1.4-bonding

• ionic bonds


  the electrostatic force of attraction between cations and anions formed when one atom gives one or more electrons to another

  (the net balance of the attractive and repulsive forces determine the intermolecular distance between ions and the strength of the bond)
• repulsive forces in ionic bonds

  1) ions of the same charge

  2) inner shells of electrons repelling the outer shell electrons

  3) the positively charged nuclei repel each other
• covalent bonds

  an electrostatic attraction between the negatively charged electron pair in the bond and the positively charged nuclei of the elements (electrons are shared between atoms)
• repulsive forces in covalent bonds

  1) the electrons in the pair between the atoms repel each other, but this is overcome by their attractions to both nuclei, and the fact that they spin in opposite directions

  2) if atoms get too close, the nuclei and their inner electrons will repel each other so the bond has a certain length
• coordinate bond


  a covalent bond in which both electrons in the shared pair come from the same atom (e.g. NH4)
• metallic bonding


  - giant lattice of positive ions

  - held together by a sea of delocalised electrons

  - electrons are free to move and carry charge/electrical current

  - strong forces/bonds between the delocalised electrons and positive ions requires a lot of heat energy to overcome
• electronegativity


  a measure of the ability of an atom in a covalent bond to attract a bonding electron
• why does electronegativity increase across a period

  - there is an increase in the atomic number (number of protons) and the charge on the nucleus

  - this attracts bonding electrons
• Pauling electronegativity index


  allows you to work out the difference in electronegativity for 2 atoms joined together by a covalent bond

  - less than 0.3 = covalent

  - between 0.3 and 1.8 = polar covalent

  - greater than 1.8 = polar
• electronegativity difference of < 0.3

  covalent
• electronegativity difference of between 0.3 and 1.8

  polar covalent (permanent dipole - electrons are not equally shared)
• electronegativity difference of > 1.8

  polar

  
• intermolecular bonding


  weak bonding between MOLECULES and governs the physical properties of the substance

  Caused by attraction between opposite charges

  (think inter-national = different countries, like the different molecules)

  
• what type of bonding governs the melting point, boiling point, density, solubility and polarity

  intermolecular bonding
• examples of intermolecular bonding

  Hydrogen bonding, van der Waal forces (dipole-dipole, induced dipole-induced dipole)
• intramolecular bonding


  strong bonding between ATOMS in the molecule and governs its chemistry

  (intrA = Atoms)

  
• examples of intramolecular bonding

  Ionic bonding, covalent bonding, metallic bonding
• what type of bonding governs a molecule's reactivity

  intramolecular bonding
• van der Waals forces

  a slight attraction that develops between the oppositely charged regions of nearby molecules

  (includes all types of intermolecular force, whether dipole or induced dipole)
• permanent dipole


  a small charge difference across a bond that results from a difference in the electronegativity of the bonded atoms
• induced dipole


  electrons are in constant motion around the nuclei so that the centres of positive and negative charge do not always coincide and give a fluctuating dipole

  one dipole can induce an opposite dipole in a nearby molecule

  
• State how a coordinate bond differs from a covalent bond

  in a coordinate bond both bonding electrons come from one atom while in a covalent bond both atoms contribute one electron in the shared pair
• state what is meant by a polar covalent bond

  a covalent bond where the electrons are not shared equally between the atoms (unequal electron density) because of differences in electronegativity between them
• explain why there is a general increase in boiling point as you descend a group

  there is an increase in atomic number, leading to an increased number of electrons. This means the strength of the Van Der Waals forces increases
• how would the boiling point of alkenes compare to the boiling point of alcohols

  - Alcohols will have a higher boiling point because they contain hydrogen bonds between O and H atoms in addition to Van Der Waals forces

  - More energy is needed to break the hydrogen bonds and separate the atoms
• why does propane C3H8, have a higher boiling point than methane, CH4

  there are more atoms and therefore more electrons, leading to more Van Der Waals forces
• why are smaller alcohols such as methanol and ethanol more readily soluble than larger alcohols

  - the smaller alcohols can dissolve readily in water because the hydroxyl group can form hydrogen bonds with water molecules

  - they also have a very small carbon chain

  - larger alcohols become more insoluble because they have a longer carbon chain, which is hydrophobic (non-polar)

  - it requires more energy to overcome the hydrogen bonds between the alcohol molecules as the molecules are more tightly packed together due to the size and mass increasing
• The boiling temperatures of 1,2-diaminoethane and 2-aminoethanol are shown below

  
  

  1,2-diaminoethane/ H2N-CH2-CH2-NH2/ 117°C

  2-aminoethanol/ H2N-CH2-CH2-OH/ 170°C

  Use these figures to comment on the strength of the intermolecular forces between the molecules in each compound suggesting reasons for your answer.


  - The forces of attraction between molecules of 2-aminoethanol are stronger than in molecules of 1,2-

  diaminoethane as the former has a higher boiling temperature

  - This suggests that intermolecular hydrogen bonding between O and H is stronger than the hydrogen bonding

  between N and H

  - This is likely due to a greater electronegativity difference

  between O and H than between N and H (O is more electronegative than N)
• Nitrogen chloride, NCl₃, is insoluble in cold water whilst the similar compound ammonia, NH₃, is very soluble. Explain this difference in behaviour

  NH₃ can form hydrogen bonds with water molecules, so it dissolves, but NCl₃ cannot form hydrogen bonds
• hydrogen bonding


  results from the attractive force between a hydrogen atom covalently bonded to a very electronegative atom such as a N, O, or F atom and another very electronegative atom

  
• Explain, in terms of bonding, why iodine is less volatile than bromine

  iodine is a larger molecule and therefore there are more electrons, leading to greater Van Der Waals forces between molecules
• what does this arrow ↷ represent

  a shift of 2 electrons
• explain why aluminium has a higher melting temperature than sodium. You should refer to the nature of the bonding

  the bonding is metallic

  there is attraction between the sea of delocalised electrons and the positive ions

  Al³⁺ has a higher charge density than Na²⁺ and therefore has more electrons in the sea

  More energy is needed to overcome the forces in Al
• isotopes are atoms with the same number of protons but different numbers of neutrons
• atomic mass = sum of protons + neutrons