topic 12

a reaction that will proceed on its own without any external influence (i.e. diffusion)

exothermic reactions result in products that are lower in energy (thermodynamically stable),

-the number of ways particles can be arranged as well as energy quanta's

gases

-much more disorderd and are able to move freely within the system

S=0, as there is no particle movement, and only one possible arrangement

ΔStotal = ΔSsystem +ΔSsurroundings

ΔStotal > 0

- change of state (solid-> gas)
- increase in number of moles between products and reactants

ΔSsystem= ΔSproducts - ΔSreactants

JK⁻¹·mol⁻¹ (joules per kelvin per mol)

ΔS₍surroundings₎ = −ΔH / T

- surroundings gain entropy as heat is released so the particles gain more kinetic energy
-ΔS₍surroundings₎ is positive.

- surroundings lose entropy as heat is absorbed by the system so the particles lose kinetic energy
-ΔS₍surroundings₎ is negative

When ΔH is negative, ΔS₍surroundings₎ = −ΔH / T is positive.
If temperature increases, the value of ΔS₍surroundings₎ decreases, because you are dividing by a larger T.

At higher temperatures, the surroundings are already more disordered, so adding the same amount of heat produces a smaller increase in entropy.

because the ΔS₍system₎ can be large enough that (ΔSₜₒₜₐₗ) is positive

T = ΔH / ΔSsystem

when a substance is in its lowest energy state

exothermic reactions form thermodynamically stable products

A substance is kinetically stable if it has a high activation energy (Ea), meaning the reaction is very slow

The enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions (298K & 100kpa)

Na(s) + 0.5Cl2(g) ---> NaCl (s) [-411.2 Kj mol]

The enthalpy change when 1 mole of gaseous atoms is formed from its element in its standard state, under standard conditions.

Na(s) --> Na(g) [+148kJ mol]
0.5O2 ---> O [+249kJ mol]

the enthalpy change required to remove one electron from one mole of gaseous atom to form one mole of gaseous 1+ ion

Mg(g)--> Mg+(g) + e-

the enthalpy change required to remove one electron from one mole of gaseous ion to form one mole of gaseous 2+ ion

Mg+(g)--->Mg2+(g) + e-

the enthalpy change when one electron is gained by one mole of gaseous atom to form of mole gaseous 1- ion

O(g) + e- ---> O-(g) [-141.1]

enthalpy change when one electron is gained by one mole of gaseous 1- ion to form one mole of gaseous 2- ion

O-(g) + e- ---> O2- [+798 KJ mol ]

Because adding one electron to a neutral atom forms a negative ion with a full outer shell, which is more stable, so energy is released (exothermic).

because energy input is required to overcome the repulsive force between the additional electron and the negative ion

standard enthalpy change when 1 mole of an ionic compound is fromed from its gaseous ions

Na+(g) + Cl-(g) --> NaCl(s) [-787Kj mol]

The larger the ions, the less negative the lattice enthalpy, because larger ions have charges more spread out → weaker electrostatic attraction between ions.

The larger the charge on the ions, the more negative the lattice enthalpy, because stronger electrostatic attraction → stronger lattice.

NaCl(s)

1- enthalpy change of atomisation (Na) (endothermic)
2- enthalpy change of atomisation (Cl) (endothermic)
3- 1st ionisation energy (Na) ([Na+] + Cl + [e-]) (endothermic)
4- 1st Electron affinity (Cl) ([Na+] + [Cl-]) (exothermic)

The enthalpy change when 1 mole of an ionic solid dissolves in water to form aqueous ions.

*It can be endothermic or exothermic depending on lattice energy and hydration enthalpies.*

KCl(s)→K+(aq)+Cl−(aq).

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

*This is always exothermic because water molecules attract and surround the ion, releasing energy.*

Na+(g)→Na+(aq)

ΔH(solution) = Σ Hydration Enthalpies − Lattice Enthalpy

Hydroxides: increase down a group

Sulfates: decrease down a group