Chem Paper 1
C1 - Atomic Structure and the Periodic Table
Atoms, Elements, Mixtures and Compounds
All Substances are made of atoms
Atoms are the smallest part of any element that can exist - but not the smallest part of matter
The periodic table consists of ~100 different elements, each represented by an individual symbol
Chemical reactions can be represented by either word or symbol chemical equations
Mixtures consist of two or more elements which aren't chemically bound together
Methods of separating mixtures are distillation, fractional distillation, crystallisation, filtration and chromatography
Structure of the Atom over Time
Simple model
Before electrons were discovered, atoms were believed to be just tiny spheres of matter
Plum pudding model
The nucleus was a ball of positive charge with negative charges embedded into the surface
Nuclear model
The nucleus was actually tiny - only 1/10,000th of the atom's diameter - surrounded by electrons (first thought to orbit at random distances) found to be orbiting at specific orbitals / energy levels
The Structure of the Atom
An atom's radius is approx. 1 ÷ 100,000,000,000 metres, of which the nucleus takes up 1/10,000
The nucleus is the majority of the mass in an atom
Electrons will occupy the lowest available orbital / energy level
Elements in the same group in the periodic table have the same number of electrons in their outer shells
Elements in the same period have the same number of electron shells but not necessarily the same number of electrons
Properties of Transition Metals
Transition metals often form two or more ions - e.g Fe2+ and Fe3+
Transition metals are less reactive than the group 1 metals, and often form ionic compounds with oxygen - metal oxides
Many of these transition metals form coloured compounds and are useful as catalysts in other reactions
C2 - Bonding, Structure and Properties of Matter
Types of Bonds
Ionic Bonding
Occurs between a metal and a non-metal. This is because metals form + ions and non-metals form - ions, so there is a force of attraction between them.
Ionic substances for giant ionic lattices which have very high melting points. These structures also have a very low electrical conductivity as there are no free electrons to carry the charge.
However, when dissolved in water or melted, the individual ions are free to move around - allowing them to carry a charge and thus conduct electricity. This is the foundation principle of electrolysis.
Covalent Bonding
Covalent bonds form solely between non-metals. There are two types of covalent molecules; small covalent molecules and giant covalent molecules.
Small covalent molecules have a low melting and boiling point as - although the intramolecular forces are strong - the intermolecular forces are weak, and it is these that are broken in the process of melting or boiling a substance. SCM's are often gases at room temperature (H2) and do not conduct electricity.
Giant covalent molecules have very high melting and boiling points as all the atoms are joined together by intramolecular covalent bonds, which are very strong. Also, they are very hard and do not conduct electricity.
Metallic Bonding
Metallic bonds form purely between metal atoms. Each atom 'donates' electrons to the sea of delocalised electrons which causes a force of attraction between all of the metal ions, forming a very strong molecule. Metals have very high melting and boiling points, as well as being malleable in their pure forms due to the uniformity of their structure.
The sea of delocalised electrons allows for metals to be great conductors of electricity, as each of the electrons can move freely, carrying a charge flow.
Fullerenes
Fullerenes are hollow molecules of carbon atoms. Each is made of rings of 5,6 or 7 carbon atoms joined together.
Buckminsterfullerene is also notated as C60 as it has 60 carbon atoms in each molecule. It is spherical in shape, meaning it has a wide variety of uses.
Carbon nanotubes are a fullerene with a very high length:width ratio. Their properties make them very useful for nanotechnology, electronics (they have 1 free electron per carbon) and materials technology. An example of the latter is bulletproof and knifeproof clothing.
The States of Matter
Types of Matter
Solids
Solids have a generally uniform structure in the particle model, with none of the particles being able to move more than jiggle a bit.
Solids are considered uncompressable.
Particles in a solid have a very low internal energy store.
Liquids
Particles in a liquid have a slightly higher internal store of energy, and are thus free to move around more as some of the bonds joining them are broken.
Liquid versions of materials take up more room than the solid version - the exception being ice and water.
Gas
Gases have a high internal store of energy and are thus completely free to move around. They mover around randomly and take up the most room per unit of mass of any of the states of matter.
Plasma
Plasma is essentially ionised gas, where the individual atoms have lost some of their electrons. This results in a mixture of positive ions and negative electrons. If the plasma loses heat, the ions and electrons reform into a gas.
Antimatter
Antimatter is a strange substance we don't yet fully understand. However, we know that each subatomic particle of an antimatter atom has the exact same mass as it's normal matter counterpart (or partner), with the opposite electric charge. So an electron is 1- charge, and a positron is 1+.
Other Stuff
Nanoscience
Anything that is 1-100nm in size is a nanoparticle
Nanoparticles have different properties to the bulk version of the same material as the surface area to volume ratio is so huge.
