Bonding

• Define ionic bond


  - the strong electrostatic forces of attraction between oppositely charged ions in an giant lattice
• Define covalent bond


  - a shared pair of electrons
• Define dative bonding


  - a lone pair of electrons from one atom is donated to another to form a dative covalent bond
• Explain the shape of CH4 (2)

  - there are 4 bonding pairs of electrons around the central atom and no lone pairs

  - these repel equally to get as far apart as possible, resulting in a tetrahedral shape of bond angle 109.5
• Explain the shape of NH3 (3)

  - there are 3 bonding pairs of electrons around the central and one lone pair of electrons

  - the lone pair of electrons repel more than the bonding pairs

  - resulting in a pyramidal shape and a bond angle 107°
• Define electronegativity


  - the power of an atom to attract the pair of electrons in a covalent bond towards itself
• What does electronegativity depend on??

  - atomic radius

  - nuclear charge

  - shelding
• How does atomic radius affect electronegativity?

  - as the radius of an atom increases, the bonding pair of electrons become further from the nucleus.

  - so they are less attracted to the positive nucleus resulting in a lower electronegativity
• How does the number of unshielding protons affect electronegativity?

  - the greater the number of protons in a nucleus, the greater the attraction to the electrons in the covalent bond, resulting in a higher electronegativity

  - full energy shells shield the electrons in the bond from the increased attraction of the greater nuclear charge, thus lowering the electronegativity
• Describe the trend in electronegativity across a period

  - increases

  - atomic radius decreases

  - charge on the nucleus increases without any extra shielding

  - new electrons don't contribute to shielding because they're added to the same energy level
• Describe the trend in electronegativity down a group

  - decreases

  - atomic radius increases

  - although charge on the nucleus increases, shielding also increases a lot.

  - this is because electrons added down a group fill new energy levels
• Describe what happens if the electronegativity of both atoms in a covalent bond is equal

  - results in a symmetrical distribution of electron density around the 2 atoms

  - bonding is always non-polar because the electronegativity is the same
• Describe what happens if the electronegativity of both atoms is different

  - one atom has a greater power to attract the bonding pair of electrons in a covalent bond

  - bonding pair of electrons is closer to the more electronegative atom's nucleus and the electron density is asymmetrical

  - this type of covalent bond is polar

  - dipole formed - one side is slightly positive charge and the other a slightly negative charge
• What is a permanent molecular dipole?

  - if a molecule contains polar bonds then there is a possibility that is has a permanent molecular dipole

  - one end is partially negative, other side is partially positive

  - molecules containing an asymmetrical arrangement of polar bonds do have a permanent dipole, but if arranged symmetrically then the dipoles cancel each other out
• Describe carbon dioxide through dipoles

  - 2 dipoles within the molecule

  - no permanent molecular dipole

  - dipoles cancel out as they're arranged symmetrically around central atom
• Describe water through dipoles

  - 2 dipoles within the molecule

  - has a permanent molecular dipole

  - dipoles don't cancel out completely as they're arranged asymmetrically around central atom
• What are the 3 main types of intermolecular forces?

  - van der waals

  - permanent dipole-dipole forces

  - hydrogen bonds
• Describe Van der Waals forces

  - the uneven distribution of electrons in one molecule INDUCES a dipole in a neighboring molecule

  - sets up a temporary dipole that is very weak

  - acts between all molecules
• How does molecular size affect the strength of the van der waals forces?

  - boiling point increases going down group 8

  - there is an increase in the number of electrons from He to Radon

  - so radon has larger van der waals forces between atoms which require more energy to overcome
• Explain why boiling points of HCl, HBr and HI increase

  - boiling point increases going down the group

  - HI has a higher boiling point because there are more electrons

  - strength of permanent dipole-dipole forces decrease going down the group due to size of electronegativities decreasing

  - this is outweighed by increasing van der waals forces between molecules due to increasing number of electrons
• Describe permanent dipole-dipole forces

  - occur in addition to van der waals forces in molecules that are POLAR - have a permanent molecular dipole

  - strength of permanent dipole-dipole forces decreases down group 7 due to electronegativities decreasing
• Describe the requirements for hydrogen bonds to form

  - hydrogen atom must be attached to nitrogen, oxygen or fluorine - these are most electronegative

  - the nitrogen, oxygen or fluorine must have a lone pair of electrons

  This makes the attraction between lone pair and hydrogen stronger that a normal dipole attraction
• Rules for drawing hydrogen bonds

