bonding and structure

• States of matter


  • Solid
  • Liquid
  • Gas
• Melting point


  Temperature at which solid and liquid states are in equilibrium
• Boiling point


  Temperature at which liquid and gas states are in equilibrium
• The amount of energy needed to change state depends on the strength of the forces between the particles of the substance
• The nature of the particles involved depends on the type of bonding and the structure of the substance
• The stronger the forces between the particles the higher the melting point and boiling point of the substance
• The simple model of particles represented by small solid spheres has limitations as atoms, molecules and ions are not solid spheres and there are forces between the particles
• Gas particles


  • Widely spaced and in constant quick motion
  • Collisions are frequent and elastic
  • Weak forces between particles
• Liquid particles


  • Closely spaced but still in constant motion, and therefore are constantly colliding
  • Forces between particles less than in solid
• Solid particles


  • Can only vibrate in a fixed position
• Pure elements and compounds melt and boil at specific temperatures
• Melting point and boiling point data can be used to distinguish pure substances from mixtures
• Melting point curve


  A pure solid will melt at a fixed temperature and the line will stay horizontal when it is melting. The temperature does not rise when the solid is melting because the heat is absorbed to break the bonds between the solid particles
• Chemical equations


  The three states of matter are shown as (s), (l) and (g), with (aq) for aqueous solutions
• Ionic bonding


  1. Metal atoms lose electrons to become positively charged ions

  2. Non-metal atoms gain electrons to become negatively charged ions
• Ionic bonding


  Occurs in compounds formed from metals combined with non-metals
• Dot and cross diagram

  • For sodium chloride

  
• Ionic bond


  The force of attraction between the oppositely charged ions
• The ions produced by metals in Groups 1 and 2 and by non-metals in Groups 6 and 7 have the electronic structure of a noble gas (Group 0)

  
• Common charges of ions

  • Mg2+
  • Ca2+
  • Cl-
  • O2-
• Formulae for ionic compounds

  The total positive charge of the ions is the same as the total negative charge of the ions – giving a neutral compound
• When to use brackets in a formula


  If there are two or more of a compound ion in a formula we put the ion in brackets and put the number outside the bracket
• Giant ionic structure


  An ionic compound is a giant structure of ions held together by strong electrostatic forces of attraction between oppositely charged ions
• Properties of ionic substances

  • High melting points and high boiling points
  • Do not conduct electricity when solid
  • Conduct electricity when melted or dissolved in water
• Empirical formula


  The simplest ratio of ions in a compound
• Compound ions


  • Carbonates CO3^2-
  • Sulfates SO4^2-
  • Hydroxides OH-
  • Nitrates NO3-
  • Ammonium NH4+
• Covalent bonding


  Occurs in non-metallic elements and in compounds of non-metals. Atoms share pairs of electrons to form covalent bonds
• Covalently bonded substances


  May consist of small molecules or giant covalent structures
• Size of molecule


  Intermolecular forces increase with the size of the molecules, so larger molecules have higher melting and boiling points
• Properties of simple molecular covalent substances

  • Usually gases or liquids with relatively low melting and boiling points
  • Have only weak intermolecular forces between the molecules
  • Do not conduct electricity
• Polymers


  Have very large molecules with atoms linked by strong covalent bonds. Intermolecular forces between polymer molecules are relatively strong so they are solids at room temperature
• Properties of giant covalent substances

  • Solids with very high melting points
  • All atoms linked by strong covalent bonds
• Giant covalent structures


  • Diamond
  • Graphite
  • Silicon dioxide
• Diamond


  Each carbon atom forms four covalent bonds with other carbon atoms in a giant covalent structure, so diamond is very hard, has a very high melting point and does not conduct electricity
• Graphite


  Each carbon atom forms three covalent bonds with three other carbon atoms, forming layers of hexagonal rings. The layers are free to slide over each other because there are no covalent bonds between the layers, so graphite is soft and slippery. One electron from each carbon atom is delocalised, allowing graphite to conduct thermal energy and electricity
• Graphene


  A single layer of graphite, one atom thick
• Fullerenes


  Molecules of carbon atoms with hollow shapes, based on hexagonal rings of carbon atoms but may also contain rings with five or seven carbon atoms

  
• Carbon nanotubes


  Cylindrical fullerenes with very high length to diameter ratios, with useful properties for nanotechnology, electronics and materials
• Metallic bonding


  Consists of a giant structure of atoms arranged in a regular pattern, with delocalised electrons in the outer shell of metal atoms that are free to move through the whole structure. The sharing of delocalised electrons gives rise to strong metallic bonds

  
• Properties of metals


  • Strong metallic bonding means most metals have high melting and boiling points
  • Metals are good conductors of electricity and thermal energy due to the delocalised electrons