unit 08 bonding concepts - harris chemistry€¦ · unit 08 bonding concepts 1 unit 8/9: ... rank...

AP Chemistry - Michalek Unit 08 Bonding Concepts

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Unit 8/9: Bonding Concepts Last revised: June 15, 2011

“One is almost tempted to say... at last I can almost see a bond. But that will never be, for a bond does not really exist at all: it is a most convenient fiction which, as we have seen, is convenient both to experimental and theoretical chemists.”

You should be able to: Differentiate between ionic, covalent, and metallic bonds.

Use Coulomb’s Law to find the energy of interactions between a pair of ions

Explain why electronegativites are important to bond polarities

Explain how charges effect ionic electrostatic forces

Define isoelectronic ion

Explain the energy needed to make a binary ionic compound from its elements

Calculate the lattice energy

Explain the ionic character of covalent compounds

Calculate ΔH from bond energies

Describe the localized electron model

Draw lewis dot structures

o Explain the octet rule

o Know exceptions to the octet rule

o Show resonance structures

o Write formal charges

o Name VSEPR shapes

Describe the differences in single, double, and triple bonds

o Length

o Strength

o Bonding

8.1 Types of Chemical Bonds Read and outline 8.1 Define

Bond Energy

Ionic Bonding

Ionic Compound

Coulomb’s Law

Bond Length


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Covalent Bonding

Polar Covalent Bond Why do Na+ and Cl- form a bond? What does the negative sign mean for energy when solving Coulomb’s Law? How does a bonding force develop between two identical atoms? What are the four postulates to the bond length theory? 1. 2. 3. 4. What is the symbol for partial charge? Why is a partial charge formed in some atoms?


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8.1 Notes Bond Energy is the energy required to break a bond. Ionic Bonding is a bonding force that forms a great thermal stability results from the electrostatic attractions of the closely packed, oppositely charged ions. Ionic compound results when cations react with anions. Coulomb’s Law The energy of interactions between a pair of ions

k = 2.31 x 10-19 J*nm, E is energy, r is the distance between ion centers, Q1 and Q2 is the ionic charge.

Example In solid sodium chloride the distance between the centers is 2.76 A. What is the energy of interaction between the pairs of ions? If the sign is negative, the ion pair has a lower energy than the separated ions. This means a bond will form. Chemicals are most stable in the lowest energy form. This law can also be used to calculate the repulsive energy when two like charged ions are brought together. A bond will form if the energy of the aggregate is lower than that of the separated atoms. Bond length The system will act to minimize the sum of the positive (repulsive) energy term and the negative (attractive) energy term. The distance where the energy is at a minimum is called bond length. In covalent bonding the electrons are being shared by the nuclei. Polar covalent bond share electron unequally; meaning that the electrons are pulled closer to the most electronegative atom.


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8.2 Electronegativity Read and outline 8.2 Define

Electronegativity

8.2 Notes Electronegativity is the ability of an atom to attract shared electrons to itself. Expected H-X bond energy = H-H bond energy + X-X bond energy 2 ∆ = (H-X) act – (H-X)exp When the bond strength is greater the actual bond energy will be larger than the expected. Example Order the following bonds according to polarity: H-H, O-H, Cl- H, S-H, and F-H.

8.5 Formation of Binary Ionic Compounds Read and outline 8.5 Define

Lattice Energy Describe the steps to form a generic ionic bond. (This means don’t use any symbols of chemicals) 1. 2. 3. 4. 5.


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What type of thermodynamic reaction brings cations and anions together? Why is the energy released greater in MgO than NaF? Why is it much greater to remove two electrons from Mg than one electron from Na? Why is O2-not stable?

8.5 Notes Lattice energy—the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid. Steps of forming an ionic solid from its elements.

1. Sublimation of the metal 2. Ionization of lithium atoms to form the ion in gas phase 3. Dissociate the diatomic atom 4. Formation of the anion 5. Formation of the compound

Lattice Energy =


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Worksheet 8.01 Name______________________________ Lattice Energy 1. Circle the compound in the pair of ionic substances that has the most exothermic lattice energy. Sodium Chloride / Potassium Chloride Lithium Fluoride / Lithium Chloride Magnesium Hydroxide / Magnesium Oxide Iron (III) Hydroxide / Iron (II) Hydroxide 2. Use the following data to estimate ΔHf for sodium chloride. Na (s) + ½ Cl2 → NaCl.

Lattice Energy -786 kJ/mol

Ionization Energy for Na 495 kJ/mol

Electron Affinity of Cl -349 kJ/mol

Bond Energy of Cl2 239 kJ/mol

Enthalpy of sublimation of Na 109 kJ/mol

3. Why are some bonds ionic and some covalent? 4. Rank the following in order of increasing ionic character: N—O, Ca—O, C—F, Br—Br, K—F. 5. Give three ions that are isoelectronic with neon. Place them in order of increasing ionic size. 6. Give three ions that are isoelectronic with Argon. Place them in order of decreasing ionization energy.


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8.9 The Localized Electron Bonding Model Read and outline 8.0 Define

Localized electron model

Lone pair

Bonding pair What are the three parts to the localized electron model? 1. 2. 3.

