Term
| Explain what the idea is behind the molecular orbitals |
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Definition
A molecular orbital is the total area in which an electron can be found. There's no mathematical calculations that we're doing yet to determine the chances of finding an e- in a particular spot but the MO is all of the area around a bonded atom in which it is possible to find an electron from that atom |
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Term
| Explain what happens when ethyne is formed in regards to a carbon's 1 s and 3 p orbitals |
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Definition
An s and a p orbital get hybridized into a sp orbital which creates 2 pi bonds and a sigma bond between the carbons and just one sigma bond between the carbons and hydrogens due to the sp-s overlap |
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Term
| What occurs when pi and sigma bonds are mixed together? |
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Definition
Hyperconjugation occurs. In this process the MO's are warped by each other and, while somewhat being intertwined, are actually fairly opposite each other |
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Term
| When are atoms, ions and molecules considered to be isoelectronic with each other? |
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Definition
They're considered isoelectric with each other when they have the same number of electrons and the same structure (see slide 31 chap 1 for examples) |
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Term
| Why does a polar bond have a dipole moment? |
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Definition
One end of the polar bond is positive and the other is negative in relative charge |
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Term
| What is the equation needed to determine the dipole moment of a bond? |
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Definition
Dipole moment (D) = e x d (when e = the magnitude of the charge and d = the distance between the charges) |
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Term
| What factors determine the overall dipole moment? |
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Definition
The vector sum of magnitudes and the directions of all the dipoles factor into it |
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Term
| What are the dissolving rules? |
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Definition
Like dissolves like. Therefore, polar compounds will dissolve polar compounds and nonpolar compounds will dissolve nonpolar compounds |
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Term
| What are the definitions for Bronsted-Lowry acids and bases, respectively? |
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Definition
| A BL acid will donate H+ ions whereas BL bases will accept the H+ ions |
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Term
| What happens to the conjugate base as an acid gets weaker and weaker? |
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Definition
The conjugate base will become stronger and stronger. Take water, for example. Its conjugate base is pure hydroxide and its conjugate acid is pure hydrogen ions |
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Term
| What determines how stable an acid or a base is? |
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Definition
The stronger an acid or base is the more unstable it is |
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Term
| What is the condition on which the equation pKa=pH dependent upon? |
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Definition
For that to work [HA] must equal [H-] |
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Term
| What effect does electronegativity have on a substance's pH? |
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Definition
A compound is more acidic when it has a larger electronegativity. |
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Term
| How does electron delocalization affect pH? |
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Definition
| Since electron delocalization spreads electrons around it decreases pH. |
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Term
| How do substituents (things bonded on in addition to the base atoms) affect pH? |
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Definition
The substituents increase pH because they induce electron withdrawal |
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Term
| What is the definition of a Lewis acid and base (respectively)? |
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Definition
A Lewis acid will accept and electron pair whereas a Lewis base will donate that electron pair |
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Term
| How do both more electronegative and more organic compound groups affect the Lewis-defined acidity of compounds with a metal? |
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Definition
More electronegative compounds will set the pH more towards neutral but more organic groups will speed up this process towards neutral even faster |
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Term
| What happens when a halogen acid differs in size? |
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Definition
Its proton will attach to the larger of the 2 atoms (which is almost always the halogen) |
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Term
| Which has greater impact on acidity: electronegativity or size? |
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Definition
Electronegativity. Size does factor into the acidity but it's the electronegativity that is the most important factor |
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