[Year 2] O

• Most UV radiation from the sun is absorbed by gasses in the earth’s atmosphere. • Some is reflected back into space. • Some reaches the surface of Earth which is absorbed. • This heats up the surface of the earth and radiates IR back towards space. • Some IR escapes through the ‘IR window’, as they are not absorbed by the gases in the atmosphere. • Some IR is absorbed by gases in the troposphere and reemitted in all directions included back towards earth.

CO₂ H₂O CH₄

• Their bonds in the molecules vibrate more. • The vibrational energy is transferred to neighbouring particles during collisions. • These particles now have more kinetic energy and their temperature starts to rise.

Using polar solvents.

• Hydration is dissolving with water as a solvent. - Whereas solvation is dissolving with a non-water solvent.

• Water is polar. • Its 𝛿+ Hs are attracted to the negative ions • Its 𝛿- O is attracted to the positive ions. • The water molecules surround the ions during hydration. • This ion-dipole interaction must be stronger/as strong as the bond broken for it to dissolve.

• The bond between Al and O ions are stronger than the ion-dipole interaction that would form.

Through hydrogen bonds. e.g. alcohols disolving in water

• they have weak dipoles. - so water will form a stronger intermolecular bond than with the haloalkane.

Use a solvent that can interact via pertinent dipole-dipole interactions.

• by forming instantaneous dipole-induced dipole forces.

The energy change when 𝟭 𝗺𝗼𝗹𝗲 of an ionic compound is formed from its 𝗴𝗮𝘀𝗲𝗼𝘂𝘀 ions under 𝘀𝘁𝗮𝗻𝗱𝗮𝗿𝗱 𝗰𝗼𝗻𝗱𝗶𝘁𝗶𝗼𝗻𝘀. (298K/100kPa)

• Always negative, as energy is released as bonds form. - more negative = stronger bonding.

The energy change when 𝟭 𝗺𝗼𝗹𝗲 of an ionic compound is dissolved in the 𝗺𝗶𝗻𝗶𝗺𝘂𝗺 𝗮𝗺𝗼𝘂𝗻𝘁 𝗼𝗳 𝘀𝗼𝗹𝘃𝗲𝗻𝘁 to ensure 𝗻𝗼 𝗳𝘂𝗿𝘁𝗵𝗲𝗿 enthalpy change upon further dilution.

The energy change when 𝟭 𝗺𝗼𝗹𝗲 of aqueous ions are formed from its 𝗴𝗮𝘀𝗲𝗼𝘂𝘀 ions under 𝘀𝘁𝗮𝗻𝗱𝗮𝗿𝗱 𝗰𝗼𝗻𝗱𝗶𝘁𝗶𝗼𝗻𝘀.

enthalpy of solution = enthalpy of hydration - lattice enthalpy.

• The enthalpy change of solution will be negative. • So the enthalpy change of solution will release energy. • So the process is exothermic.

Calorimetry.

• In a polystyrene cup add the solvent and measure the initial temperature of the solvent, • measure out your solid solute in a weighing boat using scales. • add the solute to the solvent and stir. Put a lid over the polystyrene cup (leave a hole in the lid to keep the thermometer in. • Note down the maximum temperature reached after the solute was added. • use the equation E = mcΔT to work out the enthalpy. (mass in grams, T can be in any unit). • Work out the moles of solute added by doing moles = mass/Mr • Divide this by the energy to find the enthalpy change of solution.

• heat lost to surroundigns.

• Size of ions (ionic radii) - Size of change (bigger difference in charge = stronger bond).

• The smaller the ionic radii, the stronger the electrostatic attraction between ions. • as smaller ions can pack together more closely. • and so have a larger charge density. • Thus more energy is required to overcome the stronger forces.

• Bigger the charge on an ion, the stronger the electrostatic attraction between ions. • Thus more energy is required to overcome these forces.

The measure of disorder in a system. • It is the number of ways energy can be shared out in a system. • Given the letter S.