The NSF Bonding Module NH NH2 NH3.

The NSF Bonding Module NH NH2 NH3

Introduction Thanks for participating in today’s session! We’ve been developing a new approach for teaching about bonding and molecules in General Chemistry courses like Chem 102. Our objective is to determine if the material makes sense to students. Discovering the Nanoworld: A New Module for Teaching about Molecules and Bonding in General Chemistry The material will be substantially different than what you’ve seen in Chem 102 or other classes. We hope that it helps you to understand the concepts!

Introduction To receive the compensation of $25 for participating, you need to do three things: (1) Respond to the Pre-Session Questionnaire (2) Attend this session (3) Respond to the Post-Session Questionnaire    To fulfill the third condition, please log into Lon-Capa again and respond to the second set of questions you’ll find in the “NSF Bonding Module” course. The questions will be available Thursday morning and accessible through the end of the semester, but the sooner you respond, the sooner we will be able to tell you where to pick up your $25 in cash.

Format A number of questions are interspersed through the lecture material. Please respond to them on the scantron answer sheet provided to you. There will be 3 types of questions:  Background Questions: What do you already know?  Comprehension Questions: Is a concept making sense to you?  Anticipation Questions: Can you deduce where things are going? Let’s get started!

Background Questions 1. Which of the following statements is FALSE?
(a) Light can’t move faster than 3 x 108 m/s in a vacuum. (b) The gravitational force between two masses is proportional to m1 m2 / r2, where r is distance between m1 and m2. (c) Unlike charges repel each other, while like charges attract each other (Coulomb’s law). (d) I don’t know.

Background Questions 1. Which of the following statements is FALSE?
(a) Light can’t move faster than 3 x 108 m/s in a vacuum. (b) The gravitational force between two masses is proportional to m1 m2 / r2, where r is distance between m1 and m2. (c) Unlike charges repel each other, while like charges attract each other (Coulomb’s law). (d) I don’t know. In fact, Coulomb’s law says the opposite. Like charges REPEL each other, while unlike charges ATTRACT each other, so  Two protons (p+) repel each other.  Two electrons (e–) repel each other.  An electron and a proton attract each other.

Background Questions 2. Which of these potential energy curves corresponds to attraction between unlike charges, such q1 = 1 and q2 = –1? (a) (b) (c) (d) I don’t know.

Background Questions 2. Which of these potential energy curves corresponds to attraction between unlike charges, such q1 = 1 and q2 = –1? (a) (b) (c) attraction repulsion bound minimum

Background Questions 3. Which of the following is the most accurate model of the H atom (a proton, p+, and electron, e–)? (b) (c) (a) (d) I don’t know.

Background Questions 3. Which of the following is the most accurate model of the H atom (a proton, p+, and electron, e–)? (a) (b) (c) Orbit model Atoms don’t behave like solar systems! Orbital model Quantum mechanics rules at the atomic level!

Background Questions 4. The Pauli Exclusion Principle says that two electrons cannot have the same quantum numbers. This means that: (a) Electrons always repel each other and cannot be in the same orbital. (b) An orbital can hold two electrons if they have the same spin. (c) An orbital can hold two electrons if they have opposite spins. (d) I don’t know.

The Dilemma of Quantum Mechanics
The quantum mechanical framework needed to properly describe the behavior of electrons and nuclei in atoms and molecules involves advanced math (partial differential equations and solutions to eigenvalue problems). The bad news: The math is beyond what general chemistry students are expected to know. The good news: We can use graphical results from programs that treat the quantum mechanics to learn basic concepts without getting into the math.

Sulfur Chemistry on Venus Ice Chemistry on Interstellar Dust
I use quantum chemical programs in my research to study astrochemistry for NASA.

= = = The H Atom with Multiple Representations Orbitals: H:
3D surface (isodensity) = 2D contours (cross-section) = 1D cross-section H: cloud The electron is more likely to be found near the proton than farther away from it.

