Chemistry 2e PDF Download – OpenStax

Students taking General Chemistry can download the complete textbook “Chemistry 2e” by Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson, and contributors, free as a PDF from OpenStax. It is a traditional, example-and-exercise-driven chemistry text — every section works through several fully-solved examples before a large practice-problem set, and every chapter closes with a Key Terms, Key Equations, and Key Concepts review before its Chapter Review Exercises.

This book covers the full two-semester General Chemistry sequence in 21 chapters, from the essential ideas of measurement and atomic structure through stoichiometry, gases, and bonding in the first semester, to equilibrium, thermodynamics, electrochemistry, and descriptive chemistry of the elements in the second. It is designed to stand alone as a complete two-semester course for BS Chemistry, Pre-Medical, Engineering, and other STEM majors, with no prior college chemistry required.

Book Overview

Chapter List

Chapter 1: Essential Ideas

Difficulty: Easy · General Chemistry I · Key topics: chemistry in context, phases and classification of matter, physical and chemical properties, measurement, measurement uncertainty and precision, mathematical treatment of measurements

This opening chapter sets the vocabulary and habits of mind the entire book depends on. It defines chemistry itself and the atomic/molecular perspective, then classifies matter by phase (solid, liquid, gas) and by composition (element, compound, mixture, pure substance). It distinguishes physical properties/changes (no new substance formed) from chemical ones (a new substance forms), before turning practical: SI units, unit conversion, and the difference between accuracy and precision. It closes with the mathematical treatment every later chapter assumes without re-explaining — significant figures and dimensional analysis (the factor-label method) for tracking units through a calculation.

Key Points:

• Matter is classified by phase (solid, liquid, gas) and by composition — pure substances (elements, compounds) versus mixtures (homogeneous or heterogeneous)
• A physical change alters a substance’s form without changing its chemical identity (melting ice); a chemical change produces a genuinely new substance (burning wood)
• Accuracy is how close a measurement is to the true value; precision is how reproducible repeated measurements are — a data set can be precise without being accurate
• Significant figures communicate a measurement’s precision, and a calculated answer can never be more precise than its least-precise input measurement
• Dimensional analysis (the factor-label method) converts between units by multiplying by conversion factors arranged so unwanted units cancel — the single most-used calculation technique in the entire book

Memory Tip: Before every calculation in this book, write down the units next to every number and cancel them algebraically (dimensional analysis) — if the units left over at the end don’t match what the question asked for, the setup is wrong, regardless of the arithmetic.

Common Mistake: Confusing accuracy with precision, or assuming a precise result is automatically correct. A poorly calibrated instrument can give very precise (tightly clustered) readings that are all equally far from the true value — precision alone says nothing about whether a measurement is actually correct.

Important Questions:

• What is the difference between a physical change and a chemical change? Give an example of each. A physical change alters a substance’s form or appearance without changing its chemical identity — for example, ice melting into liquid water is still H2O throughout. A chemical change produces one or more new substances with different properties — for example, burning wood produces ash, smoke, and gases that are chemically distinct from the original wood.
• Explain the difference between accuracy and precision, and why a measurement can be precise without being accurate. Accuracy measures how close a measurement is to the true, correct value, while precision measures how close repeated measurements are to each other (their reproducibility). A measurement can be precise but inaccurate if, for example, a scale is miscalibrated: every reading might cluster tightly together (high precision) while consistently being off from the true weight by the same offset (low accuracy).

Chapter 2: Atoms, Molecules, and Ions

Difficulty: Medium · General Chemistry I · Key topics: early atomic theory, evolution of atomic theory, atomic structure and symbolism, chemical formulas, the periodic table, ionic and molecular compounds, chemical nomenclature

This chapter builds matter from the atom up. It traces atomic theory from Dalton’s original postulates through Thomson, Rutherford, and the modern nuclear model, then defines atomic structure — protons, neutrons, electrons, atomic number, mass number, and isotopes — along with the standard symbolism (Z, A) used to write them. From there it introduces chemical formulas as shorthand for a compound’s composition, the periodic table as elements’ organizing chart, and the fundamental split between ionic compounds (formed by electron transfer between metals and nonmetals) and molecular/covalent compounds (formed by electron sharing). It closes with the systematic nomenclature rules for naming both classes correctly.

Key Points:

• An atom’s identity is set by its atomic number (Z, the number of protons); isotopes of the same element share Z but differ in neutron count, and therefore in mass number (A)
• Ionic compounds form between metals and nonmetals via electron transfer, producing oppositely charged ions held together by electrostatic attraction; molecular (covalent) compounds form between nonmetals via electron sharing
• The periodic table organizes elements by increasing atomic number into periods (rows) and groups (columns) that share recurring chemical properties
• A chemical formula’s subscripts give the exact ratio of atoms (or ions) in one formula unit of a compound — molecular formulas for covalent compounds, formula units for ionic compounds
• Systematic nomenclature differs for ionic versus molecular compounds: ionic names use the cation name plus an anion name ending in -ide (or a polyatomic ion’s own name); molecular names use Greek numerical prefixes (mono-, di-, tri-) to specify exact atom counts

Practice Tip: Before naming or writing the formula of any compound, first decide whether it’s ionic (metal + nonmetal) or molecular (nonmetal + nonmetal) — the two families use completely different naming rules, and picking the wrong rule set is the single most common nomenclature error.

