Presentation is loading. Please wait.

Presentation is loading. Please wait.

Ionic Compounds Chapter 6. Chapter Outcomes  At the end of this chapter you should be able to:  Describe the ionic bonding model  Use the model to.

Similar presentations

Presentation on theme: "Ionic Compounds Chapter 6. Chapter Outcomes  At the end of this chapter you should be able to:  Describe the ionic bonding model  Use the model to."— Presentation transcript:

1 Ionic Compounds Chapter 6

2 Chapter Outcomes  At the end of this chapter you should be able to:  Describe the ionic bonding model  Use the model to explain the properties of ionic compounds  Explain how ions are produced when metals and non-metals react  Write chemical formulas for ionic compounds  Describe the uses of some ionic compounds

3 Ionic Compounds  Ionic compounds are made up by the chemical combination of metallic and non-metallic elements.  Most rocks, minerals and gemstones are ionic compounds.  Ceramics, bricks and kitchen crockery are made from clays which contain ionic compounds.  While most of the above are made up of mixtures of different ionic compounds table salt is a pure ionic compound made up of sodium chloride (NaCl)

4 Properties of Ionic Compounds  Think of the properties of rocks, bricks, crockery and table salt. What properties do they share?  Have high melting and boiling temperatures.  Are hard but brittle  They also:  Do NOT conduct electricity in the solid state  They will only conduct electricity if they are melted or dissolved in water

5 Structure of ionic compounds  The physical properties of ionic compounds are very different from metals.  The structure of ionic compounds must therefore be very different from those present in metals.  What do we already know about ionic compounds.

6 What do the properties tell us?

7 Structure  From the properties we can conclude:  The forces between the particles are strong.  There are no free-moving electrons present, unlike in metals.  There are charged particles present, but in solid state they are not free to move.  When an ionic compound melts, however, the particles are free to move and the compound will conduct electricity.

8 The ionic bonding model Chemists believe that when metallic and non-metallic atoms react to form ionic compounds the following steps occur:  Metal atoms lose electrons to non-metallic atoms and become positively charged metal ions.  Non-metal atoms gain electrons from the metal atoms and so become negatively charged non-metal ions.  Large numbers of positive and negative ions formed in this way then combine to form a three-dimensional lattice.  The three dimensional lattice is held together strongly by electrostatic forces of attraction between positive and negative ions. This electrostatic force is called ionic bonding.

9 How many chlorine ions surround each sodium ion and vice versa?

10 Using the ionic bonding model to explain the properties of sodium chloride

11 High Melting Temperature  Ever noticed that when you eat fish and chips the food may be hot but the salt does not melt.  This is because to melt and ionic solid energy must be provided to allow the ions to break free and move.  NaCl has a high melting temp, this indicates a large amount of energy is needed to reduce the electrostatic attraction between the oppositely charged ions and allow them to move freely.

12 Hardness and Brittleness  Unlike metals ionic compounds are not malleable. They break when beaten.  A force can disrupt the strong electrostatic forces holding the lattice in place.  A sodium chloride crystal cannot be scratched easily but if a strong force (a hammer blow) is applied it will shatter.  This is because the layers of ions will move relative to each other due to the force.  During this movement, ions of like charge will become adjacent to each other. Resulting in repulsion

13 Hardness and Brittleness  Figure 6.4 The repulsion between like charges causes this sodium chloride crystal to shatter when it is hit sharply.

14 Electrical Conductivity  In the solid form, ions in sodium chloride are held in the crystal lattice and are not free to move so cannot conduct electricity.  When the solid melts the ions are free to move.  The movement of these charged particles to an electrode completes an electrical circuit.  In a similar way, when sodium chloride dissolves in water, the ions separate and are free to move towards the opposite charge.

15 Conducting Electricity

16 Reactions of metals with non-metals  Metallic atoms have low ionisation energies and low electronegativities.  Non-metallic atoms have high ionisation energies and low electronegativities.  In other words metallic atoms lose electrons easily and non-metallic atoms gain electrons easily.

17 Ionic Compounds  So the metal atoms lose an electron to the non-metal atoms.  In doing so, both atoms will often achieve the electronic configuration of the nearest noblest gas, which is particularly stable.

