Chemistry notes, Year 10, Mr. Blaurock
1. Structure of molecules and material properties
1.1
The orbital model
 See handout/whiteboard-pdf
Additional info:
-
1.2
Each main group element (exceptions: hydrogen, helium) can have up to 4 valence orbitals
Electrons will always spread out into as many orbitals as possible, occupying them singly
before filling them in pairs (Hund’s rule)
Atomic bonds in the orbital model
 See handout/whiteboard-pdf
Recap of important principles/terms
Octet rule:
The goal of every atom from a main group element is to reach 8 valence electrons (-> an octet) and
therefore completing its valence shell by forming chemical bonds
LEWIS-structure:
A graphical representation of molecules.
Example: H2CO
 Element symbols represent the
nucleus and inner electrons of an
atom
 valence electrons are drawn as dots (single electron) or lines (electron pair)
1.3
The VSEPR-model
The 3d geometry of molecules can easily be appraised following the principle that all electron pairs
(lone AND binding) around a molecule’s central atom repulse each other. Because of this all orbitals
occupy the spaces furthest away from each other.
Further rules:
-
For the purpose of geometry double and triple bonds are treated as single bonds
Lone electron pairs occupy slightly more space than binding electron pairs
This set of rules is known as VSEPR-model (valence shell electron pair repulsion)
Chemistry notes, Year 10, Mr. Blaurock
Possible geometries:
Instructions on how to draw a 3d tetrahedron (Natta-projection) *
These lines are on the
plane of the page
This cone goes out
behind the plane of
the page
This cone goes out in
front of the plane of
the page
Chemistry notes, Year 10, Mr. Blaurock
Advanced VSEPR-geometries
In terms of 3d structure, lone electron pairs around a central atom take up space just as binding
electron pairs would.
Example
But since lone electron pairs take up more space than binding electron pairs, the geometries that
result have smaller angles between the remaining bonds.
Example
The naming of the resulting geometries is differs from the standard naming because only the
chemical bonds are considered when identifying the molecule’s shape.
Example
Chemistry notes, Year 10, Mr. Blaurock
Overview
chemical formula
1.4
LEWIS-structure
3d geometry
*
Electronegativity *
Electronegativity (EN) is a chemical property that describes the relative tendency of an atom (of a
certain element) to pull electrons toward itself. The larger the EN-value, the stronger the pull.
In a chemical bond the binding electron pair (BEP) is pulled towards the atom with the higher ENvalue. Flourine (F) is the element with the highest EN-value, which has been deliberately set to 4.0.
1.5
Polar bonds *
A chemical bond between two atoms with differing EN-values is polar.
The atom with the lower EN-value carries a positive partial charge, because the probability of the
binding electrons being located around the nucleus is lowered.
The atom with the higher EN-value carries a negative partial charge, because the probability of the
binding electrons being located around the nucleus is higher.
The larger the difference in electronegativity (abbreviated as ΔEN) is, the stronger the polarity (and
the larger the partial charges) of the bond will be.
A chemical bond between two atoms of the same EN-value (e.g. two atoms of the same element) is
non-polar. The binding electrons are distributed equally.
Chemistry notes, Year 10, Mr. Blaurock
Examples
 Partial charges are represented by the greek letter delta and the charges sign.
A positive partial charge is represented by δ+, a negative partial charge by δJudging the polarity of a bond *
When talking about the polarity of a bond, it is often not enough to differentiate between non-polar
bonds at ΔEN=0 and polar bonds at ΔEN>0. Bonds with a small ΔEN are usually considered more
non-polar. If the ΔEN is very high the electron will completely transfer from the atom with the lower
EN-value to the atom with the higher one. This means the bond transitions from polar to ionic.
The result is a categorization of bonds as follows (the border values can be blurry):
1.6
ΔEN
type
<0,4
non-polar
0,4-1,7
polar
>1,7
ionic
Polar and non-polar molecules
If the distribution of positive partial charges and negative partial charges across a molecule (in 3d
space!) leads to a separation of the center of positive partial charges and the center of negative
partial charges (-> creating a “positive” and a “negative” side of the molecule), the whole molecule is
called polar. If the two centers of charges overlap, the molecule is non-polar.
Example (H2O)
 For a molecule to be polar there have to be polar bonds (and therefore partial charges)
 A molecule with polar bonds is not necessarily polar altogether (if the centers of charges
overlap)
Chemistry notes, Year 10, Mr. Blaurock
1.7
Intermolecular forces
a) Dipole-Dipole-forces
Polar molecules (also called dipoles because of their two - positive and negative - sides) are not only
attracted by outside electrical fields. The opposing charges of neighboring molecules are attracted
mutually. This means there is an attracting force between the different molecules.
Example: HCl
This attracting force is called dipole-dipole-force. The greater the polarity of a molecule is (-> the
greater ΔEN), the stronger the attraction becomes.
b) Hydrogen bonds
Hydrogen bonds are a specific type of dipole-dipole-force, that occur if a number of prerequisites are
met:


a hydrogen atom is bound to a nitrogen, oxygen or fluorine atom (which are all highly
electronegative)
a nearby nitrogen, oxygen or fluorine atom hat at least one lone electron pair
The resulting intermolecular forces are formed between the hydrogen atoms and lone electron pairs
described above.
Example: H2O
 Hydrogen bonds are much stronger than usual dipole-dipole forces
 Despite the name, hydrogen bonds are not the same as covalent chemical bonds
Chemistry notes, Year 10, Mr. Blaurock
c) Van-der-Waals-forces
1.8
Effects of intermolecular forces
1) The stronger intermolecular forces in a substance are, the higher it’s melting and boiling
temperatures will be.
2) Polar and non-polar substances do not mix, because the predominant intermolecular forces
(dipole-dipole-forces and Van-der-Waals-forces) are incompatible.
3) Solid polar substances (e.g. salts) can only be solved in polar fluids (e.g. water). Because of
this they are called hydrophilic (Greek: “water loving”). Solid non-polar substances (e.g.
iodine) can only be solved in non-polar fluids (e.g. oil). Because of this, they are called
lipophilic (Greek: “fat loving”).
P.S.: lipophilic
=
hydrophobic (“hates/fears water”)
hydrophilic
=
lipophobic (“hates/fears fats”)