromide ��� Ø Lead (II)bromide crystals melt form a clear, colourless liquid. The liquid is heated to 400 °C and

electrolysed using inert electrodes. A brown vapour of bromine is evolved at the positive electrode and silvery globules of lead form at the surface of the negative electrode and sinks to the bottom of the vessel.

Synthesis Reactions: Ø the formation of a compound from its elements or a more complex compound from simpler ���compounds. Ø Rusting of iron, photosynthesis.

Energy Changes Ø The stronger the chemical bonding in a compound, the more energy that is required to break the compound

into atoms. Alternatively, the stronger the chemical bonding in a compound the more energy is released when the compound is formed from its atoms

Empirical Formula is the lowest ratio formula. It does not tell you the number of atoms present For example, NaCl means that for every 1 Na there is 1 Cl (not the exact amount of each element)

¬ The empirical formula is used for ionic compounds because the size of the lattice is unknown therefore a ratio must be used

Another example is glucose C6H12O6 is its molecular formula, whereas its empirical formula (lowest ratio) is CH2O

Metals Metals through history The Copper Age was 3200BC to 2300BC. It is the period that archaeological records indicate that copper was the first metal to be extracted from its ore. Copper was heated with charcoal and globules of copper formed. Molten copper was used to make ornaments and domestic utensils. The Bronze Age was 2300BC to 1200BC. It was later discovered that heating copper with tin produces an alloy, bronze. Bronze was harder than copper and more easily melted to be molded due to its low melting point. Bronze was used for tools and weapons. The Iron Age was 1200BC to 1AD. Iron is more reactive than copper, so it need a higher temperature to melt. Hematite was mixed with charcoal in primitive furnaces by blowing air and obtaining a sufficiently high temperature. By 1000BC, iron had replaced bronze for tools and weapons because it was harder and had hard tensile strength. The Modern Age is 1Ad to present. There had been more extraction and uses of other metals such as aluminium, chromium and metal alloys. Iron is the most widely used metal today. Many other metals have come into common use due to the advancement in extraction technology. Alloys – homogenous mixture of metal with one or more other elements Alloys Properties Uses Brass (50-60% copper with zinc)

Lustrous gold appearance Hard but easily machined

Plumbing fittings Musical instruments Decorations

Bronze (80-90% copper with tin)

Hard Resists corrosion Easily cast

Ships’ propellers Casting statues

Solder (30-60% tin with lead) Low melting point Adheres firmly to other metals when molten

Joining metals together in plumbing and electronics

Steels Mild Steel

Soft, malleable

Car bodies, pipes, nuts and bolts, roofing

Structural Steel

Hard, high tensile strength

Beams and girders, railways, concrete reinforcement

High – carbon Steel

Very hard

Knives and tools such as drill bits, chisels, hammers

Stainless Steel

Hard, resists corrosion, lustrous appearance

Food processing machinery, kitchen sinks and appliances, cutlery, surgical instruments

Energy Input Energy such as electricity and heat is required to extract a metal from its ore in order to break the chemical bonds within the compounds. There is the release and absorption of heat in a form of energy. Absorption:

Ø Extracting copper from sulfide ores liberates heat Ø Extracting aluminium from alumina absorbs heat

However the energy involved in the extraction reaction is only one part of the total energy budget of extraction processes. Energy has to be supplied

Ø To mine the ore Ø Purify or concentrate the ore Ø Maintain high temp needed to make extraction reactions Ø Purify the raw metal or to form it into alloys such as steel

Why more metals are able to extracted now Many metals have been available for use to due lower cost of generating electricity and more advanced in commercial extraction techniques. Two hundred years ago, there was a lack of extraction technology and scarcity of metals and resulted in only a limited amount of metals being able to be extracted and used. Some metal ores have very high melting points and it would have been difficult to reach a very high melting point two hundred years ago with the lack of technology. Reactivity of Metals The Reactivity series is reactivity of metals with oxygen, water and dilute acids in order of decreasing reactivity. It lists in order of decreasing ease of losing electrons: Oxidation K, Na, Li, Ba……… are more easily oxidized than Cu, Ag, Pt, Au It lists in order of increasing ease of gaining electrons: Reduction Ag+, Cu+ ions are easily reduced while Na+, Al3+, Zn2+ are difficult to reduce The most reactive metals react with cold water to form hydroxide ions and hydrogen gas. K, Na, Ca Then highly reactive metals burn in the presence of oxygen. K. Na, Ca, Mg, Al, Zn, Fe The vigorousness with which metals react with dilute acids can be sued to compare. K, Na, Ca, Al, Zn, Fe, Pb As metals atomic size increases,

