An introduction to chemical bonds

Group 0 - The Noble gases:

Facts:

• make up 1% of the air.

• Ne used in lights for advertising.

• Ar used in filament bulbs

• Kr used in lasers

• Xe used in headlight bulbs

• He / Ne do not form any compounds

• Others form compounds with great difficulty

Electron configuration:

• These are the only elements that have the electron configuration to exist as single atoms:

• They have full outer shells and the electrons are paired with opposite spins fulfilling the 'octet rule'.

Bonding:

• All other elements on the periodic table will combine by sharing or transferring electrons in order to fill their outer shell.

• The atoms, when combined tend to have the noble gas configuration - fulfilling the 'octet rule', although there are some exceptions

Chemical bonding:

• The Periodic Table is made up of 2 types of elements.

                    Metals        +            Non - metals

• This means there are 3 types of bonding in which the atoms of the elements can combine:

Questions:  p49, 1 / p73, 12 /

Ionic Bonding

Ionic bonding results from the electrostatic attraction between oppositely charged ions.

·        Metals become positive ions (cations) electrons are removed.

·        Non-metals become negative ions (anions) electrons are gained. 

·        Electrons are transferred from metal atoms to non-metal atoms.  Both will have full outer shells.

·        In an ionic compound the ions are held in place by strong forces of attraction between oppositely charged ions (and repulsion between like charged ions).

·        Attraction outweighs repulsion and the forces are balanced.

Lithium Fluoride:       Click  for DEMO

·        Metals form positive ions which we call cations.

·        Non - metals form negative ions which we call anions.

·        Dot and cross diagrams is a means of electron counting.

Becomes

However we only show the ions and only the outer shells:-

Sodium oxide:

• Na atom loses 1e to empty its outer shell

• O atom must gain 2e to fill its outer shell

• This means that 2 Na atoms must lose 1e each = 2e to give O its 2e's needed to fill its outer shell.

Giant ionic lattices

• The dot and cross diagrams only show what happens in order to form the ions but in an ionic reaction lots of metal atoms lose electrons to non-metal atoms forming ions.

• This means that each ion will attract oppositely charged ions in all directions.

• The ions are held in place by strong forces of attraction between oppositely charged ions (and repulsion between like charged ions).

• Forces are non-directional and are therefore arranged in giant lattice structures (hence crystals):-

• Each sodium ion is surrounded by 6 chloride ions (and vice versa).

• These crystals (ionic compounds) have high melting and boiling points due to this.

• Magnesium oxide has a similar structure to sodium chloride but the electrostatic attraction in magnesium oxide is greater due to the higher charges on the ions. 

• This means that the melting and boiling point of magnesium oxide 3125K is greater than sodium chloride 1074K.

• Brittle as they cleave along a plane of ions.

Further examples of ionic bonding:

CaO	AlF3
Ca à Ca2+ + 2e-	Al à Al3+ + 3e-
O + 2e- à O2-	F + e- à F- 3 times

Questions:  p51, 1-2 /  p73, 3

Ions and the Periodic Table

Predicting charges:

• It is possible to predict the charges of elements using the Periodic Table:

• Be, B, C and Si do not usually form ions as too much energy is required.

Transitions metals:

• These elements can usually form more than one ion:

• The charge / oxidation number of the ion is written as a roman numeral:

            Iron (II)    Fe²⁺

            Copper (II)    Cu²⁺

Molecular ions (common ions)

• Groups of covalently bonded atoms can also gain and lose electrons.
• These are called molecular (or common ions)

Predicting ionic formula:

• The ions in a chemical formula must add up to zero.

• Use subscripts after an ion in a formula to double/triple that ion so the sum=0.  eg. CuCl₂

• If you are double/tripling ions that consist of more than one element (molecular ions) brackets must be used. eg.  Ca(OH)₂

• If Roman numeral numbers follow a transition metal ion in brackets, that tells you the positive charge on that transition metal ion. 

                Sodium Chloride -            Na⁺               Cl^-                Write the ions with charges

                                                            1+         :       1-                 Write the ratio of the charges

                                                                                                      Multiply up if necessary to =0

                                                                     NaCl                        Bring together omitting the charges

                Copper (II) Chloride          Cu²⁺              Cl^-              Write the ions with charges

                                                              2+       :         1-  (x2)      Write the ratio of the charges

                                                            Cu²⁺     :        Cl^-₂            Multiply up if necessary to =0

                                                                     CuCl₂                      Bring together omitting the charges

                Calcium Hydroxide           Ca²⁺              OH^-            Write the ions with charges

                                                             2+          :        1-  (x2)     Write the ratio of the charges

                                                           Ca²⁺       :      (OH^-)₂        Multiply up if necessary to =0

                                                                     Ca(OH)₂                 Bring together omitting the charges

Questions  p53, 1-3 / p73, 4

Covalent Bonding

• These occur between 2 or more non metals.  As neither will donate, they have to share.

• Attraction comes between the positive nucleus and the negative electrons that are shared between them.

• In each case the outer shell has to be filled (reach the noble gas configuration).

Octet rule (more accurately - full outer shell rule)

Single covalent bonds

• The single outer shell electrons are involved in a covalent bond.

• Each hydrogen has 1 single unpaired electron in its outer shell.

