A2-Level Energy and Enthalpy Change

Lattice Enthalpies

  • Solid ionic lattices are held together by strong attraction between oppositely charged ions.

  • The stronger the attraction between ions, the stronger the lattice (higher melting point)

  • Ion charge and size influence the strength of ionic attraction.

  • When two opposite charges come close together, their energies lower and energy is released.

  • Lattice enthalpy of formation is the enthalpy change when 1 mole’s worth of an ionic compound is formed from gaseous ions.

    • Lattice formation enthalpies are always negative (exothermic process).

  • Lattice enthalpy of dissociation is the enthalpy change to break apart 1 mole’s worth of an ionic compound into gaseous ions.

    • Lattice dissociation enthalpies are always positive (endothermic process).
       

Born-Haber Cycles

  • It is very difficult to find lattice enthalpies experimentally.

  • Born-Haber energy diagrams are used to find a lattice enthalpy by using Hess’ Law.

  • Hess’ Law states that the overall energy change that occurs when a product is formed from its elements is the same, regardless of the route taken to form the product.

  • Different standard reactions are used to construct a Born-Haber cycle

    • Standard enthalpy of atomisation (forming gaseous atoms)

    • Ionisation energies (removing electrons from gaseous atoms to form positive gaseous ions)

    • Electron affinities (addition of electrons to gaseous atoms to form negative gaseous ions)

    • Standard enthalpy change of formation (enthalpy change when one mole of a compound is formed from constituent elements in standard states)

Entropy

  • Enthalpy change and entropy change determines how likely a reaction is to happen.

  • Enthalpy change is the change in thermal energy that occurs during a reaction.

  • Entropy change refers to the change in disorder of a system that occurs during a reaction.

  • Entropy is a measure of disorder within in a chemical system.

    • If there a more ways of arranging particles in the system, it has a higher entropy than if there are fewer ways of arranging the particles.

  • Solids have low entropies and gases have high entropies. Going from solid to gas increases the entropy of a system as there are now more ways of arranging the particles.

Entropy Change

  • The standard entropy of a substance refers to the entropy of that substance measured in standard conditions (temperature of 298K and a pressure of 101Kpa).

  • Units for standard entropy are JK  mol

  • A substance has different entropies depending upon its state and the conditions it is in.

  • Water has a low entropy below 0°C (as it’s a solid), but a high entropy at over 100°C (as it’s a gas).

  • Standard entropies, So, enable the entropies of different substances to be compared and the overall entropy change that occurs in a reaction can be found.

  • Entropy change, ∆S = (sum of standard entropies of products) – (sum of standard entropies of reactants)

Free Energy (Gibbs)

  • The feasibility of a reaction refers to how likely it is to happen (high feasibility means a reaction is likely to happen, low feasibility means a reaction is unlikely to happen).

  • For a reaction to be feasible, energy must be released overall.

  • The overall energy change of a reaction is determined by both the change in enthalpy and the change in entropy that would occur.

  • Gibbs free energy is a value calculated that shows how energy changes overall during a reaction. The equation is:       ∆G = ∆H - T∆S
    where ∆H = enthalpy change, T = temperature (in Kelvin) and ∆S = entropy change.

  • If Gibbs free energy is negative – reaction is feasible (overall increase in stability).

  • If Gibbs free energy is positive – reaction is not feasibility (overall decrease in stability).

Total Entropy

  • During a chemical reaction, the entropy of reacting particles and the surroundings change.

  • How the entropies of the reacting particles change in a reaction is called the entropy of the system, ∆Ssystem

  • How the entropy of the surroundings changes in a reaction is called the entropy of the surroundings, ∆Ssurroundings

  • Entropy change of the system and entropy change of the surroundings can be combined to describe the total entropy change of the reaction, ∆Stotal.

  • The entropy of surroundings is linked to the temperature of the reaction and the enthalpy change that occurs during the reaction. It can be calculated by dividing the change in enthalpy (∆H) by the temperature.

