Bonding Notes
Bonding
Chemical Bonding between atoms involves interaction of electrons
in the valence shell of the atoms.
Three fundamental types
Metallic
Ionic
Covalent
Bond type depends on attraction for electrons in the atom
involved (electronegativity)
If electrons have very different electronegativities ionic
bonding If elements both have quite high electronegativities
covalent
Both have low electronegativities form a metallic bondTypes of
bonds give rise to distinct physical properties
Ionic
Between metals and non metals
Metals lose electrons form + charged ion or cations
# of e- lost = # e in valence level
non metals gain electrons form charged ions or anions
# e- gained = # e- to fill valance level
groups of atoms joined by covalent bonds can also have
electrical charge and therefore form compounds with ions of the
opposite charge by ionic bonding (SO42-)
Anions and cations have opposite electrical charges and are
attracted into a crystal lattice in which each anion is surrounded
by cations and vice versa (3 dimensional structure)
The whole substance is held together by electrostatic
attractions in all 3 dimensions.
Covalent Bond
Occurs between atoms that have high electronegativities, i. e.
non-metals
Involves two atoms sharing some of their valence electrons
The attraction of the two nuclei for these shared electrons
results in the two atoms being bonded together
Covalent bonds usually form so that the bonded atoms achieve
octet configurations.
Lewis symbols: show the number of valence electrons
H
C
Cl
N
O
P
Single covalent bond
Consists of a shared pair of electrons
Usually each atom involved contributes one electron
Usually fill their valence level so the number of bonds formed
equals the number of electrons
Lewis Diagrams (electron dot diagrams)
H
F
H C H
F C F
H
F
Cl Cl
H Cl
Shared pairs of electrons
Lone pairs
In some circumstances one atom can donate both electrons. In
this special case, the bond is known as a dative covalent bond.
(Will return to this idea during solution chemistry)
More than one pair of electrons can be shared between the same
atoms, resulting in what is called a multiple covalent bond.
In a double covalent bond (or double bond), two electron pairs
are shared between the same two atoms.
Double bonds join atoms more tightly than a single bond.
C forms four bonds
O forms two bonds
Therefore two oxygens are needed for each carbon.
O=C=O
H2C=CH2
Note that the valence levels of both the carbon and oxygen are
filled
In a triple covalent bond (triple bond) three electron pairs are
shared between the same two atoms
N N
C O
Electronegativity of elements
Covalent bonding is the sharing of electrons.
Unless the two atoms sharing the molecules are identical, the
sharing will not be equal
The more electronegative atom will attract the electrons more
strongly that the less electronegative atom.
This will result in the more electronegative atom having a
slight negative charge ((-)
The less electronegative atoms will be slightly deficient in
electrons and thus have a slight positive charge ((+)
This type of covalent bond in which atoms have slight electrical
charges are called polar bondsEg.
Electronegativity of elements can be judged from their position
on the periodic table.
All elements involved in covalent bonds do have quite high
electronegativities.
B & Si < P & H < C & S & I < Br < Cl
& N < O < F
High
Very High
Extremely
High
The greater the difference in electronegativity of the atoms
involved the greater the polarity of the bond.
Eg.
In many molecules, the polar bonds result in the molecule have a
resultant dipolei.e. there is a positive and negative end to the
molecule
Eg.
In some molecules, the effects of the polar bonds cancel out
because of the symmetry of the molecule.
Eg.
For a molecule to be polar
It must contain polar bonds
Its shape must be such that the centre of positive and negative
charges are not in the same place
Experimentally it is easy to determine if a liquid is polar by
bringing a charged rod close to a stream of liquid running out of a
burette.
Shapes of molecules
The shapes of molecules are determined by the repulsion between
the electron pairs in the valence level.
This is known as the Valence Shell Electron Pair Repulsion
Theory (VSEPR Theory)
In the most common molecules, filled valence levels contains
four pair of electrons.
In order for these electron pairs to be as widely separated as
possible, they distribute themselves so that they are pointing
towards the corners of a tetrahedron (regular triangular based
pyramid).
e.g. CH4Some molecules also contain non-bonding or lone pairs of
electrons. These lone pairs affect the shape of the molecule.
In ammonia, NH3, there are three pairs of bonded electrons and
one lone pair.
