Chapter 09: Chemical Bonding I: The Lewis Model Flashcards Preview

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Flashcards in Chapter 09: Chemical Bonding I: The Lewis Model Deck (29)
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Covalent bond

Electrons shared

Nonmetal + nonmetal

Usually polar, sometimes nonpolar


Metallic bonding

Electrons pooled ("sea of electrons")

Metal + metal

Results from attraction of cations to delocalized electrons


Ionic bond

Electrons transferred

Oppositely charged ions bond, lowering overall potential energy


Lewis structure of ionic bond

Cation + anion in squared brackets


Exceptions to octet rule

1. Duet rule (H, Li, Be, B -- first four elements)
2. Expanded octet (period 3 and below)


Lattice energy

Exothermic energy associated with formation of crystalline lattice

Electrostatic attraction is nondirectional (no direct anion-cation pair)

Hence, no ionic molecule

Depends directly on size;
inversely on distance between ions


Born-Haber cycle

Hypothetical series of steps representing the formation of an ionic compound from constituent elements

Change in enthalpy known for each step except last (for lattice energy)

Utilize Hess' law to determine lattice energy enthalpy change


Ion size & lattice energy

Force of attraction inversely proportional to distance between them

Larger ionic radius = weaker attraction = smaller lattice energy

Less exothermic with increasing ionic radius


Ion charge & lattice energy

More exothermic with increasing ionic charge

Force of attraction directly proportional to product of charges

Larger charge = stronger attraction = larger lattice energy

*Generally more important than ion size


Bonding pair

Electrons that are shared by atoms


Lone pair

Nonbonding pair

Electrons not shared by atoms


Polar covalent bond

Covlanet bond between unlike atoms

Unequal sharing of electrons

One atom pulls electrons in bond closer to its side




Ability of an atom to attract electrons to itself in a chemical bond

Measure of EN is relative

EN increases toward the right of a period and up a group

F = most EN; Fr = least EN

Noble gases = NO EN

EN inversely related to atomic size


Pure/nonpolar covalent bond

Electrons shared equally (between atoms with like EN)


Dipole moment

A measure of bond polarity

i.e. the polarity of a bond between two atoms within a molecule or ion

Directly proportional to size of particle charges and distance between charges

Represented with vector arrow

The more electrons shared & the larger the atoms are, the larger the dipole moment


Percent ionic character

Percentage of a bond's measured dipole moment compared to what it would be if the electron were completely treansferred

Indicates the degree to which the electron is transferred


Writing Lewis structures (4)

1. Write skeletal structure (H is always terminal; more electronegative elements are terminal)
2. Calculate total # of valence electrons
3. Distribute electrons
4. If any atoms lack an octet, form double or triple bonds as necessary


Formal charge (4)

Charge an atom would have if all bonding electrons were shared equally

FC of an atom = Valence e- - nonbonding e- - (1/2)(bonding e-)

1. Sum of all formal charges in neutral molecule = 0
2. Sum of all formal charges in ion = ion charge
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


Resonance structures

Based on delocalization of electrons which stablilizes molecule/ion

Multiple correct Lewis structures that average out to what occurs in nature

Same atomic structure, different bond arrangement

*Formal charges must total the same in each resonance structure


Resonance hybrid

A combination of resonance forms that best represents that actual molecule (cannot be drawn as one structure)


Free radicals

AKA odd-electron species

Molecules and ions with an odd number of electrons in Lewis structures

Reactive because of odd number of electrons



Incomplete octet

Applies to first five elements: H, He, Li, Be, B


Expanded octets

Applies to elements in third row and below

Up to 12 (sometimes 14) electrons

Stored in d orbitals


Bond energy

The amount of energy required to break one mole of a bond in the gas phase

*Average bond energies can be used to estimate ΔHrxn

ΔHrxn = Σ(ΔH bonds broken) + Σ(ΔH bonds formed)

broken = positive

formed = negative


Periodic trends in bond energies (4)

1. The more electrons shared between two like atoms, the stronger the bond
2. The shorter the bond, the stronger the bond
3. Bonds get stronger up a group
4. Bonds get stronger toward the right of a period


Bond length

Distance between nuclei of bonded atoms


Average bond length

Average for similar bonds from many compounds

Used because the actual bond length depends on the other atoms around the bond


Periodic trends in bond lengths (4)

1. The more electrons shared, the shorter the bond
2. The longer the bond, the weaker
3. Bond length increases to the left of a period
4. Bond length increases down a group


Electron sea model

Simplest theory of metallic bonding

Metal atoms release valence electrons to be shared as a pool