Topic 2: Bonding & Sturcture Flashcards
Ionic bond
Strong electrostatic attraction between two oppositely charged ions
cation +, anion -
Ionic Bond Strength Factors
Greater charge = stronger bond (more electrons transferred)
Smaller ionic radii = stronger bond as electrostatic attraction gets weaker with distance
Higher charge density = stronger bond
Ionic Radii Trends
Increases as you go down a group
Ionic radius of isoelecreonic electrons decreases across a period -> higher electrostatic attraction so more effective nuclear charge
Isoelectronic Ions
Same electronic configuration
E,g N3-, O2-
Giant Ionic Lattice
same basic unit repeating
ions are electrostatically attracted to the opposite charge in all directions
Properties of Ionic Bonds
High melting point
Soluble in water but not in non-polar solvents: proves it contains ions
Conducts electricity when aqueous or molten but not when solid as ions are free to move
Brittle: if it is same ions are on top of one another strong repulsion causes bonds to break
Covalent Bonds
Strong electrostatic attraction between a pair of shared electrons and two positively charged nuclei
Covalent Bond Length
Distance at which forces are balanced
Nuclei are attracted to area of e- density but repel each other
Covalent Bond Enthalpy
Energy needed to break one molecule of a bond, directly proportional to bond strength
More shared pairs -> higher e- density -> stronger attraction -> higher bond enthalpy
Stronger enthalpy = shorter bond
Dative Covalent Bonds
When one atom donates both electrons to a bond
Shapes of Molecules
Depend on number of electron pairs in outer shell of central atom
Lone / bonded pairs
Electron repulsion: angles
Bonded/bonded has highest bond angle, then lone/bonded then lone/lone
Bond angle is smaller because lone/lone is larger
Atoms are pushed closer to minimise repulsion
Linear
2 electron pairs, none lone
180°
Trigonal Planar
3 electron pairs, none lone
120°
Non-linear, “bent”
3 electron pairs, one lone (e.g SO2)
119°
4 electron pairs, 2 lone (e.g H2O)
104.5°
Tetrahedral
4 electron pairs, none lone
109.5°
Trigonal pyramidal
3 electron pairs, 1 lone
107°
Giant Covalent Structures
Huge lattices of covalently bonded atoms with much stronger electrostatic forces of attraction
Properties of Giant Covalent Structures
Very high melting points
Extremely hard
Good thermal conductors as vibrations can travel easily through stiff lattices
Insoluble as they don’t contain ions
Cannot conduct electricity as they have no delocalised e- or ions
Graphite
CAN CONDUCT ELECTRICITY
shares 3/4 e- so each atom has one delocalised that can move between layers (lubricant)
GRAPHENE is a sheet of carbon, can conduct as delocalised moves along sheet
Strong, light, transparent
Giant Metallic Structures
Layers of positive metal ions separated by layers of electrons
Metallic Bonding: +ve metal ions are electrostatically attracted to delocalised electrons
Properties of Giant Metallic Structures
High melting point
Good electrical & thermal conductors
Malleable, ductile
Insoluble
Electronegativity
Ability of an atom to attract the bonded pair of electrons in a covalent bond
Increases across periods (nuclear charge increases and atomic radii decreases)
Decreases down groups
Polarisation of bonds
Occurs due to differences in electronegativity, causes a dipole
Non-polar if e-ns are similar or identical e.g homonuclear molecules
Polar: uneven spread of e- in bond -> bonded pair will be closer to most electronegative
Higher difference: more ionic
Polar Molecules
In simple molecules a polar bond makes whole thing permanent dipole (polar molecule)
If there are several bonds:
1) non-polar if they are on opposite sides -> cancel out
2) polar if they are roughly on same side
London Forces
Type of intermolecular forces (weaker than covalent)
Induced dipole-dipole interaction: e- moves to one side so nucleus is attracted to e- cloud
Overall effect makes atoms attracted to each other in a lattice -> simple molecular structure
More molecular surface contact means stronger forces and higher mp
Permanent dipole-dipole bonds
Occur in addition to London Forces
Form polar molecules
Higher energy to overcome than just london
Hydrogen Bonding
Strongest type of intermolecular force
If H is covalently bonded to N, O, F (which draw e- away from H)
Very polarised bond & H has high charge density to H atoms form weak bonds to lone pairs
Group 7 hydrides
HF has high boiling point
From HCl to HI permanent dipole-dipole decrease but is overridden as number of e- increases so there are stronger London forces which cause bp to increase
Group 6 hydrides follow similar trends
Why does ice float?
lattice structure in ice aims for maximum n of hydrogen bond, “wasting space” -> ice is less dense than water
As ice melts, some bonds are broken so lattice breaks down