• Each hydrogen contributes 1 electron to the covalent bond = 2 electron.

• The electrons in a covalent bond are in between the 2 atoms therefore they are directional.

• Each bond is represented by a line: 

Questions: 1-2 P55

Further covalent bonds

Dative covalent bonds:

• A covalent bond is when each atom provides 1e each to be shared between them.

• A dative covalent bond is when both of the bonding electrons are provided by the same atom.

• In order for this to happen, an atom or atom in a molecule must have a lone pair of electrons:

The oxonium ion:

Water has 2 lone pairs. It will react with the hydrogen chloride gas to form the Oxonium ion

• Looking at the atoms for how this molecule is formed:

• A lone pair of electrons from the water molecule is used to forma dative covalent bond.

• The bond is indistinguishable from the other covalent bonds but we use an arrow to show that one of the bonds is dative:

• The oxonium ion is responsible for acid reactions and is represented as:

                                                                              H₃O⁺    or    H⁺_(aq)

The ammonium ion:

• The same can be done for ammonia.  This forms the ammonium ion:

Breaking the octet rule

• Until now we have used the 'octet' rule to work out how many covalent bonds an atom can make.

The problem with this rule is that some atoms do not have enough electrons in their outer shell to form enough covalent bonds to make 8e. Boron is in group 3 of the Periodic table.  This means it has 3e in its outer shell. This means that Boron can only form 3 covalent bonds = 6e in its outer shell:

Expanding the octet

Period 3, the electron shell is much larger and can accommodate more covalent bonds.

• Some atoms are able to have more than 8 electrons in its outer shell when covalently bonding.

• It very much depends on how many of their outer shell electrons are used in bonding / lone pairs:

Phosphorous	Sulphur	Chlorine
If all unpaired electrons are used then there will be 3 covalent bonds = 6e plus the lone pair of 2e = 8e	If all unpaired electrons are used then there will be 2 covalent bonds = 4e plus the 2 lone pair of 2e = 8e	If all unpaired electrons are used then there will be 1 covalent bond = 2e plus the 3 lone pair of 6e = 8e
If the lone pair unpairs then there are 5 unpaired electrons.  These will form 5 covalent bonds = 10e	If one of the lone pair unpairs then there are 4 unpaired electrons.  These will form 4 covalent bonds = 8e plus the remaining lone pair of 2e = 10e	If one of the lone pair unpairs then there are 3 unpaired electrons.  These will form 3 covalent bonds = 6e plus the 2 remaining lone pairs of 4e = 10e
	If both of the lone pairs unpair then there are 6 unpaired electrons.  These will form 6 covalent bonds = 12e	If 2 of the lone pairs unpair then there are 5 unpaired electrons.  These will form 5 covalent bonds = 10e plus the remaining lone pair of 2e = 12e
		If all of the lone pairs unpair then there are 7 unpaired electrons.  These will form 7 covalent bonds = 14e
Group 5 = 3, 5 bonds	Group 6 = 2, 4, 6 bonds	Group 6 = 1, 3, 5, 7 bonds

•  It all depends upon how many electrons are used in bonding.

All 6 electrons in sulphur have been used to form covalent bonds. This makes 12e in the outer shell of sulphur. Each of the fluorine's have 8e in their outer shell.

A better rule:

1. All electrons can be used to form covalent bonds from Period 3 onwards
2. The maximum number of bonds = the number of electrons in the outer shell

Questions:  1-2 P57

Shapes of molecules and ions:

Electron pair repulsion theory:

• The shape of a molecule is wholly determined by the number of electron pairs (or areas) around a central atom:

                            Bonding pairs & Lone pairs

• Electron pairs repel each other as far as possible.

• This repulsion determines the shape and bond angle.

Bonding pairs of electrons

Lone pair electrons:

	Tetrahedral
	Pyramidal
	Non - linear

Explanation:

• Count how many bonding pairs (or areas with double bonds) and lone pair electrons

• Pairs of electrons repel other pairs of electrons as far as possible and this determines the shape

• Lone pairs are closer to the central atom than bonding pair electrons.

• Lone pair electrons will repel other pairs of electrons more than bonding pairs of electrons.

• Each lone pair of electrons reduces the bond angle by 2.5^o.

• Order of repulsion:

               Lone Pair – Lone Pair       >       Lone Pair – Bonding Pair       >       Bonding Pair – Bonding Pair

Molecules with double bonds:

• Double bonds have 4e in one region.

• To work out the shape we have to think in terms of bonding regions instead of bonding pairs:

                                                                                                                            Linear

• CO₂ has 2 bonding regions which will repel each other as far as possible.

• This will give a linear molecule with a bond angle of 180^o.

Shapes of ions:

• Molecular ions follow the same rules as any molecule, remember the ammonium ion:

• It is basically the same shape as CH₄ - tetrahedral.

Summary:

State the number of bonding and lone pairs of electrons

Pairs of electrons repel as far as possible (lone and bonding)

This determines the shape

Lone pairs repel more than bonding pairs as closer to central atom

Each lone pair reducing the bond angle by 2.5^o as it is closer to the central atom

 How does the addition of H+ ion change shape of ammonia (water)

Questions:  1-3 P59 / 5, 9 P73