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, 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: 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 multiple charges. Indicate the charge with Roman numerals in parentheses (e.g., iron(II) for Fe2+, iron(III) for Fe3+).

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-

• Formulas with Polyatomic Ions: Use parentheses when more than one polyatomic ion is needed (e.g., Ca(NO3)2).

• Naming: Name the cation first, then the polyatomic ion. Polyatomic ions usually end in -ate or -ite, while 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 shared electrons and the nuclei of both atoms.

• Repulsive Forces: Between the nuclei and between the electrons.

• 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.

• Valence Electrons: Count total valence electrons using group numbers and sum for all atoms in the molecule.

• Predicting Bonds: The number of bonds = 8 - 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 stepwise diagrams showing how atoms share electrons. Some elements do not follow the octet rule strictly.

• Drawing Steps: (Reference sheet will provide steps; generally, count electrons, connect atoms, complete octets, and check for multiple bonds.)

• Octet Rule Exceptions:

  • Boron (B): Often has only 6 electrons.

  • Hydrogen (H): Only 2 electrons (duet rule).

  • Phosphorus (P) and Sulfur (S): Can have expanded octets (more than 8 electrons) in some compounds, especially acids.

Chapter 6.4: Predicting the Shapes of Molecules

VSEPR Theory and Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the 3D shapes of molecules based on electron pair repulsion around a central atom.

• VSEPR Theory: Electron pairs (bonding and lone pairs) arrange themselves to minimize repulsion, determining molecular shape.

• Predicting Geometry: Draw the Lewis structure, count bonding groups and lone pairs, and use a geometry table to assign 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 are named using prefixes to indicate the number of each atom present. The more electronegative element is named last, ending in -ide.

• Prefixes: mono-, di-, tri-, tetra-, penta-, etc. (A table will be provided on exams.)

• Namingul Steps: (Not provided; refer to reference sheet.)

Chapter 6.6: Skeletal Structures

Interpreting and Drawing Skeletal Structures

Skeletal structures are simplified line drawings representing organic molecules, showing bonds between carbon atoms and omitting hydrogen atoms bonded to carbon.

• From Lewis to Skeletal: Convert detailed Lewis structures to line-angle representations.

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

• Determining Formula: Count carbons and hydrogens to write the molecular formula.

Chapter 7.1: Bond Polarity and Electronegativity

Understanding Bond Polarity

Bond polarity arises from differences in electronegativity between bonded atoms, resulting in unequal sharing of electrons.

• Electronegativity: The ability of an atom to attract shared electrons in a bond.

• 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 (table provided) 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:

  • Presence of polar bonds

  • Shape of the molecule (geometry)

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

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 (e.g., methanol, ethanol) are polar; larger alcohols become less polar as the carbon chain increases.

Chapter 7.4: Intermolecular Forces (IMFs)

Types and Effects of Intermolecular Forces

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

• Types of IMFs (from weakest to strongest):

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

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

  • Hydrogen Bonding: A strong dipole-dipole interaction when H is bonded to N, O, or F; significant in water, proteins, and DNA.

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

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

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, temperature, or gas formation. Chemical equations represent these changes symbolically.

• 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 of the equation.

• Balancing Steps: (Refer to reference sheet for detailed steps.)

Example: The reaction of hydrogen and oxygen to form water:

$2\mathrm{H}_2 + \mathrm{O}_2 \rightarrow 2\mathrm{H}_2\mathrm{O}$

Additional info: Balancing ensures the law of conservation of mass is obeyed in chemical reactions.