Chem 101 Exam 2 Study Guide: 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: Main group metals (e.g., Na+, Ca2+) have only one possible charge and are named by the element name.

• Cations with Multiple Charges: Transition metals can have more than one charge; indicate the charge with Roman numerals (e.g., iron(II) for Fe2+, iron(III) for Fe3+).

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

Chapter 5.5: Polyatomic Ions

Understanding Polyatomic Ions

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying an overall 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 around polyatomic ions when more than one is needed (e.g., Ca(NO3)2).

• Naming: Name the cation first, then the polyatomic ion.

• Distinguishing Names: 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: Between the positively charged nuclei and the shared electrons.

• Repulsive Forces: Between the 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

Patterns in Covalent Bonding

Covalent bonds typically form between nonmetal atoms. The number of bonds an atom forms is related to its need to achieve an octet (8 valence electrons).

• Types of Atoms: Nonmetals bond covalently with other nonmetals.

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

• Predicting Bonds: The number of bonds = 8 minus the number of valence electrons (except for hydrogen, which forms one bond).

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 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).

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 the repulsion between electron groups around a central atom.

• VSEPR Theory: Electron groups (bonds 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: Prefix + first element, prefix + second element (ending in -ide).

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 structures for clarity.

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

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

Chapter 7.1: Bond Polarity and Electronegativity

Understanding Bond Polarity

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

• Polar Bond: Unequal sharing of electrons; one atom is partially negative (δ-) and the other partially positive (δ+).

• Non-polar Bond: Equal sharing of electrons.

• Dipole: A separation of charge within a bond, indicated by an arrow (→) pointing toward the more electronegative atom.

• 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.

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

Chapter 7.3: Polarity of Larger Molecules

Organic Molecules and Biomolecules

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

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

• Alcohols: Organic molecules with an -OH group; small alcohols are polar, but larger ones become less polar as the non-polar carbon chain increases.

Chapter 7.4: Intermolecular Forces (IMFs)

Types and Importance of IMFs

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

• Types of IMFs (from weakest to strongest):

Type	Description	Occurs In
London Dispersion Forces (LDFs)	Temporary dipoles due to electron movement	All molecules, especially non-polar
Dipole-Dipole Forces	Attraction between permanent dipoles	Polar molecules
Hydrogen Bonding	Strong dipole-dipole involving H bonded to N, O, or F	Molecules with H-N, H-O, or H-F bonds

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

• Recognition: Analyze molecular structure to determine possible IMFs.

Chapter 8.1: Balancing Chemical Equations

Understanding and Balancing Chemical Reactions

Chemical reactions involve the transformation of reactants into products, indicated by observable changes such as color change, gas formation, or precipitate formation.

• Chemical Equation: Uses formulas and symbols to represent a reaction: 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 of the equation.

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

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