2.2.2 bonding and structure Flashcards
Ionic bonding
Electrostatic attraction between oppositely charged ions formed by electron transfer
Giant ionic lattice
Ionic bonding between ions
Properties of ionic compounds
High melting points - strong electrostatic forces between oppositely charged ions in lattice
- higher when ions are smaller and/or have higher charges
Non conductor of electricity when solid - ions held together tightly in lattice and can’t move so no charge conducted
Good conductor of electricity when in solution or molten - ions are free to move when in solution and molten so can carry a charge
Usually soluble in aqueous solvents
Covalent bonding
Strong electrostatic attraction between a shared pair of electrons and nuclei of bonded atoms
Dative covalent bonding
When the shared pair of electrons in covalent bond come from only one of the bonding atoms
Aka co-ordinate bonding
E.g. NH4+, H3O+, NH3BF3
Average bond enthalpy
Measurement of covalent bond strength
The larger the value, the stronger the covalent bond
Covalent bond structure
Simple molecular
Intermolecular forces (induced/ permanent dipole-dipole/ hydrogen bonds) between molecules
E.g. iodine, CO2, H2O
Molecular (simple) boiling and melting points
Low as weak intermolecular forces between molecules (specify type)
Molecular conductivity
Poor as no ions to conduct and electrons are in fixed positions so cannot move to carry a charge
Linear
2 BP, 0 LP
180
Trigonal planar
3BP, 0LP, 120
Tetrahedral
4BP, 0LP
109.5
Trigonal pyramidal
3BP, 1LP
107
Bent
2BP, 2LP
104.5
Octahedral
6BP, 0LP
90
How to explain shape
State no. BP + LP
State electron pairs repel and try to get as far apart as possible
If no LPs, state electrons repel equally
If there are LPs, then state LPs repel more than BPs
State actual shape and bond angle
Lone pairs
Repel more than bonding pairs
Reduce bond angles by 2.5 per lone pair
Electronegativity
Relative tendency of an atom in a covalent bond in a molecule to attract electrons towards itself in a covalent bond
Most electronegative atoms
F, O, N, Cl
electronegativity across a period
Increase across a period as proton number increases
atom radius decreases as electrons are in the same shell
Electronegativity down a group
Decreases down a group as distance between nucleus and outer electrons increases
Shielding of inner shell electrons increases
Formation of permanent dipole
When the elements in the bond have different electronegativities
When there is an unequal distribution of electrons in the bond, a dipole forms.
Element with the larger electronegativity will have the negative dipole
Symmetric molecules
Will not be polar even if all individual bonds are polar
No net dipole
Induced dipole-dipole where
All molecular substances and noble gases, not in ionic
How London forces occur
In any molecule the electrons are moving constantly and randomly
As this happens electron density fluctuates and parts of the molecule become more or less negative
Temporary dipoles from this cause dipoles to form in neighbouring molecules (induced dipoles)
Factors affecting size of induced-dipole
The more electrons, the higher the chance the temporary dipole will form
This makes the induced dipole-dipole interactions stronger between molecules and so boiling point increases
Group 7 boiling point
Increase down the group
Increasing number of electrons, increases the size of induced dipoles interactions
More energy needed to overcome
Alkane homologous series boiling points
Increases down the group due to increasing electrons
Increase in size of induced dipole-dipole interactions between molecules
More energy needed to overcome
Shape of molecule boiling point
Longer straight chain alkanes have a larger surface area of contact between molecules for induced dipole-dipole interactions
Higher boiling point
Hydrogen bonding
Between hydrogen and N,O,F
Ice
Lower density - Molecules are held further apart than in liquid water
Hydrogen bonding - high boiling point
Iodine
Covalent bonds between I2 molecules
Crystals contain regular arrangement of I2 molecules held together by weak induced dipoles
