Unit 04: BONDING IB Topics 4 & 14 Text: Ch 8 (all except sections 4,5 & 8) Ch 9.1 & 9.5 Ch 10.1-10.7 My Name is Bond. Chemical Bond PART 5: Giant Covalent Structures, Metallic Bonding & Physical Properties
Giant Covalent Structures Allotropes of carbon: allotropes occur when an element can exist in different crystalline forms. Graphite Hybridization: sp2; bonded to 3 other C-atoms; parallel layers of hexagons
Density: 2.26 g cm-3 Conductivity: conductor (contains 1 non-bonded delocalized e- per atom)
Appearance: non-lustrous, grey solid Uses: lubricant; pencils Diamond Hybridization: sp3; bonded to 4 other C-atoms; (hardest known natural substance)
Density: 3.51 g cm-3 Conductivity: nonconductor (all es bonded; mobile es) no
Appearance: lustrous crustal Uses: jewelry; ornamentation; tools & machinery for grinding and cutting Buckminsterfullerene (C60) Look familiar???
Buckminsterfullerene (C60) Named for the architect Richard Buckminster Bucky Fuller (July 12, 1895 July 1, 1983) Buckminsterfullerene (C60)
Hybridization: sp2; bonded to 3 other C-atoms; sphere of 60 atoms (12 pentagons & 20 hexagons) Density: 1.72 g cm-3 Conductivity: semiconductor; some e- mobility; (easily accepts es to form neg. ions)
Appearance: yellow crystalline solid Uses: Reacts w/ K to make superconducting crystalline material; related forms are used to make nanotubes for the electronics industry; catalysts and lubricants Silicon Group 4 element (like C)
4 valence shell electrons In the elemental state, each silicon atom is covalently bonded to four others in a tetrahedral arrangement. This results in a giant lattice structure much like diamond. Silica
Silicon dioxide (SiO2), commonly known as silica or quartz, also forms a giant covalent structure. This is a similar tetrahedrally bonded structure, but here the bonds are between Si and O. Each Si atom is covalently bonded to four oxygen atoms, and each O to two Si atoms. SiO2
Si: geology :: C : biology Silica, quartz, glass Metallic Bonding Metallic Bonding The valence electrons in metals become detached from the individual atoms so
that metals consist of a close packed lattice of positive ions in a sea of delocalized electrons. Metallic Bonding A metallic bond is the attraction that two neighboring positive ions have for the delocalized electrons between them. Metallic Bonding
Metals are malleable --- they can be bent and reshaped under pressure. They are also ductile --- they can be drawn out into a wire. Metals are malleable and ductile
because the close-packed layers of positive ions can slide over each other without breaking more bonds than are made. Metallic Bonding Impurities added to the metal disturb the lattice and make the metal less malleable and ductile. This is why alloys are harder than the pure metals from which they are made.
Steel mostly iron & carbon (and some other elements such as Mn, Cr, V, W) TYPE OF BONDING AND PHYSICAL PROPERTIES Melting and boiling points Boiling point When a liquid turns into a
gas the attractive forces between the particles are completely broken, so boiling point is a good indication of the strength of intermolecular forces. Melting point When solids melt the crystalline structure is broken down, but there are still some attractive forces between the particles.
thermometer convection currents Thiele tube sample capillary tube Melting point
Melting points are affected by impurities. These weaken the structure and result in lower melting points. Melting points and boiling points Covalent macromolecular structures have extremely high m.pts. and b.pts. Metals and ionic compounds also tend to have relatively high b.pts. due to ionic attractions. H-bonds are in the order of 1/10 the
strength of a covalent bond. London dispersion forces are in the order of less than 1/100 of a covalent bond. Melting points and boiling points The weaker the attractive forces, the more volatile the substance. Intermolecular forces will increase with
Increasing molecular size The extent of polarity within the bonds of the structure Example: diamond m.pt. over 4000C!!! All bonds in covalently-bonded macromolecular structure
Like dissolves like Polar substances tend to dissolve in polar solvents, such as H2O. Nonpolar substances tend to dissolve in non-polar solvents, such as heptane or tetrachloromethane.
Solubility Organic molecules often have a polar head and a nonpolar carbon chain tail. For example, SOAP CH2 CH3 CH2 CH2
P CH2 CH2 CH2 O- CH2 O-
O- Soap OCH2 CH3 CH2 CH2 P
CH2 CH2 CH2 Hydrophobic non-polar end CH2 O- O-
Soap OCH2 CH3 CH2 CH2 P CH2
CH2 CH2 CH2 O- O- Hydrophilic polar end
OCH2 CH3 CH2 CH2 P CH2 CH2
CH2 CH2 _ O- O- Think of a drop of grease in water
Grease is non-polar Water is polar Soap lets you dissolve the non-polar in the polar Soap is an emulsifier
Hydrophobic ends dissolve in grease Hydrophilic ends dissolve in water Running water washes it all away. Remember Biology?
Remember something about a phospholipid bilayer making up cell walls? Solubility As the nonpolar carbon chain length increases in a homologous series, the molecules become less soluble in water. CH3CH2OH
CH3CH2CH2OH CH3CH2CH2CH2OH CH3CH2CH2CH2CH2OH Decreasing solubility CH3OH Whats a homologous series Organic compounds with a similar general formula, possessing
similar chemical properties due to the presence of the same functional group, and shows a gradation in physical properties as a result of increase in molecular size and mass. Another example of a homologous series Whats a homologous series Organic compounds with a similar general formula, possessing
similar chemical properties due to the presence of the same functional group, and shows a gradation in physical properties as a result of increase in molecular size and mass. Solubility Ethanol (C2H5OH) is a good solvent for other substances as it contains both polar and nonpolar ends. Solubility
Example: Put the following substances in order of decreasing solubility in water: methanol, butanol, propanol, ethanol. CH3OH CH3CH2CH2OH CH3CH2CH2CH2OH CH3CH2OH Decreasing solubility methanol < ethanol < propanol < butanol
Conductivity Electricity = electrons moving from atom to atom. For conductivity to occur, the substance must possess electrons or ions that are free to move. Conductivity Metals (and graphite) contain delocalized
electrons and are excellent conductors. Molten ionic salts also conduct electricity, but are chemically decomposed in the process. When all electrons are held in fixed positions, such as in diamond and in simple molecules, no electrical conductivity
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