Year 12s: what will be the first way you research your future options? >>

AQA Chemistry Unit 5 with Anon_98.

Well, these threads help me quite a lot + they seem to be helping other people too so I'm doing another one. The exam's on Wednesday + my aim is to finish by tomorrow night at the absolute latest so I can do Bio 5. (It'll probably end up going past that though bc I'm tired + blah. )': )

Anyway, let us begin.. :crazy:

Scroll to see replies

Reply 1

Anon_98

Thermodynamics - Enthalpy Change - Topic 1.

Enthalpy of formation - The enthalpy change when 1 mol of a compound is formed from it's elements under standard conditions, all reactants and products being in their standard states.

(E.g. Na(s) + 1/2Cl2(g) -------------------> NaCl(s))

Ionisation enthalpy(DHi) - The enthalpy change for the removal of 1 mol of electrons from a species in the gas phase to form a positive ion also in the gas phase.

(First ionisation enthalpy of sodium: Na(g) ---------> Na⁺(g) + e^-(g) )
(Second ionisation enthalpy of sodium: Na⁺(g) ------------> Na²⁺(g) + e^-(g))

Enthalpy of atomisation of an element - The enthalpy change when 1 mol of gaseous atoms is formed from the element in its standard state.

(E.g. Na(s) → Na(g))

Enthalpy of atomisation of a compound- The enthalpy change when 1 mol of a gaseous compound is completely atomised.

(E.g. C₂H₆(g) → 2C(g) + 6H(g))

Bond dissociation enthalpy - The enthalpy change when a covalent bond is broken in a gaseous molecule to form two gaseous free radicals.

(E.g. Cl₂(g) ---> 2Cl(g) OR CH₄(g) -----> CH₃(g) + H(g))

Electron affinity - The enthalpy change when an electron is added to an atom or ion in the gas phase.

(E.g. Cl(g) + e-(g) -----------> Cl^-(g))

Lattice enthalpy of formation - The enthalpy change when 1 mol of a solid ionic lattice is formed from its gaseous ions.
(exothermic, therefore arrow down in born haber cycle)

(E.g. Na⁺(g) + Cl^-(g) -------------> NaCl(s))

Lattice enthalpy of dissociation - The enthalpy change when 1 mol of a solid ionic lattice is dissociates into its gaseous ions.
(endothermic, therefore arrow up in born haber cycle.)

(E.g. NaCl(s) -------------->Na⁺(g) + Cl^-(g))

Enthalpy of hydration - The enthalpy change when 1 mol of gaseous ions change in to 1 mol of aqueous ions.

(E.g. Cl^-(g) -------> Cl(aq))

Enthalpy of solution - The enthalpy change when 1 mol of an ionic solid dissolves in enough water to ensure separation of the ions such that they do not interact with each other.

(E.g. NaCl(s)+ (aq) -------> Na⁺(aq) + Cl^-(aq))

Reply 2

Absent Agent

Original post by Anon_98

Writing down the information you already know, especially in an attempt to summarise them, allows them to be stored into your long-term memory, shifting and deepening your mental pathways. It's surprising that Einstein came up with one of the happiest thought of his life when he was rewriting the consequences of one of his theories. :wink:

Reply 3

Anon_98

Why is the second ionisation of an element greater than its first ionisation enthalpy?

- The second electron is removed from a smaller ion and is thus closer to the attraction of the nucleus.(The ion contracts because the same number of protons attract fewer electrons.)

- The second electron maybe less shielded than the first electron if it occupies an inner energy level.

Reply 4

Anon_98

Why is the second electron affinity of an element endothermic whereas the first electron affinity is exothermic?

- The first electron has a strong affinity(attraction) for the nucleus of the atom so that energy is released to the surroundings.

- The second electron is repelled by the negatively charged ion and energy is required to overcome this force of repulsion.

Reply 5

flyingpanda

Goooo Anon_98!

Your notes are always very nice! :biggrin:

Reply 6

Kevin De Bruyne

You can do it, Anon! :rave:

Reply 7

Anon_98

Original post by flyingpanda

Goooo Anon_98!