8.10 Lewis Dot Structures Read and outline 8.10 Define

Lewis Structure

Duet Rule

What is the most important requirement to form a stable compound? Rules for writing Lewis Structures: 1. 2. 3.


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8.10 Notes Rules for Writing Lewis Structures

1. Sum the valence electrons from all the atoms.

2. Don’t forget to add in the charges

3. Use a pair of electrons to form a bond between each pair of bound atoms.

4. Arrange the remaining electrons to satisfy the duet rule for hydrogen and the octet rule for the second-row of elements.

8.13 Molecular Structure: The VSEPR model Read and outline 8.13 Define

Molecular Structure

Valence Shell Electron Pair Repulsion

Linear Structure

Trigonal Planar Structure

Tetrahedral Structure

Trigonal Pyramidal

Trigonal bipyramidal

Octahedral Structure

Square Planar Structure


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What are the 4 steps to predict the structure of a molecule? 1. 2. 3. 4. ____________________requires more room than___________________pairs. Why? How many effective electron pairs are double and triple bonds counted as? What structure do you use when resonance is involved?


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8.13 Molecular Structure: the VSEPR Model


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9.1 Hybridization and the Localized Electron Model Read and outline 9.1 Define

Hybridization

sp3 hybridization

hybrid orbitals

sp2 hybridization

sigma bonds

pi bonds

sp hybridization

dsp3hybridization Methane

d2sp3hybridization

9.1 Notes

Hybridization The mixing of the native atomic orbitals to form special orbitals for bonding. With sp3 orbitals, the atomic orbitals of s and p reshape to form the tetrahedral shape. Whenever a set of equivalent tetrahedral atomic orbitals is required by an atom, this model assumes that the atom adopts a set of sp3 orbitals; the atom becomes sp3 hybridized. Example Describe the bonding in the ammonia using the localized electron model.


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A double bond always has a one sigma and one pi bond.

Example 8.10/8.13/9.1a Write a LDS for the following compounds:

Picture VSEPR name Hybridization

H2O

CO2


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CN-

HF

N2

NH3

CH4

CF4

NO+,

PCl5


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I3-

ClF3

XeO3

RnCl2

BeCl2

ICl4-

Example 9.1b Describe the bonding in the N2 molecule.


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Example 9.1c Describe the bonding in the triodide ion I3

- Example 9.1d How is the Xenon atom in XeF4 hybridization. Example 9.1e For each of the following molecules or ions, predict the hybridization of each atom, and describe the molecular structure. CO, BF4

-, XeF2


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Worksheet 8.02 Name_____________________________ Lewis Dot Structures Draw Lewis structures for the following compounds and write the VSEPR shape on the bottom line of the box.

HCN PH3

CHCl3 NH4+

H2CO3 SeF2

CO2 O2

HBr POCl3


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8.12 Resonance Read and outline 8.12 Define

Resonance

Resonance Structures

Formal Charge

8.12 Notes

Example 8.12a Describe the electron arrangement in the nitrite anion using the localized electron model.


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Formal Charge –charge assigned to each atom in a molecule Rules Governing Formal Charge

1. To calculate the formal charge on an atom: a. Take the sum of the lone pair electrons and .5 the shared electrons. b. Subtract the number of assigned electrons in valence shell

2. The sum of the formal charges is the charge of the atom or ion. Example 8.12b Give possible Lewis Structures for XeO3. Which Lewis Structure or structures are most appropriate according to the formal charge?


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Worksheet 8.03 Name_______________________________ Resonance Draw Lewis structures for the following compounds, write the VSEPR shape on the bottom line of the box, and place the formal charge next to the atom in a DIFFERENT COLOR.

NO2-

NO3-

OCN- SCN-

N3- C6H6 (the carbons for a ring)

B3N3H6 ( Alternating ring of B/N) SO3


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8.3 Bond Polarity and Dipole Moments Read and outline 8.3 Define

Dipolar

Dipole Moment What is the symbol for a partial negative charge? Positive charge?

8.3Notes A dipole moment is when there is a positive and negatively charged center.

Example 8.3a Which of the following molecules are polar: NH3, SF6, PF5, CH3Cl?

8.6 Partial Ionic Character Read and outline 8.6 What is the equation to find the percent ionic character of a bond? What is the operational definition of an ionic compound?


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How can we tell the difference between covalent, polar covalent, and ionic bonds? Percent ionic character of a bond = Measured dipole moment of X—Y * 100 Calculated Dipole moment of X+Y-

This will never reach 100%. This means no bond is completely ionic. More than 50 % ionic character is classified as an ionic bond.


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Worksheet 8.04 Name_______________________________ Lewis Dot Structures Draw Lewis structures for the following compounds, write the VSEPR shape on the bottom line of the box, and tell if the molecule is polar or nonpolar.

SO4-2 XeO4

PO4-3 NF3

ClO2- SO2Cl2

NO4-3 BeH2

ClF5 SF6


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8.7 The Covalent Chemical Bond: A Model—Read this section Chemical bonds can be viewed as a force that causes a group of atoms to behave as a unit. Bonds result from the tendency of a system to seek its lowest possible energy.

