Download Bonding and Nomenclature

1
Chapters
7,8
Adventures
of Oxygen
Clip
22
GOALS
1. Compare and contrast types of chemical bonds (i.e.
ionic, covalent).
2. Predict formulas for stable ionic compounds
(binary and tertiary) based on balance of charges.
3. Use IUPAC nomenclature for both chemical names
and formulas:
•Ionic compounds (Binary and tertiary)
•Covalent compounds (Binary and tertiary)
4. Apply concepts of the mole and Avogadro’s number
to conceptualize and calculate empirical/molecular
formulas, mass, moles and molecules relationships.
5. Identify substances based on chemical and physical properties
3
Why do Atoms Form Compounds?
• Stability.
• What makes an atom
stable?
• Full outer energy level.
• Eight.
4• A Chemical Bond holds atoms
together in a compound.
• Two basic types:
1-Ionic
2-Covalent
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Ionic Bonding
Electrically neutral
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OPPOSITS ATTRACT!
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8
Properties of Ionic Compounds
• Crystalline solids at room temperature.
• Arranged in repeating threedimensional patterns
• Have high melting points
• Can conduct electricity
when melted or dissolved
in water
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Covalent Bonding
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11
Hydrogen and Fluorine
Hydrogen and Chlorine
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Single, Double, Triple
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Single Covalent Bonds (2e-)
Structural Formula: dashes
Unshared pair
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Double and Triple Covalent Bonds
• Double bond- 2 pairs (for a total of 4)
• Triple bond- 3 pair (for a total 6)
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Unequal Sharing of Electrons
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Called Polar Molecules
• The element that has a
greater electronegativity
attract the electrons more
• So, the electronegativity
difference between two
atoms tells you what kinds of
bond is likely to form
Clip
Polar molecules
happen when
one atom has
a greater
positive charge
17 Unequal Sharing of Electrons
δ+
δ_
Called Polar Molecules
• Animation
• The shape may affect
the polarity of an
entire molecule
• Ex CO2 (2 polar bonds
cancel each other)
• The presence of a polar bond in a molecule
often makes the entire molecules polar.
(Water molecule)
• A molecule that has 2 poles is called a
dipolar molecules, or dipole.
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Properties of Covalent Molecules
• Many are gases
or liquids at
room
temperature
• Composed of
two nonmetals.
• Have low
melting and
boiling points
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• Ionic and Covalent Bonding Review Clip
Properties of Ionic and Covalent
Compounds/Molecules
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Covalent or Ionic?
1. CO2
2. H2O
3. NaCl
4. MgCl2
5. NO2
6. Li2S
7. NaF
9. BeO
10.HCl
11.NaF
12.KCl
13.H2O2
14.N2
15.Cl2
Metallic Bonds
• Valence electrons (1-3)
can be thought of as a
sea of electrons. They
are “mobile” and can
easily drift freely from
one part of the metal to
another.
• Metallic bonds consist of
the attraction of the
free-floating valence
electrons for positively
charges metal ions.
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Other Atomic Attractions
• Intermolecular attractions are weaker
than either ionic or covalent bonds.
• Van der Waals Forces
– Weak attraction consisting of dipole
interactions and dispersion forces
– Dipole interactions: when polar
molecules are attracted to another.
– Dispersion Forces: weakest of all
interactions. Caused by motion of
electrons. Occurs between nonpolar
molecules. Temporary polarity.
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Hydrogen bonding
• Found in many biological
molecules
• Important in the
properties of water.
• Attraction between
hydrogen (when bonded
to a very
electronegative
element) and another
molecule.
• About 5% the strength
of an average covalent
bond.
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Goals
revisited
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Ionic Bonding- Formula Units
 A formula unit is the lowest
whole-number ratio of the
ions in an ionic compound.
 A chemical Formula shows
the kinds and numbers of
atoms in the smallest
representative unit of a
substance.
 How do you figure out the
“Chemical Formula?”
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•Writing chemical formulas is a shorthand way
of indicating what a substance is made of.
•These formulas also let you know how many
atoms of each type are found in a molecule.
The chemical formula for water is H2O.
Carbon Dioxide is CO2.
Why does oxygen combine in different ratios,
in different compounds?
The chemical formula for table salt is NaCl.
Calcium Chloride is CaCl2.
Why does chlorine combine in different
ratios, in different compounds?
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The simplest compounds
are ones with only two
elements
These are called binary
KI, CO, H2O, NaCl
+1
+4
-4
Oxidation numbers
0
Tell you how many electrons an
atom must gain, lose or share to
become stable.
+2
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+3
-3 -2 -1
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Oxidation numbers
+1
-1
Cl
We can predict the ratio
of atoms in ionic
valence
compounds based on 1electron
K
their oxidation numbers
All compounds
are
neutral
Tells you how many
electrons an atom
must gain, lose or
share to become
stable.
KCl
7 valence
electron
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+1
-1
Br
Na
NaBr
+2
Ca
-1
Br
To make it
ZERO, you need
CaBr
1 Ca & 22Br.
Subscripts show the number of atoms of
that kind in the compound
Some elements have more than
one oxidation number
+3
-2
+2
-2
Fe
O
Fe
O
Fe2O3
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FeO
We call these elements- Multivalent Elements
1.
2.
3.
4.
5.
6.
7.