Without the lone pair, the molecule would look like an
equilateral triangle with a nitrogen atom at its centre.
The lone pair that is not involved in bonding repels the bonding
electrons so the molecule has the shape of a trigonal pyramid.
In water, there are two pairs of bonded electrons and two lone
pairs.
The two lone pairs cause the molecule to be bent, not linear as
would be expected without the lone pairs.
Resonance Structures
Often two or more equivalent structures can be drawn for a
molecule. These are known as resonance structures.
Resonance refers to the arrangement of valence electrons in
molecules or ions for which more than one Lewis structure can be
written.
The actual molecule that exists is often referred to as a
resonance hybrid of these structures.
Eg. CO32-
O3Resonance does NOT mean than the molecule flips from one
structure to another. The bonds actually have lengths and strengths
intermediate between those of single and double bonds or double and
triple bonds.
Metallic Bonding
In a metal all the atoms are all packed together as closely as
possible (like marbles in a box). This type of regular framework is
known as a lattice.
Valence electrons are delocalized amongst (shared by) all the
atoms, so that no electron belongs to any particular atom and they
are free to move throughout the metal.
The atoms, having lost electrons, are better described as ions
as they are positively charged. The attraction of these positive
ions towards the mobile electrons provides the force that holds the
structure together.
The attraction is between the ions and the mobile electrons, not
between the ions themselves. This allows the layers of ions to
slide past each other without the need to break the bonds in the
metal.
Hence, metals are malleable and ductile.
As electrons are free to move from one side of the lattice
structure to the other, they can also carry an electric current
(good conductors of electricity).
Mobile electrons also make them good conductors of heat.
The strength of the bond between the metal atoms depends on how
many electrons each atom shares with the other and how far from the
positive nucleus the electrons are. (ionic radius)
Intermolecular Forces
Covalent bonding can result in either a giant structure or in a
molecular structure.
In molecules, weak forces exist between the molecules.
If these forces did not exist there would never be condensing to
form liquid and gases.
Three types of intermolecular forces are, in increasing
strength:
1. van der Waals Forces
2. Dipole Dipole Forces
3. Hydrogen Bonding
1. van der Waals Forces (also known as London Forces or
dispersion forces)
act on ALL atoms and ALL molecules, whether polar or
non-polar
are responsible for the condensation, at very low temperatures,
of even the monatomic noble gases
are the result of momentary shifts in the symmetry of the
electron cloud of a molecule
causes a temporary dipole in one molecule that has an inductive
effect on neighbouring molecules, the net result being an
attractive force
the strength of van der Waals forces is influenced by the size
and the geometry of the molecules involved and by the ease of
polarization of the electron clouds
forces get stronger as
the molecule gets less spherical in shape
molecular mass increases - increasing number of electrons (among
molecules with similar geometry
van der Waals forces are the forces of attraction between
fluctuating dipoles in atoms and molecules that are very close
together.
2. Dipole Dipole interaction
Molecules with dipole moments attract each other
electrostatically (positive end of one molecule attracts the
negative end of another molecule.
The dipole-dipole attraction, along with other forces, must be
overcome in melting a solid or vapourizing a liquid.
Therefore, dipole-dipole interactions influence the melting
point, heat of fusion, boiling point, and heat of vaporization.
3. Hydrogen bonds
When hydrogen atoms are covalently bonded to electronegative
atoms that strongly attracts the shared electron pair, the small
hydrogen atoms have little electron density around them.
Hydrogen atoms will carry a partial positive charge and act as a
bridge to another electronegative atom.
Hydrogen bond the attraction of a hydrogen atom covalently
bonded to an electronegative atom for a second electronegative
atom.
Strong H bonds will form between F, N or O atoms, which are
small and have negative charges highly concentrated in a small
volume
H bonds are the strongest intermolecular forces.
Water is most greatly affected as it has two H atoms and two
non-bonding electron pairs. Thus, it can form two hydrogen bonds
per molecule.
This accounts for many of its anomolous properties, such as ice
having a density less than that of water, as well as influencing
its properties as a solvent.
H bonding is also of great biological importance.
Provides the pairing of bases in DNA and the structure of
protein molecules.