Ø Atomic radii increases (distance between nucleus and the valence electron increases) Ø Electrostatic forces decreases (attraction between nucleus and the valence electron decreases) Ø Electropositivty increases (ease with which metals are able to lose electrons and become charged ions)

Energy required to remove the electron from the valence shell decreases (easier to remove electrons) Ø Thus, reactivity of metals increases!

Generally…

Ø The most reactive metals are found in Group I, increases down the group because they will react with other elements in order to lose electrons

Ø Moderately reactive metals are in transition elements Ø Least reactive are in lower central region of transition metals

Metals are used for different purposed depending on many factors such as abundance, ease of extraction, hardness, and reactivity.

Ø Magnesium is a highly reactive metal, used in the cathodic protection of less reactive metals to protect them from corrosion. Magnesium is called a sacrificial anode. Also use in fireworks due to its high reactivity and fact that it burns bright when heated with oxygen

Ø Calcium is highly reactive, is added to steels to remove any remaining traces of oxygen, sulfur and phosphorus.

Ø The reactivity of zinc makes it suitable for use in batteries such as dry cells and button batteries. In these cells the zinc is oxidized and the electrons it loses travel through an external circuit producing as electric current.

Ø Aluminium: lack of reactivity means it can be used as drink cans, food wraps, aircraft bodies, in automobiles and in window frames

Ø Titanium: used as artificial joints, aircraft and ship bodies, pipes. Its relevant properties are a low reactivity, stable, resistance to corrosion and chemically inert nature in the human body

Ionisation energy: Ø The first ionisation energy of an element is the energy required to remove the electron from a gaseous atom

of the element. Ø Energy is required in order to overcome the attractive force that binds electrons to the nucleus of an atom. Ø The energy required to remove the loosely bound electron is called the first ionisation energy Ø Ionisation energies reflects the position of electrons in their shells

Ionisation energy increases as you move inwards- as the inner shells are closer the nucleus. The electrostatic attraction increases thus, harder to remove the electrons. Electrons in the outermost shells are further away from the attraction of the nucleus hence, easier to remove electrons. Model of Atom Bohr’s model: The nucleus is central part of the atoms which contains the protons and neutrons. The electrons move through a relatively large space outside the nucleus. The electrons are kept moving around the nucleus by attractive electrostatic forces between positively charged nucleus and negatively charged electrons. Periodic table history In the 1800s, 30 naturally occurring chemical elements were known. French chemist, Antoine Lavoiser classified the elements into two groups, metals and non – metals based on their physical properties. In 1829, a German chemist, Dobereiner recognized the similarities of several groups of three elements in which he called the triads. In 1864, an Englishman, John Newlands, proposed the law of octaves where the elements were ordered according to their atomic weight. In 1869, Mendeleev proposed the periodic law where the properties of the elements vary periodically with their atomic weight. He arranged the elements with increasing atomic weight and grouped them with elements with similar properties. Mendeleev knew that there were still more elements to be discovered and left spaces in his periodic table. In 1914, a British chemist, Henry Moseley, proposed a modified periodic law where the properties of the elements vary periodically with their atomic numbers.

Trends in periodic table … Mole

Ø amount of a substance that contains the same number of particles as there are atoms in exactly 12 g of carbon of Carbon – 12. Chemists have determined that the number of atoms in 12 g of carbon – 12 is 6.02 x 10!"

