Unit 11

COVALENT BONDING

Name: ___________________________

Period: ______ TEST: __March 19th__

Unit 12 – Covalent Bonding – Calendar

Monday

Tuesday

Wednesday

Thursday

Friday

March 2nd

Notes #1: Bond Types and Properties and

pg 3,

Notes #2: Covalent Bonding and Lewis Structures

(pg 5-7)

Hwk: Page 4, 7

3th

Timed Writing

Hwk: Finish Pages 4 & 7

4th

Notes #3: Electrical and Molecular Geometry (pg 8)

Hwk: Page 9

5th

Notes #4: Bond Polarity (pg 10)

Hwk: Page 11

6th

VSPER Lab

HWK:REVIEW Page 14 - 16

9th

10th

11th

12th

13th

16th

Notes #5: Molecule Polarity, and Intermolecular Forces (pg 1)

Hwk: Page 13

And REVIEW

Page 14 - 16

17th

Chemthink.com

Covalent Bond and Molecular Shapes

HWK:REVIEW Page 14 - 16

18th

VSPER Lab

REVIEW DUE: Page 14 - 16

19th

Covalent

Bonding Test

20th

Notes #1: Bond Types and Properties

Ionic Compounds and Properties

https://www.youtube.com/watch?v=TAVwvOUewLk&list=PLfWuBca_SOr8aVsqv27mpD7UV0aBKEOWU&index=1

List the properties of Ionic Compounds:

1.

2.

3.

4.

Draw Example:

Metallic Bonding and the Properties of Metal

https://www.youtube.com/watch?v=Oagr9xMAmfY&index=3&list=PLfWuBca_SOr8aVsqv27mpD7UV0aBKEOWU

List the properties of Ionic Compounds:

The electrons in a metallic compound are able to move throughout the entire compound. The electrons can be described as a "__________________ ____ __________________."

Draw Example:

Molecular / Covalent Compounds and Properties

https://www.youtube.com/watch?v=og8qXtEzrvA&index=4&list=PLfWuBca_SOr8aVsqv27mpD7UV0aBKEOWU

List the properties of Ionic Compounds:

1.

2.

3.

4.

.

Draw Example:

HOMEWORK: Three types of bonds - metallic, ionic and covalent - each have their own characteristics.

1. Metallic bonds are formed by pooled valence electrons of metallic atoms providing the negative charges to hold positively charged metallic ions together. This bonding structure provides for relatively high melting points and easy reshaping (bending, flattening). The delocalized electrons provide high electrical conductivity.

1. Ionic bonds are formed when metallic atoms donate valence electrons to non-metallic atoms. The resulting ions have opposite charges and attract each other into rigid lattices. This bonding structure gives high bond strength that provides brittle substances with high melting points and low conductivity. If the lattice is disrupted by being heated or dissolved in water, the ions break apart and find movement easier. Conductivity of molten or aqueous ions is much higher than that of solids.

1. Covalent bonds are formed when two non-metallic atoms approach and share valence electrons. These are the strongest of all bonds. Covalent networks form when atoms bond each to several others, making an interlocking web of atoms. Covalent networks are very hard to disrupt, giving these substances very high melting points and low conductivity in any state. Molecules form when a few covalent bonds form between a countable number of molecules, as in CO2 or H2O. While the bonds within the molecule are very strong, the molecules are so small that we commonly deal with a very large number of them. One molecule requires little energy to separate from another, so these substances have very low melting points, often below room temperature. Most liquids and gases that we are familiar with are molecular. Because molecules hold their electrons so tightly, molecules also tend to be poor conductors.

Classify the following as characteristics of a compound that contains metallic (M),

ionic (I) or covalent (C) bonding.

