Module 2: Chapter 6 (Shapes of molecules and intermolecular forces) Flashcards

1
Q

What are the 3D symbols used when drawing molecules?

A

solid line = a bond in the plane of paper
solid wedge = comes out of the plane
dotted wedge = going into the plane

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2
Q

Draw the example of two electron pairs: BeCl2, state bond angle and the shape

A

bond angle = 180

linear

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3
Q

Draw the example of 3 electron pairs: BF3, state bond angle and the shape

A

bond angle = 120

trigonal planar

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4
Q

Draw the example of 4 electron pairs: CH4, state bond angle and the shape

A

bond angle = 109.5

tetrahedral

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5
Q

Draw the example of 6 electron pairs: SF6, state bond angle and the shape

A

bond angle = 90

octahedral

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6
Q

What is the effect of lone pairs?

A
  • lone pairs are slightly closer to the central atom, occupies more space
  • results in lone pair repelling stronger than the bonded pair
  • they repel the bonding pairs slightly closer together, decreasing the bond angle 2.5 per lone pair
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7
Q

Draw the example of 4 electron pairs with one being a lone pair: NH3, state bond angle and the shape

A

bond angle = 107

pyramidal

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8
Q

Draw the example of 4 electron pairs with 2 being lone pairs: H20, state bond angle and the shape

A

bond angle = 104.5

non linear

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9
Q

How do you deal with multiple bonds or polyatomic ions?

A

multiple bonds treated as a bonding pairs, same with polyatomic ions

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10
Q

Draw the example of CO2, state bond angle and shape

A

bond angle = 180

linear

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11
Q

Draw the example of carbonate, state bond angle and shape

A

CO3^2-
bond angle = 120
trigonal planar

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12
Q

Draw the example of ammonium, state bond angle and shape

A

NH4+
bond angle = 109.5
tetrahedral

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13
Q

Draw the example of sulphate, state bond angle and shape

A

SO4^2-
bond angle = 109.5
tetrahedral

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14
Q

What is electronegativity?

A

Electronegativity is a measure of the attraction of a bonded atom for the pair of electrons in a covalent bond

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15
Q

How is electronegativity measured?

A

pauling scale

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16
Q

Describe electronegativity of elements

A
  • non metals: nitrogen, oxygen, fluorine and chlorine are the most electronegative (tend to form δ-), all above 3.0
  • group 1 metals including lithium, sodium and potassium are the least electronegative (δ+), all below 1.0
17
Q

What is the list of comparative elements for electronegativity?

A

δ+ = C,H

δ- (in descending order of electronegativity) = 
F
O
Cl, N (similar value)
Br
I
18
Q

Describe bond polarity

A

In a non-polar bond, bonded electrons are shared equally between bonded atoms. This happens when:
• bonded atoms are the same
• have a similar/equal electronegativity (for example a C-H bond is ALWAYS NON-POLAR)

In a polar bond, bonded electrons are shared unequally. This happens when:
• atoms have different electronegativity values -> more electronegative atom has a small partial negative charge and the other a small partial positive charge -> separation of opposite charges is called a dipole

19
Q

What is the exception for polar molecules?

A
  • dipoles can cancel each other out if the shape is symmetrical, results in the molecules itself being non-polar with no overall dipole
  • for example: CO2 (linear) and CCl4 (tetrahedral)
  • if there were lone pairs it would be polar
20
Q

What are intermolecular forces?

A
  • weak interactions between neighbouring molecules
  • broken when a molecular substance melts, boils or dissolves
  • significantly weaker than the covalent bonds
  • three main types: london forces, permanent dipole-dipole and hydrogen bonding
21
Q

Describe london forces

A
  • exist between all molecules, polar or non-polar
  • movement of electrons produces a change in dipole in a molecule and at any point an instantaneous dipole will exist, constantly changing
  • the instantaneous dipole induces a dipole on a neighbouring molecule, which induces further molecules and attract one another
  • they are only temporary, in the next instant could disappear
22
Q

What factors increase the strength of london forces?

A

The more electrons in each molecule
• larger the instantaneous and induced dipoles
• greater the interactions
• stronger attractive forces

For example as you go down the noble gases group, boiling point increases as there are more electrons and therefore stronger induced dipoles

23
Q

Describe permanent dipole-dipole interactions

A
  • only act between permanent dipoles in polar molecules

- more energy than London forces needed to overcome forces -> gives substances higher boiling points

24
Q

What is a hydrogen bond?

A

A hydrogen bond is an attraction between a
lone pair of electrons on a highly electronegative
atom (O, N or F) in one molecule and an
electron-deficient hydrogen atom (H attached to
a highly electronegative O, N or F atom) in
another molecule

25
Define a simple molecular substance
made up of simple molecules, i.e. small covalently bonded units with a definite molecular formula
26
Describe the melting and boiling points of simple molecular substances
- in a solid substance the molecules form the structure a simple molecular lattice, where molecules are held in place by weak molecular forces - little energy is required to break these forces and therefore typically have low melting and boiling points
27
Describe electrical conductivity of simple molecular substances
- non conductors | - no mobile charge carriers (either delocalised electrons or mobile ions)
28
Why is liquid water denser than solid water?
- In ice each water molecule is involved in 4 stable H-bonds in a tetrahedral arrangement. - Extensive H-bonds lead to a crystalline structure - This creates an open lattice with quite large gaps which lowers the density (the gaps are empty. NOT filled with air!) - When heated some of the H-bonds in ice break. H2O molecules fill some of the gaps increasing the density
29
Describe surface tension in water
Surface tension in water is caused by hydrogen bonding. - Surface tension is caused by molecules on surface experiencing unbalanced hydrogen bonding forces pulling them in. - Molecules in the bulk experience balanced forces in every direction.
30
Describe the high melting and boiling point of water
- hydrogen bonds are extra forces over and above London - more energy is needed to break the H bonds in water - when the ice lattice breaks the rigid arrangement of hydrogen bonds in ice is broken and when water boils the H-bonds break completely
31
What is the solubility of non-polar simple molecular substances?
- Soluble in non-polar solvents as intermolecular forces form between molecules and the solvent, weaken lattice forces and eventually breaking them all - tend to be insoluble in polar solvents as intermolecular bonding within polar solvents may be too strong to be broken
32
What is the solubility of polar simple molecular substances?
- soluble in polar solvents as molecules can attract each other - solubility mainly depends on the strength of the dipole
33
Explain why water molecule are polar
O is more electronegative than H (O attracts the bonded pair of electrons in the covalent bond between O and H more than H) H2O is non-symmetrical. Dipoles do not cancel.