Covalent Model

Thus far: S2.2 → covalent model & Intermolecular forces

Covalent bond

the bond between atoms as a result of sharing (pairs of) electrons

• the atoms do this to gain stability

bond order (strength and type of covalent bond)

STRENGTH: single < double < triple

LENGTH: single > double > triple

relationship between bond strength, length and bond enthalpy

⬇ bond length = ⬆  strength = ⬆ bond enthalpy

coordination bond

when both electrons in a covalent bond originate from a single atom

*Notation of an arrow in lewis structure (ex. CO)

octet rule

the tendency of atoms to gain electrons to have a valence shell of 8 electrons

how to draw lewis structure

1. ID valence electrons of each atom (periodic table group)

• H is almost always a terminal

H cannot be the central atom of a molecule why?

Hydrogen only has 1 (valence) electron (to share) thus H is a terminal atom

lewis structure of CO₂

why is carbon special in its covalent bonding

Carbon always shares it’s 4 electrons

notation of ions in lewis structure

use brackets around the entire molecule and add the charge in the top left

see OH^- example

exceptions to the octet rule

Boron & Berylium — when central atoms will be ‘electron deficient’ (will have less than 8 valence ēs)

valence shell electron pair repulsion (VSEPR) model

allowing for the prediction of a molecules shape & geometry based on the repulsion of the atoms electron (pairs)

electron domain

regions of high electron density due to a “lone pair” of electrons

electron domain geometry

the shape the electron domains form

molecular geometry

the shape formed by (just) the atoms of a molecule

electron & molecular geometry + bond angle of 2 electron domains

linear electron & molecular geometry

bond angle: 180°

• the 2 electron domains repel each other

electron geometry & bond angle of 3 domains

trigonal planar electron geometry

bond angle: 120°

molecule geometry of 3 electron domains

if 1/3 of the domains is a lone pair then the molec geo becomes ‘V-shaped’

if 3/3 of the domains are bonding domains then the molec geo is trigonal planar (like the ē geo)

bond angle of 3 electron domain (exception rule)

if there is a lone pair as one of the electron domains; they will have a larger repulsion, thus the angle will be slightly smaller than 120°

electron geometry & bond angles of 4 electron domain

tetrahedral electron geometry

however the bond angle is dependent on the # lone pairs:

0 lone pairs = 109.5°

1 lone pair = 107°

2 lone pairs = 104.5°

molecular geometry of 4 electron domains

dependent on the # of lone pairs:

0 lone pairs = tetrahedral

1 lone pair = trigonal pyramidal

2 lone pairs = V-shaped

electronegativity

the ability of an atom to attract electrons (the electrons that are shared within a covalent bond) towards itself

explain polarity (of covalent bonding)

polar covalent bonds result from a difference in electronegativity of two bonded atoms

• they have a bond dipole (opposing difference in charge)

non polar bonds in relation to electronegativity

non-polar bonds mean there is an equal sharing of electrons within the [covalent] bond

how to identify polarity of bonds from data booklet values

differences of

0.1-0.4 = weakly polar covalent bond

0.5-1.7 = (fully) polar covalent bond

what makes a molecule polar

molecular geometry/net dipole moment

+

bond polarity

dipole

when a molecule has a -ve and a +ve end, between two bonded atoms

net dipole moment

having a -ve end and a +ve end, within a whole molecule

• this net value is based on the dipole between each atom of the molecule

thus you are able to “cut” the molecule in half and evidently have a +ve and -ve end

relate net dipole moment to molecular polarity

molecules can only be polar when there is a net dipole moment

TRUE OR FALSE: some molecules can have polar and non-polar regions.

TRUE

ex. large molecules

examples of non-polar molecules

• CO₂ / carbon dioxide

examples of polar molecules

• H₂O / water

TRUE OR FALSE: If all covalent bonds in a molecule are non-polar, the entire molecule is non-polar.

TRUE

There is no net dipole moment if the terminal atoms of a molecule with the geometry of _________ , ________, or ___________ are all the same.

• tetrahedral

• linear

• trigonal planar

how to signify electronegativity on lewis structure/formula of molecule

note: difference in electronegativity is demonstrated by a difference in length of the arrows

intermolecular forces

forces in-between molecules

a change in intramolecular forces results in…

a chemical change

a change in intermolecular forces

a physical change (change in physical properties such as melting pt and boiling pt)

what pattern can be identified about group 14 hydrides, in relation to their boiling point

as you proceed ⬇ the group, the boiling point of the elements hydride ⬆

what intermolecular force applies to all molecules

London dispersion forces (LDF)

london dispersion forces

a weak force caused by the movement of electrons within a molecule (or atom)

• this movement of electrons leads to a temporary instant dipole

what is the relationship between molar mass & boiling pt

positive correlation:

as boiling pt ⬆ = molar mass ⬆

explain LDFs using H₂ as an example

1. a pair of electrons will move (slightly) closer to one atom (within the bond) 

• thus making the H₂ molecule now slightly polar

2. if another molecule is nearby, the recently moved ēs will repel a pair of ēs of the neighbouring molecule

and so on and so forth…

polar molecules can have what types of intermolecular forces

• LDFs

• dipole-dipole

• H bonding

what molecules have dipole-dipole forces

polar molecules aka molecules with a net dipole moment

explain dipole-dipole intermolecular forces

• because of the difference in EN between the atoms in the molec, the electrons are ‘pulled’ closer to the more EN atom

what type of intermolecular force is hydrogen bond(ing)

a special dipole-dipole bond

when does H bonding form

when there is a strong dipole involving H

• H is bonded with F, O or N —highly EN atoms

how to signify H bonds in a diagram

|||||||||||||||||||

rank intermolecular forces from weakest to strongest

LDFs < dipole-dipole < H bonds

explain H bonding

as a result of dipole-dipole interactions between H and atoms such as F, O or N

1. the H end of the molecule now has a partial +ve polarity

2. electron deficient H then attracts the lone pair of a nearby EN atom (that is of another molecule)