1. Bonding Special Topics

2. Metallic Bonding Model must account for metallic properties:  Malleability  Ductility  Conduction of heat and electricity in all directions  High melting points Metallic bonds are strong and nondirectional.

3. Metallic Bonding Metals have 1, 2, or 3 valence electrons. “electron sea” model – a sea of delocalized electrons surrounding a positively charged metal center Valence electrons delocalized – free to move around – shared by all atoms Positive ions arranged in regular, repeating pattern, stationary - crystal

4. A - Outermost electrons wander freely through metal. Metal consists of cations held together by negatively-charged electron "glue." B - Free electrons can move rapidly in response to electric fields, that's why metals are a good conductor of electricity. C - Free electrons can transmit kinetic energy rapidly, hence metals are good conductors of heat. D - The layers of atoms in metal are hard to pull apart because of the electrons holding them together, that's why metals are tough. But individual atoms are not held to any other specific atoms, it's why atoms slip easily past one another. Thus metals are ductile.

5. Metallic luster Metal atoms contain many orbitals separated by extremely small energy differences Absorb and emit wide range of light frequencies Emission responsible for shiny appearance

6. Metallic Bonding Metallic bond strength and melting point correlate with number of valence electrons More valence electrons = more “glue” Alkali metals softest, lowest melting point Heats of Vaporization of Some Metals (kJ/mol) periodelement second Li 147 Be 297 third Na 97 Mg 128 Al 294 fourth K 77 Ca 155 Sc 333 fifth Rb 76 Sr 137 Y 365 sixth Cs 64 Ba 140 La 402

7. Alloys Metallic bonding can occur with like atoms or with different kinds of metal atoms Alloy – a substance that contains a mixture of elements and has metallic properties

8. Types of Alloys Substitutional alloy – some of host metal atoms replaced by other metal atoms of similar size Interstitial alloy – some of interstices (holes) in metal structure occupied by smaller atoms

9. Examples of alloys Bronze (copper and tin) Brass (copper and zinc) Steel (iron and carbon) Sterling silver (silver and copper) Pewter (tin, copper, bismuth, antimony) Solder (tin and antimony)

10. Covalent Network Solids Contain strong covalent bonds Can be viewed as “giant molecule” – size limited by number of atoms Two allotropes of carbon are examples of covalent network solids.

11. Allotropes of Carbon: Diamond Each carbon bonded to 4 other carbons Each atom is sp 3 hybridized and has tetrahedral geometry mp 4500°C – C-C bonds very strong solubility – insoluble in all solvents – solvent molecules can’t penetrate lattice of strong C-C bonds hardness – very hard – rigid tetrahedral arrangement of covalent bonds conductivity – no mobile electrons so no conductivity

12. Allotropes of Carbon: Graphite Each C bonded to 3 others Each C is sp 2 hybridized – trigonal planar Flat sheets of carbon stack up – vdw forces hold sheets together

13. Allotropes of Carbon: Graphite mp - 3730°C – high – strong C-C bonds solubility – insoluble in all solvents hardness – soft and has lubricative properties – weak vdw forces between layers conductivity – conducts electricity – has delocalized electrons in network of π bonds

14. Covalent Network Solids: SiO 2 Silica – empirical formula SiO 2 Quartz (some types of sand) based on network of SiO 4 tetrahedra

15. Quartz and Glass When silica heated above mp (~1600°C) and cooled rapidly, amorphous solid called glass results

16. Glass Additives Properties of glass can be varied greatly by adding different substances to the melt before cooling. Adding B 2 O 3 produces borosilicate glass – expands/contracts very little with changes in temp – cooking and lab glassware – brand name Pyrex

17. Ionic Solids and Lattice Energy Modified Coulomb’s law gives lattice energy lattice energy = k(Q 1 Q 2 )/r Lattice energy – change in energy when separated gaseous ions are packed together to form an ionic solid

18. Born-Haber Cycle A method for calculating lattice energy from thermodynamic data Consider all energy changes that must occur when elements in standard state form one mole of ionic solid

19. Lattice Energy Problem Calculate the lattice energy for LiF(s) given the following: sublimation energy for Li(s) 166 kJ/mol bond energy for F 2 (g) 154 kJ/mol first IE for Li(g)520. kJ/mol EA for F(g)-328 kJ/mol enthalpy of formation of LiF(s) -617 kJ/mol a. 182 kJ/mol b. -1129 kJ/mol c. -1052 kJ/mol d. -105 kJ/mol e. 724 kJ/mol