306975339-igcse-chemistry-notes (1).pdf

IGCSE Chemistry Separating Solid/Liquid Mixtures: Solute – the solid which dissolves in a solvent Solvent – the liquid that the solute dissolves in Solution – formed when a solute dissolves into another solvent: solute + solvent  solution Saturated solution – a solution which contains as much dissolved solute as it can at a particular temperature Soluble – when the solute can dissolve in a solvent Insoluble – when the solute cannot dissolve in a solvent Filtration – the process of separating a solid from liquid using a fine filter paper which does not allow the solid to pass through. The solid (residue) will stay in the filter and the liquid (filtrate) will be in the container under the filter. Decanting – the process of separating a liquid from solid (which has settled) or an immiscible heavier liquid by pouring the solution into another container. The solid or the immiscible heavier liquid will stay at the bottom while the liquid will pour out. Centrifuging – the separation of the components of a mixture by rapid spinning. The denser particles are flung to the bottom of the containing tubes. The liquid can then be decanted off. Evaporation – the separation of a liquid and a dissolved solid by heating the solution. The liquid will evaporate completely leaving the solid behind.

Crystallisation – the process of forming crystals from a liquid. This occurs when a solution is saturated the salt begins to crystallise and can be removed with large scoops. Simple Distillation – the process of boiling a liquid and then condensing the vapour produced back into a liquid. It is used to purify liquids and to separate mixtures of liquids.

Separating Liquid/Liquid Mixtures: Miscible – description of liquids that form a homogeneous layer when two are mixed together Immiscible – description of liquids that form two layers when two are mixed together Separating Funnel – funnel that allows the layers in immiscible liquids to separate Fractional Distillation - process used to separate miscible liquids into liquids that have different boiling points. When the mixture is heated, liquids with low boiling points evaporate and turn to vapour and can then be separated as liquids. Those with high boiling points remain liquids Separating Solid/Solid Mixtures: Sublimation – heating of substances, in which one will sublime (I2 and CO2 will sublime) Magnetism – takes out those things which are attracted to a magnet (Like Fe, Co and Ni). Chromatography – substances dissolved in water (or other solvents) travel along chromatography paper at different speeds. This difference in properties is used to separate some chemicals in analytical laboratories. The substances move at different speeds due to their different solubilities in the solvent.

Locating Agent – used in chromatography to make the spots show when the substance is not visible Rf Values – used to identify which spot is which item Solvent Extraction – when substances are extracted from a mixture by using a solvent which dissolves only those substances required How the purity of a substance can be shown:

Melting Point Boiling Point Chromatography

Pure Sharp Melting Point; usually high Sharp Boiling Point; usually low One well-defined Spot on chromatogram

Impure Range of temperatures Range of temperatures Several spots on chromatogram

Element - the simplest building blocks of the physical world. There are 92 naturally occurring elements. The periodic table is a list of elements in order of atomic number. Atom – the defining structure of an atom which typically includes a nucleus of protons and neutrons with electrons orbiting the nucleus Molecule - Two or more atoms joined together discreetly. (Usually non-metals; can be elements or compounds) Compound - two or more elements (different types) chemically bonded. With compounds you can always give formula. Mixture - two or more substances not chemically bonded Charge + NO CHARGE

Proton Electron Neutron

Mass 1 0 1

Ion - Charged particles that are formed when an atom loses or gains an electron Cat-ion - A positive ion An-ion - A Negative ion Valence Electrons - Electrons on the outer shell of an element Element:

Bohr Diagram Example

Lewis Diagram Example

Fluorine 9P

Electron Configuration:

F

10 N 2,7 Magnesium 12

Mg 24

Atomic Number Mass Number

The Atomic Number is the number of Protons and Electrons unless it is an ion. (Protons = Electrons) The Mass Number is the number of Protons + Number of Neutrons. To work out the number of neutrons you must calculate: Mass Number - Atomic Number = Number of Neutrons Isotope - An element that occupies the same place in periodic table but has a different number of neutrons. The number of protons and electrons are the same. Alloy – a mixture of a metal and another element (usually a metal)

Bonding Metal Metal Non-Metal

+ + +

Electrons Non-Metal Non-Metal

= = =

Metallic Bonding Ionic Bonding Covalent Bonding

Metallic Bonding:     

Metallic Bond is the attraction between the metal ions and the delocalised electrons Too many atoms to count, unlike small molecules like H2O Number of delocalised electrons = number of electrons in out shell of element Metals conduct electricity because the electrons can move Metals are easily shaped (malleable) because cat-ions are in compact layers and can move

+

+ +

+ +

+ +

+

Free (delocalised) electrons, they can move anywhere

+ A

+ +

+ A

+ A Metallic Lattice – the regular arrangement in metal ions in solid metals +

+

+

+

Cat-Ions

+

Ionic Bonding:     

Ionic Bond is the electrostatic attraction between positive and negative ions Only show valence electrons in diagrams. Show transfer of electrons with arrow Write the correct formula Both atoms should end up with full valence shells

+

Na

Cl

Na

-

Cl

Polyatomic ions – ions containing more than one atom. Brackets must be used to write formulae involving more than one of these ions. E.g. Al2(SO4)3; SO4 is a polyatomic ion Covalent Bonding:  Covalent bond is the sharing between atoms to gain full valence shells  Only show valence electrons on diagrams  No arrows  Atoms will be connected

H

H

O

Polar Bonds Non-Polar Bonds – covalent bonds that involve exactly equal sharing of the bonded pair(s) of electrons (e.g. Cl with Cl, O with O) or close enough to equal sharing (e.g. C with H) that the shared pair of electrons is equidistant between the two atoms and thus the electronic charge is evenly balanced around all atoms Polar Bonds – covalent bonds involve uneven sharing of the electron pair or pairs, with one of the atoms (e.g. F, O or Cl) having a slightly stronger attraction for the shared pair of electrons in the bond than the other atom (e.g. C or H). As a result the covalent bonds are closer to that atom with stronger attraction. This gives that atom a slightly negative charge (-) and the other atom a slightly positive charge (+) Polar Molecules – molecules with at least one polar bond and an asymmetrical shape so the dipoles do not cancel  

   

Oxygen atom has stronger attractions and hydrogen has weaker attractions. Covalently bonded electrons are closer to oxygen making it slightly negative and the hydrogen atoms slightly positive creating dipoles Molecule in asymmetrical shape therefore dipoles do not cancel Two polar bonds Water is Polar

+

+

H -

H

O

-

Non-Polar Molecules – molecules where all bonds are non-polar or with a symmetrical shape causing the dipoles to cancel out  

 



Carbon and Hydrogen have the same attraction Covalently bonded electrons are shared almost equally creating no dipoles All bonds are non-polar Molecule is in symmetrical shape so if there were polar bonds they would cancel out Methane is Non-Polar

H

H

C H

H

Shapes of Molecules Molecule Example CO2 H2O BF3 NH3 CH4

Shape Linear Bent/V-Shaped Triangular Planar Pyramidal Tetrahedral

Bond Angle O-C-O = 180° H-O-H = 105° F-B-F = 120° H-N-H = 107° H-C-H = 109.5°

Radioactivity – when the nucleus in a radioactive atom is unstable and so it emits particles/waves to form a more stable atom Radioisotopes – a radioactive isotope; one having an unstable nucleus and emitting radiation during its decay to a stable form Radiation – the particles/waves emitted by radioactive substances Radioisotopes and their Uses Radioisotopes Carbon – 14

Cobalt - 60 Krypton – 81

Cesium - 137

Uses Carbon Dating; When an organism dies it stops taking in new carbon atoms so the amount of carbon-14 slowly drops as the atoms decay. By measuring the radiation from the carbon-14 atoms the age of the remains can be determined. Cancer Treatment; weak beams of radiation will kill cancer cells more easily than healthy cells. Cobalt-60 is also used to kill germs and bacteria. Tracers; a small amount of krypton is breathed in, it decays in the lungs and the radiation can be detected and viewed on a screen where it shows up as bright spots. Dark patches show where the lungs are not working properly. Kill Germs and Bacteria; a low dose of gamma radiation is used to kill bacteria in food that causes it to decay. They also kill germs and bacteria on surgical equipment.