Often, smaller amounts of nanoparticles are needed than bulk materials to be effective.
Other Particles
Fine particles have diameters between 100 and 2500nm
Coarse particles are anything between 0.0000025m and 0.00001m.
If you halve the length of the sides of a cube, you double the surface area to volume ratio.
C3 - Quantitative Chemistry
Conservation of Mass
This law states that no atoms are lost or made during a reaction; so the mass of the products has to equal the mass of the reactants.
Any apparent changes in mass can usually be attributed to faulty equipment or a gas being produced / absorbed.
In a balanced chemical equation, the Mr of the reactants should equal the Mr of the products.
Chemical Equations
The balancing numbers can be obtained from using the mass in grams of the reactants in grams and converting these to moles, then simplifying the ratio if possible.
Limiting Reactants
The reactant that runs out first is the limited reactant.
The rest are then in excess.
It is common practice to put the reactants you don't need to all react in excess, so that the reactant that determines the mass of the product (for example) is the limiting factor. This way you can control the extent to which the reaction occurs.
Yield and Atom Economy
Some atoms are lost to unwanted reactions or residue that was not collected.
Atom economy is measured by the mass of reactant that ended up as useful products.
The percentage yield is the actual yield as a percentage of the theoretical yield.
Concentration of a solution is measured in moles per dm cubed.
If the volumes of two solutions react completely and one of the concentrations is known, then we can work out the other concentration as a ratio.
C4 - Chemical Changes
The Reactivity Series
Metals react with oxygen to produce metal oxides. These are oxidation reactions.
OILRIG
There are oxidation reactions, where electrons are lost.
There are reduction reactions, where electrons are gained.
Reduction reactions involve the loss of oxygen atoms.
The Series
Copper
Iron
Hydrogen
Lead
Tin
Zinc
Carbon
Aluminium
Magnesium
Calcium
Sodium
Potassuim
The reactivity series is ordered in terms of reactivity with water and dilute acids.
A more reactive metal can replace a less reactive metal in a compound.
Metals less reactive than carbon can be removed from their compounds by reduction with carbon.
Reactions of Adics
When acids react with metals they produce salts and hydrogen.
Acids are neutralised by alkalis (e.g. soluble metal hydroxides). This produces salts and water.
The reaction between the acid and the alkali produces water and salts which are determined based off of the acid used and the positive ions in the base/alkali/carbonate.
Acids produce H+ ions in aqueous solutions. Aqueous solutions of alkalis contain OH- ions.
The pH Scale
Ranges from 0-14, with 0 being the strongest acid and 14 being the strongest alkali.
pH 7 is neutral (pure water is pH 7).
Universal indicator determines what pH a solution is based off of a colour scale.
A strong acid (pH 0) will be completely ionised in an aqueous solution - an example of this being HCl, H2SO4 and HNO3.
A weak acid is only parially ionised in an aqueous solution - examples of this being CH3COOH, C6H8O7 and H2CO3.
The pH scale is logarithmic - decrease of 1 on the scale increases the H+ ion concentration by a factor of 10.
Titrations
The volumes of acid and alkali solutions that react with eachother can be measured by titration using a suitable indicator.
Process of Titration
- Have the solution of unknown solution in a conical flask.
- Add a solution of known concentration until it has neutralised the one in the flask (can be seen using an indicator).
- Using the ratio of volumes required, the concentrations follow the same ratio so can work out the concentration of the unknown solution.
Electrolysis
General Stuff
This process is called electrolysis.
When ionic compounds are melted or dissolved in water, the ions become free to move around.
These liquids or solutions are able to conduct electricity and are called electrolytes.
Passing a current through the electrolyte causes the ions to move to the electrodes.
Positively charged ions move to the cathode, negative ions move to the anode.
In a binary ionic compound electrolysis, the negative ion will always be produced at the anode; the positive ion will always be produced at the cathode.
Electrolysis is used if the metal being extracted is too reactive for reduction by carbon.
Ionic Compounds in Aqueous Solution
Cathode
Hydrogen is produced if the metal is more reactive than hydrogen.
Otherwise the metal is always produced here.
Anode
Oxygen is produced unless halide ions are present in the solution.
Otherwise the halogen in solution will be discharged at the anode.
Reaction Types
Cathode
Always reduction reactions as ions are gaining electrons.
Anode
Always oxidation reactions as ions are gaining electrons.
Example Ionic Reactions:
2H + 2e- ---> H2
4OH- - 4e- ---> O2 + 2H2O
C5 - Energy Changes
Exo and Endothermic Reactions
Energy is conserved in chemical reactions. The total amount of energy in a closed system is the same as the total energy at the beginning of the reaction.
An exothermic reaction transfers energy to the surroundings.
Endothermic reactions draw energy from the surroundings.
