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CHM136 > Introduction to Organic Compounds > Flashcards

Introduction to Organic Compounds Flashcards

Week 1+2

1
Q

What is Organic Chemistry?

A

Study of Molecules with Carbon

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2
Q

Properties of Carbon

A
  • Unique bonding properties
  • Froms Covalent bonds with a wide diversity of elements
  • Up to 4 covalent bonds
  • Can form chains, rings and complex molecules
  • Make be flat or 3D
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3
Q

Hydrocarbons

A

Organic compounds made up of ONLY carbon and Hydrogen

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4
Q

Alkanes

A

Formula; C(n)H(2n_2)
No double or triple bonds

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5
Q

Alkenes

A

C=C

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6
Q

Alkynes

A

C triple bond C

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7
Q

4 kinds of orbitals for electrons

A

s, p, d, f

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8
Q

s orbitals

A

Spherical
1s, 2s

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9
Q

Where are you most likely to find an electron in s orbitals

A

The probability of finding an electron is highest near the nucleus and decreases as the distance from the nucleus increases.

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10
Q

p orbitals

A

dumbbell shape; 2 lobes separated by a region of sero electron density (node).
lobes are either + or - and are in different spin states
2px, 2py and 2pz are all degenerates (equal in energy)

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11
Q

probabilty of finding electrons in the p orbital

A

Equal for both lobes, but no electron density in nodal plane.

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12
Q

degenerates

A

Electron orbitals having the same energy levels

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13
Q

Valence electrons are

A

Less stable and participate in bonding.

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14
Q

Ground configuration state

A

It is the DIAGRAM oh orbital levels

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15
Q

Why do bonds form?

A

Because moleucles have a lower energy than individual atoms.

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16
Q

How do bonds form?

A

Electrons are shared or transferred to fill the valence shell (octet rule)

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17
Q

Ionic bonding

A
  • electrostatic attraction between oppositely charged ions
  • bond forms from electron transfer
  • results in salts
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18
Q

Covalent bonding

A
  • sharing of electrons between atoms
  • each bonding atom usually contributes one electron to the new bond (opposite spin)
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19
Q

What are the two types of covalent bonding?

A

Sigma and Pi bonds

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20
Q

Sigma bonds

A

Head-on orbital overlap

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21
Q

pi bonds

A

side-to-side orbital overlap

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22
Q

electronegativity

A

the ability to pull electrons towards tiself through a bond
the arrow points to the more electronegative atom.
- leads to bond dipole

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23
Q

Lewis structure

A
  • covalent bonds (2 shared electrons) are represented by a line between atoms as a single bond (1 line; 2 electrons), double bond (2 lines; 4 electrons), or triple bond (3 lines; 6 electrons)
  • non-bonded electrons (i.e. lone pairs) are shown as dots.
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24
Q

