CK - quizlet

• Ionic bonding


  The strong electrostatic forces of attraction between oppositely charged ions in a lattice
• Strength of ionic bonding increases with

  smaller ionic radius and greater charge (greater charge density)
• Trends in ionic radii down a group

  The ionic radius increases down a group as there are more electron shells
• Evidence of ions


  Migration of ions e.g electrolysis of green copper chromate.

  blue copper 2+ ions move towards the negative electrode

  Yellow chromate ions move towards the positive electrode
• Physical properties of ionic compounds

  Hard, brittle, crystalline - repulsion between ions if structure is distorted

  High melting and boiling points - strong attraction between oppositely charged ions

  Soluble in polar substances - contain positive and negative ions

  Conduct electricity in molten or aqueous - ions free to move
• Covalent compound properties - simple molecular

  Usually gases, liquids or soft solids at room temperature - weak intermolecular forces

  Low melting and boiling points

  Cannot conduct electricity in any state - no free electrons to carry charge

  Usually more soluble in non-polar solvents
• Covalent bond


  Strong electrostatic attraction between two nuclei and the shared pair of electrons between them.
• Bond length and strength in covalent bonds

  Bond strength is inversely proportional to bond length.

  large atoms make longer, weaker bonds (more shielding)

  double bonds shorter and stronger than single
• Electron pair repulsion theory

  Electron pairs in the outer shell of atoms and ions repel each other and get as far apart as possible (results in molecular shapes)

  Lone pairs repel more than bonding pairs and decrease the bond angle by approximately 2.5 degrees
• Electronegativity


  The ability of an atom to attract the bonding pair of electrons in a covalent bond
• Bonding spectrum


  Pure covalent (electrons shared halfway between two atoms)

  Polar covalent (attracted more strongly to electronegative atom - has a permanent dipole)

  Polar ionic - ions form but the electrons around the anion are attracted towards the cation resulting in a distorted anion shape (not spherical)

  Pure ionic - electron is completely removed and transferred to the more electronegative atom.
• Polar molecules


  Bonds between atoms of elements with different electronegativity values are polar

  Polar molecules are asymmetric - contain a net dipole moment.

  Not all molecules with polar bonds are polar.
• London forces


  Intermolecular forces that exist between all molecules.

  They arise from attractions between temporary instantaneous dipoles (caused by random movement of electrons) and the dipoles they induce in neighbouring molecules.
• Dative covalent bonds


  A bond in which two atoms share a pair of electrons, both of which are donated by one atom.

  Behaves the same way as a covalent bond once formed.

  E.g. Al2Cl6 (lone pair on chlorine is donated to empty orbital on aluminium) and NH4+ (ammonia molecule shares its lone pair with a hydrogen ion H+)
• Hydrogen bonds


  A strong intermolecular force between a delta +ve hydrogen covalently bonded to fluorine, oxygen or nitrogen and a lone pair of electrons on the delta -ve O,F,N of a nearby molecule.

  Accounts for;

  - high Bp of ammonia, water and HF compared to other hydrides in the same group

  - open structure and low density of water

  - solubility of alcohol in water

  **H-O..H angle is 180 degrees

  **must show lone pairs and partial charges across bonds
• Trend in BP group 7 hydrides

  HCl, HBr, and HI all have permanent dipoles and London Forces. Although the dipole in HCl is stronger than HBr and HI the increase in electrons makes LF more significant so BP increases from HCl to HI.

  HF has hydrogen bonding so has much higher BP than the other three.
• Dipole-dipole interaction


  Attractive forces between the positive pole of one molecule and the negative of another.

  Occurs in molecules with permanent dipoles

  E.g. HCl
• Property of metals


  - high mp - strong attraction

  - high density - close packing with little space

  - heat and electrical conductors - delocalised electrons move freely even in solid

  - malleable - not directional bonds. Force causes a slip, ions can resettle into close packing. repulsion of cations overcome by attraction of electron 'sea
• Giant lattice


  Regular arrangement of particles, the formula is the simplest repeating unit, no limit to number of particles that can join together (unlike molecules)

  All ionic compounds, all metallic elements are lattices. Covalent lattices are diamond, graphite and silicon dioxide.
• Lone pairs


  A pair of of electrons in the outer shell of one of the atoms in a molecule or ion which is not involved in bonding

  Reduces bond angle by 2.5 degrees

  Counts as an area of electron density in working out shape..

  E.g H20 = 109.5 - (2x2.5) = 104.5
• tetrahedral


  molecule with four bonding regions of electrons no lone pairs, bond angle 109.5, e.g. CH4
• octahedral


  six bonding regions of electrons and no lone pairs

  bond angles 90, e.g. SF6
• trigonal planar


  three bonding regions of electrons and no lone pairs

  bond angles 120, e.g. BF3, AlCl3, CH2O
• linear


  two bonding regions of electrons, no lone pairs

  angle 180, e.g. BeCl2, CO2
• non-linear


  two bonding regions of electrons, one or more lone pairs. Bond angle 109.5 - 2.5 for every lone pair e.g. SO2 107 degrees, H2O 104.5 degrees
• pyramidal


  three bonding regions and one lone pair

  bond angle 107 e.g. NH3, PH3
• trigonal bipyramidal


  five bonding regions and no lone pairs

  bond angle 120 and 90

  e.g. PCl5
• dipole moment


  when the presence of polar bonds results in an overall dipole across the molecule. The molecule must have an asymmetrical distribution or charge overall.
• Multiple bonds


  Double or triple bonds.

  Area of high electron density.

  Shorter than single bonds.

  Counts as one area of electron density when calculating bond angles.
• Metallic bonding


  The strong electrostatic attraction between metal cations and a sea of delocalised electrons
• Bond angle


  Angle between two covalent bonds on a molecule or giant covalent structure
• Shapes of molecules


  Electron d shape angle example

  2. Linear. 180. BCl2

  3. Trigonal planar 120 BeCl3

  4. Tetrahedral 109.5. CH4

  5. Trigonal bipyramid 180, 90, 120 PCl5

  6. Octahedral 180, 90 SF6
• Solubility in water


  Ionic substances and those that can form hydrogen bonds will tend to be soluble in water
• Solubility in non-polar solvents

  Non-polar substances will dissolve well in non-polar solvents, for example iodine dissolves well in hexane