OCR A Level Inorganic Chemistry

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Standard enthalpy of formation
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The enthalpy change when one mole of a compound is formed from its elements under standard conditions, with all substances in their standard states.
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Standard change of atomisation
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The enthalpy change when one mole of gaseous atoms is formed from the element in its standard state under standard conditions.
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First ionisation energy
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The enthalpy change needed to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.
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Second ionisation energy
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The enthalpy change required to remove one electron from each ion in one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions.
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First electron affinity
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The enthalpy change when one electron is added to each atom in one mole of gaseous atoms to form one mole of gaseous 1- ions.
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Second electron affinity
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The enthalpy change when one electron is added to each ion in one mole of gaseous 1- ions to form one mole of gaseous 2- ions.
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Lattice enthalpy
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The enthalpy change that accompanies the formation of one mole of a solid compound from its gaseous ions under standard conditions.
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The standard enthalpy change of solution
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The enthalpy change when one mole of a compound is completely dissolved in water under standard conditions.
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The standard enthalpy change of hydration
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The enthalpy change that accompanies the dissolving of gaseous ions in water to form one mole of aqueous ions.
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Entropy
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A measure of the dispersal of energy in a system, which is greater the more disordered a system.
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ΔS=
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S of products – S of reactants
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Effect of increasing ionic size on lattice enthalpy
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This leads to a weaker attraction between oppositely charged ions so there is a less exothermic lattice enthalpy.
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Effect of increasing ionic charge on lattice enthalpy
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This leads to a stronger attraction between oppositely charged ions and a more exothermic lattice enthalpy.
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Effect of increasing ionic size on hydration enthalpy
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This leads to a weaker attraction between the ions and water molecules so there is a less exothermic hydration enthalpy.
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Effect of increasing ionic charge on hydration enthalpy
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This leads to a stronger attraction between ions and water molecules so there is a more exothermic hydration enthalpy.
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ΔG=
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ΔH-TΔS
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A process can take place spontaneously (i.e. is feasible), when…
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ΔG<0
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Free energy change
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The overall change in energy during a reaction.
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A reaction is feasible at any temperature if…
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The sign of its ΔH is negative and its ΔS is positive, because ΔG will always be less than zero.
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A reaction is NOT feasible at any temperature if…
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The sign of its ΔH is positive and its ΔS is negative, because ΔG will always be more than zero.
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A reaction is only feasible at high temperatures if…
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The sign of its ΔH is positive and its ΔS is positive as increasing the temperature will increase the magnitude of TΔS, until the magnitude of TΔS becomes greater than the magnitude of ΔH, so ΔG will become less than zero and the reaction will become feasible.
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A reaction is only feasible at low temperatures if…
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The sign of its ΔH is negative and its ΔS is negative as decreasing the temperature will decrease the magnitude of TΔS, until the magnitude of TΔS becomes less than the magnitude of ΔH, so ΔG will become less than zero and the reaction will become feasible.
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Why do reactions with ΔG<0 sometimes not take place?
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The rate of reaction is too slow or the activation energy is too high.
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Proton relative mass and charge
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1 and 1+
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Neutron relative mass and charge
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1 and 0
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Electron relative mass and charge
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1/1836 and 1-
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Mass number=
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Number of protons + number of neutrons
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Atomic number=
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Number of protons
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Isotopes
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Atoms of the same element (with the same number of protons) but with different numbers of neutrons and different masses.
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Relative isotopic mass
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The mass of an isotope compared with one twelfth of the mass of an atom of carbon-12.
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Relative atomic mass
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The weighted mean mass of an atom compared with one twelfth of the mass of an atom of carbon-12.
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Relative molecular mass
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The weighted mean mass of a molecule compared with one twelfth of the mass of an atom of carbon-12.
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Water of crystallisation
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The water present in hydrated salts.
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To determine relative atomic mass…
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Multiply the isotopic mass by percentage abundance for each isotope then sum the results.
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Avagadro’s number
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The number of atoms in one mole of the substance. 6.02×1023
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Mole
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The mass of substance containing the same number of atoms as there are in 12g of carbon-12.
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Empirical formula
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The simplest whole-number ratio of atoms in a compound.
