Chemistry (C2)

Annotations:

• Elements react to form compounds by gaining, sharing or losing electrons.
• Elements react to form compounds by gaining, sharing or losing electrons.
• The element is ALWAYS trying to get to a noble gas configuration.
• Atoms involved in chemical bonding are in the OUTER SHELL.

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• Held together by strong forces between the oppositely charged ions.
• -Gaining or losing electrons. -When a non-metal and metal react.

1. Characteristics

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   • -Giant structure (or latice) -High melting and boiling point. -Hard and Brittle. -Conduct electricity in water.
   1. High melting and boiling points

      Annotations:

      • High temperature required to break the positive and negative attraction.
   2. Conductors

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      • In water, dissociated ions are free to conduct an electric charge.
      • MOLTEN ionic compounds (salts) also conduct electricity.
2. Representation

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   • Ionic bonding can be shown by: dot and cross diagrams.
   1. Example

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• The charges on the ions in an ionic compound always cancel each other out.

1. Brackets

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   • Used when we need to multiply ions. An example of this is:  Ca(OH)₂ - Need two hydroxide ions.

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• Sharing electrons
• -Sharing electrons -When two non-metals react.
• Substances containing covalent bonds consist of simple molecules, but some have giant covalent structures.

1. Characteristics

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   • -Low melting and boiling points. -Don't conduct electricity well.
   1. Low melting and boiling points

      Annotations:

      • Weak intermolecular forces break down easily.
   2. Conductors

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      • They do not have free electrons or an overall electrical charge.
2. Representation

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   • We can use dots and crosses to show where the electrons are shared, on each atom.
   1. Example

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• The atoms in metals are: -CLOSELY packed together. -Arranged in REGULAR layers.

1. Higher tier

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   • -Electrons in the highest energy level are delocalized. -They can move FREELY. -Produces a lattice of positive ions in a 'sea' of moving electrons -The DELOCALISED ELECTRONS strongly attract the positive ions and hold the giant structure together.
2. Characteristics

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   • -Good conductors of heat and electricity. -Malleable -Ductile
   1. Conductors

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      • Free electrons carry heat or carry a charge throughout the metal.
   2. Malleable and ductile

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      • Free electrons allow metal atoms to slide over each other.

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• -Strong electrostatic forces hold ions together. -Solids at room temperature. -High melting and boiling points

1. Melting

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   • -MELTING = free electrons/ions -Allows them to carry electrical charge. -Therefore conduct electricity.
2. Dissolving

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   • -Some ionic structures dissolve in water. -Water molecules can split up the lattice. -Ions are free to move therefore conduct electricity.

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• -Held together by strong covalent bonds. -ONLY act between atoms of molecule. -Low melting and boiling points. -No electricity conduction.

1. Intermolecular forces (H)

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   • -Intermolecular forces are the forces of ATTRACTION. -Broken when boiled or melted. -Small molecules (H₂, CH₄) have weakest intermolecular forces, thus they are gases at RT. -Larger molecules (Br₂) have stronger attractions, so may be liquids at room temperature. -I₂ is a solid at room temperature.

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• -Atoms of some elements can form several covalent bonds. -Atoms can join together to make GCS's (macromolecules). -Every atom joined to several other atoms (STRONG bonds). -Therefore, VERY high melting point.

1. Diamond

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   • -Form of carbon that has a regular 3-D giant structure. -Each carbon atom covalently bonded to four others. -Makes diamond hard and transparent. -Silicon dioxide IS SIMILAR.
2. Graphite

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   • -Form of carbon (2-D layers) -One carbon atom is covalently bonded to three others. -No covalent bonds between layers, so they slide over each other. -Graphite is slippery and grey.
   • -One electron from each carbon atom is delocalised. -Allows graphite to conduct heat and electricity. -Weak intermolecular forces between layers in graphite, meaning layers can slide over each other easily.
3. Fullerenes

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   • -Large molecules from hexagonal rings of carbon. -Rings join together to form cage-like shapes with different numbers of carbon atoms (some are NANO-SIZED).

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• When we bend and shape metals, the layers of atoms in the giant metallic structure slide over each other.

1. Alloys

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   • -Mixtures of metals or metals mixed with other elements. -Different sized atoms distort layers in metal structure therefore difficult for them to slide. -Alloys are harder than pure metals.
   1. Shape memory alloys

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      • -Bent or deformed into a different shape. -When heated, they return to original shape. -Used for many things: e.g. dental braces.
2. Metal structures

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   • -Delocalised electrons (move throughout the giant metallic lattice and transfer energy quickly). -Good conductors of heat and electricity.

