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CCP 3 > ppt 2 > Flashcards

ppt 2 Flashcards

1
Q

Most of the mass is in the ______

A

nucleus

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2
Q

containing protons and neutrons

Surrounded by an electron cloud containing electrons in a large volume of space

A

nucleus

protons + neutrons

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3
Q

Most of the volume is in the _______

A

electron cloud

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4
Q

variation of an element that possesses the same atomic number but a different mass number

any element will always have the same number of protons and electrons

They will differ in the number of neutrons held by their respective nuclei.

A

isotopes

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5
Q

the exact location of the electron cannot be known; only the electron density can be known

A

Quantum Mechanical Model

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6
Q

the probability that the electron will be in a certain region of space (orbital) at a given instant

A

Electron density

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7
Q

Electrons do not circle around the nucleus in fixed orbits but rather in _______

Each _____ has a characteristic shape and size (energy)

A

orbitals

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8
Q

used to describe orbitals

A

QUANTUM NUMBERS

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9
Q

Principal

A
  • n
  • Size (energy level, shell)
  • positive integers (1,2,3,…)
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10
Q

Azimuthal/ Angular Momentum

A
  • l
  • Shape (subshell)
  • 0 to (n-1) / s, p, d, f
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11
Q

Magnetic

A
  • ml
  • Orientation (orbital)
  • (-l to +l)
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12
Q

Electron Spin

A
  • ms
  • Direction
  • +1/2,-1/2
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13
Q

each box represents ________

A

one orbital

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14
Q

each cluster of boxes represents _________

A

one subshell

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15
Q

each row represents ________

A

one shell

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16
Q

describes the orbitals occupied by the atom’s electrons when they are all in the available orbitals with the lowest energy

A

Ground-state electron configuration

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17
Q

states that an electron occupies orbitals in order from lowest energy to highest (increasing order)

A

Aufbau Principle

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18
Q

Every orbital in a sublevel is singly occupied before any orbital is doubly occupied.

All of the electrons in singly occupied orbitals have the same spin (to maximize total spin).

A

Hund’s Rule

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19
Q

no two electrons can have the same four electronic quantum numbers.

As an orbital can contain a maximum of only two electrons, the two electrons must have opposing spins.

This means if one electron is assigned as a spin up (+1/2) electron, the other electron must be spin-down (-1/2) electron.

A

Pauli’ Exclusion Principle

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20
Q

states that we cannot know both the position and speed of a particle, such as a photon or electron, with perfect accuracy;

the more we nail down the particle’s position, the less we know about its speed and vice versa

A

Heisenberg’s Uncertainty Principle

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21
Q

electrons below the outermost shell

A

Core electrons

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22
Q

electrons in the outermost shell

A

Valence electrons

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23
Q

an attractive force between two ions or between two atoms

A

CHEMICAL BOND

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24
Q

Why do chemical bonds form?

A

Because the compound that results is more stable and lower in energy than the separate atoms

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25
Q

How do chemical bonds form?

A

Octet rule

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26
Q

atom is most stable if its outer shell is either filled or contains eight electrons, and it has no electrons of higher energy

A

OCTET RULE

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27
Q

An atom will give up, accept, or share electrons in order to achieve a filled outer shell containing eight electrons EXCEPT _______ and ________

A

hydrogen and helium

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28
Q

Chemistry of main-group elements is governed by their tendency to take on the electron configuration of the nearest

A

noble gas

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29
Q

attractive forces between opposite charges

A

Electrostatic attraction

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30
Q
  • bond formed as a result of the electrostatic attraction between ions of opposite charge
  • bond formed from the transfer of electrons
  • usually formed from the reaction of metals with nonmetals
A

