Atoms and Basic Atomic Structure

 

 

Ionic bonding

An ionic bond is the electrostatic force of attraction between the positive and negative ions in a giant ionic lattice with 6:6 coordination (surrounding each positive ion are six negative ions and surrounding each negative ion are six positive ions).

 

 
 

 
 
 

 

 

 

 

 

 

 


6:6 Coordination

 

In an ionic bond electrons are transferred. HOWEVER, this isn’t the ‘bond’, the bond is the electrostatic force of attraction, not the transferred electron!

 

Ionic bonds are between metals and non-metals.

 

 

Examples:

NaCl is an ionic bond – Sodium is in group 1 and to reach the stability of the nearest noble gas it has to lose one electron. Chlorine is in group 7 and for it to reach the stability of the nearest noble gas it must gain one electron. This results in an electron being transferred and the electrostatic force of attraction that results from the positive sodium ions and the negative chlorine ions holds the ions together in a giant ionic lattice with 6:6 coordination.

 

 

Dot and cross diagrams

Ionic compound can be represented by dot and cross diagrams similar to covalent compounds, however in ionic compounds they are shown differently:

 

An example of a dot and cross diagram in a covalent compound
                            

 

 (Outer Electrons only)

 

 

An example of a dot and cross diagram in an ionic compound

 

 

 

 

Melting/Boiling points in compounds

 

Ionic compounds have high melting/boiling points because of the strong electrostatic forces of attraction between the positive and negative ions in a giant ionic lattice.

 

Covalent compounds have low melting/boiling points because although the bonds between atoms are strong the electrostatic force of attraction between molecules is weak. It is these bonds that break when the compound is heated.

 

 

 

Ionic compound in water

When ionic compounds are dissolved in water the ions are attracted to the water molecules separating them from the lattice and becoming hydrated.

 

Some ionic compounds, however won’t dissolve in water. This occurs if there is insufficient attraction to the water molecules to break the bonds. E.g. Na+  Cl- will dissolve in water, however Mg2+  O2- will not because MgO has stronger ionic bonds.

 

 

MgO also has a higher boiling point than NaCl because of the stronger ionic bonds.

 

 

 

Ionic compounds and conductivity

Ionic compounds will not conduct electricity when solid as the ions are not free to move, however when aqueous or liquid the ions are free to move and it will therefore conduct electricity.

 

Test for ionic compounds

A simple test for an ionic compound is to see if it will conduct electricity when molten or aqueous.

 

 

Test for ions

Below are some basic chemical tests for ions. These do not have to be learnt for the exam.

 

 

Test for cations

 

Cation

Test

If present colour of precipitate formed

Ca2+

Add NaOH

Blue precipitate

Fe2+

Add NaOH

Green precipitate

Fe3+

Add NaOH

Brown precipitate

Ca2+

Add excess NaOH

White precipitate forms that then doesn’t dissolve when excess NaOH is used

Pb2+

Add excess NaOH

White precipitate forms that then dissolves when excess NaOH is used

 

Test for anions

 

Anion

Test

If present colour of precipitate formed

CO32-

Add HNO3

Blue precipitate

CO32-

Add HCl

Fizzes

SO42-

Add Barium Chloride

White precipitate

Cl-

Add Silver Nitrate

White precipitate

Br-

Add Silver Nitrate

Cream precipitate

I-

Add Silver Nitrate

Yellow precipitate

 

 

How to prepare an insoluble salt

Mix the two solutions, filter the mixture, wash the residue with distilled water and then dry the residue in an oven until mass remains constant.

 

Example: Lead (II) nitrate solution with sodium iodide solution forms a precipitate of lead (II) iodide. This is filtered and the filtrate discarded is Sodium nitrate solution. The lead (II) iodide is then washed with distilled water and dried.

 

 

Soluble salts

This is where a salt is hydrated.

 

Hydrated compound

A compound that contains water of crystallisation. Water molecules form part of the crystal lattice.

                                                                                  

Equation for a hydrated salt

Na2CO3.10H2O

 

      Sodium Carbonate                     Water of crystallisation

 

Anhydrous compound

A salt that’s left after water of crystallisation is driven off the crystal lattice.

 

A precipitate

An insoluble solid formed when two solutions mix.

 

 

Isotopes

 

Definition

An atom of the same element with the same number of protons but a different number of neutrons

 

 

Example:

Chlorine has two isotopes: 35Cl and 37Cl each with different percentage abundances. 75% of all Cl atoms have a mass number of 35 and 25% have a mass number of 37. When calculating the relative atomic mass of an element the percentage abundance of its isotopes are used.

 

The Relative atomic mass of an isotope = (% x Isotope 1) + (% x Isotope 2) + (% x Isotope 3)…

                                                                                                          100

 

Therefore:        

Isotope                         35 Cl                            37 Cl

                        %                                 75                                 25

 

Ar = (75 x 35) + (25 x 37) = 35.5

                      100

 

 

 

 

These are other definitions that you must learn for the exam:

 

 

Relative atomic mass (Ar)

Average mass of an elements atoms relative to 1/12th that of a carbon-12 atom.

 

Relative isotopic mass

Mass of an isotope relative to 1/12th of a carbon-12 atom.

 

Relative molecular mass (Mr)

Mass of its molecules relative to 1/12th of a carbon-12 atom.

 

Relative formula mass

Total mass of atoms in a formula relative to 1/12th of a Carbon-12 atom

E.g. From periodic table – H2O = (2x1) + 16 = 18