Chemical bonds, Lewis Symbols and The Octet Rule

Objectives:

• Describe the different types of chemical bonds
• Evaluate electron configurations
• Draw Lewis dot diagrams
• Apply the octet rule while being aware of its limitations.

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Chemical Bonds

Definition: When atoms are strongly attached to one another

Metallic

• Metals bonding to metals
• ‘Communal’ electrons that travel throughout the whole solid
• Highly conductive and have luster.

Ionic

• Metals bonding to Nonmetals
• Electrons are transferred
• Opposite charges attract.

Covalent

• Nonmetals Bonding to nonmetals
• Electrons are being shared.

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Bonding and Valence Electrons

• S: [Ne] 3s²3p⁴

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• O: [He]2s²2p⁴

• Regardless of what type of bonding, it tends to only be concerned with the valence electrons
• G. N. Lewis; ‘Here’s a nice way to show just the valence electrons’
• Chemical symbol
• 4 sides around the symbol
• 1 dot for each valence electron
• Up to 2 on each side
• Where to double up first is arbitrary
• AKA: Lewis Dot Diagrams
• NOTE: each group have the same number of valence electrons

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Lewis Structure Practice

• Li: 1s²2s¹ B: 1s²2s²

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• P: 1s²2s²2p⁶3s²3p³ Ar: [Ne]3s²3p⁶

The Octet Rule

• Atoms tend to gain, lose or share electrons until they have 8 valence electrons
• The Noble Gases have 8 valence electrons (except He)
• Full s and p subshell
• He only has the s, which is full with 2
• This state is relatively stable and unreactive

Ng

Ne: [He]2s²2p⁶

Ar: [Ne]3s²3p⁶

The Octet Rule

• Ionic Bonding Covalent Bonding

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Na

Cl

O

O

[Ne]3s¹ [Ne]3s²3p⁵

[He]2s²2p⁴ [He]2s²2p⁴

Exceptions to the Octet Rule

• More on this later
• For now; Know that they exist

Summarize

Can you

• Describe the different types of chemical bonds
• Evaluate electron configurations
• Draw Lewis dot diagrams
• Apply the octet rule while being aware of its limitations.

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DONE

What is ionic bonding

• Lose/gain electrons
• Electron configurations
• Lose / gain rules
• Nobel gas config and octet rule
• Transition metals and exceptions to octet
• Polyatomic Ions
• Atoms covalently bonded together
• Act as a whole as a charges species in forming ionic ompounds

8.2 ionic bonding and electron configurations

Objectives:

• Describe what happens during the formation of an ionic bond
• Describe how electron configurations change for metals, nonmetals and transition metals during the formation of ionic bonds

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Ionic Bonding and Electron Configurations

Objectives:

• Describe what happens during the formation of an ionic bond
• Describe how electron configurations change for metals, nonmetals and transition metals during the formation of ionic bonds

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Ionic bonding; What is it?

• One atom loses electron(s) and another one gains those electron(s)

Na

Cl

+

-

Ionic bonding; What is it?

• One atom loses electron(s) and another one gains those electron(s)
• Now these atoms have charges and are attracted to each other
• Ionic bond

Ionic configurations

• Na
• 2-8-1
• 1s² 2s² 2p⁶ 3s¹
• [Ne] 3s¹
• Na⁺
• 1s² 2s² 2p⁶
• [Ne]

• Cl
• 2-8-7
• 1s² 2s² 2p⁶ 3s² 3p⁵
• [Ne] 3s² 3p⁵
• Cl^-
• [Ne] 3s² 3p⁶
• [Ar]

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Na

+

Cl

-

Electrons are removed from the highest available ‘n’ number and are added to the lowest available ‘n’ number. This generally achieves a noble gas configuration

Transition metals

• You still follow the same rules
• Lost electrons come from the highest available ‘n’
• For transition metals, this will be from the ‘s’ sublevel
• If you have lost all you can from the ‘s’ sublevel and still have more electrons to lose, then you move into the ‘d’ sublevel for the lower ‘n’

