CHEMICAL BONDING

Overview Bonding IonicCovalentMetallic StructureGiant ionic Simple molecular Giant covalent Giant Metallic Example Sodium chloride WaterDiamondIron Bonding and structure explains the properties of a substance!

Giant Covalent Structures GraphiteDiamond

Structure of Covalent Substances Covalent substances may exist as:  simple molecular structures  giant covalent structures  Diamond  Graphite  Silicon dioxide

Diamond  One of the allotropes of carbon ‘ Allotropes’: Different (structural) forms of the same element ‘ Isotopes’: Atoms of the same element with the same number of protons but different number of neutrons  Diamond is one form of the element carbon.

Diamond  Has a giant covalent structure  Giant network of carbon atoms held together by covalent bonds in a tetrahedral arrangement These four carbon atoms form a tetrahedron.

Diamond  Each carbon is joined to four other carbon atoms by strong covalent bonds.  Each carbon has four covalent bonds.

Properties of Diamond  Very high M.P. and B.P. (Diamond melts at about 3500°C)  Hardest natural substance Reason: Carbon atoms are held together in a giant rigid structure by strong covalent bonds. A lot of energy is required to break these strong covalent bonds.

Properties of Diamond  Does not conduct electricity  All electrons are held in the covalent bonds.  No ions or free electrons to conduct electricity  Insoluble in water

Uses of Diamond  Used in cutting other hard solids (because of its hardness)  E.g. Diamond-tipped drills to cut through rock

Graphite  Another form of the element carbon  Another allotrope of carbon  Has a giant covalent structure  Consists of layers of carbon atoms (flat two-dimensional layers)

Graphite Within the layer  Each carbon atom is joined to three other carbon atoms by strong covalent bonds.  Arranged in rings of six atoms

Structure of Graphite Arrangement of carbon atoms in one layer Arrangement of layers Strong covalent bond Weak force between layers

Properties of Graphite  Very high M.P. and B.P.  (Within each layer) The bonds between the carbon atoms are difficult to break – strong covalent bonds  Soft and slippery  Weak forces between the layers  the layers can slide past each other Strong covalent bond Weak force between layers

Properties of Graphite  The only non-metal that conducts electricity Reason: Each carbon atom has one electron that is not used in bonding. Free to move  Able to conduct electricity

Uses of Graphite  Pencil lead: Made of graphite and clay Since it is soft, it flakes off and stick to paper when we write.  Lubricant (for hot machines) It does not decompose at high temperatures.

Silicon Dioxide  a.k.a. silica  Found in nature as sand or quartz  Consists of silicon and oxygen atoms  Has a giant covalent structure like diamond  SiO 2 tells us the ratio of silicon atoms to oxygen atoms is 1:2

Physical Properties of Giant Covalent Substances  Physical state  At room temperature, all substances with a giant covalent structure are solids.  Strong covalent bonds make it hard.  M.P. and B.P.  High M.P. and B.P. because of its strong covalent bonds

Physical Properties of Giant Covalent Substances  Solubility in water  Insoluble in water  Electrical conductivity  Do not conduct electricity (except graphite!) Graphite Diamond

Overview Bonding IonicCovalentMetallic StructureGiant ionic Simple molecular Giant covalent Giant Metallic Example Sodium chloride WaterDiamondIron Bonding and structure explains the properties of a substance!

Think about metals! What are some properties of metals? Shiny surfaces? Can it be bent? Strong materials? Can it stretch? Good electrical conductor? Good heat conductor?

Physical Properties of Metals AppearanceCan have shiny surfaces Physical stateUsually solid DensityUsually high M.P. & B.P.Usually high Ease of shapingMalleable and ductile Heat and electrical conductivity Good Exception??

Metals Have Giant Structures  Atoms are packed closely together in an orderly manner.  "Giant" implies that it is large.

Metallic Bonding  Each metal atom gives up its valence electrons to form positive ions.  These electrons no longer belong to any metal atom  they are delocalised.  They move freely in the space between the metal ions.

‘Sea’ of electrons  The number of delocalized electrons depends on the number of metal ions.  It does not mean an excessive number of electrons.

Metallic Bond: The Definition A metallic bond is… the (electrostatic) force of attraction between positively charged metal ions and negatively charged free or mobile(delocalized) electrons.

Physical Properties: Explanation  Solid, high density  Metal ions are packed tightly in layers.  High M.P. and B.P.  Metal ions and electrons are held together by strong metallic bonds.  Application? Exception? metal ion free electron

Physical Properties: Explanation  Malleable and ductile  When a force is applied to a metal, the layers of atoms can slide over each other easily.  The ‘sea’ of electrons holds the atoms together so that it does not break.

Physical Properties: Explanation  Good electrical conductor  Metals conduct electricity when solid and when molten.  Due to the movement of the delocalized electrons.  Application?

Physical Properties: Explanation  Good heat conductor  Due to the movement of delocalized electrons  When heated, the delocalized electrons gain more energy and move faster, colliding with neighbouring electrons. Heat is transferred in these collisions.  Application?

Ask yourself! 1. List the physical properties of metals. 2. Explain why metals can be malleable or ductile 3. Explain how a metal conduct electricity

So, what have you learnt?  The structure of a metal  a lattice of postive ions in a ‘sea of electrons’  Metallic bonding: the electrostatic force of attraction between positive metal ions and negative electrons.  Physical properties of metals and their explanation