Chem 101: Ionic and Covalent Compounds, Molecular Polarity, and Chemical Reactions

Chapter 5.4: Naming Ionic Compounds

Overview of Ionic Compound Nomenclature

Ionic compounds are formed from the electrostatic attraction between cations (positively charged ions) and anions (negatively charged ions). Correct naming and formula writing are essential for clear chemical communication.

• Naming from Formula: Identify the cation and anion, name the cation first, then the anion (ending in -ide for monoatomic anions).

• Writing Formulas from Names: Use the charges to determine the correct ratio of ions so the compound is electrically neutral.

• Cations with One Charge: Groups 1A, 2A metals (e.g., Na+, Ca2+) have only one possible charge.

• Cations with Multiple Charges: Many transition metals can have more than one charge; their charge is indicated in Roman numerals (e.g., iron(II) = Fe2+, iron(III) = Fe3+).

Example: FeCl2 is iron(II) chloride; FeCl3 is iron(III) chloride.

Chapter 5.5: Polyatomic Ions

Common Polyatomic Ions and Their Properties

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying a net charge. They are common in many ionic compounds.

• Examples of Polyatomic Ions:

  • Carbonate: CO32−

  • Nitrate: NO3−

  • Hydroxide: OH−

  • Phosphate: PO43−

  • Sulfate: SO42−

• Naming Compounds with Polyatomic Ions: Name the cation first, then the polyatomic ion. Parentheses are used in formulas when more than one polyatomic ion is needed (e.g., Ca(NO3)2).

• Distinguishing Polyatomic from Monoatomic Ions: Polyatomic ions often end in -ate or -ite; monoatomic anions end in -ide.

Example: Iron(II) sulfate = FeSO4; Iron(II) sulfide = FeS.

Chapter 6.1: Covalent Bonds

Formation and Representation of Covalent Bonds

Covalent bonds form when two nonmetal atoms share electrons to achieve a stable electron configuration. The shared electrons create an attractive force that holds the atoms together.

• Attractive Forces: The attraction is between the positively charged nuclei and the shared electrons.

• Lewis Dot Structures: Visual representations showing valence electrons as dots around atomic symbols. Shared pairs (bonds) are shown as lines.

• Single and Multiple Bonds: Single bonds share one pair of electrons; double and triple bonds share two and three pairs, respectively.

• Diatomic Molecules: Seven elements naturally exist as diatomic molecules: H2, N2, F2, O2, I2, Cl2, Br2.

Chapter 6.2: Covalent Bonding and the Periodic Table

Types of Atoms and Valence Electrons

Covalent bonds typically form between nonmetal atoms. The number of valence electrons determines bonding behavior.

• Counting Valence Electrons: Each bond represents two electrons; lone pairs are non-bonding electrons.

• Predicting Number of Bonds: Atoms tend to form enough bonds to achieve an octet (8 electrons) in their valence shell (except H, which needs 2).

Chapter 6.3: Drawing Covalent Compounds Using Lewis Structures

Lewis Structure Construction and Octet Rule Exceptions

Lewis structures are drawn by arranging atoms, distributing valence electrons as bonds and lone pairs, and ensuring octet fulfillment where possible.

• Steps: (1) Count total valence electrons, (2) Arrange atoms, (3) Connect with single bonds, (4) Complete octets with lone pairs, (5) Use multiple bonds if needed.

• Octet Rule Exceptions: Boron (6 electrons), Hydrogen (2 electrons), Phosphorus and Sulfur (can have expanded octets in acids).

Chapter 6.4: Predicting the Shapes of Molecules

VSEPR Theory and Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron pair repulsion around a central atom.

• Key Principle: Electron pairs (bonding and lone pairs) arrange themselves to minimize repulsion.

• Predicting Geometry: Draw the Lewis structure, count bonding groups and lone pairs, and use a geometry table to determine shape and bond angles.

• Larger Molecules: Predict geometry around each central atom individually.

Chapter 6.5: Naming Covalent Compounds

Binary Covalent Compound Nomenclature

Binary covalent compounds (two nonmetals) are named using prefixes to indicate the number of each atom.

• Prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, etc.

• Naming Order: The less electronegative element is named first; the second element ends in -ide.

Chapter 6.6: Skeletal Structures

Interpreting and Drawing Skeletal Structures

Skeletal structures are simplified representations of organic molecules, showing bonds between carbon atoms as lines and omitting hydrogen atoms bonded to carbon.

• Conversion: Given a Lewis structure, draw the skeletal structure by representing carbon chains as zig-zag lines.

• Counting Atoms: Each vertex or line end represents a carbon; hydrogens are implied to fill carbon's four bonds.

Chapter 7.1: Bond Polarity and Electronegativity

Polarity, Electronegativity, and Dipoles

Bond polarity arises from differences in electronegativity between bonded atoms. Electronegativity is an atom's ability to attract shared electrons.

• Polar vs. Non-polar Bonds: Polar bonds have unequal electron sharing; non-polar bonds share electrons equally.

• Dipole: A separation of charge within a bond, indicated by a dipole arrow (→) and partial charges (δ+ and δ−).

• Predicting Polarity: Use electronegativity values to determine if a bond is polar.

Chapter 7.2: Polarity of Smaller Molecules

Determining Molecular Polarity

The overall polarity of a molecule depends on both the polarity of its bonds and its molecular geometry.

• Key Factors: (1) Presence of polar bonds, (2) Shape of the molecule (symmetry can cancel dipoles).

• Prediction: Draw the Lewis structure, determine shape, and assess if dipoles add or cancel.

Chapter 7.3: Polarity of Larger Molecules

Organic Molecules, Hydrocarbons, and Alcohols

Organic molecules contain carbon and hydrogen, often with other elements. Biomolecules are a subset of organic molecules.

• Hydrocarbons: Compounds of only carbon and hydrogen; generally non-polar.

• Alcohols: Organic molecules with an -OH group. Small alcohols (e.g., methanol, ethanol) are polar; larger alcohols become less polar as the non-polar carbon chain increases.

Chapter 7.4: Intermolecular Forces (IMFs)

Types and Effects of Intermolecular Forces

Intermolecular forces are attractions between molecules, influencing physical properties like boiling and melting points.

• Types of IMFs (from weakest to strongest):

  • London Dispersion Forces (LDFs): Present in all molecules, especially non-polar; arise from temporary dipoles.

  • Dipole-Dipole Forces: Occur between polar molecules due to permanent dipoles.

  • Hydrogen Bonding: Strong dipole-dipole interaction when H is bonded to N, O, or F.

• Significance: Hydrogen bonds shape important biomolecules like DNA and proteins.

• Recognition: Analyze molecular structure to determine which IMFs are present.

Chapter 8.1: Balancing Chemical Equations

Chemical Reactions and Equation Balancing

Chemical reactions involve the transformation of reactants into products, indicated by observable changes (color, gas, precipitate, energy).

• Chemical Equation Structure: Reactants on the left, products on the right, separated by an arrow.

• Coefficients: Numbers placed before formulas to balance the number of atoms of each element on both sides.

• Balancing Steps: (1) Write correct formulas, (2) Count atoms, (3) Adjust coefficients to balance, (4) Check work.

Example Equation:

$\mathrm{2H_2 + O_2 \rightarrow 2H_2O}$