Percentage composition of a particular element within a substance, use the following formula: In the compound of formula AwByCz: w x (atomic weight of A) x 100 %A= ________________________ molecular weight of AwByCz Calculate percentage composition of iron in common ore, hematite, Fe2O3. Atomic weights are Fe=55.9 and O=10.0 One mole of Fe2O3 is 2x55.9+3x16 = 159.8g One Fe2O3 contains 2 atoms of iron. Hence one mole of Fe2O3 has 2 moles of iron. 2x55.9= 111.8g Fraction of Fe2O3 which is Fe= 111.8/159.8 = 0.700g % iron in Fe2O3 = 70.0% Law of mass of conservation: In a chemical change there is no gain or loss of mass

Ø The total mass of the system remains constant in a chemical reaction. When metals react with oxygen in the air, they generally form metal oxides. The oxygen combines with the metal and adds mass to the original metal.

Gay- Lussac: Law of combining Volume

Ø When measured at constant temperature and pressure, volume of gases exist in simple whole number ratios

Avogadro’s contribution:

Ø Equal number molecules of different gases occupy the same volume at the same temperature and pressure. Ø Every mole of gas will occupy the same volume

Empirical formula

Ø Expresses the ratio in which the atoms are present. It is only the formula for ionic and covalent lattice compounds

Ø NaCl, Na2CO3m MgSO4, K2O

Molecular Formula

Ø Expresses how many of each type of atom are present in the molecule. It is used for molecular compounds only- molecular formula implies the structure of the molecule

Ø H2O, CH4, N2O

Minerals and ores Minerals are naturally occurring inorganic substances, usually compounds with a particular chemical composition and a definite crystal structure. Examples of minerals include hematite, magnetite, gibbsite, boehmite, malachite and chalcopyrite. Ores are naturally occurring deposits that are mixtures of minerals from which a substance, usually a metal can be economically extracted (commercial metal). Examples of ores include bauxite and iron ore. Yield of metal - Is the mass of metal that can be obtained from a particular mass of the mineral or ore. Calculate the theoretical yield of iron from 1000t of iron ore that contains 80% hematite (Fe2O3) Fe2O3(s) + 3CO(g) à 2Fe(s) + 3CO2(g) 80% x 1000t = 800t n(Fe2O3) = 5.009moles From equation Fe2O3 : Fe = 1:2 n(Fe) = 10.019moles Thus, mass(Fe) = 559.5t

Commercial prices The commercial price of metals depends on a few factors including their relative abundances and the cost of production. The greater the abundance of a metal the lower the commercial price of the metal would be. The cost of production of the metals depends on where it is located and the amount of energy input. If the location of the ore is located in a high population zone, the mining procedure would be difficult because there would be damages done to the environment and increase the cost of production. If an ore is located in remote places, then the cost of production would increase because it would cost money to transport the raw materials to refinery plants. The more reactive the metal is, then the higher the energy input is needed for extraction and it would increase the cost of extraction.

Ores: non-renewable resource Ores are deposits of naturally occurring minerals which were formed during the evolution of the universe and the planets; therefore they are non – renewable resources.

Extraction of copper from its ores

1. Mining, crushing and grinding (physical change)

The mined ore (containing ammonium of 6.5% of copper by weight) us placed in a crusher and converted to pebbles. The pebbles are then grounded in a grinding mill to liberate the mineral crystals from the rock

2. Froth Floatation

Using froth flotation, 30% of the copper is obtained by weight.

3. Roasting and smelting

The copper is roasting in the air. 2CuFeS!! +  4O!!     →  Cu!S! +  2FeO! +  3SO!!

The mixture is then heated to a sufficiently high temperature to produce material from which the required metal can be obtained. The mixture of Copper (I) sulfide and Iron oxide with sand is heated to a sufficient high temperature where it produces two immiscible liquids.

FeO!   +  SiO!!  →  FeSiO!! The liquid FeSiO!  is removed. The copper (I) sulfide is then heated on its own to a higher temperature while air is bubbled through it. This reduces sulfide to copper metal and sulfur dioxide is produced.

