Lecture 34 Review I

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Lecture 34 Review I
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You are responsible for...

• Assignments 1 9
• Tests 1, 2
• Lectures 1 33Tutorials 1 - 12
• Chapters 1-9, 13 and 16
• except for sections
• 2.9, 2.10, 3.8, 3.9, 3.10, 3.13, 4.5, 4.9, 4.11,
  5.6, 6.4, 6.11, 6.13, 13.10

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Studying

• Chapter by chapter
• 1. Review your notes
• 2. Read the chapter
• 3. DO the problems
• 4. Get help if needed

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The Big Review

• Major Topics
• Atomic Structure, Periodicity
• Ionic vs. Covalent Bonds
• Molecular Structure, Shape
• Balancing Reactions
• Mole Calculations

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The Big Review

• Major Topics
• REDOX reactions
• Thermochemistry
• Thermodynamics
• Equilibrium
• Gases

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Atomic Structure
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electrons

nucleus protons and neutrons
A Z N
mass atomic neutrons no. no.
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Electronic Structure

• Planck Ehn (waves have
• characteristic energies)
• i.e. light behaves like matter
• de Broglie matter behaves like light (i.e.
  like waves)

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Electronic Structure

• Schrodinger atoms behave like waves
• Wave equation squared Probability of finding an
  electron in a region of space (An Orbital)

Quantum Numbers
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Quantum Numbers

• n - principle (size of shell),
• l - angular momentum (shape),
• ml - magnetic (orientation),
• ms - spin (of electrons),

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Important guys

• Pauli no 2 e-s in an atom have same quantum
  numbers
• Hund maximum un-pairing of electrons in
  degenerate orbitals

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Aufbau

• 1s period 1
• 2s 2p period 2...
• 3s 3p 3d
• 4s 4p 4d 4f
• 5s 5p 5d 5f
• 6s 6p 6d

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Atomic Structure, Periodicity

• The Periodic Table - a unifying concept
• - similar chemistry in groups (columns)
• - valence (outer shell) electrons
• determine chemistry
• - periodic trends across periods (rows)

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Atomic Structure, Periodicity

• The Periodic Table - a unifying concept
• - across a period - adding e-s to same
• valence shell
• - down a group - successively higher
• valence shells

nuclear charge
atomic size, Ip, Ea
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Size increases (atomic or ionic)
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Ionization Potential Increases
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Electron Affinity Increases (becomes more
negative)
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Octet Rule

• Atoms like to have 8 electrons in their valence
  shell
• Works for most elements
• Metals lose Non-metals
• electrons gain electrons

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Electronegativity

• Tendency of an atom in a bond to attract
  electrons to itself.

HIGH
LOW
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Ionic vs. Covalent Bonds

• Low electronegativity difference
• SHARING of bonding electrons (covalency)
• High difference IONIC bond

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Ionic Solids

• High electrostatic forces
• High melting T
• High lattice energy
• (varies with ionic radii, charges)
• Solubility in POLAR solvents

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Born Haber Cycles

• 1. Vaporize the metal (DHvap, gt0)
• 2. Ionize the metal (Ip, gt0)
• 3. Dissociate the non-metal (BE, gt0)
• 4. Ionize the non-metal (Ea, lt0)
• 5. Combine ions to form salt (Elattice, lt0)
• 6. Overall M(s) ½ X2(g) MX(s) (sum)

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Lewis Structures

• 1. Find total number of valence electrons
• 2. Connect atoms with bonds
• 3. Put octets around terminal atoms
• 4. Put remaining electrons as lone pairs on
  central atom
• 5. Make multiple bonds as necessary

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Formal Charge

• of valence e-s in free atom
• -
• of valence e-s in bound atom
• More likely structure has formal charges closer
  to 0
• Negative formal charge is most likely on the more
  electronegative atom

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VSEPR

• 1. Draw the Lewis structure
• 2. Count the charge clouds around the central
  atom (including lone pairs)
• 3. Figure out the geometry

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Valence Bond Theory

• Covalent bonds formed by OVERLAP of atomic
  orbitals
• e.g. HBr

Br
H
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Hybrid Orbitals

• 1. Take ground state electronic configuration
  e.g. C (1s22s22p2)

2p
2s
1s
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Hybrid Orbitals

• 2. Promote electrons until the atom has (4)
  unpaired electrons

2p
2s
1s
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Hybrid Orbitals

• 3. Hybridize

sp3 hybrids
2p
2s
1s
1s
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Count the Charge Clouds

• Clouds Hybridization
• 2 sp
• 3 sp2
• 4 sp3
• 5 sp3d
• 6 sp3d2

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e.g. ethylene (C2H2)
leftover p-orbital
sp2 hybrid
a sigma (?) bond
a pi (p) bond
a double bond!