Steps for writing Lewis structures
1. Write the correct arrangement of the atoms using single
bonds. Where necessary, apply the following guidelines.
a) Smaller, more electronegative non-metal atoms surround
larger, less electronegative non-metal atoms.
b) Oxygen, hydrogen, and/or halogen atoms often surround a
central metal or non-metal atom in a symmetrical arrangement.
c) Carbon atoms are usually bonded to each other.
d) Oxygen atoms are bonded to each other only in peroxides (or
superoxides)
e) In most acids, such as H2SO4, and in many other compounds
that contain both oxygen and hydrogen atoms, the hydrogen atoms are
all bonded to oxygen atoms.
2. Find the total number of valence electrons. Add together the
number of valence electrons contributed by each atom. If the
species is an ion, subtract one electron for each unit of positive
charge or add one electron for each unit of negative charge.
3. Assign two electrons to each covalent bond.4. Distribute the
remaining electrons so that each atom has the appropriate number on
nonbonded electrons. For elements from the second period, other
than beryllium and boron, this is the number of electrons needed so
that each atom is surrounded by an octet. For elements of the third
period and beyond, except aluminum, this is often the number of
electrons needed to complete on octet, although extra electrons can
also be placed around atoms if these elements when they are the
central atoms in compounds. Remember that atoms bonded to a central
atom usually obey the octet rule.
5. If there are not enough electrons to go around, change some
of the single bonds to multiple bonds. Multiple bonds can be
written between carbon, nitrogen, oxygen, sulphur, selenium and
phosphorous atoms. (Note that beryllium, boron, and aluminum do not
form multiple bonds.
Steps in using VSEPR to predict geometry1. Write the Lewis
structure of the molecule.
2. Determine the number of bonding pairs and only pairs of
electrons around the central atom.
3. Determine the ideal geometry. Then, if necessary taking into
account the presence of lone pairs, predict the actual shape of the
molecule
4. Keep in mind that lone pairs occupy large site (equatorial in
molecules derived from AB5 molecules) or, when sites are equal,
occupy sites opposite rather than next to each other.
Physical Properties
Depend on the forces between the particles
The stronger the bonding between the particles
The harder the substance
The higher the melting and boiling points (melting point also
dependant on lattice structure bonding)
Volatility (how easy the substance is converted to a gas) also
depends on the strength of these forces
Electrical conductivity dependant on free electrically charged
particles
Solubility (mixing of two substances) will only occur if the two
types of molecules in the mixture have as strong or stronger forces
than that between the particles in the two pure substances.
In metals
Hardness, volatility, melting and boiling point all depend on
number of valence electrons that individual metal contributes
Conductivity, malleability, ductility
Metals do not dissolve in other metals, but the can dissolve in
other metals to form alloys
Ionic compounds
Held together by strong electrostatic forces
Non-volatile
High melting and boiling points
Due to crystal lattice structure if one later moves a fraction,
similarly charged ions converge and then will repel, causing
substance to break. Thus ionic solids are brittle.
In solid form, ions cant move no electrical conduction
When molten or in solution, can carry an electric current
Strong forces between ions insoluble in most solvents
Very polar molecule water can bond to ions (known as hydration
of ions)
Ionic substances are more soluble in water than in non-polar
solvents
However, if forces between ions are very strong, ionic substance
will be insoluble in water.
Covalent
Giant covalent structures
All atoms of substance joined by strong covalent bonds
Very hard
Very high melting and boiling points
Insoluble in all solvents
Electrons firmly in place cannot conduct electricity
Molecular covalent
Strong covalent bonds between particles (intramolecular
forces)
Weak intermolecular forces between molecules
Weak intermolecular forces
substances are usually liquids or gases at room temp
are often quite soft (molecules of a solid)
often dissolve in non-polar solvents, such as petrol
insoluble in very polar solvents like water (water has strong H
bonds, inclusion of non-polar covalent would require breaking of
bonds)
electrons firmly held in bonds dont conduct electricity
H bonding can have large effect on properties of molecular
covalent substances.
H bonded substances have
Higher melting and boiling points than similar massed non H
bonded
Crystals being harder and more brittle
Quite soluble in water (molecule can form H bonds to water to
compensate for broken water-water bonds
In alkanals, the OH group will H bond, cut hydrocarbon chain
will disrupt H bonding in water as length of chain increases,
solubility decreases