  • The total entropy change that occurs in a reaction can be calculated using:




     

  • For a reaction to be possible, change in total entropy must be positive.

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A2-Level Rates of Reaction

Rates of Reaction

  • Kinetics is the study of reaction rates (how quickly reactions occur) and the factors that control the rate of a reaction.

  • Rates of reaction can only be found experimentally, by measuring a change in concentration of a reactant or product in a given amount of time.

  • How fast a reaction happens is called its rate of reaction and is measured with the units mol dm  s .

  • Concentration of reactants, temperature of the reaction, the pressure of a (gaseous) system and the activation energy for a reaction all determine the rate of a particular reaction.
     

Orders of Reaction

  • Changing the concentration of reactants in a reaction can affect the rate of the reaction.

  • The orders of a reaction describe how much the rate of a reaction changes when the concentrations of reactants are changed.

  • The order of a reaction with respect to a reactant describes how much the rate of a reaction change when the concentration of that reactant is changed by a given factor

    • Zero order means changing the concentration of the reactant has no change on the rate of reaction.

    • First order means changing the concentration of the reactant by a given factor will change the rate of the reaction by the same factor.

    • Second order means changing the concentration of the reactant by a given factor will change the rate of the reaction by the same factor squared.
       

The Rate Equation

  • A rate equation can be used to show how the rate of a reaction changes considering all the reactants’ concentrations.

  • A rate constant, k, is used in the rate equation to account for the influence of temperature on the rate of a reaction. A rate constant is only for a given temperature.

  • Different reactions can have different rate constant values and the units can also be different.

  • In the rate equation, each concentration of reactant is raised to the power of its order and multiplied by the rate constant, k.




     

Rate Equation and Mechanisms

  • Reactions often happen in several steps (described using mechanisms).

  • The slowest step in a reaction determines the rate of the reaction and is called the rate determining step. 

  • The rate equation for a reaction can be used to help predict the mechanism of a reaction and it identifies the reactants involved in the rate determining step.

Clock Reactions

  • Measuring the initial rate of a reaction is difficult because the concentrations of the reactants are constantly changing, meaning the rate of reaction is also changing.

  • Clock reactions are used to find the initial rate of a reaction by measuring the length of time taken to form a small amount of product.

  • The iodine clock experiment is a common example of a clock reaction at A-level chemistry.

  • In the main reaction, hydrogen peroxide reacts with iodide ions to produce iodine molecules.



     

  • In a second reaction, thiosulfate ions react with the iodine produced in the main reaction to form iodide ions again.



     

  • Starch indicator is added to the reaction mixture, which turns dark blue in the presence of iodine.  As long as there are thiosulfate ions in the mixture the starch will not cause a color change, as the thiosulfate ions are instantly converting any iodine molecules formed into iodide ions.

  • When all the thiosulfate ions are used up the starch will cause a color change, as iodine molecules will no longer be reacted back to iodide ions by the thiosulfate ions and they can react with the starch.

  • In the clock reaction, the amount of thiosulfate ions in the mixture is known and the length of time taken for the mixture to change color is recorded. This tells you how long it took to form the amount of iodine required to completely react with all the thiosulfate ions present.

  • The time taken can be used to find the rate of reaction.

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A2-Level Equilibrium

Equilibrium Constant, Kc

  • Equilibrium of a reversible reaction is achieved when the concentrations of both reactants and products do not change.

  • At equilibrium, the rate of the forward reaction is the same as the reverse reaction.

  • How much the forward reaction or reverse reaction is favored can be shown mathematically by comparing the concentrations of reactants to products at equilibrium.

  • The equilibrium constant, Kc, is a constant that describes the ratio between reactants and products at equilibrium. It is calculated by dividing the concentrations of products by the concentrations of reactants (raised to the power of their molar ratios).



     

  • Kc values are only for a specific temperature.
     

