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

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Flashcards in Module 2: Chapter 6 (Shapes of molecules and intermolecular forces) Deck (33)
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1

What are the 3D symbols used when drawing molecules?

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

2

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

bond angle = 180
linear

3

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

bond angle = 120
trigonal planar

4

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

bond angle = 109.5
tetrahedral

5

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

bond angle = 90
octahedral

6

What is the effect of lone pairs?

- 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

7

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

bond angle = 107
pyramidal

8

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

bond angle = 104.5
non linear

9

How do you deal with multiple bonds or polyatomic ions?

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

10

Draw the example of CO2, state bond angle and shape

bond angle = 180
linear

11

Draw the example of carbonate, state bond angle and shape

CO3^2-
bond angle = 120
trigonal planar

12

Draw the example of ammonium, state bond angle and shape

NH4+
bond angle = 109.5
tetrahedral

13

Draw the example of sulphate, state bond angle and shape

SO4^2-
bond angle = 109.5
tetrahedral

14

What is electronegativity?

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

15

How is electronegativity measured?

pauling scale

16

Describe electronegativity of elements

- 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

What is the list of comparative elements for electronegativity?

δ+ = C,H

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

18

Describe bond polarity

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

What is the exception for polar molecules?

- 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

What are intermolecular forces?

- 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

Describe london forces

- 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

What factors increase the strength of london forces?

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

Describe permanent dipole-dipole interactions

- only act between permanent dipoles in polar molecules
- more energy than London forces needed to overcome forces -> gives substances higher boiling points

24

What is a hydrogen bond?

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