  - show all lone pairs

  - show all partial charges

  - attraction must be show from delta+H on one molecule 180 degrees to lone pair on other molecule
• explain the patterns in boiling points between hydrogen halide molecules

  - they have van der waals forces between molecules

  - increase in boiling point from HCL to HBr due to size of molecule increasing because of more electrons

  - HF has hydrogen bonding so requires more energy to overcome
• describe the structure of ice

  - when water freezes, there is an expansion and takes up more space because the molecules get pushed further apart

  - ice has a lower density than water so ice floats on water
• What are the 4 main types of crystal structures

  - ionic

  - metallic

  - simple molecular

  - macromolecular
• describe metallic bonding

  - the attraction between a lattice of positive ions and a cloud of delocalised electrons
• what do the strength of the metallic bonding depend on

  charge of ion - the bigger the charge, the more delocalised electrons so the stronger the electrostatic forces of attraction

  size of ion - smaller the ion, the closer the electrons are to the positive nucleus so the bond is stronger, and more energy needed overcome the forces.
• describe a giant metallic lattice

  - malleable because when a force is applied, metal layers can slide and metallic bonds reform - the lattice doesn't break only changes shape

  - high melting points due to strong electrostatic attractions between positive ions and delocalised electrons

  - conducts electricity due to delocalised electrons can carry charge through structure
• describe simple covalent structures

  - eg. iodine

  - low melting points due to have weak van der waals forces between the molecules

  - only small amounts of energy needed to break the lattice

  - boiling points depend on size of molecules
• describe the giant ionic lattice

  - ionic compounds have this structure

  - has high melting points due to strong forces of electrostatic attraction between oppositely charged ions

  - melting points increase with charge density of ions due to stronger electrostatic attraction of charges

  - ionic compounds are soluble in water

  - can conduct electricity when molten because ions are free to move

  - brittle because when force is applied, layers of ions can shift, causing like-charged ions to align, repelling each other
• describe giant covalent structure

  - high melting and boiling points due to its giant covalent structure

  - lots of strong covalent bonds must be broken before it melts

  - cant conduct electricity because there are no free electrons to carry charge - they are all involved in the covalent bonds except graphite

  - graphite and diamonds are allotropes of carbon
• describe diamond


  - has a giant covalent lattice

  - each carbon atom is covalently bonded to 4 others in a tetrahedral arrangement with a bond angle of 109.5

  - strong bonds in all direction
• describe graphite

  - each carbon atom is bonded to 3 others in a layered structure

  - layers are made of hexagons with a bond angle of 120

  - so has delocalised electrons which can carry charge through the structure

  - layers are held by van der waals forces - allowing layers to slide over each other, making graphite malleable
• describe graphene

  - made of a single layer of carbon atoms that are bonded together

  - very thin
• predict the shape of and bond angles in a molecule of sulfur hexafluoride SF6 (7)

  - SF6 has an octahedral shape (1)

  - with bond angles of 90 degrees (2)

  - there are 6 bonding pairs of electrons around the central atom (3)

  - and no lone pairs of electrons (4)

  - the electrons repel equally (5)

  - so they are as far apart as possible (6)

  - draw the molecule (7)
• explain why a CH4 molecule has no overall polarity (3)

  - each of the 4 C-H bonds have a slight dipole present because the electronegativity of carbon is slightly higher than that of hydrogen (1)

  - the dipoles cancel each other out completely as they are arranged symmetrically around the central atom (2)

  - so the molecule is non-polar (3)
• explain why the boiling point of PH3 is higher than CH4

  - PH3 has more electrons than CH4 (1)

  - so greater van der waals forces between its molecules (2)

  - which require more heat energy to overcome (3)
• suggest why the strength of the C-H bond in CH4 is stronger than that of the SI-H bond in SiH4. (2)

  - the C-H bond is stronger because the electron configuration of carbon is 2,4 and of silicon is 2,8,4.

  - because more electron shells are filled by the electrons in silicon the outer shell is further away from the nucleus.

  - so the electrostatic attraction between the electrons in the shared pair and the nucleus of silicon is weaker

  - covalent bonds aren't broken so this difference in bond strength doesn't change the boiling point
• van der waals forces exist between all molecules. explain how these forces arise (3)

  - electron movement in the first molecule (1)

  - induces a dipole in another molecule (2)

  - induced temporary attraction/ slightly positive attracts slightly negative in an adjacent molecule (3)