AS-Level Atomic Structure
Structure of an Atom
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Atoms are the smallest particles of an element.
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Atoms are made of sub-atomic particles:
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protons (charge +1, relative mass 1)
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neutrons (no charge, relative mass 1)
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electrons (charge -1, relative mass 1/1840)
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The number of protons an atom has is called its ‘atomic number’.
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The centre of an atom is called a nucleus and contains just protons and neutrons.
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Electrons are attracted to the positive charge of a nucleus and exist in orbitals.
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Atoms can lose or gain electrons to become charged ions.
Electron Orbitals
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Electrons are attracted to the positive charge of a nucleus (electrostatic attraction).
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The closer electrons are to a nucleus, the lower their energy.
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Electrons want to get close to a nucleus but they repel each other.
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They can only exist in set regions called orbitals.
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Different orbitals have different shapes and therefore different energies.
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Sub-shells describe how far away an electron (and its orbital) is from a nucleus.
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Electron sub-shells are filled successively – lowest energy sub-shells and orbitals are filled first.
The Mole
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Atoms and particles are very small, so referring to actual numbers of atoms in a sample would require very large numbers that would be hard to work with.
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The mole is just a unit used to refer to the amount of atoms or particles in a substance.
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Just like one dozen = 12 of something, one mole = 6.022x10 of something.
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The number one mole refers to is known as Avogadro’s constant.
Relative Atomic Mass
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The mass of an atom is determined by the number of protons and neutrons it has.
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Different elements have different masses.
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An element’s ‘relative atomic mass’ is the mass (in grams) of one mole’s worth of its atoms.
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One mole of carbon-12 (12C) has a mass of 12g.
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Each atom of carbon-12 has 6 protons and 6 neutrons.
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The mass of each must be 1g per mole.
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Therefore, atomic masses of elements are described as being ‘relative to 1/12 the mass of an atom of carbon-12’.
Moles by Mass
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The number of moles in a sample can be calculated using the formula:
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number of moles = mass (g) / molar mass
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The molar mass (relative molecular mass or relative formula mass) of a compound is calculated by adding up all the relative atomic masses of the atoms within the compound’s formula.
Moles in a Solution
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The number of moles of a solute in a solution can be calculated using the formula:
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number of moles = concentration (mol dm ) x volume (dm ) / 1000
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1 dm is the same as 1000 cm
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To convert dm to cm x 1000
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To convert cm to dm / 1000
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Concentration refers to how many moles of solute are in 1dm of solution (units are mol dm )
Moles of a Gas
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The ideal gas equation is used to calculate the moles of gas present in a system if the volume (m ), pressure (Pa) and temperature (K) are known.
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pV = nRT
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n = moles
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R = gas constant, 8.31 Jmol K
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The equation assumes that the same number of moles of all gases occupies the same volume at a set temperature and pressure.
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At room temperature (293K and 101KPa), it is assumed that 1 mole of any gas occupies a total volume of 24000cm (24dm ).
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AS-Level Bonding
Bonding
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Atoms can bond together to create larger species.
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Bonds that hold atoms together are called atomic bonds.
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Forces between molecules are called intermolecular forces.
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Intermolecular forces are broken and formed when molecular substances change state (solid, liquid and gas)
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Covalent Bonds
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Atoms are most stable when they have a ‘full’ outer shell of electron (two electrons for hydrogen, eight electrons for other elements).
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To have access to full outer shells, atoms ‘share’ electrons between orbitals.
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A covalent bond is the sharing of one pair of electrons between two atoms.
Ionic Bonding
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To get a full outer shell, atoms lose or gain electrons.
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Metals lose electrons to get a full outer shell and become positive ions.
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Non-metals (usually) gain electrons to get a full outer shell and become negative ions.
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Ionic bonding refers to the attraction between oppositely charged ions.
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The greater the charge of the ions, the greater the attraction and the stronger the ionic bonding between the ions.
Metallic Bonding
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Metal atoms are more stable if they lose electrons to have a full outer shell.
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Electrons lost by metal atoms in solid metal have nowhere to go and stay attracted to the positive charge of the metal ions.
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Electrons are not held by the metal atom but are free to move - they are delocalised.
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In metallic bonding, positive metal ions are attracted to the negative charges of the delocalised electrons and this gives metals a dense, rigid structure.
Intermolecular Forces
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Intermolecular forces exist between molecules
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When molecular substances change state, it is the arrangement of intermolecular forces that change (not the covalent bonds inside the molecules)
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Temporary induced dipole-dipole forces (Van der Waals, London forces, Dispersion forces) exist between all molecules
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Unequal electron distribution around a molecule creates an instantenous dipole that can induce a dipole on a neighbouring molecule.