Oxygen Atom – Review              
5. Oxygen atoms (z=8) have 8 electrons, of which 6 are valence electrons that are found in the 2s and 2px, 2py, and 2pz orbitals. What’s the most stable occupation of electrons for O? (a) 2s0 2p6 (b) 2s1 2p5 (c) 2s2 2p4 (d) 2s2 2p4 (e) I don’t know. 2s 2px 2py 2pz              

Each p orbital gets one e– before being doubly occupied.
Oxygen Atom – Review 5. Oxygen atoms (z=8) have 8 electrons, of which 6 are valence electrons that are found in the 2s and 2px, 2py, and 2pz orbitals. What’s the most stable occupation of electrons for O? (a) 2s0 2p6 (b) 2s1 2p5 (c) 2s2 2p4 (d) 2s2 2p4 (e) I don’t know. 2s 2px 2py 2pz           Each p orbital gets one e– before being doubly occupied.    

Oxygen Atom – Valence Orbitals
Let’s build up O orbital by orbital: 3D 2D 2s orbital 2s orbital 3D 2s

Let’s build up O orbital by orbital: 3D 2D 2px orbital 2px orbital 3D 2s

Let’s build up O orbital by orbital: 3D 2D 2px orbital 2px orbital 3D 2s+2px

Let’s build up O orbital by orbital: 3D 2D 2py orbital 2py orbital 3D 2s+2px

Let’s build up O orbital by orbital: 3D 2D 2py orbital 2py orbital 3D 2s+2px+2py

Let’s build up O orbital by orbital: 3D 2pz orbital 3D 2s+2px+2py

Let’s build up O orbital by orbital: 3D 2pz orbital 3D 2s+2px+2py+2pz

Let’s build up O orbital by orbital: 3D 2px+2py+2pz

Let’s build up O orbital by orbital: 3D 2D 2px+2py+2pz 2px+2py+2pz

Symbolic Representations for Atoms
It’s useful to have a shorthand representation for atoms that conveys useful information about them without drawing the orbitals every time. The symbols should represent the shapes of the various orbitals and should tell us how many electrons are in them (their occupations). For example, the 1s orbital of H is spherical, and it has one electron. Here’s one representation of that: H orbital symbol We represent the occupation with a dot, but that’s not the location of the electron.

Representing Valence Electron Configurations
H He 2D H He 3D yellow or blue: singly occupied gray: doubly occupied number of dots indicates occupation 2s0 2pz0 2px0 2py0

2px 2py 2pz O (2s2 2p4)   2s0 2pz0 2px0 2py0

2px 2py 2pz O (2s2 2p4)   2py1 2pz1 2px2 2s2

Valence Electron Configurations: Li–Ne
Li (2s1)  Be (2s2)  B (2s2 2p1)   C (2s2 2p2)  

Valence Electron Configurations: Li–Ne
N (2s2 2p3)   O (2s2 2p4)   F (2s2 2p5)   Ne (2s2 2p6) 

Representations of H–Ne
clear green: unoccupied yellow or blue: singly occupied gray: doubly occupied H He Li Be B C N O F Ne

The Nature of Molecular Chemistry
Molecular chemistry is about two things: 1. The structures and properties of molecules (i.e., assemblages of atoms) 2. The interactions of molecules and atoms (reactions) Let’s explore what happens if we bring different atoms together in different ways to see what that tells us about both the nature of bonding and reactivity. For now, we’ll limit ourselves to pairings from this set: H H+ He Li Be B C N O F Ne

Nature and Types of Chemical Bonds
Consider this set: H+ H He H+ + 1s0  1s0 H+ H + 1s0  1s1 He 1s0  1s2 1s1  1s2 H + 1s1  1s1 He 1s2  1s2

Anticipation Question
6. What happens if two H+ ions approach each other? (a) They repel each other. (b) They attract each other. (c) They orbit each other.

6. What happens if two H+ ions approach each other? (a) They repel each other. (b) They attract each other. (c) They orbit each other. H+ If two protons approach each other, they will be kept apart by Coulombic repulsion.

Some More Repulsive Interactions
If H approaches He or two He atoms approach one another, the interactions are also repulsive (mostly). H He

Orbitals and Potential Energy Curves
It’s useful to watch what happens to orbitals as atoms move toward or away from one another. He 1s2 pair on He1 1s2 pair on He2

Orbitals and Potential Energy Curves
It’s useful to watch what happens to orbitals as atoms move toward or away from one another. The Pauli Principle prevents the two He 1s2 pairs from forming a bond. All they can do is try to get out of each other’s way! He 1s2 pair on He1 1s2 pair on He2

7. What do you think happens when two H atoms come together? (a) They repel each other. (b) The H2 molecule forms. (c) It depends: we need more information.

7. What do you think happens when two H atoms come together? (a) They repel each other. (b) The H2 molecule forms. (c) It depends: we need more information. There are two combinations of electron spin that can occur when two H atoms come together: H () ()  ? () ()  ? +  ?

8. Which combination do you think yields a bond? (a) When the spins are the same. (b) When the spins are different.

This is the bond length of H2.
H + H Potential Energy Curves Here are the two outcomes: Curve (1) is purely repulsive. Curve (2) is attractive at large separations but repulsive at small separations, forming a bound minimum. This is the bond energy of H2. This is the bond length of H2.

Repulsive H + H Potential Energy Curve
The repulsive curve occurs if the spins on the two H atoms are the same. 

Repulsive H + H Potential Energy Curve
The Pauli Exclusion Principle prevents the electrons from being in the same space. 

Bound H + H Potential Energy Curve
The bound curve occurs if the electrons on the two H atoms have opposite spins. 

Each electron moves toward the other nucleus as the nuclei get closer together. 

If the spins are different, the Pauli Exclusion Principle allows the electrons to overlap… 

… and H2 can form. Note that the overlap bar increases as the electrons overlap. 

The electrons are coupled together to form the bond.
Definitions A chemical bond is a bound (or stable) configuration of atoms. H2 + H  The electrons are coupled together to form the bond. A bond that forms along the axis between two nuclei is called a s bond.

Explanation Question 9. Where does most of the bond energy of H2 come from? (a) Attraction between the two electrons. (b) Attraction between the two protons. (c) Attraction between each electron and the other proton.

H + F Consider this set: H 1s1  2pz1 F  ? H 1s1  2pz2 F  ?

H + F – Anticipation Question 1
Consider this arrangement: H F 10. What do you think will happen if H and F approach each other in this arrangement? (a) A bond will form if the unpaired electrons on H and F have the same spin. (b) A bond will form if the unpaired electrons on H and F have opposite spins. (c) H and F can only interact weakly because they are too different.

H + F – Anticipation Question 2
Consider this arrangement: H F 11. What do you think will happen if H and F approach each other in this arrangement? (a) H will form a bond with one of the electrons in the 2pz2 pair of F. (b) H and F are repulsive in this arrangement.

H + F Potential Energy Curves
1s1  2pz1 () 1s1  2pz1 () 1s1  2pz2

H + F Potential Energy Curves
1s1  2pz1 () 1s1  2pz1 () 1s1  2pz2 HF

Bound H + F Potential Energy Curve
At a long separation (Re + 2 Å), we see the H 1s1 orbital and the F 2pz1 orbital. F H

As the nuclei get closer, the H orbital shifts a great deal toward F, but the F orbital barely moves. F H

This trend continues to Re. The H orbital resembles a second, distorted F orbital. F H

The animation shows that both orbitals move toward the other nucleus, but the H orbital moves much more. F H

More Definitions We call the bonds in both H2 and HF covalent bonds because the two electrons are shared between the two atoms. While the sharing is equal in H2, it is not equal in HF. We say that the pair of electrons is polarized toward F, so the bond in HF is a polar covalent bond.

An Extreme Example of Polarization
If we look at the orbitals for LiF at its minimum separation, we see that they are almost completely localized on F. Li F Li F A bond that involves the complete or nearly complete transfer of an electron from one atom to another is called an ionic bond.

Charge and Dipole Moments
When HF forms, the orbitals involved in the bond change from atom-like: … to ones that are shifted toward F: F H F H

In HF, F has more electronic charge than it did before the bond formed; it has become negatively charged. In HF, H has less electronic charge than it did before the bond formed; it has become positively charged. F H –q +q The charge shift means that HF has a dipole moment.

12. Do you think that H2 has a dipole moment? (a) Yes, because each H’s orbital moves toward the other H. (b) No, because the orbitals shift the same amount. (c) No, because the orbitals don’t shift when H2 forms.

