Chapter 9: Models of Chemical Bonding

Merged flashcards from Chapter 9 of Chemistry: The Molecular Nature of Matter and Change.

Bonding

This lowers the potential energy between positive and negative particles

• Can be imperfectly modeled with classifications

Ionic bonding

This occurs between atoms with large differences in their tendencies to lose or gain electrons

• Often occurs between reactive metals and nonmetals to form a noble gas configuration for a solid

Covalent bonding

This occurs between atoms with small differences in their tendencies to lose or gain electrons

• Often occurs with electron-sharing nonmetals as separate molecules form

Metallic bonding

This occurs between metals when electrons are shared, with electrons being “pooled” (delocalized) due to low electron affinity that creates attractions between the positive nuclei

Lewis electron-dot symbol

A way of representing an atom through drawing a symbol that represents the nucleus and core electrons with dots representing the reactive valence electrons

• Dots are drawn one at a time on each side until all valence electrons are represented, demonstrating a target towards the octet rule

  • Metals: Total number of dots is the number needed to form a cation

  • Nonmetals: Number of unpaired dots is the number needed to form an anion, or shared for covalent bonds

Octet rule

A generalization of atomic bonding that states that atoms lose, gain, or share electrons to attain a filled outer level of eight electrons

Energy transfer

This occurs during ion formation

• Energy input is required to break electron bonds

• Energy output as heat occurs as bonds are formed

Energy release

This occurs when a compound forms from its elements in ionic bonding, due to a decrease in potential energy

Energy absorption

This occurs when an atom loses an electron or changes into a greater energy state, due to an increase in potential energy from the input needed

Lattice energy (ΔH°_lattice)

The enthalpy change associated with one mol of ionic solid separating into gaseous ions

• Equal to the opposite of the last step of ionic bond formation

Born-Haber cycle

A series of steps and enthalpy changes that converts elements to an ionic compound, used to calculate the lattice energy from known enthalpies of formation

• Based on Hess’s law, stating that ΔH°_f = ΣΔH°_steps

• Requires adding individual ΔH values in order to form elements, then solving for the last step to obtain lattice energy

1. List out your given ΔH°_f values.

2. Determine the order and sign of the individual steps to add.

3. Using Hess’s law, solve for -ΔH°_lattice (the last step).

4. Take the opposite of the last step to find ΔH°_lattice.

How do I solve for an equation using the Born-Haber cycle?

Coulomb’s law

States that the electrostatic energy between particles is directly proportional to the product of their charges and inversely proportional to the distance between them

• Expressed as E = (Charge A ⋅ Charge B) / Distance

  • Greater charges, then smaller ion sizes, increase energy

• Directly proportional to ΔH°_lattice

Ionic bonding model

This explains the properties of ionic solids, as:

• Ions in specific positions make for a hard, brittle, and rigid solid

• Ions dissolve and carry a current in water

• Ions require large amounts of energy for separation due to high lattice energies, creating ion pairs in gaseous states instead of individual ions

Covalent bond

A bond where two atoms share electrons, arising from the repulsing nuclei and attracting electrons

• Results in greater electron density between nuclei

To form the bonds, energy from the atoms needs to be released, with some being released by heat through attraction.

Why does bond formation release energy?

To break the bonds, energy from the surroundings needs to be absorbed, enabling ionization.

Why does bond breakage absorb energy?