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• Depends on the monomers used to make them.  Could also depend on the reaction conditions.

1. Poly(ethene)

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   • Low density (LD) poly(ethene) and high density (HD) poly(ethene) are made using different catalysists and different reaction conditions.
   • Changing reaction conditions can also change the properties of the polymer that is produced.
2. Thermosoftening polymers

   Annotations:

   • -Poly(ethene) is an example. -Made up of individual polymer chains that are tangled together. -Heated - soft. -Cooled - hard. -Can be heated to mould into shapes and remoulded through heating again.
   • -Forces between polymer chains are weak. -When heated, these weak intermolecular forces are broken and polymer becomes soft. -When cooling down, intermolecular forces bring polymer molecules back together, thus hardening polymer.
3. Thermosetting polymers

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   • -Do not melt or soften when heated. -Polymers set hard when first moulded. -Due to strong covalent bonds forming cross-links between their polymer chains. -Strong bonds hold polymer chains in position.

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• The study of small particles that are between 1 and 100 nanometres in size.

1. Nanoparticles

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   • -1 nanometre is 10−⁹m. -Very small sizes give them: --BIG surface area. --New properties which make them very useful materials.
2. Nantechnology

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   • Uses nanoparticles to: --Highly selective sensors. --Very efficient catalysts. --New coatings. --New cosmetics. --Improvement in construction materials.
3. The Future

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   • -VERY EXCITING. -Unpredictable consequences for our health and the environment. -MORE research needed to find out effects.

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• Protons - 1 Neutrons - 1 Electrons - 0

1. Mass number

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   • Total number of protons and neutrons.
2. Atomic number

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   • (Or the proton number) The number of protons and electrons in an atom.
3. Isotopes

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   • Atoms of the same element with DIFFERENT numbers of NEUTRONS.

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• We use relative atomic masses to compare the masses of atoms.

1. Relative atomic mass

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   • Relative atomic mass (Ar) is an average value that depends on the isotopes the element contains.
2. Relative formula mass

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   • (Mr) Adding up all the relative atomic masses of the atoms in its formula.
3. Moles

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   • One mole of a substance is its relative formula mass in grams.

1. Percentage composition

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   • Divide the relative atomic mass of the element by the relative formula mass of the compound then multiply by 100.
   1. Example

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      • Percentage of C in CO₂: C = 12 , O = 16 Mr of CO₂ is: 12 + 32 = 44 So: 12/44 x 100 = 27.3%
2. Empirical formula

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   • Simplest ratio of atoms or ions in a compound.
   • To work out: -Mass in 100g -Mass / Ar -Divide by smallest (mass/Ar) -Convert to Empirical formula -REMEMBER: You can't have a decimal of an atom, so convert to integers.
   1. Example

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      • Empirical formula of the hydrocarbon that contains 80% C -Mass in 100g (C=80 , H=20) -Mass/Ar (C=80/12=6.67 , H=20/1=20) -Smallest ratio (C=6.67/6.67 = 1 , H=20/6.67=3) -Empirical formula is CH₃
3. Molecular formula

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   • -Empirical formula not always the same as molecular formula in covalent compounds. -e.g. Ethane's molecular formula is: C₂H₆ BUT its empirical formula is CH₃

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• To work out: -Calculate Mr for all compounds. -Work out how much the product weighs. -(The amount you have / mole amount) x Mr of thing you are trying to work out 

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• The yield of a chemical reaction describes how much product is made.
• The percentage yield tells us how much product is made compared with the maximum amount that could be made.

1. Formula

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   • Formula for Percentage Yield: (amount of product collected / maximum amount of product possible) x 100
2. Why?

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   • It isn't possible to collect the amounts calculated from the chemical equation because: -Incomplete reaction -Other reactions may happen. -Product may be lost when it is seperated or collected from the apparatus. 
3. High yields

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   • Good for industry as it: -Conserves resources -Reduces waste.
   • Chemical processes should also waste as little energy as possible so that: -There is a reduction in pollution. -It makes production more sustainable. 

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• In a reversible reaction, the products of the reaction can react to make the original reactants.
• Reversible reactions are shown by this sign: ⇌

1. Example

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   • Ammonium chloride ⇌ ammonia + hydrogen chloride

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• The process whereby small amounts of dissolved substances are separated by running a solvent along a material such as absorbent paper.

