John-Chapter 3 Chemistry 111
Chapter 3: Bonding - General Concepts
Learning Outcomes
Describe the differences between ionic and covalent bonding and the properties of each type of compound.
Predict relative ion sizes based on ionic radii trends.
Calculate lattice energy using Coulomb’s Law and Hess’s Law.
Name and write formulas for ionic compounds.
Name and write formulas for covalent compounds.
Understand how electronegativity of elements impacts covalent bond polarity.
Draw Lewis structures for covalent compounds.
Draw resonance structures for compounds that have more than one possible Lewis structure.
Calculate the formal charge of elements in a molecule and use it to optimize structure.
Use bond energies to calculate enthalpy changes for chemical reactions.
3-1: Types of Chemical Bonds
Atoms bond to achieve greater stability.
The most stable electron configurations belong to noble gases because they have filled outer (valence) shells:
He: 1s²
Ar: 1s² 2s² 2p⁶ 3s² 3p⁶
Other atoms achieve a stable filled valence shell by:
Sharing or transferring electrons with nearby atoms, forming chemical bonds.
Lewis Dot Symbols
Lewis Dot symbols represent an element's symbol along with its valence shell electrons.
For hydrogen and helium, a full and stable outer shell contains 2 electrons.
For larger elements, a full valence shell typically contains 8 electrons.
Example Lewis dot symbols:
Na: •
H: •
Cl: ••
Achieving Noble Gas Configuration
Ionic Bonds
Lewis Dot diagrams showcase how sodium (Na) and chlorine (Cl) can transfer electrons to achieve noble gas configurations.
The resulting electrostatic attraction between these ions is known as an ionic bond.
Covalent Bonds
Lewis Dot diagrams show how hydrogen (H) and chlorine (Cl) can share electrons to obtain noble gas configurations.
The pair of shared electrons between these two atoms is known as a covalent bond.
Ionic and Covalent Compounds
Ionic Compounds
Examples: NaCl, NaHCO₃
Comprise metals (or NH₄⁺) and nonmetal ions.
Electrons are transferred between atoms.
Features:
Solid, brittle, crystalline.
High melting points.
Dissolve in water to form ions that conduct electricity.
Conduct electricity in molten states.
Covalent Compounds (Molecular)
Examples: H₂O, C₆H₁₂O₆
Comprise 2 or more nonmetals.
Electrons are shared between atoms.
Features:
Can exist as solids, liquids, or gases.
Typically have lower melting and boiling points compared to ionic compounds.
Some dissolve in water, but do not form ions and do not conduct electricity.
3-3: Ions - Electron Configurations and Size Trends
When atoms become ions, they acquire noble gas electron configurations, stabilizing them.
Example for Calcium:
Calcium atom configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
Calcium ion configuration (lost 2 electrons): 1s² 2s² 2p⁶ 3s² 3p⁶
Formula for calcium ion: Ca²⁺
Example for Oxygen:
Oxygen atom configuration: 1s² 2s² 2p⁴
Oxide ion configuration (gained 2 electrons): 1s² 2s² 2p⁶
Formula for oxide ion: O²⁻
Relative Sizes of Ions
Ion size trends:
Increases down a group, decreases left to right across a period.
Exception occurs when cations switch to anions; anions have one more electron shell than cations.
Ranking of ions from smallest to largest:
Al³⁺ < Mg²⁺ < Na⁺ < Cl⁻
Isoelectronic Ions
Isoelectronic ions share the same electron configuration.
Size decreases as atomic number increases due to increased nuclear charge pulling electrons closer to the nucleus.
Example: Na⁺ (smaller than F⁻) despite having the same electron configuration (1s² 2s² 2p⁶).
Ionic Radii Trends
Summary of ionic radii in picometers (pm):
Li⁺: 60 pm, Na⁺: 116 pm, Mg²⁺: 72 pm, Al³⁺: 50 pm, Ca²⁺: 99 pm.
Energy Effects in Binary Ionic Compounds
Ionic bonds form through the attraction of a cation and an anion, resulting in lower energy than separated ions.
Lattice energy (LE) is defined as the energy of attraction between ions, calculable by Coulomb’s Law:
where Q represents the charges of the ions and r is the distance between the ions.
Factors increasing lattice energy:
Higher opposing ion charges (Q1 and Q2).
Decreased ionic size/distance (r).
Determining Lattice Energy with Hess’s Law
Hess’s Law states that the energy change for a reaction is independent of the pathway taken.
Lattice Energy can be inferred using Hess’s Law through a series of measurable reactions.
