Topic 4. Chemical bonding and structure

There are three types of strong bonds:

– Ionic

– Covalent

– Metallic

● Some substances contain both covalent and ionic bonding or an intermediate.

4.1 Ionic bonding

● Ionic bonding is an electrostatic attraction between oppositely charged ions.

● One ore more electrons are transferred from the outer shell of one atom to the outer shell of another atom.

● The charge of an ion depends on the number of electrons the atom needed to loose or gain to achieve a full outer shell.

2 Na(s) + F2 (g) → 2 NaF (s)

● The electrons are shown in pairs, because each pair of electrons occupies an orbital.

● The successive energy levels in the atoms and ions are shown getting closer together.

● The radius of a sodium atom is approximately twice that of a chlorine atom.

● The radius of a sodium ion is approximately half that of a sodium atom.

● The radius of a chlorine ion is approximately twice that of a chlorine atom.

Cations

● If an atom loses e-, it becomes a positively charged ion.

● Group 1:

● Group 2:

● Group 3

● Transition metals can form more than one ion, for example Cu+ and Cu2+, Fe2+ and Fe3+

Anions

● If an atom gains one or more e-, it becomes a negatively charged ion.

● Group 15:

● Group 16:

● Group 17:

Polyatomic ions

Ionic compounds

● Between metals (electropositive elements) and non-metals (elements with high electronegativity).

● The difference in electronegativity values needs to be greater than about 1.8.

Formulas of ionic compounds

● The overall charge of the compound must be zero.

● Ex. CaF2

Lattice

● When an ionic compound is formed, the ions are packed in an organized crystalline structure, a lattice.

● The sum of all the electrostatic attractions between the oppositely charged ions is called the lattice energy.

● The lattice energy has a high value and this energy is released when the ionic compound is formed.

● e.g. the formation of NaCl from Na(s) and Cl2(g) is an exothermic reaction.

● The value of lattice energy depends on:

● The charge of the ions

● The size of the ions

● The higher the value of lattice energy, the more stable is the ionic compound.

Physical properties

● Melting: The crystal structure is broken down, but there are still some attractive forces between the particles.

● Boiling: The attractive forces between the particles are completely broken.

● The stronger the bonds, the higher the boiling point.

● High melting and boiling points because of strong attractive forces between the ions in the lattice (mp of Na 801º C)

● Conducts electricity when molten or dissolved in water.

● When a salt dissolves, new bonds are formed between the water molecules and the ions.

● This process is called hydration and the ions are said to be hydrated.

Properties of ionic compounds

4.2 Covalent bonding

● Covalent bonding is the electrostatic attraction

between a pair of electrons and positively charged

nuclei.

Multiple covalent bonds

● Single bond: One shared electron pair with one electron from each atom.

● Double bond: Two shared electron pairs with two electrons from each atom.

● Triple bond: Three shared electron pairs with three electrons from each atom.

● The more pairs of electrons there are in a covalent bond:

- the shorter the bond length

- the stronger the bond

Polarity of molecules

● Molecules with polar bonds can be non-polar if they are symmetrical, that is if the central atom is symmetrically surrounded by identical atoms.

● In carbon dioxide the dipoles are exactly opposite in direction and cancel each other.

O = C = O

Non-polar molecules

● In a chlorine molecule, the difference in electronegativities of the atoms is 0.

● This means that the electronpair in the covalent bond is on average shared EQUALLY between the 2 chlorine atoms.

● The bond is called a non-polar bond, thus making the molecule a non-polar molecule.

Polar molecules

● In hydrochloric acid, the difference in electronegativities is 1.0.

● The more electronegative chlorine atom draws the bonding pair of electrons towards itself and becomes negatively charged.

● The hydrogen atom then becomes positively charged.

● The bond is polar and the molecule has a dipole moment.

4.3 Covalent structures

● Lewis symbols show the number of valence electrons of an element represented ass either dots or crosses.

Drawing Lewis structures of molecules

Draw the Lewis structures for:

a) O2 b) N2 c) CO2 d) HCN

Shapes of molecules and ions

● The shape of a molecule or ion can be predicted by the valence shell electron pair repulsion theory (VSEPR).

● The theory states that electron pairs (= electron domains) repel each other, and are therefore located as far away from each other as possible.