X–H Bonds for X={B–O} B H +  NH OH BH CH C N O

X–H Bonds for X={B–O} 13. Which statement is true?
(a) The four hydrides of B–O can form if the electrons in their 2pz orbital and the H 1s orbital have the same spin. (b) The four hydrides of B–O can form if the electrons in their 2pz orbital and the H 1s orbital have opposite spins. (c) None of these hydrides can form because they don’t satisfy the octet rule.

NeH? Ne H +  ? 14. What do you think happens if we try to add H to Ne as shown above? (a) A bond forms. (b) A bond cannot form.

Electron Waves in a Quantum Corral

The Well (Quantum Corral) – a sculpture by Julian Voss-Andreae
Electron Waves in a Quantum Corral The Well (Quantum Corral) – a sculpture by Julian Voss-Andreae

Reactivity We’re seen cases where molecules form when two atoms come together and other cases where they repel each other. The first case is an example of a reaction, while the second case is non-reactive. Any process that changes one or more atoms or molecules into something else constitutes a reaction.

Couple these two orbitals…
Multiple Bonds: NO Let’s see what happens when N and O come together: Couple these two orbitals… N O

Couple these two orbitals…
Multiple Bonds: NO Let’s see what happens when N and O come together: Couple these two orbitals… s bond N O

Multiple Bonds: NO Let’s see what happens when N and O come together:
p bond s bond N O A compound like NO with one or more singly occupied orbitals is called a radical.

Multiple Bonds: NO Here’s an animation of the s bond forming: O N

Multiple Bonds: NO Here’s an animation of the p bond forming:
For both bonds, the N orbital shifts more toward O than the O orbital shifts toward N, so both bonds are polarized toward O.

What Do You Think? 15. How reactive do you think NO is?
(a) It’s not very reactive, because it has a double bond. (b) It’s reactive, because it has an unpaired electron.

Coupled Pair Bonds All of the bonds we’ve seen so far have involved two electrons: the covalent s bond in H2, the polar covalent s bond in HF, the ionic bond in LiF, and the slightly polar s and p covalent bonds in NO. We can call all of these coupled pair bonds because in each case both entities involved in the bond contribute one electron to each of the bonds. Bonding occurs be-cause each pair of electrons in question have opposite spins.

Something Different 16. What do you think happens if a H atom (1s1) interacts with a Be atom (1s2 2s2)? Every other time we’ve tried this, the interaction has been repulsive. (a) H–Be is also repulsive. (b) Something probably happens. Why else would the instructor bring this up again?

The Bond in BeH Here are potential energy curves for the interactions of H with He and Be. They are quite different! Be H + Be has no unpaired electrons. It can’t form a normal coupled pair bond.

The Bond in BeH Here are potential energy curves for the interactions of H with He and Be. They are quite different! Be H + Unlike He, Be has an unoccupied 2p orbital it can use to its advantage.

BeH Potential Energy Curve and Orbitals
Viewing Be as a 2s2 valence configuration is too simplistic. Be mixes 2s2 and 2pz2 character to form lobe-like orbitals on either side of the nucleus: 2s2 2p0 2s2+l2p2 2s2–l2p2

At large R, the pair that is coupled together is on Be, and H has its normal 1s orbital. Be H electron pair singly occupied

In BeH, the coupled pair is a bond pair between the inner Be orbital and H, and the outer Be orbital is leftover. Be H electron pair singly occupied

Here’s the animation. Note how much rearrangement happens. Be H electron pair singly occupied

Recoupled Pair Bonding
When a singly occupied orbital is able to break apart an existing electron pair, a recoupled pair bond is formed. This is a conditional form of bonding, since there are many cases (such as He+H and Ne+H) where it doesn’t occur.

Old Representations for H through Ne
clear green: unoccupied yellow or blue: singly occupied gray: doubly occupied H He Li Be B C N O F Ne

Corrected Representations for Be, B, and C
clear green: unoccupied yellow or blue: singly occupied gray: doubly occupied H He Be B C Li N O F Ne

Building Molecules with More Than 2 Atoms
So far we’ve explored the ways in which two atoms chosen from the elements H through F can come together to form very simple molecules. Now we’ll explore what we can build if we start with one atom and keep adding other atoms to it.

