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Chem.200-11.Covalent
General Chemistry Ch. 11 - Theories of Covalent Bonding
| Question | Answer |
|---|---|
| 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. |
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