Bonding, structure, and the properties of matter

• Chemical bonds 


  Bonds between atoms in an element or compound
• Ionic bonding 


  -Between metal and non-metal atoms

  -the particles are oppositely charged ions

  -involves transfer of electrons
• Covalent bonding 


  -Between non-metal & non-metals

  -involves sharing an electron

  -can be two of the same element
• Metallic bonding 


  -in metal elements and alloys

  -involves sharing delocalised electrons

  -between atoms in a metal
• Ionic bonding 


  -electrons are transferred from the metal

  -metal atoms loose electrons and become positively charged ions

  -non-metal atoms gain an electron and become negatively charged ions
• Explain what happens when sodium and chlorine react 


  -there is a transfer of electron from sodium to chlorine -this is so both atoms gain a full outer shell -the sodium atom becomes a sodium ion and the chlorine becomes a chloride ion
• - because the sodium ion has a positive charge and the chloride ion has a negative charge those two will attract - this is called an electrostatic force of attraction
• Ionic compounds 


  -giant lattice structures

  -held together by strong electrostatic forces of attraction between oppositely charged ions

  -these forces act in all directions in the lattice
• Properties of ionic compounds

  -high melting & boiling points

  -because a lot of energy is needed to break the many strong ionic bonds

  -cannot conduct electricity when they're in solid form because ions are in fixed positions and not free to move (no free electrons either)
• When can ionic compounds conduct electricity?

  -when they are dissolved in water or molten -this is because ions are free to move so charge can flow
• When can ionic compounds conduct electricity?


  -when dissolved water or melted (in molten form)

  -because ions are free to move do charge can flow
• Covalent bonds 


  -bonds between atoms are strong

  -all molecules have covalent bonds
• Chlorine gas (Cl2) •Cl--Cl Represents single covalent bond Hydrogen chloride (HCl) H--Cl
• Water (H2O) O--H--O
• Water (H2O) & Oxygen (O2)


  --share 2 pairs of electrons (double covalent bond)
• Nitrogen (N2) and Ammonia (NH3)

  -shared 3 pairs of electrons (triple bond) -3 single bonds between 3 hydrogen atoms and 1 Ammonia atom
• Metallic bonding 


  -consist of giant structures of atoms arranged in a regular pattern

  -electrons in the outer shell of metal atoms are delocalised and so are free to move through the whole structure

  -
• Solids

  -particles are closed together in rows and columns -they vibrate in fixed positions
• Liquids

  -particles are in a less regular arrangement and free to move around
• Gases 


  Particles are far apart, moving in random directions at different speeds
• Melting and freezing take place at the melting point, boiling and condensing take place at the boiling point.
• Change in states of matter

  Solid to liquid - Melting Liquid to gas -boiling Gas to liquid - condensing Liquid to solid - freezing
• -particles gain energy when a substance is heated and loose energy when it's substance is cooled -the stronger the forces between the particles in a substance, the higher the melting & boiling points
• Limitations 


  -in a model the particles are shown as solid spheres where as in reality they are atoms molecules & ions -in a model the particles don't show the forces between them. In reality particles have forces between them or are bonded together -particles aren't shown as bonds in the model but different bonds exist
• Melting & boiling points of halogens 


  Iodine MB- 114•C BP-184•C Solid Bromine MB -7•C BP 58•C Liquid Chlorine MB -101•C BP -35•C Gas Flouring MB -220•C BP -118•C Gas
• Small molecules 


  -substances made of small molecules have relatively low melting and boiling points

  E.g oxygen melts at -219 degrees and boils at -183 degrees
• Why do small molecules have low melting & boiling points?

  -because the intermolecular forces of attraction between the molecules are weak

  -so only a small amount of energy is needed to overcome these forces
• Intermolecular forces increase with molecule size

  -larger molecule have higher bp
• -Intermolecular forces are weak -covalent bonds are strong

  -intermolecular forces are between molecules

  -covalent bonds are between atoms in the molecules
• -Small molecules have no overall electrical charge & no free electrons

  -so the substance e.g oxygen cannot conduct eke o
• Polymers 


  -have very large molecules

  -strong covalent bonds between atoms in polymers

  -intermolecular forces are relatively strong

  -making these substances solids at room temperature
• Giant covalent structures 


  -these are solids with very high melting and boiling points because

  -all the atoms are linked together by strong covalent bonds

  -requires a lot of energy to break bonds

  Examples include: Diamond and graphite & silicone dioxide
• Diamond 


  -made of only carbon atoms

  -very hard because it has a husky lattice structure
• Diamond properties 


  -each carbon atom forms 4 covalent bonds with other carbon atoms so diamond is very hard

  -has a high melting point and does not conduct electricity (no free electrons)
• Graphite 


  -each carbon atom forms 3 covalent bonds with 3 other carbon atoms

  -this forms layers of hexagonal rings which have no covalent bond between the layers
• properties of graphite

  
  

  -there weak forces of attraction between the layers of hexagonal rings

  -this allows the layers to slide over each other very easily

  -this makes graphite soft & slippery

  -conducts electricity because it has delocalised electrons that can flow through the structure and carry charge
• Silicon dioxide 


  -hard because it has a giant lattice structure

  -each silicon atom is covalently bonded to 2 oxygen atoms

  -has no free electrons so cannot conduct electricity
• Structure and properties of metals

  -metals have giant structures of atoms

  -electrons in the outer shell are delocalised so are free to move

  -sea of delocalised electrons holds metal ions together & makes strong metallic bonds
• Metals 


  -can conduct electricity

  -because the delocalised electrons can carry charge through the metal

  -also good conductors of heat energy because energy is transferred by those delocalised electrons
• Pure metals

  -atoms are arranged in layers which can slide over each other

  -this allows metals to be bent and shaped