Review – Bond Categories
We found two categories of bonds that could occur (there are more types!): 1  1 coupled pair bonds: covalent, polar covalent, ionic condition: opposite electron spins 1  2 recoupled pair bonds condition: reorganization must be favorable (not repulsive)

Consider An Oxygen Atom…
Let’s add H atoms to O atom: or

Nominal Valence & Nominal Bond Angles
It is straightforward to infer how many coupled pair bonds an atom can form: just count the unpaired electrons in its most stable atomic configuration. We call this the nominal valence of the element. The nominal bond angles of an atom are the angles between the singly occupied orbitals used to form the bonds.

17. Given the configuration diagram shown below, what is the nominal valence of O? 1 (b) 2 (c) 3 (d) Not enough information.

18. Given the configuration diagram for O shown below, what is the nominal bond angle of H2O? 180 (b) 109 (c) 90 (d) Not enough information.

18. Given the configuration diagram for O shown below, what is the nominal bond angle of H2O? 180 (b) 109 (c) 90 (d) Not enough information. Sulfur has a valence configuration of 3s2 3p4, analogous to oxygen, and thus also has a nominal valence of two. Like water, H2S has a nominal bond angle of 90.

Beyond the Nominal The nominal geometry of a molecule is only a starting point. We can determine the actual geometries of molecules with experiments or accurate calculations. We will soon compare the actual geometries of H2O and H2S to see how close they are to the nominal ones. Some elements, including Be, B, and C, can exceed their nominal valence by undergoing recoupled pair bonding.

The s bond is polarized toward O.
Adding H to O: First H Here is an animation of the s bond forming in OH from the H 1s orbital and one of the singly occupied O 2p orbitals: The s bond is polarized toward O.

What Do You Think? OH forms if H(1s) interacts with one of the singly occupied 2p orbitals of OH, as long as the two electrons have opposite spins. 19. What do you think will happen if H interacts with the doubly occupied 2p orbital instead? Oxygen is more like F & Ne than Be, so this will be repulsive. Oxygen is more like Be than F & Ne, so recoupling can occur.

Unlike Be, O doesn’t have an unused 2p orbital.
What Do You Think? OH forms if H(1s) interacts with one of the singly occupied 2p orbitals of OH, as long as the two electrons have opposite spins. 19. What do you think will happen if H interacts with the doubly occupied 2p orbital instead? Oxygen is more like F & Ne than Be, so this will be repulsive. Oxygen is more like Be than F & Ne, so recoupling can occur. Unlike Be, O doesn’t have an unused 2p orbital.

Adding H to OH: Second H OH(2p) H(1s) first bond pair

Adding H to OH: Second H formation of second bond pair
first bond pair is largely unperturbed

Adding H to H2O: Third H? We saw previously that H can’t recouple the 2p2 pair on O. We can’t form H3O by adding a third H to H2O because it would require recoupling the 2p2 pair on O.

Bonding Diagrams – 2D + O H OH H2O H + OH

Lewis Structures from Bonding Diagrams
+ O H O H OH (by inspection) O H + OH H H2O

Bonding Diagrams – 3D O

Bonding Diagrams – 3D O H

Bonding Diagrams – 3D OH O H

Bonding Diagrams – 3D H H2O O H

H2O vs. H2S H2O and H2S both have nominal bond angles of 90°. Here are the actual structures of water and hydrogen sulfide: 104.2 92.3 O S Note that the bond angle of H2S is very close to the nominal bond angle of 90. The bond angle of water is larger because there is greater repulsion between the bonds and between the H atoms than there is in H2S.

Test Your Understanding
20. The O atom and the OH radical are far more reactive than H2O. Why? (a) Neither O nor OH have an octet in their outer shells, but H2O does. (b) O and OH both have singly occupied orbitals that can easily form new bonds, while H2O has no singly occupied orbitals.

21. Based on its 2s2 2px 2py configuration, what is the nominal valence of carbon? (a) 1 (b) 2 (c) 3 (d) 4

21. Based on its 2s2 2px 2py configuration, what is the nominal valence of carbon? (a) 1 (b) 2 (c) 3 (d) 4 On the basis of the number of unpaired electrons, C would only be expected to form two coupled pair bonds.