CuS! +  O!! →  2Cu! +  SO!! 4. Electrolytic refining (99.9% pure)

Anode Cu(s) à Cu2+(aq) + 2e- Cathode Cu2+(aq) + 2e- à Cu(s)

Recycling of Aluminium

Ø Collect the used products from homes, shopping centres and factories Ø Transport the collected material to a central processing plant Ø Separate the required metal from the impurities by chemically cleaning (use acid) Ø Pieces are compressed in block forms to reduce aluminum exposed to oxygen hence minimizing reactivity Ø Re – melt the metal into stock ingots Ø Consider the issue of H2 being dissolved as reacted with Al from the air during the melting phase, hence

reacted with Cl … H2 + Cl2 à 2HCL Ø The aluminium is extremely brittle thus additives are added such as Mg, Cu, Zn to make alloys Ø Let it set into the shapes and molds desired Ø And transport them to product manufacturers

Justify need for increasing recycling in terms of energy

Recycling Processing Ores

1.Collecting scraps (energy of transporting waste products)

1.Finding the ores (energy of transportation)

2.Melting scrap (heat energy used to melt the scrap metal in furnace)

2.Mining ores (energy of mining equipment)

3.Light Purification (almost 0, negligible energy involved in cleaning and purifying with metals with acids)

3.Extracting materials (heat energy from the 5 steps of extraction and purification of ore)

4.Pouring into ‘ingots’ for shaping (negligible energy) 4.Purifying (highest electrical energy consumed by electrolysis)

5.Pouring into ‘ingots’ (negligible energy)

Solving limiting reagents

1. Write balanced equation 2. Convert the quantities of the reactants given into moles 3. Determine the number of moles of reach reactant required using ratios 4. The limiting reagent will be present in insufficient amounts 5. Use number of moles of limiting reagent to determine number of moles of unknown using ratios 6. Convert number of moles to mass, volume or number of molecules required

Hydrogen gas can be produced by reaction of Zinc with HCL. If 2.74g of HCL in solution is added to 3.27g of Zinc, calculate the mass of H2 prodcued

Zn(s) + 2HCL(aq) à ZnCl2(aq) + H2(g)

n(Zn) = 0.050moles

n(HCL) = 0.075moles

Ratio of Zn:HCL = 1:2

Therefore, HCL is present in insufficient amounts and is the limiting reagents

Ratio HCL:H2 = 2:1

n(H2) = 0.0375moles

thus, m(H2) = 0.076g

Water Role of water -Necessity for all living matter in the form of

Ø Raw material that is used in chemical reactions that constitute life

Ø Solvent in which life processes occur Ø Transport medium - Cells and

metabolism: water transports nutrients to and wastes away from cells; metabolic water is produced by respiration; water is needed for photosynthesis

Ø Thermal regulator - Temperature moderator: water can absorb large amounts of heat without causing a significant change in temp, which benefits ectothermic aquatic life

Ø Agent of weathering: Rain and waves physically wear away rocks, freezing water creates ice wedging; minerals are dissolved by groundwater

-Natural resource for humans Ø Drinking, food prep Ø Irrigation Ø Fluid in electricity generating stations and as

coolant Ø Hydro-electricity Ø As reactant, solvent and cleaning agent and

for waste disposal Ø Mode of transport such as ships Ø Recreational purposes such as swimming

Sphere of Earth Physical state % Hydrosphere Solid: polar ice caps, glaciers

Liquid: Oceans, rivers, seas 95-99

Biosphere Liquid: Cells, internal transport systems Gas: Water vapour in cells and organs

60-95

Lithosphere Solid: frozen soil, hydrated minerals Liquid: aquifers. Groundwater

<10

Atmosphere Solid: ice crystals, hail Liquid: water droplets in clouds Gas: Water vapour

0.5-5

Solute is the substance dissolved in a solution. Solvent is the substance, which does the dissolving. Solution is a homogenous mixture in which the dispersed particles are so small (molecules or ions) that they never settle out.

Saturated – no more solute will dissolve in current conditions Unsaturated – more solute is able to be dissolved in current conditions / does not contain maximum amount of solute

Expansion on freezing

Ø Large open spaces between water molecules in ice increases the volume and so decreases the density, since mass remains unchanged allowing to expand.

