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