Common Mistake: Using Greek numerical prefixes (di-, tri-, tetra-) when naming an ionic compound, or omitting them when naming a molecular compound. Prefixes belong only to molecular/covalent compound names (carbon dioxide, not “carbon monoxide-two”); ionic compound names never use them because the formula is already fixed by the ions’ charges (sodium chloride, never “monosodium monochloride”).

Important Questions:

• What determines whether two elements form an ionic compound or a molecular (covalent) compound? The combination of a metal and a nonmetal typically forms an ionic compound, where the metal atom loses one or more electrons (becoming a positively charged cation) and the nonmetal atom gains them (becoming a negatively charged anion); the resulting oppositely charged ions attract electrostatically. Two nonmetals typically form a molecular compound instead, sharing electrons in covalent bonds rather than transferring them outright.
• What is an isotope, and how do isotopes of the same element differ from one another? Isotopes are atoms of the same element (same atomic number, Z, meaning the same number of protons) that have different numbers of neutrons, and therefore different mass numbers (A). Because chemical behavior is governed almost entirely by the number and arrangement of electrons (which matches the proton count), isotopes of an element behave nearly identically chemically despite their differing masses.

Chapter 3: Composition of Substances and Solutions

Difficulty: Medium · General Chemistry I · Key topics: formula mass and the mole concept, empirical and molecular formulas, molarity, other units for solution concentration

This chapter introduces the mole — the central counting unit of all quantitative chemistry — and everything built on it. It defines formula mass (the sum of atomic masses in a formula) and Avogadro’s number as the bridge between a lab-scale mass in grams and a countable number of individual atoms or molecules. From there it covers determining a compound’s empirical formula (simplest whole-number atom ratio) from percent composition data, then scaling up to the true molecular formula using molar mass. The chapter’s second half turns to solutions, introducing molarity (moles of solute per liter of solution) as the standard concentration unit, plus alternative units (mass percentage, ppm, molality) used in specific contexts.

Key Points:

• One mole of any substance contains exactly Avogadro’s number (6.022 × 10^23) of its constituent particles, and a substance’s molar mass (in g/mol) is numerically equal to its formula mass
• The mole is the essential bridge between the macroscopic quantities chemists actually measure (grams, milliliters) and the microscopic particle counts chemical equations are written in terms of
• An empirical formula gives the simplest whole-number ratio of atoms in a compound; the molecular formula is always a whole-number multiple of the empirical formula, found by comparing the compound’s actual molar mass to the empirical formula’s mass
• Molarity (M) is defined as moles of solute per liter of solution — the standard, most widely used concentration unit throughout the rest of the book
• Percent composition (by mass), parts per million, and molality are alternative concentration/composition units used when molarity isn’t the most convenient choice, such as for very dilute solutions or temperature-sensitive applications

Memory Tip: Think of the mole as chemistry’s version of “a dozen” — just a fixed, very large counting number (6.022 × 10^23) that lets you convert between a countable number of particles and a mass you can actually weigh on a balance.

Common Mistake: Confusing empirical formula with molecular formula, or assuming they’re always identical. The empirical formula is only the simplest whole-number ratio — glucose’s molecular formula (C6H12O6) and its empirical formula (CH2O) are different, and you must know the compound’s actual molar mass to determine the correct multiplier between them.

Important Questions:

• What is Avogadro’s number, and why is it essential for connecting mass measurements to particle counts? Avogadro’s number, 6.022 × 10^23, is the number of particles (atoms, molecules, or ions) in exactly one mole of a substance. It’s essential because chemists measure mass in the lab (grams), but chemical reactions occur atom-by-atom and molecule-by-molecule; the mole, defined using Avogadro’s number, is the conversion bridge that lets a measured mass be translated into a countable number of particles and vice versa.
• How do you determine a compound’s molecular formula once you know its empirical formula? Divide the compound’s actual (experimentally determined) molar mass by the empirical formula’s mass to find a whole-number multiplier n. Then multiply every subscript in the empirical formula by n to get the molecular formula — for example, if the empirical formula CH2O has a mass of 30 g/mol but the actual molar mass is 180 g/mol, n = 6, giving the molecular formula C6H12O6.

Chapter 4: Stoichiometry of Chemical Reactions

Difficulty: Medium · General Chemistry I · Key topics: writing and balancing chemical equations, classifying chemical reactions, reaction stoichiometry, reaction yields, quantitative chemical analysis

This chapter turns a balanced chemical equation into a quantitative tool. It starts with writing and balancing equations (conserving atoms of every element on both sides), then surveys the common reaction classes — precipitation, acid-base, and oxidation-reduction — as a way to predict what products a given reaction will form. The chapter’s core is reaction stoichiometry: using a balanced equation’s mole ratios, combined with the mole concept from Chapter 3, to calculate exactly how much product forms from a given amount of reactant, or how much reactant is needed for a target amount of product. It introduces the limiting reactant (the reactant that runs out first, capping how much product can form) and distinguishes theoretical, actual, and percent yield, closing with quantitative analytical techniques such as gravimetric and titrimetric analysis.