High temp = larger volume, lower density Low temp = smaller volume, higher density Density = mass Volume

Number of Electron Groups

Electron-Group Geometry

Number of Lone Pairs

Molecular Geometry Examples

2 linear 1

BeH2 CO2

3 trigonal-planar 0

NH3

1

O3 BCl3

4 tetrahedral 0

Tetrahedral

CH4

1

H3O+

2

H2O

Intramolecular: (Within) Chemical bonds Ø Covalent bonds Ø Ionic bonds Ø Metallic

Intermolecular: (between) “physical” force Ø Dispersion / van der vaal forces Ø Hydrogen bonding Ø Dipole-dipole

Types of intermolecular bonding Intermolecular forces exist between molecules, they strengthen the binding force between molecules and contribute to the increase in strength of physical features Dispersion Forces

Ø Weakest Ø Every molecule has dispersion forces Ø Due to very rare “temporary” instabilities in

electrostatic forces

Hydrogen Bonding Ø Strongest Ø Between H and FON (chemistry if FON)-

some of the most electronegative (ability to become negative, gain electrons) atoms

Ø Attraction that arises from the significant disparity between electronegativity of the hydrogen atom and FON

Dipole-Dipole Ø Exists between polar molecules which have a

‘net dipole’ Ø Negative attracts positive end Ø Polar molecular have net dipole which arises

from significant imbalance in electrostatic attraction between the atoms of that molecule

More reactive molecule: δ- (electronegative more) Less reactive molecule: δ+ (electropositive more)

Impact of intermolecular forces Melting and boiling point

Ø Mass crucial factor in strength of intermolecular forces, dispersions forces

∴ size, mass increases, surface area increases, intermolecular forces increases, dispersion forces increase, melting and boiling point increases ∴ water’s MP and BP are considerably higher due to H-bond Surface tension: water’s ability to resist an increase in surface area

Ø Since H2O possesses relatively strong intermolecular forces, because of this, when certain parts on water are exposed and unsupported by a ‘structure’, there will be an imbalance in attractive forces

In the container/water droplet: Ø The stronger attractive forces between molecules near the center (core) of the structure pulls, as the

exposed water molecules are drawn inwards Ø ∴ an imbalance of attractive forces is created, this imbalance forms the “resistant” against an increase in

surface area. Ø A film like structure is formed at the exposed surface

Viscosity: the cause of which molecules can be poured or flowed (measure of liquid’s resistance to flow) Ø Complexity of molecule Ø Size Ø ∴ when complexity increases, size increases, intermolecular forces increases ∴ viscosity increases

When a soluble ionic compound (such as sodium chloride) interacts with water, they break up into positive and negative ions (dissociation). These ions interact with ion-dipole interactions between slightly positive hydrogen atom and negative ion and the slightly negative oxygen atom and the positive ion. When a soluble molecular compound (such as sucrose) interacts with water, the crystals of the solid break up and disperse throughout the solvent (water) and they break down to the molecular level. When a soluble or partially soluble molecular element or compound (such as iodine, oxygen or hydrogen chloride) interacts with water, the solvent-solute interactions are weak dispersion forces and this is why the solubilities of such substances are quite low. When a covalent network structure substance (such as silicon dioxide) interacts with water, nothing happens because water is not able to break the strong covalent bonds between the particles (atoms) in these lattice solids. When a substance with large molecules (such as cellulose or polyethylene) interacts with water, nothing happens because water is not able to break the strong covalent bonds between the particles (molecules) in these solids. However some large molecules such as amylose and glycogen are soluble in water as it contains F, O or N atoms which form hydrogen bonding with the water. When acids dissolve in water they form hydronium ion (H3O+ which is acidic) and hydroxide ion (OH+ which is alkaline) Rule: “Like dissolves Like” Dissolution of salts

Ø when solvent dissolves a solute… the intermolecular forces that hold together solvent (H2O) > intermolecular forces of the solute

Ø ∴ solute dissolves- breaks up into smaller particles as the physical attractive forces of solute have been overcome.

Equilibrium exists between Ø Ions in solid lattices dissolving and Ø Ions in solution precipitating

Forward reaction: precipitation (ions in solution precipitate into solids) Backward reaction: dissolution (solids break away into ions)

Dynamic balance between dissolution and precipitation; both are occurring simultaneously, but at equal rates so there is no overall chance in concentration in solution ----- DYNAMIC EQUILBRIUM GENERALLY SOLUBLE EXCEPTIONS (ions) GENERALLY INSOLUBLE EXCEPTIONS All group I Ammonium NH4