18 Sodium Chloride  When sodium reacts with chlorine:  Na atom (1s 2 2s 2 2p 6 3s 1 ) loses an electron to become 1s 2 2s 2 2p 6 (the same as Neon)  Cl atom (1s 2 2s 2 2p 6 3s 1 3p 5 ) gains an electron to become 1s 2 2s 2 2p 6 3s 1 3p 6 (the same as argon)

19 Electron Configuration

20 Your Turn  Page 96  Questions 2 - 5

21 Electron Transfer Diagrams  When sodium and chloride react together sodium loses an electron and chlorine gains an electron.

22 Sodium Chloride What is happening:  Chlorine molecules splitting into separate chlorine atoms  Electrons being transferred from sodium atoms to chlorine atoms – positively charged sodium and negatively charged chlorine ions are being formed.  Sodium and chloride ions combining to form a three dimensional lattice.

23 Notes:  When a non-metal atom gains one or more electrons, the name of the negative ion ends in –ide.  When a metal atom loses one or more electrons the name of the positive ion is the same as the metal and is always named first.  For example: sodium chloride

24 Electrovalency  The charge on an ion is known as its electrovalency.  That is the little positive or negative number to the top right of a chemical symbol.  Sodium has an electrovalency of +1 whilst chlorine has an electrovalency of -1  Na +1 and Cl -1

25 Magnesium Oxide  What are the electron configurations for Magnesium and Oxygen?  How many electrons does magnesium need to lose to get a full outer shell?  How many electrons does oxygen need to gain to get a full outer shell?  Draw an electron transfer diagram.  What is the electrovalency of a magnesium ion and an oxide ion?

26 Magnesium Chloride  What are the electron configurations for Mg and Cl?  So a Mg atom will have a stable outer shell if 2 electrons are removed.  A Cl atom only needs to gain one electron.  So how can this work?

27 MgCl 2

28 Your Turn  Page 100  Question 6  Question 7  Question 8

29 Chemical Formulas  Almost every compound in which a metal is combined with a non-metal displays ionic bonding.  The formulas of simple ionic compounds, such as NaCl and MgCl 2 can be predicted from the electron configurations of the atoms.

30 Electrovalencies  Elements in groups 1 all have an electrovalency of +1 (they all have only one electron to lose)  Elements in group 17 all have an electrovalency of -1  What about groups 2 and groups 16?  Does this formula work for all atoms?

31 Writing Formulas: Rules  Chemical formulas are part of the language of chemists. To understand and use this language, you need to follow a number of fules.

32 Simple Ions  The positive ion is place first in the formula, the negative ion is second.  For example, Kf, CuO  Positive and negative ions are combined so that the total number of positive charges is balanced by the total number of negative charges.  For example, CuS, CuCl 2, AlCl 3 and Al 2 O 3  When there are two or more of a particular ion in a compound, then in the chemical formula the number is written as a subscript after the chemical symbol. For example, Al 2 O 3 Writing Formulas: Rules


34 Polyatomic ions  Some ions contain more than one atom.  These are called polyatomic ions.  They include nitrate (NO 3 - ) and hydroxide (OH - ). What else?  If more than one of these ions is used to balance the charge of a compound, then it is placed in brackets with the required number written as a subscript after the brackets. For example Mg(NO 3 ) 2 and Al(OH) 3  Brackets are not required for the formula of sodium nitrate NaNO 3, where there is only one nitrate ion present for each sodium ion.

35 Different Electrovalencies  Some elements form ions with different charges.  Iron ions can have a charge of +2 or +3.  In this situation you need to specify the electrovalency when naming the compound.  This is done by placing a Roman numeral representing the electrovalency of the ion immediately after the metal in the name of the compound.  For example  Iron(II) chloride contains Fe 2+ ions and so the formula is FeCl 2  Iron(III) chloride contains Fe 3+ ions and so the forumla is FeCl 3

36 Your Turn  Page 102  Question


Download ppt "Ionic Compounds Chapter 6. Chapter Outcomes  At the end of this chapter you should be able to:  Describe the ionic bonding model  Use the model to."

Similar presentations

Ads by Google