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 combination of cations (positively charged ions) and anions (negatively charged ions). Proper naming and formula writing are essential for clear chemical communication.

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

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

• Cations with One Charge: Group 1A, 2A metals, Ag+, Zn2+, Al3+ have only one possible charge.

• Cations with Multiple Charges: Many transition metals can have more than one charge; indicate the charge with Roman numerals (e.g., iron(II), iron(III)).

Example: NaCl is sodium chloride; FeCl2 is iron(II) chloride.

Chapter 5.5: Polyatomic Ions

Understanding Polyatomic Ions

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

• Common Polyatomic Ions:

Name	Formula	Charge
Carbonate	CO3	2-
Nitrate	NO3	1-
Hydroxide	OH	1-
Phosphate	PO4	3-
Sulfate	SO4	2-

• Writing Formulas: Use parentheses and subscripts to indicate multiple polyatomic ions (e.g., Ca(NO3)2).

• Naming: Name the cation first, then the polyatomic ion. Polyatomic ions usually 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 stable electron configurations. The shared electrons create an attractive force that holds the atoms together.

• Attractive Forces: Between the nuclei of each atom and the shared electrons.

• Repulsive Forces: Between the positively charged nuclei and between the electrons themselves.

• Lewis Dot Structures: Visual representations showing valence electrons as dots around atomic symbols.

• 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 Bonds: The number of bonds an atom forms is often determined by the number of electrons needed to complete an octet (8 electrons).

Example: Oxygen (6 valence electrons) typically forms two bonds to complete its octet.

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, and ensuring most atoms achieve an octet. Some elements are exceptions.

• 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.

• 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.

• VSEPR Theory: 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.

Example: CH4 (methane) is tetrahedral; H2O (water) is bent.

Chapter 6.5: Naming Covalent Compounds

Binary Covalent Compound Nomenclature

Binary covalent compounds are named using prefixes to indicate the number of each atom present.

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

• Naming: The first element keeps its name; the second element ends in -ide. Prefixes are used for both elements (except "mono-" is often omitted for the first element).

Example: CO2 is carbon dioxide; PCl3 is phosphorus trichloride.

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.

• From Lewis to Skeletal: Convert detailed Lewis structures to line-angle (skeletal) diagrams.

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

• Determining Formula: Count carbons and hydrogens based on skeletal structure conventions.

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 sharing of electrons; non-polar bonds share electrons equally.

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

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

Example: H–Cl is polar; Cl–Cl is non-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).

• Predicting Polarity: Draw the Lewis structure, determine shape, and assess if dipoles cancel.

Example: CO2 has polar bonds but is non-polar overall due to linear shape; H2O is polar.

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):

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

  2. Dipole-Dipole Forces: Occur between polar molecules; result from permanent dipoles aligning.

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

• Physical Properties: IMFs affect boiling/melting points, solubility, and viscosity.

• Biological Importance: Hydrogen bonds shape proteins and DNA.

Example: Water exhibits hydrogen bonding, leading to high boiling point.

Chapter 8.1: Balancing Chemical Equations

Chemical Reactions and Equation Balancing

Chemical reactions involve the transformation of reactants into products, indicated by observable changes such as color, temperature, or gas formation. Chemical equations use formulas and symbols to represent these changes.

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