Example Calculation for Lattice Energy
For lithium fluoride (LiF): From the Hess’s Law equation,
Result:
Naming Simple Compounds
Binary Ionic Compounds
Contain a simple type I cation (same name as element) and an anion (named with “ide”).
No prefixes for subscripts existing; overall compound charge is neutral.
Examples:
LiF: Lithium Fluoride.
CaCl₂: Calcium Chloride.
Al₂O₃: Aluminum Oxide.
Common Monoatomic Ions
Cation Name | Anion Name |
|---|---|
H⁺ Hydrogen | H⁻ Hydride |
Li⁺ Lithium | F⁻ Fluoride |
Na⁺ Sodium | Cl⁻ Chloride |
K⁺ Potassium | Br⁻ Bromide |
Cs⁺ Cesium | I⁻ Iodide |
Be²⁺ Beryllium | O²⁻ Oxide |
Mg²⁺ Magnesium | S²⁻ Sulfide |
Ca²⁺ Calcium | N³⁻ Nitride |
Ba²⁺ Barium | P³⁻ Phosphide |
Al³⁺ Aluminum |
Type II Binary Ionic Compounds
Type II Metals: metals that can have different charge states indicated with Roman numerals following the cation.
Examples:
CuCl₂: Copper (II) Chloride.
Only transition metals except Ag⁺, Zn²⁺, and Cd²⁺ fall under this definition.
Polyatomic Ions in Naming Compounds
Oxyanions: ions with an atom of a given element and varying amounts of oxygen.
Two Oxyanions:
Lesser oxygen: ends in “-ite”.
Greater oxygen: ends in “-ate”.
Example: Chlorate = ClO₃⁻ and Chlorite = ClO₂⁻.
Naming Polyatomic Ions
Prefixes used in multi-oxyanion series:
Hypo- for the fewest oxygens.
Per- for the most oxygens.
Examples:
Perchlorate = ClO₄⁻ and Hypochlorite = ClO⁻.
Common Polyatomic Ions to Memorize
Name | Formula |
|---|---|
Nitrate | NO₃⁻ |
Nitrite | NO₂⁻ |
Carbonate | CO₃²⁻ |
Bicarbonate | HCO₃⁻ |
Chlorate | ClO₃⁻ |
Chlorite | ClO₂⁻ |
Perchlorate | ClO₄⁻ |
Hypochlorite | ClO⁻ |
Sulfate | SO₄²⁻ |
Sulfite | SO₃²⁻ |
Phosphate | PO₄³⁻ |
Phosphite | PO₃³⁻ |
Naming Acids
Binary acids
Contain hydrogen and one other element:
Use the “hydro” prefix, the ion name, and “-ic acid”.
Example: HF = hydrofluoric acid.
Oxoacids
Contain hydrogen, oxygen, and another element.
Use ion name:
“-ate” becomes “-ic acid”; “-ite” becomes “-ous acid”.
Examples: HNO₃ = nitric acid, HNO₂ = nitrous acid.
Strong and Weak Acids
Strong acids dissociate fully in water producing a high concentration of H₃O⁺ ions. Examples include:
HCl, HBr, HI, HNO₃, HClO₄, H₂SO₄
All other acids are classified as weak acids.
Strategy for Naming Compounds
Classify the compound before naming:
Ionic: metal and nonmetals.
Covalent: all nonmetals.
Acids: those with hydrogen or acid in the name.
Hydrates: those with . H₂O present.
Lewis Structures
To predict molecular structure, use the Localized Electron Model, which assumes atoms are bound by shared pairs of electrons.
Procedure for Lewis Structures
Sum the valence electrons.
Draw skeletal structure based on electronegativity (lowest in center).
Use pairs of electrons to form bonds.
Satisfy outer atoms’ octets first, followed by inner.
Use double/triple bonds if necessary to satisfy octets.
Check and minimize formal charges across all atoms.
Formal Charge Calculation
Formal Charge = (Group #) - (Non-bonding electrons) - 1/2 (Bonding electrons)
Exceptions to the Octet Rule
Some elements like Beryllium and Boron do not follow the octet rule, typically holding 4 or 6 electrons.
Elements in period 3 and beyond can expand their octets, utilizing available d orbitals.
Resonance Structures
Certain structures allow for multiple valid Lewis structures; the actual molecule is a resonance hybrid averaging these structures.
Bond Energies
Bond Energy: The energy required to break a chemical bond.
Tabulated values rank bonds: single < double < triple.
Example Bond Energies: