CHAPTER 6 – CHEMICAL BONDING
Chemical Bond – a link between atoms that holds them together in a compound.
Why Bonding Occurs – usually to get to a lower energy state. Most atoms drop in
energy when they form bonds with other atoms, because bonding allows them to
get to an octet.
**OCTET RULE – most atoms will gain lose, or share enough e- to form an
octet (8 e-) in the outer shell.
Exceptions: 1. H and He want 2 e- since their outer shell only holds 2 e-.
2. Some atoms form less than or more than an octet due to
unusual bonding patterns. Don’t need to know details.
**Since forming a bond causes a drop in energy, most bond formation is exothermic
(the atoms release E as they bond).
Similarly, bond breaking is usually endothermic.
I. TYPES OF BONDS
A. Ionic bond – one atom loses e- to get to an octet (producing a cation),
another atom gains e- to get to an octet (producing an anion).
The ions are then attracted to each other by their opposite
charges.
EX: A + B  A+ + B-  AB
B. Covalent Bond – results from 2 atoms sharing electrons, so both atoms can
have an octet.
Ex: A + B  AB
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Bonds are rarely completely ionic or completely covalent. They are somewhere in a
range.
The type of bond formed between any 2 atoms can be determined by subtracting
their electronegativities.
|---------------------------------------|----------------|---------|
completely
completely
ionic
covalent
Nonpolar covalent bond – the bonding e-‘s are shared equally between the two
atoms; there is a balanced distribution of charge.
A
B
Polar covalent bond – e- are shared, but somewhat unequally. One atom pulls the eharder than the other does.
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PRACTICE
Cs – F bond: ____________________________________________________
C – O bond: ____________________________________________________
P – Cl bond: ____________________________________________________
H – Br bond: ___________________________________________________
Li – O bond: ____________________________________________________
S – Se bond: ___________________________________________________
Trends:
metal
+
nonmetal = _____________________________
nonmetal
+
nonmetal = _____________________________
Quiz #1 – Types of bonds
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II. Metallic Bonding – occurs in solid or liquid metals (Cu, Hg, Ag, Au, Ni, Na, etc.)
These atoms give up e- to get to an octet, but the e- don”t go to another
atom. Instead, the e- move freely around the sample of metal.
A. “Electron-Sea Model” – the piece of metal is held together by the
attraction between the cations and the “sea” of electrons.
Ex:
B. Properties of Metals – all explained by the electron-sea model of metallic
bonding.
1. Good conductors of heat and elec. – because the e- are free to
move, transporting the charge (or the heat) to the other end of
the substance. (True in all phases).
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2. Malleable and ductile – because the e- sea operates within a “plane”
or layer. So the bonding is:
Metal
Ionic Cmpd
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3. Shiny – because the e- sea allows each e- to have a wide range of
possible energies (they aren’t locked into the energy levels of a
particular atom), so they can absorb and re-emit light of many
wavelengths.
4. Bond Strength – directly proportional to the # of e- each metal
atom gives up. (A cation with higher charge, such as +2 or +3, pulls
harder on the sea of e- than a cation with charge of +1).
Ex: Na+1
(soft, cut with a knife)
vs.
Fe+2 or Fe+3
(very hard metal)
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III. IONIC BONDING AND IONIC COMPOUNDS
Ionic bonds: metal + nonmetal
A. Formation:
1. ions form
2. they attract each other due to charge
**Determining what ion an element will form:
1. Determine the # of valence e2. Decide how many it needs to gain or lose to get to eight (octet
rule). It always chooses the path (gaining or losing) that
involves the fewest electrons.
3. Since the + and – charges are no longer balanced, the neutral atom
has become an ion (charged). If it gained e-, the final
charge will be negative. If it lost e-, the charge will be
positive.
Practice with disks, using Na, Cl as examples.
Na: has _____electrons. Has _____ valence e-.
To get to octet, could gain _____e-, or lose _____e-.
Will __________, giving a charge of _____. The ion is __________.
Cl: has _____ valence e-.
To get an octet, can gain ____ e- or lose _____ e-.
After gaining 1 e-, the ion formed is: __________.
Al: Will __________ to get to octet. The ion formed is __________.
O: Will ___________ to get to octet. The ion formed is _________.
N: Will __________ to get to octet. The ion formed is __________.
Ca: Will __________ to get to octet. The ion formed is __________.
S: Will __________ to get to octet. The ion formed is __________.
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A. Ionic Cmpds – composed of positive and negative ions drawn together by their
charges in a ratio that balances the + and – charges to give a
neutral compound.