Key Points:

• A balanced chemical equation must have the identical number of atoms of every element on both the reactant and product sides — balancing coefficients never change subscripts, only the number of whole molecules/formula units
• The coefficients in a balanced equation give the exact mole ratio between any two reactants or products, which is the conversion factor for every stoichiometry calculation
• The limiting reactant is whichever reactant runs out first, and it alone determines the maximum theoretical amount of product a reaction can form — any excess of the other reactant(s) is simply left over unreacted
• Theoretical yield is the maximum amount of product stoichiometry predicts; actual yield is what’s really obtained in the lab (always less, due to side reactions, incomplete reactions, or losses); percent yield = actual/theoretical × 100%
• Gravimetric analysis (measuring mass) and titrimetric analysis (measuring volume of a reacting solution of known concentration) are both quantitative techniques that use stoichiometric relationships to determine an unknown sample’s composition

Practice Tip: Every stoichiometry problem follows the same three-step conversion chain: grams of A → moles of A → (using the balanced equation’s mole ratio) moles of B → grams of B. Identify which of the three steps a given problem is asking for before trying to solve it.

Common Mistake: Using the given masses of two reactants directly in a ratio, instead of first converting both to moles. Reaction stoichiometry is always based on MOLE ratios (from the balanced equation’s coefficients), not mass ratios — skipping the gram-to-mole conversion step is the most common stoichiometry error, especially when identifying the limiting reactant.

Important Questions:

• What is a limiting reactant, and how does it determine the amount of product formed in a reaction? The limiting reactant is the reactant that is completely consumed first in a chemical reaction, while some amount of the other reactant(s) remains unreacted (in excess). Because a reaction stops once any one reactant runs out, the limiting reactant alone determines the maximum theoretical amount of product that can form — calculating from the excess reactant instead would overestimate the actual yield.
• What is the difference between theoretical yield, actual yield, and percent yield? Theoretical yield is the maximum amount of product a reaction could produce, calculated purely from stoichiometry assuming the limiting reactant reacts completely with no losses. Actual yield is the amount of product genuinely obtained when the reaction is carried out in the lab, which is virtually always less than the theoretical yield due to side reactions, incomplete reactions, or material lost during purification. Percent yield expresses actual yield as a percentage of theoretical yield (actual/theoretical × 100%), quantifying how efficiently the reaction actually performed.

Chapter 5: Thermochemistry

Difficulty: Medium · General Chemistry I · Key topics: energy basics, calorimetry, enthalpy

This chapter connects chemical reactions to energy, since roughly 85% of energy used in modern society comes from combustion reactions. It starts with energy basics — kinetic versus potential energy, heat versus work, and the units (joules, calories) used to measure them — then covers calorimetry, the experimental technique for measuring the heat released or absorbed by a reaction using a calorimeter’s temperature change. The chapter’s second half introduces enthalpy (H), a state function representing heat flow at constant pressure, and Hess’s Law, which lets you calculate a reaction’s overall enthalpy change by adding up the enthalpies of a series of steps that sum to the same overall reaction, regardless of the actual pathway taken.

Key Points:

• Energy exists as kinetic energy (motion) or potential energy (stored, position-dependent); heat is energy transferred due to a temperature difference, while work is energy transferred by a force acting through a distance
• Calorimetry measures a reaction’s heat flow indirectly, by measuring the resulting temperature change of a known mass of surrounding material (often water) with a known specific heat capacity
• Enthalpy (H) is a state function representing heat flow at constant pressure — ΔH is negative for an exothermic reaction (releases heat) and positive for an endothermic reaction (absorbs heat)
• Because enthalpy is a state function, ΔH for a reaction depends only on the initial and final states, not on the path taken — this is the basis for Hess’s Law
• Hess’s Law lets you calculate an unmeasurable or hard-to-measure reaction’s ΔH by algebraically combining the known ΔH values of a series of other reactions that sum to the same overall equation

Memory Tip: “Exo” means exit — an exothermic reaction’s heat exits the system into the surroundings, so ΔH is negative. “Endo” means enter — an endothermic reaction absorbs heat entering from the surroundings, so ΔH is positive.

Common Mistake: Forgetting to reverse the sign of ΔH when a Hess’s Law step’s reaction must be flipped (reversed) to sum correctly to the target overall equation. If a step is run in reverse, its enthalpy value must also be multiplied by −1, and if a step’s coefficients are scaled up or down, ΔH must be scaled by the identical factor.

Important Questions:

• What does it mean for enthalpy to be a “state function,” and why does that matter for Hess’s Law? A state function’s value depends only on a system’s current state (its initial and final conditions), not on the specific path or process taken to get there. This matters for Hess’s Law because it means a reaction’s overall ΔH can be found by adding up the ΔH values of any convenient series of steps that sum to the same net reaction — regardless of whether that particular multi-step pathway is how the reaction actually happens in practice.
• How does a calorimeter measure the heat released or absorbed by a chemical reaction? A calorimeter contains a known mass of a surrounding material (commonly water) with a known specific heat capacity. As the reaction proceeds inside the calorimeter, the heat it releases or absorbs changes the surrounding material’s temperature; measuring that temperature change, combined with the known mass and specific heat, lets you calculate the amount of heat transferred using q = mcΔT.