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Atom economy=
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(Molar mass of desired product/Molecular mass of all products) x100
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Hydrated salts
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Contain some water as part of their structure.
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Anhydrous salts
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Hydrated salts after being heated so the water has been lost.
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Molar volume
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The volume occupied by one mole of a substance at a given temperature and pressure.
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Number of moles of gas at room temperature and pressure=
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V(dm3)/24
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Ideal gas equation
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PV=nRT P is in pascals V is in m3 n is in mol T is in Kelvin
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Concentration of a solution
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The amount of solute per dm3.
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Standard solution
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A solution of a known concentration.
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Concentration=
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n/V
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Stoichiometry
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The relationship between relative quantities of substances taking part in a reaction.
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Limiting reagent
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The reactant that is first to be used up so stops the reaction.
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Apparatus uncertainty=
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(Number of readings x instrumental error/reading) x 100
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Base
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A proton acceptor.
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Acid
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A proton donor.
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Strong acid
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An acid that completely dissociates in aqueous solution.
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Weak acid
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An acid that partially dissociates in aqueous solution.
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Neutralisation equation
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H+(aq) + OH-(aq) –> H2O(l)
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Acid + base –>
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Salt + Water
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Acid + carbonate –>
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Salt + water +carbon dioxide
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Acid + alkali –>
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Salt + water
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Method of preparation of a standard solution
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1. Weight the solid accurately. 2. Dissolve the solid in a beaker. 3. Transfer solution to a volumetric flask. Rinse traces of solution from the beaker into the flask. 4. Carefully fill the volumetric flask to the graduation line (read the meniscus at the bottom). 5. Invert the flask several times to ensure the solution is well mixed.
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Method of an acid-base titration
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1. Using a pipette, we measure a volume of solution into a conical flask. 2. The other solution solution is placed in a burette and the volume recorded. 3. A few drops of indicator are added to the solution in the conical flask. 4. The solution in the burette is then added to the solution in the conical flask until the reaction has just completed. This is called the end point and is when we see a permanent change in colour. The volume in the burette is then recorded. 5. The volume of solution from the burette is then calculated and this is known as the titre. 6. The experiment is repeated until two titres are concordant to 0.10cm3 of each other. 7. The mean titre is then calculated (to 1 d.p.). Only concordant titres are used in the mean.
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Atomic orbital
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A region within an atom that can hold up to 2 electrons, with opposite spins.
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The 4s sub-shell fills before…
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The 3D sub-shell.
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Ionic bonding
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The electrostatic attraction between oppositely charged ions.
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Why do ionic compounds have high melting and boiling points?
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Lots of energy is required to break the strong electrostatic forces between oppositely charged ions.
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Explain why molten or dissolved ionic compounds are able to conduct electricity.
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The ions are free to move throughout the substance and carry the charge.
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Ionic lattices dissolve in…
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Polar solvents, such as water.
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Why are some ionic substances less soluble than others?
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If a compound has very big ions then not many water molecules can fit round them. Also, strong charges are difficult to overcome so 3+ or 3- ionic compounds do not dissolve very well.
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Metallic bonding
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The strong electrostatic attraction between positive Mel=tal ions and delocalised electrons.
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Delocalised electrons
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Electrons that are not associated with a particular atom and are free to move throughout the structure.
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Why do metallic compounds have high melting and boiling points?
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The strong electrostatic forces between positive metal ions and delocalised electrons need a lot of energy to be overcome.
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Why are metallic compounds able to conduct electricity as solids and liquids?
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The delocalised electrons are free to move throughout the structure and can pass on the electrical current.
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Why do metallic compounds have poor solubility?
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The electrostatic attraction is too strong to be broken by polar solvents surrounding each ion in solution.
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Covalent bonding
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The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.
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Lone pair
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An outer shell pair of electrons that are not involved in chemical bonding.
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Dative (co-ordinate) covalent bonds
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This is when one of the atoms supplies both of the electrons that form the bond.
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Bond enthalpy
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A measure of the energy required to break the bond. A larger enthalpy means a stronger bond.
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Lone pairs repel…
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More than bond pairs.
