A Molecular approach

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Chapter 9 – Bond Length and Energy – Part 2
We can use bond energies (in kJ/mole) to estimate the enthalpy change of a reaction (Hrxn):
Hrxn = (H’s bonds broken) – (H’s bonds formed)
Using the equation above and the table of bond energies, calculate Hrxn (in kJ) for the following reactions:
(hint: you may want to draw the Lewis Structures of all reactants and products so you can see all of the bonds)
All moelcules are in the gas phase.
CBr3H + Cl2 → CBr3Cl + HCl
N2H4 + Cl2 → 2 NH2Cl
2 HCl3 → N2 + 3 Cl2
N2 + 3 H2 → 2 NH3
Chapter 9 – Bond Length and Energy – Part 1
The bond energy of a chemical bond is the energy required to break 1 mole of the bond in the gas phase.
Cl2 (g) → 2 Cl (g)
Bond Energy = H = 243 kJ
HCl (g) → H (g) + Cl (g)
Bond Energy = H = 431 kJ
1. Why do you think bond energies are always positive?
2. Which bond is stronger, Cl−Cl or H−Cl? How do you know?
Given the following information:
3. What happens to the bond length as you go from a single bond to a double bond to a triple bond?
4. Which bond do you think has the larger bond energy, a double bond or a single bond? Why?
5. Which bond do you think has the larger bond energy, a triple bond or a double bond? Why?
Chapter 9 – Bond Polarity and Electronegativity
Bond Types:
1. What do pure covalent bonds and polar covalent bonds have in common? How are they different from each
other? How are they different from an ionic bond? Do polar covalent bonds contain ions?
Polar covalent bonds occur when there is a difference in electronegativity between the two atoms of a bond.
Electronegativity is the ability of an atom to attarct electrons to itself in a chemical bond.
2. Which do you think would be more electronegative, a large atom or a small atom? Why? (hint: think about size
and metallic and nonmetallic character).
3. Where would you find the most electronegative elements on the periodic table? Where would you find the least
electronegative elements? (hint: think about atomic size)
4. What is the general trend in electronegativity as you move from left to right along a row on the periodic table?
What is the general trend in electronegativity as you move down a group on the periodic table?
Consider the bond between two different elements: X – A
The difference in electronegativity (EN) between X and A determines the type of bond that forms:
small EN = covalent bond
intermediate EN = polar covalent bond
large EN = ionic bond
5. Using the information provided above (and your knowledge of periodic trends), predict whether the following
bonds will be ionic, covalent or polar covalent:
K and I
H and F
Br and Br
Ca and N
O and S
P and Cl
Chapter 9 – Lewis Structures and Expanded Octets
H can only have 2 valence electrons (no more)
C, N, O and F can only have 8 valence electrons (no more)
Halogens never form double or triple bonds
Any central, nonmetal atom in the 3rd period or below can have MORE than 8 valence electrons. These atoms are
said to have an expanded octet.
These compounds are stable because the central atoms (As, S and Xe) all have d orbitals than can accommodate the
extra valence electrons.
1. How many valence electrons are surrounding the As atom in AsF5? How many valence electrons are surrounding
the S atom in AsF6? How many valence electrons are surrounding the Xe atom in XeF2?
2. Why can’t C, N, O and F have more than 8 valence electrons?
3. Draw correct Lewis Structures for each of the following molecules or ions (all with have a central atom with an
expanded octet):
ClF5 (Cl is central atom)
Chapter 9 – Ionic Bonding
When solid potassium reacts with gaseous chlorine, solid potassium chloride is formed.
1. Write a balanced equation for this reaction.
2. What type of reaction is it?
3. Does potassium chloride contain ionic bonds or covalent bonds? How do you know?
4. Using the Lewis Structures of potassium and chlorine, show the movement of electrons in this reaction.
Where do the electrons go?
What are the charges AND electron configurations of the resulting ions?
What are the Lewis Structures of the resulting ions?
5. Use Lewis Structures to diagram the reaction between calcium and oxygen. Again, where do the electrons go?
What are the charges, electron configurations and Lewis Structures of the resulting ions?
6. Do the same thing for the reaction between potassium and oxygen and then again for the reaction between
calcium and chlorine.
The lattice energy (E) of an ionic compound is the energy associated with forming a crystalline lattice of alternating
cations and anions from the gaseous ions:
Na+ (g) + Cl (g) → NaCl (s)
H = lattive energy (E) = -788 kJ
E  Q1Q2
The lattice energy is proportional to Coulomb’s Law:
Q1 = charge of cation
Q2 = charge of anion
r = distance between ions
7. Why do you think lattice energies are always negative values?
8. Which compound do you think has the larger (more negative)
lattice energy, LiCl or CsCl? Why?
9. Which compound do you think has the larger (more negative)
lattice energy, NaF or CaO? Why?
10. Which compound do you think has the larger (more negative)
lattice energy, KCl or CaCl2? Why?
11. Arrange the following compounds in order of increasing (more
negative) lattice energy: LiF, SrS, RbBr, and NaBr
Chapter 9 – Lewis Structures of Atoms
In a Lewis Structure we represent the valence electrons of main-group elements as dots surrounding the symbol for
the element:
e- configuration
# valence e-
Lewis Structure
The dots are placed around the element’s sylbol with a maximum of two dots per side. Dots are placed singly before
Using the information above, complete the following table:
e- configuration
# valence e-
Lewis Structure
Chapter 9 – Lewis Structures of Molecular Compounds
The following is the Lewis Structure of Water:
which can also be written as:
Using the rules for writing Lewis Structures, draw the correct structure for each of the following molecules or ions:
Neutral Molecules:
H3COH (hint: look at how the formula is written)

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