Unit #4 Pt2 Ionic and Molecular Compounds: Covalent Compounds
CHM 100 - Unit #4 Pt2 Ionic and Molecular Compounds: Covalent Compounds
Learning Objectives
Describe the nature of covalent bonds and how they are formed.
Differentiate between ionic and covalent bonds.
Predict the number of covalent bonds an atom will form based on its position in the periodic table and the octet rule.
Distinguish structures, compositions, and properties of molecular compounds from those of ionic compounds.
Interpret molecular formulas and draw Lewis structures for molecules using their molecular formula and the octet rule.
Use Lewis structures to predict molecular geometry and molecular shape.
Distinguish between polar covalent, nonpolar covalent, and ionic bonds using electronegativity.
Predict polarity of molecules using electronegativity and molecular geometry (VSEPR).
Distinguish between intra- and intermolecular forces and determine the forces between different compounds.
Compare relative strength of inter- and intermolecular forces and relate to melting and boiling points.
Covalent Bonds
Covalent bond: A bond formed by sharing electrons between atoms.
Molecular Compounds Form When:
Atoms of two or more nonmetals share electrons to form a covalent bond.
Valence electrons are shared to achieve noble gas electron configuration.
A molecule forms when two or more atoms share electrons, thus being a type of compound.
Covalent Bonding Examples
Hydrogen (H₂):
Spherical 1s orbitals overlap, each contributing 1 electron leading to a 1s² configuration of He for each H atom.
Notations: H-H, H:H, and H₂ represent a hydrogen molecule.
Chlorine (Cl₂):
Overlap of p orbitals with 1 electron contributed by each atom.
Configuration for Cl: 1s² 2s² 2p⁶ 3s² 3p⁶ of Argon.
Notations: Cl-Cl, Cl:Cl, and Cl₂ represent a chlorine molecule.
Diatomic Elements:
Seven diatomic elements include: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂, where electrons are shared equally.
Types of Covalent Bonds
Typical Covalent Bond: Each atom donates an electron to form the bond.
Coordinate Covalent Bond: Both electrons are donated by the same atom. Examples include:
Bonding in polyatomic ions (e.g., NH₄⁺, H₃O⁺).
Formation of acids (e.g., Cl⁻ + H⁺ → HCl).
Characteristics of Coordinate Covalent Bonds: Once formed, they behave like standard covalent bonds but can lead to unusual bonding patterns such as nitrogen with four bonds or oxygen with three (e.g., H₃O⁺).
Covalent vs Ionic Bonds
Covalent Bonds:
Form between nonmetals with shared valence electrons.
Result in molecular compounds or molecules.
Are non-electrolytes.
Ionic Bonds:
Form mainly between metal and nonmetal atoms with a transfer of valence electrons.
Result in the formation of ionic compounds or salts.
Most salts are electrolytes and acids/bases consist of ions.
Molecular Compounds and the Periodic Table
A molecular compound consists of molecules rather than ions.
Covalent Bond Formation: The number of covalent bonds formed by a nonmetal atom usually equals:
The number of electrons it needs to achieve a stable configuration (octet).
Exceptions to the Octet Rule:
Boron: Only shares three electrons to form three covalent bonds leading to six valence electrons.
Elements in rows three and below can utilize vacant d orbitals for bonding, allowing for expanded octets.
Practice Problems
Determine if the following molecules are likely to exist based on covalent bonding rules.
Hydrogen and Fluorine: Show their reaction using electron-dot symbols.
Likely formulas for molecules SiH₂Cl, HBr, PBr.
Examples of unreasonable formulas: N₂, He₂, F₂, O₂.
Discerning least likely existing molecules: BCl₂, CH₂Cl₂, NH₂Cl, H₂Se.
Identify most likely value of x in CHClₓ.
Multiple Covalent Bonds
Single Bond: One pair of shared electrons (2 electrons) - represented as H—H.
Double Bond: Two pairs of shared electrons (4 electrons) - represented as O═O.
Triple Bond: Three pairs of shared electrons (6 electrons) - represented as N≡N.
Certain atoms like C, N, and O frequently form multiple bonds to achieve outer-shell electron octets.
Characteristics of Molecular Compounds
Molecular compounds are neutral and exhibit no strong electrostatic attraction between molecules unlike ionic compounds.
