O Chem Exam 1 Raheim

Front 1

Organic compounds are very complex and are only obtained from what?

Back 1

Living sources (vitalism)

Front 2

What is vitalism?

Back 2

Belief that a "magic" vital force, present in plants and animals, is necessary for synthesis of organic compounds

Front 3

Who observed that organic compounds are composed primarily of carbon and hydrogen?

Back 3

Antoine Laurent Lavoisier

Front 4

What organic compound was synthesized from and inorganic compound in 1828? Who did it?

Back 4

Urea; Friedrich Wohler

Front 5

Modern organic chemistry is the study of __________ compounds.

Back 5

carbon

Front 6

The study and practice of O chem can be subdivided into three areas:

Back 6

1. Structure

2. Mechanism

3. Synthesis

Front 7

Why is Carbon the main element studied in organic chem?

Back 7

Carbon forms a variety of strong covalent bonds to itself and other atoms (O, N, P, S , F, Cl, Br, I)

Carbon can bond to itself in many different configurations, including chains, rings, and branched structures of varying size and complexity

Front 8

What is the nucleus made up of?

Back 8

Protons and Neutrons

Front 9

Protons have a ___________ charge

Back 9

positive

Front 10

Neutrons are negatively charged. TRUE OR FALSE

Back 10

FALSE; no charge

Front 11

Electrons are ___________ charged and are in __________.

Back 11

negatively; motion

Front 12

Where does all the mass of an atom come from?

Back 12

The nucleus

Front 13

Atomic number (Z) = ?

Back 13

# of protons in nucleus, # of electrons surrounding nucleus

Front 14

Electrons surrounding the nucleus exist in what?

Back 14

Shells of increasing energy and at increasing distances from the nucleus

Front 15

Define valence electrons

Back 15

Exist in the valence shell, and are the outermost electrons that are used in making chemical bonds

Front 16

Electron cloud is about ________ times the diameter of its nucleus.

Back 16

10,000

Front 17

"An atom is most stable if its outer shell is either filled or contains 8 electrons, and it has no electrons of higher energy," What is this?

Back 17

Lewis bonding theory

Front 18

Group # = # of ?

Back 18

Valence electrons

Front 19

Define the octet rule

Back 19

An atom will give up, accept, or share electrons in order to achieve a filled outer shell o ran outer shell that contains 8 electrons

Front 20

Define Electronegativity

Back 20

a measure of an atom's attraction for electrons that it shares in a chemical bond with another atom

Front 21

How does electronegativity increase across the periodic table?

Back 21

Left to right

Bottom to top

Front 22

Electrons are completely transferred from one atom to another. What is this?

Back 22

Ionic bond

Front 23

Define covalent bonds

Back 23

electrons are shared between two atoms

Front 24

EN between bonded atoms >1.9

Back 24

Ionic

Front 25

EN between bonded atoms 0.5 to 1.9

Back 25

Polar covalent

Front 26

EN between bonded atoms <0.5

Back 26

Nonpolar covalent

Front 27

Atoms of widely different electronegativity

(EN>2); usually a metal and a nonmetal, what is a result?

Back 27

Formation of ions, a cation(+) and an anion (-)

Front 28

What are atoms held together by?

Back 28

Electrostatic attraction

Front 29

Covalent bonding usually occurs between atoms of similar what? What are they usually?

Back 29

Electronegativity; non-metallic

Front 30

Lewis dot structure vs Kekule drawing

Back 30

Lewis dot includes dots whereas Kekule includes lines

2 dots = 1 line

Front 31

Define polar covalent

Back 31

Unequal sharing of electrons

Front 32

How does an atom get a partial negative charge?

Back 32

The more electronegative atom gains a greater fraction of the shared electrons and acquires a partial negative charge.

Front 33

How does an atom get a partial positive charge?

Back 33

The less electronegative atom has a smaller fraction of the shared electrons and acquires a partial positive charge

Front 34

In CCl4, what happens?

Back 34

Each Cl has three lone pairs of electrons

Front 35

Define lone pairs

Back 35

unshared pair of electrons; non-bonding pair of electrons

Front 36

What is the HONC-1234 Rule?