21. Based on its 2s2 2px 2py configuration, what is the nominal valence of carbon? (a) 1 (b) 2 (c) 3 (d) 4 However, we’ll see that C can undergo recoupled pair bonding of the 2s2 pair like Be does, so its real valence is greater than two. These are better representations for carbon.

three unpaired electrons
Carbon Compounds: CH There are two ways H can approach C. It can form a covalent bond in this configuration: one unpaired electron + However, we can rotate C so that the H approaches the 2s2 pair: then it behaves like Be, and the 2s2 pair is recoupled: three unpaired electrons +

Carbon Compounds: CH Potential Energy Curves
The recoupled pair bond state of CH is almost as stable as the covalently bonded state. CH 1st bond: (kcal/mol) 1st bond:

The state of CH2 with recoupled pair bonding is more stable!
Carbon Compounds: CH + H  CH2 We can add H to each of these forms of CH. In the first case, we get this bonding diagram and optimized structure: 102 1st bond: 82.4 2nd bond: 99.5 Net: 181.9 (kcal/mol) If we add a second H to the form of CH with three unpaired electrons, we get this structure: 134 1st bond: 66.0 2nd bond: 121.3 Net: 187.3 (kcal/mol) The state of CH2 with recoupled pair bonding is more stable!

Carbon Compounds: CH + H  CH2
We can use the following 2D bonding diagram for this form of CH2. Note that it has one unpaired electron in the plane of the screen and one perpendicular to the screen. This is the in-plane orbital. (The out-of-plane orbital is a 2p orbital.)

Carbon Compounds: CH2 + H  CH3
If we add a third H, it forms a covalent bond to the unpaired electron in the in-plane orbital to yield a trigonal planar structure. 120 This is the remaining singly occupied orbital.

Carbon Compounds: CH3 + H  CH4
When the 4th H adds to CH3, its bond pair pushes the other three bond pairs into a tetrahedral arrangement as the bond forms.

3D Bonding Diagrams for CHn Compounds
Start with a C atom.

Add H.

Uncouple the C pair.

Add H to form CH2. Recouple to C to form CH. Add H to form CH3. Add H to form CH4.

Three Bonding Motifs for Carbon Compounds
Carbon has four bonds in most stable organic compounds. It can adopt three distinct motifs.

Sometimes C is found in a linear configuration connected to two other atoms. Here is one example, acetylene (C2H2): +

Sometimes C is found in a linear configuration connected to two other atoms. Here is one example, acetylene (C2H2): + Another example is HCN, shown here with a 3D diagram: H C N

Another motif is trigonal planar, when C is connected to three atoms, as we saw with the methyl radical. Another example is ethylene, C2H4: +

The final motif adopted by carbon when it is bonded to four other entities is the tetrahedral structure found in methane and compounds with a methyl group, such as methanol:

Wrapping Up  Atomic Configuration Symbol Orbital Animations
O (2s2 2p4)   C (2s2 2p2)  

Wrapping Up   Atomic Configuration Symbol Bond Types
Coupled Pair Bonds Recoupled Pair Bonds  O (2s2 2p4)   C (2s2 2p2)  

nominal valence & bond angles (and beyond)
Wrapping Up Atomic Configuration Symbol Bond Types   O (2s2 2p4)   Coupled Pair Bonds   C (2s2 2p2)   Recoupled Pair Bonds nominal valence & bond angles (and beyond)  H2O OH O CH2 CH C CH3 CH4

nominal valence & bond angles (and beyond)
Wrapping Up Atomic Configuration Symbol Bond Types   O (2s2 2p4)   Coupled Pair Bonds   C (2s2 2p2)   Recoupled Pair Bonds nominal valence & bond angles (and beyond)  H2O OH O CH2 CH C CH3 CH4 most reactive species marked in RED

Thank you very much for participating!
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Fun with Atoms

One Last Question 22. This is a symbol for the 2s2 2px 2py 2pz configuration of N. What is the maximum number of coupled pair bonds N can form? (a) 1 (b) 2 (c) 3 (d) 6

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