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Gen Chem > Ch. 3: Bonding and Chemical Interactions > Flashcards

Ch. 3: Bonding and Chemical Interactions Flashcards

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1
Q

what is the basis behind why covalent bonds form?

A
  1. when two or more atoms with similar EN interact, the energy required to form ions through the complete transfer of one or more electrons is greater than the energy that would be released among upon formation of an ionic bond

when two atoms of similar tendency to attract electrons form a compound, it is energetically unfavorable to create ions

  1. instead of transferring electrons to form octets, they share electrons

there is an attraction that each electron in the shared pair has for the two positive nuclei of the bonded atoms

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2
Q

why do covalent compounds have lower melting and boiling points?

A

because they contain discrete molecular units with relatively weak intermolecular interactions

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3
Q

why are covalent compounds poor conductors of electricity in the liquid state or in aq. solns?

A

because they do not break down into constituent ions

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4
Q

why do single, double, or triple covalent bonds form?

A

the formation of one covalent bond may not be sufficient to fill the valence shell for a given atom

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5
Q

defn: bond order

A

the number of shared electron airs between two atoms

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6
Q

what is the bond order for a single bond? double bond? triple bond?

A

single bond: bond order one

double bond: bond order two

triple bond: bond order three

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7
Q

what are the three important characteristics of a covalent bond?

A

bond length

bond energy

polarity

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8
Q

defn: bond length

A

the average distance between the two nuclei of atoms in a bond

as the # of shared electron pairs increases, the two atoms are pulled closer together, resulting in a decrease in bond length

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9
Q

defn: bond energy

A

the energy required to break a bond by separating its components into their isolated, gaseous atomic states

the greater the number of pairs of electrons shared between the atomic nuclei, the more energy required to break the bonds holding the atoms together

the greater the bond energy, the stronger the bond (generally)

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10
Q

defn: polarity

A

occurs when two atoms have a relative difference in electronegativities

when these atoms come together in covalent bonds, they must negotiate the degree to which the electron pairs will be shared

the atom with the higher electronegativity gets the larger share of the electron density

a polar bond creates a dipole, with the positive end of the dipole at the less EN atom and the negative end at the more EN atom

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11
Q

what is the relationship between the number of covalent bonds, bond length, bond energy, and bond strength?

A

1 bond = longest, weakest, lowest bond energy

2 bonds = medium length, medium strength, medium bond energy

3 bonds = shortest, strongest, highest bond energy

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12
Q

defn: nonpolar covalent bond

A

when atoms that have identical or nearly identical EN’s share electron pairs, they do so with an equal distribution of electrons

there is no separation of charge across the bond

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13
Q

what set of molecules contain perfectly nonpolar covalent bonds?

A

bonds between atoms of the same element (have exactly the same EN) –> exhibit a purely equal electron distribution

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14
Q

what are the 7 common diatomic molecules?

A

H2, N2, O2, F2, Cl2, Br2, I2

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15
Q

what is the EN difference that qualifies as nonpolar? polar? ionic?

A

nonpolar: 0 - 0.5
polar: 0.5 - 1.7
ionic: 1.7 - 2.0

for the MCAT: if molecule in this range, has a metal and a nonmetal it is effectively ionic, otherwise it is polar covalent

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16
Q

defn: polar covalent bonds

A

atoms that differ moderately in their EN’s share their electrons unevenly

the difference in their EN’s is not enough to form an ionic bond, but is sufficient to cause a charge separation across the bond

the more EN element takes on a greater portion of the electron density (a partial negative charge)

the less EN element takes on a smaller portion of the electron density (a partial positive charge)

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17
Q

defn + eqn: dipole moment

A

the dipole moment of a polar bond or polar molecule is a vector quantity given by the equation

p = qd

p = the dipole moment
q = the magnitude of the charge
d = the displacement vector separating the two partial charges

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18
Q

defn: polar molecule

A

a molecule that has a significant separation of positive and negative charges

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19
Q

unit: dipole moment vector

A

Debye’s (coulomb-meters)

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20
Q

defn: coordinate covalent bond

A

both of the shared electrons originated on the same atom

generally: a lone pair of one atom attacked another atom with an unhybridized p-orbital to form a bond

indistinguishable from any other covalent bond

typically found in Lewis acid-base reactions

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21
Q

defn: bonding vs. nonbonding electrons

A

bonding electrons = the electrons involved in a covalent bond and are in the valence shell

nonbonding electrons = the electrons in the valence shell not involved in covalent bonds

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22
Q

defn: lone pairs

A

unshared electron pairs (only associated with one atomic nucleus)

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23
Q

what was the Lewis structure system developed for?