Partial Pressures

  • 1 mole of all gases are assumed to occupy the same volume in space (24 000cm at room temperature and pressure).

  • The total pressure of a gaseous system is directly related to the number of moles in the system.

  • How much pressure one gas contributes to the total pressure of a system is called its partial pressure.

  • All partial pressures of gases in a system add up to give the total pressure of the system.

  • The mole fraction of a gas describes how many moles of a gas there are compared to all other gases in the system.

Equilibrium Constant, Kp

  • How much the forward or reverse reaction is favored at equilibrium can be found by comparing the amounts of reactants to products at the point of equilibrium.

  • As the partial pressures of a gas in a closed system are directly linked to the moles of that gas, partial pressures of reactants and products can be used to determine how much the forward or reverse reaction is favoured. 

  • The equilibrium constant, Kp, describes the ratio of reactants compared to products. It is found by dividing the partial pressures of all products by the partial pressures of all reactants (raised to the power of their molar ratios).

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A2-Level Electrochemistry

The Basics

  • Metals react by losing electrons and forming positive ions.

  • The reactivity of a metal refers to how easily it can lose electrons (and form positive ions).

  • When a reactive metal is placed in water, metal atoms lose electrons and form positive ions. The lost electrons have nowhere to go, so they build up on the surface of the metal. Eventually, an equilibrium is established between the ions in solution and the metal itself.

  • A build-up of charge on a surface creates an electrical potential. The size of the electrical potential is determined by the equilibrium between the ions in solution and the metal solid – this is called an electrode potential.

  • Highly reactive metals will lose electrons easily and there will be a greater build-up of electrons on the surface of the metal.

  • The more electrons on the surface of the metal, the more negative the electrode potential.

  • Zinc is more reactive than iron, so it would have a greater buildup of electrons on its surface – giving it a more negative potential.

  • The solution, metal strip (electrode) and container are called a half-cell.
     

Electrochemical Cells

  • If two different half-cells are connected by a wire (and a salt bridge), an electrical current is created and electricity flows.

    • This is called an electrochemical cell. 

  • Electrical current is the net movement of charge in one direction. As electrons are charged, if they are moving in one direction then an electrical current is produced.

  • The electrode that has the most electrons on its surface (more negative electrode potential) will force its electrons to flow to the electrode with fewer electrons on its surface (more positive electrode potential).

  • The greater the difference in potential (voltage) between the two half cells, the greater the amount of electricity produced.

  • It is the relative difference between potentials that is important in electrochemistry. It is not possible to measure the actual potential of a half-cell – just how different it is compared to other half-cells.
     

Electrode Potentials

  • If a half-cell is connected to another half-cell that is the same, the potential difference between the two will be zero. If one of the half-cells is changed to a different type of half-cell, there will now be a difference in potential between the two – enabling a comparison to be made.

  • The relative potentials of half-cells are found by comparing them all to the same type of half-cell. The half-cell used is called the standard hydrogen electrode.

  • All potential differences measured are referred to as standard electrode potentials.

  • For an electrode potential to be considered standard, all ions in solution for the half-cell being used must be 1mol dm , temperature must be 298K and, for gaseous half-cells, the pressure must be 101kPa – standard conditions.

  • The potential difference for a cell (Ecell) made of two different half-cells can be calculated by finding the overall difference between their standard electrode potentials.



     

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A2-Level Acids and Bases

Acids and Bases

  • Acids are species that donate protons and bases are species that accept protons.

    • This is the Bronsted-Lowry definition of acids and bases.

  • In solution, acids actually release protons to water molecules, forming hydroxonium ions (H O⁺).

  • Hydroxonium ions then react with bases, giving the extra proton to the base and reforming water.

  • In equations, H⁺ (aq) is used to represent a proton, assuming it is understood that the proton actually exists in solution with water as hydroxonium ions.

  • Acids lose protons to form their conjugate base.

  • Bases gain protons to form their conjugate acid.
     