States of Matter Solid

Liquid

Gas

Fixed shape and volume Particles are held together by relatively strong forces Incompressible Particles do not have free movement but can vibrate around fixed positions

No fixed shape, fixed volume Particles have weaker forces so are further apart Slight compressibility Particles can move throughout bulk of liquid

No fixed shape or volume Particles are very far apart Compressible Particles are very spread out and move in random fashion

Deposition Solidification

Condensation

Melting

Evaporation Sublimation

Matter – all the substances and materials from which the physical universe is composed Kinetic Particle Theory – a theory which accounts for the bulk properties of matter in terms of the constituent particles. It states that:   

All matter is made up of tiny, moving particles, invisible to the naked eye. Different substances have different types of particles (atoms, molecules or ions). The particles move all the time. The higher the temperature, the faster they move and the forces of attraction weaken Heavier particles move more slowly than lighter particles at a given temperature

Diffusion – the process by which different substances mix as a result of the random motions of their particles. Particles with smaller Mr diffuses faster. Intimate mixing – when diffusion takes place between a liquid and a gas Brownian Motion – random motion of particles caused by smaller and faster moving water particles constantly colliding with them and moving them around Gas Vaporising Liquid Melting

When the object is melting or vaporising, heat energy is being added but temperature is not changing. The average kinetic energy stays the same. Energy changes to potential energy by separating.

Solid Absolute Zero – the theoretical temperature (which can never be reached) at which all particle motion stops. Absolute zero is -273°C or 0 Kelvin. To calculate Kelvin = °C + 273. A 1K change equals a 1°C change. Pressure of a Gas  The free moving particles of a gas will spread evenly within a container and collide with the walls. This will exert a force on the wall when it bounces off.  When this happens on a large scale (billions of particles) there is an average force exerted on the wall. This creates a pressure due to Pressure = Vapour Pressure – particles that gain enough energy to become gaseous at the top of a liquid

Ionic Compounds  

Solid ionic compounds have no moving charged particles, they do not conduct electricity. Liquid and aqueous ionic compounds have free moving charged particles (ions) in solution which can carry charge under the influence of an electric filed.

Bonding

Properties

Ionic Bond - Results from the attraction between positive and negative ions - Occurs when a metal reacts with a non-metal - Does form ions so it will conduct electricity as liquid or aqueous - Forms a crystal lattice structure - Has a high melting point

Covalent Bond - Formed by the sharing of electrons between atoms - Occurs when two non-metals react - Does not form ions so does not conduct electricity - Forms a shared electron structure - Has a low melting and boiling point

Intermolecular – these are attractions between molecules when they are close together and are broken when substances melt or boil. Intramolecular – this refers to the covalent bonding. These bonds are only broken during a chemical reactions and never when melting or boiling. There are three varying degrees of strength of these bonds that depend on the type of molecules: 1. Temporary dipole attractions – the weakest attraction between non-polar molecules 2. Permanent dipole attractions – the next strongest attraction between polar molecules 3. Hydrogen bonding – strongest attraction between polar molecules and extremely reactive nonmetals (like F, O, or N) Physical Changes – when the appearance/form of the substance changes but the actual identity and characteristics of the substance remain the same Chemical Change/Reaction – when substances chemically combine and alter one another forming new substances with different properties and characteristics. Conductors  

To conduct, charged particles must be present and these charged particles must be free to move. There are two types of conductors: o Elements which conduct in both solid and liquid because their outer shell electrons are mobile e.g. metals o Electrolytes conduct because they contain positive and negative ions. In electrolytes, the mobile ions carry the current under the influence of an electric field, and the electrolyte is decomposed/discharged as the ions gain or lose electrons at the electrodes e.g. Sodium Chloride solution

Allotropes – different forms of the same element. Allotropy – when an element can exist in more than one physical form in the same state

Graphite

Silicon Dioxide (Silica)

Each carbon atom has four covalent bonds with other carbon atoms

Each silicon atom has four covalent bonds with oxygen atoms.

Arrangement

Carbon atoms link together to form a giant lattice structure

Each carbon atom has three covalent bonds with other carbon atoms. Van der Waal’s forces hold the layers together. Arranged in hexagons and are arranged in layers on top of each other. Electrons move throughout layers

Prop.

Does not conduct electricity, high melting point, insoluble

Conducts Electricity, high melting point, insoluble

Does not conduct electricity, high melting point, insoluble

A hard, colourless, transparent crystal which sparkles in light

A soft dark grey, shiny solid with a slippery feel

A hard, colourless, transparent crystal which sparkles in light (Quartz)

Use

Bonding

Structure

Diamond

Look

Giant covalent structures – structures with a network of covalent bonds throughout it. They take a lot of energy to break and have high boiling and melting points.

Atoms link together to form giant lattice structure

Jewellery, Glass Cutters, Pencils, Electrodes, Polishers Lubricants Calculating Ar (Relative Atomic Mass/RAM):

Cement

 The Ar of an element is defined as the average mass of its isotopes compared with onetwelth of the mass of one atom of Carbon 12  Carbon 12 has a mass of 12 therefore, one-twelth of the mass of one atom of Carbon 12 is 1.  To calculate Ar you use the equation: Ar =

.

Average mass of isotopes of the elements = Calculating Mr (Relative Formula Mass/Relative Molecular Mass): To calculate the Mr of a compound, you add the mass number of each of the elements e.g. NaCl = 23 + 35 = 58 or MgBr2 = 24 + 80 + 80 = 184

The Mole Concept Avogadro’s Constant – equal volumes of all gases measured under the same conditions of temperature and pressure contain equal numbers of molecules Moles – the amount of a substance which contains 6x1023 atoms, ions or molecules. This number is called the Avogadro’s constant. One mole of atoms has a mass equal to the relative atomic mass (Ar) in grams. One mole of molecules has a mass equal to the relative molecular mass (Mr) in grams.

The calculation of no. of moles is moles =

m

The calculation of Molar Mass is Molar Mass =

n

The calculation of mass is mass = number of moles x Molar Mass

The calculation of concentration is concentration =

M n

The calculation of Volume is Volume =

c

The calculation of no. of moles is moles = concentration x Volume

V V

The calculation of no. of moles is moles = One mole of any gas at rtp = 24

n

The calculation of Volume = Volume x 24

24

Empirical Formula     

Percentage composition by mass gives a ratio by mass of the elements contained in a compound This ratio by mass can be converted to a ratio by moles if the % figures (or the known mass) are divided by the respective relative atomic masses Ratios of moles will seldom be whole number ratios To calculate the ratio of numbers or atoms in the empirical formula divide all the mole ratios by the smallest calculated mole value The amount of water associated with a particular salt (water of crystallisation) can be calculated using this method as well.

Step 1 - % By Mass (or Mass) 2 – Mole Ratio 3 – Ratio by atoms 4- Empirical Formula

Calcium 20% 20/40 = 0.5 0.5/0.5 = 1 Ca = 1

Bromine 80% 80/80 = 1 1/0.5 = 2 Br = 2

Percentage Yield – the percentage of the reactants that are converted to products. In some reactions this will be 100% but others it is much less than 100%. Theoretical Yield – the amount of a substance that should be produced through a chemical reaction Actual Yield – the amount of a substance that is actually produced through a chemical reaction Percentage Yield =

e.g.

