A-Level Chemistry AQA Physical Chemistry

Standard enthalpy of formation

The enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions

Standard enthalpy of combustion

The enthalpy change when 1 mole of a substance is completely burnt in (excess) oxygen under standard condition

Enthalpy change

Heat energy change at constant pressure

Standard States

100kPa and 298K

Define electronegativity

A measure of the tendency of an atom to attract a bonding pair of electrons in a covalent bond

Stages of TOF spectrometer

Ionisation: spray/impact

Acceleration

Drift

Detection

Ionisation (stage in mass spectrometry)

Electrospray: sample dissolved and pushed through nozzle. High voltage applied, particles gain H+

Electron impact: sample vaporised and electron gun fires electrons knocking off electron creating 1+ ions

Acceleration (stage in mass spectrometry)

Ions accelerated by electric field so all have the same KE

Ion Drift (stage in mass spectrometry)

Region of no electric field

Detection (stage in mass spectrometry)

Detector detects particles and mass spectrum produced. m/z

Metallic Bonding

The attraction between delocalised electrons and the positive metal ions in a lattice

Hydrogen Bonding

Strong type of intermolecular dipole-dipole attraction. Occurs between hydrogen and F, O or N. Creates a covalent bond Hydrogen bonding occurs due to difference in electronegativity leading to bond polarity and there is an attraction between ∂+ on one molecule and ∂− on another

Permanent dipole-dipole bonding

Difference in charge between atoms due to difference in electron density

Van der Waal Forces

Weak attractive forces between molecules resulting from the formation of temporary dipoles due to electron charge clouds moving

Covalent Bonding

A bond formed when atoms share one or more pairs of electrons, with both nucleus' electrostatically attracted to electrons

Dative Bonding

Form of covalent bonding in which the electrons being shared are supplied by only one of the participating atoms . This type of bonding occurs when one of the atoms has a lone pair of electrons .

Mean bond enthalpy

The average value of the bond dissociation enthalpy for a given type of bond taken from a range of different compounds.

Activation Energy

The minimum amount of energy required to start a chemical reaction

Behaviour of ionic compounds

-Conduct electricity molten or dissolved in an aqueous solution as ions are free to move and carry their charge. Not possible solid, as in fixed position by ionic bonds

-High melting point due to giant ionic lattices, held together by strong electrostatic force requiring lots of energy to overcome

-Dissolve in water due to the -ve part of the polar H2O molecule attracting the +ve part of the ionic compound whilst the +ve part of the polar H2O molecule attracting the -ve part of the ionic compound. Ions in the compound get pulled away from the lattice so dissolves

Creating a standard solution

Work out moles needed

Measure mass of solid on boat, find precise mass by re-weighing boat after pouring mass in beaker with distilled water and stirring

Pour solution into volumetric flask to make up to volume

Rinse beaker and rod, add this to flask

Add stopper then shake upside down

Titration process

Determine if unknown concentration solution is acid or base, using litmus paper. Use a measured volume of the unknown concentration solution. Add one or two drops of indicator solution. Put standard solution into burette. Add this by drop by drop unknown concentration solution, swirling after each addition. When the colour in the flask with unknown solution changes, stop adding drops, and record the volume of the known concentration solution added

Ionic bonding

The electrostatic attraction between oppositely charged ions in a lattice

% yield

actual yield/theoretical yield x 100

% atom economy

mass of desired product / total mass of reactants x 100

Catalyst

Chemical agents that selectively speed up chemical reactions without being consumed by the reaction

Exothermic

Releases heat

Makes bonds

-ve

Endothermic

Absorbs heat

Breaks bonds

+ve

Explain how VDWFs arise between molecules

Within a covalent bond, the electrons move randomly resulting in temporary dipoles leaving one delta +ve, the other delta -ve. This induces a dipole in another molecule

Explain the structure of a metal

Ordered rows +ve metal ions surrounded by a sea of delocalised electrons with strong electrostatic force of attraction between them

Oxidising Agent

Electron acceptor, is reduced

Reducing Agent

Electron donator, is oxidised

What affects a Kc value?

