Chapter 10: Chemical Bonding I - The Lewis Model

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Lewis model

A simple model of chemical bonding using diagrams that represent bonds between atoms as lines or pairs of dots. According to this model, atoms bond together to obtain stable octets (eight valence electrons).

Lewis electron-dot structures

A drawing that represents chemical bonds between atoms in molecules as shared or transferred electrons; the valence electrons of atoms are represented as dots in a Lewis structure.

Why do chemical bonds form?

Because they lower the potential energy between the charged particles that compose atoms.

What are the three types of chemical bonds (depending on the kind of atoms involved in the bonding)?

1. Ionic (metal and nonmetal)
2. Covalent (nonmetal and nonmetal)
3. Metallic (metal and metal)

ionic bond

A chemical bond formed between two oppositely charged ions, generally a metallic cation and a nonmetallic anion, that are attracted to each other by electrostatic forces.

covalent bond

A chemical bond in which two atoms share electrons that interact with the nuclei of both atoms, lowering the potential energy of each through electrostatic interactions.

metallic bonding

The type of bonding that occurs in metal crystals, in which metal atoms donate their electrons to an electron sea, delocalized over the entire crystal lattice.

electron sea model

Proposes that all metal atoms in a metallic solid contribute their valence electrons to form a "sea" of electrons, and can explain properties of metallic solids such as malleability, conduction, and ductility.

Lewis symbol

The symbol of an element surrounded with dots representing the element's valence electrons.

octet

The eight dots around atoms in a Lewis structure that signify a filled outer electron shell for s and p block elements.

duet

A Lewis symbol with two dots, signifying a filled outer electron shell for the elements H and He.

chemical bond

The sharing or transfer of electrons to attain stable electron configurations for the bonding atoms.

octet rule

The tendency for most bonded atoms to possess or share eight electrons in their outer shell to obtain stable electron configurations and lower their potential energy.

Do solid ionic compounds contain distinct molecules? Why or why not?

They do not. They are composed of alternating positive and negative ions in a three-dimensional crystalline array.

lattice energy

The energy associated with forming a crystalline lattice from gaseous ions.

Born-Haber cycle

A hypothetical series of steps based on Hess's law that represents the formation of an ionic compound from its constituent elements.

Describe the relationship between ion size and lattice energy.

As ionic radii increase, the ions cannot get as close to each other. Therefore, the lattice energy decreases.

Describe the relationship between ionic charge and lattice energy.

Lattice energies increase with increasing magnitude of ionic charge.

bonding pair

a pair of electrons shared between two atoms

lone pair

a pair of electrons that is not involved in bonding and that belongs exclusively to one atom

nonbonding electrons

pairs of valence electrons on an atom that are not involved in electron sharing

double bond

The bond that forms when two electron pairs are shared between two atoms.

triple bond

The bond that forms when three electron pairs are shared between two atoms. Triple bonds are even shorter and stronger than double bonds.

polar covalent bond

A covalent bond between two atoms with significantly different electronegativities, resulting in an uneven distribution of electron density.

electronegativity

An atom's ability to attract electrons to itself in a covalent bond.

Describe the electronegativity trend across a period in the periodic table.

Increases

Describe the electronegativity trend down a column in the periodic table.

Decreases

What is the most electronegative element?

Fluorine

What is the least electronegative element?

Francium

nonpolar

When two atoms with identical electronegativities form a covalent bond, they share the electrons equally, forming a purely covalent bond.

What is the electronegativity difference that results in a pure covalent bond?

0.0 to 0.4

What is the electronegativity difference that results in a polar covalent bond?

0.4 to 2.0

What is the electronegativity difference that results in an ionic bond?

2.0+

dipole moment (μ)

A measure of the separation of positive and negative charge in a molecule.

percent ionic character

The ratio of a bond's actual dipole moment to the dipole moment it would have if the electron were transferred completely from one atom to the other, multiplied by 100%.

resonance structure

Two or more valid Lewis structures that are shown with double-headed arrows between them to indicate that the actual structure of the molecule is intermediate between them.

resonance hybrid

The actual structure of a molecule that is intermediate between two or more resonance structures.

resonance stabilization

the stabilization associated with the delocalization of electrons via resonance

formal charge

The charge that an atom in a Lewis structure would have if all the bonding electrons were shared equally between the bonded atoms.

Formal Charge formula

FC = #of valence electrons - [lone pair electrons + 1/2bonding electrons]

4 Rules of Formal Charges

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 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.

free radicals

A molecule or ion with an odd number of electrons in its Lewis structure.

incomplete octets

molecules or ions with fewer than eight electrons around an atom

common examples: BF3, BeH2

expanded octets

Occur for non metals in rows 4-7 where the central atoms can hold more than 8 electrons b/c the extra electrons expand/fill into the d-orbital.

Expanded octets NEVER occur in second-period elements.

bond energy

The energy required to break 1 mol of the bond in the gas phase.

What is the relationship between bond strength and bond length?

The stronger the bond, the shorter the bond.

Calculating ΔHrxn from average bond energies

ΔHrxn = (The sum of the ΔH's of bonds broken) + (The sum of the ΔH's of bonds formed)

bonds broken = a positive value
bonds formed = a negative value

Describe an exothermic reaction in terms of bonds broken and bonds formed.

A reaction is exothermic when weak bonds break and strong bonds form.

Describe an endothermic reaction in terms of bonds broken and bonds formed.

A reaction is endothermic when strong bonds break and weak bonds form.

bond length

The average length of a bond between two particular atoms in a variety of compounds.

In general, triple bonds are shorter than double bonds, which are shorter than single bonds.

malleability of metals

their ability to be pounded into sheets

ductility of metals

their ability to be drawn into wires

Which pair of elements is most likely to form an ionic bond?

a. nitrogen and oxygen
b. carbon and hydrogen
c. sulfur and oxygen
d. calcium and oxygen

calcium and oxygen

Which set of elements is arranged in order of increasing electronegativity?

a. O<S<As<Ge
b. Ge<As<S<O
c. S<O<As<Ge
d. As<O<Ge<S

Ge<As<S<O

Which compound is likely to have an incomplete octet?

a. NH3
b. SO3
c. N2O
d. BH3

BH3

Which compound has the lattice energy with the greatest magnitude?

a. MgS
b. CaS
c. SrS
d. BaS

MgS

Which set of compounds is arranged in order of increasing magnitude of lattice energy?

a. CsI<NaCl<MgS
b. NaCl<CsI<MgS
c. MgS<NaCl<CsI
d. CsI<MgS<NaCl

CsI<NaCl<MgS

Which pair of atoms forms the most polar bond?

a. N and O
b. C and O
c. C and F
d. N and F

carbon and fluorine

Which pair of atoms forms a nonpolar covalent bond?

a. C and S
b. C and O
c. B and O
d. Na and Cl

carbon and sulfur

Use bond energies from Table 9.3 to determine ΔHrxn

for the reaction between ethanol and hydrogen chloride.
CH3CH2OH(g)+HCl(g)→CH3CH2Cl(g)+H2O(g)

a. −1549 kJ
b. 1549 kJ
c. −12 kj
d. 12 kj

-12 kJ

Consider the halogenation of ethene, where X is a generic halogen:
H2C=CH2(g)+X2(g)→H2XC—CH2X(g)

Use bond energies to determine which halogen produces the most exothermic halogenation reaction with ethene.

The C—F, C—Br, and C—I bond energies are 552 kJ/mol, 280 kJ/mol, and 209 kJ/mol, respectively. Look up all other necessary bond energies in Table 9.3.

a. fluorine
b. chlorine
c. bromine
d. iodine

fluorine