Bonding pair (shared pair)

The pair of electrons shared by atoms in a bond

• Represented by a pair of dots or a line

Lone pair (unshared pair)

A pair of outer-level electrons not involved in bonding

• Generally shown as dots

Bond order

The number of electron pairs being shared by a given pair of atoms

• Higher numbers result in greater bond energy and lower bond length

Single bond

One bonding pair of electrons with two electrons shared

• Has a bond order of 1

Double bond

Two bonding pairs of electrons with four electrons shared

• Has a bond order of 2

Triple bond

Three bonding pairs of electrons with 6 shared electrons

• Has a bond order of 3

Bond energy (BE)

The input energy needed to break apart the bond of 1 mol of a gaseous compound

• Stronger bonds have higher levels due to lower potential energy present (spent during formation)

• Weaker bonds have lower levels due to higher potential energy present (reserved with fewer formations)

Bond length

The distance between the nuclei of two bonded atoms at a lowest energy point

• Equal to the sum of radii

• Inversely correlated to bond order and energy

Covalent bonding model

This explains the properties of covalent substances, with:

• Low melting and boiling points, given weak intermolecular forces and strong bonding forces

• Strong networked bonds in solids like quartz and diamond which do use covalent strength

• Little conductivity due to localized electrons and few ions

Intermolecular forces

Unlike in ionic bonds, these are weak in covalent bonds, due to bonding forces being strong instead

• Quartz and diamond are two exceptions

Electrical conductivity

Unlike in ionic bonds, this is weak for covalent substances, due to localized electrons and a lack of ions

Infrared spectroscopy (IR spectroscopy)

The observation of IR wavelengths passing through molecules as a fingerprint

• Captures the vibration of atoms in different bonds as they stretch and bend bonds

Enthalpy of reaction (ΔH°_rxn)

This arises from the differences between reactant and product bond energies

• Defined as ΣΔH°_{bonds broken} + ΣΔH°_{bonds formed}

  • Also defined as ΣBE_{bonds broken} - ΣBE_{bonds formed}

• Results from absorption during breakage, and emittance during formation, using Hess’s Law

Hess’s Law

States that the enthalpy of a reaction is equal to the sum of the enthalpies taken to get to an end point regardless of method

• Enables the calculation of ΔH°_rxn through the breakage and formation of all bonds to obtain an average

Fuel

Any material that reacts with O₂ to release energy, featuring weak bonds

• Less energy is absorbed in breakage than released in formation

• Energy yield depends on number of O bonds present

Strong bonds

The presence of these results in a stronger material as more energy is needed to break them

• Seen in forever chemicals, preventing their breakdown in nature

Electronegativity (EN)

The relative ability of a bonded atom to attract shared electrons

• Inversely related to atomic size due to greater nuclear attraction

• Can be compared in 0.5 increments across period two, starting with Li (1.0) to F (4.0)

  • Hydrogen (H) has a value of 2.1

  • Halogens descend: F (4.0), Cl (3.0), Br (2.8), and I (2.5)

4.0

What is the EN value of F?

3.0

What is the EN value of Cl?

2.8

What is the EN value of Br?

2.5

What is the EN value of I?

Electronegativity difference (ΔEN)

The difference between the electronegativity values of the bonded atoms

• Greater levels equate to a higher partial ionic character due to charge separation

  • Nonpolar covalent bonds have a value of 0.0

  • Mostly covalent bonds have a value from 0.0 to 0.4

  • Polar covalent bonds have a value from 0.4 to 2.0

  • Mostly ionic bonds have a value from 2.0 to 3.3

Nonpolar covalent bond

The equal sharing of electrons in a covalent bond with an electronegativity difference of 0.0

• Seen with identical atoms or atoms with close electronegativity values

Mostly covalent bond

The roughly equal sharing of electrons in a covalent bond with an electronegativity difference between 0.0 and 0.4

Polar covalent bond

The unequal sharing of electrons in a covalent bond with an electronegativity difference between 0.4 and 2.0

• Denoted by a polar arrow pointing towards the partially negative pole, or with symbols δ+ and δ-

Mostly ionic bond

The unequal sharing of electrons in a bond with an electronegativity difference greater than 2.0

• Range is capped at 3.3, between Cs (0.7) and F (4.0)

Percent ionic character

A number quantifying the level of sharing between atoms, increasing with electronegativity difference

• A value at 50% divides ionic and covalent bonds

• A value of 100% is not possible due to some electron sharing always occurring