Chem- Bonding, structure and energy changes

• The four types of solids are ionic, metallic, molecular and covalent network.
• Ionic solids are made up of a metal and a non-metal. For example, sodium chloride.
• Properties of ionic compounds:

  • moderately high melting point
  • soluble in water
  • conduct when molten and in solution
  • does not conduct when solid
  • brittle
• Ionic structure is a lattice with positive and negative ions and strong electrostatic forces of attraction between them.
• Ionic solids have a moderately high melting points because the electrostatic forces of attraction between the bonds is strong
• Ionic substances do not conduct when solid because the ions are held in a lattice and not free to move (no free electrons)
• Ionic substances conduct when molten because the forces of attraction between the ions are broken and the ions are free to move.
• Ionic substances are brittle because attempts to move ions in the structure would bring like charges closer together, and the repulsion caused by this would cause the structure to crumble or split.
• Ionic substances conduct when in solution because the ions are surrounded by water molecules and are free to move.
• Ionic substances are soluble in water because the water molecules are able to disrupt the attractive forces between the ions. The ions become surrounded by water molecules and the lattice is disrupted.
• Metallic solids are metals, e.g. aluminium, zinc, lead, copper, iron.
• Properties of metallic solids:

  • ductile
  • malleable
  • shiny
  • conduct electricity
  • conduct heat
• In metallic structures, the metal atoms are held together by electrostatic forces between the nuclei and the sea of electrons. The delocalised electrons are the valence electrons.
• In metallic structures there are metallic bonds and non-directional forces.
• Metals conduct electricity because they have electrons that are free to move.
• Covalent network solids are made up of atoms, and contain covalent bonds.
• Diamond is a covalent network solid made up of carbon atoms. There are four covalent bonds between each carbon atom and four others.
• Diamond is insoluble in water and all solvents, very hard, has a very high melting point, and does not conduct electricity (because there are no free electrons or ions)
• SiO2 is a covalent network solid, with a high melting point that does not conduct
• Graphite is a covalent network solid with a layer structure. There are three covalent bonds between eat carbon atom, and so there is one electron from each atom that is free to move (delocalised), so graphite can conduct electricity. Graphite has a high melting point because the energy required to break the covalent bonds is very high. Graphite is insoluble and soft and slippery.
• Molecular substances are made up of molecules, e.g. iodine or carbon dioxide.
• Properties of molecular substances:

  • Weak intermolecular force
  • Low melting points
  • Do not conduct electricity
  • Insoluble
• In ionic bonds, electrons have been transferred and positive and negative ions have been formed. The ionic bond is the electrostatic attractive force between the negative and positive ions.
• In covalent bonds, electrons are shared and positive nuclei are attracted to the electrons between them.
• Electronegativity is the ability of an atom in a covalent bond to attract electrons towards itself. Fluorine is the most electronegative element.
• When there is a difference in electronegativity, electrons in a covalent bond are not shared equally and the more electronegative atom attracts the electron density towards itself. One end of the bond is more negative than the other- a molecular dipole is formed.
• In polar bonding, there is an unequal sharing of electrons and partial charges.
• The shapes of molecules are determined by the way clouds of electrons are arranged around the central atom in the molecule.
• When there are two clouds of electron density, the clouds are arranged on opposite sides of the central atom, so the shape is linear and the bond angle is 180 degrees.
• VSEPR: 'Regions of electron density are all negatively charged, so they will repel each other therefore they will be arranged as far away from each other as possible, putting them in the lowest energy, most stable stat possible."
• Bonding electrons and lone pairs both affect the shape of a molecule. Lone pairs occupy more space= greater repulsion, so lone pairs push bonding electrons closer together.
• Enthalpy, H, is the energy change in a chemical reaction.
• Exothermic reactions produce heat. In exothermic reactions, the change in enthalpy is negative.
• Endothermic reactions absorb heat. In endothermic reactions, the change in enthalpy is positive.
• In exothermic reactions the products are more stable and have a lower potential energy (than the reactants).

  In endothermic reactions, the products are less stable and have a higher energy.
• Increasing the energy by heating causes the intermolecular forces to weaken and break, so melting and evaporation are endothermic processes.
• Decreasing the energy by cooling causes intermolecular forces to form, so condensation and solidification are exothermic processes.
• Bond breaking requires energy (endothermic) and bond making releases energy (exothermic).
• Bond energies measure how tightly a pair of atoms are held together. The bond energy is influenced by the size of the nuclei, the number of electrons involved in the bond and the type of atoms.
• The change in enthalpy is the sum of energy of all of the bonds broken - sum of energy of all bond formed