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Electronegativity
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A measure of the attraction of a bonded atom for the pair of electrons in a covalent bond.
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Permanent dipole
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A small charge difference that doesn’t change across a bond.
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How are induced dipole-dipole forces formed?
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The movement of electrons in shells causes an uneven distribution of charge this results in an instantaneous dipole. This induces a dipole in a neighbouring molecule and small induced dipoles attract one another.
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Explain why ice is less dense than water.
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Water expands when it freezes because the hydrogen bonds become rigid forming a fixed open lattice with space in between making it less dense.
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Explain why water has higher than expected melting and boiling points
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Hydrogen bonds are the strongest intermolecular forces so more energy is ended to break the forces between water molecules.
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Diamond
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Has a tetrahedral structure held together by strong covalent bonds throughout the lattice.
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Graphite
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Has a structure composed of hexagonal layers, each layer is held by weak intermolecular forces. It has delocalised electrons, so is a good conductor of electricity.
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Graphene
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A single layer of graphite, made of hexagonally arranged carbon atoms. Each carbon atom has 3 strong covalent bonds to 3 other carbon atoms. It has delocalised electrons so is a good conductor of electricity.
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Simple molecules
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Molecules that contain only a few atoms held together by strong covalent bonds (e.g. Iodine or Bromine)
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Why do simple molecules have low melting and boiling points?
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There are only London forces between molecules which are weak and don’t need much energy to be overcome.
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Non-polar molecules are soluble in…
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Non-polar solvents.
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Polar molecules are soluble in…
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Polar solvents.
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Oxidation
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Loss of electrons and increase in oxidation number.
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Reduction
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Gain of electrons and decrease in oxidation number.
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Redox reaction
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A reaction involving both reduction and oxidation.
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Reducing agent
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A reagent that reduces (adds electrons to) another species.
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Oxidising agent
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A reagent that oxidises (takes electrons from) another species.
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What are the factors affecting ionisation energy?
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Atomic radius, electron shielding and nuclear charge.
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Electron shielding
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The repulsion between inner shell electrons and outer shell electrons which reduces the attraction between the nucleus and outer electrons.
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Trends in ionisation energy down a group
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The first ionisation energies decrease since electron shielding and atomic radii increase and thus there is a decrease in attraction between outer electrons and the nucleus.
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Trends in ionisation energy across a period
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The first ionisation energy increases as nuclear charge increases while atomic radii decreases and electron shielding remains the same and thus there is an increase in attraction between outer electrons and the nucleus.
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Exceptions to the rule of ionisation energy
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1. Slight drop between group 12 and 13 as the outermost electron in group 12 is in the 2s sub-shell while the outermost electron in group 13 is in the 2p sub-shell. The 2p sub-shell is in a higher energy level so it is easier to remove and ionisation energy is lower. 2. Slight drop between groups 15 and 16. Outermost electron in both groups is in 2p sub-shell but it is unpaired in 15 and paired in 16. The pairing in group 16 causes repulsion so the outermost electron is easier to remove and ionisation energy is lower.
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Trend in boiling points down group 17
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Boiling point increases because the number of electrons increases so there are more London forces and more energy is needed to break them.
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Reaction of Cl2 or Br2 with I- observation
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Orange/brown solution OR purple in organic.
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Reaction of Cl2 with Br- observation
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Yellow solution OR orange in organic.
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Disproportionation
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Where an element is both oxidised and reduced in the same redox reaction.
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Chloride observation with AgNO3
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White precipitate forms.
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Chloride observation with NH3
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White precipitate dissolves on the addition of either dilute or concentrated NH3.
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Bromide observation with AgNO3
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Cream precipitate forms.
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Bromide observation with NH3
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Cream precipitate doesn’t dissolve with dilute NH3 but does dissolve with concentrated NH3.
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Iodide observation with AgNO3
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Yellow precipitate forms.
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Iodide observation with NH3
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Yellow precipitate doesn’t dissolve with either dilute or concentrated NH3.
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Cl2 + H2O –>
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HClO + HCl
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Equation for making bleach
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Cl2 + 2NaOH —> NaCl + NaClO + H2O
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Down group 2…
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Solubility and pH increases.