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The two opposite dipoles from two molecules are attracted to each other and a weak force occurs
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They are the weakest intermolecular forces
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Permanent Dipole-Dipole
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If a molecular substances is polar (has polar bonds and an overall uneven distribution of electron density) permanent dipole-dipole interactions exist between molecules
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The slightly negative part of one molecule is attracted to the slightly positive part of another molecule.
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The attractions are permanent and are stronger than temporary, induced dipole-dipole forces
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The greater the polarity of a bond or molecule, the stronger the permanent dipole-dipole interactions
Hydrogen Bonding
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Hydrogen bonds are a type of permanent dipole-dipole force that are very strong
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They exist between molecules that have a N-H, O-H or F-H bond
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Nitrogen, Oxygen and Fluorine are the most electronegative elements and the covalent bonds between them and hydrogen are highly polar
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Highly polar bonds result in strong attractive forces between molecules (the slightly negative N, O or F forms hydrogen bonds with slightly positive hydrogen atoms from other molecules)
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Substances that can form hydrogen bonds have higher melting points than those unable to
AS-Level Energy and Enthalpy
Enthalpy Changes
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During chemical reactions, energy is exchanged with the surroundings.
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In exothermic reactions, products are lower energy than the reactants and the temperature of the surroundings increases as energy is released overall.
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In endothermic reactions, products are higher energy than the reactants and the temperature of the surroundings decreases as energy is absorbed overall.
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The actual amount of energy in a compound or atom is not measured in chemistry but the change in energy during a reaction can be measured and is called enthalpy.
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Standard enthalpy changes refer to the thermal energy change that occurs during a reaction, where 1 mole of product is formed in standard conditions (100 KPa and 298K).
Bond Enthalpies
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The energy levels of two individual atoms lower when they bond together.
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A certain amount of energy is released by two atoms when a bond is formed, and that same amount of energy is absorbed by the two atoms when the same bond is broken.
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The amounts of energy released or absorbed for a particular bond is called its bond enthalpy.
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Mean bond enthalpies are used to account for the different environments a bond may be in.
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The overall enthalpy change that occurs in a reaction can be found by subtracting the sum of all bond enthalpies formed from the sum of all bond enthalpies broken.
Calorimetry
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The energy change that occurs during a reaction can be found by measuring how the temperature of the surroundings changes during a reaction, this is called calorimetry.
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Specific heat capacity refers to the energy required to raise the temperature of 1kg of a substance by 1°C.
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The temperature change that a reaction forces on a substance can be used with the specific heat capacity to find the energy that has been exchanged.
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Change in energy = mass x specific heat capacity x change in temperature
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Q = mc∆T
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Energy is often lost to the surroundings from the substance being used to find the temperature change (for example water).
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To increase the accuracy of calorimetry, heat loss must be minimised (extra insulation can help).
Hess' Law
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There is often more than one way to turn a substance into a product.
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The energy change that occurs when a substance is changed into a product is the same, regardless of the route taken. This is called Hess’ Law.
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Hess’ Cycles are ways of drawing out the energy changes that occur when a substance is turned into a product via more than one route.
AS-Level Kinetics
Collision Theory
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Kinetics is the study of how fast reactions occur.
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How quickly a reaction happens is called its reaction rate.
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Reactions occur when reactant particles collide together with enough energy (activation energy), creating a successful collision.
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Increasing the frequency of successful collisions increases the rate of reaction.
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Increasing the temperature, concentration and pressure of reacting particles increases the rate of a reaction.
Boltzmann Distribution Curves
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Boltzmann distribution curves show the number of particles with a given amount of energy in a system.
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Effects of changing concentration, temperature and the use of a catalyst on the energies of particles can be shown.
Equilibrium
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In a reversible reaction, reactants react together to produce products (forward reaction) and those products also react to produce the original reactants (reverse reaction).
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At equilibrium, the rates of both the forward and reverse reactions are the same – the concentrations of all reactants and products do not change.
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The position of equilibrium describes how much one reaction is favoured compared to the other.
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Changing the conditions of an equilibrium system causes the position of equilibrium to move to oppose this change (Le Chatelier’s Principle).
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Increasing the temperature of a system forces the position of equilibrium to move to favor the endothermic reaction.
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Increasing the concentration of a reactant forces the position of equilibrium to move to favor the reaction that uses up the reactant.
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Increasing the pressure of a system forces the position of equilibrium to move to favor the reaction that produces the fewest moles of gas.