+ Nitrates NO3

-

All Carbonates CO32- All solubles

Sulphates SO42-

Sr, Ba, Pb ions insoluble Ca, Ag slightly soluble

Phosphates PO43- All solubles

Chloride Cl- Ag insoluble Pb slightly soluble

Oxides O2- All solubles , Ba soluble Ca slightly soluble

Bromine Br- Ag insoluble Pb slightly soluble

Hydroxides OH- All solubles , Ba soluble Ca slightly soluble

Iodine I- Ag, Pb insoluble

Concentration measures Mass % = Weight of solute X100 Weight of solution Mass-volume % = Weight of solute x100 Volume of solution Volume-Volume % = Volume of solute X100 Volume of solution ppm (parts per million) = weight of solute (mg) or (mg) weight of solution (L) or (Kg) Specific Heat capacity:

Ø amount of energy required to increase the temperature of 1g of substance by 1OC/Kelvin Ø High MP and BP, stronger H-bonds ∴ high specific heat capacity

Exothermic Endothermic

Energy is released Energy is absorbed (from surroundings)

Decrease in enthalpy of system (-) Increase in enthalpy of system (+)

Increase in surrounding temp Decrease in surrounding temp

Combustion reactions heat + light released Photosynthesis, thermal decomp

Dissolution energy release NaCl à Na+ + Cl-

ΔH<0 1. Breaking of bonds between NaCl (endothermic- absorbs energy) 2. Formation of bonds between ions and water (exothermic- releases energy) 3. Breaking bonds between atoms of water molecules (endothermic- absorbs energy)

∴ if energy released in step 3 is greater than energy absorbed in steps 1 & 2, the dissolution will be exothermic and increase the surrounding temp ∴ if energy released in step 3 is less than energy absorbed in steps 1 & 2, the dissolution will be endothermic and decrease surrounding temp

Importance of water high specific heat capacity Ø Temp remain fairly constant (high specific heat capacity) Ø Land located near large masses of water will experience cooling cycle (high thermal conductivity) Ø Given that most living things are composed of high % of water, it also means their cells and bodies do not

increase in temp too much when exposed to high temps

Thermal Pollution: discharge into river or lake quantities of hot water that are large enough to increase significantly the temp of water body

Ø If increase in temp 3-5 degrees Celsius, then solubility of gases decreases Ø Lower dissolved oxygen concentration by 10-15% Ø Lower concentration of dissolved CO2

o Affecting lifeforms that rely on photosynthesis for energy o If these suffocate and die, it will further affect the food chain

Ø Affects organisms metabolic rates Ø Affects organisms breeding cycles, migration and spawning cycles

Calorimetry: limitations Ø Unavoidable loss of heat energy to surrounding

o Thus, use of Styrofoam cups to trap heat o Thus, stirring so heat is dispersed more evenly

Ø Non-homogenous change in temp o Use better insulation o Stir water o Use digital thermometer

Ø Inaccurate measurements o Use digital thermometer o Use more accurate devices

Energy Role of photosynthesis in transforming light energy Photosynthesis: process that plants use in order to generate glucose for energy Raw materials

o CO2 à Glucose o Light à O2

6CO2 (g) + 6H2O (l) à C6H12O6 (aq) + 6O2 (g) ΔH= +2803kJ/mol (endothermic) Light energy transformed into glucose Role of production of high-energy carbohydrates from CO2 Carbohydrates produced are glucose molecules

o However, chlorophyll from plant leaves are required to convert the light energy into chemical energy (glucose)

o The glucose produced provides energy to cells of photosynthetic organisms as well as those that consume them. The glucose is converted into CO2 and H2O through exothermic process called respiration

o Plants convert excess glucose into polymers such as starch and cellulose o Animals convert excess glucose and store as glycogen

C6H12O6 (aq) + 6O2 (g) à6CO2 (g) + 6H2O (l) ΔH= -2803kJ/mol (exothermic) Light energy -----cholorphyll-------> Chemical Energy Photosynthetic origins of fossil fuels

o When organisms die out they undergo 2 processes in decomposition (by decomposers- bacteria)

1.Complete Decay o Rot, scattered throughout soil o Energy now forms in the state of fertilisation compost o Plants reuse this energy as sustenance for growth and continued photosynthesis o Thus, energy cycle repeats

2.Trasnformation into fossil fuels Ø High temp + high pressure Ø Decayed organisms become compressed and form solid stores of energy (coal, fossil fuels)

