Chemistry study guide
This guide is designed to help you understand the core ideas that commonly appear in introductory college Chemistry courses. Read the explanation, follow the example, then try the practice questions without looking back. The goal is to understand the reasoning—not just memorize an answer.
Course requirements vary by school and instructor. Use this as a learning and review resource alongside your assigned textbook, lectures, and instructor guidance.
Atomic structure
Concept: Atoms contain protons and neutrons in the nucleus and electrons around it. Atomic number equals proton count and identifies the element. Isotopes have the same protons but different neutrons. Ions form when electrons are gained or lost.
Worked example: Carbon has atomic number 6, so every carbon atom has 6 protons. Carbon-14 has 8 neutrons because 14−6=8. A C⁴⁻ ion would have gained electrons relative to neutral carbon.
Remember: Protons determine element identity; mass number=protons+neutrons; charge depends on protons versus electrons.
Common mistakes: Using mass number as proton count; changing proton number and still calling it the same element; confusing isotope with ion.
Practice it
- How many neutrons are in Na-23 if Z=11?
- What changes when an atom becomes an ion?
- Which particle determines element identity?
Periodic trends
Concept: Periodic trends arise from electron arrangement and nuclear attraction. Across a period, effective nuclear charge generally increases, pulling electrons closer. Down a group, added energy levels increase size and shielding. These patterns help predict atomic radius, ionization energy, and electronegativity.
Worked example: Across Period 2, atomic radius generally decreases from left to right, while ionization energy generally rises. Down Group 1, atoms generally get larger.
Remember: Trends have exceptions; compare positions on the table; shielding increases down groups.
Common mistakes: Memorizing arrows without understanding attraction; treating trends as absolute with no exceptions; confusing ionization energy with electronegativity.
Practice it
- Which is generally larger: Li or Na?
- Which generally has higher ionization energy: C or F?
- Why does atomic radius grow down a group?
Chemical bonding
Concept: Chemical bonds involve valence electrons. Ionic bonding is electrostatic attraction between oppositely charged ions, often after electron transfer. Covalent bonding shares electron pairs. Molecular shape and bond polarity influence overall molecular polarity and properties.
Worked example: NaCl forms from Na⁺ and Cl⁻ ions. In H₂O, O–H bonds are polar covalent and the bent shape gives the molecule a net dipole.
Remember: Use valence electrons; distinguish bond polarity from molecular polarity; Lewis structures help track bonding and lone pairs.
Common mistakes: Calling every metal-nonmetal compound a molecule; ignoring molecular shape; confusing intermolecular forces with bonds inside molecules.
Practice it
- Classify bonding in MgO.
- Why is water polar?
- How many valence electrons does oxygen have?
Balancing reactions
Concept: A balanced chemical equation has the same number of each type of atom on both sides, reflecting conservation of matter. Balance by changing coefficients in front of formulas, never by changing subscripts inside correct chemical formulas.
Worked example: H₂ + O₂ → H₂O becomes 2H₂ + O₂ → 2H₂O. Both sides then have 4 H atoms and 2 O atoms.
Remember: Count atoms on both sides; adjust coefficients; reduce coefficients to smallest whole-number ratio.
Common mistakes: Changing subscripts; forgetting atoms inside parentheses; stopping before checking every element.
Practice it
- Balance N₂ + H₂ → NH₃.
- Why must subscripts not be changed?
- Check whether 2Na+Cl₂→2NaCl is balanced.
Moles & stoichiometry
Concept: The mole connects microscopic particles to measurable amounts. One mole contains Avogadro’s number of particles, and molar mass converts grams to moles. Stoichiometry uses balanced-equation coefficients as mole ratios to relate reactants and products.
Worked example: For 2H₂ + O₂ → 2H₂O, 2 mol H₂ react with 1 mol O₂. If 4 mol H₂ react with enough O₂, the ratio predicts 4 mol H₂O.
Remember: Balance first; convert given quantity to moles; use mole ratio; convert to requested units.
Common mistakes: Using unbalanced coefficients; using gram ratios directly from coefficients; skipping units.
Practice it
- How many mol O₂ react with 6 mol H₂?
- What is the first step in stoichiometry?
- Why is molar mass needed for gram-to-mole conversion?
Acids, bases & pH
Concept: Acids donate H⁺ in Brønsted–Lowry chemistry and bases accept H⁺. pH is a logarithmic measure related to hydrogen-ion concentration; lower pH means greater [H⁺]. A change of one pH unit corresponds to a tenfold change in [H⁺].
Worked example: A solution at pH 3 has ten times the hydrogen-ion concentration of a solution at pH 4 and 100 times that of pH 5.
Remember: pH below 7 is acidic and above 7 basic at room-temperature conventions; strong and concentrated are different ideas; neutralization involves acid-base reaction.
Common mistakes: Treating pH as linear; assuming every strong acid is highly concentrated; confusing pH with pOH.
Practice it
- Which is more acidic: pH 2 or pH 5?
- How many times greater is [H⁺] at pH 3 than pH 6?
- What is the difference between strength and concentration?