Temperature

Define Le Chatelier's principle

When a system under dynamic equilibrium is subjected to a change, the equilibrium will shift so as to oppose the change

Effect of concentration

A+2B⇌C+D

Increasing A shifts to right

Increasing C shifts to left

Decreasing A shifts to left

Decreasing C shifts to left

Effect of pressure

A+2B⇌C+D

Increasing pressure shifts to side with fewer moles

Decreasing pressure shifts to side with more moles

Effect of temperature

A+2B⇌C+D

Increasing temperature shifts to endothermic to absorb heat

Decreasing temperature shifts to exothermic to release heat

Effect of a catalyst

Has no effect, but helps increase rate of reaction equally

Define relative atomic mass

The weighted mean mass of an atom of an element relative to 1/12th of the mass of an atom of carbon-12

Define rate of reaction

The change in concentration of a reactant or a product in a given time.

Define dynamic equilibrium

Concentrations of reactants and products remain constant in a closed system. Forwards and backwards have same rate

Define 1st ionisation energy

The energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions

Define 2nd ionisation energy

The energy required to remove 1 mole of electrons from 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions

Define Hess' Law

The total enthalpy change of a reaction is independent of the route taken.

What is an isotope?

Atoms with the same number of protons but different number of neutrons

Do isotopes have the same chemical properties?

Yes, because they have the same configuration of electrons but not same physical as different masses

Rate equation

Rate = k[A]^m[B]^n. units for k mol^-1dm^-3s^-1

Orders of reaction

Zero order: changes in concentration has no effect on rate

First order: changes in concentration has a proportional change on rate

Second order: changes in concentration hs a squared proportional change on rate

Overall order is m + n

What is a rate constant?

Is a reflection of the probability that a reaction will occur relates the rate of the reaction to the concentration of reactants. Fixed at a particular temperature, so will change if temperature does. The larger the k value, the faster the reaction. Increasing temperature, increases k

Iodine clock reaction

Reaction for observing concentration effects.

Monitor rate of reaction by disappearance of bisulfite by adding more IO3- than HSO3- at the start of reaction. When the bisulfite is all used up there will be some iodate left.

Can detect the appearance of iodine w/ aid of start indicator, formas a blue complex w/ iodine. The time it takes for blue color to suddenly appear indicates when all the bisulfite is used up.

H2O2 + 2I- +2H+ -> 2H2O + I2 2S2O32- + I2 -> S4O62- + 2I-

How can rate be measured in experiments?

-Change in pH in a reaction

-Amount of mass lost

-Volume of gas produced

-Colour change (using colorimeter)- calibration curve

Which step is the rate determining step?

Slowest step

Arrehnius equation

k = Ae^(-Ea/RT)

What happens when Ae gets smaller in k = Ae^(-Ea/RT)?

k gets bigger, so rate of reaction increases

Define the term overall order of reaction

The sum of powers to which the concentrations are raised in the rate equation

Mole fraction equation

Number of moles of a gas / total number of moles of all gases

Partial pressure of gas equation

Mole fraction of gas x total pressure of the mixture

What affects Kp value?

Temperature

What affect does increasing temperature have on the rate constant?

Means particles have more kinetic energy, so have more energy or energy equal to activation energy required. So more successful collisions, increasing rate of reaction thus increasing rate constant

What is a half cell?

One half of an electrochemical cell.

Which direction do electrons flow to in an electrochemical cell?

From most reactive to least. Meaning most reactive is oxidised, and least reactive is reduced. The most negative value ion gets oxidised

Why is the electrode made out of platinum?

As it is inert but electrically conductive

Define standard electrode potential

The voltage measured under standard conditions when a half cell is connected to a standard hydrogen electrode.

Standard conditions for electrode potential

-Solutions must have concentrations of 1.00 mol dm^-3

-Temperature must be 298K

-Pressure must be 100kPa

What is the electrochemical series?

A list of electrode potentials in order of decreasing or increasing potential.

