• Ionic Bonding
• Ions


  Ions are positively or negatively charged atoms/molecules
• Ion formation


  Ions form when electrons are transferred from one atom to another

  
  

  The atom gains or loses electrons to make its outer shell full

  
  

  Cations (positive ions) are formed when an atom loses 1 or more electrons

  
  

  Anions (negative ions) are formed when an atom gains 1 or more electrons
• Ionic Bonding


  An ionic bond is the strong electrostatic attraction between oppositely charged ions
• Effect of ionic charge on ionic bonding

  Generally, greater charge → stronger ionic bond

  therefore greater charge means higher melting/boiling points

  e.g. NaF (Na⁺ and F⁻) has melting point 993°C

  and CaO (Ca²⁺ and O²⁻) has melting point 2572°C
• Effect of ionic radii on ionic bonding

  Smaller ions can pack closer together; electrostatic attraction decreases with distance so small, tightly packed ions have stronger ionic bonding than larger, loosely packed ions

  
  

  e.g. ionic radius of Cs⁺ > ionic radius of Na⁺;

  CsF has melting point 683°C, NaF has melting point 993°C because Na⁺ and F⁻ can pack closer together than Cs⁺ and F⁻ can
• Ionic radius increases down a group


  As you go down a group, the ionic radius increases as the atomic number increases; this is because an extra shell is added each time
• Isoelectronic ions


  Ions of different atoms with the same number of electrons e.g. N³⁻, O²⁻, Mg²⁺ and Al³⁺ all have 10 electrons

  
  

  Ionic radius of a set of isoelectronic ions decreases as atomic number increases because of increased nuclear attraction on same number of electrons
• Giant Ionic Lattices


  Ionic crystals, e.g. NaCl are giant ionic lattices

  Lattice = regular structure

  Giant = lots of the same repeat unit

  
  

  Form because ions are attracted to oppositely charged ions in all directions
• Theory of Ionic Bonding fits with evidence from Physical Properties

  1. High melting points: shows that ions are held together by strong attractions; +ive and -ive ions are strongly attracted; fits evidence

  
  

  2. Are usually soluble in water but not in non-polar solvents; shows that the particles are charged because they can be pulled apart by polar molecules like water

  
  

  3. Do not conduct electricity when solid; ions and electrons are strongly held in place when solid, but free to move when dissolved or molten

  
  

  4. Ionic compounds are brittle; moving layers over each other would cause like-charged ions to go near each other, which have strong repulsion, causing the compound to break apart
• Migration of ions is evidence for presence of charged particles

  Electrolysing green copper(II) chromate(VI) on wet filter paper, paper turns blue at cathode (negative electrode) and yellow at anode (positive electrode)

  
  

  Copper(II) ions are blue in solution

  Chromate(VI) ions are yellow in solution

  Copper(II) chromate(VI) is green in solution because it contains both

  
  

  When a current is passed through, positive ions move to cathode and negative ions move to anode
• Covalent Bonding
• Covalent Bond


  A covalent bond is formed when two atoms share electrons so they both have a full outer shell

  
  

  A covalent bond is the strong electrostatic attraction between the two positive nuclei and the shared electrons in the bond
• Bond Enthalpy relates to Bond length


  In a covalent bond there is strong attraction, but there is also a degree of repulsion because the positive nuclei repel each other, as do the shared electrons

  
  

  Bond length is the distance at which the attraction and repulsion are balanced

  
  

  Greater electron density (more electrons in the bond); greater attraction; shorter bond length; higher bond enthalpy
• Dative Covalent (coordinate) Bonding is where Both Electrons come from the same atom


  Both electrons in a bond come from the same atom e.g. CO or ammonium (NH₄⁺) which forms when the N atom in ammonia shares an electron pair with H⁺

  
  

  Ammonium can then form ionic bonds with other ions

  
  

  e.g. Al₂CL₆; AlCl₃ is a stable covalent compound where the central atom does not have a full outer shell, but sometimes 2 AlCl₃ can join to make Al₂Cl₆; the lone pair on a Cl atom is donated to the Al atom in each molecule, forming a coordinate bond
• Shapes of Molecules
• Molecular Shape


