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CHEM1110 > SECTION 4 - CHEMICAL BONDING II > Flashcards

SECTION 4 - CHEMICAL BONDING II Flashcards

1
Q

VSEPR

A

Valence shell electron repulsion theory that tells us the 3 dimensional shape of a molecule

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2
Q

Electron group geometry (egg)

A

Based on the number of electrons surrounding central atom (bonded atoms + lone pairs)

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3
Q

5 basic electron group geometry shapes

A
  1. Linear (2 e)
  2. Trigonal planar (3 e)
  3. Tetrahedral (4 e)
  4. Trigonal bypyramidal (5 e)
  5. Octahedral (6 e)
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4
Q

Dash lines

A

Extend BEHIND plane

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5
Q

Solid thick lines

A

Protrude OUT plane

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6
Q

Molecular shape (mg)

A

Only count atoms bonded to central atom

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7
Q

egg = mg

A

When all electron groups are atoms

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8
Q

egg doesnt equal mg when

A

Lone pairs present

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9
Q

5 molecular geometry shape

A
  1. Bent
  2. Seesaw
  3. T shape
  4. Square pyramidal
  5. Square planar
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10
Q

BLT

A

BP LP
2 1
2 2
BENT

2 3
LINEAR

3 1
TRIGONAL PYRAMIDAL
3 2
T SHAPE

4 1
SEESAW
4 2
SQUARE PLANAR
5 1
SQUARE PYRAMIDAL

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11
Q

Axial bonding position

A

Up/down

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12
Q

Equatorial bonding position

A

Horizontal/diagnol
- LP go on equatorial

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13
Q

Dipole moment

A

Measure of molecular polarity with unit of Debye (D). 1D = 3.34x10^-30cm

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14
Q

Dipole moment in diatomic molecule

A

There is only 1 bond so the dipole determined by bond

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15
Q

Dipole moment in polyatomic molecule

A

The total polarity determined by bond polarity and bond angle

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16
Q

Arrows facing out from one another

A

U = 10 (not polar)

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17
Q

Arrows facing towards each other

A

Polar

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18
Q

Valence bond theory

A

Tells us how electrons interact to hold atoms together into molecules and defines covalent bonds as electron density between 2 atoms as a result of constructive interference

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19
Q

Ideal distance of valence bond theory

A

Occurs with overlap between partially filed atomic bonds on each atom and the overlap creates new covalent bond

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20
Q

Inphase

A

Atoms that are in phase will overlap constructively and form bond

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21
Q

Outphase

A

Atoms that are out phase creates region of 0 electron density (node)

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22
Q

Node

A

N-1 (region of 0 electron density)

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23
Q

Sigma bond (3)

A
  • Formed by overlap of atomic orbitals head to head
  • All single bonds are sigma bonds
  • Free rotation allowed
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24
Q

Pi bond (5)

A
  • Formed by overlap of p orbitals side to side
  • Multiple bonds consists of one sigma bond and some pi bonds
  • Contains 2 regions of electron density
  • Rotation not allowed
  • Overlapping must occur in similar phase orbitals
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25
Q

Sp hybridization

A

2 electron groups

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26
Q

Sp2 hybridization

A

3 electron groups

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27
Q

Sp3 hybridization

A

4 electron groups

28
Q

Sp3d hybridization

A

5 electron groups

29
Q

Sp3d hybridization

A

6 electron groups

30
Q

How to valence theory (4)

A
  1. Identify valence orbitals of bonding atoms
  2. Sketch valence orbital
  3. Bring in atoms to be bonded to central atom and sketch overlap
  4. Describe structure
31
Q

Hybrid orbitals

A

The “mixing’ of orbitals

32
Q

Relationship between the # of atomic orbitals and the # of hybrid orbitals produced

A

Must equal

33
Q

Molecular orbital theory

A

States that electron in molecules are not assigned to individual chemical bonds, but rather moving under the influence of atomic nuclei in molecule. Electrons are assigned to molecules as a whole

34
Q

Constructiveness interference

A

Forms a bonding molecular orbital with high region of electron density between nuclei

35
Q

Destructiveness interference

A

Forms an anti bonding molecular orbital with a node between nuclei

36
Q

Molecular orbital rules (6)

A
  1. # of atomic orbitals combined = # of molecular orbitals
  2. Bonding orbitals hold molecules together with low energy
  3. Antibonding orbitals do not help hold molecules together and are high energy
  4. Electrons fill lowest energy first
  5. Degenerate orbitals fill singly before pairing
  6. Maximum 2 electrons per orbital
37
Q

*

A

High energy antibonding molecular orbital

38
Q

Bonding molecular orbital

A

Force goes towards atom

39
Q

Antibonding molecular orbital

A

Forces repel

40
Q

Bond order

A

Bonded electrons-antibonded electorns/2

41
Q

Bond order 1

A

Single bond

42
Q

Bond order 2

A

Double bond

43
Q

Bond order 1/2

A

Single bond with one electron

44
Q

Bond order 0

A

Indicates that a bond is unlikely to form

45
Q

Higher bond order

A

Indicates stronger bond, high stability, low energy, smaller bond length

46
Q

Bond order 3

A

Triple bond

47
Q

Paramagnetic

A

Contains unpaired electrons

48
Q

Diamagnetic

A

Contains paired electrons

49
Q

HOMO

A

Highest energy orbitals (has electrons)

50
Q

LUMO

A

Lowest energy orbital (no electrons)

51
Q

3 states of matter

A

Solid, liquid, gas

52
Q

Condensed phases

A

Solid and liquids. This is because they are all packed closer together

53
Q

How does temperature impact kinetic energy

A

Higher temperature = higher kinetic energy

54
Q

Intermolecular forces

A

Attractive/repulsive force between molecules of a substance that impact physical properties of matter. Without this, all matter would exist as a gas

55
Q

Strong intermolecular force

A

Requires more thermal energy to break the molecules apart = higher boiling/melting point

56
Q

3 types of intermolecular force

A
  1. London dispersion
  2. Dipole-dipole
  3. Hydrogen bonding
57
Q

Intramolecular force

A

Found within the molecule that keeps it together

58
Q

London dispersion

A

Present in all neutral molecules/atoms and the weakest bond. Caused by when electrons are not evenly distributed inside orbital and results in small partial charges

59
Q

Polarizability

A

Tendency of an atom/molecule to form a distorted electron cloud based on 3 conditions

60
Q

Polarizability dependancy (3)

A
  1. Larger atoms = more polarizable
  2. Larger molecules = more polarizable
  3. Larger surface area = stronger dispersion force
61
Q

Dipole-dipole (3)

A
  • Occurs when a molecule has a permanent dipole and partial negative lines up with partial positive
  • The more polar, the stronger the dipole
  • Stronger the dipole = higher the boiling point (harder to pull)
62
Q

Hydrogen bonding

A

Strongest intermolecular force that occurs when a partial positive hydrogen atom is covalently bonded to small electronegative partial negative atom N, O, F

63
Q

Order in strength of intermolecular force

A

London dispersion –> dipole-dipole –> hydrogen bond

64
Q

Water physical property

A

Each H2O molecule can form up to 4 hydrogen bonds at one time which makes water liquid at room temperature

65
Q

Why is liquid water more dense than solid water

A

Because of the orientation of hydrogen bonds that cause the molecules to push further apart in solid state, lowering the density

CHEM1110 (18 decks)

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