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CCP 3 LEC > Bonding > Flashcards

Bonding Flashcards

2- Structure and Bonding

1
Q

an attractive force between two ions or between two atoms

A

Chemical Bond

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2
Q

Why do chemical bonds form?

A

The compound resulting from this is more stable and lower in energy than the separate atoms

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3
Q

How do chemical bonds form?

A

Octet Rule

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4
Q

principle which states that atoms tend to gain, lose, or share electrons in order to achieve a full outer shell with eight electrons

except hydrogen and helium for the outer shell (1s) can occupy only 2

A

Octet Rule

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5
Q

An atom is most stable if its outer shell is either filled or contains ___, and it has NO electrons of higher energy.

A

eight electrons

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6
Q

Main-group elements in chemistry tend to adopt the electron configuration of the nearest ___.

A

noble gas

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7
Q

a chemical bond may be ___ and ___.

A

ionic and covalent

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8
Q

types of covalent bonds

A

polar and nonpolar

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9
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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10
Q

compounds formed by ionic bonds

A

Ionic Compound

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11
Q

attractive forces between opposite charges

A

Electrostatic Attraction

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12
Q

bond formed as a result of sharing electrons between two nuclei

A

Covalent Bond

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13
Q

compounds formed by covalent bonds

A

Molecular Compounds

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14
Q

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

A

Electronegativity

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15
Q

a covalent bond between atoms with the same electronegativity

A

Nonpolar Covalent Bond

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16
Q

a covalent bond between atoms with different electronegativities

A

Polar Covalent Bond

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17
Q

electronegativity difference of a pure covalent bond

A

< 0.4

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18
Q

electronegativity difference of a polar covalent bond

A

between 0.4 and 1.8

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19
Q

electronegativity difference of an ionic bond

A

> 1.8

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20
Q

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

A

Dipole

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21
Q
  • measure of dipole
  • magnitude of the charge on either atom x distance between the two charges
A

Dipole Moment

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22
Q

what molecules have dipoles

A

Polar Molecules

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23
Q

electron-dot structures

A

Lewis Structures

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24
Q

line-bond structures

A

Kekulé Structures

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25
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

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26
Q

formula for getting the formal charge

A

FC = (no. of valence e- in free atom) - (number of valence e- in bonded atom)

FC = (no. of valence e- in free atom) - (no. of bonding e- / 2) - (no. of nonbonding e-)

27
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 1 of a particular atom)

A

Condensed Structures

28
Q
  • Carbon atoms are NOT 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 are NOT shown.
  • Atoms other than carbon and hydrogen are shown.
A

Skeletal Structures

29
Q

3D structures for drawing chemical structures

A
  • Perspective Drawing
  • Ball-and-Stick Model
  • Space-Filling Model
30
Q

shows 3D shape

A

Perspective Drawing

31
Q

shows bond angles accurately

A

Ball-and-Stick Model

32
Q

shows atoms in scale

A

Space-Filling Model

33
Q

bond lies in the plane of the paper (or screen, when viewed electronically)

straight line —————-

A

“normal” bond

34
Q

bond extends backwards, away from the viewer, so effectively “into” the paper (or screen)

A

dashed bond

dashed line - - - - - - - - -

35
Q

bond protrudes forward, towards the viewer, so effectively “out of” the paper (or screen)

A

wedged bond

horizontal triangle

36
Q
  • geometry based on the arrangement of atoms in a molecule
  • defined by bond angles
A

Molecular Geometry

37
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

38
Q

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

39
Q

The best arrangement of a given number of electron pairs (bonding and nonbonding) is the one that ____ the repulsions among them.

A

minimizes the repulsions

40
Q

arrange the pair repulsions

A

lone-pair vs lone-pair repulsion > lone-pair vs bonding pair repulsion > bonding pair vs bonding pair repulsion

41
Q

Without lone pairs:
electron pair geometry =

A

molecular geometry

42
Q

Molecular geometry can
be described in terms of
___.

A

bond angles

43
Q

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

A

Valence Bond Theory

44
Q

assumes the formation of molecular orbitals from the atomic orbitals

A

Molecular Orbital Theory

45
Q

type of bond formed when atomic orbitals on neighboring atoms overlap one another

A

Covalent Bond

46
Q

head-on overlap of atomic orbitals; stronger

A

Sigma (σ) Bonds

47
Q

sideway overlap of atomic orbitals; weaker

A

Pi (π) Bonds

48
Q

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

A

Hybrid Orbitals

49
Q

produces the molecular geometry of the molecule

A

Hybridization

50
Q

shape of any hybrid orbital is [same/different] from the shapes of the original atomic orbitals

A

different

51
Q

as s character increases, bond length _ and bond strength _

A

bond length: decreases
bond strength: increases

inversely proportional

52
Q

the greater the electron density in the region of orbital overlap, the _ and _ the bond

A

stronger and shorter

53
Q

as s character increases, bond angle _

A

bond angle: increases

directly proportional

54
Q

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

A

Molecular Orbital

55
Q

2 ways to combine atomic orbitals into molecular orbitals

A
  1. Additive
  2. Subtractive
56
Q

lower energy, bonding molecular orbital

A

Additive

57
Q

higher energy, antibonding molecular orbital

A

Subtractive

58
Q

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

A

π Electron Delocalization

59
Q

π electron delocalization structures are represented by

A

resonance contributors

60
Q

refers to two individual line-bond structures and and their special resonance relationship is indicated by the double-headed arrow between them

A

Resonance Forms

61
Q

The only difference between resonance forms is the placement of the _ and _.

A

π bond and nonbonding valence electrons

62
Q

A compound with resonance forms _ jump back and forth between two resonance forms.

A

does NOT

63
Q

has a single unchanging structure of the two individual forms and has characteristics of both

A

Resonance Hybrid

64
Q

slide 80

A

CCP 3 LEC (4 decks)

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