SCH4U Molecular Shape and Polarity: Notes and Practice

Carbon dioxide and water are both small molecules made of three atoms, yet water is a liquid that dissolves salt and carbon dioxide is a gas that does not. A large part of the difference is shape: carbon dioxide is a straight line and water is bent. Shape decides whether the charges in a molecule cancel out or add up, and that in turn controls boiling point, solubility and how molecules fit together, including how drug molecules fit into the proteins in your body. The valence shell electron pair repulsion (VSEPR) model lets you predict shape from a Lewis structure using one idea: groups of electrons around a central atom repel each other and spread out as far apart as possible.

Count the total valence electrons. For an ion, add one electron for each negative charge and subtract one for each positive charge. For $NH_{4}^{+}$ : $5 + 4 (1) - 1 = 8$ .
Draw the Lewis structure: put the least electronegative atom in the centre (never hydrogen), connect each outer atom with a single bond, complete the outer atoms' octets (hydrogen needs only 2), and place any electrons left over on the central atom as lone pairs.
Count the electron groups on the central atom. Each lone pair is one group, and each bond to another atom is one group, whether it is single, double or triple.
Use the number of groups to find the electron-group arrangement, then look only at where the atoms are to name the molecular shape.

Lone pairs take up more room

A lone pair is held by only one nucleus, so it spreads out closer to the central atom and pushes harder than a bonding pair. Repulsion decreases in the order lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair. That is why the angle shrinks from 109.5° in

CH_{4}

to about 107° in

NH_{3}

(one lone pair) and about 104.5° in

H_{2} O

(two lone pairs).

Groups	Lone pairs	Electron-group arrangement	Molecular shape	Bond angle	Examples
2	0	linear	linear	180°	$CO_{2}$ , $HCN$
3	0	trigonal planar	trigonal planar	120°	$BF_{3}$ , $SO_{3}$
3	1	trigonal planar	bent	slightly less than 120°	$SO_{2}$ , $O_{3}$
4	0	tetrahedral	tetrahedral	109.5°	$CH_{4}$ , $NH_{4}^{+}$
4	1	tetrahedral	trigonal pyramidal	about 107°	$NH_{3}$ , $H_{3} O^{+}$
4	2	tetrahedral	bent	about 104.5°	$H_{2} O$

Shapes for central atoms that obey the octet rule.

Fig 1Linear, trigonal planar, tetrahedral, trigonal pyramidal and bent shapes.

Expanded Octets: Five and Six Electron Groups

Central atoms from period 3 and below, such as phosphorus, sulfur, chlorine, bromine, iodine and xenon, are large enough to hold five or six electron groups around them, so they can have more than eight valence electrons. Five groups form a trigonal bipyramid: three equatorial positions in a triangle 120° apart, plus two axial positions straight above and below. Six groups form an octahedron, with every position 90° from its four neighbours.

Groups	Lone pairs	Electron-group arrangement	Molecular shape	Bond angles	Example
5	0	trigonal bipyramidal	trigonal bipyramidal	90°, 120° and 180°	$PCl_{5}$
5	1	trigonal bipyramidal	seesaw	less than 90° and less than 120°	$SF_{4}$
5	2	trigonal bipyramidal	T-shaped	slightly less than 90°	$ClF_{3}$
5	3	trigonal bipyramidal	linear	180°	$XeF_{2}$ , $I_{3}^{-}$
6	0	octahedral	octahedral	90° and 180°	$SF_{6}$
6	1	octahedral	square pyramidal	slightly less than 90°	$BrF_{5}$ , $IF_{5}$
6	2	octahedral	square planar	90°	$XeF_{4}$

Shapes for central atoms with expanded octets.

Fig 2The six shapes built on the trigonal bipyramid and the octahedron, each with an example molecule.

Where the lone pairs go

In a trigonal bipyramid, lone pairs always take equatorial positions. An equatorial position has only two neighbours at 90°, while an axial position has three, so a lone pair there is squeezed less. Filling equatorial positions with 1, 2 and 3 lone pairs gives seesaw, T-shaped and linear. In an octahedron, all positions are equivalent for the first lone pair; a second lone pair goes directly opposite the first (180° away), which leaves the four atoms in a square plane.

Bond polarity. Electronegativity is an atom's ability to attract the shared electrons in a bond. When two atoms with different electronegativities bond, the electrons spend more time near the more electronegative atom, giving it a partial negative charge ( $δ^{-}$ ) and leaving the other atom partially positive ( $δ^{+}$ ). This separation of charge is a bond dipole, drawn as an arrow pointing toward the more electronegative atom with a cross on its tail at the positive end. The bigger the electronegativity difference ( $Δ EN$ ), the more polar the bond. Bond types form a continuum, but a common guideline is below.

Electronegativity difference	Bond type	Example (Pauling values)
less than about 0.5	nonpolar covalent (electrons shared about equally)	C–H: 2.55 and 2.20, difference 0.35
about 0.5 to 1.7	polar covalent (electrons shared unequally)	O–H: 3.44 and 2.20, difference 1.24
greater than about 1.7	mainly ionic (electrons transferred)	Na–Cl: 0.93 and 3.16, difference 2.23

Guideline ranges only. Electronegativity values differ slightly between tables.

Molecular polarity: shape decides. A molecule is polar when its bond dipoles do not cancel, leaving one end partially positive and the other partially negative. Bond dipoles are vectors, so shape matters as much as bond polarity. In $CO_{2}$ the two C=O dipoles are equal and point in exactly opposite directions, so they cancel: the molecule is nonpolar even though both bonds are polar. In bent $H_{2} O$ the two O–H dipoles both point partly toward oxygen, so they add up and water is strongly polar.

Fig 3Bond dipole arrows cancel in symmetrical carbon dioxide but add up in bent water.

AUsually nonpolar (identical outer atoms, no net dipole)

Linear with no lone pairs on the centre: $CO_{2}$
Trigonal planar: $BF_{3}$ , $SO_{3}$
Tetrahedral: $CH_{4}$ , $CCl_{4}$
Trigonal bipyramidal: $PCl_{5}$
Octahedral: $SF_{6}$ ; square planar: $XeF_{4}$ ; linear with three lone pairs: $XeF_{2}$

BPolar (bond dipoles do not cancel)

Bent: $H_{2} O$ , $SO_{2}$
Trigonal pyramidal: $NH_{3}$
Seesaw: $SF_{4}$ ; T-shaped: $ClF_{3}$ ; square pyramidal: $BrF_{5}$
Any symmetrical shape with different outer atoms: $CH_{3} Cl$ , $CH_{2} Cl_{2}$
Any diatomic molecule with a polar bond: $HCl$

VSEPR and McMaster

The VSEPR model was developed in 1957 by Ronald Gillespie and Ronald Nyholm at University College London. Gillespie then moved to McMaster University in Hamilton, where he spent a long career extending and teaching the model that chemistry students around the world now use to predict molecular shapes.

Molecular Shape and Polarity

Key things to remember

Expanded Octets: Five and Six Electron Groups

Worked examples

Where marks get lost

Flashcards 9

Practice