● The order of repulsion strength is:

lone pair-lone-pair > lone pair-bond pair > bond pair-bond pair

● If one or more of the negative charge centres is a non-bonding pair, this will influence the final shape of the molecule.

● e.g NH3 and H

2O

● For some molecules it is possible to write more than one correct Lewis structure.

● These structures are called resonance structures and true structure is an intermediate form known as a resonance hybrid.

Resonance structures

● Ex. All of the C-C bonds in benzene have the same bond length:

Coordinate covalent bonds

● In coordinate covalent bonds (dative covalent bonds) the shared pair of electrons comes from the same atom.

Covalent network solids

● Pure carbon has several different structural forms:

● These forms have different physical properties and they are called allotropes.

● Allotropes are crystalline forms of the same element, in which the atoms are bonded differently.

Silicon

● Tetrahedral arrangement

Silicon dioxide, SiO2 (quartz)

● Strong

● Insoluble in water

● High melting point

● Non-conductor of electricity

● A common impure form of silicon dioxide is sand, which is colored yellow by the presence of iron (III) oxide.

● Delocalized valence electrons move freely through the metal.

● The attraction between these electrons and the cations holds the piece of metal intact.

Metallic bonding

Electrical conductivity

● The delocalization electrons enables free movement in

response to electric fields.

● Tight packing of cations and delocalized electrons transmit kinetic energy rapidly.

Thermal conductivity

Malleability

● Individual atoms are not held to any other specific

atoms, hence atoms slip easily past one another.

Row 1 Row 2 Row 3 Row 40

2

4

6

8

10

12

Column 1

Column 2

Column 3

4.4 Intermolecular forces

● Intramolecular forces:

- holds the atoms together within a molecule

- affects molecular geometry and reactivity

● Intermolecular forces:

- between the molecules within a compound

- affects melting and boiling points

● Attractive forces that exist between ALL atoms and molecules.

● These forces are only temporary and very weak.

● Compounds that only have London forces have very low boiling points (they are gaseous at room temperature)

London forces (dispersion forces)

1. Number of electrons in an atom

– The more electrons, the stronger the London forces.

The more electrons, the further they are from the nucleus = less attraction → the electron cloud is more easily polarized

Factors that affect the magnitude of the London forces

2. Size of the electron cloud

- The longer the carbon chain, the larger the electron cloud → the stronger the London forces and the higher the boiling point

3. Shapes of molecules

- The more contact area for the molecules, the stronger the forces.

● Van der Waal´s forces are due to the motions of electrons, which causes temporary dipoles.

● These forces generally increase in strength as the number of electrons in a molecule increases or if the surface area between the molecules increases.

● These forces are so weak that non-polar molecules have low boiling-points (many of them are gases at room temperature).

http://www.youtube.com/watch?v=3t1Jn_jrsQk

Dipole- dipole bonding

● Between permanent dipoles

● The negative pole of one polar molecule is attracted to the positive pole of another polar molecule.

Hydrogen bonding

● In molecules where hydrogen is directly bonded to a small highly electronegative element such as oxygen, nitrogen or fluorine.

● Small molecules can have surprisingly high boiling points due to hydrogen bonds.

The lattice structure of ice

14.1 Further aspects of covalent bonding and structure (HL)

● The octet is the most common electron arrangement because of its stability.

● Exceptions:a) Fewer electrons (incomplete octet) if the central atom is a small atoms, e.g. Be and B

b) More than eight electrons (expanded octet) if the central atom is a 3rd row element or below, e.g. P and S

Species with five negative charge centres

● If a molecule has five charge centres and they all are bonding electrons, the shape is triangular bipyramidal.

● If one or more of these five negative charge centres is a non-bonding pair, this will influence the final shape of the molecule.

● One: Tetrahedron

● Two: T-shaped ClF3

● Three: Linear I3

-

Species with six negative charge centres

● Molecules with six charged centres that are all bonding have an octahedral shape, e.g. SF

6.

● One non-bonding pair: square pyramidal BrF5

● Two non-bonding pairs: square planar XeF4

Formal Charge

● Formal charges are assigned to atoms that have an “abnormal” number of bonds.

Formal charge

● Ex. For the nitrogen in ammonium:

formal charge = 5- 8/2 – 0 = +1