A

to keep track of bonded and nonbonded electron pairs

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24
Q

defn: formal charge

A

the difference in the number of valence electrons assigned to an atom in a Lewis structure of a particular molecule and the number of valence electrons in the neutral atom (normally found in the atom’s valence shell)

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25
Q

what do Lewis structures represent if they differ in bond connectivity or arrangement? if they differ in arrangement of electron pairs?

A
  1. different possible compounds
  2. different resonance forms of a single compound
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26
Q

do Lewis structures represent actual geometry? theoretical geometry?

A

no no no!

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27
Q

what is the most stable arrangement of a compound as viewed through Lewis dot diagrams?

A

the arrangement that minimizes the number and magnitude of formal charges

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28
Q

defn: Lewis dot diagram

A

the chemical symbol of an element surrounded by dots which each represent one of the s or p valence electrons of the atom

the number of dots comes from group members (group 1 has 1 dot)

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29
Q

what are the steps for drawing a Lewis structure? HCN is your example ifneeded.

A
  1. draw out the compound’s backbone (the arrangement of the atoms) –> the least EN atom is the central atom, H always is terminal, F, Cl, Br, and I are usually terminal
  2. count all the valence electrons of the atoms (molecule is sum of all the atoms present)

HCN has 10 total

  1. draw single bonds between the central atom and the surrounding atoms (each single bond = a pair of electrons)
  2. complete the octets of all atoms bonded to the central atom, using remaining valence electrons left to be assigned (H is an exception)
  3. Place any extra electrons on the central atom
  4. If the central atom has less than an octet, try to write double or triple bonds between the central and surrounding atoms using the lone pairs on the surrounding atoms
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30
Q

what is the purpose of calculating the formal charge of an atom?

A

to determine if a lewis structure is representative of the actual arrangement of atoms in a compound

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31
Q

what assumption must be made when calculating the formal charge of an atom?

A

assume a perfectly equal sharing of all bonded electron pairs, regardless of actual differences in EN

(assume that each electron pair is split evenly between the two nuclei in the bond)

32
Q

what are the actual and mnemonic equations to calculate formal charge?

A

actual:

formal charge = V - Nnonbonding - 0.5 Nbonding

(V = normal # of electrons in the atom’s valence shell; Nnonbonding = # of nonbonding electrons, Nbonding = # bonding electrons (double the number of bonds because each bond has two electrons))
mnemonic:

formal charge = valence electrons - dots - sticks

33
Q

what is the relationship between the charge of an ion or compound and the formal charges of the individual atoms?

A

the charge of an ion or compound = the sum of the formal charges of the individual atoms comprising the ion or compound

34
Q

what is the difference between formal charge and oxidation number?

A

formal charge UNDERESTIMATES the effect of EN differences

oxidation numbers OVERESTIMATE the effect of EN differences, assuming that the more EN atom has 100% share of the bonding electron pair

in reality, the distribution of electron density lies somewhere between the extremes predicted by the formal charge and the oxidation states

35
Q

defn: resonance structures

A

the same arrangement of atoms that differ in the specific placement of electrons

it may be possible to draw 2+ Lewis structures that demonstrate this

the actual electronic distribution in the compound is a hybrid of all possible resonance structures (resonance hybrid)

the more stable the structure, the more it contributes to the character of the resonance hybrid

36
Q

what are the guidelines to follow to use formal charge to assess the stability of resonance structures?

A
  1. a lewis structure with small or no formal charges is preferred over a Lewis structure with large formal charges
  2. a lewis structure with less separation between opposite charges over a lewis structure with a large separation of opposite charges
  3. a lewis structure in which negative formal charges are placed on more electronegative atoms is more stable than one in which the negative formal charges are placed on less electronegative atoms
37
Q

describe exceptions to octet rule

A
  1. H, He, Li, Be, B –> cannot or do not usually reach the octet
  2. in or beyond the third period –> can take on more than 8 electrons in their valence shells (can be placed into orbitals of the d subshell) –> atoms of these electrons can form more than 4 bonds
38
Q

defn + long name: VSEPR theory

A

valence shell electron pair repulsion theory

uses Lewis dot structures to predict the molecular geometry of covalently bonded molecules

the 3-D arrangement of atoms surrounding a central atom is determined by the repulsions between bonding and nonbonding electron pairs in the valence shell of the central atom

these electron pairs arrange themselves as far apart as possible, thereby minimizing repulsive forces

39
Q

what are the steps used to predict the geometrical structure of a molecule using VSEPR theory?

A
  1. Draw the Lewis dot structure of the molecule
  2. Count the total number of bonding and nonbonding electron pairs in the valence shell of the central atom
  3. Arrange the electron pairs around the central atom so that they are as far apart as possible
40
Q

defn: electronic geometry vs. molecular geometry

A

ELECTRONIC GEOMETRY = the spatial arrangement of all pairs of electrons around a central atom (including both bonding and lone pairs)

MOLECULAR GEOMETRY = the spatial arrangement of only the bonding pairs of electrons

41
Q

defn: coordination number

A

the number of atoms that surround and are bonded to a central atom

relevant to molecular geometry

42
Q

what is a unique aspect of electronic geometry that molecular geometry does not entail?

