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CHEM 119 > Unit 3: Molecular Structure > Flashcards

Unit 3: Molecular Structure Flashcards

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1
Q

how do ionic bonds form

A
  • metal transfers electrons to a nonmetal
  • there is no sharing of electrons
  • the cation and anion each have noble gas configuration
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2
Q

what keeps ionic compounds together

A

electrostatic forces

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3
Q

what is the equation for lattice energy of ionic compounds

A

lattice energy = k(charge of cation * charge of anion/distance between ions)

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4
Q

what causes ionic compounds to have a higher or lower lattice energy

A
  • higher: larger charges, smaller distance
  • lower: smaller charges, larger distance
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5
Q

how do covalent bonds form

A
  • orbitals on one non-metal atom overlap orbitals on another non-metal atoms
  • sharing of valence electrons
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6
Q

is a covalent bond with higher or lower energy more stable

A

lower energy is more stable

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7
Q

do atoms in covalent compounds share electrons equally

A

not always

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8
Q

define electronegativity

A
  • ability of covalently bonded atom to attract electrons to itself
  • higher electronegativity = electrons pulled towards that atom
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9
Q

what is the most electronegative atom in the periodic table

A

fluroine

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10
Q

describe the periodic trend of increasing electronegativity

A
  • increasing moving up
  • increasing moving right
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11
Q

define non-polar covalent bonds in terms of electronegativity

A
  • atoms have similar electronegativities
  • electrons shared equally
  • difference in EN is less than 0.4
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12
Q

define polar covalent bonds in terms of electronegativity

A
  • atoms have different electronegativities
  • electrons shared unequally
  • difference in EN is between 0.4 and 1.9
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13
Q

define ionic bonds in terms of electronegativity

A
  • atoms have highly different electronegativities
  • electrons are transferred to the ion with a higher electronegativity
  • difference in EN is greater than 1.9
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14
Q

describe metallic bonding

A
  • metal atoms donate electrons to an electron sea
  • electrons are delocalized across all of the metal atoms
  • causes electrical conductivity of metals
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15
Q

what do lewis structures represent/show

A
  • valence electrons
  • bonding electrons
  • nonbonding electrons
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16
Q

do lewis structures resemble to actual shape of a molecule

A

no

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17
Q

describe how to draw a lewis structure

A
  • determine the total number of electrons needed
  • draw the central atom and outer atoms connected by single bonds
  • add nonbonding electrons to outer atoms
  • change some nonbonding electrons to bonding electrons to use the exact number of total electrons and make sure all atoms have full valence shells
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18
Q

generally, which atom is the central atom in a lewis structure

A
  • the least electronegative one
  • never hydrogen
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19
Q

define resonance structures

A
  • lewis structures that are different but both work
  • often with double bonds in different places
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20
Q

if there are two resonance structures, does that mean that half of the molecules are built like one and the other half are built like the other

A
  • no
  • resonance structures combine to create 1.5 strength bonds instead of a single and double strength bond
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21
Q

explain how we know that resonance structures combine using O3 as an example

A
  • in O3, we know that the bond lengths are the same between each oxygen
  • we would expect that a shorter length would be observed if there is a double bond
  • we would expect there to be different lengths due to different bonds but that is not what we see
  • bonds are the same length, so we know that there aren’t 1 single and 2 double bonds, there are 2 1.5 strength bonds
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22
Q

what are the 3 elements that form an incomplete octet when bonded

A
  • Be
  • B
  • Al
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23
Q

how many bonds does Be like to form

A

2

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24
Q

how many bonds does B like to form

A

3

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25
Q

how many bonds does Al like to form

A

3

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26
Q

define expanded octet

A

occurs when elements have more than 8 valence electrons surrounding them when bonded

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27
Q

does fluorine like to make a double bond

A

no

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28
Q

what is the equation for formal charge

A
  • valence electrons - (lone electrons + 1/2 bonding electrons)
  • difference between how many valence electrons an atom started with and how many it ended with
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29
Q

explain how formal charge is related to ions

A
  • the sum of the formal charges of all atoms in a molecule in the charge on that molecule
  • neutral molecules have a sum of formal charges equal to 0
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30
Q

are electronegative elements assigned more positive or negative charges in terms of formal charge

A

negative

31
Q

what is the best structure related to formal charge

A
  • one that minimizes the formal charges
  • one with the greatest number of elements having a formal charge of 0
32
Q

describe the valence shell electron pair repulsion theory

A

electron pairs will repel one another so they are placed as far apart in space as possible to minimize the repulsion