(30.7g) CaCO3  (11.7g) CO2 + CaO; 100 = Mr of CaCO3, 44 = Mr of CO2 n(CaCO3) =

= 0.307 moles

44 x 0.307 = 13.508 = 86.6% Types of Reactions Single Displacement - one element displaced Most Metals + Water  Metal Hydroxide + Hydrogen Gas Most Metals + Acid  Metal Salt + Hydrogen Gas e.g. Cu + HCl  CuCl2 + 2H A + BC  A + B Double Displacement - two elements displaced Acid + Base  Salt + Water e.g. 2KI + Pb(NO3)2  PbI2 + 2KNO3 AB + CD  CB + AD Synthesis - two elements combine to form one compound Most Metals + Oxygen Gas  Metal Oxides e.g. 2Mg + O2  2MgO A + B  AB Decomposition - one compound becomes two (breaks apart/decompose) Most Metal Hydroxides  Metal Oxide + Water Liquid Most Metal Nitrates  Metal Oxide + Nitrogen Dioxide Gas + Oxygen Gas Group One Metal Nitrates  Group 1 Metal Nitrate + Oxygen Gas Most Metal Carbonates  Metal Oxide + Carbon Dioxide Gas e.g. CuCO3  CuO + CO2 AB  A + B Combustion – oxygen combines with another compound to form water and CO2 Organic Molecule + excess Oxygen Gas  Carbon Dioxide + Water e.g. C10H8 + 12 O2  10 CO2 + 4 H2O AB + C  AC + BC

Redox Reactions Oxidation – the gain of oxygen; the loss of hydrogen; the loss of electrons; increase in the oxidation number Reduction – the loss oxygen; the gain of hydrogen; the gain of electrons; decrease in the oxidation number Oxidising Agents (Oxidants) – the molecule that is reduced (lost oxygen, gained hydrogen etc.) Reducing Agents (Reduced) – the molecule that is oxidised (gained oxygen, lost hydrogen etc.) Oxidation Number – numerical bookkeeping system to help keep track of electron movements between substances. The oxidation number of ions is there electric charge. Elements have an oxidation number of 0. Photochemical Reactions Photosynthesis – the conversion of water and carbon dioxide into glucose and oxygen for energy. The glucose produced is used to make sugars and starch as carbohydrates. Carbon Dioxide + Water  Glucose + Oxygen Gas 6 CO2 + 6 H2O  C6H12O6 + 6 O2 Respiration – reverse of photosynthesis; combustion reaction Glucose + Oxygen Gas  Carbon Dioxide + Water C6H12O6 + 6 O2  6 CO2 + 6 H2O Photography – this photochemical reaction is used as the basis of black and white photography. The photographic film is made of flexible plastic, coated with a layer of gelatine with millions of particles of silver bromide spread through it which is changed to silver when light falls on the exposed parts of the film. Silver Bromide  Silver + Bromide 2AgBr  2Ag + Br2 Hydrogen as a fuel:   

 

Hydrogen is considered to be the fuel of the future and is being trialled by motor manufacturers as an alternative to fossil fuels such as petrol. Hydrogen is non toxic, produces more energy per gram than any other fuel and burns cleanly to form water so there is no exhaust pollution. It has a lower flammability than fossil fuels. Hydrogen is obtained from the electrolysis of water which is plentiful. So far this is not a cost efficient alternative to using fossil fuels and non-renewable if produced using fossil fuels or nuclear energy. Hydrogen is difficult to transport and store. Because it is too light to liquefy easily, a large fuel tank would be needed. It is explosive in correct proportions with air.

Energy Exothermic Reactions – when energy is lost to the surroundings during a reaction. Energy Profile for an Exothermic Reaction – the graph shows the variation in energy during the course of a chemical reaction where heat is released. Energy Level

Ea

Reactants Energy Lost (ΔH)

Products

Reaction Progress Endothermic Reactions – when energy is absorbed by the products from the surroundings during a reaction. Energy Level

Ea

Reactants

Products Energy Absorbed (ΔH)

Reaction Progress Activation Energy (Ea) – the initial energy that is required for a reaction to begin Calorimeter – determines the amount of heat generated in a chemical reaction by the rise in temperature of the reaction chamber and the water jacket around the reaction vessel. Bond Energy – the amount of energy needed or released to break or form a bond. Bond breaking is endothermic. Bond forming is exothermic. ΔH = Total bond energy of all bonds broken – Total bond energy of all bonds formed Equilibrium Reactions Irreversible Reactions – reactions that has products that cannot turn back into their reactants. Reversible Reactions – reactions that has products that can react back into the original reactants. Dynamic Equilibrium – when there is no overall change in the amount of products and reactants even though the reaction is ongoing. Dynamic Equilibrium can only take place in a closed system. The position of dynamic equilibrium is not always at a half-way point, as in it may be at a position where there are more products than reactants. Le Chatelier’s principle – if a closed system at equilibrium is subject to a change then the system will adjust in such a way as to minimise the effect of the change.

Factors affecting Equilibrium Factor Temperature

Concentration

Pressure

Catalyst

Increase of Factor Equilibrium shifts to decrease the temperature so it shifts to the endothermic direction Equilibrium shifts to decrease the concentration

Decrease of Factor Equilibrium shifts to increase the temperature so it shifts to the exothermic direction Equilibrium shifts to increase the concentration

Equilibrium shifts to decrease the pressure so it shifts in the direction of the least molecules Speeds up the time it takes to reach equilibrium but does not change the position

Equilibrium shifts to increase the pressure so it shifts in the direction of the most molecules -

Haber Ammonia Process Haber Process – the process by which ammonia is made from nitrogen and hydrogen. Nitrogen is obtained from air and hydrogen is obtained from methane. It follows the following equation: N2 + 3 H2  2 NH3 



ΔH = -92 kJ/mol

Increasing the temperature will produce less ammonia because this will use up the added heat. Lowering the temperature will produce a greater yield of ammonia but will decrease the rate of the overall reaction. Increasing pressure should move the equilibrium to the right to produce more ammonia. However this will increase the cost because of the thickness of the walls of the plant needed to contain the reaction and it means the temperature will increase and its disadvantages.

Conditions of the Haber Process    

Pressure of 200 atm and Temperature between 380 and 450 °C Ground Iron catalyst to increase the rate of reaching equilibrium at the lower temperature The equilibrium mixture is cooled, allowing ammonia to liquefy and be removed. Unused Nitrogen and Hydrogen is continuously recycled back into the system.