_____ 1. Contains a metal and a nonmetal ion

(and possibly polyatomic ions)

_____ 2. Contains 'sea' of delocalized electrons

_____3. The smallest particle is the formula unit

_____ 4. The smallest particle is the molecule

_____ 5. Particles are held together by strong

electrostatic forces

_____ 6. Particles are held together by relatively

weak intermolecular forces

_____ 7. The elements in the compound share

valence electrons

_____ 8. The elements in the compound gain and

lose valence electrons

_____ 9. Are usually soluble in water

_____ 10. Are electrolytes when dissolved in water

or molten (liquid) and conduct electricity

_____11. Have low melting and boiling points

_____ 12. Have high melting and boiling points

(2 answers)

_____ 13. Often exist as gases or vaporize easily at

room temperature

_____ 14. Are hard crystalline solids at room

temperature

_____ 15. Malleable and able to be flattened

_____ 16. Are rigid and brittle

_____ 17. An example is MgCl2

_____ 18. An example is CCl4

_____ 19. Conduct electricity in solid state

______20. Its structure looks like this:

_____ 21. Its structure looks like this:

_____22. Its structure looks like this:

Ionic Bonding -

Generally occurs between _____________________________.

Can also occur with __________________________.

Involves _________________________ electrons, followed by electrostatic attraction.

Covalent Bonding -

Generally occurs between ________________________________________________

Involves ____________ of electrons, rather than transfer.

Octet Rule - Atoms will acquire, through sharing or transfer, the electron configuration of a _________________. This happens in order for the atoms to gain stability.

· Most noble gases have ___ valence electrons.

· _____ is the exception.

· These elements do not need a full octet: _____________.

Valence Electrons – The electrons in the highest occupied energy level. There are two ways to determine the number of valence electrons.

A. Using the electron configuration: 1s22s22p63s23p2, how many valence electrons are in this element?

B. Look at the group number to determine valence electrons.

Dot Models - The number of dots is equal to the number of ______________ electrons.

Example: Phosphorus

Examples of Ionic compounds:NaClMgBr2

Lewis Structures for Molecules -

Each atom in the molecule is connected by bonds.

Covalent bonds are shared paired of electrons, and they are represented with a ____________.

Pairs that are not shared are called __________________________.

A single bond is created by _______ shared pair of electrons. *Draw it: _____

A double bond is created by _______ shared pairs of electrons. *Draw it: _____

A triple bond is created by _______ shared pairs of electrons. *Draw it: _____

Steps for Drawing Lewis Structures:

Example: CH4

1. Determine the number of electrons pairs (bonds) you must place in the structure

· Add up all the valence electrons for all atoms in compound.· Count up the total number of electrons to make all atoms happy (to fulfill the octet rule…usually 2 or 8).· Subtract the 2 totals.· Divide by 2· This tells you how many bonds (-) connect the central atom.

2. Draw a skeleton of the molecule, connecting atoms in the molecule with ________________________ (1 shared pair).

· Usually the ____ atom in the molecular formula is ____________.

· __________________ is never central – ______________________________________________

· Nature likes symmetry.

3. From the total pairs you counted in step 1, _________________ ______________. This will determine the number of unpaired electrons pairs you have left to distribute.

4. Distribute remaining pairs as _____________________around the atoms in the molecule.

· Hydrogen always only gets ________ of electrons.

· Each atom should have ________________________________ unless it is an exception to the octet rule.

· If there ___________________________ to fulfill the octet rule, you will need a multiple bond.

Practice:

Bond calculation: Lewis StructureBond calculation: Lewis Structure

H2O HCP

CO2N2

Remember: Exceptions to the octet rule:

· Atoms with less than an octet.

Day 2 Homework

Complete the tables below

Compound

Bond Calculation

Lewis Dot Structure

Drawing

Compound

Calculation

Lewis Dot Structure

Drawing

CCl4

O2

SiH4

BeF2

H2O

HCN

NH3

CH3Cl

HCl

N2O

Day #3: Electron and Molecular Geometry

Electron Geometry:

· The geometry based on the number of electron groups on the central atom.

· It does not matter if the groups are a single bond, a multiple bond, or lone pairs.