Back 36

Hydrogen and Halogens make 1 bond

Oxygen makes 2 bonds

Nitrogen makes 3 bonds

Carbon makes 4 bonds

Front 37

Ethylene

Back 37

C2H4

Front 38

Double bond:

Back 38

two pairs of electrons

Front 39

Triple bond:

Back 39

three pairs of electrons

Front 40

Acetylene

Back 40

C2H2

Front 41

Hydrogen Cyanide

Back 41

HCN

Front 42

CH4

Back 42

Methane

Front 43

What charges will you ONLY see in second row atoms?

Back 43

1, 0, and -1

Front 44

Formal charges help chemists do what?

Back 44

To keep track of the placement of electrons in molecules

Front 45

How do you determine the formal charges?

Back 45

(Group #) - (# of nonbonding electrons) - 1/2(# shared electrons)

Front 46

PRACTICE DOING FORMAL CHARGES

Back 46

PRACTICE

Front 47

For line angle formulas, how should carbon atoms in a straight chain be drawn?

Back 47

In a zigzag formation

Front 48

You have to write Hydrogens in line angle formula when they are bonded to heteroatoms. TRUE OR FALSE

Back 48

TRUE

Front 49

When parentheses are used in condensed formulas, what does that mean?

Back 49

The part of the formula in parentheses is directly attached to the Carbon atom previously listed.

Front 50

Get a piece of paper for practice:

Convert following condensed formula to a line angle formula

(CH3CH2)2CCHCH2CHOHCH3

Back 50

YOU GOT THIS

Front 51

What are the angle in a tetrahedral formation?

Back 51

109.5 degrees

Front 52

What do lines mean when you are talking about the three dimensional shape of an organic molecule?

Back 52

Lines are in the plane of the paper

Front 53

What do dashes mean when you are talking about the three dimensional shape of an organic molecule?

Back 53

The dashes mean the atoms are going back into the paper (away from you)

Front 54

What do wedges mean when you are talking about the three dimensional shape of an organic molecule?

Back 54

The wedges mean the atoms are coming out of the paper (toward you)

Front 55

Each ____________________ corresponds to a different energy state for an electron.

Back 55

wave function

Front 56

What is the symbol for wave function? (What does it look like)

Back 56

A small pitch fork

Front 57

Why do we discuss wave function in organic chemistry?

Back 57

Moving particles exhibit the properties of a wave

Front 58

How can the relative probability of finding an electron in a given region of space be calculated?

Back 58

Wave function

Front 59

What does the phase sign of a wave function indicate?

Back 59

Whether the solution is positive or negative when calculated for a give point in space relative to the nucleus

Front 60

Each energy state is a sublevel where _____________ electrons can reside.

Back 60

one or two

Front 61

When does constructive interference occur?

Back 61

When wave functions with the same phase sign interact. There is a reinforcing effect and the amplitude of the wave function increases.

Front 62

When wave functions with opposite phase signs interact. What is this describing? What kind of effect does it have?

Back 62

Destructive interference; there is subtractive effect and the amplitude of the wave function goes to zero or changes sign

Front 63

What is the symbol for wavelenth? Symbol for amplitude?

Back 63

Looks like an upside down v with a tail; amplitude = a

Front 64

Back 64

Study this image

Front 65

A region of space where the probability of finding an electron is high. What is this defining?

Back 65

Orbital

Front 66

Electrons are waves! So how do they exist?

Back 66

As 3-D standing waves (orbitals)

Front 67

What are atomic orbitals?

Back 67

Unhybridized orbitals on an atom (s,p, d, f)

Front 68

What is the difference between a hybird atomic orbital and a molecular orbital?

Back 68

A hybrid atomic is a combination of atomic orbitals from the same atom whereas a molecular orbital is comes from different atoms

Front 69

What atomic orbitals are the most important in organic compounds?

Back 69

the s and p orbitals

Front 70

What happens when you add orbitals together? What is the called?

Back 70

You always end up with the same # of orbitals you started with; conservation of orbitals

Front 71

Individual wave functions (orbitals) combine to form ______________ and ____________________.

Back 71

Hybrid atomic orbitals and molecular orbitals

Front 72

Name the types of hybrid atomic orbitals

Back 72

sp3, sp2, and sp

Front 73

Name the types of molecular orbitals

Back 73

sigma, sigma*, pi, pi*

Front 74

Describe s orbitals

Back 74

spherical, electrons held close the nucleus, one sign

Front 75

Describe p orbitals

Back 75

two lobes with opposite signs, electrons further from nucleus. There are three degenerate (orbitals of equal energy) 2p orbitals that are higher in energy than the 2s orbital

Front 76

The sign of the orbital indicates the charge. TRUE OR FALSE

Back 76

FALSE; it does not indicate the sign of the charge

Front 77

Why is the Hund's Rule important?