Rates of Reaction Rates of reactions can be measured by the:    

Time for a solid to dissolve or form Loss in Mass (gas given off) over time Volume of Gas collected per time Time for a colour to appear or disappear

Collision Theory - in order for a chemical reaction to occur the particles must collide with each other and have sufficient energy to react. The rate of reaction depends on the number of successful collisions there are in a given time. When particles move faster, they have more kinetic energy. Rates of reaction are affected by: 1. Concentration - adds more particles so they can collide with each other. At the beginning of the reaction, the concentration is at its highest, so the initial rate is the fastest as there are a large number of reactant particles per unit volume and more collisions will occur. As the reactants decrease in concentration, there will be less collisions and the rate slows 2. Particle Size/Surface Area larger area means there is more room for the particles to roam and will collide easier. The amount of product remains the same but the surface area is different 3. Gas Pressure – an increase in pressure forces the particles to come closer together and increases the chance of successful collisions. 4. Temperature – a higher temperature gives particles more energy for collisions and makes the particles move faster so they are more likely to collide. 5. Catalysts – provides particles an alternate way of reacting if the activation energy of the particle is too low, without itself being consumed. E.g. Iron in Haber process, Enzymes in human body

Acids and Bases Acid – a substance that acts as a donor of hydrogen ions Base – a substance that acts as an acceptor of hydrogen ions Alkali – soluble bases Acid Sour Taste pH less than 7 In solution, contains hydronium ions (H3O+) Turns blue litmus red Turns phenolphthalein colourless Corrosive Reacts with metals to produce salt and hydrogen Examples of Acids Hydrochloric Acid HCl Nitric Acid HNO3 Sulphuric Acid H2SO4 Ethanoic Acid CH3COOH

Base Bitter Taste pH greater than 7 In solution, contains hydroxide ions (OH-) Turns red litmus blue Turns colourless phenolphthalein pink Soapy feel Cannot react with metals Examples of Bases Sodium Hydroxide NaOH Potassium Hydroxide KOH Calcium Hydroxide Ca(OH)2 Ammonia Solution NH3 (aq)

Hydronium Ion – same as a single proton because when a hydrogen atom loses an electron, only a proton remains. H+ is irresistibly attractive to water molecules and therefore it would form H3O+. Dissociation – breaking apart Strong Acids – in aqueous solutions, strong acids donate all their protons to water molecules. Weak Acids – there is only a slight tendency to donate protons to water molecules, therefore an aqueous solution of a weak acid contains mainly undissociated molecules and a low concentration of H3O+.

Dissociation in Aqueous Solution Equilibrium Electrolyte Electrical conductivity [H3O+] pH value Examples

Strong Acids Completely dissociate None (forward only) Good Good Higher Lower HCl, HNO3, H2SO4

Weak Acids Partially dissociate Equilibrium reaction Poor Poor Lower Higher CH3COOH, NH4+

Amphiprotic – substances can act as both an acid and a base e.g. H2O, HCO3-, HSO4Amphoteric – substances will undergo chemical reactions with both acids and bases Neutralisation – an alkali or base can neutralise an acid by removing the H+ ions and converting them to water. Neutralisation always produces a salt.

Concentration – a measure of the amount of acid per dm3, refers to the proportion or ratio of acid to water in the solution Concentrated Acids – high proportion of acid to water Dilute Acids – low proportion of acid to water Monoprotic – having one transferrable proton Diprotic – having two transferrable protons Titration – an indicator shows when the acid properties are just destroyed by the alkali. The salt can then be recovered by evaporating the water away allowing the salt to crystallise. This method is used when the base, acid and salt are all soluble.

Oxides     

Oxides of metals are bases (they will react with acids to form salts) Oxides of non-metals are acids (they will react with acids and bases) Some metal oxides are amphoteric (they will react with acids and bases) Some non-metal oxides are neutral Oxide ions immediately react with water and then dissolve to form hydroxide ions. Although potassium hydroxide solution exists, potassium oxide solution does not exist

Metal Oxides – compounds of metal cations and the oxide anion O2-. Few metal oxides react or dissolve in water. The main metal oxides which are considered soluble are potassium and sodium oxides, as well as barium, calcium, and magnesium oxides in decreasing amounts. Metal oxides are either basic or amphoteric. The basic oxides will only react with acids, while the amphoteric oxides will react with both acids and bases. Non-metal Oxides – covalently bonded compounds of a non-metal with oxygen. They are either acidic or neutral oxides. The acidic oxides react with water immediately and dissociate to form acid solutions while the neutral oxides do nothing when placed in water. The acidic oxides will react only with bases, while the neutral oxides are unreactive with both acids and bases.

Solubility of Ionic Compounds in Water Precipitate – a solid formed in a solution. Sparingly Soluble (SpSol) – materials have very low solubilities Hydrolyse (Hyd) – reacts with water Always Soluble All NO3All NH4+ Group 1

Usually Soluble All SO42- EXCEPT Ba, Pb, Ag, Ca All Cl- EXCEPT Ag, Pb All I- EXCEPT Ag, Pb All Br- EXCEPT Ag, Pb

Usually Insoluble All CO32- EXCEPT Group 1 All O2- EXCEPT Group 1 and 2 All OH- EXCEPT Group 1; Ca and Ba are slightly soluble

NH4+ Na+ K+ Al2+ Zn2+ Ca2+ Cu2+ Ag+ Fe2+ Fe3+

NO3-

Cl-

Br-

SO42-

CO32-

OH-

O2-

I-

White Soluble White Soluble White Soluble White Soluble White Soluble White Soluble Blue Soluble White Soluble Green Soluble Violet Soluble

White Soluble White Soluble White Soluble White Soluble White Soluble White Soluble Blue Soluble White Insoluble Yellow Soluble Brown Hyd

White Soluble White Soluble White Soluble White Soluble White Soluble White Soluble Black Soluble White Insoluble Green Soluble Red Hyd

White Soluble White Soluble White Soluble White Soluble White Soluble White Insoluble Blue Soluble White Insoluble Green Soluble Yellow Soluble

White Soluble White Soluble White Soluble -

-

-

White Soluble White Soluble White Insoluble White Insoluble White SpSol Blue Insoluble -

White Hyd White Hyd White Insoluble White Insoluble White Hyd Black Insoluble Brown Insoluble Black Insoluble Red Insoluble

White Soluble White Soluble White Soluble White Hyd White Soluble White Soluble -

White Insoluble White Insoluble Green Insoluble White Insoluble Grey Insoluble -

Green Insoluble Brown Insoluble

Yellow Insoluble Grey Soluble -

Identification of Cations NH4+ 2+

Cu Fe2+ Fe3+ Al3+ Ca2+ Zn2+

Add a few drops of NaOH No precipitate formed Blue precipitate formed Green precipitate formed Orange/Brown precipitate White precipitate White precipitate White precipitate

Add excess NaOH When warm = litmus blue Precipitate dissolves Precipitate remains Precipitate dissolves

Add Ammonia to fresh sample Insoluble white precipitate Soluble white precipitate

Identification of Anions Test with red litmus CO3SO42ClINO3-

Litmus turns blue No change No change No change No change

Add HNO3 and Ba(NO3)2 Precipitate forms No precipitate No precipitate No precipitate

Add dilute HNO3 and AgNO3 to fresh sample CO2 released (HNO3 only) White precipitate Yellow precipitate No precipitate

Add Al and NaOH to fresh sample; warm NH3 produced

Identification of Gases Ammonia Carbon Dioxide Chlorine Hydrogen Oxygen

Turns damp litmus blue; forms white smoke when in contact with HCl fumes Turns limewater milky Bleaches damp litmus paper ‘Pops’ with lighted splint Relights a glowing splint

Physical Properties of Metals, Non-Metals and Metalloids Metals Lustre (shiny) Good conductor of heat Good conductor of electricity Malleable Ductile High Density High Melting Point

Non-Metals No lustre Poor conductor of heat Poor conductor of electricity Not Malleable Not Ductile Low Density Low Melting Point

Metalloids Can be shiny or dull Fair conductor of heat Fair conductor of electricity Malleable Ductile Solids

Chemical Properties of Metals and Non-Metals Metals Easily loses electrons Oxides generally basic and amphoteric Corrodes easily

Non-Metals Tends to gain electrons Oxides generally neutral

Alkali Metals      

Group One Metals Very low density and therefore floats on water. The densities increase down the group. Silvery and shiny when freshly cut, however they quickly tarnish Low melting point Low boiling point The reactivity increases down the group. Since the valence electron is further from the nucleus, the attractive force holding it is weaker and therefore other stronger forces can easily remove it. Transition Metals

Physical Properties (compared to Group 1) Much harder Higher tensile strength Higher density Higher melting point and boiling point Many of their compounds are coloured

