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CHEM1601 > Chapter 5: Chemical Bonding I > Flashcards

Chapter 5: Chemical Bonding I Flashcards

1
Q

What does VSEPR stand for?

A

Valence shell electron pair repulsion.

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2
Q

Define electronegativity.

A

Electronegativity is the ability of an atom to attract electrons to itself in a chemical bond. An atom that is more electronegative than another takes a greater share of the electron density.

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3
Q

How does electronegativity change across a period?

A

It increases.

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4
Q

How does electronegativity change down a column?

A

It decreases.

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5
Q

What is the most electronegative element?

A

Fluorine (F).

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6
Q

What is the least electronegative element (the most electropositive)?

A

Francium.

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7
Q

In general, how is electronegativity related to atomic size?

A

Electronegativity is inversely related to atomic size. The larger the atom, the less ability it has to attract electrons to itself in a chemical bond.

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8
Q

How is the polarity of a bond quantified?

A

The polarity of a bond is quantified by the size of its dipole moment. A dipole moment occurs anytime there is a charge separation and can be calculated using the equation DM = (magnitude of charge)(distance)

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9
Q

What is the percent ionic character?

A

The percent ionic character is the ratio of a bond’s actual dipole moment to the dipole moment it would have if the electron were completely transferred from one atom to the other, multiplied by 100%. A bond in which an electron is completely transferred from one atom to another would have 100% ionic character.

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10
Q

What are the four steps for writing a Lewis structure?

A
  1. Write the correct skeletal structure for the molecule.
  2. Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule.
  3. Distribute the electrons among the atoms, giving octets (or duets in the case of hydrogen) to as many atoms as possible.
  4. If any atoms lack an octet, form double or triple bonds as necessary to give them octets.
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11
Q

What are the two guidelines for writing the skeletal structure of a molecule?

A
  1. Hydrogen atoms are always terminal.
  2. The more electronegative elements are usually more terminal, and the less electronegative elements (other than hydrogen) are more central.
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12
Q

How are electrons shared in metallic bonds?

A

A crystal of metal atoms exist in a “sea of electrons.” This means the electrons are delocalized throughout the entire structure, and no one electron can be attributed to one atom.

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13
Q

What are the exceptions to the octet rule within the main group elements?

A
  1. Hydrogen can only hold two electrons.
  2. Beryllium is satisfied with a quartet (4 electrons)
  3. B & Al are satisfied with a hextet (6 electrons)
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14
Q

How do electrons shared in a covalent bond differ from lone pairs?

A

Lone pairs are localized to a single atomic orbital. Electrons pairs shared by two atoms in a covalent bond are located in a molecular orbital. Essentially, it is a matter of atomic vs molecular orbitals.

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15
Q

What is the electronegativity of H?

A

2.1

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16
Q

What is the electronegativity of C?

A

2.5

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17
Q

What is the electronegativity of N?

A

3.0

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18
Q

What is the electronegativity of O?

A

3.5

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19
Q

What is the electronegativity of F?

A

4.0

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20
Q

What is the electronegativity of Cl?

A

3.0

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21
Q

A difference in electronegativity over what range results in a pure covalent bond?

A

0.0-0.4

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22
Q

A difference in electronegativity over what range results in a polar covalent bond?

A

0.4-2.0

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23
Q

A difference in electronegativity over what range results in an ionic bond?

A

> 2.0

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24
Q

Define resonance structure.

A

A resonance structure is one of two or more Lewis structures that have the same skeletal formula (the atoms are in the same locations) but different electron arrangements.

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25
Q

Define resonance hybrid.

A

A resonance hybrid is the actual structure of the molecule intermediate between the two or more resonance structures.

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26
Q

Why do electrons delocalize in a molecule?

A

The delocalization of electrons lowers their energy and stabilizes them.

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27
Q

What is the formal charge of an atom in a Lewis structure?

A

The charge it would have if all bonding electrons were shared equally between the bonded atoms. In other words, the calculated charge for an atom in a molecule if we completely ignore the effects of electronegativity.

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28
Q

How do you calculate formal charge?

A

Formal charge = number of valence electrons - (number of nonbonding electrons + 1/2 number of bonding electrons)

29
Q

What are the four rules for determining formal charge?

