Organic Chemistry I: Chapter 1

Front 1

Why is life carbon-based?

List the 5 reasons.

Back 1

1) Tetravalent (i.e. four bonds which allows carbon-based molecules to have 3-D shape for life's building blocks)

2) C-C bonds are strong

3) C-X bonds are possible

4) Carbon makes double and triple bonds

5) Exchange reactions (e.g. AB + CD > AC + BD)

Front 2

Define "Atom"

Back 2

Atom:
The basis for organic molecules

Front 3

List the three types of atoms and its composition.

Back 3

1) Neutral:
p+ = e-

2) Cation:
p+ > e-

3) Anion
p+ < e-

Front 4

Define "orbital"

Back 4

Orbital:
The standing wave around nuclei where the e- can be found.

Front 5

Heisenberg Principle

Back 5

An e-''s location OR trajectory can be identified, but not both at the same time.

Front 6

Back 6

Draw an "s-orbital".

How many types are there?

Front 7

Back 7

Draw a "p-orbital"

How many types is/are there?

Front 8

List the "principal quantum numbers"

Back 8

Principle quantam numbers: n, l, ml, ms

Front 9

Define principle quantum number "n".

Back 9

Principle quantum number "n":

Tells you the size and energy of the e-.

Front 10

What does the principle quantum number "l" describe?

What is its equation?

Back 10

describes: shape of the orbital

l = 0,1,2....(n-1)

Front 11

Define "n".

Back 11

"n" = principal quantum #

Tells you the size and energy of the e-.

Front 12

Define "electron selection rules".

Back 12

Electron selection rules:

Determines what orbital an e- is in.

Front 13

How many types of d-orbitals are there?

Back 13

There are 5 types of d-orbitals.

Front 14

Define "l"

Back 14

Principal quantum "l":
Angular momentum.

Tells you the shape ofthe orbital.

Front 15

When "n" equals the following what does "l" equal? Identify what types of orbitals are in each.

n=1
n=2
n=3

Back 15

n=1 L = 0
"s" orbital

n=2 L= 0, 1
"s" and "p" orbitals

n=3 L=0, 1, 2
"s", "p", "d" orbitals

Front 16

Can "n" equal zero? Why/why not?

Back 16

Principal quantum "n" cannot equal zero because it describes the size and energy of the e-. Size and energy of an e- cannot equal zero.

Front 17

What does principal quantum "m(l)" describe?

What is its equation?

Back 17

Principal quantum "m(l)" describes the orientation {i.e. p(x), p(y), p(z)} of the e-.

equation: -L...0...+L

Front 18

When "m(L)" equals the following, what does each mean?

m(L) = -1
m(L) = 0
m(L) = +1

Back 18

"m(L) = -1" means the e- can be found on the "p(x)" axis.

"m(L) = 0" means the e- can be found on the "p(y)" axis.

"m(L) = +1" means the e- can be found on the "p(z)" axis.

Front 19

What does it mean when "m(L)" equals the following?

Back 19

"m(L) = -2, -1, 0, +1, +2" describes the location of the e- within the 5 different types of d-orbitals.

Front 20

Describe each:

1s
2s
2p
3s
3p
3d
3d

Back 20

"1s"
n=1, L=0, m(L) = 0, m(s) = -1/2 and +1/2

"2s"
n = 2, L=0, m(L) = -1, 0, +1, m(s) = -1/2 and +1/2

"3s"

Front 21

Pauli Exclusion Principal

Back 21

Pauli Exclusion Principal:
No two e-'s can have the same set of principal quantum numbers {i.e. n, l, m(L), m(s)}.

Front 22

What are the two atom bonding types? Describe each.

Back 22

Atom bonding types:

1) Ionic:
Complete charge separation of atoms. (e.g. Na+ and Cl-) The electrostatic attraction between the positive and negative atoms hold the atoms together.

2) Covalent:
Sharing of e-'s.
2)

Front 23

Define "formal charge".

Back 23

"Formal charge":
The observed charge on each atom within the molecule.

Front 24

How do you calculate "formal charge"?

Back 24

(# of valence e-'s) - ({# of bonded e-'s} / 2) - (# of e-'s within lone pairs)

Front 25

Go online to practice calculating "formal charges" until it becomes second nature.

Back 25

..

Front 26

Define "electronegativity".

What is its trend in the periodic table?

Back 26

Electronegativity:
The degree by which an atom wants to attract e-'s.

Electronegativity trend:
1) Diagonally toward Fluorine: increasing electronegativity
2) Upward: Increasing
3) Downward: Decreasing
4) Away from fluorine: Decreasing

Front 27

Fill in the blanks:

____(1)_____ is important for breaking and formation of _____(2)____ because it leads to ___3___ and ____(4)___ .

Back 27

____(1, Electronegativity)_____ is important for breaking and formation of _____(2, covalent bonds)____ because it leads to ___(3, polarity)___ and ____(4, dipole moments)___ .

Front 28

What is the symbol for dipole moments?

Back 28

Greek letter "mu".

Front 29

Define "hybridization".

Back 29

Hybridization:
A phenomenon observed why leads to a molecule's geometry.

Front 30

Why do atoms want to hybridize?

Back 30

To lower energy.

Front 31

How many sigma bonds are there?

What types of orbitals are in each?

Back 31

Types of sigma bonds:
3 types

Orbitals found in each type of sigma bond:
1) s+s
2) s+sp
3) sp+sp

Front 32

How many types of pi bonds are there?

What types of orbitals are in each?

Back 32

Number of pi bonds:
There is only one type of pi bond.

Orbitals within a pi bond:
p + p

Front 33

Find hybridization problems online and practice until they become second nature.

Back 33

..

Front 34

What is a function of antibonding orbitals?

Back 34

Antibonding orbitals help lock the geometry of a molecule in place.

Front 35

Back 35

Explicitly draw each type of pi bond.

Front 36

Back 36

Explicitly draw each type of sigma bond.

Front 37

Back 37

Draw each of the d-orbitals.

Front 38

Back 38

Draw all of the p-orbital orientations.