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

Chapter 5.4: Naming Ionic Compounds

Identifying and Naming Ionic Compounds

Ionic compounds are formed from the electrostatic attraction between cations (positively charged ions) and anions (negatively charged ions). Correctly naming and writing formulas for these compounds is essential in chemistry.

• Naming from Formula: Name the cation first, followed by the anion. For metals with only one possible charge (e.g., Na+, Ca2+), simply use the element name. For metals with multiple possible charges (transition metals), indicate the charge with Roman numerals in parentheses (e.g., iron(II), copper(II)).

• Writing Formulas from Names: Write the symbol for the cation and anion, then balance the charges to determine subscripts.

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

• Cations with Multiple Charges: Many transition metals (e.g., Fe, Cu, Pb) can have more than one charge; specify with Roman numerals.

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:

  • Carbonate: CO32−

  • Nitrate: NO3−

  • Hydroxide: OH−

  • Phosphate: PO43−

  • Sulfate: SO42−

• 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; 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.

• Single and Multiple Bonds: Single bonds share one pair of electrons, double bonds share two, and triple bonds share three pairs. Represented by lines: one line for a single bond, two for double, three for triple.

• 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) forms two bonds to complete its octet.

Chapter 6.3: Drawing Covalent Compounds Using Lewis Structures

Lewis Structures and Octet Rule Exceptions

Lewis structures help visualize the arrangement of atoms and electrons in a molecule. Most atoms follow the octet rule, but there are exceptions.

• Drawing 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 the 3D shape of molecules based on electron pair repulsion.

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

• 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: Methane (CH4) is tetrahedral; water (H2O) is bent.

Chapter 6.5: Naming Covalent Compounds

Binary Covalent Molecules

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

• Naming: Prefix + first element name + prefix + second element name (ending in -ide).

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

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

Chapter 6.6: Skeletal Structures

Drawing and Interpreting Skeletal Structures

Skeletal structures are simplified representations of organic molecules, showing the carbon backbone and functional groups.

• From Lewis to Skeletal: Convert detailed Lewis structures to lines representing carbon chains; hydrogens on carbons are usually omitted.

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

Example: A zig-zag line with three vertices represents propane (C3H8).

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.

• Electronegativity: The greater the difference, the more polar the bond.

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

Example: H–Cl is polar; Cl is more electronegative, so H is δ+ and Cl is δ−.

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 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 (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 due to electron movement.

  • Dipole-Dipole Forces: Occur between polar molecules; positive end of one molecule attracts negative end of another.

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

• Physical Properties: Stronger IMFs lead to higher boiling/melting points.

Example: Water's high boiling point is due to hydrogen bonding.

Chapter 8.1: Balancing Chemical Equations

Chemical Reactions and Equation Balancing

Chemical reactions involve the transformation of reactants into products, often indicated by color change, gas formation, or precipitate formation. Chemical equations represent these changes using formulas and symbols.

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

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