A

determination of the ideal bond angle

43
Q

describe how the polarity of molecules works

A
  1. the presence of bond dipoles does not necessarily result in a molecular dipole (the overall separation of charge across the molecule)
  2. we must consider the molecular geometry and the vector addition of the bond dipoles based upon that molecular geometry

compound with nonpolar bonds = always nonpolar

compound with polar bonds = may be polar or nonpolar depending on the spatial orientation of the polar bonds in the molecule

if vector sum is 0 (the bond dipole moments cancel each other out) = result is nonpolar compound (CCl4)

if molecular geometry is such that bond dipoles do not cancel each other out –> molecule has a net dipole moment, molecule is polar (H2O)

44
Q

describe the model of an atom in regards to orbitals

A

atom = a dense, positively charged nucleus surrounded by a cloud of electrons organized into orbitals

orbital = regions in space surrounding the nucleus within which there are certain probabilities of finding an electron

45
Q

describe the characteristics of s and p subshells

A

S SUBSHELL
- l = 0
- one spherical orbital

P SUBSHELL
- l = 1
- 3 orbitals shaped like barbells along the x, y, and z axes at right angles to each other

46
Q

defn: molecular orbital

A

forms when atomic orbitals interact when two atoms bond to form a compound

describes the probability of finding the bonding electrons in a given space

obtained by combining the wave functions of the atomic orbitals

the overlap of the two atomic orbitals describes this quantitatively

47
Q

defn: bonding orbital vs. antibonding orbital

A

both types of molecular orbitals

BONDING = the signs of the two atomic orbitals are the same

ANTIBONDING = the signs of the two atomic orbitals are different

48
Q

what are the two types of bonds that are observed in the formation of molecular orbitals? describe their characteristics.

A

these are both two different patterns of overlap

HEAD TO HEAD OVERLAP = sigma bond (allow for free rotation about axes because electron density of the bonding orbital is a single linear accumulation between the atomic nuclei)

TWO PARALLEL ELECTRON CLOUD DENSITIES = pi bond (do not allow for free rotation because the electron densities of the orbitals are parallel, cannot be twisted in a way that allows for continuous overlapping of electron density clouds)

49
Q

CH 3.1

A
50
Q

defn: molecules

A

the atoms of most elements can combine to form molecules

51
Q

defn: chemical bonds

A

strong attractive forces that hold together the atoms within molecules

52
Q

how are chemical bonds formed?

A

via the interaction of the valence electrons of the combining atoms

53
Q

what rule do many atoms follow when joining together to form molecules?

A

the octet rule

54
Q

what does the octet rule state? is this truly a rule?

A

an atom tends to bond with other atoms so that it has eight electrons in its outermost shell, thereby forming a stable electron configuration similar to that of the noble gases

the desire of all gases to achieve noble gas configuration

it is a rule of thumb (there are many exceptions)

55
Q

diagram: electron configuration of Argon as an example of a complete octet in its valence shell

A
56
Q

what are 5 notable exceptions of the octet rule? and how many valence electrons do they have?

A
  1. Hydrogen – 2 valence electrons
  2. Lithium – 2 valence electrons
  3. Beryllium – 4 valence electrons
  4. Boron – 6 valence electrons
  5. All elements in period 3 or greater – can expand the valence shell to include more than 8 electrons by incorporating d-orbitals
57
Q

what are the three types of exceptions to the octet rule which are easier to remember than specific elements?

A
  1. incomplete octet = these elements are stable with fewer than 8 electrons in their valence shell and include hydrogen (2), helium (2), lithium (2), beryllium (4), and boron (6)
  2. expanded octet = any element in period 3 and greater can hold more than 8 electrons, including phosphorus (10), sulfur (12), chlorine (14), and many others
  3. odd numbers of electrons = any molecule with an odd number of valence electrons cannot distribute those electrons to give eight to each atom (e.g. NO has 11)
58
Q

defn: ionic bonding

A

one or more electrons from an atom with a low ionization energy (typically a metal) are transferred to an atom with a high electron affinity (typically a nonmetal)

the resulting electrostatic attraction between opposite charges is what holds the ions together

59
Q

what type of structure does ionic bonding contribute to?

A

a lattice structure of repeating rows of cations and anions, rather than individual molecular bonds

60
Q

defn: covalent bonding

A

an electron pair is shared between two atoms, typically nonmetals, that have relatively similar values of EN

61
Q

what determines the degree of polarity of a covalent bond?

A

the degree to which the pair of electrons is shared equally or unequally between the two atoms

equally: nonpolar
unequally: polar
if both of the shared electrons are contributed by only one of the two atoms: coordinate covalent

62
Q

does covalent bonding form a structure like how ionic bonds form a lattice?

