Reactivity of Elements
How many valence electrons do the noble gases have?
8 electrons = octet
Compounds form when two or more elements come
together, ideally so each can have a full octet
Ionic compound:
electron(s) transferred from one atom to another
Covalent compound:
electron(s) shared between two atoms
Electron interactions generate chemical bonds, which hold
together atoms in a compound
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Ions
• Metals
– often lose e-
– CATION
• Nonmetals
– often gain e– ANION
– “-ide”
Practice
Predict the ion each of the following elements
becomes:
• O
• I
• S
• Mg
• Cs
• Al
• N
Ionic Compounds
• Composed of ions
• Transfer of electrons
• Neutral (no net charge)
• Charges must balance
• Subscripts may be used
Binary ionic compounds
– Two elements
– Usually metal + nonmetal
– Solids at room temp
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Ionic Compound: Sodium Cloride
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Naming Binary Ionic Compounds
Cation
name of metal
+ Anion
+ root of nonmetal-ide
Example: Strontium and Iodine
strontium iodide
Example: Nitrogen and Cesium
cesium nitride
Naming Binary Ionic Compounds
• How would you name the following compound?
– Calcium and Fluorine
Calcium Fluoride
– Lithium and Bromine
Lithium Bromide
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Ionic Compounds
• Net ionic charge = 0
• Examples:
– Sodium chloride
– Calcium bromide
– Magnesium sulfide
– Lithium oxide
Metals that form more than one ion
• Usually in B groups (transition metals)
• Ex. Iron
Fe2+
Fe3+
FeCl2
FeCl3
Iron (II) chloride
Iron (III) chloride
chromium
copper
gold
iron
lead
tin
Cr2+
Cu+
Au+
Fe2+
Pb2+
Sn2+
Cr3+
Cu2+
Au3+
Fe3+
Pb4+
Sn4+
Polyatomic Ions
• Polyatomic ions:
• ions made up of more than one atom
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Polyatomic Ions
• Nitrate: NO3
• Sulfate: SO4
Examples:
Potassium nitrate
Calcium nitrate
Potassium sulfate
Calcium sulfate
charge = -1
charge = -2
Covalent Compounds
Atoms share electrons to form molecules
H
+
H
H H
H
H
H2
Bonding electrons
F
+
F
F F
F
F
F2
Lone pair electrons
Covalent bonds usually between nonmetal + nonmetal
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C + H
C H
C
+
H
H
C
+4 H
H
C
H
Carbon atom:
• 4 valence electrons
• Wants to form 4 bonds
H
H
C
H
H
C
H
H
N + H
N H
N
H
+
H
N
+3 H
H
N
H
H
N
H
H
Nitrogen atom:
• 5 valence electrons
• 3 unpaired electrons
• Wants to form 3 bonds
N
H
H
O + H
O
+2 H
O H
O
H
Oxygen atom:
• 6 valence electrons
• 2 unpaired electrons
• Wants to form 2 bonds
H
O
H
+
H
O
H
H
O
H
H
Multiple Bonds
N +
N
O + C + O
N
N
N
O C O
O
C
O
N
Unpaired electrons left!
O
C
O
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Resonance Structures
When two or more electron-dot formulas can be
drawn, these are called resonance structures
Example: O3
O
O
O
18 electrons total
O
O
O
O
O
O
O
O
O
Each bond is a
“one and a half” bond
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Covalent Compounds
• Combination of two different elements
(usually nonmetal + nonmetal)
Rules for Naming
-Element with lower group number is first word;
higher group number element is second word.
-Same group? element with higher period number first
-Second element: “root”-ide
-Add Greek numerical prefixes
Naming Covalent Compounds
Example: CS2
1. Which element is in lower group?
2. (Not in same group so doesn’t apply)
3. Second element is root-ide
4. Add prefix
2 sulfurs
Carbon
Disulfide
Naming Covalent Compounds
Example: N2O
1. Which element is in lower group?
2. (Not in same group so doesn’t apply)
3. Second element is root-ide
4. Add prefix
2 nitrogens
Dinitrogen oxide
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Electronegativity
Electronegativity: ability of an atom to attract electrons
(EN)
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Types of Bonding / Variations
• Ionic bond: large difference in electronegativity; polar
• Covalent bond: small difference in electronegativity
• Some difference in electronegativity: polar covalent
• No difference in electronegativity: nonpolar
• Polar bonds creates a dipole
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Look up electronegativities (Fig. 4.5)
Find the Difference between the two
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Shapes of Molecules
VSEPR: Valence Shell Electron Pair Repulsion
Minimize electron repulsion
Electron groups pushed as far apart as possible
Electron groups:
• Bonding electrons (bonds)
– Note: single, double, triple bonds all count as ONE
electron group
• Lone pairs of electrons
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Basic Molecular Shapes
2 electron groups
linear
O=C=O
3 electron groups
bent
trigonal planar
F
B
F
F
S
O
O
Basic Molecular Shapes: 4 electron groups
H
tetrahedral
C
H
H
H
bent
pyramidal
O
H
H
Polarity of Molecules
Polar bonds: differences in negativities of bonded atoms
H-H
F-F
H-F
Polar
(F more EN than H; lone pairs around F)
O
H
Non-polar (atoms equal EN)
H
Add polar bond together to see if
molecule is polar overall (head to tail)
Polar
O partially negative
H partially positive
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Polarity of Molecules
H
H
C
Add polar bonds (head to tail)
H
H
Non-polar
Molecules with polar bonds may not be polar overall!
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