Bonding models

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Bonding models

• Metallic bonding model
• We will look at the properties of metals and how
  those properties can be used to develop a model
  that explains the arrangement of metal atoms in
  the solid state.
• We will also consider how the properties of
  metals can be modified.
• Lets discuss the uses of metals

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Bonding models

• Properties of metals
• Good conductors of electricity and heat,
• Malleable and ductile
• Have relatively high melting and boiling
  temperatures
• Generally have high densities
• Are lustrous
• Are often hard

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Bonding models

• Not all metals have all of those properties
• Mercury is a liquid at room temperature
• chromium is brittle and not malleable.
• Alkali metals are soft enough to be cut with a
  knife. They react vigorously with water to give
  hydrogen gas
• Answer questions 1-4, p80

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Metallic bonding model

• A satisfactory model for the structure of metals
  will explain the properties we have discussed
• Chemists believe that in a solid sample of a
  metal
• Positive ions are arranged in a closely packed
  structure described as a 3-dimensional lattice.
• The outer shell electrons are free to move
  throughout the lattice.

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Metallic bonding model

• The positive ions occupy fixed positions in the
  lattice
• The negatively charged electrons move freely in
  the lattice.
• They are called delocalised electrons because
  they belong to the whole lattice
• Electrons in the inner shells are not free to
  move. They are said to be localised.

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Metallic bonding model

• The positively charged ions are attracted to the
  negatively charged delocalised electrons.
• This electrostatic force of attraction is what
  holds the metal lattice together.
• It is called the metallic bonding

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Metallic bonding model

• Based on the metallic bonding model, try to
  explain in groups, why metals
• Have relatively high boiling temperatures
• Are good conductors of electricity
• Are malleable and ductile

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Limitations of the metallic bonding model

• This model cannot explain metals properties such
  as
• High boiling temperatures and densities
• Good electrical conductivity.
• The magnetic nature of metals such as cobalt,
  iron and nickel.
• Answer questions 5-2, p82

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Modifying metals

• Most metals need to be modified in order to
  produce the desired properties for particular
  uses
• The properties of metals can be significantly
  altered by adding small amounts of another
  substance, usually a metal or carbon
• The substances are melted together, mixed and
  allowed to cool. The resultant solid is called an
  alloy.

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Modifying metals

• Table 5.3, p86 gives some features of common
  alloys
• The way a metal has been prepared affects how it
  behaves.
• The way a metal behaves-its malleability and
  brittleness- will depend on the size of crystals
  making up metals. (Textbook, p87)
• The smaller the crystals, the harder the metal

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Modifying metals

• When a metal is heated, the rate at which it is
  cooled has a significant effect on the properties
  of the solid.
• Table 5.4, p88 gives three methods of
  heat-treating metals

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Exercise, p90

• Do the following questions on page 90 15, 16,
  17, 18, 19, 21
• Homework Questions on p60-61, Full report of
  the experiment

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Ionic compounds

• Ionic compounds are formed from a reaction
  between metals and non-metals
• Elements will either loose or gain outer shell
  electrons to become stable.
• Metal atoms loose electrons to non-metals atoms
  to form cation
• Non-metal atoms gain electrons to metal atoms to
  form anion

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Ionic compounds

• E.g. In NaCl, Sodium will loose an electron to
  become Na and Chlorine will gain that electon to
  become Cl
• Why do metal ions form form positive ions and non
  metals atoms form negative ions?
• Think of first ionisation energy and
  electronegativity

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Ionic compounds

• Properties of ionic compounds
• High melting point
• Hard and brittle
• Does not conduct electricity in a solid state
• Can only conduct electricity in a molten state

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The Ionic bonding model

• An ionic compound can be represented as a three
  dimensional lattice. See fig 6.3, p94
• The lattice is held together by the electrostatic
  forces of attraction between the cation and anion
• This force is called the ionic bonding
• Explain the properties of ionic compounds based
  on the ionic bonding model
• Exercise 6.1, p96

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Electron transfer diagram

• When a metal reacts with a non-metal, the
  electrons that the metal looses to form a stable
  ion are gained by the non metal to also form a
  stable ion.
• E.g. NaCl, MgO, MgCl2
• When a non-metal gains one or more electrons, the
  name of the non-metal ends in ide.
• Exercise 6.2, Q6, 7, 8

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Covalent bonds

• Covalent bonds form between atoms of non-metals.
• E.g. H2, O2, Cl2
• When 2 atoms of hydrogen form a H molecule, 2
  electrons, one from each atom are shared between
  the 2 atoms,
• The strong force of attraction between the shared
  electrons is called a covalent bond.