A
  1. The sum of all formal charges in a neutral molecule is zero.
  2. The sum of all formal charges in an ion must equal the charge of the ion.
  3. Small (or zero) formal charges on individual atoms are better than large ones.
  4. When formal charge cannot be avoided, negative formal charge should reside on the most electronegative atom.
30
Q

In general, where should the least electronegative atom be in a Lewis structure?

A

It should be in the central position.

31
Q

Define free radicals.

A

Free radicals are a type of odd-electron species that have an odd number of electrons in their Lewis structures.

32
Q

Why are 10- and 12-electron expanded octets common in third-period elements and beyond?

A

The d orbitals in these elements are energetically accessible (they are not much higher in energy than the orbitals occupied by the valence electrons) and can accommodate the extra electrons. Expanded octets never occur in second-period elements because they do not have energetically accessible d orbitals and therefore never exhibit expanded octets.

33
Q

Define bond energy.

A

The bond energy of a chemical bond is the energy required to break 1 mole of the bond in the gas phase. Bond energies are positive because energy must be put into a molecule to break a bond (the process is endothermic).

34
Q

In VSEPR theory, what determines the geometry of a molecule?

A

The repulsions between electron groups on interior atoms (or the central atom) of a molecule determine its geometry. The preferred geometry of a molecule is the one in which the electron groups have the maximum separation (and therefore, the minimum energy, based on Coulomb’s law) possible.

35
Q

Why do lone pairs of electrons exert greater repulsions than bonding electrons?

A

A lone electron pair is more spread out in space than a bonding electron pair because a lone pair is attracted to only one nucleus while a bonding pair is attracted to two. The lone pair occupies more of the angular space around a nucleus, exerting a greater repulsive force on neighboring electrons and compressing the bond angles.

36
Q

How is a bond angle related to the number of lone pairs on a molecule?

A

As lone pairs are added to the molecule, the bond angle gets progressively smaller.

37
Q

Why do bond angles vary from their idealized angles sometimes?

A

Double and triple bonds occupy more space than single bonds (they are bulkier even though they are shorter), and lone pairs occupy more space than bonding groups. The presence of lone pairs usually makes bond angles smaller than the ideal angle for the particular geometry.

38
Q

How do polar bonds affect the overall polarity of a diatomic molecule?

A

If a diatomic molecule has a polar bond, the entire molecule is polar. If the bond is nonpolar, the molecule as a whole will be nonpolar.

39
Q

How do polar bonds affect the overall polarity of a polyatomic molecule?

A

If the molecular geometry is such that the dipole moments of individual polar bonds sum together to a net dipole moment, then the molecule is polar. But if the molecular geometry is such that the dipole moments of the individual polar bonds cancel each other out, then the molecule is nonpolar. Polarity is dependent on the geometry of the molecule.

40
Q

Why can dipole moments cancel each other out?

A

Dipole moments cancel out because they are vector quantities, meaning they have both a magnitude and a direction. If you think of each polar bond as a vector pointing in the direction of the more electronegative atom, the length of the vector is proportional to the electronegativity difference between the bonding atoms.

41
Q

How are odd electron molecules accommodated in a Lewis structure?

A

If a molecule contains an odd number of electrons, all octets cannot be satisfied. The odd electron must be put on the more electropositive atom.

42
Q

Which atoms can form molecules that are electron deficient?

A

Be, B, and Al. Be = 4, B and Al = 6

43
Q

What is the prerequisite for valence shell expansion?

A

For atoms whose valence shell is greater than or equal to 3, they can expand their valence to hold up to 12 electrons. This requires energetically accessible d orbitals.

44
Q

When minimizing the formal charge of a molecule, where should the negative formal charge be placed?

A

On the most electronegative atom.

45
Q

Define bond energy and give its two important characteristics.

A

Bond energy is the total energy required to break 1 mole of a bond in a molecule in the gas phase. Bond energy is always a positive number (endothermic), and the greater the value, the stronger the bond.

46
Q

What is the electronic and molecular geometry of steric number 2 molecules?

A

The geometry is linear for both electronic and molecular geometry.

47
Q

What is the bond angle between atoms for steric number 2 molecules?

A

180.

48
Q

What is the electronic and molecular geometry of steric number 3 molecules?