1. Paper Chromotography

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   • Used to identify food additives, such as artificial colours. 
   • The process: -A spot of colour is put onto paper and a solvent can move through it. -The colours move different distances depending on their solubility. 
2. Gas Chromotography

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   • Used to separate mixtures so that compounds can be identified.
   • The process: -Mixture is carried by a gas through a long column packed with particles of a solid. -Individual compounds travel at different speeds and come out at different times. -Amount of substance leaving column at different times is recorded against time and shows: --Number of compounds --Retention times
   1. Mass spectrometer

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      • This is connected to the gas chromatography column to give further data.
      1. Relative molecular masses

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         • A GC-MS can give the relative molecular mass of a compound. It can do this by reading off the molecular ion peak, which is furthest to the right on the mass spectrum. 

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• Measures the speed of a reaction or how fast it is.

1. Equation

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   • Rate of reaction = amount of reactant used / time  
   • Rate of reaction = amount of product formed / time
2. Gradient

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   • The gradient of the line on a graph of amount of reactant or product against time tells us the reaction at that time. The steeper the gradient. the faster the reaction.
   • The faster the rate, the shorter the time it takes for the reaction to finish. So rate is INVERSELY PROPORTIONAL to time. 

Annotations:

• -Reactions can only happen if particles collide. -Particles must collide with enough energy to create new substances. -This minimum energy is called the activation energy.
• IMPORTANT to learn each node word by word.

1. The effect of temperature

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   • Increasing the temperature increases the kinetic energy of particles therefore they will collide more frequently with enough activation energy and so the rate of reaction will increase.  
   • A rise of 10°C will roughly double the rate of many reactions, so they go twice as fast.
2. The effect of surface area

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   • The rate of a chemical reaction increases if the surface area of any solid reactant is increased. This increases the frequency of collisions between reacting particles.
3. The effect of catalysts

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   • -Catalysts change the rates of chemical reactions. -Catalysts that speed up reactions lower the activation energy required, therefore more successful collision occur.
   • Catalysts do not get used up during a chemical reaction.
   • Different catalysts are needed for different reactions.
4. The effect of concentration

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   • Increasing concentration increases the frequency of successful collisions (these are the ones with the activation energy). This therefore will increase the rate of reaction.
5. The effect of pressure

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   • Increasing pressure increases the frequency of collisions because particles are closer together since they occupy a smaller volume. This will increase the frequency of collisions with enough activation energy and so increases the rate of reaction. 

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• -Catalysts are used in industry to increase the rate of reactions and reduce energy costs. -This means that less fossil fuels are burned. therefore conserving resourcing and reducing pollution

1. Traditional catalysts

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   • Often transition metals or their compounds. However, some of these are toxic and may cause harm to the environment. 
2. Modern catalysts

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   • Being developed in industry which result in less waste and are safer for the environment. 

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• In a reaction, energy may be transferred to or from the reacting substances.

1. Exothermic

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   • Reactions that transfer energy to the surroundings.
   1. Examples

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      • -Combustion (burning fuels) -Oxidation reactions (respiration) -Neutralization reactions (involving acids and bases)
   2. Uses

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      • Used in hand warmers and self-heating cans.
2. Endothermic

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   • Reactions which take in energy from the surroundings.
   1. Examples

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      • Some cause a decrease in temperature and others require a supply of energy.
      • Solid compounds that are mixed with water cause temperature decrease because endothermic changes happen as they dissolve
      • Thermal decomposition reactions need to be heated continuously to keep the reaction going.
   2. Uses

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      • Used in instant cold packs for sports injuries.

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• In reversible reactions, the reaction in one direction is exothermic and in the other, endothermic.

1. Example

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   • Hydrated copper sulfate ⇌ anhydrous copper sulfate CuSO₄⋅5H₂O ⇌ CuSO₄ + 5H₂O
   • The hydrated copper sulfate must be heated continuously for the reaction to complete, because it is an endothermic reaction.
   • Adding water to the anhydrous copper sulfate causes the mixture to get hot, because it is an exothermic reaction.

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• Acids are substances that produce H⁺(aq) ions, when they are added to water.
• Alkalis are bases that dissolve in water to make alkaline solutions. They produce OH⁻ ions.

1. Bases

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   • React with acids and neutralizes them. 
2. pH scale

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   • Potential Hydrogen scale - it shows how acidic or alkaline a solution is. 

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• Acids will react with metals that are above hydrogen in the reactivity series.

