1 Chemistry 111 Sections 11.1 – 11.8 Chemistry 111.
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Transcript of 1 Chemistry 111 Sections 11.1 – 11.8 Chemistry 111.
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Chemistry 111
Sections 11.1 – 11.8 Chemistry 111
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Outline
• Recap of Chapter 10– Quantum Mechanics Rules– Valence Electrons & Noble Gas
Configuration
• Chapter 11.1 – 11.7:– Chemical Bonding with Ionic &
Covalent compounds.
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Quantization – A Concept.
• Quantum means that a property (e.g. electron’s energy, altitude in the classroom) can only have certain values.
• Example:– going up stairs (is)/(is not) quantized– going up a ramp is (is)/(is not) quantized.– length of string is (is)/(is not) quantized– amount of flour in cookies (is)/(is not)
quantized– number of eggs (is)/(is not) quantized
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Quantum Mechanics Steps
• I have a 6-step program to review quantum mechanics rules.
• Step #1: Keep “n” from the Bohr Model
– Quantum Mechanics refines Bohr’s Model– “n” is the “Principle Quantum Number”– Matches row number on the periodic table.
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Quantum Mechanics Steps
• Step #2: Add “sublevels”– s sublevels are spherical– p sublevels are pear
shaped (2 flower petals)– d sublevels are mostly
4-petal flowers– f sublevels are mostly
8-petal flower
s p
d f
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Quantum Mechanics Steps
• Step #3: Add “orbitals”
Sublevel # of Orbitals
s 1
p 3
d 5
f 7
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Quantum Mechanics Steps
• Step #4: Electron Capacity, 2 e– per orbital
Sublevel # of Orbitals e- Capacity
s 1 12= 2
p 3 32= 6
d 5 52= 10
f 7 72= 14
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Quantum Mechanics Steps• Step #5: Energy level e– capacity.
n # Sublevels Names e– Capacity
1 1 s 12= 2
3 2 s, p 12= 2+32= 6 8
5 3 s, p, d 12= 2 32 = 6+52= 10 18
7 4 s, p, d f … = 32
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Quantum Mechanics Steps
• Step #6: Orbital Filling / Building an Atom– Recipe for an atom:
• Choose Element, add protons & neutrons• Add electrons into orbitals until # e– = #p
– Filling Order:• Start at lowest energy level• Start at lowest suborbital• Add electrons 1 at a time, remember to “use all empty
seats on the bus”
– Be capable of drawing the energy level diagram– Write the electron configuration (1s22s22p63s2…)
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Valence Electrons
• Defined as electrons in outer shell– We will mostly work with “s” & “p” electrons
• How many valence electrons in a neutral atom?– Column number on periodic table.
• I A = 1 valence e–
• II A = 2 valence e–
• VI A = 6 valence e–
• VIII A = 8 valence e–
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Noble Gas Electron Configuration
• Noble gases are special:– They have full outer shells.– Both s & p sublevels are full.– They don’t want to react / bond with other atoms
• Everyone wants to be like a noble gas:– Atoms form ions to get a full outer shell.– Atoms share electrons to get a full outer shell.
• As we learned last time, the noble gas electron configuration is: ns2np6
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Periodic Trends – Atomic Radii
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Periodic Trends – Ionization Energy
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Chapter 11: Chemical Bonds
• Chapter 11 Topics– Monatomic (1
atom) Ions– Ionic Bonds– Covalent Bonds– Covalent Bond
Polarity– Multiple
(double/triple) Bonds
– Simple Molecules– Metal Bonds
• What we’ll do:– Discuss Ions– Ionic Bond Movie
– Covalent Bonds– Examples of
Covalent Bonds– Polarity
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Ions• We use the periodic table to predict what
ions an element forms.• Noble Gas Configuration:
– Elements ionize to get a full or empty outer shell (which ever is faster)
• Column Numbers– Use the IA – VIIIA columns.
• IA (1A) = 1+ VIIA (7A) = 1-• IIA (2A) = 2+ VIA (6A) = 2-• IIIA (3A) = 3+ VA (5A) = 3-• IVA (4A) = ?
– Ignore Transition Metals for now.
• Examples: Na, Al, Se, As, Ba, Sb, I, Xe
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Movie
• Movie on ionization:2 Na (s) + Cl2 (g) 2 NaCl
(s)
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Covalent Bonds
• Recall from Previous Chemistry Class:– Ionic = Electrons Transferred,
Covalent = Electrons shared.– Lewis Diagrams Use:
• Elemental Symbols (1 for each atom in the molecule)
• Dots Represent Electrons (usually paired)• Lines Represent Bonds (2 electrons per line)
– Things bond until they get 8 valence electrons (except Hydrogen)
• Each atom take credit for all electrons in its bonds.
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Lewis DiagramEach H counts 2 e- Fluorines each
count: LP 23=6 8 e- BP 21=2
Oxygens each count: LP 22=4 8 e- BP 22=4
NH3
H H F F
O ONH
H
H
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Big, Ugly Lewis Diagram
H gets 2 e-, C, N & O each get 8 e-
CC
CC
C
C
H
O
H
H
H
NH
H
H
O
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Exceptions
• Sometimes we can’t make everything work – and the octet rule gets broken.– Not enough Electrons
These “radicals” are quite reactive / Toxic
N O N OO
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Bond Polarity
• Covalent Bonds involve Sharing Electrons but not all Sharing is equal!
• Hydrogen Fluoride is a good example.– Fluorine is grabby and pulls electrons to it.– Hydrogen isn’t as grabby and loses its
electrons• We can use “Electronegativity” to
decide.
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Electronegativity Table
H2.1
He
Li1.0
Be1.5
B2.0
C2.5
N3.0
O3.5
F4.0
Ne
Na0.9
Mg1.2
Al1.5
Si1.8
P2.1
S2.5
Cl3.0
Ar
K0.8
Ca1.0
Ga1.8
Ge1.8
As2.0
Se2.4
Br2.8
Xe
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Using Electronegativity (EN)
1. Look up the EN values for each Element:H = 2.1, F = 4.0
2. Compute the difference (make it >0)EN = 4.0 – 2.1 = 1.9
3. Rate the difference:< 0.4 Non-Polar0.4 – 0.9 Slightly Polar0.9 – 1.7 Polar> 1.7 Probably Ionic
Note: Hydrogen cannot form Ionic Bonds!
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Who cares about Polarity?
• Polar molecules (e.g. H2O) dissolve other polar molecules.
• Polarity governs the TLC lab.• Extreme Polarity holds DNA
together.• Polarity governs how proteins fold.
• Mayonnaise depends on polarity…