Thermochemistry

Thermochemistry

1. FORCES AND PROPERTIES

   • LATENT HEAT OF FUSION
     • Or enthalpy of fusion: the energy needed to break enough of the weakest forces between the particles in one mole of a solid at its melting point to form a liquid
   • LATENT HEAT OF SUBLIMATION
     • Or enthalpy of sublimation: the energy needed to break the forces between the particles to convert one mole of solid to a gas
   • LATENT HEAT OF VAPORISATION
     • Or enthalpy of vaporisation: is the energy needed to break enough of the weakest forces between the particles in one mole of a liquid at its boiling point to form a gas

2. PERIODIC TRENDS

   • ATOMIC RADIUS
     • Average distance between neighbouring atoms
     • Across a Period - atomic number increases (because the number of protons increases) and atoms are smaller (the charge on the nucleus increases so the effective nuclear attraction on the atoms increases, the valence electrons are pulled closer to the nucleus and the atomic radius decreases)
     • Down a Group - the number of energy levels increases, which results in effective nuclear attraction on the valence electrons being weaker - the outer electrons are further away from the nucleus, with increased shielding from inner shells of electrons, reducing the electrostatic attraction between the
   • IONIC RADIUS
   • IONISATION ENERGY
     • Measured in kJmol-1
     • The minimum energy needed to remove one electron from each atom in a mole of atoms in the gaseous state
     • If H is + the process is endothermic, because energy is required to overcome the electrostatic attraction between the nucleus and the valence electrons and hence, to remove the electron
       • Process of ionisation can be continued (except for H) so another electron can be removed
     • Down a Group: decrease in ionisation energies. The electron lost is further away from the nucleus than the element above and there is more shielding by the increased number of inner shells of electrons
       • This results in a decrease in the electrostatic attraction between the nucleus and valence electron, and less energy is needed to remove the electron
     • Across a Period: increase in ionisation energies. As atomic number increases, there are more protons in the nucleus and electrons are being added to the same level with no additional shielding
       • This results in stronger electrostatic attraction between the nucleus and valence electrons, and more energy is needed to remove the electron
     • The more reactive a metal, the lower its IE
       • Metal atoms react by losing electrons, so the more easily they can lose electrons, the lower their IEs the more reactive they will be
     • Non-meals have a very high IE
       • These elements have atoms with more protons, giving a larger nuclear attraction on their valence electrons. The valence electrons will be closer to the nucleus and harder to remove, resulting in higher IEs
     • The first ionisaion energy of any atom always has a smaller value than its second, third etc, which have successively greater values. The second IE is the energy required to remove a second electron
       • Once the outer shell of electrons has been removed, the next ionisation energy recorded becomes proportionally much greater. This is because the attractoin exerted by the nucleus is much stronger on electrons in shells closer to the nucleus

3. ELECTRONEGATIVITY

   • A measure of the attraction between a nucleus and a bonded pair of electrons. This attraction that an atom has for its bonded electrons depends on: nuclear charge (no. of protons); distance between nucleus and bonded electrons (i.e atomic radius); shielding effect of inner shells of electrons
     • The greater the electronegativity value, the greater is the pull of the nucleus on the bonded electrons
   • Down a Group: decrease in electronegativity as the bonded electrons are in an energy level further from the nucleus, with increased shielding from inner shells
     • This results in a decreased electronegativity, despite the increase in nuclear charge
   • Across a Row: increase in electronegativity as the bonded electrons are in the same energy level but there is increased positive charge on the nucleus (more protons) so the attraction increases

4. POLARITY

   • Bond Polarity
     • Covalent bonds involve the sharing of at least one pair of electrons between two atoms. If the two atoms are different, they will have a different attraction for the bonded pair of electrons ie. two different enegs
     • When two atoms have different electronegativities, an uneven sharing of the bond pair of pairs of electrons occurs. This bond polarity occurs because: the more eneg atom has a stronger pull on the electrons - this atom becomes slightly negatively charged; or the less eneg atom has a weaker pull on the electrons - this atom becomes slightly positively charged
       • As a result of this charge separation, the bond forms a permanent dipole
   • Ionic-Covalent Bond Continuum
     • If the two atoms have very similar enegs, then the bond is a non-polar covalent bond, and the electron pair is shared equally between the two atoms
     • As the eneg difference increases, the polarity of the bond increases
     • If the difference between the enegs of two atoms is high enough, the atoms do not share electrons - the atoms transfer electrons to form ions
   • Diatomic Molecules
     • Non-polar if the two atoms have the same eneg and the bonded pair of electrons is shared equally
     • If the two atoms have different enegs the molecules will be polar
   • Polyatomic Molecules
     • For molecules containing more than two atoms and polar bonds, the molecule will be non-polar if the polar bonds are arranged symmetrically around the central atom so that the bond dipoles cancel out
     • Molecules with more than two atoms will be polar if the molecule has polar bonds arranged asymmetrically around the central atom so that the bond dipoles do not cancel. Overall, the polar molecule will have two oppositely charged regions - the molecule acts as a dipole

5. ENTHALPY

   • EXOTHERMIC
     • -
     • Energy released to the environment
     • More energy is given out by bonds forming than is taken in to break the bonds
   • ENDOTHERMIC
     • +
     • Energy is absorbed from the environment
     • More energy is needed to break bonds than is given out by bonds forming
   • In catalysed reactions, overall enthalpy changes are the same for uncatalysed reactions. It is the reaction pathways that are different
   • ENTHALPY OF REACTION
     • Change in enthalpy during a chemical reaction where the mole quantities of the chemicals and their states are specified by the reaction equation
       • Requires: pressure of 100kPa or 1 atm; all substances in their standard state
   • ENTHALPY OF FORMATION
     • The energy released or absorbed when: one mole of a product is formed from its constituent elements; or all measurements are made under standard conditions
   • ENTHALPY OF COMBUSTION
     • The energy released when: one mole of a substance is completely burnt in oxygen; or all measurements are made under standard conditions
   • Hess's Law
     • States that energy change due to a chemical reaction is independent of the route taken
       • H = products - reactants

6. BOND ENERGIES

   • The minimum energy (jJmol-1) required to break one mole of a specified bond in a gaseous molecule
   • Chemical reactions involve the breaking of bonds (in reactants) followed by the formation of bonds (in products): breaking chemical bonds requires energy and is endothermic; forming chemical bonds releases energy and is exothermic; the total energy of the bonds broken is the sum of the energies of all the individual bonds that break; the total energy of the bonds that form is the sum of the energies of all the individual bonds that form
   • The overall heat energy or enthalpy of a reaction can be expressed as: H = energy of the bonds broken - energy of the bonds formed