Bonding in Compounds

Lesson 1: Bonding in Compounds

(chapter 8-9 of text)

Definition of a chemical bond:

Purpose of a chemical bond:

Ionic / covalent
What group(s)
END?
Share or transfer VE
Why those group(s)
Form ions?
Information given by formula
Structure formed
Strength determined by
General properties
Common mistake I when comparing properties of ionic and covalent compounds
Common mistake II

Lesson 2: Ionic Bonding

Why do ionic compounds have particular properties?

How are the formulas for ionic compounds different than the formulas for covalent compounds?

What is meant by “formula unit” and “formula mass”

If ionic compounds are neutral, how come they can by dissolved by water—which is polar and is attracted to another “polar” substance—or in general, charges?

**When drawing ions ALWAYS have a

Key that shows SIZE!!!! Look at

Bottom of AgCl crystal to see what

I mean.

What is meant by Lattice energy? What is the general formula?

How do we use this information to make comparison of strength between ionic compounds?

Look at the charge and size of the ions involved in the lattice structure:

Fe2O3 and Al2O3 have similar LE, give two factors that would explain similarity.

Lesson 3: Covalent Bonding and Lewis Structure

·  Nonmetals

·  Sharing of electrons

o  Equal, nonpolar (identical or C—H bond)

o  Unequal polar (any other nonmetal—nonmetal bond)

·  Low END

·  Discrete 3D molecules---Master these—and other aspects become apparent!!

·  Formula gives atom total—use Greek prefixes

·  No ions form

·  Neutral, non-conductor

·  Liquid or gas at room temp

·  Strength of bond depends on bond length (size of atoms) bond type (single, double, or triple)

·  Experience external attractions to other molecules called IMF’s (intermolecular forces)

Lewis Structures

Purpose: give a general idea of bonding

Steps:

1.  Determine number of VE (MUST do this step on my tests!!!!)

2.  determine central atom (common sense, and H and F can NEVER be central)

3.  draw a skeletal structure

4.  make sure periphery atoms obey octet (or duet if hydrogen)

5.  count VE, if match step 1 make sure everyone happy

6.  too few—fix with multiple bond

7.  have extra? Place on central atom

practice:

level one: all obey octet rule

NH3 CH4 H2O CO2 GeF4 CH2O HCN

Diatomics:

H2

N2

O2

F2

Level 2:

Extended octet (central atom only atom that can exceed octet or be octet deficient)

Why can some atoms expand octet? Which ones cannot?

SF4

PF5

XeF2

I3-

Lesson 4: Bond Order, Formal Charge and Lewis Structures

Summary of bond types in covalent bonding

Bond length / Bond strength / Bond order / σ/π
Single
Double
triple

Explanation:

Bond length:

Bond order:

Sigma/pi bond:

What context is bond order used in exam?

NO2- CO32-

What if there is more than one valid Lewis structure?

Formal charge: VE – ∑oe (oe: owned electrons an atom owns all the lone pairs and half the bonded pairs)

Goal: have no formal charge (all equal zero), formal charge makes sense from EN

SNC-

OSCl2

CNO-

SO2

Lesson 5: Molecular Geometry

Observation: molecules arrange themselves in specific, 3D ways that follow a particular arrangements with specific bond angles

VSEPR Theory: on the central atom electron pairs will orient themselves as far away from each other as possible

Molecule / Lewis structure/BP:LP / Electron geometry / Molecular geometry / Bond angles
CO2
VE
CH2O
VE
SO2
VE
CH4
VE
NH3
VE
H2O
VE
Axial vs equatorial electron pairs (difference between, and position LP will go first and why)
PCl5
VE
TeCl4
VE
BrF3
VE
XeF2
VE
SF6
VE
IF5
VE
XeF4
VE

Lesson 6: Sigma bond formation and overlapping orbitals

We can sometimes predict the bond by looking at the electron configuration. Write the orbital notation for two hydrogen atoms

Using the orbital notation as your guide show the sigma bond formation:

Do the same with F2 and HF

Compare the sigma bond of HF to HI to explain why bond length is a such a factor in determining bond strength

Lesson 7: Hybridization and bond angle anomalies

can’t explain the actual shape seen, with the shape we SHOULD see ex: CH4

What we see:

What we SHOULD see:

Solution:

Because hybridization requires energy, will only hybridize the amount of orbitals needed to accomplish the bonding.

