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Ch 1 - Chem Review Flashcards

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1
Q

Organic compounds contain

A

carbon atoms

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2
Q

Constitutional isomers

A

share the same molecular formula but have different connectivity of atoms and different physical properties.

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3
Q

Each element will generally form a predictable number of bonds.

A

Carbon is generally tetravalent, nitrogen trivalent, oxygen divalent, and hydrogen and the halogens monovalent.

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4
Q

covalent bond

A

results when two atoms share a pair of electrons.

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5
Q

Covalent bonds are illustrated using Lewis structures,

A

in which electrons are represented by dots.

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6
Q

Second-row elements generally obey the octet rule,

A

bonding to achieve noble gas electron configuration.

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7
Q

lone pair.

A

A pair of unshared electrons

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8
Q

formal charge

A

occurs when atoms do not exhibit the appropriate number of valence electrons; formal charges must be drawn in Lewis structures.

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9
Q

Bonds are classified as

A

(1) covalent, (2) polar covalent, or (3) ionic.

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10
Q

Polar covalent bonds exhibit induction, causing the formation of partial positive charges and partial negative charges

A

Electrostatic potential maps present a visual illustration of partial charges.

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11
Q

Quantum mechanics

A

describes electrons in terms of their wavelike properties.

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12
Q

wave equation

A

describes the total energy of an electron when in the vicinity of a proton. Solutions to wave equations are called wavefunctions(W) , where W^2 represents the probability of finding an electron in a particular location.

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13
Q

Atomic orbitals

A

are represented visually by generating three-dimensional plots of W^2; nodes indicate that the value of W is zero.

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14
Q

electron density

A

An occupied orbital can be thought of as a cloud

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15
Q

Electrons fill orbitals following three principles:

A

(1) the Aufbau principle, (2) the Pauli exclusion principle, and (3) Hund’s rule. Orbitals with the same energy level are called degenerate orbitals.

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16
Q

Valence bond theory

A

treats every bond as the sharing of electron density between two atoms as a result of the constructive interference of their atomic orbitals. Sigma (σ) bonds are formed when the electron density is located primarily on the bond axis.

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17
Q

Molecular orbital theory

A

uses a mathematical method called the linear combination of atomic orbitals (LCAO) to form molecular orbitals. Each molecular orbital is associated with the entire molecule, rather than just two atoms.

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18
Q

The bonding MO of molecular hydrogen results from

A

constructive interference between its two atomic orbitals

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19
Q

The antibonding MO results from

A

destructive interference.

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20
Q

atomic orbital

A

is a region of space associated with an individual atom, while a molecular orbital is associated with an entire molecule.

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21
Q

Two molecular orbitals are the most important to consider:

A

(1) the highest occupied molecular orbital, or HOMO, and (2) the lowest unoccupied molecular orbital, or LUMO.

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22
Q

Methane’s tetrahedral geometry can be explained using

A

four degenerate sp^3-hybridized orbitals to achieve its four single bonds.

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23
Q

Ethylene’s planar geometry can be explained using

A

three degenerate sp2-hybridized orbitals. The remaining p orbitals overlap to form a separate bonding interaction, called a pi bond. The carbon atoms of ethylene are connected via a sigma bond, resulting from the overlap of sp2-hybridized atomic orbitals, and via a pie bond, resulting from the overlap of p orbitals, both of which comprise the double bond of ethylene.

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24
Q

Acetylene’s linear geometry is achieved via

A

sp-hybridized carbon atoms in which a triple bond is created from the bonding interactions of one sigma bond, resulting from overlapping sp orbitals, and two pie bonds, resulting from overlapping p orbitals.