Chapter 6: Electronic Structure and Periodic Properties

Difficulty: Hard · General Chemistry I · Key topics: electromagnetic energy, the Bohr model, development of quantum theory, electronic structure of atoms (electron configurations), periodic variations in element properties

This chapter explains why the periodic table’s shape isn’t arbitrary — it’s a direct consequence of how electrons are arranged around the nucleus. It starts with electromagnetic radiation’s wave properties (wavelength, frequency) and its connection to atomic emission spectra, then traces the Bohr model’s quantized electron energy levels before quantum theory replaced it with the modern picture: orbitals described by a wavefunction, characterized by a set of quantum numbers rather than a fixed orbit. The chapter builds up electron configurations using the Aufbau principle, Pauli exclusion principle, and Hund’s rule, then explains how those configurations directly produce the periodic table’s structure and its recurring trends — atomic radius, ionization energy, and electron affinity — across periods and down groups.

Key Points:

• The Bohr model correctly explained hydrogen’s line spectrum by proposing quantized electron energy levels, but modern quantum mechanics replaced its fixed circular orbits with orbitals — probability regions described by a wavefunction and a set of quantum numbers
• Electron configurations are built using three rules together: the Aufbau principle (fill lowest-energy orbitals first), the Pauli exclusion principle (no two electrons in an atom share all four quantum numbers), and Hund’s rule (electrons fill degenerate orbitals singly before pairing up)
• The periodic table’s structure directly reflects electron configuration — elements in the same group share the same valence electron configuration, which is why they share similar chemical properties
• Atomic radius generally decreases across a period (increasing nuclear charge pulls electrons in tighter) and increases down a group (each new period adds an entire electron shell)
• Ionization energy (energy to remove an electron) generally increases across a period and decreases down a group — the opposite trend direction from atomic radius, since a smaller, more tightly held atom is harder to ionize

Practice Tip: To write an electron configuration quickly, follow the periodic table’s own shape as a map — each row corresponds to filling a new principal energy level, and each block (s, p, d, f) corresponds to which subshell type is being filled, so the table itself is a built-in Aufbau-principle diagram.

Common Mistake: Assuming atomic radius and ionization energy trends both increase (or both decrease) in the same direction across the periodic table. They move in OPPOSITE directions: atomic radius decreases left-to-right across a period while ionization energy increases, because a smaller atom holds its electrons more tightly and therefore requires more energy to remove one.

Important Questions:

• Why do elements in the same group (column) of the periodic table share similar chemical properties? Elements in the same group have the identical valence electron configuration (the same number and arrangement of electrons in their outermost shell), even though their total number of electrons differs. Since chemical behavior — how readily an atom gains, loses, or shares electrons — is determined primarily by the valence electron configuration, elements sharing that configuration behave similarly in reactions.
• How does atomic radius change across a period (left to right) and down a group, and why? Atomic radius generally decreases moving left to right across a period, because each added proton increases the nuclear charge pulling all the electrons (which are still in the same principal shell) in more tightly. Atomic radius generally increases moving down a group, because each successive period adds an entirely new, larger electron shell farther from the nucleus, which outweighs the increased nuclear charge.

Chapter 7: Chemical Bonding and Molecular Geometry

Difficulty: Medium · General Chemistry I · Key topics: ionic bonding, covalent bonding, Lewis symbols and structures, formal charges and resonance, strengths of ionic and covalent bonds, molecular structure and polarity (VSEPR)

This chapter explains how atoms connect into the compounds surveyed conceptually back in Chapter 2, and predicts the resulting three-dimensional shapes. It contrasts ionic bonding (electron transfer, electrostatic attraction between ions) with covalent bonding (electron sharing between atoms), then introduces Lewis symbols and structures as the standard way to represent bonding electrons and lone pairs on paper. It covers formal charge as a tool for choosing the best Lewis structure among several valid options, and resonance for molecules whose bonding can’t be captured by any single Lewis structure alone. The chapter’s second half applies VSEPR theory (valence shell electron pair repulsion) to predict a molecule’s actual three-dimensional geometry from its Lewis structure, and uses that geometry plus bond polarity to determine whether the overall molecule is polar or nonpolar.

Key Points:

• Ionic bonds form through complete electron transfer, creating oppositely charged ions held together by electrostatic attraction; covalent bonds form through electron sharing between two atoms, which can be nonpolar (equal sharing) or polar (unequal sharing) depending on electronegativity difference
• A Lewis structure shows all valence electrons in a molecule as either bonding pairs (shared between two atoms) or lone pairs (unshared, belonging to one atom)
• Formal charge helps choose the most reasonable Lewis structure among several valid possibilities — the best structure generally has formal charges as close to zero as possible, with any negative formal charge on the more electronegative atom
• Resonance structures are used when a single Lewis structure can’t accurately represent a molecule’s actual bonding (such as ozone or the carbonate ion) — the real molecule is a blend (resonance hybrid) of all valid resonance structures, not any one of them alone
• VSEPR theory predicts a molecule’s three-dimensional shape by minimizing repulsion between electron groups (bonding pairs and lone pairs) around a central atom — lone pairs take up more space than bonding pairs and compress bond angles
• A molecule’s overall polarity depends on BOTH its individual bond polarities AND its molecular geometry — a molecule with polar bonds can still be nonpolar overall if its geometry causes the individual bond dipoles to cancel symmetrically (like CO2)

Memory Tip: Lone pairs push harder than bonding pairs in VSEPR theory — a lone pair is held by only one nucleus and spreads out more, so it repels neighboring electron groups more strongly than a bonding pair (shared between two nuclei), compressing bond angles below the ideal geometric value whenever lone pairs are present.