**Most ionic cmpds are 3-D crystals.
--this is the lowest energy position
--So, the formula for an ionic compound does not represent one molecule (an
independent particle), but represents the simplest ratio of ions that
gives electrical neutrality.
Ex: “NaCl” doesn’t mean one molecule of NaCl, but shows that a ratio of one
Na+ and one Cl- gives neutrality.
Ex: What is the most likely formula for a compound of Al and Br?
Al  forms __________ion
Br  forms __________ion
To balance the charges takes _____ Al+3 and _____ Br-1 ions.
Formula: __________(**always write the cation first)
Practice: What is the most likely formula for a compound of Mg and S?
_____________________________________
A compound of I and Na?
______________________________________
A compound of Ca and N?
______________________________________
A compound of Pb+4 and O-2?
_______________________________________
Quiz #2 – Ionic and Metallic compounds
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B. Comparing Ionic and Covalent Compounds
1. Comparing Bond Strength
a. ionic bonds – between ions within the crystal. Are fairly strong
due to the attraction between oppositely charged ions.
b. covalent bonds – between atoms within a mcl. Are very strong,
because breaking them apart causes each atom to lose its octet.
c. intermolecular forces (IMF’s) – forces of attraction from one
covalent molecule to another. Are fairly weak. (ex: water
molecules hold together into a drop, but not strongly).
2. Comparing Properties
a. Solubility – Ionic cmpds are generally more soluble in water than
covalent cmpds.
b. Melting Point/Boiling Point
For a solid to melt, the particles have to be separated so they can
move freely in the container.
Is that easier for an ionic cmpd, or a covalent one?
--For an ionic cmpd, you must break the bonds between ions in
an ionic crystal; these forces are __________________.
-- For a covalent cmpd, you must break the IMF’s between
adjacent covalent mcls. These forces are ___________.
So covalent cmpds generally have lower MP and BP than ionic cmpds.
(melt and boil more easily)
c. Hardness –
1. Ionic – ions are held strongly in place by ionic bonds. Very hard.
2. Covalent – mcls held in place by weak IMF’s. Not as hard as ionic
cmpds.
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d. Conductivity – how easily electric current passes through a substance.
Electric Current – caused by the flow of charged particles through a
substance. So, for current to be conducted, you must:
(1)
have charged particles, and
(2)
they must be able to flow.
1. Ionic cmpds –
Solid: _______________________________________________
Liquid: ______________________________________________
Dissolved in H2O: ______________________________________
2. Covalent Cmpds –
Solid: _______________________________________________
Liquid: ______________________________________________
Dissolved in H2O: ______________________________________
IV. COVALENT BONDING AND MOLECULAR COMPOUNDS
(“Molecular cmpd” is the same as “covalent cmpd”)
A. Molecule (mcl) – the smallest particle of a cmpd that can exist while still having
the properties of that cmpd.
Chemical Formula – tells the relative numbers of each kind of atom found in a
certain compound.
Ex: H2O: 2H atoms, 1 O atom per mcl
C2H5OH:
Ca(OH)2:
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B. Formation of Covalent Bonds
Ex: H-H bond
Each H atom:
H
H
As they come close,
1. the e- repel each other
2. each nucleus attracts the other atom’s e(stronger force)
So, they are drawn together. The e- clouds merge.
H H
If they get too close, the nuclei start to repel each other. Finally the attractive
and repulsive forces balance each other, and the atoms settle into being a
molecule.
1. Bond length – the average distance between the nuclei of two bonded
atoms. Bond lengths generally increase when larger atoms
are involved.
2. Bond energy – the E required to break a chemical bond and form neutral
atoms.
**CLASS ACTIVITY : Gumdrop molecules
Needed per group: 5 toothpicks
Ruler
3 small gumdrops (same color)
3 large gumdrops (diff. colors)
Molecule #1: 1 toothpick with 1 small gumdrop at one end, 1 large gumdrop at other end.
Molecule #2: 2 toothpicks with 2 cm broken off one end of each; hold them together, put 1
small and 1 large gumdrop on ends.
Molecule #3: 2 toothpicks broken in half; hold all four pieces in a bundle together, put 1
small and 1 large gumdrop on ends.
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You have built HBr, HI, and HCl. Based on their positions on the P.T. and the new
info about bond length, which mcl is which? Why?
Try to break the bonds joining the atoms.
Trend: As bond length increases, bond energy _____________________.