Ionic bond

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31
Q

compounds formed by ionic bonds

A

Ionic compound

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32
Q

bond formed as a result of sharing electrons between two nuclei

A

COVALENT BOND

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33
Q

compounds formed by covalent bonds

A

Molecular compounds

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34
Q

measure of the ability of an atom to pull the bonding electrons toward itself

A

ELECTRONEGATIVITY

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35
Q

a covalent bond between atoms with the same electronegativity

A

nonpolar covalent bond

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36
Q

covalent bond between atoms with different electronegativities

A

polar covalent bond

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37
Q

bond where no electrons are shared; opposite charges attract each other

A

ionic bond

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38
Q

bond where electrons are shared equally

A

nonpolar covalent bond

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39
Q

Polar molecules have _______

A

dipoles

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40
Q

a pair of equal and oppositely charged poles separated by a distance

A

Dipole

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41
Q
  • measure of dipole
  • magnitude of the charge on either atom x distance between the two charges
A

Dipole Moment

42
Q

electron-dot structures

A

Lewis Structures

43
Q

line-bond structures

A

Kekulé Structures

44
Q

DRAWING CHEMICAL STRUCTURES

A
  • Chemical symbols
  • Covalent bond
  • Nonbonding/lone-pair electrons
  • Formal charges
45
Q
  • not an actual charge
  • used for bookkeeping of electrons
  • the charge the atom would have if each bonding electron pair in the molecule were shared equally between atoms
  • the difference between the number of valence electrons an atom has when it is not bonded to any other atoms and the number it “owns” when it is bonded
A

Formal Charge

46
Q

formal charge equation

A

formal charge = no of valence electrons in free atom - no of valence electrons in bonded atoms

47
Q

omitting of the covalent bonds and listing
atoms bonded to a particular carbon (or nitrogen or oxygen) next to it (with a subscript if there is more than one of a particular atom)

A

Condensed Structures

48
Q
  • Carbon atoms aren’t usually shown.
  • A carbon atom is assumed to be at each intersection of two lines (bonds) and at the end of each line.
  • Hydrogen atoms bonded to carbon aren’t shown.
  • Atoms other than carbon and hydrogen are shown
A

Skeletal Structures

49
Q

3D Structures

A

Perspective Drawing
Ball-and-Stick Model
Space-Filling Model

50
Q

shows 3D shape

A

Perspective Drawing

51
Q

shows bond angles accurately

A

Ball-and-Stick Model

52
Q

shows atoms in scale

A

Space-Filling Model

53
Q

bond lies in the plane of the paper

A

normal bond

54
Q

bond extends backwards, away from the viewer; “into” the paper

bond behind a page

A

dashed bond

55
Q

bond protrudes forwards, towards the viewer; bond out of page; “out of” the paper

A

wedged bond

56
Q
  • Geometry based on arrangement of atoms in a molecule
  • Defined by bond angles
A

MOLECULAR GEOMETRY

57
Q
  • Geometry based on valence electron pairs (bonding and non
    bonding) around a central atom
    • Bonding electron pairs
    • Nonbonding/lone electron pairs
  • Defined by bond angles
A

ELECTRON PAIR GEOMETRY

58
Q

A model for the prediction of molecular geometry based on the minimization of electron repulsion between regions of electron density around an atom

A

VALENCE-SHELL ELECTRON-PAIR
REPULSION (VSEPR) MODEL

59
Q

The best arrangement of a given number of electron pairs (bonding and nonbonding) is the one that

A

minimizes the repulsions among them.

60
Q

Without lone pairs:
electron pair geometry __ molecular geometry

A

=
equal

61
Q

assumes that the electrons in a molecule occupy overlapping atomic orbitals of the individual atoms

A

Valence Bond theory

62
Q

assumes the formation of molecular orbitals from the atomic orbitals

A

Molecular Orbital Theory

63
Q

formed when atomic orbitals on neighboring atoms overlap one another

A

Covalent bonds

64
Q

head-on overlap of atomic orbitals; stronger

A

σ

65
Q

sideway overlap of atomic orbitals; weaker

A

π

66
Q

atomic orbitals obtained when two or more nonequivalent orbitals of the same atom combine for covalent bond formation

A

Hybrid orbitals

67
Q

Produces the molecular geometry of the molecule.