Common Mistake: Assuming a molecule with polar bonds must always be a polar molecule overall. Molecular polarity depends on geometry as well as bond polarity — carbon dioxide (CO2) has two polar C=O bonds, but its linear, symmetric geometry causes the two bond dipoles to point in exactly opposite directions and cancel, making the overall molecule nonpolar.

Important Questions:

• What is the difference between an ionic bond and a covalent bond? An ionic bond forms when one atom (typically a metal) transfers one or more electrons completely to another atom (typically a nonmetal), creating two oppositely charged ions that are then held together by electrostatic attraction. A covalent bond forms instead when two atoms (typically both nonmetals) share a pair of electrons between them, rather than one atom giving them up entirely to the other.
• How does VSEPR theory predict a molecule’s three-dimensional shape, and why do lone pairs affect bond angles? VSEPR theory predicts molecular geometry by assuming that electron groups (both bonding pairs and lone pairs) around a central atom arrange themselves to minimize mutual repulsion, spreading out as far apart as possible. Lone pairs occupy more space and repel neighboring electron groups more strongly than bonding pairs do (since a lone pair is held by only one nucleus rather than shared between two), which compresses the bond angles between the actual bonded atoms below the ideal geometric angle.

Chapter 8: Advanced Theories of Covalent Bonding

Difficulty: Hard · General Chemistry I · Key topics: valence bond theory, hybrid atomic orbitals, multiple bonds, molecular orbital theory

Lewis structures and VSEPR theory (Chapter 7) predict shape well but can’t explain everything — famously, why liquid oxygen is attracted to a magnet while liquid nitrogen is not, despite both having simple-looking Lewis structures. This chapter introduces two more advanced bonding models to close that gap. Valence bond theory explains covalent bonds as overlapping atomic orbitals, and hybrid orbital theory (sp, sp2, sp3, and beyond) reconciles the simple s and p atomic orbital shapes with the actual observed bond angles in real molecules. The chapter then covers how multiple bonds (double, triple) form from combinations of sigma and pi bonds, before introducing molecular orbital theory — a fundamentally different approach that combines atomic orbitals into new molecular orbitals spanning the whole molecule, correctly predicting oxygen’s magnetic behavior where valence bond theory cannot.

Key Points:

• Valence bond theory describes a covalent bond as the overlap of two atomic orbitals, each contributing one electron to the shared bonding pair
• Orbital hybridization (mixing atomic s and p orbitals into new hybrid orbitals like sp3, sp2, or sp) explains observed bond angles that pure, unmixed atomic orbitals could not produce on their own
• A single bond consists of one sigma bond (direct, head-on orbital overlap); a double bond adds one pi bond (sideways overlap of unhybridized p orbitals); a triple bond has one sigma and two pi bonds
• Molecular orbital theory combines atomic orbitals from ALL atoms in a molecule into new molecular orbitals (bonding and antibonding) that span the entire molecule, rather than treating each bond as an isolated two-atom overlap
• Molecular orbital theory correctly predicts that O2 is paramagnetic (attracted to a magnetic field, due to unpaired electrons in its molecular orbitals) even though its simple Lewis structure suggests all electrons are paired — a prediction valence bond theory and Lewis structures alone cannot make

Practice Tip: Count sigma and pi bonds separately when analyzing a multiple bond: every bond (single, double, or triple) contains exactly one sigma bond, and any ADDITIONAL bonds beyond the first are always pi bonds — a triple bond is one sigma plus two pi bonds, never three equivalent bonds.

Common Mistake: Assuming Lewis structures and valence bond theory can explain every observable molecular property. They correctly predict bonding patterns and shapes for most simple molecules, but they fail to predict magnetic properties like O2’s paramagnetism — only molecular orbital theory, which considers electrons delocalized across the whole molecule, gets this right.

Important Questions:

• What is orbital hybridization, and why is it needed to explain real molecular bond angles? Orbital hybridization is the mixing of an atom’s atomic orbitals (such as one s and three p orbitals) into a new set of equivalent hybrid orbitals (such as four sp3 orbitals) with different shapes and orientations than the original unmixed orbitals. It’s needed because the bond angles predicted by simply overlapping unhybridized s and p orbitals don’t match the bond angles actually observed in real molecules (like methane’s 109.5° tetrahedral angles); hybrid orbitals, oriented to minimize repulsion, correctly reproduce these observed geometries.
• Why does molecular orbital theory correctly predict that O2 is paramagnetic, while a simple Lewis structure does not? A Lewis structure for O2 shows a double bond with all electrons apparently paired, suggesting no magnetic attraction. Molecular orbital theory instead combines the atomic orbitals of both oxygen atoms into molecular orbitals spanning the whole molecule, and filling these molecular orbitals according to their actual relative energies places two electrons individually into separate, degenerate antibonding pi orbitals rather than pairing them — those two unpaired electrons are what make O2 paramagnetic, correctly explaining its attraction to a magnetic field.

