M2 - Chapter 6 Flashcards
why do scientists use the electron repulsion theory
for explaining and predicting shapes of molecules and polyatomic ions.
what is the electron repulsion theory
- An electron has negative charge
- Electron pairs repel one another so they occupy a position that will minimise this repulsion – as far apart as possible
different number of electrons…
different shapes
what does a solid line represent in a 3D diagram
represents a bond in the plane of the paper
what does a solid wedge represent in a 3D diagram
comes out of plane of paper
what does a dotted wedge represent in a 3D diagram
goes into plane of paper
pattern in bond angles
Bond angle is reduced by about 2.5 degrees for each lone pair
difference between bonding pair and bonding region
can be used interchangeably
bonding pair is a single covalent
bonding region is a double covalent
greater number the electron pairs… (relate to bond angle)
smaller the bond angle
define lone pair
any pair of electrons that is NOT involved in a covalent bond - real definition
define bonded pair
one bonded pair of electrons in a covalent molecule
why repels more strongly, a bonded pair or a lone pair
a lone pair
why does a lone pair bond more strongly
- Lone pair of electrons is slightly closer to central atom and
- occupies more space than a bonded pair.
= Lone pair repels MORE strongly than bonding pair.
bond pairs + lone pairs of carbon dioxide
2 bonding pairs + 0 lone pairs
describe the shape of carbon dioxide
2 bonding pairs + 0 lone pairs • Linear shape
• 180 degrees angle
• SYMMETRICAL
bond pairs + lone pairs of boron trifluoride
3 bonding pairs + 0 lone pairs
shape of boron trifluoride
3 bonding pairs + 0 lone pairs • Triagonal planar shape
• 120 degrees angle
bond pairs + lone pairs of methane
4 bonding pairs + 0 lone pairs
shape of methane
4 bonding pairs + 0 lone pairs • - Tetrahedral shape
• 109.5 degrees angle
• SYMMETRICAL
bond pairs + lone pairs of ammonia
3 bonding pairs + 1 lone pair
shape of ammonia
3 bonding pairs + 1 lone pair • Pyramidal shape
• 107 degrees
bond pairs + lone pairs of water
2 bonding pairs + 2 lone pairs
shape of water
2 bonding pairs + 2 lone pairs • Non-linear shape
• 104.5 degree angle
bond pairs + lone pairs of sulphur hexaflouride
6 bonding pairs + 0 lone electrons
shape of sulphur hexafluoride
6 bonding pairs + 0 lone electrons
- Octahedral shape
- 90 degrees
shape of 2 bonding pairs + 0 lone pairs
• Linear shape
• 180 degrees angle
• SYMMETRICAL
eg. CO2
shape of 3 bonding pairs + 0 lone pairs
• Triagonal planar shape
• 120 degrees angle
eg. BF3
shape of 4 bonding pairs + 0 lone pairs
• Tetrahedral shape
• 109.5 degrees angle
• SYMMETRICAL
eg. CH4
shape of 3 bonding pairs + 1 lone pair
• Pyramidal shape
• 107 degrees
eg. NH3
shape of 2 bonding pairs + 2 lone pairs
• Non-linear shape
• 104.5 degree angle
eg. H20
shape of 6 bonding pairs + 0 lone electrons
- Octahedral shape
- 90 degrees
eg. SF6
angles of linear shape
180
angles of trigonal planar
120
angles of tetrahedral shape
109.5
angles of pyramidal shape
107
angles of non-linear
104.5
angles of octahedral shape
90 degrees
Table summary of all bonding pairs
number of BP + LP // diagram // shape name // angle degree
important note when counting bond pairs/lone pairs
Around the central atom!