· Know the following Geometries:

Electron Geometry:

Linear

Trigonal Planar

Tetrahedral

Number of electron regions on the Central Atom:

Picture:

Molecular Shape (Geometry):

· In order to predict the specific molecular shape, we assume the valence electrons ____________ each other. The molecule adopts whichever 3D geometry ________________this repulsion.

· We call this process __________________________________________________.

· Some of the names for the Molecular Geometry are the same as the Electron Geometry

· Know the following Molecular Geometries:

Molecular Geometry

Linear

Trigonal Planar

Tetrahedral

Pyramidal

Bent

Number Of Lone Pairs On The Central Atom

Number Of Atoms Bonded To The Central Atom

Bond Angle

Picture

Day 3 Homework: Geometry and VSPER shapes

1. Explain the VSEPR theory in your own words:

2. Use the picture to the right to answer the next 3 questions:

a. How many electron groups are in the molecule?

b. What electron geometry will the molecule have?

c. What molecule geometry will the molecule have?

Complete the table below

Formula

Total Pairs

Lewis Structure Sketch

# lone pairs of electrons around central atom

# shared pairs of electrons around central atom

Electron Geometry

Molecule Geometry

HF

AlBr3

BF3

CO2

C2H4

Day #4: BOND Polarity

In ________________ bonds, shared pairs of electrons are pulled between the nuclei of atoms sharing them. Sometimes electrons are pulled equally and sometimes they are not. This has to do with electronegativity. *Recall: electronegativity is the ______________________ electrons.

1. Nonpolar covalent bonds: When electrons are shared ______________.

a. The molecule will be electrically neutral – little to no difference in electronegativity

b. Ex:

2. Polar covalent bonds: The electrons are shared ________________due to electronegativity

a. The more electronegative atom will have a _______________ attraction for the bonded electrons and will have a slightly ______________ charge.

b. The less electronegative atom will have a slightly ______________ charge.

c. Ex: HCl – Look up the electronegativities on the chart.

H: ____________, Cl: ___________

· There are two ways to show polarity: HCl has a polar bond

or

Ex: Water

Electronegativity Difference

Type of Bond

Example

0.0-0.3

0.4-1.0

1.0-2.0

>2.0

The difference in electronegativities indicates the type bond the atoms will form.

Question: What type of bond will form between:

· N and H?

· F and F?

· Ca and O?

· Br and Cl?

Day 4: Bond Polarity Homework

1-3 True/False

1) In a polar covalent bond, the more electronegative atom has a slightly positive charge.

TF

2) In general, non-metallic elements have greater electronegativities than metallic elements.

TF

3) If the electronegativity difference between two atoms is greater than 2, they will form an ionic bond.

TF

4) What is the difference between a polar and a non-polar bond? (answer in terms of electrons)

5) If an element with an electronegativity of 0.3 bonds with an element with an electronegativity of 2.9. What type of bond is between these elements?

6) State whether the following contain polar, non-polar or ionic bonds.

a)

b) KF

c) SO2

d) NO2

e) Cl2

f) Na2O

g) O2

h) P2O5

7) Draw the Lewis structure for the following compounds. Label any polar bonds and identify their shape.

a) H2Sd) CHCl3

b) CS2e) OF2

c) HBrd) CF4

Day #5: MOLECULE Polarity and Intermolecular Forces

Molecule Polarity:

If a molecule has all nonpolar bonds then the molecule is __________________.

If a molecule has a polar bond then the whole molecule is usually ______________, but not always.

To determine if a molecule is polar or nonpolar, look at the central atom.

1. If the central atom has ___ lone pairs and has ________________ of atoms attached to it, then the molecule is ________________.

2. If the central atom has __ lone pairs but ________ atoms attached to it, the molecule is _____.

3. If the central atom has lone pairs, the molecule is _______.

*In other words: If the molecular geometry is _____________, the bond polarities cancel, and the molecule is ___________.

Example: methane

Example: carbon dioxide

Example: water

In a polar molecule, one end has a positive charge and the other has a negative charge. A molecule that has poles is called a ________________________ or a ______________.

