Chapter 8: Covalent Bonding

Ms. Nguyen

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Warm-up Questions...

Please discuss and answer the following questions:

1) What type of bond occurs between Sodium (Na) and Chlorine (Cl)? Why do you think this?

2) Is Si and O an Ionic Bond? Why or why not?

With Your Partners!

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8.1 Molecular Compounds

In the previous chapter you learned how metals and nonmetals lose or gain electrons in an ionic bond.

This chapter will introduce another type of bonding that occurs between only nonmetals.

» Covalent Bonding: Atoms held together by sharing electrons between nonmetals.

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MoleculesIn the previous chapter you learned that a metal cation and a nonmetal anion are joined together by an ionic bond (called a salt).

A neutral group of atoms joined together by a covalent bond is called a molecule.

Examples: Water (H2O), Carbon Dioxide (CO2),

Ammonia (NH3)

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Monatomic vs. Diatomic Molecules

Molecules can be monatomic or diatomic

Monatomic Molecule: A molecule that consists of one atom.

Example 1: Noble Gases (He, Ne, Ar, Kr, Xe, Rn)

Under the right conditions (STP), they can exist in the form of a monatomic molecule because they are already stable.

Example 2: Na+ (Sodium ion)

This is a monatomic ion because there’s a positive charge after the elemental symbol, and there’s only one atom.

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Diatomic Molecule: A molecule that consists of two atoms.

There are seven diatomic molecules: H2, N2, O2, F2, Cl2,

Br2, I2 .

They exist as diatomic molecules because they cannot exist alone, to stabilize themselves they combine with another atom of the same element.

Monatomic vs. Diatomic Molecules

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When looking at the periodic table to find the 7 diatomic molecules, remember that “H” is one of them and the other six form a number seven and start at (N) Nitrogen.

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Molecular Compounds

Atoms of different elements can combine chemically to form compounds.

A compound composed of molecules is called a molecular compound.

The molecules of a given molecular compound are all the same.

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Properties of Molecular Compounds

Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds.

Many molecular compounds are gases or liquids at room temperature.

They involve nonmetals.

The bonds don’t involve ions, so they are nonconductors of electricity.

Given what you know about ionic bonds, discuss with

your partners two different properties between ionic and covalent bonds?

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Molecular FormulasMolecular formulas: The chemical formula of a molecular compound. It shows how many atoms of each element a molecule contains.

The subscripts written after the symbol indicate the number of atoms of each element in the molecule.

Examples:

H2O

C2H6

3 atoms (2 Hydrogens; 1 Oxygen)

8 atoms (2 Carbons; 6 Hydrogens)

• 2BF3 8 atoms [(1 Boron; 3 Fluorines) x 2]

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Molecular Formula Practice

How many atoms total and of each element do the following molecular compounds contain?

1. H2

2. 2CO

3. CO2

4. 3NH3

5. C2H6O

2 atoms (2 Hydrogen atoms)

4 atoms [(1 Carbon atom; 1 Oxygen atom) x 2]

3 atoms (1 Carbon atom; 2 Oxygen atoms)

12 atoms [(1 Nitrogen atom; 3 Hydrogen atoms) x 3]

9 atoms (2 Carbon atoms; 6 Hydrogen atoms; 1 Oxygen atom)

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Section 8.2:

The Nature of Covalent Bonding

Ms. Nguyen

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Warm-Up Questions With Your Partners!

Please discuss and answer the following questions:

1. How many atoms total and of each element does PCl5 (Phosphorus Pentachloride) contain?

2.How many atoms total and of each element does 2BF3 contain?

3.What are the seven diatomic molecules and how can you identify them on the periodic table?

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8.2 The Nature of Covalent Bonding

Remember that ionic compounds either gain or lose electrons in order to attain a noble gas electron configuration.

Covalent compounds form by sharing

electrons to attain a noble gas electron configuration.

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Ionic Bonding

-->

(Nonpolar) Covalent Bonding

<--

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Review of Octet Rule and Valence Electrons

Octet Rule: Atoms react by gaining or losing electrons so as to acquire the stable electron structure of a noble gas, usually eight electrons.

The octet rule still applies to covalent bonds. Atoms usually acquire a total of eight electrons by sharing electrons.