define covalent bond
- The strong electrostatic attraction between a shared pair of electrons and the nuclei of bonded atoms
define electronegativity
- A measure of the tendency of an atom to attract a bonding pair of electrons in a covalent bond
what does it mean to have a greater electronegativity
- the more it attracts electrons towards it
if the atoms in a bonded molecule are the same, what is the polarity / electronegativity like
the electrons shared equally - non-polar
will the more electronegative atom have a smaller to larger share of the electrons
larger
relationship between electronegativity and dipoles
- The more electronegative – more electrons – slightly negative charge – negative dipole
- The less electronegative – less electrons – slightly positive charge – positive dipole
Electronegativity of group 0
There is none - 0
3
Factors affecting electronegativity
- Nuclear charge
- Atomic radius
- Electron shielding
what is the Pauling scale
- Compares electronegativity
How does electronegativity change across a period
Increases
Why does electronegativity increase across a period
- Because nuclear charge increases
- (One more proton on each element)
- And atomic radius decreases
- (Increased charge – pulls in shells
- – decreasing atomic radius)
- Attracts a bonding pair more strongly
How does electronegativity change down a group
Decreases
Why does electronegativity decrease down a group
- Atomic radius increases
- More shells = more shielding
- The bonding pair of electrons are attracted less strongly to the nuclei of an atom
Electronegativity difference in covalent
0
Electronegativity difference in Polar covalent
0 - 1.8
Electronegativity difference in Ionic
>1.8
Spectrum of bonds
What is a non-polar bond
- The electron pair is shared equally between atoms
When is a bond non-polar
- The bonded atoms are the same
- The bonded atoms have similar electronegativities
What is a polar bond
- The electron pair is shared unequally between atoms
- This forms a permanent dipole
What does a polar bond form
A permanent dipole
What is a dipole
A separation of opposite charges
When is a bond polar
The bonded atoms have different electronegativities
What is a polar molecule
- A molecule is polar if the individual dipoles do NOT oppose each other and cancel each other out
When would there be a polar molecule
- the shape of the molecule is NOT symmetrical - dipoles don’t cancel out
What is a non-polar molecule
- A molecule is non-polar if the individual dipoles oppose each other and cancel each other out
When would a molecule be non polar
If the shape is symmetrical - dipoles cancel each other out
Example of a non-polar molecule
CO2
Example of a polar molecule
H2O
Solubility of polar and non polar substances
- ‘Like’ dissolves ‘like’
- Polar solutes dissolve in polar solvents
- Non-polar solutes dissolve in non-polar solvents
What is an intermolecular force
- Weak interactions between dipoles of different molecules
What are intermolecular forces responsible for
- Responsible for physical properties
- E.g. melting/boiling point
3
Types of intermolecular forces
- Induced dipole-dipole interactions – London forces
- Permanent dipole-dipole interactions
- Hydrogen bonding
What can some intermolecular forces be referred to as
- Both London forces and induced dipole-dipole interactions can be referred to as Van der Waal’s forces
Table of strengths for intermolecular forces - don’t need to memorise exact numbers
Strongest type of intermolecular force
Hydrogen bonds
Weakest type of intermolecular force
London forces
How do induced dipole-dipole interactions (London forces) occur
- The constant movement of electrons produces an instantaneous dipole – the position of this is constantly changing
The instantaneous dipole induces a dipole on a neighbouring molecule - The induced dipole induces further dipoles on neighbouring molecules, which then attract one another
- The more electrons in a molecule, the stronger the London forces
- Because more electrons can be on one side/atom at once
How can you increase the strength of a London force
- more electrons = more electrons can be at one side at once = larger charge difference
How should we think of electrons
- Think of electrons as a cloud and is not in one shell but constantly moving in an area, so at one point it will be closer to one atom then another
What are London forces - indused dipole-dipole
Weak intermolecular forces that exist between ALL molecules
What are permanent dipole-dipole interactions
- Act between the permanent dipoles in different polar molecules
If something has a permanent dipole-dipole interaction, what else will it have
London force - acts between ALL molecules
More specifically what do London forces act between
ALL SIMPLE COVALENT
What are hydrogen bonds
- Special type of permanent dipole-dipole interaction found between molecules containing:
- An electronegative atom with a lone pair (O/N/F)
- i.e. H – O, H – N, H – F
- A hydrogen atom attached to an electronegative atom
How do you draw hydrogen bonds
- Dipoles must be shown
- Lone pairs must be shown
- The hydrogen bond is represented by a dashed line
What is a simple molecule
- Small units containing a definite number of atoms with a definite molecular formula
What do simple molecules form in the solid state
- a regular structure called a simple molecular lattice
How are molecules and atoms held together in a simple molecular lattice
- The molecules ae held together by weak intermolecular forces
- The atoms within each molecule are bonded together strongly by covalent bonds
3
Properties of simple molecular substances
- Low Melting and boiling point
- Solubility
- Electrical conductivity
How are simple molecular substances bonded
Covalently
why do simple molecular substances have a low melting and boiling point
- The weak intermolecular forces can be broken even by the energy present at low temps
- Therefore they have low boiling/melting points
What attraction is broken when a simple molecular lattice / substance melts/boils
- Only the weak intermolecular forces break
- The covalent bonds do NOT break
Example of a non-polar solvent
Hexane
Why are non-polar simple molecular substances soluble in non-polar solvents
- When a simple molecular compound is added to a non-polar solvent (e.g. Hexane), intermolecular forces form between the molecules and the solvent
- The interactions weaken the intermolecular forces in the simple molecular lattice.
- The intermolecular forces break and the compound dissolve
why is a non-polar simple molecular substance not soluble in a polar solvent
- When a simple molecular substance is added to a polar solvent, there is little interaction between the molecules in the lattice and the solvent molecules
- The intermolecular bonding within the polar solvent is too strong to be broken
Why may polar substances dissolve in polar solvents
- as the polar solute molecules and the polar solvent molecules can attract each other
- Similar to dissolving an ionic compound
Example of dissolving polar in polar
- E.g. sugar dissolves in water (a polar solvent)
- Sugar = polar covalent compound with many O-H bonds
- These attract and bond with polar water molecules
Why is solubility of a polar solute difficult to predict
depends on strength of the dipole
How can a substance dissolve in both polar and non-polar solvents and examole
If they have both polar and non-polar bonds
eg. ethanol, has both O-H polar bonds and non-polar C-C bonds = dissolve in both polar and non-polar substances eg.
define hydrophilic
- polar and contain electronegative atoms (usually O2) that can interact with water
Define hydrophobic
- non-polar and comprised of a carbon chain
What gives water anomalous properties
Hydrogen bonds
Anomalous properties of water
Density
high mp/bp
surface tension - cohesion
What is anomalous about the density of water
ice is less dense than water
Why is ice less dense than water
- Hydrogen bond hold water molecules apart in open lattice structure
- Water molecules in ice are further apart than in water
- Solid ice is less dense than liquid water and floats
- With two lone pairs on oxygen atom and two hydrogen atoms, each water molecule can form 4 hydrogen bonds. The hydrogen bonds extend outward, holding water molecules slightly apart, forming an open tetrahedral lattice full of holes. Bond angle about hydrogen atom in hydrogen bon, is close to 180 degrees.
- Holes in open lattice structure decrease density of water on freezing, when ice melts the ice lattice collapses and molecules move closer together so liquid water is denser than ice.
Consequences of ice being less dense than water
- Forms an insulating layer in ponds and lakes, preventing water from freezing solid, so fish can survive.
Why does water have a relatively high mp/bp
- Water has London forces between molecules
- Hydrogen bonds are extra forces over and above London ones
- Fair amount of energy need to break hydrogen bonds in water, so water has much higher mp/bp than expectedWater has London forces between molecules
- Hydrogen bonds are extra forces over and above London ones
- Fair amount of energy need to break hydrogen bonds in water, so water has much higher mp/bp than expected
What’s the difference between ice melts and water boils
- When ice lattice breaks down, rigid arrangement of hydrogen bonds in ice is broken.
When water boils, hydrogen bonds break completely.
What reduces surface tension
Detergents
Does water have a high or low viscosity
High
How do you increase the strength of London forces
More electrons = more electrons are on any given side at one point = instantaneous dipole is larger
More branching
Weaker London forces….less points of contact
Trend in group 5 hydrides b.p
Decreases then it increases
draw a dot and cross for ozone