Comparative Properties of Ionic and Molecular Compounds:
Property
Ionic Compounds
Molecular Compounds
Smallest Component
Ions (e.g., Na⁺, Cl⁻)
Molecules (e.g., CO₂, H₂O)
Composition
Metals combined with nonmetals
Nonmetals combined with nonmetals
State at Room Temp
Crystalline solids
Gases, liquids, low-melting solids
Melting Points
High (e.g., NaCl = 801 °C)
Low (e.g., H₂O = 0.0 °C)
Boiling Points
High (e.g., NaCl = 1413°C)
Low and variable
Conductivity
Conducts in molten state or solution
Does not conduct electricity
Solubility
Many are water-soluble
Relatively few are water-soluble
Molecular Formulas and Lewis Structures
A molecular formula represents the number and kinds of atoms in a molecule.
An ionic formula represents the ratio of ions present in an ionic compound.
Drawing Lewis Structures
Lewis Structure: A 2D representation showing atom connections and locations of lone pair valence electrons.
Steps for Drawing Lewis Structures:
Count total valence electrons (subtract for + ions, add for - ions).
Draw lines for bonds between atoms (each line = 2 electrons).
Add lone pairs to complete octets for atoms connected to the central atom (excluding H).
Place any remaining electrons in lone pairs on central atom.
If central atom lacks an octet, shift lone pairs from adjacent atoms to form multiple bonds.
Example Problem:
Draw Lewis structure for hydrogen cyanide (HCN) following above steps including counting electrons, drawing bonds, and forming multiple bonds if needed.
Molecular Geometry and VSEPR Theory
VSEPR Theory (Valence Shell Electron-Pair Repulsion Theory): Determining molecular orientation by minimizing repulsion between electron clouds around central atoms.
Molecular Geometry: Defined by count of atoms attached to the central atom.
Predicting Molecular Shapes Using VSEPR
Draw the Lewis structure.
Identify charge clouds around the atom of interest (shared + lone pairs).
Predict geometry by maximizing distance between charge clouds.
Common Structures: Linear (2 charge clouds), Trigonal Planar (3 charge clouds), Tetrahedral (4 charge clouds).
**Examples:
CO₂ - linear with bond angle 180°.
H₂CO - trigonal planar with bond angle 120°.
H₂O - bent shape due to lone pairs with bond angle about 109°.**
Polar Covalent Bonds and Electronegativity
Electronegativity: Atom's ability to attract shared electrons in covalent bonds.
In a bond between different atoms, electrons may be shared unequally, creating polar covalent bonds.
Polar bonds exist when significant differences in electronegativity between atoms occur, leading to dipoles within the bond structure (e.g., HCl).
Dipole: Separation of charge in a polar bond with directionality toward the more electronegative atom.
Polar molecules may exhibit overall polarity if their shapes do not cancel out the polar bonds, leading to molecular dipoles.
Intermolecular vs Intramolecular Forces
Intermolecular Forces: Weaker forces that influence physical properties such as boiling/melting points.
Types include dipole-dipole, hydrogen bonding, and London dispersion forces.
Intramolecular Forces: Stronger forces that maintain bond within molecules (ionic and covalent bonds).
Higher energy is often needed to break ionic bonds compared to covalent bonds.
Main Types of Intermolecular Forces:
Dipole-Dipole Attraction: Attraction between polar molecules.
Hydrogen Bonds: Strongest for molecules with H atoms bonded to F, O, or N.
London Dispersion Forces: Present in all molecules, predominantly attractive force in nonpolar molecules, and influenced by molecular size.
Implications of Intermolecular Forces on Physical Properties
The strength of intermolecular forces directly correlates to the physical properties (melting/boiling points) of compounds.
Compounds with strong hydrogen bonds (like water) show higher boiling points relative to those with only dispersion forces (like methane).
Practice Problems: To evaluate the intermolecular forces and the correlation of boiling points for different compounds (e.g., HCl, CH₃COOH, CH₄).
Summary of Key Concepts
Interactions among atoms in a compound dictate their behavior and properties.
Understanding types of bonding and how geometry impacts molecule polarity assists in predicting the chemical nature of substances.
Familiarity with intermolecular forces provides insight into the physical properties and reactivity of molecular compounds.
In-Class Questions and Checkpoints
Determine which compounds exhibit specific bonding types (ionic, polar/nonpolar covalent).
Analyze examples of molecular shapes and predict geometries using VSEPR.
Apply electronegativity values to assess bond polarity.
Reference: © 2017 Pearson Education, Inc.