Now You Try writing Binary
Ionic formulas
K + Br
8. Ga + Br
Mg + Cl
9. Fe+2 + O
Ca + I
10.Fe+3 + O
K+O
11.Cu+2 + F
K+I
12.Cr+3 + O
Sr + Br
13.Mg + O
Na + O
14.Al + P
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Polyatomic Ions:
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Cations:
ammonium, NH4+
Anions:
nitrate, NO3-
-a tightly bound
group of covalently
bonded atoms that
has a positive or
negative charge
and behaves AS A
UNIT.
sulfate, SO42hydroxide, OHphosphate, PO43carbonate, CO32chlorate, ClO3permanganate, MnO4chromate, CrO42-
Polyatomic Ions
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-Compounds containing polyatomic ions include both ionic and
covalent bonding
Writing Formulas Examples:
Sodium and Nitrate
Magnesium and Chlorate
Ammonium and Sulfate
Try these
1.Na + SO4
2.Mg + PO4
3.Ca + CO3
4.Na + OH
5.Mg + OH
6.NH4 + OH
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7.K + PO4
8.NH4 + NO3
9.H + SO4
10.Ca + SO4
11.K + NO3
12. Na + PO4
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Naming Binary Compounds
and Molecules
• Steps:
Example:
• NaCl
– If it is Binary1. Decide if it is an ionic or covalent
bond.
– Metal- nonmetal…..
» Ionic
– Nonmetal- nonmetal….
» Covalent
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If ionic …….
2. Only 2 elements
3. Check to see if any
elements are
multivalent.
4. If all single valent,
write the name of
the positive ion
first.
5. Write the root of
the negative ion
and add –ide.
Examples:
1. NaCl
2.K2O
3.AlCl3
4.BaF2
5.KI
6.Li2O
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If ionic …….
Examples:
5. Check to see if any
elements are multivalent. 1.FeO
6. If multivalent ions,
2.Fe2O3
determine the oxidation
3.CuO
number of the element.
7. Use Roman numerals in
4.Cu2O
parentheses after the
5.PbCl4
name of the element.
8. Write the root of the
6.PbI2
negative ion and add –ide.
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If Covalent... (Molecular Formula)
2. Use Greek prefix to
Greek Prefixes
indicate how many atoms
1- monoof each element are in
2- dithe molecule
3- tri3. Add -ide to the more
4- tetraelectronegative element
5pentaExample:
6- hexa•NO
7- hepta•Nitrogen Monoxide
•PCl3
8- octa•Phosphorous trichloride
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If it contains a polyatomic ion...
2. Write the name
Examples:
of the positive
1. NaCO3
ion.
2. KNO3
3. Write the name
of the polyatomic 3. NaC2H3O2
ion.
Example:
•KOH
•Potassium Hydroxide
•CaCO3
•Calcium Carbonate
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Name the following:
1. KBr
2. HCl
3. MgO
4. CaCl2
5. H2O
6. NO2
7. CuSO4
8. CaSO4
9. NH4OH
10.CaCO3
11.Cu(ClO3)
12.Cr2O 3
2
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Drawing Lewis Structures
Step #1: Add up the number of
valence electrons that should be
included in the Lewis Structure.
(TVE)
Step #2: Calculate # of bonds.
– Determine TOE: Theoretical Octet
Electrons
– TOE- TVE from step1
– Divide by 2 ( 2 electrons for each bond)
Step #3: Draw the “skeleton
structure” with the central atoms
and the other atoms, each
connected with a single bond.
Step #4: Any “leftover” electrons
so that all elements meet octet rule
(or full outer energy).
NH3
1. 5 + 3(1) = 8 (nitrogen
has five; each hydrogen
has one)
2. . N-8, H (2 each x 3=)
6…
–
–
–
3. .
4. .
so TOE=14
14-8= 6
6/2= 3 bonds
Drawing Lewis Structures
Double, triple bonds.
• Same as last except…
• Step #4: If there are no
electrons left, move
electrons from a different
atom to form another
bond…double
• Side note: When more
than one Lewis structure
can be drawn, the
molecule or ion is said to
have resonance.
CO32-
1. CCl4
2. NF3
3. SH2
4. H2O
5. CH4
6. CO2
7. BF3
8. F2O
9. SO2
10. SO3
11.NF3
12.N2
13.NH4+ (notice the + charge)
14.NO3- (notice the - charge)
Try these…
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Molecules Have Shapes.
• VSEPR theory proposes that the geometric arrangement of terminal
atoms, or groups of atoms about a central atom in a covalent compound,
or charged ion, is determined solely by the repulsions between electron
pairs present in the valence shell of the central atom.
• The number of electron pairs around the central atom can be determined
by writing the Lewis structure for the molecule. The geometry of the
molecule depends on the number of bonding groups (pairs of electrons)
and the number of nonbonding electrons on the central atom.
Molecular Shapes
• VSEPR Theory:
(Valence electron-pair
repulsion theory)
• The repulsion between
electron pairs causes
molecular shapes to adjust so
that the valence-electron
pairs stay as far apart as
possible
• Lone pairs have more
repulsive force than do
shared electron pairs, and
thus they force the shared
pairs to squeeze more closely
together.
Linear
Tetrahedral
Trigonal Planar
Pyrimidal
Bent
Practice: Go back to Lewis
Structure Practice, and predict
shapes.
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Shapes and Polarity
• Molecules can be polar, and
when they are polar, they are
called dipoles.
• Dipoles are molecules that
have a slightly positive charge
on one end and a slightly
negative charge on the other
• Shape can help determine
polarity
• Molecules that are
symmetrical tend to be
nonpolar. Molecules that are
asymmetrical tend to be polar
Practice: Go back to
Lewis Structure
Practice, and predict
polarity.
H2O- Bent-Polar
SO2 -bent-polar
SO3 -trigonal planarnonpolar.
BF3- trigonal planarnonpolar.
Molecules in
Motion Website
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