8. 8 Covalent Bond Energies and Chemical Reactions Read and outline 8.8 Define

Single Bond

Double Bond

Triple Bond What is the equation for ΔHf from unit 6? What process is the formation of a bond? When a bind breaks is energy absorbed or releases from the molecule? What process is this?

Bond Type Bonds Shared electron pairs

Single 1 1

Double 2 2

Triple 3 3

ΔHBond energies = nBE reactant – nBE products Example 8.8a Using the bond energies Calculate ΔH for the reaction of methane with chlorine and Fluorine for give CF2Cl2 with hydrofluoric and hydrochloric acid.


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Worksheet 8.05 Name______________________________ Bond Energies Estimate ΔH for the following reactions using bond energies.

1. C2H5OH + O2 → CO2 + H2O

2. H2CCH2 + O3 → CH3OH + O2

3. Glucose → carbon dioxide + ethanol

4. Ethane + fluorine gas → C2F2H4

5. H2 + Cl2 → HCl


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9.2 Read and outline 9.2 Define

Molecular Orbital model

Molecular Orbitals

Sigma Molecular Orbitals

Bonding molecular Orbitals

Antibonding molecular orbitals

Bond order Take note on this section!!!


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9.2 Notes The molecular orbital model is used to describe bonding. Just as atomic orbitals are solutions to the quantum mechanical treatment of atoms, molecular orbitals (MOs) are solutions to the molecular problem. They can hold two electrons with opposite spin and that the square of the molecular orbital wave function indicates electron probability.

1. The electron probability of both molecular orbitals is centered along the line passing through the two nuclei.

For MO1 the greatest electron probability if between the nuclei, and for MO2 it is on either side of the nuclei. This type of electron distribution is described as sigma, as in the LEM. Accordingly, we refer to MO1 and MO2 as sigma molecular orbitals.

2. In the molecule only the molecular orbitals are available for occupation by electrons. The 1s atomic orbitals of the hydrogen atoms no longer exist, because the H2 molecule –a new entity__ has its own set of new orbitals.

3. MO1 is lower in energy than the 1s orbital of free hydrogen atoms, while MO2 is higher in energy. Since H2 has a lower energy in the MO1, they will have a lower energy than they do as two separate atoms. This situation favors molecule formation, because nature tends to seek the lowest energy state. That is, the driving force for molecule formation is the molecular orbital available to the two electrons has lower energy than the atomic orbitals. This situation is favorable to bonding. A bonding molecular orbital is lower in energy than the atomic orbitals of which it is composed. An antibonding orbital is higher in energy than the atomic orbitals of which it is composed. Electrons in this type of orbital will favor the separated atoms.

4. Since the greatest probability of electrons are between the orbitals this bonding force is greater than the antibonding force.

5. MO1 = σ1s MO2 = σ1s*

6. Each molecular orbital can hold two electrons, but the spins must be opposite 7. Orbitals are conserved. The number of molecular orbitals will always be the same as the number of atomic

orbitals used to construct them. When using H2

- the electrons use the antibonding orbital. Two electrons are lowered in energy and one is raised, producing a new lowering of the energy of only one electron. Thus the model predicts that H2 is twice as stable as H2

-. With respect to their separated components H2 is twice as strong as H2

-. Bond Order Bond order is the difference between the number of bonding electrons and the number of antibonding electrons divided by two. Bond order = number of bonding electrons – number of antibonding electrons 2


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Example 9.2a What is the bonding order of hydrogen and helium?

9.3 Bonding in Homonuclear Diatomic Molecules


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Paramagnetism Paramagnetism causes the substances to be attracted into the inducing magnetic field. Diamagnetims causes the substance to be repelled from the inducing magnetic field. Example 9.3a For the species O2, O2

+, O2-, give the electron configuration and the bond order for each. Which has the strongest bond?

Example 9.3b Use the molecular orbital model to predict the bond order and magnetism of each of the following molecules: Ne2 and B2.

9.4 Bonding in Heteronuclear Diatomic Molecules Example 9.4a Use the molecular orbitals model to predict the magnetism and bond order of the NO+ and CN- ions.

9.5 Combining the Localized Electron and Molecular Orbital Models

Delocalized pi bonding- spreading of electrons over multiple bonds


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Worksheet 9.01 Name ___________________________________ MOs and Bond Order

1. Use the molecular orbital model, write electron configurations for the following diatomic species and calculate the bond orders. Are they paramagnetic or diamagnetic? Li2, O2, C2

2. Use the molecular orbital model, write electron configurations for the following diatomic species and calculate

the bond orders. Are they paramagnetic or diamagnetic? CO, CO+, CO+2.

3. Place the species in number two in order of increasing bond length and increasing energy. Explain your answer.

4. Describe the bonding in O3 molecule and the NO2- ion using the localized electron model. How would the

molecular orbital model describe the pi bonding in these two species?

unit 08 bonding concepts - harris chemistry€¦ · unit 08 bonding concepts 1 unit 8/9: ... rank...

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