Neutral states of important molecules

A

Carbon - no lone pairs
Nitrogen - 1 long pair
Oxygen - 2 long pairs

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25
Formal Charge
based on the bonding of an atom, includeing the number of shared (bonded) and non-shared (non-bonded) electrons to the number of valence electrons
26
Formal Charge formula
Formal Charge = (group #)-(# of bonds) - (# of non-bonded electrons)
27
Formal Charge sum
Sum of all formal charge must equal the charge of the molecule or ion.
28
VSEPR Theory acronym
Valence Shell Electron Pair Repulsion Theory
29
VSEPR Theory
The most stable structure is the one which valence electron pairs (bonding pair or lone pair of electrons) are as far apart as possible to minimize electron-electron repulsion.
30
Tetrahedral
- Four groups around carbon are equally distributed in a pyramid arrangement - Bond angle of about 109° between atoms
31
Trigonal Planar
- Three groups around carbon are equally distributed in the same plane (flat) - Bond angle of about 120° between atoms
32
Linear
- Two groups around carbon are equally distributed in a line. - Bond angle of 180° between atoms
33
Valence Bond Theory
Covalent bond involves the sharing of two spin-paired electrons through the overlap of atomic orbitals. Better overlap = stronger bond
34
What are the two theories that explain bonding
Valence bond theory and molecular orbital theory
35
What are hybrid orbitals
atomic orbitals that mix to give the observed geometry of atoms - They are ''hybrid'' meaning that they are found between the ground state electron configuration orbitals Ex. sp^3 = 1 s orbital + 3 p Obitals ## Footnote https://www.youtube.com/watch?v=pdJeQUd2g_4&t=13s 2:24
36
Energy level mixing means
the energy level is in between s and p
37
sp^3 hybrid orbitals
sp3 = 2s + 2p + 2p + 2p mix to from four equivalent tetrahedral orbitals with a bond angle of 109.5
38
sp^2 hybrid orbitals
sp2 = 2s + 2p + 2p mix to form three equivalent trigonal planar orbitals with 120 degree bond angles
39
Line-Bond Diagrams
Bonds are shown as Zig-Zags
40
delocalized pi bond
Sharing of 2 electrons among more than 2 atoms
41
Resonance
Sharing of electrons among many atoms (delocalization) stabilizes molecules.
42
Resonance forms have
- identical placement of atoms - Different placement of electrons (lone pairs and pi bonds) - A double-headed arrow separating them
43
Resonance Hybrid
Blend of all the resonance forms
44
For a molecule to have resonance it must have
- pi bond with atoms of different electronegativities - pi bond beside an atom with a p-orbital - Atoms with incomplete octets next to an atom with a long pair of electrons
45
Highest quality resonance forms have the following characteristics These are used to order them from most to least significant...not all resonance forms contribute equally to the hybrid
1. The most atoms with full octets 2. The fewest number of formal charges 3. If formal charge exists, the negative charges are on the most electronegative atoms while positive charges on the most electropositive atoms 4. like charges separated by the minimum distance possible, while opposite charges as close together as possible
46
Functional Groups
groups that give molecules distinct reactivity and properties.
47
Inductive effect
Distribution of electron density in a bond due to differences of electronegativity.
48
What determines bond polarity
Degree of electron sharing depends on differences in electronegativity of the elements involved in the bond.
49
An electrostatic Potential Map
Shows regions of positive (electron deficient) and negative (electron rich) charge distribution.
50
C-H bonds are polar or non-polar
Non-polar
51
Electrostatic Interaction
Oppositely charged molecules attract
52
Dipole-Dipole Interaction
Attractive forces between the negative end of a permanent dipole in a molecule and the positive end of a permanent dipole in a neighboring molecule (or vice versa)
53
Hydrogen Bonding
Attractive force between a N or O lone pair and a H in OH or NH
54
Dispersion Forces
Attractive interactions that exist between all molecules in close proximity to each other. - Important for non-polar molecules
55
Boiling points (bp) and melting points (mp) are a reflection of the relative strength of intermolecular forces. So What would have a higher melting or boiling point?
Polar groups impart higher mp/bp values due to increased dipole-dipole interactions (strong), hydrogen bonding (stronger), or electrostatic effects (strongest). Size of a molecule also increases mp/bp (bigger size = higher mp/bp)
56
What are the 3 types of solvents?
Polar protic solvents Polar aprotic solvent Non-polar solvent
57
Polar protic solvents
H-bond donors (OH, NH), very polar
58
Polar aprotic solvents
Strong dipoles (polar), most are H-bond acceptors
59
Non-polar solvents
Mainly non-polar bonds or no significant net dipole
60
Like-dissolves-like
Strongly polar solvents dissolve strongly polar/ionic substances; weakly polar solvents dissolve weakly polar/non-polar substances.
61
Constitutional Isomer
Two molecules which have the same molecular formula but different structural formulas, or bonding arrangements.

CHM136 (9 decks)

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