1.4 b: Polarity and Shapes of Molecules

Polarity and Shapes of Molecules

Year 11

​Syllabus Dot points:

• ​1.4 b investigate the differences between ionic and covalent compounds through:

– using nomenclature, valency and chemical formulae (including Lewis dot diagrams) (ACSCH029)

– examining the spectrum of bonds between atoms with varying degrees of polarity with respect to their constituent elements’ positions on the periodic table

– modelling the shapes of molecular substances (ACSCH056, ACSCH057)

​What we will cover:

• ​what is polarity

• polar bonds and molecules

• shapes of molecules ​

​What is polarity?

• a separation of electric charge that allows a molecule to have a dipole moment (a small negative and a small positive charge)

• ​molecules can be polar, and this means they must have at least one polar bond

• these polar bonds occur due to a difference in electronegativity between bonded atoms

​Electronegativity and Polarity of Bonds

• electronegativity is the power of an atom to attract an​ electron to itself

• if two atoms forming a bond have different electronegativities, then the one with the high electronegativity will attract the electron slightly more, and the sharing of the electrons negative charge will be uneven - this results in a polar bond

​Difference in electronegativity:

<0.4 = ​non-polar

0.4-1.4 = polar

>1.4 = ionic (complete polar) ​

Multiple Choice

The electrons in a polar covalent molecule are shared...

• Evenly

• Unevenly

• Electrons are not shared

• None of the Above

​Examples:

Cl₂ - two identical atoms

Therefore, the electrons are shared evenly. ​

Therefore, non-polar.

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​HCl - two different atoms, with different electronegativities.

Electrons are shared unevenly, with the Cl attracting the electrons more strongly. The Cl 'end' of the molecule is therefore slightly more negative than the H. We display with with small delta symbols (to imply a small or partial charge). Therefore, polar.

Multiple Choice

The electrons in a nonpolar covalent molecule are shared...

• Evenly

• Unevenly

• Electrons are not shared

• None of the Above

Polar bonds and polyatomic molecules

Sometimes, in polyatomic molecules, having a polar bond does not always mean the entire molecule will be polar. The shape of the molecule may determine whether dipole moments cancel each other out. ​​

For example:

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It is the shape that determines whether the molecule has a net dipole or not.

A lot of this is based on angular symmetry. ​

​How do we actually determine whether a molecule is polar or not?

1. identify whether there are any polar bonds

   • differences in electronegativity​

2. determine whether the polar bonds add up to give a polar molecule, or whether they cancel out to give a non-polar molecule

   • shape of the molecule is needed for this

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Multiple Choice

Is this molecule polar or non-polar?

• Polar

• Non-polar

Multiple Choice

Is this molecule polar or non-polar?

• Polar

• Non-polar

Multiple Choice

In this Lewis structure, the symbol above F means...

• electrons are being transferred to Fluorine
• electrons are less attracted to F than H
• electrons are more attracted to F than H
• Fluorine has formed a cation

​Predicting Shapes of Molecules

• ​main principle for predicting shape is based on valence shell electron pair repulsion theory - in that electrons around the atom arrange themselves to be as far away from each other as possible - repulsed by each other

• ​this means it is dependent on how many bonds it makes from the central atom, and how many lone pairs the central atom has (not always, but often)

​Examples:

BeCl₂ ​- the Be has 2 electron pairs, and no lone pairs.

Therefore, the bonds want to be as far away from each other as possible.

Therefore the molecule will be linear. ​

The bonds are 180° away from each other

​BCl₃ - the B has 3 bonded electrons, and no lone pairs.

Therefore, the bonds want to be as far away from each other as possible, which is at 120​° away from each other.

Therefore, the molecule will be planar/trigonal planar. ​

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​More examples:

CH₄ (methane) - four pairs of ​bonded electrons. They want to point as far away from each other as possible, so to the four corners of a shape, 109.5∘ away from each other.

Therefore, the shape is tetrahedral. ​

​NH₃ (ammonia) - three pairs of bonded electrons, and one lone pair. As far away from each other gives them a pyramidal shape. Similar to tetrahedral, but the lone pair still pushes the 3 hydrogens to the bottom of the molecule.

​More examples:

H₂O (water) - has 2 bonded electron pairs and two lone pairs. These four pairs are again directed towards the corners of a tetrahedron to stay as far apart as possible, ​but the final structure is in the shape 'bent'.

​Main takeaway's:

• electro​ns want to be as far away from each other as possible

• knowing t​he lewis structure of molecules assists in knowing the amount of bonded pairs and lone pairs

• after some time, you will come to know many of these off by heart - practice makes perfect here

Bringing it all together: ​

CCl₄

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What is the name of this compound?

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Is it ionic or covalently bonded?

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Are there polar bonds present?

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Is the molecule overall polar or non-polar?

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What shape would the molecule be? Draw the Lewis dot diagram and the spatial diagram (including partial charges).

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Bringing it all together: ​

CCl₄

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What is the name of this compound?

Carbon tetrachloride

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Is it ionic or covalently bonded?

Covalent

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Are there polar bonds present?

Yes, because carbon and chloride have differing electronegativities, and therefore there will be an uneven distribution of charge within the bonds, with the electron's being attracted more strongly to the chlorine atoms.

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Is the molecule overall polar or non-polar?

Overall the molecule will be non-polar as the polar bonds cancel each other out as the molecule is 100% symmetrical. ​

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What shape would the molecule be? Draw the Lewis dot diagram and the spatial diagram (including partial charges).

The shape is tetrahedral. ​

​Replace this with a header

PH₃​

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What is the name of this compound?

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Is it ionic or covalently bonded?

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Are there polar bonds present?

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Is the molecule overall polar or non-polar?

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What shape would the molecule be? Draw the Lewis dot diagram and the spatial diagram (including partial charges).

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PH₃​

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What is the name of this compound?

Phosphorus trihydride ​

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Is it ionic or covalently bonded?

Covalent

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Are there polar bonds present?

No, because phosphorus and hydrogen have the same electronegativity value, so therefore there is an equal sharing of electrons between the two atoms and the bonds are therefore non-polar.

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Is the molecule overall polar or non-polar?

Overall the molecule will be polar, this is because there is asymmetry in the atom, as there are three bonded pairs and a lone pair of electrons at the top of the molecules. When there is asymmetry in the molecule, then the overall molecule is polar. ​

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What shape would the molecule be? Draw the Lewis dot diagram and the spatial diagram (including partial charges).

The shape is pyramidal. ​

Multiple Choice

In a polar covalent bond, the electrons gather around...

• The atom with the Greatest Electronegativity

• The atom with the Lowest Electronegativity

• Each atom Equally

• None of the Above

Multiple Choice

Is this molecule polar or non-polar?

• Non-polar

• Polar

Multiple Choice

What geometry will this molecular structure have?: H₂S

• bent
• tetrahedral
• linear
• trigonal pyramidal

Multiple Choice

A diatomic molecule like O2 is always_______ because electrons are shared ________.

• nonpolar; equally
• polar; equally
• nonpolar; unequally
• nonpolar; unequally

Multiple Choice

What geometry will this molecular structure have?: CCl₄

• bent
• tetrahedral
• linear
• trigonal pyramidal

Next up:​

Intramolecular and intermolecular bonds ​

Allotropy ​