• Fe

[Ar] 3d⁶ 4s²

• Fe⁺²

[Ar] 3d⁶

• Fe⁺³

[Ar] 3d⁵

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A word about the transition metals

• Their electron configurations get weird
• Sometimes its easier to start putting electrons into the d sublevel before having the s sublevel full
• Ag
• You might expect: [Kr] 4d⁹ 5s²
• But its actually: [Kr] 4d¹⁰ 5s¹
• You still remove electrons from the highest available ‘n’ first
• Noble gas configuration wont be achieved
• Ag⁺ [Kr] 4d¹⁰
• The octet rule has its limitations

Summarize

• Describe what happens during the formation of an ionic bond
• Describe how electron configurations change for metals, nonmetals and transition metals during the formation of ionic bonds

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DONE

8.2b Born Haber Cycle: Energy of ionic bonding

Objectives:

• Interpret and create Born-Haber cycles to represent the energetics involved in ionic bonding
• Explain what Lattice Energy is and it’s role in ionic bonding

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Born Haber Cycle: The Energy of Ionic Bonding

Objectives:

• Interpret and create Born-Haber cycles to represent the energetics involved in ionic bonding
• Explain what Lattice Energy is and it’s role in ionic bonding

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Ionic bonding; What is it?

• One atom loses electron(s) and another one gains those electron(s)

Na

Cl

+

-

Ionic bonding; What is it?

• One atom loses electron(s) and another one gains those electron(s)
• Now these atoms have charges and are attracted to each other
• Ionic bond

Ionic Bonding and Energy: An Exothermic Process

1) Removing electron(s) from an atom

• Ionization energy, ALWAYS an ENDOTHERMIC process
• Na 🡪 Na⁺ + e^- ΔH=496 kJ/mol

2) Gaining electron(s)

• Electron Affinity, usually exothermic
• Cl + e^- 🡪 Cl^- ΔH= -349 kJ/mol

*If this was the whole picture, then ionic bonds would be rarely exothermic

• Ionization energy + Electron affinity
• +496 kJ/mol – 349 kJ/mol =+147 kJ/mol

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Na

+

Cl

-

Then What Are We Missing?

• We have just created two oppositely charged ions
• They are very attracted to each other
• The bigger the difference in charge, the more attracted they are
• Those ions coming together to form a crystal lattice (solid) is VERY exothermic
• This is why ionic compounds are brittle, hard and have VERY high melting points

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Lattice Energy

• “The energy required to completely separate a mole of a solid ionic compound into its gaseous state”
• Positive value for separating ions
• Negative value for ions coming together
• NaCl(s) 🡪 Na⁺ (g) + Cl^- (g) ΔH =+788 kJ/mol
• Na⁺ (g) + Cl^- (g) 🡪 NaCl(s) ΔH =-788 kJ/mol

Born-Haber Cycle

• Thermochemical cycle that shows the energy associated with each step in forming an ionic bond
• Things to consider:
• Electron affinities and Ionization energies are for atoms in the gaseous state
• Diatomic molecules need to be split up into separate atoms before reacting
• Need to refer to their bond enthalpies

Energy

Na (s) + ½ Cl₂ (g)

Na (g) + ½ Cl₂ (g)

Na (g) + Cl (g)

Na⁺ (g) + e^- + Cl (g)

Na⁺ (g) + Cl^- (g)

NaCl (s)

ΔH_sub

ΔH_bond

I₁

E_{aff Cl}

Lattice Energy

ΔH^o_f

Born-Haber Cycle

• Lets take a look at forming NaCl(s)
• ΔH^o_f is exothermic and can be looked up
• Need to get Na(s) to become Na(g).
• Heat of sublimation
• Need to split Cl₂ into separate atoms
• ½ of the bond enthalpy since we are using only ½ of the Cl₂
• Remove an electron from Na(g)
• Ionization energy
• Electron gained by Cl(g)
• Electron affinity
• Na⁺(g) and Cl^-(g) come together to form a solid
• Lattice energy

