1. Identify the elements present in each of the following compounds and the number of each element atom present.
Common salt: NaCl
Nitric acid: HNO3
Carbon dioxide: CO2
Sodium Hydroxide: NaOH
Sodium carbonate: Na2CO3
Copper sulfate: CuSO4
Sodium bicarbonate: NaHCO3
Ammonium sulfate: (NH4)2SO4
Potassium permanganate: KMnO4
Sodium thiosulfate: Na2S2O3
Calcium bicarbonate: Ca(HCO3)2
Calcium Nitrate: Ca(NO3)2

2. Ionic compounds

3. Properties of Ionic Compounds
They have high melting and boiling points
They are hard and brittle
They do not conduct electricity in the solid state
Will conduct electricity when in solution (dissolved in water) or molten (liquid state)
Will often dissolve in water to form ions

4. Ions – Charged atoms or molecules
An ion is an atom or molecule which has lost or gained one or more electrons, making it positively or negatively charged.
Anions are negatively charged ions because they have more electrons in its electron shells than it has protons in its nuclei. They gain electrons e.g. Cl atom Cl-
Cations are positively charged ions because they have more protons than electrons. Cations lose electrons e.g. sodium Na+
The number of electrons an atom gains or loses is called its valency
Cations and anions only form when cations transfer their outershell electrons to the anions

6. 11 protons
11 electrons
17 protons
17 electrons
11 protons
10 electrons
17 protons
18 electrons

7. Getting stable – Bonding
The noble gases have a full outer shell and are very unreactive or stable
Atoms are most stable when their outer shell is full like the noble gases
How do you get a full outer shell? gain or lose or share electrons when forming compounds and molecules
A covalent bond is formed when electrons are shared between atoms.
An ionic bond is formed when electrons are transferred from one atom to the other

8. Ionic Bonding
Where does the Chlorine ion get that extra electron from?
From Sodium who wants to give one away!
In ionic bonding, electrons are completely transferred from one atom to another.
It happens between a metal and a non metal Why?
Because metals form positive ions (cations) and non metals form negative ions (anions)
The oppositely charged ions are attracted to each other by electrostatic forces

9. During a reaction of sodium with chlorine, sodium loses its one valence electron to chlorine.

10. Resulting in a positively charged sodium ion and negatively charged chlorine ion.

11. Some important things to know
Where do you find the charges for ions? Look at your valency table?
Subscript is set below a word HelloYr 11
Superscript is set above the Hello Yr 11

12. Rules for writing chemical formulas
Write the symbols and charge for ions involved. Write charge as superscript.
Determine the subscripts that will produce no charge
Check that the total positive charge and negative charge equal 0
Write final chemical formula without the charges just subscripts
Mg2+ Cl-
Mg Cl-
= 0
MgCl2

13. Polyatomic Ions Na+ SO42- Al3+ SO42- 2 Na+ SO42- 2 Al3+ 3SO42-
For polyatomic ions you cannot change the formula of the ion.
For polyatomic ions if there is more than one of them you must put them in brackets and place the subscript outside
Na+ SO Al SO42-
2 Na+ SO Al3+ 3SO42-
= = 0
Na2 SO Al2(SO4)3

14. Fill in the table Iodide I- Oxide O2- Copper Cu2+ Potassium K+
Aluminium Al3+

15. Form an ionic bond between the following
Na and Cl
Cl and F
Na and Mg
Mg and Br
NH4+ OH-
NH4+ SO42-

16. Naming Ionic Compounds
1. Name the positive ion first by writing full name of metal
For metals that can have more than one charge (valency) the name of the metal is followed by the valency in capital Roman numerals in brackets e.g. Iron(II) Chloride
2. Name the anion second.
Negatively charged elements have the suffix or ending -ide Examples are oxide (O2-), sulfide (S2-), fluoride (F-), chloride (Cl-), bromide (Br-), iodide (I-), nitride (N3-), hydride (H-)

17. Naming Ionic Compounds
3. Polyatomic ions which include oxygen in the anion have the suffixes -ate or -ite.
"ate" means there is more oxygen in the anion than one ending in "ite"
Examples: sulfate (SO42-) has more oxygen than sulfite (SO32-), nitrate (NO3-) has more oxygen in the anion than nitrite (NO2-)
Other examples are carbonate (CO32-), phosphate (PO43-) and permanganate (MnO4-) Exception: OH- is named hydroxide

18. Fill in the names of the following compounds
Iodide I-
Oxide O2-
Copper Cu2+
Copper (II) Iodide
Potassium K+
Aluminium Al3+

19. The ionic bonding model
Positive metal ions (cations) and negative non metal ions (anions) are arranged in a regular arrangement, called a lattice
The 3D lattice is held together strongly by electrostatic forces
Electrostatic forces are the forces between particles that are caused by their electric charges

20. Explaining properties of ionic compounds
Why do ionic compounds have relatively high melting points?
Remember the melting point of a compound is when it goes from being a solid to a liquid
The bond between the cations and anions are very strong and hence a large amount of energy is required to separate the bonds.

21. Why are ionic compounds brittle?
When an external force distorts the crystal this causes ions of like charge (e.g. ++) to come close together (align) and the repulsion between these ions shatters the crystal.
Why are ionic compounds hard?
The bonds between the cations and anions are very strong.

22. Why do ionic compounds conduct electricity when in solution or molten?
When molten (liquid) ions are able to slide past each other. While in solution (meaning in water) ions dissociate (move out) from the lattice and can move freely so are able to conduct electricity
Why do ionic compounds not conduct electricity when solid?
When solid, the ions are not
able to move.

23. Why do ionic compounds dissolve in water?
Water molecules are able to move between ions and free them by disrupting the rigid crystal structure

24. Uses of Ionic compounds
Read the pages 103 – 104 and summarize
Yr 11 Camp Homework Due Date: Wednesday Week 3
Complete Ionic compound sheets.
Add words to the glossary at the back of your workbook (look at pg 51, 70, 90 and 107 of your Chemistry textbook at key terms).
Read Chapter 7: Covalent molecules, networks and layers (pages 110 – 131).