Topic 2: Structure and Bonding Flashcards
Definition of ionic bonding
The electrostatic force of attraction between oppositely charged ions
Factors influencing the strength of ionic bonding
Ionic radius - Smaller ions are more closely compact in a lattice so stronger electrostatic force
Ionic charge - Greater ionic charge = greater electrostatic force
Why are ionic compounds soluble in water?
Water is a polar molecule and so are ionic compounds. Partial charges of water break apart lattice and attract the ions
Definition of a covalent bond
The strong electrostatic force between shared pair of e- and the positive nuclei
Relationship between bond length and bond strength
Generally as bond strength increases, bond length decreases
How are covalent bonds formed?
Head on overlap = sigma bond
Sideways overlap = pi bond (above and below plane of molecule) less effective than head on overlap and therefore weaker
What do single, double and triple bonds comprise of?
Single (longer bond length) = sigma
Double (Shorter bond length) = sigma + pi bond
Triple = 1 sigma + 2 pi
Definition of dative covalent bond
The pair of electrons is supplied by one atom
Explanation for shape of molecules
Electron pairs repels and move as far apart to minimise the repulsion.
Bonded pairs repel more than lone pairs
Definition of electronegativity
The ability of an atom to attract the bonding electrons in a covalent bond
Trend of electronegativtiy
Across the period - electronegativity increases as no protons increase but no shielding e- stays the same so atomic radius decreases
Down the group -Electronegativity decreases as atomic radius increases as more shells are occupied
Difference between a polar bond and a polar molecule.
Polar bond is when there’s an electronegativity difference between the atoms
Polar molecules need a positive and negative end (dipole moment). Symmetry in a polar bond = non-polar molecule
Nature of different types of IMF
Induced Dipoles (London forces) - Present in all molecules and are instantaneous and temporary dipoles. Increase e- + increase in bonds = stronger LF
Permanent dipoles - Present in polar molecules caused by forces forming between dipoles
Hydrogen Bonds - Electrostatic force of attraction between Hydrogen and lone pair of e- on O, N or F and a hydrogen d+ on another molecule
Definition of Metallic Bonding
The electrostatic force between metal ions and sea of delocalised electrons
Ability of an cation to distort an anion
Polarising power - Small radius and high charge
Increasing polarising power = more covalent character showing as electron density region is getting closer