  Depends on Electron Pairs around the central atom; the number of pairs and the type of pairs
• Electron pairs repel each other (effect on molecular shape)

  
  

  Predicting shapes = electron pair repulsion theory


  All electrons are negatively charged so repel strongly

  
  

  Lone pairs repel more than lone pairs; greatest angles are between lone pairs, smallest angles are between bonding pairs, but these are often reduced further by lone-pair repulsion

  
  

  Greatest angle → smallest angle

  Lone/lone > lone/bonding > bonding/bonding

  
  

  Shape of molecule depends on type and number pairs
• Electron Pair Repulsion Theory

  Steps:

  1. Find the central atom

  2. Calculate number of electrons in outer shell of central atom

  3. Molecular formula shows how many atoms the central atom is bonded to

  4. Add up electrons and divide by 2 to find number of pairs

  5. Compare number of pairs to number of bonds to find number of lone pairs (double bonds count as 2 bonds)

  6. Use the number of bonding pairs and lone pairs to find shape of molecule
• 2 Electron pairs around central atom


  Shape: Linear

  Angle: 180°

  e.g. BeF₂ and CO₂

  F-Be-F or O=C=O

  
• 3 Electron pairs around central atom

  No lone pairs:

  Shape: trigonal planar

  Angles: 120°

  e.g. BCl₃

  
  

  1 lone pair:

  Shape: non-linear/bent

  Angle: 119°

  e.g. SO₂
• 4 Electron pairs around central atom

  No lone pairs:

  Shape: tetrahedral

  Angles: 109.5°

  e.g. NH₄⁺

  
  

  1 lone pair:

  Shape: trigonal pyramidal

  Angles: 107°

  e.g. PF₃

  
  

  2 lone pairs

  Shape: non-linear/bent

  Angle: 104.5

  e.g. H₂O
• 5 Electron pairs around central atom

  No lone pairs:

  Shape: trigonal bipyramidal

  Angles: 120° and 90°

  e.g. PCl₅

  
  

  1 lone pair:

  Shape: seesaw

  Angles: 2x 87° and 1x 102°

  e.g. SF₄

  
  

  2 lone pairs:

  Shape: distorted T

  Angle: 87.5°

  e.g. ClF₃
• 6 Electron pairs around central atom

  No lone pairs:

  Shape: octohedral

  Angles: all 90°

  e.g. SF₆

  
  

  1 lone pair:

  Shape: square pyramidal

  Angles: 90° and 81.9°

  e.g. IF₅

  
  

  2 lone pairs:

  Shape: square planar

  Angles: 90°

  e.g. XeF₄
• Giant Covalent and Metallic Structures
• Some covalently bonded substances have giant structures


  Covalent bonds form when atoms share electrons with other atoms; often leads to small molecules like CO₂ but can often lead to formation of huge lattices containing billions of atoms

  
  

  These giant covalent lattices are made up of many atoms all covalently bonded to each other; the bonds holding these structures together are much stronger than the electrostatic attractions between simple covalent molecules

  
  

  Carbon and silicon can form giant covalent lattices

  e.g. diamond is carbon atoms bonded in a tetrahedral arrangement
• Properties of giant structures provide evidence for covalent bonding

  They have very high melting points; lots of very strong bonds must be broken

  
  

  Are extremely hard; strong bonds throughout lattice arrangement

  
  

  Good thermal conductivity; vibrations travel easily through stiff lattices

  
  

  Insoluble; covalent bonds are stronger than attractions between solvent molecules and insolubility in polar solvents shows lack of ions

  
  

  Cannot conduct electricity; (in most) no charged ions or free electrons to carry a charge as all the bonding electrons are held in localised covalent bonds
• Graphite can conduct electricity


  This is an exception to the 'unable to conduct electricity' rule

  
  

  In graphite, each carbon atom is bonded to 3 others, arranged in sheets that are weakly held together, leaving one electron free to move between the sheets; allows graphite to conduct electricity

  
  

  Graphene is a single layer of graphite; it is 1 atom thick meaning it is a 2D compound. Graphene is a conductor, incredibly strong, transparent and light
• Giant metallic structures