Composition of fossil fuels Ø “Hydrocarbons” Ø However many also comprise of impurities such as sulphur, oxygen, nitrogen depending on source of fossil

fuel

Coal (pure-raw source) Ø Mainly impurities “siliceous material) and carbon

Petroleum Ø Mainly hydrocarbons Ø Some impurities

Natural gas o Light hydrocarbons

Allotropes of carbons

o Different physical form/structure of the same elements o Also possess different physical properties

Diamond Ø Infinite crystal lattice, covalent network Ø Hard (covalent network atoms are tightly bonded by covalent forces) , shiny (high refractive index) Ø Non-malleable Ø Thermal conductivity: Excellent

o In covalent network structure the transfer of heat energy is extremely high o Atoms are also closely compressed together o Heat energy readily spreads throughout the structure

Ø Electrical conductivity: Bad o Because no free moving delocalised electrons to transfer the electrical energy throughout the

covalent network Ø BP/MP = high

Graphite o Not hard, quite brittle

o Not compressed and tightly bonded as diamond o Intermolecular forces are weaker as they are not compresses and spaced out

o Dull, lack of shine o Thermal conductivity: Good

o Regardless, vibrations of atoms would allow a similar transfer of heat energy throughout the graphite molecule to occur

o Electrical conductivity: High o Pretence of free moving delocalised electrons between the individual graphite layers o Each carbon is only attached to 3 other carbons thus valence shell is not filled with 4 electrons to

satiate carbon o ∴ one free moving unpaired electron for every carbon present o ∴ free moving carbon electrons move to the “conduction bond” (space between each layer) o ∴ electrical energy/ electricity can be transferred through different layers

Alkanes Alkenes Alkynes Single bonds

o Relatively stable Double bonds

o up, large bond “Pi” bond (stronger)

o down, small bond “Sigma” bond (weaker)

o Difference in strength between molecules create instability

o More reactive (volatile) o Greater tendency to react

with other substances

Triple bonds o More reactive then the

alkenes due to increased instability

Saturated molecules (no free bonds for one carbon atom)

Unsaturated Unsaturated

CnH2n+2

CnH2n

CnH2n-2

“Isomers”

o Sets of different compounds with the same molecular formula but different structural formula (condensed formula)

Fractional distillation Ø Separation based on the substances boiling points Ø Petroleum is heated to 350oC and is passed into base of fractioning column Ø Vapours move up the column, and begin to cool and condense Ø Then they are collected in trays at different positions through the column Ø Since boiling point increases as molecular weight increases, the separation is roughly in order of increasing

molecular weights Ø The least volatile (highest BP and highest molecular weight) condense near the bottom Ø The most volatile do not condense until they reach the top of column

Fraction Boiling Point (ºC) Carbon atoms per molecule

Major uses

Gases Less than 30 1 to 4 Liquefied petroleum gas (LPG)

Petroleum ether 30 – 80 5 to 6 Industrial solvents Gasoline 70 – 200 6 to 12 Motor fuel Kerosene 175 – 250 12 to 16 Jet fuel, domestic heating

Gas oil 250 – 350 15 to 18 Diesel fuel, industrial and domestic heating

Lubricating oil Greater than 350 18 to 25 Motor oils Greases Greater than 350 Greater than 20 Lubrication Asphalt and tar Residue Greater than 25 Road-making, roofing

In terms of molecular structure… Ø Non-polar covalent bonds (C – C is non-polar, C – H is slightly polar but geometry (shape) tends to cancel

it out). Ø Only intermolecular forces are dispersion forces. Ø Dispersion forces increase as molecular weight (no. of C atoms) increases.