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Energy

Na (s) + ½ Cl₂ (g)

Na (g) + ½ Cl₂ (g)

Na (g) + Cl (g)

Na⁺ (g) + e^- + Cl (g)

Na⁺ (g) + Cl^- (g)

NaCl (s)

ΔH_sub

ΔH_bond

I₁

E_{aff Cl}

Lattice Energy

ΔH^o_f

Lattice Energy

• A MAJOR driving force for forming ionic bonds
• Accounts for the majority of the exothermic processes
• The bigger the differences in charge, the greater the lattice energy is
• Na⁺ Cl^- < Mg⁺² O^-2 < Al⁺³ N^-3

Energy

Na (s) + ½ Cl₂ (g)

Na (g) + ½ Cl₂ (g)

Na (g) + Cl (g)

Na⁺ (g) + e^- + Cl (g)

Na⁺ (g) + Cl^- (g)

NaCl (s)

ΔH_sub

ΔH_bond

I₁

E_{aff Cl}

Lattice Energy

ΔH^o_f

Born-Haber Cycle

• Applications:
• Generally, you may be asked to solve for the lattice energy by creating a Born-Haber cycle for an ionic bond
• The rest of the process will be given in the problem or can be looked up

Energy

Na (s) + ½ Cl₂ (g)

Na (g) + ½ Cl₂ (g)

Na (g) + Cl (g)

Na⁺ (g) + e^- + Cl (g)

Na⁺ (g) + Cl^- (g)

NaCl (s)

ΔH_sub

ΔH_bond

I₁

E_{aff Cl}

Lattice Energy

ΔH^o_f

Born-Haber Cycle

• Example:
• What is the lattice energy for NaCl given the following:
• Heat of sublimation for Na is 108 kJ/mol
• Bond enthalpy for Cl₂ is 244 kJ/mol
• First ionization energy of Na is 496 kJ/mol
• Electron affinity of Cl is -349 kJ/mol
• The standard heat of formation for NaCl(s) is -411 kJ/mol
• ΔH^o_f = ΔH_sub + ΔH_bond + I₁ + E_{aff Cl} + (-)Lattice Energy
• -411= 108 + (½)*244 + 496 – 349 + (-)Lattice Energy
• Lattice energy = 788 kJ/mol

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Energy

Na (s) + ½ Cl₂ (g)

Na (g) + ½ Cl₂ (g)

Na (g) + Cl (g)

Na⁺ (g) + e^- + Cl (g)

Na⁺ (g) + Cl^- (g)

NaCl (s)

ΔH_sub

ΔH_bond

I₁

E_{aff Cl}

Lattice Energy

ΔH^o_f

Summarize

Can you:

• Interpret and create Born-Haber cycles to represent the energetics involved in ionic bonding
• Explain what Lattice Energy is and it’s role in ionic bonding

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DONE

Covalent bonding

Objectives:

• Describe what covalent bonding is
• Draw Lewis structures for covalent compounds
• Describe what occurs for multiple covalent bonds

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Covalent Bonding

• Most substances don’t behave like Ionic substances
• If they behave differently, maybe its because they are bonded together differently
• SHARE electrons instead of transfer
• Still achieve Noble Gas configuration.

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Na

Cl

[Ne]3s¹ [Ne]3s²3p⁵

H

H

Covalent Bonding

• Positive Nuclei repel each other
• Electron clouds repel each other
• Why do atoms bond?!
• Nuclei are attracted to electron densities of OTHER atoms

Lewis Structures: Single Bonds

• For atoms that only need 1 more electron to achieve a Noble Gas configuration, single bonds form
• Each atom shares 1 of its electrons with the other atom, and ‘keeps the rest for themselves’
• The shared electrons are placed between the two elemental symbols
• A shared pair of electrons usually shown as a line and unshared electrons as dots

H

H

Cl

Cl

Lewis Structures: Multiple single bonds

• Atoms can make multiple single bonds with multiple atoms to complete their octet
• Each single bond is still made up of 1 electron from EACH atom in the bond
• Each bond ‘adds’ 1 electron
• Carbon:
• 4 valence electrons
• needs 4 more to reach the octet
• will make 4 single bonds