(atoms are further apart, so they can’t hold as tightly)
C. Drawing Covalent Molecular Structures (Lewis Structures)
--use the octet rule and e- dot notation.
--e- dots can be shifted around the atom if needed to help arrange the atoms.
Ex: Structure of H2 mcl
H
H
H H
Ex: Cl2 mcl
Cl
Cl
Cl Cl
unshared
pairs
shared pair
Unshared pair – two e- that are not involved in the bond; instead they remain
exclusively with one atom.
In a Lewis Structure, the shared pair can be replaced by a dash.
H-H
or
Cl-Cl
(this is a timesaver in complex mcls)
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Steps in Drawing Lewis Structures
1. Determine the type and # of atoms in the mcl (may be given, or may need
to find based on # of bonds each wants to form).
2. Draw the e- dot symbols separately for each atom.
**Each unpaired e- dot should be thought of as an “attachment point”
for another atom.
3. Arrange the atoms next to each other to produce an octet in each one
(except H and He).
*Note: In a multi-atom mcl, the atom that is closest to having four valence
e- is placed in the center, because it needs the most bonds to get an octet.
Ex: Draw the Lewis structure for CH3I
Step 1:
Step 2: put e- dots around them.
Step 3:
C is closest to having 4 valence e-, so put it in the center.
Ex: Draw the Lewis Structure for NH3
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Ex: Draw athe L.S. of a cmpd of Si and F
Step 1: Si has _____ valence eF has _____ valence e-
Si
(4 attschment pts)
F
(1 attachment pt)
So, need:
Independent Practice: Draw the L.S.’s for:
1. PCl3
2. A cmpd of H and S
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Single bond – a covalent bond in which one pair of e- is being shared.
(all covalent bonds shown so far are single bonds)
D. Multiple Covalent Bonds
--some elements, especially C, N, and O, can share more than one pair of ewith another atom.
Double bond – when two pairs of e- are shared between 2 atoms.
Ex:
Ethene
Triple bond – when three pairs of e- are shared between 2 atoms.
Ex:
N2
Drawing Lewis Structures w/Multiple Bonds
--Remember that C, N, and O will often form multiple bonds if needed to get
an octet.
Ex: Draw the L.S. for CH2O (formaldehyde)
So, the unpaired e- join to form a double bond.
Independent Practice: Draw L.S.’s for:
1. CO2
HCN
Quiz #3 – Drawing Lewis Structures
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E. Diatomic Elements – always found in nature as a double-atom mcl; covalently
bonded. Never found as a single, free atom. There are seven diatomic elements –
know these.
H2
H-H
N2
N N
O2
O=O
F2
F-F
Cl2
Cl-Cl
Br2
Br-Br
I2
I-I
V. MOLECULAR GEOMETRY
--the properties of mcls depend not only on bonding (ionic/covalent/metallic), but
also on the shape of the mcls.
A. VSEPR Theory – used to predict the shapes of covalent mcls.
Valence
Shell
Electron
Pair
Repulsion Theory  Electron pairs (shared or unshared) will repel each
other (same charge). So, they orient themselves in
space so they will be as far apart as possible.
The five most common shapes:
1. Linear – ex:
H2
HCl
CO2
2. Tetrahedral – ex:
(H-H)
(H-Cl)
(O=C=O)
CH4
Lewis Structure:
To get the e- pairs as far apart as possible, they form a tetrahedral
shape. (Model)
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3. Triangular Pyramidal – like tetrahedral, with an unshared pair at the 4th point of
the tetrahedron.
Ex: NH3
Lewis Structure:
Shape:
4. Bent – like a tetrahedral, but with two unshared pairs at two points of
the tetrahedron.
Ex: H2O
Lewis Structure:
Shape:
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5. Triangular Planar – ex: H2CO (formaldehyde)
Lewis Structure:
Model?
To get these as far apart as possible, they form a triangle.
B. Steps for Predicting What Shape a Molecule Will be:
1. Draw the Lewis Structure
2. Count the # of atoms and the # of unshared pairs coming off the central
atom.
3. If the total is :
1 or 2: the molecule is linear.
3: the basic shape is triangular planar.*
4: the basic shape is tetrahedral.*
*If the basic shape is triangular planar or tetrahedral, check if any of the items
counted in Step #2 were unshared pairs. If so mentally cover those corners of
the tetrahedral (or triangle) and see what shape remains (triangular pyramidal or
bent).
Practice: Determine the shaped of the following molecules:
1. SF2
2. AsI3
3. HCN
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