Your notes are always very nice! :biggrin:

Original post by SeanFM

You can do it, Anon! :rave:

thanks guys! <3

Reply 8

Anon_98

Okay, so as an example I'm going to draw the Born-Haber cycle for Calcium oxide below:

Reply 9

Anon_98

Drawing this was such a struggle, I couldn't reach. ) ':

Posted from TSR Mobile

1466211863367.jpg15.6KB

Reply 10

Anon_98

Alright so, the general steps to drawing these are:

- Formation.
- Atomisation of metal.
- Ionisation of metal.
- Atomisation of gas.
- Affinity of gas.

Reply 11

Anon_98

Why does calcium oxide has a larger lattice enthalpy value than sodium oxide?

Well it's bc there are two factors that produce a higher lattice enthalpy – smaller size and higher charge.

Although the calcium ion will be larger than the sodium ion and thus less attracting to the oxide ion, its higher charge increases its attraction for oxide ions and produces a higher lattice enthalpy

Reply 12

Anon_98

The perfect Ionic model.

The perfect ionic model assumes that the ions are perfect spheres (point charges) with no distortion due to polarisation which would give rise to some covalent character (overlap of ions).

~

How do calculations of lattice enthalpies based on the perfect ionic model provide evidence for covalent character in ionic compounds when compared with values obtained from the Born-Haber Cycle?

A larger lattice enthalpy value, obtained from the Born-Haber cycle, compared with a value calculated using the theoretical model would indicate that the compound has a degree of covalent character. The forces in the lattice are therefore stronger than in pure ionic attraction.

Reply 13

Anon_98

Okay so here is the cycle for the enthalpy of solution of magnesium chloride with its lattice enthalpy and the hydration enthalpies of its ions:

Reply 14

Anon_98

Posted from TSR Mobile

1466213823336.jpg15.3KB

Reply 15

Anon_98

Bond enthalpy - The energy needed to break 1 mol of a particular covalent bond in the gaseous state averaged over several compounds where the bonds exists.

Bond breaking is an endothermic process, Bond forming is an exothermic process.

Bond enthaply formula.

DHr = Σ (enthaply bonds broken) - Σ (enthalpy bonds formed)

~

Why would the enthalpy change for a reaction calculated from mean bond enthalpies differ from that determined using enthalpy of formation data?

It's bc the echanges obtained using mean bond enthalpies use values that are averages of the bond enthalpies – not the actual bond enthalpies - whereas enthalpy of formation data has no such approximation.

Reply 16

Anon_98

Thermodynamics - Free energy change(DG) and entropy change (DS) - Topic 1.

Okay so here is the equation you need to know:

DG = DH - TDS

For a reaction to be just feasible: DG ≤ 0

~

An exothermic reaction is feasible when:
(DH is negative then for DG to be negative either....)

- TDS must be positive (I.E. increasing disorder)

- or if DS is negative then TDS must be ≤ DH

An endothermic reaction is feasible when:
(If DH is positive then for DG to be negative...)

- TDS must be negative (I.E. decreasing disorder) and TDS must be ≥ DH

~

If DS is negative then the reaction is not feasible at any temperature

(edited 7 years ago)

Reply 17

Anon_98

Alright, so now I'll draw a graph of entropy, S, against temperature, T, for water showing the melting and boiling points then I'll explain the points on the graph:

Reply 18

Anon_98

GRAPH OF ENTROPY AGAINST TEMPERATURE.

Posted from TSR Mobile

1466216299265.jpg24.3KB

Reply 19

Anon_98

Explanation of the graph above:

- At 0K, S = 0 because particles have no disorder.

- As temperature of solid increases particles vibrate about a fixed point with increasing speed and therefore become slightly more disordered.

- At 273K water melts and there is a jump in disorder as the particles move from the relatively ordered solid state to the less ordered liquid state where particles are free to move.

- As temperature increases the particles move more rapidly and with increasing disorder.

- At 373K water boils and there is a big increase in disorder as particles have much more random movement in the gaseous state

AQA Chemistry Unit 5 with Anon_98.

Quick Reply

Related discussions

Latest

Trending

Trending