Chemical Properties (compared to Group 1) Much less reactive Many have excellent corrosion resistance Show more than one valency (e.g. Fe2+ or Fe3+) Them and their compounds are useful catalysts Some are strongly magnetic Alloys

Alloy Solder Brass Bronze Mild Steel Hard Steel Stainless Steel Alnico

Mixture 70% Tin, 30% Lead 60-95% Copper, 5-40% Zinc 90% Copper, 10% Tin 99.5% Iron, 0.5% Carbon 99% Iron, 1% Carbon 74% Iron, 18% Chromium, 8% Nickel Iron/Aluminium/Nickel/Cobalt

Use Joining wires and pipes Taps, hose/pipe fittings, zips, screws Ornaments, bells, bearings General structural purposes, cars Blades Corrosion resistance Permanent Magnets

Metals Uses Metal Aluminium

Zinc Iron

Copper

Property Does not corrode Low density, unreactive Low density, strong, conducts Low density, strong, cheap Low density, conducts heat Reactive More reactive than iron Similar expansivity Strong, cheap Strong and abundant Good conductor of electricity Unreactive, workable Unreactive Unreactive

Uses Food containers Containers and packaging buildings Long distance wiring Transport vehicles Car Engines Dry cells (“batteries”) Galvanising Iron Reinforcing concrete Nails Ship building Electrical wiring Alloys – brass and bronze Coinage (with Nickel) Hot water piping

Reactivity Series Most Reactive

Least Reactive

K Na Ca Mg Al C Zn Fe Sn Pb H Cu Ag Au Pt

Potassium Sodium Calcium Magnesium Aluminium Carbon Zinc Iron Tin Lead Hydrogen Copper Silver Gold Platinum

Any metal higher on the reactivity series will displace another lower metal’s ions from solution. e.g. Ca (s) + Cu2+ (aq)  Ca2+ (aq) + Cu (s) BUT Cu (s) + Ca2+ (aq)  No Reaction The more reactive metals are difficult to extract from their ores in compound form as they are stable. The less reactive metals have the greater tendency to form atoms and therefore their compounds are less stable. Corrosion

Corrosion – when metals react with water and oxygen. The metal ions lose electrons to form ions. Rusting – the corrosion of iron metal to form a red-brown compound (hydrated iron (III) oxide) Covering with Protective Coat Painting Greasing of metal parts Oiling of bike chains Tin Plating – In cans Plastic covering on electric wires Galvanising – zinc coating for galvanised steel Chromium plating of car parts

Preventing Oxidation of Metal Galvanising – zinc atoms react before the iron Sacrificial protection – a more reactive metal reacts before the metal that it is protecting Carthodic protection – an electric power source pushes electrons into the metal to prevent the loss of electrons

Reduction of Metal Oxides By Hydrogen: only the metals below Hydrogen in the reactivity series are reduced by using this method (mainly only CuO) CuO (s) + H2 (g)  Cu (s) + H2O (l) By Carbon: only the metals below Carbon in the reactivity series are reduced by using this method 2 PbO (s) + C (s)  2 Pb (l) + CO2 (g) By Carbon Monoxide: only metals below Carbon in the reactivity series are reduced by using this method CuO (s) + CO (g)  Cu(s) + CO2 (g) Blast Furnace    

Iron is extracted from Haematite or Ironsand in a Blast Furnace A charge is a mixture of limestone, coke (carbon) and iron oxide (as well as it’s impurities, mainly consisting of SiO2) The charge is placed in the top of the blast furnace and hot air is blasted through at the bottom, making the charge glow white hot. The following reactions take place:

C (s) + O2 (g)  CO2 (g)

CaCO3 (s)  CO2 (g) + CaO (s)

CO2 (g) + C (s)  2 CO (g) 3 CO (g) + Fe2O3 (s)  2 Fe (l) + 3 CO2 (g) If the iron from the blast furnace solidifies, it is called cast iron and is mostly turned into steel. Steel is manufactured the following way: 





Unwanted impurities are removed in an oxygen furnace where the molten metal is poured into a furnace along with some scrap iron (to recycle it). Calcium oxide is added and a jet of oxygen is blasted into it. The calcium oxide reacts with the impurities forming slag that can be skimmed off. Oxygen reacts with the excess carbon, burning most of it away as CO2, leaving some to mix with the iron to make the metal hard but not brittle. Other elements are then added to gain the desired steel properties.

CaO (s) + SiO2 (s)  CaSiO3 (l) This is known as slag.

Zinc from Zincblende 1. The ore zincblende (made mostly from Zinc Sulphide) is crushed and put into water through which air is blown. Rock particles sink and the zinc sulphide floats in a froth which is skimmed off and dried. The product of this stage is 55-75% Zinc Sulphide. 2. The Zinc Sulphide is converted to Zinc Oxide by strong heating in a furnace: ZnS + 3 O2  2 ZnO + 2 SO2 3. Zinc Oxide is mixed with coke in a furnace and heated to 1400 °C where it is reduced to zinc: ZnO + C  Zn + CO 4. The zinc metal produced cools and the carbon monoxide is burnt, with the heat given out to help reduce costs of the furnace.

Heating Metal Compounds

Potassium Sodium Calcium Magnesium Aluminium Carbon Zinc Iron Lead Hydrogen Copper

Hydroxide Stable No Reaction

Nitrate Decomposes 2NaNO3  2NaNO2 + O2

Carbonate Stable No Reaction

Decomposes Cu(OH)2  CuO + H2O

Decomposes 2Ca(NO3)2  2CaO + 4NO2 + O2

Decomposes MgCO3  MgO + CO2

Halogens    

    

Group VII elements are known as Halogens They are non-metals They are poisonous Melting point and Boiling point will increase as it goes down the group because the size increases, meaning an increase in the strength of the Van der Wall forces holding them together, causing a higher temperature to be needed to break them Colour goes darker as it goes down the group Less reactive as it goes down the group because the bigger the atom, the smaller attraction between the nucleus and incoming electron All have similar properties because they all have seven electrons in the outer shell Reacts with metals to form ionic compounds, containing halide ions A more reactive halogen will displace a less reactive one from solution

Sulfur Sulfur – mined from solid underground deposits of elemental sulphur, extracted from fossil fuels, received from metal sulphide ores when the metal is extracted. It is used in the manufacture of sulfuric acid. Sulfur Dioxide – is prepared from when sulphur burns in air or oxygen (burns with a blue flame): S(s) + O2(g)  SO2(g) 

Sulfur dioxide dissolves in water to form sulfurous acid (a weak acid) which can lead to the problem of acid rain H2O(l) + SO2(g)  H2SO3(aq)

   

Sulfur dioxide can cause bronchiospasm in asthmatics It is used as a bleaching agent when paper is made from wood pulp It is used as a preservative for food by killing bacteria. Sulphites and hydrogen sulphites are also used as preservatives because they liberate SO2 in solution

Sulfuric Acid – a typical acid used in fertilisers, paints, pigments, dyestuffs, chemical manufacture, soaps and detergents and fibres. Contact Process – the industrial preparation of sulfuric acid. All reactions in it are exothermic. 1. Sulfur is burned in air: S(s) + O2(g)  SO2(g) 2. The SO2 is reacted with further oxygen over a catalyst bed (vanadium (V) oxide). The vanadium (V) oxide is a catalyst which speeds up the reaction without being used up. It melts at 400 °C, spreading to give a larger area. The yield of SO2 is sufficiently high for this stage to be carried out at atmospheric pressure. V2O5 2 SO2(g) + O2(g)  2 SO3(g) 3. The sulfur trioxide is reacted with 98% H2SO4 to form oleum, which then reacts with water to form more sulfuric acid. The SO3 must be reacted with sulfuric acid first and not immediately with water as it is too exothermic/violent to carry out directly. SO3(g) + H2SO4(l)  H2S2O7(l) H2S2O7(l) + H2O(l)  2 H2SO4(l)

Cells Electrolyte – molten or dissolved metal compounds that conduct electricity. When electrolytes conduct electricity, ions move. Electrode – most electrodes are metals or graphite. When metals conduct electricity, valence electrons move from ion to ion, from the negative to positive electrode. When graphite conducts, the delocalised electrons between the layers can flow. The two electrodes are called the cathode (which is negatively charged and attracts cations) and the anode (which is positively charged and attracts anions). Reduction occurs at the cathode and oxidation occurs at the anode. The mobile ions of the electrolyte carry the current between the electrodes. Graphite electrodes must be replaced periodically because graphite will react with oxygen to form CO2. Electrochemical Cells – produce electricity spontaneously via a chemical reaction (redox) between two metals. 







