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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 (I 2 and CO 2 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:
Pure Impure Melting Point Sharp Melting Point; usually high Range of temperatures Boiling Point Sharp Boiling Point; usually low Range of temperatures Chromatography One well-defined Spot on chromatogram Several spots on chromatogram
Bonding
Metal + Electrons = Metallic Bonding Metal + Non-Metal = Ionic Bonding Non-Metal + Non-Metal = 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 H 2 O 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
Metallic Lattice – the regular arrangement in metal ions in solid metals
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
Polyatomic ions – ions containing more than one atom. Brackets must be used to write formulae involving more than one of these ions. E.g. Al 2 (SO 4 ) 3 ; SO 4 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
A
A
A
Cat-Ions
Free (delocalised) electrons, they can move anywhere
Na Cl Na Cl
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
Non-Polar Molecules – molecules where all bonds are non-polar or with a symmetrical shape causing the dipoles to cancel out
H H
O
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
C
H H
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
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
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.
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
Melting Evaporation
Sublimation
Solidification
Deposition
Solid
Melting
Liquid
Vaporising
Gas
Condensation
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.
Ionic Bond Covalent Bond Bonding - Results from the attraction between positive and negative ions
- Occurs when a metal reacts with a non-metal - Formed by the sharing of electrons between atoms
- Occurs when two non-metals react
Properties - Does form ions so it will conduct electricity as liquid or aqueous
- Forms a crystal lattice structure
- Has a high melting point
- 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:
- Temporary dipole attractions – the weakest attraction between non-polar molecules
- Permanent dipole attractions – the next strongest attraction between polar molecules
- Hydrogen bonding – strongest attraction between polar molecules and extremely reactive non- metals (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
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 6x10^23 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 ( A r) in grams. One mole of molecules has a mass equal to the relative molecular mass ( M r) in grams.
The calculation of no. of moles is moles =
The calculation of Molar Mass is Molar Mass =
The calculation of mass is mass = number of moles x Molar Mass
The calculation of concentration is concentration =
The calculation of Volume is Volume =
The calculation of no. of moles is moles = concentration x Volume
The calculation of no. of moles is moles =
One mole of any gas at rtp = 24
The calculation of Volume = Volume x 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 Calcium Bromine 1 - % By Mass (or Mass) 20% 80% 2 – Mole Ratio 20/40 = 0.5 80/80 = 1 3 – Ratio by atoms 0.5/0.5 = 1 1/0.5 = 2 4 - Empirical Formula Ca = 1 Br = 2
n M
m
c V
n
n 24
V
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) CaCO 3 (11.7g) CO 2 + CaO; 100 = Mr of CaCO 3 , 44 = Mr of CO 2 n(CaCO 3 ) = = 0.307 moles 44 x 0.307 = 13. = 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 CuCl 2 + 2H A + BC A + B
Double Displacement - two elements displaced Acid + Base Salt + Water e.g.
2KI + Pb(NO 3 ) 2 PbI 2 + 2KNO 3 AB + CD CB + AD
Synthesis - two elements combine to form one compound Most Metals + Oxygen Gas Metal Oxides e.g.
2Mg + O 2 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.
CuCO 3 CuO + CO 2 AB A + B
Combustion – oxygen combines with another compound to form water and CO 2 Organic Molecule + excess Oxygen Gas Carbon Dioxide + Water e.g.
C 10 H 8 + 12 O 2 10 CO 2 + 4 H 2 O AB + C AC + BC
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.
Endothermic Reactions – when energy is absorbed by the products from the surroundings during a reaction.
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.
Energy Level
Reaction Progress
Reactants
Products
Reaction Progress
Products
Reactants
Energy Level
Energy Lost (ΔH)
Ea
Ea
Energy Absorbed (ΔH)
Factors affecting Equilibrium
Factor Increase of Factor Decrease of Factor
Temperature
Equilibrium shifts to decrease the temperature so it shifts to the endothermic direction
Equilibrium shifts to increase the temperature so it shifts to the exothermic direction
Concentration
Equilibrium shifts to decrease the concentration
Equilibrium shifts to increase the concentration
Pressure
Equilibrium shifts to decrease the pressure so it shifts in the direction of the least molecules
Equilibrium shifts to increase the pressure so it shifts in the direction of the most molecules
Catalyst
Speeds up the time it takes to reach equilibrium but does not change the position
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:
N 2 + 3 H 2 2 NH 3 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.