Valence Electrons: The electrons in the highest occupied energy level of an element’s atoms.

Example: Water (H2O)

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Single Bonds

Two atoms held together by sharing a pair of electrons are joined by a single bond.

Example: Hydrogen gas (H2) consists of diatomic molecules whose atoms share only one pair of electrons, forming a single covalent bond.

An electron dot structure such as H H for Hydrogen gas represents the shared pair of electrons of the covalent bond by two dots.

Hydrogen gas can also be represented by a dash H-H in a structural formula.

The molecular formula for Hydrogen gas is H2 which indicates the number of atoms in the molecule.

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Practice With Single Bonds

Example 1: F2

Solution: We know that the Fluorine atom has 7 valence electrons and it only needs 1 more to fulfill the octet rule (usually 8 electrons).

Notice that the two fluorine atoms share only one pair of valence electrons.

Important note: When drawing structural formulas for molecules, we want to draw symmetrical structures.

A lone pair is a pair of valence electrons that is not shared between atoms.

Example 2: F2

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Let’s Practice Drawing!

Please draw the structures for each of the covalently bonded atoms. Also list the amount of lone pairs that may exist for each?

1. Cl2

2.H2

3.I2

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Multiple Covalent BondsAtoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two pairs or three pairs of electrons.

Double bond: A bond that involves two shared pairs of electrons.

Example: CO2

Triple bond: A bond formed by sharing three pairs of electrons.

Example: N2

Question: Will Hydrogen ever form double or triple bonds? Why or why not?

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Let’s Review!

Please draw the electron dot structure for each of the following atoms:

B N Al

He Mg Be

Li Ar P

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Lewis (Dot) Structures- 5 Steps

Step 1: Determine the number of valence electrons in the Lewis structures

Step 2: Determine the number of bonds in the structure

# of electrons needed = # of bonds in molecule 2 electrons per bond

Step 3: Draw a simple and symmetrical Lewis structure

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Step 4: Fill the valence shells

Look at the chart from step 1 and see if the atoms have the correct number of valence electrons.

Step 5: Check your structure

A) Do we have the same number of total valence electrons?

B) Does each atom fulfill the octet rule?

Lewis (Dot) Structures- 5 Steps

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Example: H2O

Step 1: Atoms Valence e- # e- needed

H 1 1

H 1 1

O 6 2

Total 8 4

Step 2: 4 e- needed/ 2 e- per bond = 2 bonds

Step 3: Draw the symmetrical structures.

H—H—O H—O—H O—H—H

Lewis (Dot) Structures

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Step 4: Fill the valence shells

Look at the chart from step 1, do the atoms have the correct number of valence electrons? If not, remember to include lone pairs to correct the structure.

H — O — H

Step 5: Check your structure by answering the following questions:

a) Do we have the same number of total valence electrons as step 1?

2 bonds + 2 lone pairs = 8 electrons

b) Does each atom fulfill the octet rule?

Each Hydrogen has 2 valence electrons around it, and Oxygen has 8.

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Let’s Practice Drawing!

Example 1: Cl2

Example 2: CO2

Example 3: NF3

Example 4: CH4

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Polyatomic IonsPolyatomic Ions: A tightly bound group of atoms that has a positive or negative charge and behaves as one unit.

Example: NH4+ (Ammonium Ion)

The ammonium ion forms when a positively charged hydrogen ion (H+) attaches to the unshared electron pair of an ammonia molecule (NH3).

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Bond Dissociation Energies

Bond Dissociation Energy: The energy required to break the bond between two covalently bonded atoms.

A strong bond dissociation energy corresponds to a strong covalent bond.

Example: Carbon- Carbon has a strong bond dissociation so it’s not very reactive. It’s unreactive because the dissociation energy for each of these bonds is high.

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Section 8.4: Polar Bonds and Molecules

Ms. Nguyen

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Warm-Up Questions With Your Partners!

Please discuss and answer the following questions:

How many valence electrons does oxygen have?

How do you draw the electron dot

structure for methane (CH4)?

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8.4 Polar Bonds and Molecules

• There are two types of covalent bonds:

a) Nonpolar covalent bonds

b) Polar covalent bonds (or polar bonds).

• The bonding pairs of electrons in (nonpolar) covalent bonds are pulled equally.