A

no, covalent compounds consist of individually bonded molecules

63
Q

example + diagram: formation of a covalent bond with Fluorine

A

F has 7 valence electrons

by sharing one electron from each atom, both flourines achieve an octet formation

64
Q

CH 3.2

A
65
Q

defn + char: ionic bonds

A

form between atoms that have significantly different electronegativities (greater than 1.7 on the Pauling scale)

the atom that loses the electrons becomes a CATION
the atom that gains electrons becomes an ANION

66
Q

mnemonic: cations and anions in ionic bonds

A

meTals lose electrons to become caTions = posiTive (+) ions

Nonmetals gain electrons to become aNions = Negative (-) ions

67
Q

char: ionic compounds

A
  1. very high melting points and boiling points (bc of the strength of the electrostatic force between the ionic constituents of the compound)
  2. many readily dissolve in water and other polar solvents
  3. many are good conductors of electricity in the molten or aqueous state
  4. in the solid state: ionic constituents of a compound form a crystalline lattice consisting of repeating positive and negative ions –> attractive forces between oppositely charged ions are maximized and repulsive forces between ions of like charge are minimized
68
Q

defn: intermolecular forces + 3 types

A

weak electrostatic interactions between atoms and compounds

  1. London-dispersion forces
  2. dipole-dipole interactions
  3. hydrogen bond

even the hydrogen bond, which is the strongest of these, has only about 10% of the strength of a covalent bond (thus they can be overcome with small to moderate amounts of energy)

69
Q

defn + expln: London dispersion forces

A
  1. at any point in time, bonding electrons in nonpolar covalent bonds will be located randomly throughout the orbital (in a given moment, the electron density may be unequally distributed between the two atoms)
  2. this results in a rapid polarization and counterpolarization of the electron cloud and the formation of short-lived dipole moments
  3. these dipoles interact with the electron clouds of neighboring compounds, inducing the formation of more dipoles
  4. the momentarily negative end of one molecule will cause the closest region in any neighboring molecule to become temporarily negative, which in turn induces other molecules to become temporarily polarized, and the cycle begins again

THE ATTRACTIVE OR REPULSIVE INTERACTIONS OF THESE SHORT-LIVED AND RAPIDLY SHIFTING DIPOLES ARE KNOWN AS LONDON DISPERSION FORCES (a type of van der Waals force)

70
Q

why are dispersion forces the weakest of all intermolecular interactions?

A

because they are the result of induced dipoles that change and shift moment to moment

71
Q

why are dispersion forces significant only when molecules are in close proximity?

A

because they do not extend over long distances

72
Q

explain how the strength of the London force depends on the degree and ease by which the molecules can be polarized (how easily the electrons can be shifted around)

A

large molecules are more easily polarizable than comparable smaller molecules and thus possess greater dispersion forces

73
Q

give one example in which dispersion forces are crucial

A

without them: noble gases would not liquefy at any temperature because no other intermolecular forces exist between the noble gas atoms

74
Q

defn + expln + implications: dipole-dipole interactions

A
  1. polar molecules tend to orient themselves in such a way that the oppositely charged ends of the respective molecular dipoles are closest to each other: the positive region of one molecule is close to the negative region of another molecule
  2. this arrangement is energetically favorable because an attractive electrostatic force is formed between the two molecules
  3. this attractive force is denoted by dashed lines in most molecular notations and indicates a temporary bonding interaction
  4. present in the solid and liquid phases but become negligible in the gas phase bc of the significantly increased distance between gas particles
  5. polar species tend to have higher melting and boiling points than nonpolar species of comparable molecular weight due to these interactions
75
Q

what is the main difference between London forces and dipole-dipole interactions?

A

they are different not in kind but in duration

both are electrostatic forces between opposite partial charges; the difference is only in the transience or permanence of the molecular dipole

76
Q

defn + expln + char: hydrogen bond

A

a specific, unusually strong form of dipole-dipole interaction that may be intra- or intermolecular

  1. not actually bonds (no sharing or transferring of electrons between two atoms)
  2. when hydrogen is bonded to one of three highly EN atoms (N, O, F), the hydrogen atom carries only a small amount of the electron density in the covalent bond
  3. the hydrogen atom basically acts as a naked proton –> the positively charged hydrogen atoms interacts with the partial negative charge of fluorine, oxygen, or nitrogen on nearby molecules
  4. unusually high boiling points compared to compounds of similar molecular weights that do not exhibit hydrogen bonding (the difference derives from the energy required to break the hydrogen bonds)

SUPER IMPORTANT

77
Q

mnemonic: hydrogen bonds

A

pick up the FON

hydrogen bonds exist in molecules containing a hydrogen bonded to Fluorine, Oxygen, or Nitrogen

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