C

H

H

H

H

C

H

H

H

H

Lewis Structures: Multiple single bonds

• Nitrogen:
• 5 Valence Electrons
• Needs 3 bonds to reach Octet/Noble Gas configuration

N

H

H

H

Lewis Structures: Multiple single bonds

• Oxygen:
• 6 valence electrons
• Needs 2 bonds to reach octet/noble gas configuration

O

H

H

Lewis Structures: Multiple Bonds

• Atoms can make multiple bonds with 1 other atom!
• Instead of each pitching in 1 electron each, they can pitch in 2 or 3 each!
• Double Bond:
• 2 electron pairs shared
• Two lines drawn
• Triple Bond:
• 3 electron pairs shared
• 3 lines drawn

O

O

Hey! That was great, but I still need to make another bond!

Hey! Me TOO! Want to make ANOTHER bond with me?

That’s a GREAT IDEA!

Now we BOTH have our octet!

N

N

Multiple Bonds and Bond Length

The General Pattern:

• As the number of electrons being shared between 2 atoms increases, the bonds get shorter (nuclei get closer together)
• Single bonds are the longest
• Triple bonds are the shortest

N

N

N

N

N

N

Single Bond

1.47 Å

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Double Bond

1.24 Å

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Triple Bond

1.10 Å

Summarize

Can you

• Describe what covalent bonding is
• Draw Lewis structures for covalent compounds
• Describe what occurs for multiple covalent bonds

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DONE

8.4 Bond polarity and electronegativity

• Objectives:
• Describe what is meant by bond polarity
• Predict bond polarities based on electronegativities of the bonding elements
• Describe bond polarity in terms of dipole moments

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• How electrons can be shared
• Equally
• Unequally
• How that affects bond polarity
• Polar v nonpolar
• Electronegativity
• Def
• Analogy
• Role in determining bond polarity
• Dipole moments
• What are they
• Why do they occur
• Visual representations
• Notations
• Wont get into quantifying them in this video

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Bond Polarity and Electronegativity

Objectives:

• Describe what is meant by bond polarity
• Predict bond polarities based on electronegativities of the bonding elements
• Describe bond polarity in terms of dipole moments

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Covalent bonding

• Electrons are being shared between atoms
• Electrons spend time around both atoms
• This can be done in a ‘fair’ manner, equal times around both atoms
• OR it could be unfair, where the electrons are ‘hogged’ by one atom more than the other

H

H

H

F

Bond Polarity

• Bonds that share electrons evenly are considered NONPOLAR
• There is no magnetic ‘poles’ in the bond
• Bonds that share electrons unevenly are considered POLAR
• They have a ‘north’ and ‘south’ pole, like a magnet
• One region where the negative electrons spend more time (negative) and a region where they don’t (positive)

https://phet.colorado.edu/en/simulation/legacy/molecule-polarity

What determines bond polarity?

• Electronegativity: the attraction for electrons in a bond
• It’s a measure of how ‘greedy’ an atom is for the shared electrons
• More ‘greedy’ = greater electronegativity = hogs the electrons more
• 0 - 4.0 scale
• If the atoms have the same electronegativity, electrons are share evenly
• Nonpolar bond
• If one atom has a greater electronegativity, the electrons are shared unevenly
• Polar bond

H

H

H

F

2.1

4.0

https://phet.colorado.edu/en/simulation/legacy/molecule-polarity

2.1

2.1

Dipole moments

• Di Pole: Two poles
• North and south
• Positive and negative
• “Two electrical charges of equal magnitude but opposite sign are separated by a distance”
• Dipole Moment: the quantitative measure of the magnitude of a dipole
• We wont cover the quantification of dipoles in this video

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Indicating Dipoles

• With showing partial charges

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H – F

2.1 4.0

• The less electronegative element ends up with the positive charge
• The negative electrons spend less time around it

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• With drawing special arrows

