Busy. Please wait.

show password
Forgot Password?

Don't have an account?  Sign up 

Username is available taken
show password


Make sure to remember your password. If you forget it there is no way for StudyStack to send you a reset link. You would need to create a new account.

By signing up, I agree to StudyStack's Terms of Service and Privacy Policy.

Already a StudyStack user? Log In

Reset Password
Enter the associated with your account, and we'll email you a link to reset your password.

Remove ads
Don't know
remaining cards
To flip the current card, click it or press the Spacebar key.  To move the current card to one of the three colored boxes, click on the box.  You may also press the UP ARROW key to move the card to the "Know" box, the DOWN ARROW key to move the card to the "Don't know" box, or the RIGHT ARROW key to move the card to the Remaining box.  You may also click on the card displayed in any of the three boxes to bring that card back to the center.

Pass complete!

"Know" box contains:
Time elapsed:
restart all cards

Embed Code - If you would like this activity on your web page, copy the script below and paste it into your web page.

  Normal Size     Small Size show me how


General Chemistry Ch. 11 - Theories of Covalent Bonding

What does the VSEPR theory NOT explain? It doesn’t explain how the shapes come from the interactions of atomic orbitals. Moreover, it doesn’t explain the magnetic and spectral properties of molecules; only an understanding of their orbitals and energy levels can do that.
Valence bond theory (VB theory) A covalent bond forms when orbitals of two atoms overlap and the overlap region, which is between the nuclei, is occupied by a pair of electrons.
How should the spins of the overlapping orbitals behave? The electron pair must have opposing spins. The maximum capacity is two electrons.
The greater the overlap… The stronger (more stable) the bond. The extent of the overlap depends on the shapes and directions of the orbitals.
An s orbital is spherical, but p and d orbitals have more electron density in one direction than in another. How does this affect the orientation of the overlapping orbitals? A bond involving p or d orbitals will be oriented in the direction that maximized overlap. E.g. in HF, H is oriented along the long axis of the F’s 2p orbital (e.g. pg 411, Figure 11.1).
It’s hard to imagine four orbitals overlapping in CH4 while maintaining the ideal bond angle of 109.5. How does that work? The orbitals change. The process of orbital mixing is called hybridization, and the new atomic orbitals that are created are called hybrid orbitals.
Two key points about the number and the type of hybrid orbitals are that… (1) The number of hybrid orbitals obtained equals the number of atomic orbitals mixed, and (2) the type of hybrid orbitals obtained varies with the types of atomic orbitals mixed.
List 5 common types of Hybridization sp Hybridization, sp2 Hybridization, sp3 Hybridization, sp3 d Hybridization, sp3 d2 Hybridization
sp hybrid orbital An orbital formed by the mixing of one s and one p orbital of a central atom
How does VB theory explain sp Hybridization By proposing that mixing two nonequivalent orbitals of a central atom, one s and one p, gives rise to two equivalent sp hybrid orbitals that lie 180 degrees apart. Both shape and orientation maximize overlap with each other.
sp2 hybrid orbital An orbital formed by the mixing of one s and two p orbitals of a central atom
sp3 hybrid orbital An orbital formed by the mixing of one s and three p orbitals of a central atom
sp3 d hybrid orbital An orbital formed by the mixing of one s, three p, and one d orbital of a central atom
sp3 d2 hybrid orbital An orbital formed by the mixing of one s, three p, and two d orbitals of a central atom
“Planar” in VSEPR theory refers to The shape being essentially 2 dimensional
Trigonal bipyramidal is so named because It’s composed of two triangle based pyramids stacked on top of each other
Why is octahedral so named? Because it’s comprised of eight “hedras”, or hedrons, i.e. faces
Do elements use their unpaired electrons to make bonds? No, they transform their orbitals to new ones to do the bonding.
Electron domain The region of space around an atom occupied by a lone pair or a chemical bond to another atom (the bond may be single, double, or triple).
How to determine the hybridization of an element? Count the electron domains. 2 = sp, 3 = sp2, 4 = sp3, 5 = sp3d, 6 = sp3d2
All single bonds are ____ bonds Sigma
Sigma bonds When elements overlap in an end-to-end fashion. (They can be hybridized or s or p or whatever).
In a double or triple bond, the first bond is a sigma bond, the others must be ____ bonds Pi
Pi bonds (why are pi bonds so named?) When p orbitals overlap in a side-to-side fashion. (They’re so named because pi is the Greek letter for P, and only P orbitals—which are side by side—can create pi bonds)
When do pi bonds come into play? Double or triple bonds.
What is stronger, pi or sigma? Sigma
Does the concept of hybridization always apply to every situation? No. For some molecules the observation doesn’t support the theory, e.g. H2S.
In layman terms explain the process of the formation of a pi bond in ethylene Each C atom forms two hybridized sigma bonds with the H atoms (thus sp-s bonds) and one sigma bond with each other (sp-sp bond). The double bond forms because the remaining unhybridized p orbitals overlap.
A triple bond consists of … One sigma bond and two pi bonds.
How do sigma and pi bonds affect rotation of elements within a molecule? A sigma bond allows free rotation of parts of the molecule with respect to each other. A pi bond does not (the p orbitals overlap at two points around the molecule, locking it in place)
How do pi bonds allow for cis and trans structures to exist? Cis (on same side) vs trans (on opposite sides): initial config depends on how the molecule is initially formed. Since they’re locked in place they stay in that form for the remainder of their existence.