Examples
Exothermic
Combustion
Oxidation
Neutralisation
Endothermic
Thermal Decomposition
Reaction of citric acid with sodium hydrogen carbonate
Reaction Profiles
Chemicals reactions can only occur when particles collide with sufficient force.
This requires a certain amount of energy - the activation energy of the reaction.
Reaction profiles show the relative energy levels of the reactants and products, as well as the activation energy.
Catalysts can reduce the activation energy as well as providing an alternate pathway for the reaction.
Bond Energy
Energy is required when bonds are broken.
Energy is released when bonds are formed.
The sum of these two determines whether a reaction is endo or exothermic.
If a reaction has an overall bond energy of less than 0, then it is an endothermic reaction. If it is positive, it is exothermic.
Chemical and Fuel Cells
Chemical Cells
Non-rechargeable have reactions that are not reversible.
Batteries consist of many cells joined together for a greater voltage.
Voltage produced by a cell is determined by many factors including the electrodes and the electrolyte.
Cells contain chemicals that react to produce electricity.
Fuel Cells
Fuel cells are supplied from an external source.
Fuel is oxidised electrochemically within the cell to produce a potential difference.
Hydrogen fuel cells include the oxidation of hydrogen to produce water and may be an alternative to other fuel cells in the future.
Required Practicals
Atomic Structure and the Periodic Table
Title: Analysis and purification of water samples from different sources.
Method
1) Use a pH indicator (universal indicator) to analyse the pH of the water sample.
2) Heat up a set volume of the water to 100 degrees centigrade so the water boils.
3) Collect the now pure water in a liebig condenser and have it output into a beaker.
4) When all the water has evaporated from the conical flask, measure the mass of the solid that remains behind to find the amount of dissolved solids in the sample (and thus its initial purity).
Hazards and Risks
Hot equipment (bunsen burner) so tie up hair and be careful around the flames and hot beakers.
Quantitative Chemistry
Title: Determination of the reacting volumes of a strong acid and alkali by titration.
Method
1) Wash and rinse a graduated pipette with the known alkali and use it to measure out an accurate amount of the known alkali solution into a conical flask.
2) Add a suitable indicator (e.g. phenolphthalein) and place flask on a white tile so the colour can be easily seen.
3) Place the acid into a burette that has been carefully washed and rinsed with the acid. Fill it up to the zero centimetre cubed mark.
4) Slowly add the acid and swirl until the indicator changes colour (just), this is called the end point. Take a note of the amount of acid used to neutralise the alkali.
5) By using the ratio of volumes of the acid and alkali, you can work out the ratio of the concentrations; working out the acid's concentration.
Hazards and Risks
Using corrosive chemicals so need to use eye and skin protection and clear up any spillages immediately.
Chemical Changes
Required Prac 1
Method
1) Add metal oxide/carbonate to a warm acid solution until no more will react.
2) Filter the excess metal oxide/carbonate off to leave a salt solution.
3) Gently warm the solution in a crucible until it starts to evaporate. Evaporate about 1/2 of the solution this way then leave it overnight to allow the rest to evaporate at room temp.
4) Harvest the new soluble salt crystals left behind.
Hazards and Risks
Corrosive materials/chemicals so need PPE such as safety glasses and clear up any spills instantly.
Hot water bath or a bunsen burner so don't have hair down and be careful around hot equipment.
Title: Preparation of a pure, dry sample of a soluble salt from an insoluble oxide or carbonate.
Required Prac 2
Title: Investigating what happens when an aqueous solution is electrolysed using inert electrodes.
Method
1) Set up two electrodes over a beaker with the aqueous solution in it.
2) Pass an electrical current through the electrodes (through the aqueous solution).
3) Observe what products are formed at each of the electrodes.
Hazards and Risks
A low voltage must be used to prevent a dangerous electric shock.
The room must be well ventilated as some of the products may be dangerous if exposed to for too long (such as chlorine gas). For the same reason, the experiment must only be conducted for a short space of time.
Energy Changes
Title: Investigate the variables that affect temperature changes in reacting solutions.
Method
1) Set up a beaker with HCl in it and a thermometer.
2) Add the metal powder and stir until the reaction is complete and record the highest temperature the reaction reached.
3) Calculate the temperature change for the reaction.
4) Repeat with different metals.
Considerations, Mistakes and Errors.
There should be a correlations between the reactivity of the metal and the temperature change.
There will always be a degree of uncertainty in the recording of the results - should repeat each reaction three times and take an average (excluding any anomalies).
Hazards and Risks
Low risk of corrosive acids so PPE such as safety glasses must be worn and spills immediately cleaned up.
Variables
The independent variable is the metal used in the reaction.
The dependant variable is the temperature change over the course of the reaction.
The control variables are: the type of acid, the concentration of the acid, the volume of the acid and the external temperature in the lab.