CHAPTER 6 - SHAPES OF MOLECULES AND INTERMOLECULAR FORCES Flashcards
What is the electron-pair repulsion theory
Electrons have a negative charge, so they repel each other
Determines she of molecule or ion
Repulsion causes the electrons to move as far away as possible
Electron pairs minimise repulsion so the bonds hold a definite shape
Different numbers of electron pairs result in different shapes
What is an example of a molecule that forms a tetrahedral shape
Methane (CH4)
What is the bonding angle in a tetrahedral shape?
109.5 degrees
What lines are drawn to show what dimension
Solid line - Straight
Solid wedge - Out of paper (towards viewer)
Dashed Wedge - into paper (away from viewer)
What is the increasing repulsion of lone pairs and bonded pairs
Bonded pair to bonded pair least repulsive
Bonded pair/lone pair - Moderate Repulsion
Lone pair/lone pair - High repulsion
For each Lone pair of electrons, ho much does the bond angle decrease by?
2.5 degrees
How many bonded pairs and Lone pairs does a tetrahedral molecule have?
4 bonded pairs (bonding regions)
0 Lone pairs
What bonding angle does a Pyramidal have and what is an example of a pyramidal molecule
107 degrees
NH3
How many bonded pairs and Lone pairs does a Pyramidal molecule have?
3 Bonded Pairs (bonding regions)
1 Lone Pair
What bonding angle does a non-linear molecule have and what is an example o a non-linear molecule?
104.5 Degrees
Water (H2O)
How many bonded pairs and Lone pairs does a non-linear molecule have?
2 Bonded Pairs (bonding regions)
2 Lone Pairs
What bonding angle does a linear molecule have and what is an example o a linear molecule?
180 degrees
CO2
How many bonding regions does a linear molecule have?
2 Bonded regions
What is the bonding angle, number of bonding regions and an example of a Trigonal Planar
120 degrees
3 bonding regions
BF3
What is the bonding angle, number of bonding regions and an example of an octahedral molecule
90 degrees
6 bonding regions
SF6
(forms shape with 8 sides)
What is the most electronegative element?
Fluorine - 4.0
electronegativity decreases further away
What happens to nuclear charge and atomic radius across the periodic table
Nuclear charge increases
Atomic radius decreases
What is the difference in electronegativity for a covalent bond
none
What is the difference in electronegativity for a polar covalent bond
0-1.8
What is the difference in electronegativity for an ionic bond
Greater than 1.8
What is a non-polar bond
Bonded electron pair is shared equally between the bonded atoms
Bonded atoms are the same or the atoms have same/similar electronegativity
What is a polar bond?
A difference in electronegativity, which leads to partial charges and a permanent/induced dipole - variations of electron inhabitance time
eg H - Cl 2.1 - 3.0
Define Electronegativity
The measure of the attraction of a bonded atom for the pair of electrons in a covalent bond
Define Polar covalent bond
Shared pair of electrons where the pair is not shared equally between the two bonded atoms
Define Dipole
Charge separation across a bond with one atom having a slightly positive charge and another with a slightly negative charge
What are the 3 categories of Intermolecular forces?
Induced dipole-dipole interactions (London forces)
Permanent dipole-dipole interactions
Hydrogen Bonding
What are London forces (induced dipole-dipole interactions)
Weak IMF that exist between all molecules, whether polar or non-polar
Movement of electrons produces changing dipole, instantaneous
Induces dipole in neighbouring molecule
Induced dipole induces further dipoles in neighbouring molecules which then attract one another
temporary
What determines the strength of induced dipole-dipole interactions
Number of electrons
Strength of attractiveness
number of interactions
What is the difference between LDFs and Van Der Waals forces?
LDF - only induced/temporary dipole-dipole interactions
VDW - Both induced and permanent dipole-dipole interactions
What are simple molecular substances
Molecules made up of small units containing a definite number of atoms with a definite molecular formula, such as Ne, H2, H2O or CO2
What structures do simple molecules form?
Simple molecular lattice
Held by weak IMF, atoms in each molecule held strongly together by covalent bonds
What are properties of simple molecular substances
Low MP + BP
soluble in non-polar non-simple substances
why are non-polar simple molecular substances soluble in non-polar non-simple solvent (eg hexane)
IMF forces form between molecules and solvent
IMF are weakened by interactions and therefore break and compound dissolves
why are non-polar simple molecular substances insoluble in polar non-simple solvent
Little interaction occurs
Hence bonds are too strong to be broken
Are simple molecular substances conductors of electricity
No mobile charged particles
Nothing to move electrical current
Explain how an induced dipole forms
Fluctuation in the electron density around a molecule creates an instantaneous dipole in a molecule. The instantaneous dipole induces a dipole in a neighbouring molecule
Explain why simple molecular compounds:
a) have low MP + BP
b) Doesn’t dissolve in water
c) Poor Electrical conductivity
a) Weak IMF are broken by the energy present at low temperatures
b) Little interactions between the molecules and lattice
c) no mobile charged particles
What are Hydrogen bonds?
A permanent dipole-dipole interaction from a Hydrogen to either an Oxygen, Nitrogen or Fluorine atom
Why is solid (ice) less dense than water (liquid)
Hydrogen bonds hold water molecules further apart in ice than in liquid
in an open lattice structure
Therefore it is less dense - so it floats
Why does Water have a relatively high MP+BP
Contain Hydrogen bonds over and above the London Forces
Requires more energy to break bonds
Water, Hydrogen Fluoride and ammonia do not follow the trend shown by the other hydrides in each group.
a) Estimate what the boiling points of water, hydrogen fluoride, and ammonia would be if they were to follow the group trends.
b) explain why water, hydrogen, fluoride, and ammonia, do not follow the group trends
(pg 82)
a)
water: -75 degrees C
HF: - 90 degrees C
NH3: -100 degrees C
b)
Three. Hydrogen bonding, which is a stronger intermolecular force than other dipole interactions. Greater energy needed to overcome the intermolecular forces, so boiling points are much higher
Explain all groups show an increase in boiling point from period 3 to period 6
Increase electrons increases the strength of the London forces
What can be drawn about the relative strengths of London forces and permanent dipole-dipole interactions for the hydrides of group 14 to 17
From period 3 to period six, difference in electronegativity between hydrogen and element decreases, decreasing the permanent dipole-dipole interactions. Number of electrons increases, increasing London forces. Boiling point increases, so London forces are stronger and more significant than a permanent dipole-dipole interactions
Suggest why pairing doesn’t take place between two purine bases or between two Pyramidine bases
Two bases would be too close together.
Two Pyramidine bases will be too far apart.
State and explain two anomalous properties of water
ICE is dense than liquid water because hydrogen bonds hold molecules apart in open lattice structure.
Higher milk and boiling points than expected because appreciable energy is needed to break the hydrogen bond