Let’s go back and look at the molecules in the chart, and apply hybridization to explain the shape we see

Back to sigma (σ) and pi (π) bonding. You can only hybridize a single bond, or a orbital occupied by lone pair. Multiple bonds are formed by UNHYBRIDIZED p-orbitals---so sp3 will never explain a molecule that has a multiple bond in it

Limited use of hybridization as an explanation

Bond angles aren’t always what they tell us they are---exact numbers are not important. Conceptually understanding the effect is.

Lesson 8: survival Organic

Note about organic molecules and the exam (as it pertains to what we’ve discussed so far.) Carbons form long chains and rings—looking at geometry, bond angles, and IMF to a chain.

Methane / CH4 / Hexane / C6H14
Ethane / C2H6 / Heptane / C7H16
Propane / C3H8 / Octane / C8H18
Butane / C4H10 / Nonane / C9H20
pentane / C5H12 / decane / C10H22

Ex: draw the structural formula for hexane

How organic molecules tested:

1.  given all right angles, determine the shape, bond angles around each carbon and the IMF between similar molecules

2.  Identify isomers given two or more groups of molecules

3.  Recognize functional groups and predict change in shape, bond angles, and IMF

4.  Filling in bonds, given the structural formula without bond lines drawn in (hint: carbon MUST obey octet with no lone pairs)

Lesson 9: Molecular polarity and IMF’s

Earlier mentioned that covalently bonded molecules, unlike ionically bonded compounds, have external attractions for each to other. These attractions are collectively called intermolecular forces (IMF’s) are more classically called Vander Wall forces

LDF’s:

Dipole-dipole:

Hydrogen bonding:

Properties related to IMF’s

Odd terms related to dipoles and what they mean:

Ion—dipole:

Dipole moment:

Water solubility, miscible and immiscible:

Deviation from ideal gas behavior:

Application to molecular geometry:

All about comparison of two

Lesson 10: Metallic and covalent network bonding

Explain the two ways in which metallic bonding is unique from ionic and molecular bonding

What property of metals would allow this special type of bonding?

List unique properties because of this special bonding

Alloys

Covalent Network Solids

Semiconductors and doping

Type of solid / Description of lattice structure / What happens during phase change
Electrical conductivity / Implications for mp
Molecular
Covalent-network
Ionic
Metallic
Summary of properties in metals, metalloids, nonmetals going across periodic table
Metals / Metalloids / nonmetals
IE/EA
Conductors
·  electricity
·  heat
Bonding
Nature of oxide formed

Homework for This Unit:

Molecular geometry worksheet

Year / Question # / Description
1999 / 8 (all) / Lewis structure and comparison of bond length
Lewis structure on molecular polarity
2002B / 6(a-d) / Comparison of bond length
Comparison of shape and polarity of molecule
Comparison of shapes and planar vs 3D
2008 / 5(d-f) / Lewis structure than geometry, hybridization of molecule and polarity of molecule (sub XeO3 for XeF4 in e ii)
2010 / 5(all) / a-c geometry: bond length, bond angle
d-f IMF’s and polarity
2014 / 5 (all) / Comment on student hypothesis of bonding options in halogens—good one on seeing the new direction the exam is taking on explaining things
2001 / 8(all) / Comparison of ALL bonding types—(c) is usually missed the most of any
2004 / 7(all) / Comparison of properties
Comparison of molecular shape
Comparison of solubility (based on shape/polarity)
2006 / 6 (all) / Comparison of molecules by IMF
Discussion of what is going on in a phase change/decomposition reaction
(c) iii perennially students screw this up—egad!!!
2009 / 6(c-d) / Geometry and strength of LDF and a-d **Look at note on (c)
2012 / 5 (all except b)) / Comparison of molecules and the energy required to vaporize—you do not need to understand thermo to answer!
2013 / 6(all) / Lewis and isomer! Discuss of comparison of properties based on IMF’s

·  PDF of all questions on web page!!!!

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