Standard cell potential equation

E°cell = E°cathode+ve(reduced) - E°anode-ve(oxidised)

What makes an electrode reaction feasible?

If the E°cell is positive

What is an electrolyte?

A liquid containing free-moving ions which conducts electricity

Fuel cells advantages

High efficiency - get more energy out of the same amount of fuel than less efficient devices

Only by-product is water (no CO2 produced)

Don't need to be recharged - they will work as long as the fuel is supplied

Fuel cells disadvantages

Hydrogen is flammable and must be stored very carefully

Energy is required to produce the reactants of hydrogen and oxygen

The energy often comes from fossil fuels so some CO2 emissions

Fuel cell equation

2H2 + O2 → 2H2O

State the substances and conditions needed in a standard hydrogen electrode

H2(g) AND 100kPa

1 mol dm−3 AND HCl/HNO3/H+

Pt electrode AND temperature of 298 K

What is the purpose of a salt bridge?

Completes the circuit and allows ions to flow without the solutions mixing.

Are acids proton donors or acceptors?

Proton donors. When mixed with water hydrogen ions are released, and form H3O+ ions

Are bases proton donors or acceptors?

Proton acceptors. When mixed with water they grab hydrogen ions from water, forming OH- ions

How easily do acids and bases dissociate?

Strong acids/bases dissociate almost completely

Weak acids/bases dissociate only slightly

Which side does equilibrium lie on for acids and bases?

Strong acids/bases on the left

Weak acids/bases on the right

Ionic product of water

The product of the H+ and the OH - concentrations in water (and all aqueous solutions), or, Kw = [H+] [OH-]. Used to find pH for strong base

pH equation

pH = -log[H+]

[H+] = 10^-pH

[H+] is the hydrogen ion concentration of a solution. For diprotic acids this needs to be doubled

Acid dissociation constant of a weak acid (Ka)

Ka = [H+][A-]/[HA]

or Ka = [H+]^2/[HA]. Used to find pH of weak acid

Ka, Kc and Kw are all affected by what?

Temperature

What is the equivalence point?

The point in a titration when neutralisation is reached (i.e. when moles H+ = moles OH-). Vertical part in graph

When is phenolphthalein used?

Weak acid/strong base. Or for strong acid/strong base. Colourless at low pH, pink at high pH

When is methyl orange used?

Strong acid/weak base. Or or strong acid/strong base

Red at low pH, yellow at high pH

What is a buffer solution?

A solution that resists changes in pH when small quantities of acids or bases are added to it or when its diluted

Uses of buffers

Shampoos- makes hair silky

Washing powders- keeping pH at right level for enzymes

Biological buffers- (tissue stays at right pH) or blood

Acidic buffer solution

Made from a weak acid and its salt

Basic buffer solution

Made up of a weak base and its salt

Explain why [H2O] is not shown in the Kw expression

Because it is constant

Why is the pH probe is washed with distilled water between each of the calibration measurements?

So the different solutions don't contaminate each other

Strong acids

HCl, HBr, HI, HClO3, HClO4, HNO3, H2SO4

Strong bases

LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

Define enthalpy of lattice formation

The enthalpy change required to form 1 mole if solid ionic compound from its gaseous ions.

Define entropy

A measure of the disorder of a system. More disordered the particles are the higher the entropy is.

Gibbs Free Energy Change equation

ΔG= ΔH - TΔS

What makes a reaction feasible?

If the value of ΔG is negative or zero

Define enthalpy of hydration

The enthalpy change when 1 mole of gaseous ions becomes aqueous ions.

Enthalpy of solution equation

Enthalpy of solution = Enthalpy of lattice dissociation + Enthalpies of hydration

ΔHsoln = ΔH L.D. + ΣΔH hyd

Define enthalpy of solution

The enthalpy change when 1 mole of an ionic solid dissolves in a large enough amount of water to ensure that the dissolved ions are separated and don't interact with each other.

When is the oxidation state of oxygen different?

When it is in a peroxide it is then-1, not -2

Define homogeneous reaction

A reaction in which all of the reactants and products are in the same physical phase