Ionic Compounds - PowerPoint

							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 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
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)
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
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.
What do the properties tell
us?
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.
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.
How many chlorine ions surround
each sodium ion and vice versa?
Using the ionic bonding
model to explain the
properties of sodium chloride
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.
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
Hardness and Brittleness




                 Figure 6.4 The repulsion
                  between like charges causes
                  this sodium chloride crystal to
                  shatter when it is hit sharply.
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.
Conducting Electricity




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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.
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.
Sodium Chloride

  When sodium reacts with chlorine:
  Na atom (1s2 2s2 2p6 3s1) loses an
   electron to become 1s2 2s2 2p6 (the same
   as Neon)
  Cl atom (1s2 2s2 2p6 3s1 3p5) gains an
   electron to become 1s2 2s2 2p6 3s1 3p6
   (the same as argon)
Electron Configuration
Your Turn

  Page 96
  Questions 2 - 5
Electron Transfer
Diagrams
  When sodium and chloride react together
   sodium loses an electron and chlorine
   gains an electron.




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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.
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
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
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?
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?
     http://www.yenka.com/freecontent/item.action?quick=so#
MgCl2
Your Turn

    Page 100
    Question 6
    Question 7
    Question 8
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 MgCl2 can be
   predicted from the electron configurations
   of the atoms.
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?
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.
Writing Formulas: Rules
 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, CuCl2, AlCl3 and Al2O3
  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, Al2O3
Polyatomic ions
 Some ions contain more than one atom.
 These are called polyatomic ions.
 They include nitrate (NO3-) 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(NO3)2 and Al(OH)3
 Brackets are not required for the formula of
  sodium nitrate NaNO3, where there is only one
  nitrate ion present for each sodium ion.
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 Fe2+ ions and so the formula
   is FeCl2
  Iron(III) chloride contains Fe3+ ions and so the forumla
   is FeCl3
Your Turn

  Page 102
  Question 9 - 12