1. Equation

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   • Acid + Base → Salt + Hydrogen
   • When this happens, a neutralisation reaction take place.
2. Crystals

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   • Salts can be crystalllised from solutions by evaporating off water.
3. Examples

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   • HCl → Chlorides HNO₃ → Nitrates H₂SO₄ → Sulfates

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• We can make soluble salts by: Acid + Alkali → Salt + Water

1. Neutralisation

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   • H⁺(aq) + OH⁻(aq) → H₂I(l)
2. Insoluble salts

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   • Mixing solutions of soluble salts that contain the ions needed, in order to form a precipitate (insoluble salt).
3. Precipitation

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   • An important way of removing some metal ions from industrial waste water.

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• -Electricity is used to break down ionic compounds into elements. -When electricity is passed through molten ionic compounds or a solution containing ions, electrolysis takes place. -Substance broken down is an electrolyte.

1. Electrodes

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   • -Made of an inert substance that doesn't react with the products. -Ions which move to either electrode are discharged to produce elements.
   1. Anode

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      • Negatively charged ions are attracted to the anode, losing their charge and forming non-metallic elements.
   2. Cathode

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      • Positively charged ions are attracted to the cathode, losing their charge and forming either hydrogen or metals, depending on the electrolyte. 

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• When an ion reaches an electrode, they lose or gain electrons to become neutral atoms. 

1. OILRIG

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   • Oxidation  Is  Loss Reduction Is Gain
   1. Oxidation

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      • Negative ions lose electrons to become neutral atoms, with some non-metals forming molecules. 
   2. Reduction

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      • Positive ions gain electrons.
2. Half equations

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   • Anode: 2Br⁻ → Br₂ + 2e⁻ Cathode: Pb²⁺ + 2e⁻ → Pb Anode - Losing electrons. Cathode - Gaining electrons.
3. Aqueous solutions

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   • Water contains hydrogen and hydroxide ions.
   1. Negative electrode

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      • Hydrogen may be produced at the negative electrode if the positive ions in the solution are those of a metal more reactive than hydrogen. 
   2. Positive electrode

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      • Oxygen is usually produced. If the solution has a high concentration of a halide ion, then a halogen will be produced.

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• Aluminium oxide is electrolysed, because aluminium is more reactive than carbon. We do this to manufacture aluminium.

1. Cryolite

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   • Added to aluminium oxide to reduce the melting temperature (from 2000°C to 850°C)
2. Electrodes

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   • Different products are formed at either electrode.
   1. Anode

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      • Oxide ions are oxidised to oxygen atoms by losing electrons and these oxygen atoms form oxygen molecules.
      1. Substances

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         • -Anode is made of carbon. -At high temperatures, the oxygen reacts with the carbon to form carbon dioxide. -All the carbon electrodes gradually burn away, therefore they must be replaced regularly. 
   2. Cathode

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      • Aluminium ions are reduced to aluminium atoms by gaining electrons. The molten aluminium is collected from the bottom of the cell.

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• Brine is a solution of sodium chloride in water.

1. Brine ions

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   • Brine contains: -Na⁺ sodium ions -H⁺ hydrogen ions -Cl⁻ chloride ions -OH⁻ hydroxide ions
2. Electrodes
   1. Anode

      Annotations:

      • Chlorine is produced from the chlorine ions. 2Cl⁻ → Cl₂ + 2e⁻ 
   2. Cathode

      Annotations:

      • Hydrogen ions. 2H⁺ + 2e⁻ → H₂
   3. Leaving

      Annotations:

      • A solution of sodium ions and hydroxide ions.
3. Products

   Annotations:

   • Sodium hydroxide is a strong alkali and has many uses: -Soap -Paper -Bleach -Neutralising acids -Controlling pH
   • Chlorine is used to: -Kill bacteria in drinking water and swimming pools. -Make bleach. -Make disinfectant. -Make plastics.
   • Hydrogen is used to: -Make margarine. -Make hydrochloric acid.

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• Uses electrolysis to put a thin coating of metal onto an object. 

1. Uses

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   • Electroplating is used to: -Make the object look more attractive. -Protect a metal object from corroding. -Increase the hardness of a surface. -Reduce costs by using a thin layer rather than pure metal.
2. Process

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   • Object to be electroplated is made the negative electrode. 
   1. Electrolyte

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      • A solution containing ions of the plating metal. 
   2. Electrodes
      1. Anode

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         • -Made from the plating metal. -Atoms of the plating metal lose electrons to form metal ions which go into the solution.
      2. Cathode

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         • Metal ions from the solution gain electrons to form metal atoms which are deposited on the the object to be plated.