Chapter 9: Gases

Difficulty: Medium · General Chemistry I · Key topics: gas pressure, the ideal gas law, stoichiometry of gaseous substances, effusion and diffusion, kinetic-molecular theory, non-ideal gas behavior

This chapter is entirely about how gases behave, from the everyday (why a hot air balloon rises) to the rigorously mathematical. It starts with gas pressure and its measurement, then builds up the ideal gas law (PV = nRT) from the individual gas laws (Boyle’s, Charles’s, Avogadro’s) that each hold one variable constant. The ideal gas law becomes a stoichiometry tool in its own right, letting you relate a gas’s volume directly to moles in a reaction. The chapter covers effusion and diffusion (how gases move through small openings or spread through space) via Graham’s Law, then explains WHY gases behave this way at all through kinetic-molecular theory — treating a gas as vast numbers of particles in constant random motion — and closes with where the ideal gas law breaks down for real gases at high pressure or low temperature.

Key Points:

• The ideal gas law, PV = nRT, combines Boyle’s Law (P and V inversely related), Charles’s Law (V and T directly related), and Avogadro’s Law (V and n directly related) into a single equation relating all four gas properties at once
• STP (standard temperature and pressure) provides a fixed reference condition where one mole of any ideal gas occupies the same molar volume (22.4 L at the traditional STP definition), useful for gas stoichiometry calculations
• Graham’s Law states that a gas’s rate of effusion (escape through a tiny opening) is inversely proportional to the square root of its molar mass — lighter gas molecules escape faster than heavier ones
• Kinetic-molecular theory explains gas laws from first principles: gas particles are in constant, random motion, their average kinetic energy is directly proportional to absolute temperature, and pressure results from particle collisions with container walls
• Real gases deviate from ideal behavior at high pressure (molecules are forced close enough together that their own volume becomes significant) and at low temperature (intermolecular attractive forces become significant relative to kinetic energy) — conditions the ideal gas law’s assumptions ignore

Practice Tip: Always convert temperature to kelvin before using PV = nRT or any related gas law — using Celsius directly is one of the most common gas-law errors, since the equation depends on temperature’s true proportional relationship to kinetic energy, which only holds on the absolute (kelvin) scale.

Common Mistake: Applying the ideal gas law without checking whether the conditions (very high pressure or very low temperature) are realistic for ideal behavior. Under those conditions, real gas molecules’ own volume and their intermolecular attractions become significant, and the ideal gas law’s predictions can differ substantially from the gas’s actual measured behavior.

Important Questions:

• State the ideal gas law and explain what each variable represents. The ideal gas law is PV = nRT, where P is pressure, V is volume, n is the number of moles of gas, R is the universal gas constant, and T is absolute temperature (in kelvin). It combines the separate relationships of Boyle’s Law, Charles’s Law, and Avogadro’s Law into one equation that relates all four gas properties simultaneously for a sample behaving ideally.
• Under what conditions does a real gas deviate most from ideal gas behavior, and why? Real gases deviate most from ideal behavior at high pressure and low temperature. At high pressure, gas molecules are forced close enough together that their own physical volume (which the ideal gas law assumes is negligible) becomes a significant fraction of the container’s volume. At low temperature, gas molecules move more slowly, giving intermolecular attractive forces (which the ideal gas law assumes are zero) enough relative influence to noticeably affect the gas’s pressure and volume.

Chapter 10: Liquids and Solids

Difficulty: Medium · General Chemistry I · Key topics: intermolecular forces, properties of liquids, phase transitions, phase diagrams, the solid state of matter, lattice structures in crystalline solids

This chapter explains why matter exists in different physical states at all, and why liquids and solids behave so differently from gases. It starts with intermolecular forces — dispersion forces, dipole-dipole forces, and hydrogen bonding — as the attractions responsible for holding condensed phases together, then connects those forces to observable liquid properties like viscosity, surface tension, and vapor pressure. It covers phase transitions (melting, boiling, sublimation) and the energy involved in each, then introduces phase diagrams as a map of which phase is stable at a given temperature and pressure. The chapter’s second half turns to solids, distinguishing crystalline solids (long-range ordered lattice structure) from amorphous ones, and covers the geometry of crystal lattices and unit cells that describe how atoms, ions, or molecules pack together in a crystal.

Key Points:

• Intermolecular forces (dispersion, dipole-dipole, hydrogen bonding), collectively weaker than intramolecular covalent/ionic bonds, are what hold molecules together in liquids and molecular solids — stronger intermolecular forces generally mean higher boiling and melting points
• Hydrogen bonding is an unusually strong type of dipole-dipole attraction that occurs specifically when hydrogen is bonded directly to a highly electronegative atom (N, O, or F), and explains water’s unusually high boiling point relative to its molar mass
• A phase diagram maps which phase (solid, liquid, gas) is thermodynamically stable at any given combination of temperature and pressure, including the triple point (all three phases coexist) and critical point (liquid and gas phases become indistinguishable)
• Crystalline solids have long-range, repeating order at the atomic/molecular level (described by a unit cell that repeats to build the whole crystal); amorphous solids (like glass) lack this long-range order
• The energy required for a phase transition (heat of fusion for melting, heat of vaporization for boiling) reflects the strength of the intermolecular forces being overcome — stronger forces require more energy to separate the particles into a less-ordered phase

Memory Tip: Rank intermolecular force strength from weakest to strongest as a fixed order: dispersion forces (present in ALL molecules) < dipole-dipole forces (polar molecules only) < hydrogen bonding (only N-H, O-H, or F-H bonds) — and a substance's actual boiling point reflects the STRONGEST force type present, not the weakest.