Strong and Weak Acids

  • Strong acids fully dissociate in solution – all molecules of the acid split, releasing the maximum number of protons, H⁺(aq), possible.

  • Weak acids partially dissociate in solution – an acid molecule that has released a proton (to form its conjugate base) can re-gain a proton to form a molecule of the acid again.

  • Monoprotic acids release one H⁺(aq) ion for every one molecule of acid.

  • Diprotic acids release two H⁺(aq) ions for every one molecule of acid.

  • The concentration of a strong, monoprotic acid is the same as the concentration of H⁺(aq) ions present in a solution of that acid.
     

Acid Dissociation Constant, Ka

  • Weak acids only partially dissociate in solution, forming an equilibrium between the acid and its conjugate base.

  • Acid dissociation constants, Ka, are used to show how much either the forward or reverse reaction is favoured and to give an indication of the acid’s strength. 

  • If a weak acid is described as HA and its conjugate base as A⁻, then the Ka for that acid can be found using:

     ​



  • The greater the value of Ka, the stronger the acid. 

  • The smaller the value of Ka, the weaker the acid.

pH Calculations

  • pH (potential of hydrogen) is a scale to show how acidic or alkaline a solution is.

  • Acidity is measured by the concentration of H⁺(aq) ions in a solution, meaning pH is a way of describing the H⁺(aq) concentration in a solution.

  • pH of 1 means solution is highly acidic; pH of 14 means solution is highly alkaline.

  • The scale is logarithmic, meaning a change in pH value of one refers to a change in concentration of H⁺(aq) ions of 10. It can be calculated using
                                                   
         pH = - log10[H⁺]
         [H⁺] = 10
     

  • pH of 0 means a solution has a H⁺(aq) ion concentration of 1 mol dm .

  • pH of 1 means a solution has a H⁺(aq) ion concentration of 0.1 mol dm .

Ionic Product of Water, Kw

  • To find the pH of an alkaline solution, the concentration of H⁺(aq) ions in solution can be found using the ionic product of water, Kw.

  • Water dissociates to release H⁺(aq) ions and can be considered a (very) weak acid mathematically.

  • The ionic product of water, Kw, is a constant that links the concentration of OH⁻(aq) ions in solution to the concentration of H⁺(aq) ions in solution


     

  • At 298K, Kw has a fixed value (1 x 10  ).

Buffer Solutions

  • Buffer solutions are a mixture of a weak acid and its conjugate base.

  • They are used to minimise a change in pH of a solution. 

  • An equilibrium is established between the weak acid and conjugate base.






     

  • The weak acid and conjugate base are in excess, meaning that the position of the equilibrium established between them will not be sensitive to changes in their concentrations, but it will be very sensitive to changes in the concentration of H⁺(aq) ions.  

  • When H⁺(aq) ion concentration increases, the position of equilibrium moves to oppose the change. The concentration of HA increases by the same amount as the concentration of A⁻ decreases.
     

Titration Curves

  • Titration curves show how the pH of an acidic (or alkaline) solution changes as a certain volume of alkali (or acid) is added.

  • The end point of a titration is when an indicator changes colour at a certain pH.

  • The equivalence point in a titration is the point when the concentration of H⁺(aq) ions  is the same as (or ‘equivalent’ to) the concentration of OH⁻(aq) ions.

  • Equivalence points are not always pH 7, which means different indicators that change colour at a pH very close to the equivalence point must be used.

Finding Ka using a Titration Curve

  • The acid dissociate constant (Ka) for a weak acid (HA) can be found using a titration curve.

  • At the equivalence point, the concentration of HA equals zero and the concentration of A⁻ will be the same as the original concentration of HA.

  • At the half-equivalence point, the concentrations of HA and A⁻ will be the same, meaning in the Ka expression for a weak acid, they cancel each other out. Ka will be the same as the concentration of H⁺ ions.

  • This value for H⁺ concentration will be the same as the value for Ka, meaning Ka = 10 

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