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25
Triple bonds are stronger and shorter than
double bonds, which are stronger and shorter than single bonds.
26
The geometry of small compounds can be predicted using valence shell electron pair repulsion (VSEPR) theory, which focuses on
sigma bonds and lone pairs exhibited by each atom. The total, called the steric number, indicates the number of electron pairs that repel each other.
27
A tetrahedral arrangement of orbitals indicates
sp^3 hybridization (steric number 4). A compound's geometry depends on the number of lone pairs and can be tetrahedral, trigonal pyramidal, or bent.
28
A trigonal planar arrangement of orbitals indicates
sp^2 hybridization (steric number 3); however, the geometry may be bent, depending on the number of lone pairs.
29
Linear geometry indicates
sp hybridization (steric number 2)
30
Dipole moments (U)
occur when the center of negative charge and the center of positive charge are separated from one another by a certain distance; the dipole moment is used as an indicator of polarity (measured in debyes).
31
The percent ionic character of a bond is determined by
measuring its dipole moment. The vector sum of individual dipole moments in a compound determines the molecular dipole moment.
32
The physical properties of compounds are determined by
intermolecular forces, the attractive forces between molecules.
33
Dipole-dipole interactions occur between two molecules that possess
permanent dipole moments. Hydrogen bonding, a special type of dipole-dipole interaction, occurs when the lone pairs of an electronegative atom interact with an electron-poor hydrogen atom. Compounds that exhibit hydrogen bonding have higher boiling points than similar compounds that lack hydrogen bonding.
34
London dispersion forces result from the interaction between
transient dipole moments and are stronger for larger alkanes due to their larger surface area and ability to accommodate more interactions.
35
Polar compounds are soluble in
polar solvents
36
nonpolar compounds are soluble in
nonpolar solvents.
37
Soaps are compounds that contain both
hydrophilic and hydrophobic regions. The hydrophobic tails surround nonpolar compounds, forming a water-soluble micelle.
38
Organic Chemistry
the study of carbon based molecules
39
constitutional isomers
same molecular formula but differ in the way the atoms are connected - different names and physical properties
40
each element will generally form a predictable number of bonds
- tetravalent(4 bonds) - trivalent(3 bonds) - divalent(2 bonds) - monovalent(1 bond)
41
covalent bond
two atoms sharing a pair of electrons
42
Lewis Structure
a drawing style in which electrons take center stage
43
Lewis structures are based on
the valence electrons pairing up to be shared between atoms
44
Identifying Formal Charge
2 steps - determine the appropriate number of valence electrons for an atom - determine whether the atom exhibits the appropriate number of electrons
45
formal charge =
valence – nonbonded-1/2(bonded) - N is 5A but has 4 electrons = + - O is 6A but has 7 electrons = -
46
three bond categories
- covalent(sharing) - polar covalent(uneven sharing) - ionic(electron transfer)
47
the bond type is based on
the electronegativity(measure of an atoms ability to attract electrons) difference of two atoms
48
General guideline for electronegativity(not definitive)
- 0.0-0.5 is covalent(C-C) - 0.5-1.7 is polar covalent bond(C-O) - 1.7+ is ionic(NaOH)
49
induction
the withdrawal of electron density that occurs when a bond is shared by to atoms of differing electronegativity
50
electrons are particles with
wavelike properties
51
quantum mechanics
a mathematical description of an electron that incorporates its wavelike properties
52
wave equation
describes the total energy of a hydrogen atom(one proton plus one electron - takes into account the wavelike behavior of an electron that is in the electric field of a proton
53
wave functions(W)
a series of solutions to the wave equation - each wavefunction corresponds to an allowed energy level for the electron - suggest electrons only exists at discrete energy levels(W1,W2,W3 etc) - energy of an electron is quantized
54
each wavefunction is a function of spatial location
- provides information that allows numerical value to be assigned - W^2 indicates the probability of finding the electron in that location - these are SPDF orbitals
55
orbital
a region of space that can be occupied by an electron | - remember: an electron is part particle and part wavelike
56
electron clouds only come in a
small number of shapes and sizes(defined by the orbitals)
57
an electron cloud is a
single entity
58
electron density
a term associated with the probability of finding an electron in a particular region of space
59
the “shape” of an orbital refers to a region of space that contains 90-95% of the electron density
- 100% density would literally be the entire universe
60
atomic orbital(AO)
a 3-D plot of W^2 of a wavefunction - region of space that can accommodate electron density - SPDF
61
for a wave function(W):
- Red means above average(+) - Blue means below average(-) - node – in atomic and molecular orbitals a location where W is zero - does NOT denote positive or negative charge - simply denotes the phase of a wave - a node has an electron density of zero
62
node
in atomic and molecular orbitals a location where W is zero
63
Orbitals are filled by electrons guided by 3 principles
- Aufbau principle – the lowest energy orbital is filled first - Pauli exclusion principle – each orbital can accommodate a maximum of two electrons that have opposite spin(up and down spin) - Hunds Rule – when dealing with degenerate orbitals, such as p orbitals, one electron is placed in each degenerate orbital first before electrons are paired up
64
Aufbau principle
the lowest energy orbital is filled first
65
Pauli exclusion principle
each orbital can accommodate a maximum of two electrons that have opposite spin(up and down spin)
66
Hunds Rule
when dealing with degenerate orbitals, such as p orbitals, one electron is placed in each degenerate orbital first before electrons are paired up
67
a covalent bond is formed from the overlap of
atomic orbitals
68
constructive interference
two waves create a wave with larger amplitude
69
destructive interference
two waves interferences cancels each other out producing a node
70
valence bond theory
a bond is simply the sharing of electron density between two atoms as a result of the constructive interference of their atomic orbitals
71
sigma bond
circular symmetry with respect to the bond axis
72
all single bonds are ____bonds(head on the axis)
sigma
73
molecular orbital(MO) theory
describes a bond in terms of the constructive interference between two overlapping atomic orbitals - linear combination of atomic orbitals(LCAO)
74
molecular orbitals
mathematically derived orbitals, hybrids
75
atomic orbits are a region of space associated with
the atom
76
molecular orbitals are an orbital associated with
the entire molecule
77
bonding MO
the lower energy MO filled first via constructive interference
78
antibonding MO
higher energy MO filled after bonding MO via destructive interference
79
important MOs
- highest occupied molecular orbital(HOMO) | - lowest unoccupied molecular orbital(LUMO)
80
dipole moment(U)
an indicator of polarity - U = (partialcharge)(d) - partial + and – are usually in order of 10^-10esu(electrostatic units) - U will generally have a magnitude of 10^-18 esu - 1 debye(D) = 10^-18 (esu)(cm) - %ionic is (actual D/theoretical 100% D)(100%)
81
molecular dipole moment
the vector sums of multiple dipole moments in a molecule to get the molecule dipole moment
82
stronger bonds =
higher boiling point
83
longer straight chain has more surface area =
higher boiling point
84
branching results in less surface area and
a lower boiling point
85
hydrophilic
water loving
86
hydrophobic
water hating
87
micelle
polar heads and nonpolar tails ball u and become water soluble(polar heads out)

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