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Covalent bonds

• Lets look at small molecules made of the same
  type of atoms E.g. H2, Cl2
• Bonding electrons (covalent bond), and lone pairs
  (Non-bonding electrons)

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• Lets look at molecules with 2 atoms E.g.?
• HCl
• Molecules with more than 2 atoms E.g.?
• H2O, NH3, .

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• H2O and NH3 form covalent bonding with more than
  1 covalent bonds
• Exercise 7.2, p119 Q1,2,5,6

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Valence structure

• Electron pairs, both bonding and non bonding, in
  the outer shell of an atom will repel and assume
  a position as far away from one another as
  possible
• This is called valence shell electron pair
  repulsion model (VSPR)

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Valence structure

• When there are 4 separate electron pairs, these
  will form a tetrahedron. This way electrons are
  as far apart from each other as possible.
• E.g. H2O, HCl, CCl4

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Valence structure

• When there are 3 separate electron pairs, these
  will arrange themselves at to one another.
• E.g. BF3
• When there are 2 separate electron pairs, these
  will arrange themselves at to one another
• E.g. BeH2

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Shapes of molecules

• The shape of atoms is determined by the position
  of the bonding atoms only. The lone pairs are not
  taken into account.
• E.g. CH4 4 bonding pairs form a tetrahedron
• E.g. H2O Although we have 4 pairs of electrons
  surrounding a H2O molecule, only 2 of those are
  bonding pairs.

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Shapes of molecules

• We only include the bonding pairs in the shape of
  a water molecule. Therefore the shape of a water
  molecule is angular or V shaped.
• Draw the valence structure and shape of NH3

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Shapes of molecules

• Shape of HCl 1 bonding pair. Will assume a
  linear geometry
• 2 double bonds will assume a linear geometry
• Exercise 7.3, Q7, Q8

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Forces between molecules

• However the covalent bonds between hydrogen and
  oxygen still holds the water molecule together.
• The force between atoms in a molecule is called
  Intermolecular force

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Forces between molecules

• Polarised bonds When 2 different non metals form
  a covalent bond, one atom usually attracts the
  bonding electrons more strongly than the other
  atom.
• Usually the element with the greater
  electronegativity will attract the bonding
  electron more strongly than the other.
• E. g. HCl.

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Forces between molecules

• The bonding between covalent molecules is weak
  relative to the strength of a covalent bonding
  holding together the atoms inside the molecule.
• E.g. When you heat H2O, the bonds between the
  water molecules break when water turns into
  gases, and the molecules become free to move.

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Forces between molecules

• Table of electronegativity.

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Forces between molecules

• Lets look at the forces between the HCl molecule
• Cl has a stronger electronegativity than H.
• For that reason, it will attract the bonding
  electrons more towards itself.

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Forces between molecules

• The chlorine atom ends up with a partial negative
  charge and the hydrogen atom ends up with a
  partial positive charge.
• The partial negative and positive charges form
  what we call positive and negative poles.
• For that reason, the HCl molecule is said to be a
  dipole, or a polar molecule.

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Identifying polar molecules

• The molecule must have polar bonds.
• The molecule must not be symmetrical.
• E.g. CO2, CF4.

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Identifying polar molecules

• Homework Worksheet p64 Q1, p65 Q1, 2
• Experiment 9, 10, p16, 17 Modeling structures
  and making molecular models

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Types of forces between molecules

• Dipole-dipole attraction Lets use the HCl
  molecule to explain the type of intermolecular
  force called Dipole-dipole attraction.
• Hydrogen bonding Occurs between molecules that
  contain hydrogen bonded to F, O or N.
• Use NH3, H2O and HF to explain the intermolecular
  force called Hydrogen bonding.

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Types of forces between molecules

• The last type of intermolecular force is called
  Dispersion force They hold non covalent and
  covalent molecules.
• E.g. Forces holding a H molecule
• They are generally weak and only considered in
  the absence of other forces
• The bigger the molecule, the grater the
  dispersion forces between the molecules.

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Types of forces between molecules

• You need to remember that the intermolecular
  forces are weak compare to metallic, ionic and
  covalent bonding.
• Exercise 7.4, p127

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