 

When two metals of different reactivity are connected electrically in a complete circuit with a conducting wire and an electrolyte, electrons flow from the more reactive metal to the least reactive metal. The electron flow is called current, and the energy transfer from the higher to lower reactivity metals is called the voltage. The greater the difference in reactivity between the two metals making up the electrodes, the greater the energy transfer and therefore greater voltage of the cell. A dry cell uses a damp paste of ionic material (salt bridge) between the electrodes instead of a liquid electrolyte. More than one cell connected together is called a battery. The more reactive metal is always the negative electrode.

Hydrogen Fuel Cell Hydrogen at the anode and Oxygen at the cathode combine to form water. The reduction of Hydrogen at the anode causes the lost electrons to form a current on their way to reducing oxygen at the cathode. 2 H2  4 H+ + 4 eO2 + 4 e-  2 O22 O2- + 4 H+  2 H2O Advantages Only product is H2O (no CO2) Hydrogen is very abundant (in compounds) It is renewable

Disadvantages Risk of explosions/gases takes up a lot of volume

Electrolysis Electrolysis - the passing of a direct current through a conducting solution or liquid and the resultant decomposition of the electrolyte. It uses electricity from a power source in order to cause a chemical reaction (redox). It will force the oxidation or reduction of substances that are high in reactivity that do not naturally oxidise or reduce using normal chemical processes.

Direction of electron flow

-

+

-

   

+

Cations are attracted to the cathode and are discharged (converted to a new substance) Anions are attracted to the anode and are discharged The electrolyte is decomposed The electrodes are usually made of graphite or a completely unreactive metal (e.g. platinum) Selective Discharge Rules

The ease of discharge of an ion depends on several factors, including the nature of the electrode, and the nature of the electrolyte (molten/aqueous, concentrated/dilute). 1. Cations always discharge at the cathode 2. Anions always discharge at the anode 3. The ions of the more reactive metals are more difficult to discharge than those of less reactive metals 4. The sulphate and nitrate ions are never discharged (but not when altered) 5. Halide ions will be difficult to discharge 6. The more concentrated the solution, the more chance the ions will be discharged Note: if chloride ion is present it is most likely always going to be the anode product. The following equation represents the discharging of Hydroxide ions: 4 OH- O2 + 2 H2O + 4 e-

Production of Caustic Soda Caustic Soda – a common name of sodium hydroxide. The electrolysis of a concentrated sodium chloride solution (brine) produces three products: hydrogen, caustic soda and chlorine. The Membrane Cell method – the membrane is permeable to cations so only the sodium ions can flow through and the hydroxide ions cannot flow back to the anode. The membrade is a porous, thin, flexible sheet.

The m em brane cell for m aking N aOH chlorine

hyd rogen m em brane

brine

d ep leted brine

w ater

anod e + ve (Ti or grap hite)

     

35 % N aOH

N a+

cathod e - ve (steel or graphite)

Anode: 2Cl-  Cl2 + 2eCathode: 2H2O + 2e-  2OH- + H2 At the cathode, Hydroxide ions are formed (produced together with the hydrogen) As the sodium ions move towards the cathode, a solution of sodium hydroxide is thereby formed The water molecules are reduced at the cathode The sodium ions are not reduced and apart from moving through membrane do not change Chlorine is used in PVC and in water Sodium Hydroxide is used in pulp and paper, and soap. Production of Aluminium

1. Mine the bauxite ore (a mixture of Al2O3 and SiO2 as well as other impurities such as Fe2O3) 2. Purify the ore by dissolving it in sodium hydroxide solution. This dissolves the alumina (aluminium oxide) which is amphoteric but not the basic impurities e.g. Fe2O3 3. At the aluminium smelter, the alumina is dissolved in cryolite (Na3AlF6) because this gives the mixture a much lower melting temperature (900 °C) and it conducts electricity better. The mixture is 95% alumina and 5% cryolite. 4. Large amounts of electrical energy are passed through the mixture. The anodes are carbon rods. The cathode lining is graphite in steel casing. The passing of electric current causes electrolysis to occur. Anode: Al3+ + 3 e-  Al Cathode: 2 O2-  O2 + 4eOverall: 4 Al3+ + 6 O2-  3 O2 + 4 Al 5. The molten aluminium is poured into ingots and used for many purposes.

Purification of Copper

      

Copper metal is readily extracted by roasting copper ores malachite (impure CuCO3) and copper pyrites (CuFeS2) to obtain copper. Impure copper was at the anode and pure copper was at the cathode Aqueous copper(II) sulphate was the electrolyte At the anode, metals more active than copper are oxidised to their cations and remain as cations and must be removed as they accumulate. Copper is oxidised to Copper(II), while metals less reactive than copper are not oxidised but instead fall to the bottom of the cell and are removed through filtration of the electrolyte. The impurities are mainly Ag and Au and are called the anode sludge. At the cathode, copper ions are reduced to copper metal (almost 100% pure). Cu2+ (aq) + Cu (s)  Cu (s) + Cu2+ (aq) Electroplating Metals

Electroplating – the process involving electrolysis to coat one metal with another. Often the purpose of electroplating is to give a protective coating to the metal beneath or as a decorate coat.

  

To plate an object with a metal, the object to be electroplated is made the cathode in an electrolysis cell. The anode is made from the metal that is to be the coating The electrolyte will be a salt solution of the metal to be electroplated

Hydrocarbons Hydrocarbons - a substance made of Hydrogen and Carbon Homologous Series – a series of carbon compounds differing from each other only by the addition of more CH2 groups to increase the length of the carbon chains. Isomers – different forms of the same molecular formula with different structural formulae         

    

Alkanes Non-polar molecules Weak intermolecular attractions Low melting point and boiling point (but increases as size increases) Lower density than water Saturated Hydrocarbons Single bonds only Formula = CnH2n +2 Gas state between 1-4 C’s Liquid state between 5-17 C’s e.g. Ethane (C2H6)

Alcohols Polar molecules (decreases as more carbons added) Strong intermolecular attractions Formula = CnH2n +1OH Colourless volatile liquids Burn cleanly and efficiently, but with less energy from presence of oxygen e.g. Ethanol (C2H5OH)

       

     

Alkenes Non-polar molecules Weak intermolecular attractions Low melting point and boiling point (but increases as size increases) Unsaturated Hydrocarbons Contain a double bond Formula = CnH2n Extremely reactive due to double bonds breaking Turns bromine water from red to colourless e.g. Ethene (C2H4)

Carboxylic Acids Polar molecules Formula = CnH2n+1COOH Strong intermolecular attractions Weak acids Formed through the oxidation of Alcohols Reacts with Alcohols to form Esters

e.g. Ethanoic Acid (CH3COOH)

