Chapter 6 (2.2.2) Flashcards
What is the electron pair repulsion theory?
- explains the shapes of molecules + polyatomic ions
- electron pairs surrounding a central atom determine the shape of the molecule or ions by repelling one another so they are arranged as far apart as possible
- arrangement minimises repulsion + thus holds the atoms in a definite shape
Wedges for 3D shapes
solid line: bond in plane of paper
solid wedge: comes out of plane of paper
dotted wedge: into the plane of paper
Bond and lone pair repulsions
- lone pairs are slightly closer to the central atom + occupy more space than a bonded pair
- thus a lone pair repels more strongly than a bond pair
- bonded/bonded < bonded/lone < lone/lone (relative strength of repulsions)
What does adding lone pairs do?
since lone pairs repel more strongly than bonded pairs, they repel the bonded pairs slightly closer together, decreasing the bond angle by 2.5° for each lone pair.
What are the molecular shapes + bond angles for methane, ammonia, water?
Methane: tetrahedral, 109.5°
Ammonia: Pyramidal, 107°
Water: Non-linear (bent) 104.5°
What happens to the molecular shape with multiple bonds?
- each multiple bond is treated as a bonding region
- e.g. 2 double bonds = 2 bonding regions
carbon dioxide, CO2
linear, 180°
boron trifluoride, BF3
trigonal planar, 120°
phosphorus pentachloride, PCl5
trigonal bipyramidal, some 120° and some 90°
sulfur hexafluoride, SF6
octahedral, 90°
ammonium, NH4+
tetrahedral, 109.5°
carbonate, CO32-
trigonal planar, 120°
nitrate, NO3-
trigonal planar, 120°
sulfate, SO42-
tetrahedral, 109.5°
What is the basis for electronegativity?
- in bonds like H2 or O2, the atoms are the same element and the bonded electron pair is shared evenly
- this changes when the atoms are different elements: since nuclear charges are different, atoms may be different sizes, the shared pair may be closer to one nucleus than the other
- shared pair may experience more attraction from one atom
What is electronegativity?
The ability of an atom to attract the bonding electrons in a covalent bond
Measure of electronegativity
The Pauling scale
Describe the Pauling scale values
Across the periodic table: - nuclear charge increases - atomic radius increases Thus are more electronegative Noble gases not included
Patterns on electronegativity
- non metals: N, O, F, Cl = most electronegative
- G1 metals: Li, Na, K = least electronegative
bond types based on electronegativity difference
covalent = 0
polar covalent = 0-1.8
ionic: >1.8
When does a bond become ionic?
If the EN difference is large, one bonded atom will have a much greater attraction for the shared pair than the other bonded atom. the more electronegative atom will have gained control of the electrons + the bond becomes ionic.
Describe non-polar bonds
- electron pair shared equally between the bonded atoms
- occurs when atoms have the same or similar electronegativity
- when bonded atoms come from the same element it is a pure covalent bond
- non-polar solvents do not mix with water
Describe polar bonds
- electron pair shared unequally
- occurs when bonded atoms are different + have different EN values - polar covalent bond
- a dipole in a polar covalent bond does not change + is a permanent dipole
What is a dipole
The separation of opposite charges
What makes a molecule polar?
- if there are 2 or more polar bonds, their direction is important
- depending on the shape of the molecule, the dipoles may reinforce one another to produce a larger dipole - making the molecule polar
- or cancel out if the dipoles act in opposite directions - making the molecule non-polar
What are London forces?
- induced dipole - dipole interactions
- weak intermolecular forces that exist between all molecules, whether they are polar or non-polar
- act between induced dipoles in different molecules
The origin of induced dipoles
- quick movement of electrons means that at any instant, the electrons in an atom are likely to be more to one side than the other. thus an instantaneous dipole will exist, but will be constantly shifting
- this dipole induces a dipole on a neighbouring atom, the 2 are then attracted
- the induced dipole induced further dipoles on neighbouring molecules, which then attract one another
Overall effect of London forces
- induced dipoles are temporary
- electrons are constantly moving so dipoles are being created + destroyed all the time
- even though the dipoles keep changing, the overall effect is for the atoms to be attracted to each other
Why does boiling point increase down a group?
down a group there are more electrons meaning:
- larger instantaneous + induce dipoles
- greater induced dipole-dipole interactions
- stronger attractive forces between molecules
Thus more energy needed so higher boiling point
When do permanent dipole-dipole interactions occur?
- they act between permanent dipoles in different polar molecules
- in addition to London forces
Boiling points in polar molecules
polar molecules have additional permanent dipole-dipole interactions between molecules so extra energy needed to break them, thus higher boiling points
What is a hydrogen bond?
Special type of permanent dipole-dipole interactions found between molecules containing:
- an electronegative atom with a lone pair of electrons, N, O, F
- a hydrogen atom attached to an electronegative atom
Strongest type of intermolecular attractions
What does a hydrogen bond act between?
A lone pair of electrons on an electronegative atom in one molecule + a hydrogen atom in a different molcule
