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chem 101 > chapter 9 > Flashcards

chapter 9 Flashcards

1
Q

why do chemical bonds form?

A

Chemical bonds form in
order to lower the
potential energy of an
atom

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2
Q

when does a chemical bond form?

A 
A chemical bond forms
when the potential
energy of the bonded
atoms is lower than the
potential energy of the
atoms alone
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3
Q

when do bonds form?

A

Bonds form when
electrons are
simultaneously attracted
to two nuclei

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4
Q

what is lewis theory based on?

A

-Simplest of bonding theories
-Based on the filling of valence shell by sharing
electrons

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5
Q

what 2 things does valence bond theory consider?

A
  • Considers orbital overlap

- Hybridization of molecular orbitals

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6
Q

what is molecular orbital theory based on?

A
  • Most complex

- Based on the mixing of molecular orbitals from atoms

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7
Q

types of atoms for ionic bonds? characteristic?

A
  • metal and nonmetal

- electrons transferred

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8
Q

types of atoms for covalent bonds? characteristic?

A

nonmetal and nonmetal

electrons shared

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9
Q

types of atoms for metallic bonds? characteristic?

A

metal and metal

electrons pooled

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10
Q

how does lewis theory describe bonding?

A

Chemical bonds form to share or transfer electrons to
attain a stable electron configuration
• We need to consider valence electrons, those in the
outermost principle energy level

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11
Q

how are valence electrons represented in lewis models?

A

We represent valence electrons with dots surrounding

the atom

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12
Q

what is the octet rule?

A

The octet rule states that bonded atoms want to have a
stable electron configuration with 8 electrons in the
outermost shell to lower their potential energy

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13
Q

how do lewis structures exist for ionic compounds?

A

Electrons are not shared in ionic compounds

• Need to transfer the electron(dot) to the non-meta

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14
Q

what is lattice energy?

A

There is energy associated with forming a crystalline

lattice of alternating cations and anions

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15
Q

how can we calculate lattice energy?

A

the Born-Haber cycle

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16
Q

ion size trend in lattice energy

A

Lattice energies become less exothermic with increasing

ionic radius

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17
Q

ion charge trend in lattice energy?

A

Lattice energies become more exothermic with

increasing magnitude of ionic charge

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18
Q

coloumbs law is

A

the reason for these trends

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19
Q

compare ions by size etc

A

qualitative

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20
Q

what is the relationship between q1q2 and lattice energy?

A

as q1q2 increases, lattice energy becomes more exothermic

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21
Q

melting and boiling points of ionic bonds? tendency to conduct?

A

!high melting and boiling points
!tendency to not conduct electricity when solid
!tendency to conduct electricity when dissolved

22
Q

what is a bonding pair?

A

Bonding pair: shared pair of electrons

23
Q

non bonding pair?

A

Non-bonding pair: lone pair, pair of electrons associated

with only one atom - non-bonding electrons

24
Q

how can two atoms attain an octet?

A

Two atoms can share more than one electron pair to
attain an octet
!can form single, double, triple bonds (2,4,6 shared
electrons or 1,2,3 shared pairs)

25
Q

what does the length of a bond depend on?

A

atoms involved and type of bond

26
Q

bond lengths for bonds. what can we predict from bond length?

A

triple bond

27
Q

which bonds are stronger?

A

!Generally, the shorter the bond, the stronger it is

28
Q

5 steps to draw lewis structures

A
  1. Calculate the total number of electrons by summing the
    valence electrons of each atom in the molecule
  2. Write the correct skeletal structure for the molecule drawing
    a single bond between each bonding atom
    • Central atom is generally the one with lower group
    number
  3. Distribute the remaining unaccounted-for electrons in pairs,
    giving octets to as many atoms as possible (except
    hydrogen)
    • Start with the terminal atoms first
  4. If any non-hydrogen atoms do not have an octet form
    double or triple bonds
  5. Charged molecules have square brackets around them with
    the net charge as the superscript
29
Q

3 guidelines for drawing lewis structures

A

!Hydrogens are always in terminal positions
! More electronegative atoms are also terminal
! Less electronegative atoms are in central positions

30
Q

which element is the most electronegative?

31
Q

what bond type has small (0-0.4) electronegativity difference?

32
Q

what bond type has medium (0.4-2.0) electronegativity difference?

A

polar covalent

33
Q

what bond type has large (2.0+) electronegativity difference?

34
Q

what are resonance structures?

A

have more than one Lewis
structure, which has the same connectivity, but different
electron arrangements

35
Q

what is formal charge?

A

charge an atom would have if all

electrons were shared equally

36
Q

why do we need to determine formal charges?

A

Need to determine the formal charge if you are to

predict the correct Lewis structure

37
Q

formal charge formula

A

Formal Charge= # of valence e-

in free atom - (number of nonbonding e - + 1/2 # of bonding e -)

38
Q

what is electronegativity?

A
  • the ability of an element to draw
    electrons to itself in a covalent bond
    • Electronegativity is measured on a scale, where 4.0 is
    the most electronegative (F) and all other elements are
    measured relative to it
39
Q

are electrons shared equally between atoms? what is a polar covalent bond?

A

no
Polar covalent bond - a bond with unequally shared
electrons

40
Q

how many electrons can an expanded octet have? which elements are likely to have one? what orbitals are used?

A

up to
12 electrons
• Usually elements in the 3rd period and below
• Use empty d-orbitals

41
Q

what are the exceptions to the octet rule?

A

Odd electron species

Incomplete octets

42
Q

describe an odd electron species

A

will have one unpaired electron,

called a radical

43
Q

describe an incomplete octet

A

When elements have less than 8

valence electrons they may not form a complete octet

44
Q

formal charge rules

A
  1. The sum of all formal charges in a neutral
    molecule must be zero
  2. The sum of all formal charges in an ion must be
    equal to the charge of the ion
  3. Small (or zero) formal charges on atoms are
    better than large ones
  4. When formal charge cannot be avoided,
    negative formal charges should reside on the
    most electronegative atom
  5. Structures having formal charges of the same
    sign on adjacent atoms are unlikely
45
Q

are all resonance structures equal?

A

all resonance structures are equally good. equally contribute to the actual bonding structure of NO3-

46
Q

what is the rule for the sum of all formal charges in a neutral molecule

A

The sum of all formal charges in a neutral

molecule must be zero

47
Q

what is the rule for the sum of all formal charges in an ion?

A

The sum of all formal charges in an ion must be

equal to the charge of the ion

48
Q

which charges on atoms are better?

A

Small (or zero) formal charges on atoms are

better than large ones

49
Q

what happens when formal charge cannot be avoided?

A

When formal charge cannot be avoided,
negative formal charges should reside on the
most electronegative atom

50
Q

which structures having formal charges are unlikely?

A

Structures having formal charges of the same

sign on adjacent atoms are unlikely

chem 101 (13 decks)

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