The theories used to explain molecular geometry, hybrid-orbital analysis, and magnetic and spectral properties Molecular geometry: VSEPR theory; hybrid-orbital analysis: VB theory; magnetic and spectral properties: molecular orbital (MO) theory
MO theory A model that describes a molecule as a collection of nuclei and electrons in which the electrons occupy orbitals that extend over the entire molecule.
How does MO theory contrast to VB theory In VB theory, a molecule is pictured as a group of atoms bound together through localized overlap of valence-shell atomic orbitals. In MO theory, a molecule is pictured as a collection of nuclei with delocalized electrons.
Just as an atom has atomic orbitals (AOs) with a given energy and shape that are occupied by the atom’s electrons, a molecule has _____ Molecular orbitals, with a given energy and shape that are occupied by the molecule’s electrons
The drawback of MO theory MOs are more difficult to visualize than the easily depicted shapes of VSEPR theory or the hybrid orbitals of VB theory
Roughly, how are properties of MOs derived? The most common approximation combines (adds/subtracts) the atomic orbitals (atomic wave functions derived from Schrodinger equation) of nearby atoms to form MOs (molecular wave functions)
Adding wave functions together (and analogy) This combinations forms a BONDING MO, which has a region of high electron density between the nuclei (analogy: light waves reinforcing each other, making the resulting amplitude higher and the light brighter).
Subtracting the wave functions from each other This combination forms an ANTIBONDING MO, which has a region of zero electron density (a node).
With electron waves, the probability that the electrons lie between the nuclei when an antibonding MO is formed, is _____ Decreasing to zero
The number of AOs combined always equals… The number of MOs formed.
The bonding MO is _____ in energy and the antibonding MO is _____ in energy than the AOs that combined to form them Lower; higher
Why is the bonding MO lower in energy than the AOs that combined to form them? The bonding MO is spread mostly between the nuclei. An electron in this MO can delocalize its charge over a much larger volume. Because the electron-electron repulsions are reduced, the bonding MO is lower & more stable.
Why is the antibonding MO higher in energy than the AOs that combined to form them? The antibonding MO has a node between the nuclei and most of its electron density outside the internuclear region. The electrons do not shield one nucleus from the other, which increases the nucleus-nucleus repulsion, less stable.
Both the bonding and antibonding MOs of H2 are _____ MOs Sigma MOs, because they are cylindrically symmetrical about an imaginary line that runs through the two nuclei. The lower energy (bonding) MO is sigma_1s, and the higher energy (antibonding) MO is sigma*_1s
To interact effectively and form MOs, atomic orbitals must have Similar energy and orientation.
MOs are filled in order of… (Principle?) Increasing energy (starts at lower energy then moves to higher energy). Aufbau principle.
An MO has a maximum capacity of two electrons with… (Principle?) Two electrons with OPPOSITE spins. Pauli exclusion principle
Orbitals of equal energy are…. (Principle?) Half filled, with spins parallel, before any of them is completely filled. Hund’s rule.
Molecular orbital (MO) diagram Shows the relative energy and number of electrons in each MO, as well as the AOs from which they formed.
MO Bond order The number of electrons in bonding MOs minus the number in antibonding MOs, divided by two
A bond order greater than zero indicates that… The molecular species is stable relative to the separate atoms, whereas a bond order of zero implies no net stability and, thus, no likelihood that the species will form.
Why does He2+ exist but not He2? Using MO theory the bond order of He2+ = 1/2 (meaning it’s a really weak bond, but still exists), however, the bond order for HE2 = 0, so a covalent form of He2 shouldn’t exist.
Recap: naturally occurring homonuclear diatomic molecules? (Have No Fear Of Ice Cold Beer) H2, N2, F2, O2, I2, Cl2, B2
Recap: standard physical states of all elements Two liquids: Hg & Br, 13 gases: the 6 inerts and 7 others: H2, N2, F2, O2, Cl2 (have no fear of cold), and ALL THE REST ARE SOLID.
Homonuclear diatomic molecules Molecules composed of two identical atoms.
When creating MO diagrams, which AOs are considered? Only valence orbitals interact enough to form molecular orbitals
Side-to-side combinations of AOs yield… pi MOs
MOs formed from 2s orbitals are ____ in energy than MOs formed from 2p orbitals. Why? Lower because 2s AOs are lower in energy than 2p AOs.
Energy order for MOs derives from 2p orbitals Sigma_2p < pi_2p < pi*_2p < sigma*_2p
For MOs when you’re overlapping p orbitals, you expect __ sigmas and __ pis 1 sigma, 2 pis
What levels should the sigma and pi p orbitals be in the MO diagram? You have to look them up, they’re different for two groups: (O2, F2, and Ne2) and (B2, C2, N2). Check pg 426 for clarification if needed. Not sure if need to memorize.
Recall: paramagnetic vs diamagnetic Paramagnetic: attracted to an external magnetic field due to having unpaired electrons. Diamagnetic: unaffected by magnetic field, no unpaired electrons.
Heteronuclear diatomic molecules Those composed of two different atoms. They have asymmetric MO diagrams because the atomic orbitals of the two atoms have unequal energies. (As opposed to homonuclear diatomic molecules)
Degenerate orbitals Orbitals at the same energy levels
Nonbonding MOs Orbitals not involved in bonding. They have the same energy as the isolated AOs.
Note: Study MOs with videos and other resources to fully understand and have intuition.
Created by: Intellex_