A

Electronic: Trigonal planar
Molecular: trigonal planar (no lone pair) or bent (1 lone pair)

49
Q

What is the bond angle for steric number 3 molecules?

A

Trigonal planar (no lone pair): 120
Bent (1 lone pair): <120

50
Q

What is the electronic and molecular geometry of steric number 4 molecules?

A

Electronic: Tetrahedral
Molecular: tetrahedral (no lone pairs)
Trigonal pyramidal (1 lone pair)
Bent (2 lone pairs)

51
Q

What is the bond angle for steric number 4 molecules?

A

Tetrahedral (no lone pairs): 109.5
Trigonal pyramidal (1 lone pair): <109.5
Bent (2 lone pairs): <109.5

52
Q

What is the electronic and molecular geometry of steric number 5 molecules?

A

Electronic: Trigonal bipyramidal
Molecular: trigonal bipyramidal (no lone pairs)
Seesaw (1 lone pair)
T-shaped (2 lone pairs)
Linear (3 lone pairs)

53
Q

What is the bond angle for steric number 5 molecules?

A

Trigonal bipyramidal (no lone pairs): 120 (equatorial) and 90 (axial)
Seesaw (1 lone pair): <120 (equatorial) and <90 (axial)
T-shaped (2 lone pairs): <90
Linear (3 lone pairs): 180

54
Q

What is the electronic and molecular geometry of steric number 6 molecules?

A

Electronic: octahedral
Molecular: octahedral (no lone pairs)
Square pyramidal (one lone pair)
Square planar (two lone pairs)

55
Q

What is the bond angle for steric number 6 molecules?

A

Octahedral (no lone pairs): 90
Square pyramidal (one lone pair): <90
Square planar (two lone pairs): 90

56
Q

For steric number 5 molecules, which plane are lone pairs placed in?

A

The equatorial plane.

57
Q

For steric number 6 molecules, where is the second lone pair located relative to the first lone pair?

A

The second lone pair is always located opposite from the first.

58
Q

What do straight lines, hacked wedges, and solid wedges represent when drawing a molecular structure?

A

Straight line: bond in plane of paper.
Hacked wedge: bond going into the page.
Solid wedge: bond coming out of the page.

59
Q

Describe how charge influences polarity.

A

Charge does NOT influence polarity.

60
Q

Use CO2 and H2O as an example to illustrate how polarity affects the physical properties of a molecule.

A

CO2 is three times heavier per molecule than H2O, yet it “boils” at -78.5 whereas H2O boils at 100. The polarity of H2O keeps it a liquid at much higher temperatures than its mass would suggest.

61
Q

What is the relationship between electronegativity and atomic size?

A

Electronegativity is inversely related to atomic size.

62
Q

Describe the relationship between the magnitude of the charge separation, the distance between the charges, and the magnitude of the dipole moment.

A

The smaller the magnitude of the charge separation, and the smaller the distance between the charges, the smaller the dipole moment.

63
Q

Why don’t individual resonance structures exist?

A

The individual resonance structures are just a way of describing the structure. The only structure of a molecule that exists is the resonance hybrid.

64
Q

How does the bond energy of triple, double, and single bonds compare?

A

The bond energy of triple bonds is greater than that of double, which is greater than that of single bonds.

65
Q

Why do the actual bond angles of a molecule deviate from the theoretical bond angle more in molecules with double bonds?

A

Double bonds have greater electron density than single bonds, meaning they repel the other electron groups in a molecule with greater force.

66
Q

Contrast the repulsions of lone pair and bonding electrons.

A

Lone pair electrons exert slightly greater repulsions than bonding electrons.

67
Q

Why do steric number 5 molecules with three lone pairs of electrons exhibit linear molecular geometry?

A

Steric number 5 molecules with three lone pairs are linear because the lone pairs occupy the three spots for electrons in the equatorial plane. This leaves the remaining two electron spots in a single plane–the axial plane–generating a linear geometry with 180 degree bond angles.

68
Q

How do mercury (I) and mercury (II) exist?

A

mercury (II) = Hg2+
mercury (I) = Hg2 2+ (exists as a dimer)

69
Q

How is lattice energy related to charge and position on a group of the periodic table?

A

Lattice energy is proportional to the product of charges and decreases in magnitude down a group.

CHEM1601 (17 decks)

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