Chapter 6 Shapes and intermolecular forces Flashcards
Electron-pair repulsion theory
- Electron pairs surrounding a central atom determine the shape of the molecule
- The electron pairs repel one another so that they are arranged as far apart from each other as possible
- The arrangement of electron pairs minimises repulsion -> holds the bonds in a definite shape
- Different numbers of electrons result in different shapes
Wedges
- Solid line
- Solid wedge
- Dotted wedge
Solid line
Bond in the plane of the paper
Solid wedge
Bond comes out of the plane of the paper
Dotted wedge
Bond goes into the plane of the paper
Bonding pair and lone pair repulsion comparison
- Lone pairs repel more than bonding pairs
- for each lone pair, bonding angle decreases by 2.5 degrees
Tetrahedral
- 4 bonding pairs
- 0 lone pairs
Pyramidal
- 3 bonding pairs
- 1 lone pair
Non linear
- 2 bonding pairs
- 2 lone pairs
4 electron pair shapes
- Tetrahedral
- Pyramidal
- Non-linear
3 electron pairs shapes
- Trigonal planar
- Bent
Trigonal planar
- 3 bonding pairs
- 0 lone pairs
Bent
- 2 bonding pairs
- 1 lone pair
2 electron pair shapes
- linear
linear
2 bonding pairs + 0 lone pairs
octahedral
- 6 bonding pairs
- 0 lone pairs
Shapes of ions
- Ammonium (NH4+)
- Sulfate (SO4 2-)
- Carbonate (CO3 2-)
- Nitrate (NO3 - )
Ammonium ion
- 4 bonding pairs + 0 lone pairs
- Tetrahedral
Sulfate ion
Tetrahedral
Carbonate ion
Trigonal planar
Nitrate ion
Trigonal planar
Electronegativity
The ability of an atom to attract a shared pair of electrons in a covalent bond
Pauling electronegativity scale
In periodic table left to right:
* Nuclear charge increases
* Atomic radius decreases
What happens to electronegativity as Pauling value increases?
Electronegativity of the atoms of that element increases
The most electronegative atoms
N
O
F
Cl
Least electronegative atoms
Group 1
What happens when electronegativity difference of atoms is large
The more electronegative bonding atoms will gain control of the electrons, and the bond becomes ionic (not covalent)
Diatomic molecules
- Atoms are of the same element
- Bonded electron pair is shared evenly
When bonded atoms are different:
- Atoms are different sizes
- Nuclear charges are different
- Shared pair of electrons maybe closer to one nucleus than the other
Non-polar bonds definition
Bonded electron pair is shared equally between the bonded atoms
Characteristics of non-polar bond
- Pure covalent bond
Example of non-polar bond
Hydrocarbon liquids (e.g. hexane)
When do non-polar bonds occur
- Bonded atoms are the same OR
- Bonded atoms have same or similar electronegativity
Shape of non-polar molecules
Symmetrical
Polar bonds
bonded electron pair is shared unequally between the bonded atoms
Characteristics of polar bonds
Polar covalent bond
Example of polar bond
e.g. hydrogen chloride
* Cl is more electronegative than H so Cl has ,ore attraction to bonded pair of electrons = polar covalent bond
δ signs
- δ+ : Atoms with smaller electronegativity
- δ- : Atoms with larger electronegativity
Dipole
Seperation of opposite charges
When do polar bonds occur
- Bonded atoms are different OR
- Bonded atoms have different electronegativity
Shape of polar bonds
Assymetrical
Solubility of ionic compounds
- Positive ions are attracted to oxygen (δ-) in water molecules
- Negative ions are attracted to hydrogen (δ+) in water molecules
Intermolecular forces
Weak attraction between dipoles of different molecules
What do intermolecular forces determine?
physical properties
What do covalent bonds determine?
identity + chemical reactions of molecules
London forces
Induced dipole-dipole interactions
Strength of intermolecular forces
Weakest -> strongest
- Induced dipole-dipole interactions
- Permanent dipole-dipole interactions
- Hydrogen bonding
- Single covalent bonds
Process of formation of induced dipole-dipole interactions
- Movement of electrons produces a changing dipole in a molecule
- At any instant, an instantaneous dipole will exist, but its position is constantly shifting
- The instantaneous dipole induces a dipole on neighbouring molecules
- The induced dipole induces further dipoles on neighbouring molecules, which then attract one another
State of induced dipole-dipole interactions
- Temporary
- Can disappear at any moment
What happens as the number of electrons in each molecule increases?
1.The larger the instantaneous-induced dipole
2. The larger the induced-induced dipole
3. The stronger the attractive force between molecules (higher b.p.)
What molecules do London forces occur in?
Both polar and non-polar molecules
What molecules do permanent dipole-dipole interactions occur in?
polar molecules ONLY
Simple molecular substance
- Made up of simple molecules
- Small units containing a definite number of atoms with a definite molecular formula
Structure of simple molecular substances as solids
Regular structure of simple molecular lattice
Simple molecular lattice
- Weak intermolecular forces hold the molecules together
- Atoms within each molecule are held together by strong covalent bonds
Properties of simple molecular substances
- Low m.p. and b.p. -> weak intermolecular forces are easily overcome by small amount of heat energy
- The strong covalent bonds do not break
Solubility of simple molecular substances
Different for
* Polar simple molecular substances
* Non-polar simple molecular substances
Solubility of non-polar simple molecular substances: Non-polar molecule + non-polar solvent
SOLUBLE
intermolecular forces form between the molecules and the solvent
These interactions weaken the intermolecular forces in the simple molecular lattice -> and break
Solubility of non-polar simple molecular substances: Non-polar molecule + polar solvent
INSOLUBLE
Little interaction between molecule and solvent
Intermolecular forces in solvent are too strong to be broken
Solubility of polar simple molecular substances: Polar molecule + polar solvent
SOLUBLE
Attract one another (similar to dissolving of ionic compounds)
Electrical conductivity of simple molecular substances
Do not conduct electricity due to:
* No mobile charged particles
Hydrogen bond
Special type of permanent dipole-dipole interaction between molecules
What does a hydrogen bond contain?
- An electronegative atom with a lone pair of electrons (O, F, N)
- A H atom attached to an electronegative atom (H-O, H-N, H-F)
What is a hydrogen bond shown by?
Dashed line
What gives water unique properties?
Hydrogen bonding
Unusual properties of water
- Ice (solid) is less dense than water (liquid)
- Water has relatively high m.p. and b.p.
Why is ice less dense than water
- Hydrogen bonds hold water molecules apart in an open lattice structure
- The water molecules in ice are further apart than in water
- Solid ice is less dense than liquid water and floats
Structure of ice
- Water has 2 lone pairs so can bond to 4 hydrogen atoms
- These hydrogen bonds extend outwards, forming an open tetrahedral lattice, full of holes
- Bond angle is close to 180 degrees
What happens when ice melts
The ice lattice collapses and the molecules move close together
Why does water have a relatively high melting point and boiling point?
- Hydrogen bonds are extra forces
- Energy is needed to break the hydrogen bonds
- When the ice lattice breaks, the rigid arrangement of bonds in ice is broken. When water boils, the hydrogen bonds break completely
Ionic lattice
Repeated pattern of oppositely charged ions
Covalent bond
shared pair of electrons