• Example: Cl2, N2, H2

Diatomic molecules are always nonpolar

• Polar Covalent Bonds (or polar bonds): A covalent bond between atoms in which the electrons are shared unequally.

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Electronegativity

Electronegativity: The tendency of an atom in a molecule to attract electrons to itself.

The more electronegative atom attracts electrons more strongly and gains a slightly negative charge.

The less electronegative atom has a slightly positive charge.

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Electronegativity values generally increase across a period table (from left to right), and also decreases going down a group.

Nonmetals have the highest electronegativity values.

Metals have the lowest electronegativity values.

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Classification of Bonds

You can determine the type of bond artificially by calculating the difference in electronegativity between elements

Important note: Consider what you know about ionic and covalent bonds as well!

Type of Bond! ! Electronegativity Difference

Nonpolar Covalent! 0 ! 0.4

Polar Covalent!! 0.5 ! 1.9

Ionic! ! ! 2.0 ! 4.0

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Electronegative Differences

Example: HCl

Step 1: Using your chart, find the electronegative values for Hydrogen and Chlorine.

Step 2: Subtract the larger electronegative value from the smaller electronegative value.

Step 3: Use the “Electronegativity Difference Table” to understand the value difference.

Cl - H = Electronegativity Difference

! ! 3.0 - 2.1 = 0.9 indicating a polar covalent bond

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Let’s Practice!1. N and H

2. H and H

3. Ca and Cl

4. Al and Cl

5. Mg and O

6. H and F

3.0 - 2.1 = 0.9 (Polar Covalent Bond)

2.1 - 2.1 = 0 (Nonpolar Covalent Bond)

3.0 - 1.0 = 2.0 (Ionic Bond)

3.0 - 1.5 = 1.5 (Polar Covalent Bond)

3.5 - 1.2 = 2.3 (Ionic Bond)

4.0 - 2.1 = 1.9 (Ionic Bond)

Which is more polar N—H or H—F?

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Let’s Practice!Using electronegativity values, place the following bonds in order of increasing polarity:

N—N O—H Cl—As O—K

Bond Electronegativity Diff. Type of Bond N-N 3.0 - 3.0 = 0.0 Nonpolar Covalent O-H 3.5 - 2.1 = 1.4 Polar Covalent Cl-As 3.0 - 2.0 = 1.0 Polar Covalent O-K 3.5 - 0.8 = 2.7 Ionic

Least Polar --------------------------> Most Polar (ionic)

N—N Cl—As O—H O—K

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DipolesDipole: A molecule that has two poles

In looking at the previous example (HCl), we learned that there is a significant difference in electronegative values.

This means that the chlorine atom acquires a slightly negative charge (minus sign). The hydrogen atom acquires a slightly positive charge (plus sign).

When there is unequal sharing of electrons a dipole exists

It can also be represented by an arrow pointing to the more electronegative atom.

H Cl H Cl!+ !"

or

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Let’s Practice Drawing Dipoles!

Please identify the polar covalent bonds by

assigning slightly positive and slightly negative

symbols to the atoms below

Si (1.8) – Br (2.8)

Se (2.4) – F (4.0)

N (3.0) – H (2.1)

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Attractions Between Molecules

Intermolecular attractions are weaker than either ionic or covalent bonds.

These attractions are responsible for determining whether a molecular compound is a gas, liquid, or solid at a given temperature.

Van der Waals forces – consists of the two weakest attractions between molecules

dipole interactions – polar molecules attracted to one another

dispersion forces – caused by motion of electrons (weakest of all forces)

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Hydrogen Bonding

Hydrogen Bonds - attractive forces in which

a hydrogen covalently bonded to a very

electronegative atom is also weakly bonded

to an unshared electron pair of another

electronegative atom

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This other atom may be in the same

molecule or in a nearby molecule, but

always has to include hydrogen

Hydrogen Bonds have about 5% of the

strength of an average covalent bond

Hydrogen Bond is the strongest of all

intermolecular forces

Hydrogen Bonding

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Intermolecular Attractions

A few solids that consist of molecules do

not melt until the temperature reaches

1000ºC or higher called network solids

(Example: diamond, silicon carbide)

Network Solid – solids in which all of the

atoms are covalently bonded to each other

Melting a network solid would require

breaking covalent bonds throughout the

solid

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