Common Mistake: Assuming all polar molecules exhibit hydrogen bonding. Hydrogen bonding requires a very specific structural feature — hydrogen bonded DIRECTLY to nitrogen, oxygen, or fluorine — not just any polar bond. A molecule can be quite polar (have significant dipole-dipole forces) without qualifying for hydrogen bonding at all if it lacks an N-H, O-H, or F-H bond.

Important Questions:

• What is hydrogen bonding, and why does it make water’s boiling point unusually high compared to similarly sized molecules? Hydrogen bonding is an especially strong form of dipole-dipole attraction that occurs when a hydrogen atom bonded directly to a highly electronegative atom (nitrogen, oxygen, or fluorine) is attracted to a lone pair on a nearby electronegative atom in a neighboring molecule. Water’s O-H bonds allow extensive hydrogen bonding between water molecules, requiring significantly more energy to separate them into the gas phase than would be predicted from water’s small molar mass alone if only weaker dispersion forces were present.
• What is the difference between a crystalline solid and an amorphous solid? A crystalline solid has a highly ordered, repeating three-dimensional arrangement of its atoms, ions, or molecules, described by a repeating unit cell that extends throughout the entire crystal. An amorphous solid lacks this long-range order — its particles are arranged relatively randomly, more similar to the disordered structure of a liquid that has simply become rigid, as in ordinary glass.

Chapter 11: Solutions and Colloids

Difficulty: Medium · General Chemistry II · Key topics: the dissolution process, electrolytes, solubility, colligative properties, colloids

This chapter examines solutions — homogeneous mixtures central to chemistry ranging from seawater to blood — at both the conceptual and quantitative level. It explains the dissolution process through intermolecular forces (“like dissolves like”), and classifies solutes as strong, weak, or non-electrolytes based on how completely they ionize in solution. It covers solubility and the factors affecting it (temperature, pressure for gases), then introduces colligative properties — boiling point elevation, freezing point depression, vapor pressure lowering, and osmotic pressure — properties that depend only on the NUMBER of dissolved particles, not their identity. The chapter closes with colloids, mixtures with particles larger than typical dissolved ions/molecules but too small to settle out, like milk or fog.

Key Points:

• “Like dissolves like” — polar solvents dissolve polar/ionic solutes well, and nonpolar solvents dissolve nonpolar solutes well, because the intermolecular forces between solute and solvent must be comparable in strength to the forces being broken
• Strong electrolytes ionize completely in solution (strong acids, strong bases, soluble ionic salts); weak electrolytes ionize only partially; nonelectrolytes don’t ionize at all — this classification determines a solution’s electrical conductivity
• Colligative properties (boiling point elevation, freezing point depression, vapor pressure lowering, osmotic pressure) depend ONLY on the concentration of dissolved particles, not on what the solute actually is chemically
• For colligative property calculations, an electrolyte’s van’t Hoff factor (i) must be included, since it dissociates into multiple particles per formula unit — one mole of NaCl produces roughly two moles of dissolved particles, doubling its colligative effect compared to a nonelectrolyte at the same molar concentration
• A colloid contains dispersed particles larger than individual molecules/ions but too small to settle out under gravity, distinguishing it from both a true solution (particles fully dissolved at the molecular/ionic level) and a suspension (particles large enough to eventually settle)

Practice Tip: Before any colligative-property calculation involving an ionic solute, first multiply the molar concentration by the van’t Hoff factor (the number of ions the formula unit dissociates into) — forgetting this step for electrolytes is the most common colligative-properties mistake.

Common Mistake: Treating an ionic solute’s molar concentration as directly equal to its effective particle concentration for colligative-property purposes. A 1 M solution of NaCl produces roughly 2 M of total dissolved particles (Na+ and Cl- ions combined), not 1 M, so its effect on freezing point depression or boiling point elevation is roughly double what the same molar concentration of a nonelectrolyte (like sugar) would produce.

Important Questions:

• What does “like dissolves like” mean, and how does it predict whether a solute will dissolve in a given solvent? “Like dissolves like” means solutes dissolve readily in solvents with similar polarity and comparable intermolecular forces — polar or ionic solutes dissolve well in polar solvents (like water), while nonpolar solutes dissolve well in nonpolar solvents (like oil or hexane). This happens because the solvent’s attraction to the solute’s particles must be strong enough to overcome the attractive forces holding the solute’s own particles together, which is far more likely when the two substances’ intermolecular force types are compatible.
• Why do colligative properties depend on the NUMBER of dissolved particles rather than their chemical identity? Colligative properties arise from how dissolved solute particles physically disrupt the solvent’s own intermolecular interactions and the number of solvent molecules available at a surface or interface, regardless of what the solute molecules or ions actually are. Because the disruption scales with how many separate particles are present, a solute that dissociates into multiple ions (like NaCl, forming Na+ and Cl-) has a proportionally larger colligative effect than the same molar amount of a solute that stays as one intact particle (like sugar).

