bonding

Annotations:

• an ionic bond is an electrostatic force of attraction between oppositely charged ions caused by electron transfer.

1. electrostatic attraction holds the positive and negative ins together very strongly.
2. the formula for this compound tells you exactly what's in it.
   1. the positive charges balance the negative charges so the overall charge is zero.
3. giant ionic lattices
   1. ionic crystals are giant lattices of ions, a lattice is just a regular structure.
      1. the structures giant as its made up of the same repeated unit over and over.
4. behaviour of ionic compounds
   1. electrical conductivity
      1. ionic cmpounds only conduct when molten or dissolved but not when solid. this is because theres no ions that are free to carry the charge
   2. melting point
      1. ionic compounds have high melting point. the giant ionic lattices ar held together by strong electrostatic forces and it requires a lot of enrgy to overcome these hence the higher melting point.
   3. soluabiity
      1. ionic compounds tend to dissolve in water. water molecules are polar. the water molecules pull the ions away from the lattice causing it to dissolve

1. molecules
   1. they're the smallest part of a compound that can take part in chemical reactions.theyre formed when two or more atoms bond together. the atoms can be the same (cl2) or different (CO)
   2. molecules are held together by very strong covalent bonds
2. single bonds
   1. in covalent bonding two atoms share electrons, so theyve both got a 'full' outer shell of electrons. both the positie nuclei are attracted electrostatically to the shared electrons.
3. double or triple bonds
   1. double bond can be formed by the one carbon atom (c) which can bond to two oxygens (O) each oxygen atom shares two pairs of electrons with the carbon atom, so each molecule contains two double bonds.
   2. triple bonds
      1. a triple bond occurs in nitrogen (n2) the nitrogen atom shares three pairs of electrons so each nitrogen contains one triple bond
4. behaviour between simple covalent compounds
   1. simple covalent compound
      1. they have STRONG bonds WITHIN molecules but weak forces BETWEEN molecules. their physyical properties such as conductiving electricity are determined by the bonding in the compound
   2. electricial conductivity
      1. simple covalent compounds dont conduct electricity because there are no free ions to carry a charge
   3. melting points
      1. they have low melting points as the weak forces between moleules are easy to break
   4. soluability
      1. some of these compounds dissolve in water depending on how polarised the molecules are.
5. giant covalent structures
   1. graphite
      1. carbon atoms are arranged in flat hexangonal sheets. each carbon atom is bonded to THREE OTHERS and the fourth 'free' electron goes inbetween the layers causing an delocalised electron.
      2. what does graphites structure mean?
         1. the weak bonds between the layers allow the sheets to slide over eachother- it feels slippery and acts as a lubricant
         2. the delocalised electrons are free to move between the sheets and therefore can carry a charge
         3. as the layers are quite far apart it ha a low density and is used to make strong lightweight sport material
         4. graphites insoluable in any solvent as the covalent bonds are to difficult to break
   2. diamond
      1. diamonds made up of carbon atoms, each carbons covalently bonded to 4 others.
      2. the atoms arrange themselves in a tetrahedral shape.
      3. because of the strong covalent bonds...
         1. diamond has a very high melting point.
         2. diamonds extremely hard
         3. vibrations can travel easily through the stiff lattice, so its a god thermal conducltor.
         4. it cant conduct electricity as all the outer electrons are held in ionsied bond.
         5. and ike graohite it doent dissolve,

1. a dative bond is a normal single covalent bond , however instead of one atom donating each electron, both electrons are fro the same atom
2. you may need to idenitfy which atom in the question is the 'donor' and which is the 'reciever' . the donor has a LONE PAIR OF ELECTRONS and the reciever has a VACANT ORBITAL.
3. FOR EXAMPLE;
   1. the hydroxonium ion (H30+)
      1. first look at the reactants involves; in this case its H20 and H+
         1. a water molecule contains 2 hydrogens and 1 oxygen molecule
            1. oxygen has six electons in its outer shell and is filled with one electron each from the hydrogen atom
               1. only four of these 6 electrons take part in covalent bonding
                  1. therefore the extra 2.. otherwise known as a 'LONE PAIR' are donated to the hydrogen ion. forming a DATIVE COVALENT/ CO-ORDINATE BOND

1. metal elements exist as giant metallic lattice sturctures. the outerost shell of electrons of a metal atom is delocalised ( the electrons are free to move about the metal)
2. this leaves a positive metal ion which are attracted to the delocalised negative electrons. they form a lattice of closely packed positive ions in a sea of delocalised electrons.
3. BEHAVIOUR OF METALLIC COMPOUNDS
   1. melting point
      1. the number of delocalised electrons per atom effects the melting point. the more electrons the stronger the bonding will be and the higher the melting point, for example mg2+ has two delocalised electrons per atom so its melting point will be higher than na+
      2. the size of the metal ion and the lattice structure also effects melting point
   2. ability to be shaped
      1. as there are no specific bonds holding the atoms together the metals ions can slide over eachother when the structures pulled, so the metals a malleable and ductile

         Annotations:

         • malleable means they can be hammered into shape and ductile means the can be drawn into thin wires
   3. conductivity
      1. as the delocalised electrons can pass kinetic energy to each other it makes metas good thermal conductors and the electrons can carry a current.
   4. soluability
      1. metals are insoluable except in liquid metals due to the strength of the metallic bonds