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H – F

2.1 4.0

• The arrow points towards the (negative) more electronegative element, showing where the directions electrons are shifting towards
• Forms + sign over the atom that is less electronegative (where the partial + charge would be)

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δ+

δ-

Practice

• Determine the bond type and indicating the dipole where applicable
• H-Cl F₂

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• ClF IBr

Summarize

• Describe what is meant by bond polarity
• Predict bond polarities based on electronegativities of the bonding elements
• Describe bond polarity in terms of dipole moments

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DONE

8.5b Formal Charge

• Objective
• Calculate the formal charge on a bonded atom
• Evaluate Lewis structures using formal charge to determine which is the best.

Formal Charge

Objectives

• Calculate the formal charge on a bonded atom
• Use the Formal Charges to evaluate Lewis structures to determine which structure is the best.

So you have some Lewis structures

• All of them satisfy the octet rule and you don’t know which one is better
• What are you going to do?!

FORMAL CHARGE

Formal Charge

• The charge an atom would have if all the atoms in the molecule had the same electronegativity
• THESE ARE NOT ACTUAL CHARGES ON THE ATOM, but its helpful to pretend in order to figure out which Lewis structure is best

Calculating it:

• F.C. = (# of valence electrons) – (all unshared electrons + ½ shared e^-)

Formal Charge

• Formal Charges should add up to be the overall charge of the molecule/ion

Formal Charges

• How does this help us figure out which Lewis Structure is the best one?
• The best Lewis structure will:
• Have atoms with formal charges closest to zero
• Any negative charges will be on the more electronegative atoms

F.C. = (# of valence electrons) – (all unshared electrons + ½ shared e^-)

Try it out!

And this one!

Summarize

Can you:

• Calculate the formal charge on a bonded atom
• Evaluate Lewis structures using formal charge to determine which is the best.

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DONE

8.5a How to Draw Lewis Structures

Objectives

• Draw Lewis structures for covalent compounds

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How to Draw Lewis Structures

Objectives

• Draw Lewis structures for covalent compounds

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The Steps PCl₃

1. Sum the valence electrons for all of the atoms
2. Write the symbols for all the elements and connect with a single bond
1. Usually written in the order of bonding OR, central atom first.
2. Least electronegative atom goes into the center
3. Complete the octets for the non-central atoms (except H)
4. Place leftover electrons on the central atom
5. If there aren’t enough e^- for the central atom to have an octet, try making multiple bonds

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• 5+ 3*7 = 26

The Steps CO₂

1. Sum the valence electrons for all of the atoms
2. Write the symbols for all the elements and connect with a single bond
1. Usually written in the order of bonding OR, central atom first.
2. Least electronegative atom goes into the center
3. Complete the octets for the non-central atoms (except H)
4. Place leftover electrons on the central atom
5. If there aren’t enough e^- for the central atom to have an octet, try making multiple bonds

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The Steps N₂H₂

1. Sum the valence electrons for all of the atoms
2. Write the symbols for all the elements and connect with a single bond
1. Usually written in the order of bonding OR, central atom first.
2. Least electronegative atom goes into the center
3. Complete the octets for the non-central atoms (except H)
4. Place leftover electrons on the central atom
5. If there aren’t enough e^- for the central atom to have an octet, try making multiple bonds

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What to do when there are multiple possible Lewis Structures?

• NCS^-1
• Which structure is ‘the right one’?
• Formal Charge
• Worth its own video

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Summarize

Can you:

• Draw Lewis structures for covalent compounds

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DONE

8.6 Resonance structures

• Objective
• Describe what Resonance is
• Determine when resonance occurs in a molecule

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Resonance Structures

Objectives

• Describe what Resonance is
• Determine when resonance occurs in a molecule

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Drawing Lewis Structures: O₃

1. Sum the valence electrons for all of the atoms
2. Write the symbols for all the elements and connect with a single bond
1. Usually written in the order of bonding OR, central atom first.
2. Least electronegative atom goes into the center
3. Complete the octets for the non-central atoms (except H)
4. Place leftover electrons on the central atom
5. If there aren’t enough e^- for the central atom to have an octet, try making multiple bonds

But which one is it?

