Studied by 1 person
heat loss to the surroundings (use a draught shield)
non standard conditions
incomplete combustion
heat loss by the metal calorimeter (use a lid)
stirring of water (even distribution of heat)
generally speaking, substances won't dissolve if the bonds to be broken are stronger than those that will be formed
if less energy is released in making the solute-solvent intermolecular forces than required to break the intermolecular forces within the solute or solvent then it does not dissolve
polar (e.g. water)
non polar (e.g. hexane)
most ionic substances only dissolve in polar solvents
the ion-dipole bonds formed between ionic substances and polar solvent molecules are strong enough to dissolve the solid
non-polar solvents
non-polar solvents don't interact strongly enough with ions to pull them away from an ionic lattice
electrostatic forces between the ions are way stronger than any bonds that could form between the ions and non-polar solvent molecules
an ion in solution and the dipole of a polar solvent interact
the ions becomes surrounded in the polar solvent via the attractive force between the ion and the molecule with a permanent dipole
e.g. NaCl and Water
Hδ+ attracted to Cl-
Oδ- attracted to Na+
water is polar
ion dipole bonds form between the water molecules and the dissolved ions
the ions separate from the ionic lattice and become surrounded by water molecules (hydration)
unless the ionic bonds in the lattice are really strong, the ion dipoles will be strong enough to pull the lattice apart which is why most ionic substances are soluble in water
most only dissolve in non-polar solvents
intermolecular bonds between covalent molecules tend to be pretty weak and can be broken by non-polar solvent molecules
don't dissolve in polar solvents
e.g. iodine doesn't dissolve in water much
the hydrogen bonds between water molecules are stronger than the bonds that would form between water molecules and iodine molecules
lattice enthalpy of formation
always negative because energy is released
the more negative the lattice enthalpy, the stronger the bonding/ionic bond strength
Na+ (g) + Cl-(g) --> NaCl(s) -∆H1
enthalpy of hydration
going from gaseous ions to aqueous ions
releases energy making this process exothermic
bonds between ions and water are made (ion-dipole?)
∆H2
enthalpy of solution
the overall effect on the enthalpy when something dissolves
the net effect of ∆LeH and ∆HydH
NaCl(s) --> NaCl (aq)
∆H3
gaseous ions at the top
solid ions in the middle
aqueous ions at the bottom
reaction solution in polystyrene cup with lid and thermometer
measure temperature change
calculate the number of moles of what's reacting
need mass of total solution
q=mcat ÷1000
N=mass/me ∆SolH= q/n kjmol-¹
the ions with higher charge density are better at attracting each other in ionic lattices than those with lower charge densities
this means the ionic bonds are stronger, meaning more energy is released when the bonds are formed, giving them a more exothermic LeH
ions with a higher charge density are better at attracting water molecules than those with lower charge densities
this means the intermolecular bonds are stronger, meaning more energy is released when the bonds are made, giving them a more exothermic HydH
an acid that releases one H+ per molecule
a base that can only react with/accept one proton per molecule
ionic product of water
kw=[H+] [OH-]
kw=10^-14
(conc X original vol)÷new total volume = X
pH= -log10(X)
pair made of acid + base in an reaction
acid reactant pairs with base product
base reactant pairs with acid product
acid dissociation constant
Ka= [H+] [A-] ÷ [HA]
check if monoprotic/diprotic etc
calculate moles H+
calculate moles OH-
which is in excess
work out excess
convert to concentration by dividing by total volume in dm³
calculate pH
if excess H+ calculate pH using -log[H+]
if excess OH- calculate pH using 14-(-log10 [OH-])
pH higher less than 7
[HA] and [A-(salt)] are much higher than [H+]
pH more than 7
[Base] and [salt] much higher than [OH-]
when small amounts of acid or alkali are added, the ratio remains roughly constant and so the pH hardly changes
if large amounts of acid or alkali are added, the ratio would change and so the pH would change significantly (- moves or release hydrogen ions)
adding H+ ions
CH3COO- reacts with the H+ ion to form CH3COOH
equilibirum shifts left
reduces [H+] to close to original value
[CH3COOH] increases slightly
[CH3COO-] decreases slightly
pH doesn't change much
[CH3COOH]/[CH3COO-] >> [H+]
CH3COOH reacts with OH- to form CH3COO-
equilibrium shifts right
[H+] increases to close to original value
[CH3COO-] increases slightly
[CH3COOH] decreases slightly
pH doesn't change much
[CH3COOH]/[CH3COO-] >> [H+]
OH- reacts with the H+ to form water
equilibrium shifts to the right to increase the [OH-] to close to its original value
pH doesn't change much
[NH3]/[NH4+] >> [OH-]
NH4+ reacts with OH- to form NH3 and water
equilibrium shifts to the right
extra OH- is consumed
[NH3] increases slightly
[NH4+] decreases slightly
pH doesn't change much
[NH3]/[NH4+] >> [OH-]
shampoos and soaps (most shampoos contain a pH 5.5 buffer as soap is very alkaline makes your hair dull)
biological washing powders (to keep the pH at the correct level for enzymes to work)
food products (such as sodium citrate)
fermentation
rearrange Ka= [H+][A-]/[HA] to find [H+]
pH= - log[H+]
physical state
amount of energy a substance has
amount of particles
solid particles wobble about a fixed point so have hardly any disorder and have the lowest entropy
gas particles have the most disordered arrangements of particles and so have the highest entropy
the more particles, the more ways they and their energy can be arranged
e.g. N2O4 --> 2NO entropy increases because the number of moles increases
-ΔH/T
ΔH in jmol-1 (convert from kjmol-1 by x1000)
T in kelvin (convert from C by + 273)
the sun emits electromagnetic radiation mainly as visible light, UV radiation and infrared radiation
most the UV and IR are absorbed by the earth's atmospheric gases and some radiation is reflected back into space from clouds
the energy that reaches the earth's surface is mainly visible light and UV. Some is reflected into space by light coloured,shiny surfaces like ice and snow. The rest is absorbed by the earth causing it to heat up
the earth then radiates infrared radiation
some of the IR escapes through the IR window
greenhouse gases in the troposphere absorb some of the remaining IR in the IR window and re-emit in all directions
