lect7_2007_metals

lect7_2007_metals - Properties of Metals Malleable: Can be...

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Unformatted text preview: Properties of Metals Malleable: Can be hammered into sheets Ductile: Can be drawn into wires Planes of atoms are capable of slipping with respect to each other Close Packed Structures Electrical Conductors Thermal Conductors Collective sharing of delocalized valence electrons Structure vs. Electron Count Increasing number of electrons per atom 4 valence e− [Ne] 3s23p2 5 valence e− [Ne] 3s23p3 6 valence e− [Ne] 3s23p4 Silicon 4 bonds (Black) Phosphorous 3 bonds + 1 nonbonding electron pair Sulfur 2 bonds + 2 nonbonding electron pairs 7 valence e− [Ne] 3s23p5 Chlorine 1 bonds + 3 nonbonding electron pairs Increasing coordination number 1 Close Packed Metals Atoms pack together as closely as possible. In cubic and hexagonal close packing each atom has 12 closest neighbors. Density Aluminum = 2.7 g/cm3 Density Silicon = 2.3 g/cm3 Equal bonding to all neighbors (in all directions) which makes it easy for planes of atoms to slip over one anonther → malleable, ductile Metallic Bonding There are not enough valence electrons to go around so they are collectively shared (kind of like communism but with bonding). We can think of a fixed lattice of positive metal ions surrounded by a sea of delocalized electrons. Delocalized electrons can move easily which makes metals good conductors of electricity and heat 2 Melting Point & Hardness Group 1 Group 3 4 5 6 Rb Sr Y Zr Nb Mo 39 °C Melting Point 2 777 °C 1522 °C 1854 °C 2477 °C 2623 °C 8 9 10 11 12 13 Tc Melting Point 7 Ru Rh Pd Ag Cd In 2157 °C 2334 °C 1964 °C 1554 °C 962 °C 321 °C 156 °C N2 MO Diagram σ∗2p Half filled 2p orbitals – maximum bonding π∗2p π2p 2p orbitals 2p orbitals σ2p σ∗2s 2s orbital 2s orbital σ2s 3 σ∗1s σ∗1s 1s orbital 1s orbital 1s orbital σ1s 1s orbital H-H Distance = 2 Å σ1s Splitting between bonding and antibonding orbitals decreases as orbital overlap decreases H-H Distance = 1 Å N2 MO Diagram σ∗2p Half filled 2p orbitals – maximum bonding π∗2p π2p 2p orbitals 2p orbitals σ2p σ∗2s 2s orbital 2s orbital σ2s 4 Orbital Overlap – Transition Metal s-orbitals Most Antibonding Overlap of s-orbitals s• s-orbitals are relatively large • Stong overlap • Wide Band • 1 orbital per atom – the band holds 2 electrons per atom Most Bonding Orbital Overlap – Transition Metal d-orbitals Overlap of d-orbitals d- Most Antibonding • d-orbitals are relatively small • Weak overlap Metal d-orbitals Metal d-orbitals Most Bonding • Narrow Band • 5 orbital per atom – the band holds 10 electrons per atom 5 Band Diagram – Rubidium (Rb, Tm = 39 °C) 39 Most Antibonding d-band – narrow, holds 10 electrons s-band Wide, holds 2 electrons Rb 4d-orbitals Rb 4d-orbitals Rb 5sorbital Bonding d-states are almost empty. Bonding is relatively weak. Melting point is low. Rb 5sorbital Most Bonding Band Diagram – Molybdenum (Mo Tm = 2623 °C) 2623 Most Antibonding d-band – narrow, holds 10 electrons s-band Wide, holds 2 electrons Mo 4d-orbitals Mo 4d-orbitals Mo 5sorbital Bonding d-states are almost full. Antibonding is relatively strong. Melting point is high. Mo 5sorbital Most Bonding 6 Alloys Substitutional Alloys Brass (Cu, Zn), Interstitial Alloys Steel (Fe, C) Sterling Silver (Ag, Cu), White Gold (Au, Pt) Common Oxidation States of Metals d-orbitals are filling up, but due to ineffective shielding they experience a stronger attraction to the nucleus The +2 oxidation state is common among first row transition metal elements metal because it corresponds to emptying the valence s-orbitals, i.e. orbitals, Ni: [Ar] 4s23d8 → Ni2+ [Ar] 3d8 Ar] [Ar] 7 Magnetism Diamagnetism No unpaired electrons Paragnetism Ferrognetism There are unpaired Unpaired electrons line electrons, but the up so that they all directions of the point in the same unpaired electrons are direction randomly oriented 8 ...
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