Alkane Reactions 1. Combustion – alkanes burn in oxygen to form carbon dioxide and water as long as sufficient oxygen is present; if insufficient, carbon monoxide or carbon will be produced instead of carbon dioxide 2. Substitution – alkanes will react with halogen molecules in a substitution reaction e.g. C2H6 + Br2  C2H5Br + HBr Alkene Reactions 1. Hydrogenation - Addition by hydrogen. Alkanes are formed when the H2 adds to the alkene molecule. A catalyst of nickel or platinum is used at a temperature of about 150 °C e.g. CH2 = CH2 + H2  CH3CH3 2. Halogenation – Addition of bromine or other halogens. Halogen alkenes are formed when halogens attach to the carbons in the double bond by covalent bonds e.g. CH2 = CH2 + Br2  CH2BrCH2Br 3. Hydration – Addition of water. Alcohols form when water is added to alkene molecules. A catalyst of dilute H2SO4 or H3PO4 is used. CH2 = CH2 + H2O  CH3–CH2–OH

Number of Carbons 1 2 3 4 5 6 7 8 9 10

Alkane Methane Ethane Propane Butane Pentane Hexane Heptane Octane Nonane Decane

Alkene --Ethene Propene Butene Pentene Hexene Heptene Octene Nonene Decene

Alcohol Methanol Ethanol Propanol Butanol Pentanol Hexanol Heptanol Octanol Nonanol Decanol

Carboxylic Acids Methanoic Acid Ethanoic Acid Propanoic Acid Butanoic Acid Pentanoic Acid Hexanoic Acid Heptanoic Acid Octanoic Acid Nonanoic Acid Decanoic Acid

Crude Oil Fossil Fuel - organic matter (once living) e.g. coal (dead plant matter), oil (dead sea creature remains). Crude oil - made of hydrocarbons. It is the result of heat and pressure on plant and (sea) animal remains over millions of years in the absence of air. This oil (and gas) rises up through permeable rocks and becomes trapped under impermeable rocks, so they have to be extracted by drilling. The oil is called crude oil because it is unrefined which makes it of little use as it is hard to transport. Fractional Distillation - process used to separate a mixture of liquids that have different boiling points. When the mixture is heated, liquids with low boiling points evaporate and turn to vapour and can then be separated as liquids. Those with high boiling points remain liquids.

Cracking – allows large hydrocarbon molecules to be broken down into smaller, more useful hydrocarbon molecules. Fractions containing large hydrocarbon molecules are vaporised and passed over a hot catalyst. This breaks chemical bonds in the molecules, and forms smaller hydrocarbon molecules. Cracking is an example of a decomposition reaction. e.g.

Polymers Polymers - very large molecules made when hundreds of monomers join together to form long chains. They have no double bonds. Synthetic Polymers (Plastics) – man-made polymers Monomers - a molecule that can be bonded to other identical molecules to form a polymer. Polymerisation - the combining of monomers to form polymers Addition Polymers – the monomer is thousands of the same alkene molecules, whose double bond is broken to join the molecules together in one long chain e.g. If you had n (number) of this monomer:

Then n of the monomers would join together, by breaking the double bond and connecting to the other monomers to form a long chain:

This would then be written as a repeated unit:

Condensation Polymers – a polymer formed by a condensation reaction (one in which water is given out). Artificial Polymers – where a product is formed from two different types of monomers arranged alternately and linked together. In this case the monomers usually contain a minimum of two of the same/different functional groups at the end of their molecules. When the polymer forms, a small molecule such as H2O or HCl is lost at each junction. Artificial polymers include: 

Ester Linkage – a polymer found between di-carboxylic acids and a diols. (-COO-) e.g. Terylene n

HO

H

H

C

C

H

H

O

OH H

+

HO O H

H

H

C

C

H

H

C

C

O

O

O

OH

C

C

O

O

n H2O

+

n

Ester Link 

Amide Linkage - a polymer formed between di-carboxylic acids and diamines (-CONH-) e.g. Nylon

n

H

H

H

N

C

C

N

H

H

H

H

H

+ HO

C

C

O

O

OH

H

H

N

C

C

N

C

C

H

H

H

H

O

O

+

n H2O

n

Amide Link Natural Polymers – those polymers found in nature and are usually condensation polymers 

Peptide (Amide) Linkage – a protein (polypeptide) formed between amino acid molecules

H

N

C

C

H

H

O

OH

+ H

N

C

C

H

H

O

H



n

OH

N

C

C

N

C

C

H

H

O

H

H

O

OH

+ n H2O

Amide Link Polysaccharides – complex carbohydrate molecules made my polymerizing simple sugar molecules such as glucose HO

OH

O

O

O

O

+ n H2O

Proteins – polymers of amino acids formed by condensation reactions. Amino Acids – naturally occurring organic compounds which possess both an –NH2 group and – COOH group on adjacent carbon atoms. There are 20 naturally occurring amino acids, of which glycine is the simplest. Carbohydrates – a group of naturally occurring organic compounds which can be represented by the general formula Cx(H2O)y Sugars – any of the class of soluble, crystalline, typically sweet-tasting carbohydrates found in living tissues and exemplified by glucose and sucrose. They are tested by warming with Benedicts or Fehlings solutions; if the sugar is present, the colour changes from a blue solution to an orange-red suspension or precipitate. Starch – made up of 200-300 glucose monomers. Starch turns iodine solution from red-brown to blue-black colour. Cellulose – made up of about 3000 glucose monomers. Fats and Oils Fats – naturally occurring polyesters with the same link between ester monomers as Terylene. The chains typically contain 12 to 20 carbon atoms. Fats and oils are rich in energy and this is their normal function to us. They are also important in soap and detergent products. A fat molecule is made of two components, a glycerol (the “backbone” of the molecule) and fatty acids (which are attached to the backbone. Glycerol

Fatty Acids

Formation of Triglyceride

R is the fatty acid chain Soaps – long chains derivatives of fatty acids. The fatty acid is reacted with a base such as caustic soda. This causes the formation of the sodium salt of the fatty acid which is used as soap. Soap (sodium stearate) is an ionic compound and can remove dirt with the covalent end (non-polar) attracting to the dirt and the ionic end attracting to water molecules. Soap molecules can make oils and water form a stable emulsion. However, it forms a scum with hard water by reacting with Ca2+ or Mg2+ present.

Air Gas Nitrogen Oxygen Argon Carbon Dioxide Others Water

Percentage in Air 79 20 0.9 0.03 0.07 -

Percentage in Inhaled Air 79 20 1 (with other Inert Gases) Trace Variable

Percentage in Exhaled Air 79 17 1 (with other Inert Gases) 4 Saturated

Fractional Distillation of Air – the main industrial method of preparation of pure oxygen and nitrogen   

Air is liquefied by compression and cooling to below the boiling point of both oxygen and nitrogen, so that most of the “air” becomes a liquid. The liquid “air” is allowed to warm slowly and the nitrogen (b.p. -195°C) boils off first and thus can be extracted. The oxygen (b.p. -183°C) boils off after the nitrogen and can then also be extracted.

Pollutants Sulphur Dioxide

Nitrogen Oxides

Carbon Monoxide

Lead

Cause Combustion of Sulphur (found in fossil fuels)

Effect Solution Forms acid rain when Scrubbing – converts reacted with water in SO2 to H2SO4 clouds S(s) + O2(g)  SO2(g) SO2(g) + H2O(l)  H2SO3(aq) Reaction between Contributes to Acid Rain Catalytic Converters nitrogen and oxygen at and are a major high temp and pressure component of in motor vehicle engines photochemical smog 2N2(g) + 3O2(g)  2NO(g) + 2NO2(g) Incomplete combustion Strong bonds with Catalytic Converters in motor vehicle engines haemoglobin, decreasing the amount of oxygen distributed around the body Combustion of petrol Brain damage, nervous Catalytic Converters, containing lead system problems using unleaded fuel

Catalytic Converter –catalysts that convert poisonous exhaust fumes into harmless gasses in cars. 