33
Q

define electron domain

A
  • any side of an atom where there are electron: either pairs or bonds
  • lone pairs = 1 electron domain
  • single bonds = 1 electron domain
  • double bonds = 1 electron domain
  • triple bonds = 1 electron domain
34
Q

what is the molecular geometry name for a molecule with 2 electron domains

A

linear

35
Q

what is the molecular geometry name for a molecule with 3 electron domains

A

trigonal planar

36
Q

what is the molecular geometry name for a molecule with 4 electron domains

A

tetrahedral

37
Q

what is the molecular geometry name for a molecule with 5 electron domains

A

trigonal bipyramidal

38
Q

what is the molecular geometry name for a molecule with 6 electron domains

A

octahedral

39
Q

what are the bond angles in a linear geometry molecule

A

180 degrees

40
Q

what are the bond angles in a trigonal planar geometry molecule

A

120 degrees

41
Q

what are the bond angles in a tetrahedral geometry molecule

A

109.5 degrees

42
Q

what are the bond angles in a trigonal bipyramidal geometry molecule

A

120 degrees and 90 degrees

43
Q

what are the bond angles in an octahedral geometry molecule

A

90 degrees

44
Q

does molecular geometry take into account atoms, nonbonding pairs, or both

A
  • only atoms
  • doesn’t take into account nonbonding pairs
45
Q

is the electron-domain geometry always the shape of the molecule

A
  • no
  • doesn’t consider nonbonding pairs which can change the shape of the molecule
46
Q

what is the molecular geometry of a trigonal planar electron-domain shape with 2 atoms and 1 nonbonding pair

A

bent

47
Q

why would the two atoms on a bent shaped molecule (trigonal planar with 2 atoms and 1 nonbonding pair) be closer together than the expected 120 degrees

A
  • nonbonding pairs are physically larger than bonding pairs
  • the added nonbonding pair will push the two atoms slightly closer together
48
Q

what is the molecular geometry of a tetrahedral electron-domain shape with 3 atoms and 1 nonbonding pair

A

trigonal pyramidal

49
Q

what is the molecular geometry of a tetrahedral electron-domain shape with 2 atoms and 2 nonbonding pair

A

bent

50
Q

what are the top and bottom atoms called in the trigonal bipyramidal electronic geometry

A

axial bonds

51
Q

what are the middle atoms called in the trigonal bipyramidal electronic geometry

A

equatorial bonds

52
Q

will lone pairs in the trigonal bipyramidal electronic geometry be in the equatorial or axial position and why

A
  • equatorial
  • gives electrons the most space
53
Q

what is the molecular geometry of a trigonal bipyramidal electron-domain shape with 4 atoms and 1 nonbonding pair

A

seesaw

54
Q

what is the molecular geometry of a trigonal bipyramidal electron-domain shape with 3 atoms and 2 nonbonding pair

A

t-shaped

55
Q

what is the molecular geometry of a trigonal bipyramidal electron-domain shape with 2 atoms and 3 nonbonding pair

A

linear

56
Q

what is the molecular geometry of a octahedral electron-domain shape with 5 atoms and 1 nonbonding pair

A

square pyramidal

57
Q

what is the molecular geometry of a octahedral electron-domain shape with 4 atoms and 2 nonbonding pair

A

square planar

58
Q

what is special about the linear molecular geometry originating from the trigonal bipyramidal electronic geometry and the square planar molecular geometry originating from the octahedral electronic geometry and why

A
  • both have the same ideal bond angles as the electronic geometry (180 for linear and 90 for square planar
  • because the lone pairs are all opposite of each other and create equal repulsion
59
Q

how should you go about determining the geometry of a larger molecule

A

determine the geometry around each particular atom

60
Q

what is the overall dipole moment of a nonpolar molecule

A

0

61
Q

what is the overall dipole moment of a polar molecule

A

not 0

62
Q

explain how to determine if a molecule is polar

A
  • look at individual bonds and identify bond polarity
  • draw arrow toward the more electronegative atom in a polar bond and add a + to the positive end
  • determine molecular polarity by looking at arrows and seeing if they cancel out (cancel out=nonpolar)
63
Q

define hybrid orbital

A

proposed that s and p orbitals mix together to form molecular geometries

64
Q

describe how to determine the hybridization of an atom

A
  • determine the number of electron domains around the atom
  • add orbitals starting from s1 until you have the same number of orbitals as electron domains
  • there is 1 s orbital and 3 p orbitals
  • hybridization examples are: s, sp, sp2, sp3
65
Q

what must the orbitals of atoms do when covalent bonds are formed

A

overlap

66
Q

what are the 2 types of internuclear bonds

A
  • sigma
  • pi
67
Q

describe the sigma internuclear bond

A
  • occurs when orbitals overlap along the bonding axis, one point of overlap
  • the first covalent bond is always sigma
68
Q

describe the pi internuclear bond

A
  • occurs when orbitals overlap parallel to the bonding axis, two points of overlap
  • the second and third covalent bonds are always pi bonds (one bond in a double bond, two bonds in a triple bond)
69
Q

what is the relationship between bond type (single, double, triple) and bond length

A
  • more bonds (more electrons shared) = shorter bond lengths
  • more elections have stronger pull between atoms to make them closer together
70
Q

what is the relationship between bond type (single, double, triple) and bond energy

A
  • more bonds (more electrons shared) = higher bond energy
  • more electrons makes the bond stronger which means it requires more energy to break
71
Q

what is the relationship between the atoms bonded and bond length

A
  • as you move to the right of the periodic table, the bond lengths get smaller
  • C-C is longer than C-O
72
Q

define constructive inference

A
  • bonding interaction
  • two things combine to create something bigger
  • stable, low energy
  • denoted as sigma
73
Q

define destructive inference

A
  • antibonding interaction
  • two things combine and cancel out
  • unstable, high energy
  • denoted as sigma *

CHEM 119 (8 decks)

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