CHEM 1070 – General Chemistry I: Practice Exam Study Notes

Atomic Structure and Ions

Charge of Ions in Compounds

Atoms can gain or lose electrons to form ions. The charge of an ion in a compound is determined by the number of electrons lost (positive charge, cation) or gained (negative charge, anion).

• Example: In SnO, tin (Sn) forms a +2 ion because oxygen (O) typically forms a -2 ion, and the compound must be electrically neutral.

Intermolecular Forces

Types of Intermolecular Forces

Intermolecular forces are attractions between molecules that affect physical properties like boiling and melting points.

• Dispersion forces: Weakest, present in all molecules, especially nonpolar.

• Dipole-dipole attractions: Occur between polar molecules.

• Hydrogen bonding: Strong type of dipole-dipole, occurs when H is bonded to N, O, or F.

• Ionic bonding: Strongest, between ions of opposite charge.

Example: The strongest force in liquid NF3 is dipole-dipole attractions, as it is a polar molecule but does not have hydrogen bonding.

Covalent Bonds and Polarity

Non-Polar Covalent Bonds

A covalent bond is non-polar if the two atoms have similar or identical electronegativities, resulting in equal sharing of electrons.

• Examples of non-polar bonds: C-H, C-S, F-F

• Polar bonds: O-O is non-polar, but O-H is polar due to a large difference in electronegativity.

Stoichiometry and Chemical Equations

Balancing Chemical Equations

Balancing ensures the same number of each atom on both sides of the equation, following the Law of Conservation of Mass.

• Example: 4NH3 + 3O2 → 2N2 + 6H2O

Stoichiometric Calculations

Stoichiometry involves using balanced equations to calculate amounts of reactants and products.

• Mole-to-mole conversions: Use coefficients from balanced equations.

• Mole-to-mass or mass-to-mole: Use molar mass ($\text{g/mol}$).

• Mole-to-particle: Use Avogadro's number ($6.022 \times 10^{23}$ particles/mol).

Boiling Points and Molecular Structure

Factors Affecting Boiling Point

Boiling point increases with stronger intermolecular forces and larger molecular mass.

• Order of boiling points: CH3CH3 < CH4 < CO2 < CH2O < CH3OH (hydrogen bonding gives CH3OH the highest boiling point).

Mole Calculations and Avogadro's Number

Converting Moles to Molecules

To find the number of molecules in a given amount of moles, multiply by Avogadro's number:

• $\text{Number of molecules} = \text{moles} \times 6.022 \times 10^{23}$

• Example: 0.0180 mol CO2 × $6.022 \times 10^{23}$ = $1.08 \times 10^{22}$ molecules

Naming Ionic Compounds

Formulas and Names

Ionic compounds are named by combining the cation (metal) and anion (nonmetal or polyatomic ion).

• Example: Calcium hydroxide is Ca(OH)2

Mole Calculations and Avogadro's Number

Drawing Lewis Structures

Lewis structures show the arrangement of valence electrons around atoms in a molecule.

• Count total valence electrons.

• Arrange atoms and distribute electrons to satisfy the octet rule.

Bond Angles and Geometry

The geometry of a molecule is determined by the number of electron groups around the central atom.

Electron Group Geometry	Molecular Geometry	Bond Angle
Tetrahedral	Tetrahedral	109.5°
Tetrahedral	Trigonal pyramidal	~107°
Tetrahedral	Bent	~104.5°
Trigonal planar	Bent	~120°

Bonding and Electrons

Total Bonding Electrons

Count the number of electrons involved in bonds in a molecule.

• Example: H–N=N–H has 8 total bonding electrons (4 bonds × 2 electrons each).

Trends in the Periodic Table

Formation of Cations

Metals tend to lose electrons to form cations due to their low ionization energies and larger atomic radii.

• Key point: Metals have lower ionization energies, making it easier to lose electrons.

Triple Bonds in Lewis Structures

Identifying Triple Bonds

Triple bonds involve three shared pairs of electrons between two atoms, commonly found in molecules like N2 and C2H2.

• Example: CO does not have a triple bond in its Lewis structure.

Unit Conversions

Metric and English Units

Common conversion factors:

• 1 inch = 2.54 cm

• 1 ft = 12 inches

• 1 L = 1000 mL

• 1 m3 = 1000 L

Significant Figures

Rules for Significant Figures

Significant figures reflect the precision of a measurement. When performing calculations, the result should have the same number of significant figures as the least precise measurement.

• Example: $\frac{12.315 \times 6.085}{29.0 + 0.9 \times 7.243} = 0.085$ (rounded to 2 significant figures)

Classification of Chemical Reactions

Types of Reactions

• Combination (Synthesis): Two or more substances form one product.

• Decomposition: One substance breaks down into two or more products.

• Single Replacement: One element replaces another in a compound.

• Double Replacement: Exchange of ions between two compounds.

• Combustion: Substance reacts with O2, producing heat and light.

Redox Reactions

Oxidation and Reduction

Oxidation is the loss of electrons; reduction is the gain of electrons.

• Example: In Ni(s) → Ni2+(aq) + 2e-, Ni is oxidized.

Free Response Practice

Lewis Structure and Polarity of SCl2

• Draw all valence electrons and lone pairs.

• Determine molecular geometry (bent) and polarity (polar due to lone pairs on S).

Stoichiometry with Balanced Equations

• Calculate molar mass using atomic masses.

• Use mole ratios from balanced equations to determine required amounts of reactants or products.

• Convert moles to molecules using Avogadro's number.

Additional info:

• These notes cover fundamental concepts in General, Organic, and Biological (GOB) Chemistry, including atomic structure, bonding, stoichiometry, molecular geometry, and chemical reactions.

• Practice problems are representative of typical exam questions in a college-level GOB Chemistry course.