The reduction catalyst uses platinum and rhodium to help reduce the NOx emissions by ripping the nitrogen atom out of the molecule, freeing the oxygen. 2NO  N2 + O2 / 2NO2  N2 + 2O2 The oxidation catalyst removes the unburned hydrocarbons and carbon monoxide by burning them over a platinum and palladium catalyst. 2CO + O2  2CO2 2NO2 + 4CO  4CO2 + N2

Oxygen Laboratory Preparation of Oxygen 1. Heating Potassium Manganate (VII) 2 KMnO4 (s)  K2MnO4 (s) + MnO2 (s) + O2 (g) purple  green + black + colourless

2. Decomposing hydrogen peroxide by a catalyst of manganese (IV) oxide 2 H2O2 (l)  2 H2O (l) + O2 (g) Industrial Uses of Oxygen      

Steel making In oxy-acetylene welding In hospitals and ambulances for treatment of trauma patients In rockets to combine with the fuel In deep sea diving helium-oxygen mixtures Compressed air is avoided in deep sea diving because the nitrogen causes ‘nitrogen narcosis’ when the diver surfaces too rapidly and the nitrogen bubbles out of the blood due to the rapid decrease in pressure. Nitrogen Properties

     

Very slightly soluble in water Very unreactive because of its strong intramolecular nitrogen to nitrogen triple bond Will react with some substances under the right conditions Has low boiling point of -196°C Non-polar molecule with weak intermolecular attractions Has a low melting point of -210°C Uses

   

Filling spaces in food packaging and oil tanks Liquid nitrogen is used for freezing food, gametes, and other delicate materials Production of Ammonia in the Haber Process Creating a non-oxidising environment for fruit storage

Ammonia Preparation 1. Ammonia gas (NH3) can be conveniently prepared in the laboratory by heating together an ammonium salt with a strong alkali e.g. Ca(OH)2 (s) + 2 NH4Cl (s)  CaCl2 (s) + 2 H2O (l) + 2 NH3 (g) 2. Ammonia can be collected by the downward displacement of air since it is lighter than air 3. Ammonia cannot be collected by the displacement of water because it is very soluble in water 4. Ammonia can be collected through the Haber Process Properties of Ammonia     

Colourless Strong choking smell Less dense than air Liquefies at -33°C. This makes it easy to transport and store as a liquid Extremely soluble in water, as it is a polar molecule and can hydrogen bong with itself and water molecules, to produce an alkaline solution  The only common alkaline gas Uses of Ammonia     

Making nitric acid Making fertilisers such as urea, ammonium nitrate, and ammonium sulphate Household cleaners Dyes Explosives Urea Production

Urea – an important nitrogenous (nitrogen containing) fertiliser. It is a white, water-soluble solid. Urea is less soluble than inorganic fertilisers and so releases the nitrogen slowly to plants; CO(NH2)2. The process occurs at 190°C and 230 atm. CO2 (g) + 2 NH3 (g)  CO(NH2)2 (s) + H2O (g) Ammonium Compounds Ammonia – NH3; polyatomic molecule Ammonium – NH4+; polyatomic charged ion, only found in ammonium compounds Ammonium Salts – formed through the reaction of ammonia with the appropriate acid. They are used as fertilisers to supply nitrogen to plants. Fertilisers are given an NPK rating (Nitrogen, Phosphorus and Potassium are best used in fertilisers). Ammonia can be displaced from its salts through decomposition by heating or by the action of strong bases.

Limestone Limestone – CaCO3; various forms of lime are used to put on pastures to raise the pH because many soils are naturally acidic. Intensive cropping also lowers the pH. The lime is basic so it neutralises the soil and brings the pH closer to 7. Quicklime – CaO; formed from limestone in a lime kiln (oven with extremely high temperatures). CaCO3 (s)  CaO (s) + CO2 (g) Slaked Lime – Ca(OH)2; the solid product, in a form of white powder, of the exothermic reaction that occurs when a minimal amount of water is added to quicklime. CaO (s) + H2O (l)  Ca(OH)2 (s) Limewater – a solution of slaked lime in excess water. The slaked lime is only sparingly soluble but produces an alkaline solution containing calcium and hydroxide ions. Ca(OH)2 (s)  Ca2+ (aq) + 2 OH- (aq) Mortar – a mixture of slaked lime, sand and water and is a thick paste. It sets when it dries, then over a long period of time becomes hard due to the formation of calcium carbonate as it absorbs carbon dioxide from the atmosphere. Ca(OH)2 (s) + CO2 (g)  CaCO3 + H2O (l) Cement – made from heating limestone with sand and silicates such as clay. It is a mixture of calcium silicates and aluminates. When water is added, a complex series of reactions occur which make it set. Carbon Cycle Methane – can be sourced from natural gas trapped in oil-bearing rocks, partial decomposition of plant materials under anaerobic conditions, waste product of digestion in animals.

Water Properties      

Colourless, odourless liquid 0°C Melting Point 100°C Boiling Point V shaped or bent; 105° bond angle Molecules join via hydrogen bonding Turns anhydrous copper sulphate from white to blue CuSO4 (s) + 5 H2O (l)  CuSO4.5H2O (s) + heat  Turns anhydrous cobalt chloride from blue to pink CoCl2 (s) + 6 H2O (l)  CoCl2.6H2O (s) + heat  Water is used in generating electricity, cleaning, cooling, dissolving, manufacturing products, making concrete, and consumption in the house (e.g. toilets and showers) Desiccator – sealed glass basins used to keep substances and papers (e.g. cobalt chloride paper) dry The Hoffman’s Voltamater

Since water is not ionic, we cannot electrolyse pure water as there are no charged particles to carry the current. If we add some dilute sulphuric acid this dissociates in the water and allows electrolysis to occur. The apparatus is known as Hoffmann’s Voltameter. Cathode: 4 H2O (l) + 4 e-  2 H2 (g) + 4 OH- (aq) Anode: 2 H2O (l)  O2 (g) + 4 H+ (aq) + 4 OH- (aq) Full: 2 H2O (l)  2 H2 (g) + O2 (g)

Water of Crystallisation Ionic solids often have molecules of water bonded into their ionic crystal lattice. This water is called water of crystallisation and often has a consistent, simple ratio in the formula. Formulae are often quoted with the water of crystallisation included. E.g.Copper sulphate pentahydrate, Sodium carbonate decahydrate. Agents Drying Agents – various drying agents are used to absorb water out of air or gas mixtures. Common drying agents are:    

Concentrated sulphuric acid is used as a drying agent in the preparation of some neutral or acidic gases Calcium oxide which is used to dry alkaline ammonia gas Anhydrous calcium chloride is used to prepare dry hydrogen Silica gel is used as a drying agent in equipment which is sensitive e.g. cameras

Dehydrating Agents – drying agents that are so powerful that they will remove all atoms required to make up water from certain substances in the solution of solid form. The most common example of a dehydrating agent is concentrated sulphuric acid which will dehydrate sucrose and hydrated copper sulphate. It will also extract the elements of water from cloth or skin. Purification of Water   

Water is stored in dams and reservoirs. It is never completely pure and may contain bacteria, dissolved substances and solid material which need to be removed. Concerns over levels of pesticides in river water have led to improvements in water purification. Water treatment essentially involves the stages, filtration, sedimentation, (ozone treatment) and chlorination.

Laboratory Equipment

Pestle and Mortar Grinding or breaking