Back 77

Partially filling orbitals as much as possible minimizes electrostatic repulsion between electrons

Front 78

In molecular orbitals, what happens if the electrons get to be too close together?

Back 78

Nuclear repulsion

Front 79

What is the Aufbau principle?

Back 79

Orbitals fill in order of increasing energy from lowest to highest

Front 80

Only two electrons can occupy an orbital an their spins must be paired. What is this defining?

Back 80

Pauli exclusion principle

Front 81

A _____________________ results when two orbitals of the same phase overlap.

Back 81

Bonding molecular orbital

Front 82

What happens when two orbitals of opposite phase overlap?

Back 82

Antibonding molecular orbital

Front 83

For bonding, where is the electron density centered?

Back 83

Between nuclei

Front 84

For antibonding, what is generally between nuclei?

Back 84

Node

Front 85

What is a node?

Back 85

Area of zero electron density

Front 86

when there is a sigma bond, where is the electron density centered?

Back 86

Along the axis of the bond

Front 87

What is the simplest sigma bond

Back 87

H2

Front 88

Which is more stable and lower in energy? less stable and higher in energy?

Back 88

Sigma (bonding) more stable and lower in energy

Sigma* (antibonding) less stable and higher in energy

Front 89

How many nodes are in a sigma bond between an s and a p orbital

Back 89

2 nodes

Front 90

Pi bonding orbitals are cylindrically symmetrical and sigma bonding orbitals are not. TRUE OR FALSE.

Back 90

FALSE; Pi bonding orbitals are NOT cylindrically symmetrical whereas sigma bonding orbitals are.

Front 91

Where is electron density located for a pi bond?

Back 91

Above and below the axis of the bond

Front 92

What bonds constitute as pi bonds?

Back 92

Double and triple bonds

Front 93

Orbitals must have the correct ___________ to overlap.

Back 93

symmetry

Front 94

What types of orbitals do not overlap? Why?

Back 94

Orthogonal; because they cannot have bonding interaction

Front 95

Di-substituted = ? (? degrees)

Back 95

linear; 180 degrees

Front 96

Tri-substituted = ? (? degrees)

Back 96

trigonal planar; 120 degrees

Front 97

Tetra-substituted = ? (? degrees)

Back 97

tetrahedral; 109 degrees

Front 98

What is the Vesper theory?

Back 98

Getting all electrons as far apart as possible

Front 99

What shape is this

Back 99

Tetrahedral

Front 100

Shape?

Back 100

Trigonal planar

Front 101

Shape?

Back 101

Linear

Front 102

What theory did Linus Pauling come up with that combined VSEPR with quantum mechanics (orbitals)?

Back 102

HYBRIDIZATION

Front 103

Atomic orbitals on the same atom combine to form hybrid atomic orbitals. WHY?

Back 103

Hybrid orbitals are more directional, so they have more effective bonding interactions

Front 104

When looking at an sp3 hybridization (tetrahedral), what are the large orbitals vs the smaller orbitals?

Back 104

Large are bonding orbitals whereas small are antibonding orbitals

Front 105

Electrons in pi bonds are always __________ in energy.

Back 105

higher

Front 106

Sigma bonds are always made from what?

Back 106

Hybridized orbitals

Front 107

Pi bonds involved ______________ p-orbitals

Back 107

unhybridized

Front 108

How many pi bonds are in a triple bond?

Back 108

two

Front 109

How many pi bonds are in a double bond?

Back 109

two

Front 110

How do you calculate the number of hybrid orbitals?

Back 110

# of sigma bonds + # of lone pairs

Front 111

Which orbital is the longest?

Back 111

sp orbital

Front 112

Which orbital is the shortest?

Back 112

sp3 orbital

Front 113

Sigma bonds rotate freely. TRUE OR FALSE

Back 113

TRUE

Front 114

Pi bonds can rotate like sigma bonds. TRUE OR FALSE

Back 114

FALSE; cannot rotate

Front 115

How do we represent electrons in motion?

Back 115

Curved arrow formalism (arrow pushing)

Front 116

Double arrow:

Back 116

2 electrons moving

Front 117

Fishhook arrow:

Back 117

1 electron moving