  Metal elements exist as giant metallic lattices:

  1. Electrons in the outermost shell of the metal atoms are delocalised, leaving a positive metal ion

  2. These positive metal ions are all electrostatically attracted to the delocalised electrons; forming a lattice of closely packed positive ions in amongst a 'sea' of delocalised electrons; this is metallic bonding

  3. Overall lattice structure is of metal ions separated by layers of electrons
• Metallic bonding explains properties of metals

  1. High melting points; number of delocalised electrons per atom, ionic charge and size of metal ion affect melting point; more electrons, higher charge and smaller ions all increase the melting point

  2. No specific bonds; all the ions can slide over each other, making metals malleable and ductile

  3. Good thermal conductors as the delocalised electrons can pass kinetic energy to each other

  4. Good electronic conductors due to vast numbers of delocalised electrons

  5. Metals are in soluble, apart from in liquid metals, due to the strength of the metallic bonding
• Electronegativity and Polarisation
• Electronegativity


  The ability of an atom to attract bonding electrons in a covalent bond is called electronegativity; some atoms attract bonding electrons more than others

  
  

  Usually measured using the Pauling scale

  Fluorine is the most electronegative element with a value of 4.0 on the Pauling scale

  
  

  More electronegative elements have high nuclear charges and small atomic radii; electronegativity increases across periods and up groups (ignore noble gases)
• Covalent bonds can be polarised by differences in electronegativity


  If both atoms in the covalent bond have the same or similar electronegativities, electrons will sit around midway e.g. homonuclear diatomic gases like H₂. C-H bonds are essentially non-polar as they have similar electronegativities

  
  

  If the bond is between 2 atoms with different electronegativities, the more electronegative atom will pull the bonding electrons towards itself, causes a polar bond. Each atom will have a partial charge (δ+ or δ-)

  In a polar bond, the difference in electronegativity causes a dipole; a difference in charge between the 2 atoms in the bond
• Use the Pauling scale to work out Percentage ionic character

  Only bonds between atoms of the same element can be purely covalent so most bonds have some ionic character. Similarly very few compounds are purely ionic

  
  

  Data booklet contains a table showing how ionic a bond is against the difference in electronegativities
• Non-polar molecules can contain polar bonds

  Polarity of a molecule depends on its shape and the polarity of its bonds; a polar molecule has an overall dipole; a permanent charge across the molecule

  
  

  E.g.

  HCl is polar as it has 1 polar bond

  CO₂ is non-polar as it has 2 polar bonds pulling in opposite directions
• Intermolecular forces
• Intermolecular forces are very weak attractions between molecules

  London forces: instantaneous dipole-instantaneous dipole bonds

  
  

  Permanent dipole-permanent dipole bonds

  
  

  Hydrogen bonding (strongest intermolecular force)
• London forces (overview)


  Cause all atoms/molecules to be attracted to one-another

  1. Electrons in charge clouds are always moving; at any moment there could be more electrons at one side of an atom/molecule than the other; this is an instantaneous dipole

  2. This instantaneous dipole can induce a dipole in the opposite direction in a neighbouring atom; these dipoles are then attracted to each other

  3. Domino effect; second dipole can induce a third, etc

  4. Electrons are constantly moving so dipoles are being created and destroyed all the time
• London forces (lattices)


  e.g. Iodine;

  Iodine atoms are covalently bonded to form I₂; these I₂ molecules are then held together in a molecular lattice by London forces; this is a simple molecular structure
• Strength of London forces affects melting and boiling points

  Larger molecules/atoms have more electrons and so have stronger London forces

  Molecules with larger surface areas also have stronger London forces because they have a larger exposed electron cloud

  
  

  Liquids with stronger London forces have higher boiling points because it takes more energy to overcome the intermolecular forces; the same idea explains higher melting points

  
  

  e.g. alkanes; long-chain alkanes have fair higher melting/boiling points than short-chain alkanes

  
  

  but

  branched alkanes have lower boiling points than straight chain alkanes because they cannot pack as closely together so molecular surface contact is reduced