A

HYBRIDIZATION

68
Q

one 2s and three 2p orbitals into four sp3 orbitals

four sp3 orbitals are 109.5° away from each other → tetrahedral geometry

A

sp3 Hybrid Orbitals: CH4 (Methane)

69
Q

one 2s and two 2p orbitals into three sp2 orbitals + one unhybridized p orbital

three sp2 orbitals are 120° away from each other: trigonal planar geometry + unhybridized p orbital perpendicular to it

A

sp2 Hybrid Orbitals: C2H4 (Ethylene)

70
Q

one 2s and one 2p orbitals into two sp orbitals + two unhybridized p orbital

two sp orbitals are 180° from each other → linear geometry + two unhybridized p orbitals mutually perpendicular to it

A

sp Hybrid Orbitals: C2H2 (Acetylene)

71
Q

bond strength _____ as bond length ______

A

increases, decreases

72
Q

As s character increases, bond angle ________

A

increases

73
Q

three sp3 orbitals w/ bonding pairs + one sp3 orbital w/ lone pair (trigonal pyramidal)

A

N

74
Q

two sp3 orbitals w/ bonding pairs + two sp3 orbitals w/ lone pairs (bent)

A

O

75
Q

Single bonds (no π bonds): sp3 w/ lone pairs in the hybrid orbitals

A

N & O

76
Q

one sp3 orbital with bonding pair + three sp3 orbitals w/ lone pairs

A

X

77
Q

describes a region of space in a molecule
where electrons are most likely to be found

A

molecular orbital

78
Q

2 ways to combine atomic orbitals into molecular orbitals

A

additive and subtractive

79
Q

lower energy, bonding molecular orbital

A

Additive

80
Q

higher energy, antibonding molecular orbital

A

Subtractive

81
Q

Occurs when an electron is shared by more than 2 atoms with π bonds (p orbitals)

A

π ELECTRON DELOCALIZATION

82
Q

Structures are represented by resonance contributors

A

π ELECTRON DELOCALIZATION

83
Q

Some substances can’t be represented by a single line-bond structure and must be considered as a _______ of two or more structures.

A

resonance hybrid

84
Q

The two individual line-bond structures are called ________, and their special resonance relationship is indicated by the double headed arrow between them.

A

resonance forms

85
Q

It has a single unchanging structure that is a resonance hybrid of the two individual forms and has characteristics of both.

A

resonance

86
Q

Individual resonance forms are ______

A

imaginary, not real

87
Q

Resonance forms obey normal rules of _______

A

valency

88
Q

The _________ is more stable than any individual resonance form

A

resonance hybrid

89
Q

resonance leads to

A

stability

90
Q

The larger the number of resonance forms, the ____ stable a substance is

A

more

91
Q

vacant p orbital, single electron, or lone pair of electrons on atom Z (resonance forms)

A

asterisk *

92
Q

from electron-rich species (negative charge, lone pair, double bond)

A

tail

93
Q

to electron-deficient species (positive charge, electronegative atom)

A

head

94
Q

between resonance forms

A

double headed arrow

95
Q

movement of electrons

movement of two electrons
movement of one electron

A

curly arrow

curly arrow (full)
curly arrow (half)

96
Q

Double bonds separated by single bonds

Refers to a system that has a p orbital adjacent to a π bond allowing delocalization of π electrons

A

conjugation

97
Q

Compounds with conjugated double bonds are thermodynamically ____ stable (lower energy) than compounds with isolated double bonds (higher energy).

A

more

98
Q

Special kind of conjugation

Cyclic unsaturated molecules with unusual stability

Originally used to classify benzene and its derivatives because many of them have distinctive odors

A

AROMATICITY

99
Q

Six p orbitals align and overlap to form six molecular orbitals

A

benzene

100
Q

TWO CRITERIA FOR AROMATICITY

  1. It must have an uninterrupted cyclic cloud of p electrons (called a ____) above and below the plane of the molecule.
A

p cloud

101
Q

huckel’s rule for aromaticity

A

Hückel’s Rule: 4n + 2

102
Q

Not all _______ (carbon ring systems with alternating single and double
bonds) are aromatic.

A

annulenes

CCP 3 (8 decks)

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