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Test for carbonates
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Add nitric acid and bubble the gas given off through lime water. The lime water should go cloudy if a carbonate is present.
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Test for sulfates
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Add barium nitrate/ carbonate/ chloride and if a sulfate is present then BaSO4 will form which is a white precipitate.
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Test for ammonium
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Add aqueous sodium hydroxide and heat the solution to release the soluble ammonia gas produced. Test the gas with moist litmus paper which will turn blue if ammonium was present.
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What order do we carry out the qualitative analysis tests?
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Carbonates Sulfates Halides
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Standard electrode potential
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The tendency to be reduced and gain electrons compared with a standard hydrogen half-cell at 298K, 100KPa and with concentrations at 1.00 mol dm3.
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E°Cell=
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E° most positive – E° most negative
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The more negative the E°Cell, the…
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Greater the tendency to release electrons.
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When using E°Cell values to predict feasible reactions, the most negative half equation always goes in…
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Reverse.
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Non-rechargeable cells (primary cells) provide electrical energy until…
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The chemicals are used up.
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Rechargeable cells (secondary cells) contain chemicals that react which can be reversed by…
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Recharging (e.g. Lithium ion or lithium-polymer batteries used in laptops.
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Fuel cells continue to provide electricity provided there is a…
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Continuous supply of fuel and oxygen.
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What is the overall cell reaction equation for fuel cells?
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H2 (g) + 1/2 02 (g) —> H2O (l)
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Transition element
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A d-block element that forms an ion with an incomplete d sub-shell.
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Which d block elements are not transition elements?
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Scandium and zinc.
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Electron configurations of chromium and copper.
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Chromium: [Ar]4s13d5 (to minimise repulsion) Copper: [Ar]4s13d10
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Chemical properties of transition elements.
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1. Variable oxidation states. 2. The have coloured compounds as they have coloured ions. 3. They act as catalysts.
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Transition metal catalyst in the Haber process.
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Iron
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Transition metal catalyst in the contact process.
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Vanadium oxide (V2O5)
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Transition metal catalyst in the reaction of zinc and acids.
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Cu2+ ions
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Transition metal catalyst in hydrogenation.
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Nickel
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Transition metal catalyst in the decomposition of hydrogen peroxide.
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MnO2
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Complex ion
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A transition metal ion bonded to one or more ligand by coordinate bonds.
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Ligand
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A molecule or ion that can donate a pair of electrons to the transition metal.
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Coordination number
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The total number of coordinate bonds formed between a central metal ion and ligands.
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Complexes containing nickel, platinum, gold or silver form…
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Square planar complexes.
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Describe the use of cis-platin as an anti-cancer drug.
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Cis-platin binds to the cells’ DNA and prevents replication of cancer cells.
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Bidentate ligand
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The ligand has 2 atoms which each have a lone pair of electrons that are donated to a central metal ion so it forms two coordinate bonds.
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Optical isomers
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Non-superimposable mirror images of each other.
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Observation: Cu2+ + 2OH- —> Cu(OH)2
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Pale blue solution forms pale blue precipitate.
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Observation: Cr3+ + 3OH- —> Cr(OH)3
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Purple solution forms grey green precipitate.
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Observation: Fe2+ + 2OH- —> Fe(OH)2
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Pale green solution forms green precipitate that turns rusty brown in air.
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Observation: Fe3+ + 3OH- —> Fe(OH)3
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Pale yellow solution forms rusty brown precipitate.
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Observation: Mn2+ + 2OH- —> Mn(OH)2
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Pale pink solution forms light brown precipitate that darkens in air.
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Cr(OH)3 is soluble in excess sodium hydroxide and forms a…
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Dark green solution.
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What is the observation when excess ammonia is added to [Cu(H2O)6]2+?
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Dark blue solution forms.
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What is the observation when excess ammonia is added to [Cr(H2O)6]3+?
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Purple solution forms.
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Ligand substitution reaction
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When one ligand in a complex ion is replaced by another ligand.
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Observation when Cu2+ and concentrated HCl are reacted.
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Pale blue solution turns to a green solution and finally a yellow solution.
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Explain, in terms of ligand substitution, why carbon monoxide is so dangerous to humans.
