Physical chemistry

Define relative atomic mass:

Mean mass of 1 atom of an element relaitve to 1/12 mass of one atom of C12

How do you calculate the abundance of an ion:

Abundance is relative to the size of the current produced

Define electronegativity:

Ability of an atom to attract the electron density in a covalent bond

Explain why the C-Cl bond is polar

Cl atom is more electronegative

Therefore Cl becomes delta negative and C becomes delta positive

Define Van der Waals forces and how they are caused:

• forces caused from the random movement of electrons creating a temporary dipole

• Induces a dipole in another molecule

• Dipoles in different molecules attract one another

Shape of molecule with 2 bp:

Linear

180

Shape of molecule with 3bp:

Trigonal planar

120

Shape of molecule with 2bp 1lp:

Bent

118

Shape of molecule with 4bp:

Tetrahedral

109.5

Shape of molecule with 3bp 1lp:

Trigonal pyramidal

107

Shape of molecule with 2bp 2lp:

Bent

104.5

Shape of molecule with 5 bp:

Trigonal bipyramidal

120, 90

Shape of molecule with 4bp 1lp:

Trigonal pyramidal/ see saw

119, 89

Shape of molecule with 3 bp 2lp:

Trigonal planar

120

Shape of molecule with 6 bp:

Octahedral

90

Shape of molecule with 5 bp 1lp:

Square pyramid

89

Shape of molecule with 4 bp and 2 lp:

Square planar

90

What are the four crystal structure types:

• ionic

• Metallic

• Macromolecular/ giant covalent

• Molecular

Ionic crystal characteristics:

• high melting and boiling point

• Strong electrostatic force of attraction holding the lattice

• Molten or in solution can conduct electricty

• Brittle

Metallic crystal characteristic:

• sea of delocalised electron

• Malleable

• High melting point

• Strong and many force of attraction between positive ions and delocalised electrons

Simple molecular crystal structure:

• covalently bonded molecules held together by weak van der waals forces

• Low melting point and boiling point

• E.g. water is simple molecular but has high boiling point due to hydrogen bonds

• Poor conductors, no charged particles

Macromolecular crystal characteristics:

• covalently bonded into a giant lattice structure

• Each atom has multiple covalent bonds

• Very high melting point

• Rigid

• E.g. diamond (4C) And graphite (3C)

Difference between malleable, ductile, and brittle:

• malleable: ability to be hammered or pressed into sheets under compressive stress

• Ductile: ability to be stretched into wires under tensile stress

• Brittle: tendency of material to break or shatter under stress

What is enthalpy change?

• heat energy change measured under conditions of constant pressure

Delta H = energy to break bonds + energy to make bonds

What are the standard conditions:

100 kPa

1 mol dm³

298K

Define Hess’s law:

enthalpy change of a reaction is independent of the route taken

Define mean bond enthalpy:

Enthalpy change needed to break the covalent bond into gaseous atoms averaged over different molecules

Define enthalpy of formation:

Enthalpy change when one mole of a substance is formed from its constituent elements with all substances in their standard states under standard conditions

-ve (exothermic)

Define enthalpy of combustion:

Enthalpy change when one mole of a substance undergoes complete combustion in sufficient oxygen with all substances in standard states under

Exothermic (-ve)

Define enthalpy of neutralisation:

• enthalpy change when one mole of water is formed in a reaction between an acid and alkali under standard conditions

• (-ve) exothermic

Define ionisation enthalpy:

• enthalpy change when each atom in one mole of gaseous atoms loses one electron to form one mole of gaseous +1 ions

• +ve endothermic

Define electron affinity:

• enthalpy change when each atom in one mole of gaseous atoms gains one electron to form one mole of gaseous 1- ions

• First ea = exothermic

• Second ea = endothermic

Define enthalpy of atomisation:

Enthalpy change when one mole of gaseous atoms is produced fom an element in it standard states

Endothermic + ve

Define enthalpy of hydration

• enthalpy change when one mole of gaseous ions become hydrated (dissolved in water)

• Exothermic (-ve)

Define enthalpy of solution:

Enthalpy change when one mole of an ionic solid dissolves in an amount of water so that the dissolved ions are well separated and do not interact with each other

Varies

Define lattice enthalpy of formation:

• enthalpy change when one mole of solid ion compounds is formed from its constituent ions in the gas phase

• exothermic -ve

Define lattice enthalpy of dissociation:

• enthalpy change when one mole of solid ionic compound is broken into its constituent ions in the gas phase

• Endothermic +ve

Define enthalpy of vapourisation:

• enthalpy change when one mole of a liquid is turned into a gas

• Endothermic +ve

Define enthalpy of fusion:

• enthalpy change when one mole of a solid is turned into a liquid

• Endothermic +ve

Define activation energy:

• minimum energy required to start a reaction

Define catalyst and how they work:

• substance that increases the rate of a chemical reaction without being changed in chemical composition or amount

• Work by providing an alternative reaction route of lower activation energy

Conditions for dynamic equilibrium:

• closed system

• Rates of forward and backward reactions are equal

• Constant concentration of reactions and products

• Reactants and products are dynamic (constantly moving)

Define Le Chatelier’s principle:

• if a system at equilibrium is disturbed, the equilibrium moves in the direction that tends to reduce the disturbance

What is entropy:

• measure of how energy is spread out over a course of a reaction

What is Gibbs free energy?

• measure of how much available energy there is in a system

• dG = dH - T x dS

• Units = J mol -1

• If reaction is not spontaneous : dG > 0

• If reaction is spontaneous : dG < 0

• If reaction is in equilibrium : dG = 0

• If reaction is not spontaneous, backward reactions is feasible

What is the salt bridge for and what is it made of?

• completing the circuit and allowing ions to move

• Made of potassium/ sodium nitrate

How would the EMF of the cell change if the surface area of the platinum electrode is increased:

No change

What effect does increasing pressure have on the Kp value:

No change

Explain why the enthalpy of hydration of fluoride ions is more negative than the enthalpy of hydration of chloride ions

Smaller ion

Stronger attraction D+ on H

Why do fuel cells not need to be electrically recharged?

Reactants supplied continuously

Give the equation of the formation of Sr + ions from Sr atoms by electron impact

Sr + e- → Sr+ + 2e-

A student plans to titrate butanoic acid with a solution of ethylamine. Explain why this titration could not be done using an indicator.

Weak acid and weak base

pH change is too gradual

What is a bronsted-lowry acid?

Proton donor

What is a bronsted-lowry base?

Proton acceptor

Electrode reactions in a lithium cell:

Li+ + CoO2 + e- → Li+[CoO2]- (positive)

Li + + e- → Li (negative)

Electrode reactions in an alkaline hydrogen-oxygen fuel cell:

2H2(g) + 4OH- → 4H2) + 4e- (negative)

O2 + 2H2O + 4e- → 4OH- (positive)

Overall: 2H2 + O2 → 2H2O

Units for Kp

KPa

How are buffer solutions made?

Adding excess acid to base

How do buffer solutions work?

A- reacts with H+ from acid

As the electrode potential gets more negative,

• reducing power increases

• Goes on the right side of a standard cell representation

• Overall should always be positive

Which half cell is reduced?

The more positive

What might a salt bridge look like?

• filter paper soaked with solution of unreactive ions (potassium nitrate)

• Tube containing unreactive ions in an agar gel

How is the full potential difference measured under zero-current conditions?

• Using a high resistance voltmeter