Shapes of Molecules and Ions | A-level Chemistry | OCR, AQA, Edexcel

Описание: Shapes of Molecules and Ions in a Snap!

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The key points covered in this video include:

1. Shapes of Molecules and Ions
2. Electron Pair Repulsion
a) Shrinking
b) Expanding
4. Specific Examples of Shapes

The Shapes of Molecules and Ions

The shape of a compound or ion is dictated by Number of electron pairs around the central atom, The nature of these pairs: Bonding Pair, Lone Pair.

Electron Pair Repulsion

Bonding Pairs. Pairs of electrons that are involved in bonding. These pairs repel each other equally. Lone Pairs. Pairs of electrons that are not involved in bonding. Lone pairs repel other pairs more than bonding pairs. More electron dense. Each lone pair reduces the bond angle by about 2.5°.

3D Shapes of Molecules and Ions

Molecules have a 3D shape. This is challenging to represent on paper! There is a simple convention when drawing 3D shapes: Normal Line. Bond in the plane of the paper. Dotted Wedge. Bond is going into the paper away from you. Bold Wedge. Bond is coming out of the paper towards you.

The Octet Rule

Through chemical bonding, elements usually achieve a Noble Gas configuration. Achieve a full outer shell. ‘Octet Rule’. 8 electrons in pairs. 4 electrons form covalent bonds. However this is not always possible. There may be: Too few electrons to form an octet. More than enough electrons to form an octet.

Shrinking the Octet

In their outer shells, neither Beryllium nor Boron have enough electrons to pair and form an Octet. The unpaired electrons will pair up. The element will not achieve an octet. Example Boron Trifluoride.

Expanding the Octet

Some elements can expand their octet. As a result, the bonding atom will more than 8 electrons in the outer shell. This can occur in Group 15-17, from Period 3 downwards. Group 15, Non-Metals. 3/5 covalent bonds. Group 16 Non-Metals. 2/4/6 covalent bonds. Group 17 Non-Metals. 1/3/5/7 covalent bonds. e.g. Sulfur - Group 16, Period 3.

1 Pair : Linear

1 bonding pair. Example: H2-Hydrogen.

2 pairs : Linear

2 bonding pairs. Example: CO2 - Carbon Dioxide.

3 Pairs: Trigonal Planar

3 bonding pairs. Example: BF3 - Boron Trifluoride.

4 Pairs: Tetrahedral

4 bonding pairs. Example: CH4- Methane.

4 Pairs: Pyramidal

3 bonding pairs. 1 lone pair. Example: NH3- Ammonia.

4 Pairs: Non-Linear

2 bonding pairs. 2 lone pairs. Example: H2O - Water.

5 Pairs: Trigonal Bipyramid

5 bonding pairs. Example: PCl5 - Phosphorous Pentachloride.

6 Pairs: Octahedral

6 bonding pairs. Example: SF6 - Sulfur Hexafluoride.

Summary

Electron Pair Repulsion
a. Electron pairs repel each other
b. Pairs take specific positions in order to minimise repulsion
c. Results in specific shapes of molecules and ions
2
a. Linear
i. 2 bonding pairs
ii. 180°
iii. CO2, BeCl2
3
a. Trigonal Planar
i. 3 bonding pairs
ii. 120°
iii. BCl3
4
a. Tetrahedral
i. 4 bonding pairs
ii. 109.5°
iii. CH4
b. Pyramidal
i. 3 bonding pairs, 1 lone pair
ii. 107°
iii. NH3
c. Non-Linear
i. 2 bonding pairs, 2 lone pairs
ii. 104.5°
iii. H2O
5
a. Trigonal Bipyramid
i. 90° and 120°
ii. PCl5
6
a. Octahedral
i. 90°
ii. SF6