Properties of bonding

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Properties of bonding

• Mrs. Kay

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Properties of Ionic bonding

• Determined by their crystalline structures (how
  the crystals form)
• Solid at room temperature (no movement)
• High melting points strong bonds
• Very hard and brittle

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• Molten compounds conduct electricity
• Solid structure does not conduct electricity
  because of rigidity and no ionic movement for
  electricity to pass through.

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Strength

• Depends on the radii and charges on the ions
• Increasing metallic charge stronger bonds
  highest melting points
• Which would produce stronger bonds Na, Ca2, or
  Al3?
• Na lt Ca2 lt Al3

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Solubility

• Most are soluble (can dissolve into ions) in
  polar solvents (ex water, ammonia)
• These solutions do conduct electricity because of
  mobile ions (electrolytes)

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Hydration

• The process which polar solvent molecules
  interact with ions in the crystal lattice and
  cause the ionic crystal to dissolve, releasing
  ions into solution
• Water surrounds the ion (ion-dipole interactions)

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Properties of simple covalent molecules

• Covalent molecules exist as s, l, or g
• Usually soft
• Evaporate easier than ionic
• Low melting and boiling points

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• Do not conduct electricity in liquid or solid
  state
• Not soluble in polar solvents, but may be soluble
  in nonpolar solvents (CCl4 or gasoline)
• Napthalene (smells like moth balls)?

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Note

• Molecules any electrically neutral group of
  atoms that are bonded tightly together to be
  considered one particle.
• Ex Cl2, NH3, H2O
• Ionic compounds are not molecules!!!
• NaCl is not one molecule but a crystal lattice
  structure with attractive forces holding them
  together.

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Metallic Bonding
Name 4 Characteristics of a Metallic Bond.
What is a Metallic Bond?
- A metallic bond occurs in metals. A metal
consists of positive ions surrounded by a sea
of mobile electrons.

1. Good conductors of heat and electricity
2. Great strength
3. Malleable and Ductile
4. Luster

This shows what a metallic bond might look like.
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Metallic bond

• Occurs between atoms with low electronegativities
• Metal atoms pack close together in 3-D, like
  oranges in a box.
• Close-packed lattice formation

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• Many metals have an unfilled outer orbital
• In an effort to be energy stable, their outer
  electrons become delocalised amongst all atoms
• No electron belongs to one atom
• They move around throughout the piece of metal.
• Metallic bonds are not ions, but nuclei with
  moving electrons

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Physical Properties

• Conductivity
• Delocalised electrons are free to move so when a
  potential difference is applied they can carry
  the current along
• Mobile electrons also mean they can transfer heat
  well
• Their interaction with light makes them shiny
  (lustre)

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Malleability

• The electrons are attracted the nuclei and are
  moving around constantly.
• The layers of the metal atoms can easily slide
  past each other without the need to break the
  bonds in the metal
• Gold is extremely malleable that 1 gram can be
  hammered into a sheet that is only 230 atoms
  thick (70 nm)

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Melting points

• Related to the energy required to deform (MP) or
  break (BP) the metallic bond
• BP requires the cations and its electrons to
  break away from the others so BP are very high.
• The greater the amount of valence electrons, the
  stronger the metallic bond.
• Gallium can melt in your hand at 29.8 oC, but it
  boils at 2400 oC!

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Alloys

• Alloying one metal with other metal(s) or non
  metal(s) often enhances its properties
• Steel is stronger than pure iron because the
  carbon prevents the delocalised electrons to move
  so readily.
• If too much carbon is added then the metal is
  brittle.
• They are generally less malleable and ductile
• Some alloys are made by melting and mixing two or
  more metals
• Bronze copper and zinc
• Steel iron and carbon (usually)

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Network Covalent MoleculesAllotropes of carbon

• elements can exist in two or more different forms
  because the element's atoms are bonded together
  in a different manner
• Carbon has 3 allotrophes
• Diamond
• Graphite
• Fullerenes (C60)
• Nanotubes
• Buckminster Fullerene

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Diamonds

• carbon atoms are bonded together in a tetrahedral
  lattice arrangement (3D framework)
• Giant covalent structure
• Very strong, so they require a lot of energy to
  break them
• M.P is 3820 K
• Does NOT conduct electricity
• 4x harder than any other natural mineral

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Graphite

• has a sheet like structure where the atoms all
  lie in a plane and are only weakly bonded to the
  sheets above and below. (2D framework)
• Much softer, conducts electricity. (delocalised
  electron)
• The C-C bonds are still quite strong.
• Each carbon bonded to 3 other carbon.

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Fullerene C60

• consists of 60 carbon atoms bonded in the nearly
  spherical configuration
• C60 is highly electronegative, meaning that it
  readily forms compounds
• Low solubility, low conductivity (greater than
  diamond, but much lower than graphite)
• Buckminster Fullerene (photo) made up of hexagon
  and pentagon carbon formations
• Also includes nanotubes (cylindrical) Made up of
  hexagons of carbon