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Carbon monoxide binds to the same site as oxygen would but forms a stable complex so doesn’t break away. This makes the haemoglobin unable to transfer anymore oxygen.
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Buffer
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A system that minimises pH changed when small amounts of an acid or a base are added.
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Equivalence point
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The exact volume of solution required to neutralise a given volume of the other solution.
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Enthalpy of neutralisation
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The enthalpy change that occurs when an acid and a base react to form one mole of water.
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Half-life
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The time taken for the concentration of a reactant to reduce by half.
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Gradient
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-Ea/R
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Y intercept
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ln A
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Ea (activation energy)
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-gradient x R
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A (pre-exponential factor)
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e to the y intercept
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Rate of reaction
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The decrease in concentration of reactants in a given unit of time.
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Calculating gradient from a graph
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Difference in y/ difference in x
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Experimental techniques used to measure rate of reaction
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1. Monitoring by gas collection 2. Monitoring by mass loss 3. Monitoring with a colorimeter 4. Iodine clocks
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Overall order of a reaction
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The overall effect of the concentrations of all reactants on the rate of reaction.
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Rate constant
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The number that mathematically converts between the rate of reaction, concentration and orders.
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Rate=
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k[A][B]
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Zero order rate-concentration graph
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Rate is unaffected by changes in concentration (straight horizontal line).
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First order rate-concentration graph
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Rate is directly proportional to the reactant concentration (straight diagonal line).
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Second order rate-concentration graph
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Rate is proportional to the reactant concentration squared (curved line getting steeper).
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Half-life of first order reactants are..
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Constant.
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Zero order concentration-time graph
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Concentration decreases at a constant rate (straight diagonal line).
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First order concentration-time graph
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Concentration halves in regular time intervals (curved line decreasing).
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Second order concentration-time graph
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Concentration decreases rapidly, but the rate of decrease then slows down (curved line steeply decreasing).
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Reaction mechanism
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The series of steps that make up an overall reaction.
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Rate determining step
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The slowest step in the sequence of a multi-step reaction.
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Initial rate is approximately proportional to
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1/t
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Rate constant of a first order reaction calculation
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k=ln2/half-life
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pH=
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-log[H+]
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Ka=
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[H+][A-]/[HA]
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Weak acids: [H+]=
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Square root of [HA]xKa
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pKa=
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-log(Ka)
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Kw=
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[H+][OH-]
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Buffer solutions: [H+]=
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Ka x [HA]/[A-]
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How does a buffer solution work when an acid is added?
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[H+] increases. The H+ ions react with the conjugate base, A-. The equilibrium position shifts to the left, removing most of the H+ ions.
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How does a buffer solution work when a base is added?
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[OH-] increases. The small concentration of H+ reacts with the OH- ions. H+ + OH- —> H2O. HA dissociates, shifting the equilibrium to the right to restore most of the H+ ions.
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Carbonic acid-hydrogencarbonate buffer system
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H2CO3 —> H+ + HCO3-
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Kc=
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[products]/[reactants]
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Dynamic equilibrium
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Exists in a closed system when the rate of the forward reaction is equal to the rate of the reverse reaction and concentrations do not change.
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Homogenous equilibrium
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An equilibrium with species that all have the same state.
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Heterogeneous equilibrium
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An equilibrium with species that have different states.
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Solid or liquid
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When you write the Kc for a heterogeneous reaction, don’t include anything that is…
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Mole fraction=
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Number of moles of gas A/total number of moles in gas mixture
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Partial pressure=
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Mole fraction x total pressure
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Kp=
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p(A) x p(B)/ p(C)
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Bigger
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If more product is formed, K is…
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Smaller
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If more reactants are formed, K is…
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Increase
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If the reaction is endothermic, a rise in temperature will make K…
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Decrease
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If the reaction is exothermic, a rise in temperature will make K…
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Le Chatelier’s principle
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If a system in dynamic equilibrium is subject to change, the position of equilibrium will shift to minimise the effect of this change.
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Kc can only be affected by
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Temperature
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Ozone depletion
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Cl + O3 —> ClO + O2 ClO + O —> Cl + O2 O3 + O —> 2O2

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