• Which structure is ‘the right one’?
• What makes one of the oxygens more special than the other?
• The answer: Both are ‘right’, and neither oxygen is more special than the other
• What the molecule actually looks like is a mixture of both of these structures
• RESONANCE

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Resonance Structures

• When there are 2 equally good Lewis structures, resonance exists
• The actual molecule isn’t going to be 1 or the other, or switching between each Lewis structure
• The observed structure is going to be a blend of the two
• If the left structure looked red, and the right structure looked yellow, the actual structure is going to look orange (a mix of the two)
• It’s not switching back between red and yellow, its just orange!
• This resonance tends to make a molecule more stable

Resonance Structures

Bond Lengths

• O – O 147 pm
• O = O 121 pm
• Observed bond lengths in Ozone (O₃) 128pm

Observed length is between the length of a single and the length of a double bond

• It’s a blend of the two

Resonance Structures

• Then where are the electrons?
• Typically: Double bond means 4 e^- shared between those two atoms
• Delocalized Electrons
• Electrons aren’t trapped between the 2 atoms, but are shared and move around the entire molecule

How Do We Indicate Resonance In Our Lewis Structures?

• You draw the Lewis structures and draw a double arrows between them
• The observed bonds are between a single and a double bond, but we have no way of drawing that.
• There may be more than 2 resonance structures; draw them all!

Practice NO₃^-

Practice SO₃

Benzene C₆H₆

• Closed Ring Hydrocarbon
• Alternating double bonds between the carbons
• The 2x bonded electrons are delocalized across the entire molecule
• C-C 1.54 Å
• C=C 1.34 Å
• Observed in benzene: 1.40 Å
• Shorthand looks like a nut for a bolt

© Tomas Castelazo, www.tomascastelazo.com / Wikimedia Commons

Summarize

• Describe what Resonance is
• Determine when resonance occurs in a molecule

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DONE

8.7 exceptions to the octet rule

Objectives

• Describe the limitations to the octet rule and explain why sometimes having less than, or more than an octet is appropriate

Exceptions To The Octet Rule

Objectives

• Describe the limitations to the octet rule and explain why sometimes having less than, or more than an octet is appropriate
• Yes, exceptions to the octet rule. Still, chemistry isn’t as bad as English with all of its exceptions. At least that’s what I leisurely tell my eight overweight neighbors

The Octet Rule

• Atoms tend to bond or transfer electrons in such a way to achieve the stable configuration of 8 valence electrons
• Stable because ‘s’ and ‘p’ sublevels are full
• The Octet Rule has its limitations however

• Na 2-8-1 Cl 2-8-7
• Na⁺ 2-8 Cl^- 2-8-8
• CO₂

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BH₃

The Octet Rule Exceptions

• Odd number of electrons
• Some molecules have an odd number of electrons, making it impossible to complete an octet

NO

The Octet Rule Exceptions

• Less than an octet
• Rare to occur, but common in Boron and Beryllium (and Hydrogen)
• Check the formal charges

BF₃

The Octet Rule Exceptions

• More than octet
• Elements in period 3 or greater
• Access to the ‘d’ sublevel
• 2^nd period elements don’t have a ‘d’ sublevel to work with
• Generally want to choose the Lewis Structure that satisfies the octet rule when possible

PCl₅

Summarize

• Describe the limitations to the octet rule and explain why sometimes having less than, or more than an octet is appropriate

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DONE

8.8a Strengths of Covalent Bonds

Objectives: Describe bond strength in terms of bond enthalpy and describe factors that affect bond enthalpy.

Strengths of Covalent Bonds

Objectives:

• Describe bond strength in terms of bond enthalpy and describe factors that affect bond enthalpy.

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What are bonds?

• Two atoms that are ‘stuck’ together because of how their electrons interact
• Transferred from one atom to the other OR
• Both atoms can share some of their electrons with each other

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Ionic Covalent

Cl

Na

H

Cl

How strong are the bonds?