These bonds explain the following properties… Melting and Boiling Point

Ø Melting & boiling point increases as number of C atoms increases (as molecular weight increases, dispersion forces increase)

Ø Melting & boiling points of the alkanes are higher than the corresponding alkene (lower molecular weight of alkene à lower dispersion forces)

Solubility Ø Insoluble in water (non-polar), soluble in

non-polar solvents

Volatility Ø Volatility is the ease at which a substance

becomes gaseous or be converted to vapour Ø Volatility decrease, intermolecular forces

increase (Molecular weight increases à dispersion forces increases)

Other

Ø Density < 1 g/cm3 (ie. floats on water) Ø Do not conduct electricity (no free electrons) Ø Generally unreactive (except combustion)

Homologous series: Ø Family of compounds that can be represented by one/same general formula Ø Common functional group Ø Similar structures & chemical properties Ø Gradation in physical properties based on order of molecular weight (eg. boiling point)

Safety precautions: Most of the fuels we employ in everyday life are small fuels

Ø ∴ extremely volatile- may ignite and explode Ø Fuels may be carcinogenic and toxic ∴ when gaseous and breathed in, can cause harm Ø Alkanes are non-polar molecules ∴ dissolve non-polar substances such as plastic (cannot be stored in

plastic containers)

Modes of storage: o Well-maintained cylinders and fittings for gaseous hydrocarbons (high pressure metal containers- keeps

fuel liquefied and mitigates issue of volatility) o Well ventilated to prevent respiratory issues and accumulation of explosive gaseous fuels o Added odours for early detection of leaks o Sturdy containers for liquids o Minimise the quantity in everyday use o Keep alkanes/hydrocarbons away from naked flames or sparks o Erect warning signs o Do not handle these liquids in confined places o Use fume hoods

Indicators of chemical reactions: Ø Gas evolved Ø Precipitate formed Ø Significant temperature rise/ release of

energy

Ø New substance formed Ø Odour produced Ø Light produced Ø Colour change

Combustion reaction Ø Combustion is an exothermic chemical reaction because it releases heat and new products are formed. It is

a self-sustaining chemical reaction that occurs at temp above surroundings. Burning and explosions are forms of combustion.

Process of chemical reaction Bond Breaking (ΔH>0 ∴ endothermic)

o Energy is applied to break the initial molecules and transform them into different/distinct products

o Intramolecular (covalent, ionic) bonds are broken

o Application of energy = the system’s initial molecules are absorbing energy from surrounding environment (enthalpy is +)

o Eg: Cool pack

Bond formation (ΔH<0 ∴ exothermic)

o Energy is released when new substance is formed

o During the formation process, reactants broken up are “destabilised”

o The energy absorbed makes them “excited” ∴ need to become stabilized in order to form new stable products

o New intramolecular bonds are formed during the reaction as energy is released to surrounding environment (enthalpy is -)

o Eg: heat pack The state of the “net reaction” will determine the overall enthalpy of a system

Ø If bond breaking energy > bond formation energy (rxn is exothermic) Ø If bond breaking energy < bond formation energy (rxn is endothermic)

Activation energy o Regardless of whether a reaction is endothermic or exothermic, it will always require an input of energy,

to initiate the reaction o All molecules need to overcome an “energy barrier” in order for the reaction to start o ie: reach state whereby the molecules are excited and are destabilised o ∴ greater frequency of collisions increases and likelihood of collisions increases

Combustion reactions: burning of fuel in presence of oxygen Complete combustion- presence of excess of oxygen Fuel + Oxygen à H20 + CO2 Incomplete combustion- presence of insufficient supply of oxygen (results in pollution) Fuel + Oxygen à H20 + C Fuel + oxygen à CO

Ignition temperature o the minimum amount of energy required to ignite a potion of fuel o A substance with high activation energy ∴ high ignition energy o Combustion reactions are self-sustaining ∴ once ignition occurs, it will trigger combustions (will burn until

all fuel is consumed assuming complete combustion)

Pollution from Combustion CO2

o Greenhouse has that causes global warming o Causes toxic to humans, respiratory issues when consumed at high quantities

CO o Toxic and harmful to respiratory system o CO is similar to O2 o Haemoglobin in the blood system/stream has high affinity to CO2 (attaches easily to CO2) causing blood

poisoning o Acid rain C02(g) + H20(l) à H2CO3(aq)

Internal combustion of an engine: NO2 o Respiratory problems o H20 + NO2 à HNO3 + HNO2 o Acid rain nitric acid + nitrous acid

S02 smelting iron/copper ore o S + O2 à SO2 (burning of sulphur impurities from ore) o O2 + 2SO2 à 2SO3 (burning of SO2 to form SO3) o SO3 + H2O à H2SO4 (formations of sulphuric acid rain)