Recall: Intermolecular Forces are the ________________________________ neighboring particles.

Dispersion Forces – occurs between __________molecules, ___________________ charges

Dipole interaction – occurs when ___________ ___________ are attracted to one another

Hydrogen bonding – occurs when a __________ atom involved in an ______________________ bond is strongly attracted to an adjacent molecule. This is the __________ intermolecular force.

Extremely polar bonds:

Hydrogen

Bonding

Day 5 Homework

1) Label the following structures. Also include the bond angles for each structure.

a) Which of the above structures would be polar molecule(s) no matter what elements were bonded in it?

2) State whether the following contain polar, non-polar or ionic bonds.

a) NaF

b) SeO2

c) PO2

d) Br2

e) Li2O

f) N2

3) Draw the structural formula for the following compounds. Label any polar bonds and determine whether the molecule itself is polar or non-polar.

a) HOOH

b) CO2

c) H20

d) BF3

4) What are the attractive forces between adjacent water molecules called?

5) In your own words, explain how or why a molecule can have polar bonds, but the molecule is nonpolar.

6) Draw the Lewis structures for CF4 and CH2Cl2 and then explain which molecule is the more polar.

Study Guide-Covalent Bonding

1. For each of the following elements, determine the number of valence electrons and draw the Lewis Dot Diagram (Electron Dot Structure) for that element.

A.

B. sodium

C. helium

D. oxygen

E. neon

2. Describe each type of bond. For each type, provide at least one example.

A.

B. nonpolar covalent bond

C. polar covalent bond

D. ionic bond

3. Complete the table below.

Type of Bond

Number of Electrons

Number of Electron Pairs

Single

Double

Triple

4. What is electronegativity?

Element

Electronegativity

H

2.1

Br

2.8

C

2.5

N

3.0

O

3.5

F

4.0

Cl

3.0

I

2.5

Na

0.9

P

2.1

5. Determine if the following bonds are ionic, polar, or nonpolar using the electronegativity values in the table below.

A. H – H

B. N – H

C. C – O

D. Na – F

E. N - Br

F. C - F

6. For the polar bonds in #5, which bond was the most polar?

7. For the POLAR bonds (and only the polar) bonds in #5, assign the positive pole(+) and negative pole(+) for the bond.

8. What shape occurs if there were three atoms and one lone pair of electrons?

9. What electron geometry does BCl3 have?

10. Draw the shape of water. The two sets of lone pairs of electrons cause the atoms to repel and create what shape?

If the lone pairs of electrons were bonded to atoms, what would this shape then be called?

11. Which two molecular shapes are polar no matter what elements are involved? Explain why.

12. In the Lewis structure for CCl4, how many unpaired electrons would appear on the structure?

13. According to VSEPR theory, why do molecules adjust their shapes?

14. Complete the large-ish content frame on the back.

Covalent Bonding Content Frame

Compound

Total Pairs

Lewis Structure (Drawing)

Polar or Non-Polar Bonds?

Polar or Non-Polar Molecule?

A. H2

N-

A-

S-

U-

B-

(Label each bond with a “P” or “NP”)

B. N2

N-

A-

S-

U-

B-

(Label each bond with a “P” or “NP”)

C. H2S

N-

A-

S-

U-

B-

(Label each bond with a “P” or “NP”)

D. CO2

N-

A-

S-

U-

B-

(Label each bond with a “P” or “NP”)

E. NCl3

N-

A-

S-

U-

B-

(Label each bond with a “P” or “NP”)

F. NH3

N-

A-

S-

U-

B-

(Label each bond with a “P” or “NP”)

G. PCl3

N-

A-

S-

U-

B-

(Label each bond with a “P” or “NP”)

H. CHCl3

N-

A-

S-

U-

B-

(Label each bond with a “P” or “NP”)

I. HCN

N-

A-

S-

U-

B-

(Label each bond with a “P” or “NP”)