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NEU Warner Fall 2010 End Chapter 1
These flashcards cover all main points made during lecture off the slides and from what Warner said was important
20
Organic Chemistry
Undergraduate 1
09/14/2010

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Cards

Term
Explain what the idea is behind the molecular orbitals
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

Term
Explain what happens when ethyne is formed in regards to a carbon's 1 s and 3 p orbitals
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

Term
What occurs when pi and sigma bonds are mixed together?
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

Term
When are atoms, ions and molecules considered to be isoelectronic with each other?
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)

Term
Why does a polar bond have a dipole moment?
Definition

One end of the polar bond is positive and the other is negative in relative charge

Term
What is the equation needed to determine the dipole moment of a bond?
Definition

Dipole moment (D) = e x d (when e = the magnitude of the charge and d = the distance between the charges)

Term
What factors determine the overall dipole moment?
Definition

The vector sum of magnitudes and the directions of all the dipoles factor into it

Term
What are the dissolving rules?
Definition

Like dissolves like. Therefore, polar compounds will dissolve polar compounds and nonpolar compounds will dissolve nonpolar compounds

Term
What are the definitions for Bronsted-Lowry acids and bases, respectively?
Definition
A BL acid will donate H+ ions whereas BL bases will accept the H+ ions
Term
What happens to the conjugate base as an acid gets weaker and weaker?
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

Term
What determines how stable an acid or a base is?
Definition

The stronger an acid or base is the more unstable it is

Term
What is the condition on which the equation pKa=pH dependent upon?
Definition

For that to work [HA] must equal [H-]

Term
What effect does electronegativity have on a substance's pH?
Definition

A compound is more acidic when it has a larger electronegativity.

Term
How does electron delocalization affect pH?
Definition
Since electron delocalization spreads electrons around it decreases pH.
Term
How do substituents (things bonded on in addition to the base atoms) affect pH?
Definition

The substituents increase pH because they induce electron withdrawal

Term
What is the definition of a Lewis acid and base (respectively)?
Definition

A Lewis acid will accept and electron pair whereas a Lewis base will donate that electron pair

Term
How do both more electronegative and more organic compound groups affect the Lewis-defined acidity of compounds with a metal?
Definition

More electronegative compounds will set the pH more towards neutral but more organic groups will speed up this process towards neutral even faster

Term
What happens when a halogen acid differs in size?
Definition

Its proton will attach to the larger of the 2 atoms (which is almost always the halogen)

Term
Which has greater impact on acidity: electronegativity or size?
Definition

Electronegativity. Size does factor into the acidity but it's the electronegativity that is the most important factor

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