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 Base Sour Taste Bitter Taste pH less than 7 pH greater than 7 In solution, contains hydronium ions (H 3 O+) In solution, contains hydroxide ions (OH-) Turns blue litmus red Turns red litmus blue Turns phenolphthalein colourless Turns colourless phenolphthalein pink Corrosive Soapy feel Reacts with metals to produce salt and hydrogen Cannot react with metals Examples of Acids Examples of Bases Hydrochloric Acid HCl Sodium Hydroxide NaOH Nitric Acid HNO 3 Potassium Hydroxide KOH Sulphuric Acid H 2 SO 4 Calcium Hydroxide Ca(OH) 2 Ethanoic Acid CH 3 COOH Ammonia Solution NH 3 (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 H 3 O+.
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 H 3 O+.
Strong Acids Weak Acids Dissociation in Aqueous Solution Completely dissociate Partially dissociate Equilibrium None (forward only) Equilibrium reaction Electrolyte Good Poor Electrical conductivity Good Poor [H 3 O+] Higher Lower pH value Lower Higher Examples HCl, HNO 3 , H 2 SO 4 CH 3 COOH, NH 4 +
Amphiprotic – substances can act as both an acid and a base e.g. H 2 O, HCO 3 - , HSO 4 -
Amphoteric – 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 dm^3 , 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 Usually Soluble Usually Insoluble All NO 3 -^ All SO 42 -^ EXCEPT Ba, Pb, Ag, Ca All CO 32 -^ EXCEPT Group 1 All NH 4 +^ All Cl-^ EXCEPT Ag, Pb All O^2 -^ EXCEPT Group 1 and 2 Group 1 All I-^ EXCEPT Ag, Pb All OH-^ EXCEPT Group 1; All Br-^ EXCEPT Ag, Pb Ca and Ba are slightly soluble
Physical Properties of Metals, Non-Metals and Metalloids
Metals Non-Metals Metalloids Lustre (shiny) No lustre Can be shiny or dull Good conductor of heat Poor conductor of heat Fair conductor of heat Good conductor of electricity Poor conductor of electricity Fair conductor of electricity Malleable Not Malleable Malleable Ductile Not Ductile Ductile High Density Low Density Solids High Melting Point Low Melting Point
Chemical Properties of Metals and Non-Metals
Metals Non-Metals Easily loses electrons Tends to gain electrons Oxides generally basic and amphoteric Oxides generally neutral Corrodes easily
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) Chemical Properties (compared to Group 1) Much harder Much less reactive Higher tensile strength Many have excellent corrosion resistance Higher density Show more than one valency ( e.g. Fe2+^ or Fe3+) Higher melting point and boiling point Them and their compounds are useful catalysts Many of their compounds are coloured Some are strongly magnetic
Alloys
Alloy Mixture Use Solder 70% Tin, 30% Lead Joining wires and pipes Brass 60 - 95% Copper, 5-40% Zinc Taps, hose/pipe fittings, zips, screws Bronze 90% Copper, 10% Tin Ornaments, bells, bearings Mild Steel 99.5% Iron, 0.5% Carbon General structural purposes, cars Hard Steel 99% Iron, 1% Carbon Blades Stainless Steel 74% Iron, 18% Chromium, 8% Nickel Corrosion resistance Alnico Iron/Aluminium/Nickel/Cobalt Permanent Magnets
Metals Uses
Metal Property Uses Aluminium Does not corrode Food containers Low density, unreactive Containers and packaging buildings Low density, strong, conducts Long distance wiring Low density, strong, cheap Transport vehicles Low density, conducts heat Car Engines Zinc Reactive Dry cells (“batteries”) More reactive than iron Galvanising Iron Iron Similar expansivity Reinforcing concrete Strong, cheap Nails Strong and abundant Ship building Copper Good conductor of electricity Electrical wiring Unreactive, workable Alloys – brass and bronze Unreactive Coinage (with Nickel) Unreactive Hot water piping
Reactivity Series
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 Preventing Oxidation of Metal Painting Galvanising – zinc atoms react before the iron Greasing of metal parts Sacrificial protection – a more reactive metal Oiling of bike chains reacts before the metal that it is protecting Tin Plating – In cans Carthodic protection – an electric power source pushes electrons into the metal to prevent the loss of electrons
Plastic covering on electric wires Galvanising – zinc coating for galvanised steel Chromium plating of car parts
Most Reactive K Potassium
Na Sodium
Ca Calcium
Mg Magnesium
Al Aluminium
C Carbon
Zn Zinc
Fe Iron
Sn Tin
Pb Lead
H Hydrogen
Cu Copper
Ag Silver
Au Gold
Least Reactive Pt 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.
