TOPIC 2-BONDING

• Ionic bonding is the strong electrostatic force of attraction between oppositely charged ions in a giant lattice.
• strenght of ionic bond depends on:

  • smaller size of ion
  • higher charge
  • so more energy required to overcome attraction - high m.p
• properties of ionic b:

  • hard, brittle (layers slide and ++ions aligned and repel -so structure breaks)
  • high m.p bc strong attraction
  • soluble in water
  • do not conduct when solid bc ions can't move from fixed position in the giant lattice- can't carry charge
• Ionic radii:

  • radius of +ve ion is smaller than its atomic radius bc loses e-. so it has more protons than e- (so greater pull over e-)
  • radius of -ve ion is larger bc an e- is added. so more e- than protons
• ionic radius down a group: increases bc n' od e- shells increases

  ionic radius in isoelectric ions: attractive force from nucleus increases - this pulls outer e- more

  A) protons

  B) e-

  C) increases

  
• Ionic Bonding - Evidence for the existence of ions

  A) e-

  B) 0

  C) radius

  D) regular

  E) larger

  
• Migration of ions

  green copper chromate on wet filter paper with electricity - ions starts to separate

  A) Cu2+

  B) CrO4 2-

  C) +ve

  D) -ve

  
• A covalent bond is a strong electrostatic attraction between the two nuclei and a shared pair of e-.
• dative covalent bond- a bond where one atom donates both e- in a shared pair of e-.

  A) deficient

  B) dimer

  C) tetrahedral

  D) dative

  
• Bond strength and length- Nuclei joined by (i.e. double and triple) bonds have a greater electron density between them. This causes a greater force of attraction between the nuclei and the electrons between them, resulting in a shorter bond length and greater bond strength.

  As the n' of covalent b increases, the bonds becomes shorter and stronger.
• bond strength

  A) attraction

  B) below

  C) further

  D) less

  E) stronger

  F) overlap

  
• bond strength/length

  A) shells

  B) shielding

  C) stronger

  
• Q. on bond strength/length

  A) decreases

  B) increases

  C) nuclear

  D) close

  E) repel

  F) reduce

  
• What is the arrangement of carbon atoms in diamond?

  Tetrahedral arrangement
• How many covalent bonds does each carbon atom form in diamond?

  4 covalent bonds
• Why can't diamond conduct electricity?

  Because the electrons in covalent bonds are in a fixed position and cannot move
• What is the reason for diamond's high melting point and hardness?

  It has many strong covalent bonds, requiring a lot of energy to break the lattice
• How many covalent bonds does each carbon atom form in graphite?

  3 covalent bonds
• What happens to the fourth outer electron in graphite?

  It becomes delocalized
• Why can graphite conduct electricity?

  Because it has delocalized electrons that can move
• Why is graphite insoluble in water?

  Because the covalent bonds are too strong to break
• What contributes to graphite's high melting point?

  It has many strong covalent bonds that require a lot of energy to break
• What type of forces exist between the layers of graphite?

  Weak intermolecular forces
• How thick is graphene?

  1 atom thick
• How many covalent bonds does each carbon atom form in graphene?

  3 covalent bonds
• Why does graphene have high tensile strength?

  Because it has many strong covalent bonds
• What type of bonding is present in simple molecular structures?

  Weak London forces
• Why do simple molecular substances have low boiling points?

  Because less energy is needed to break the weak bonds
• Why can't simple molecular substances conduct electricity?

  Because they have no free electrons
• How do the properties of diamond differ from those of graphite?

  Diamond is hard and does not conduct electricity, while graphite is soft and can conduct electricity
• How does graphene compare to graphite in terms of conductivity and strength?

  Graphene conducts electricity better and has higher tensile strength than graphite
• What distinguishes simple molecular structures from giant covalent structures?

  Simple molecular structures consist of discrete molecules with weak intermolecular forces, while giant covalent structures have a continuous network of strong covalent bonds
• Metallic bonding is the electrostatic force of attraction between the nuclei of metal ions and the delocalised e-

  • along a period m.p increases bc: ionic radius decreases, n' of delocalised e- per atom increases, so attraction btw nucleus of ions and deloc e- increases, so more E needed to break
  • down a group m.p decreases bc: ionic radius increases, n'of del e- stays the same, so attraction btw nucleus of ions and deloc e- decreases, so less E needed
• properties of metallic lattice:

  • high m.p bc strong electrostatic forces btw +ve ions and sea of deloc e- require a lot of E to break
  • conduct electricity bc e- can move through the lattice and carry a charge
  • malleable bc layers can slide over each other when force is applied
• Q abt metal

  Mg has stronger metallic b and higher m.p than Na bc in Mg more e- are in the sea of deloc e-. Mg is smaller and has 1 more proton (higher charge ) than Na. so stronger eletrostatic attraction btw +ve metal ion and del e- , so more E needed
• Electronegativity is the ability of an atom to attract electrons the bonding e- in a covalent bond to itself.

  • EN increases across a period: nuclear charge increases(n' of protons inc), atomic radius decreases, shielding stays the same (so e- in the same shell are pulled more)
  • EN decreases down a group: nuclear charge increases,atomic radius increases (distance btw nucleus and outer e- inc - so less attraction), shielding increases.
• ELECTRONEGATIVITY: the bigger difference in EN, the more ionic and more polar. A difference of zero=covalent
• POLAR BONDS: Cl is more electronegative than H, so pulls/attracts more e- towards itself. a polar covalent bond it has an unequal distribution of electrons in the bond and produces a charge separation, (dipole) δ+ δ- ends.

  • bonds in CO2 are polar bc O is more electronegative than C, nute the molecule is not polar bc polarities/dipole moments cancel out.

  
• investigating polarity

  A) electrostatic force

  B) align

  C) deflection

  D) charged

  
• LONDON FORCES: as you go down the group increases:

  • n' of e- increases
  • polarisability increases
  • so the strength of LF increases and more E needed to break - m.p increases

  temporary dipoles form when e- move from one end to another (uneven distribution). this formes induced dipoles in neighbouring molecules. --- temporary dip only exist when 2 molecules are near by