• Strength of a bond = how difficult is it to break those atoms apart
• How much energy would it take to break it?
• Known as Bond enthalpies
• They are ALWAYS positive values: it takes energy to break bonds

H

Cl

CH₄ breaking up

CH₄ 🡪 C + 4 H ΔH = +1660 kJ

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Bond Enthalpy of C-H bond: (+1660 kJ) / (4 C-H bonds) = 415 kJ/mol

C

H

H

H

H

C

H

H

H

H

Factors Affecting Bond Strength

• What elements are bonding together
• H bonding to C has a different bond enthalpy than H bonding to F

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C-H = 415 kJ/mol H-F = 567 kJ/mol

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C

H

H

H

H

Factors Affecting Bond Strength

• BOND TYPE
• Single: The longest bond and the easiest to break
• Double: The in between
• Triple: The shortest bond and the hardest to break

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348 kJ/mol 614 kJ/mol 839 kJ/mol

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1.54 Å 1.34 Å 1.20 Å

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C

C

C

C

C

C

Factors Affecting Bond Strength

• The structure of the rest of the molecule
• What is going on in the rest of the molecule may affect a particular bond’s strength
• Resonance
• Polar bonds

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H

C

O

O

H

C

H

H

H

H

O

How strong are the bonds? It Varies

• The bond strength varies between different molecules, but usually only slightly
• As a result, we use the average bond enthalpies
• You look these up on a chart somewhere, this isn’t information that you will need to memorize

Average bond enthalpies

• Values found on a chart

A pet peeve from Biology

“Energy is released when bonds are broken” or “Energy is stored in chemical bonds”

WRONG

• ENERGY IS ALWAYS ABSORBED TO BREAK A BOND
• Energy is released when new bonds are made.
• Sometimes the energy being released by the new bonds is more than you had to invest to break the old bonds. Overall, energy is released (exothermic)
• Other times, the energy being invested to break the bonds is greater than the energy released when new bonds are made. Overall, energy is absorbed (endothermic).
• Easier to say “Respiration releases energy because energy is stored in chemical bonds” than actually explaining all of this.

To summarize

• Bond enthalpy is the energy required to break a bond
• Varies slightly based on a few factors, so we use an average
• Energy is invested to break bonds and released when new bonds are made
• Biology teachers may have been lying to you.

DONE

8.8b Bond Enthalpies and Enthalpies of reactions

• Objectives: Use bond enthalpies to determine the enthalpy of a reaction

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Bond Enthalpies and Enthalpies of Reactions

Objectives:

• Use bond enthalpies to determine the enthalpy of a reaction

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What is going on in a chemical reaction?

• Bonds are being broken
• Investment of energy
• New bonds are being made
• Release of energy
• The difference between the energy invested in breaking the bonds and the energy evolved in creating new ones is the ΔH for that reaction

H-H

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H-H

O=O

H H

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H H

O O

H₂O

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H₂O

Mathematically

• ΔH = Σenthalpies of bonds broken – Σenthalpies of bonds made

Positive value because the energy is being invested (added) to break the bonds

Negative value (subtracted) because energy is being given off (lost) when new bonds are being made

Think of it as playing with magnets

• Need to put energy into separating them
• Energy is given off when they come together and bond

What is the ΔH for the following reaction:

C + 2 H₂ 🡪 CH₄

_{Need to look at the structure of each molecule}

What is the ΔH for the following reaction based on bond enthalpies

• CH₄ + 2 O₂ 🡪 CO₂ + 2 H₂O

Do you get it and wish I would stop talking?

1: Stop yelling

2: Then stop watching and skip over the additional practice that I will show! ☺

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H₂CCH₂ + H₂O ---> CH₃CH₂OH

To summarize

• Look at the structure
• Determine what bonds are being broken and made
• Determine how many of each (account for the coefficient in front)
• Take the difference between the energy invested in breaking the bonds of all the reactants and the energy released by making new bonds for the products

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