Chapter 12: Kinetics

Difficulty: Hard · General Chemistry II · Key topics: chemical reaction rates, factors affecting reaction rates, rate laws, integrated rate laws, collision theory, reaction mechanisms, catalysis

Chemical kinetics answers a question thermodynamics can’t: not whether a reaction happens, but how fast, and by what step-by-step molecular pathway. This chapter defines reaction rate and the experimental factors affecting it — concentration, temperature, surface area, and the presence of a catalyst — then introduces rate laws, which express rate as a mathematical function of reactant concentrations raised to experimentally determined powers (the reaction orders). Integrated rate laws convert this into equations tracking concentration over time, letting you calculate a reaction’s half-life. The chapter’s theoretical core is collision theory (molecules must collide with sufficient energy and correct orientation to react) and activation energy, which together explain WHY temperature and catalysts affect rate so strongly, and it closes with reaction mechanisms — the actual sequence of elementary steps a reaction proceeds through, only some of which the overall rate law can reveal directly.

Key Points:

• A rate law expresses reaction rate as k[A]^m[B]^n, where the exponents (reaction orders) are determined EXPERIMENTALLY, not from the balanced equation’s coefficients — this is a frequent point of confusion since stoichiometric coefficients and reaction orders are unrelated quantities
• Collision theory requires two conditions for a successful reactive collision: sufficient kinetic energy (at least the activation energy, Ea) and correct molecular orientation — most collisions fail one or both conditions, which is why reactions don’t happen instantaneously even when every collision is theoretically “available”
• Higher temperature dramatically increases reaction rate primarily because it increases the fraction of molecules with kinetic energy exceeding the activation energy, not primarily because it increases collision frequency
• A catalyst speeds up a reaction by providing an alternative pathway with a lower activation energy, without being consumed itself and without changing the reaction’s overall thermodynamics (ΔH, equilibrium position)
• A reaction mechanism is the sequence of elementary steps a reaction actually proceeds through at the molecular level; the slowest step (the rate-determining step) controls the overall observed rate, and the overall rate law must be consistent with (though not always directly obtainable from) the proposed mechanism

Practice Tip: Never assume a rate law’s exponents match the balanced equation’s coefficients — reaction order must always be determined from experimental rate data (such as the method of initial rates), never read directly off a balanced chemical equation, except for a genuine single-step elementary reaction.

Common Mistake: Confusing reaction order (the exponents in the experimentally determined rate law) with the balanced equation’s stoichiometric coefficients. These are only the same for a single elementary step reaction; for any multi-step reaction (which is most real reactions), the rate law’s exponents must be found from experimental data and often do NOT match the overall balanced equation’s coefficients at all.

Important Questions:

• Why does increasing temperature increase a reaction’s rate so dramatically, according to collision theory? According to collision theory, a collision between reactant molecules must have both sufficient kinetic energy (at least the activation energy) and the correct orientation to actually produce products. Raising the temperature significantly increases the fraction of molecules possessing kinetic energy above the activation energy threshold (following the Maxwell-Boltzmann distribution), which dramatically increases the fraction of collisions that succeed — this effect on the ENERGY distribution, more than any modest increase in raw collision frequency, is what drives the strong temperature dependence of reaction rate.
• How does a catalyst increase a reaction’s rate without being consumed or changing the reaction’s thermodynamics? A catalyst provides an alternative reaction pathway (mechanism) with a lower activation energy than the uncatalyzed pathway, which allows a much larger fraction of collisions to have sufficient energy to react successfully. The catalyst itself participates in an intermediate step of this alternative pathway but is regenerated by the end of the overall reaction, so it isn’t consumed; because it only changes the pathway’s activation energy barrier and not the reactants’ or products’ actual energies, it has no effect on the reaction’s overall ΔH or its equilibrium position.

Chapter 13: Fundamental Equilibrium Concepts

Difficulty: Medium · General Chemistry II · Key topics: chemical equilibria, equilibrium constants, Le Chatelier’s principle, equilibrium calculations

This chapter introduces chemical equilibrium — the dynamic state where forward and reverse reaction rates become equal, so concentrations stop changing even though the reaction hasn’t stopped. It defines the equilibrium constant K, calculated from an equilibrium reaction quotient expression, and explains what K’s magnitude reveals about a reaction’s extent (a large K favors products; a small K favors reactants). The chapter covers Le Chatelier’s principle — predicting how an equilibrium shifts in response to a change in concentration, pressure/volume, or temperature — as a qualitative tool, then turns quantitative with ICE table calculations for finding equilibrium concentrations from initial conditions and K, plus using the reaction quotient Q versus K to predict which direction a non-equilibrium system will shift.

Key Points:

• Chemical equilibrium is a DYNAMIC state — the forward and reverse reactions are both still occurring continuously and at equal rates, not a static condition where the reaction has stopped
• The equilibrium constant K is calculated from the equilibrium concentrations (or partial pressures) of products divided by reactants, each raised to their stoichiometric coefficients; K’s magnitude indicates how far a reaction proceeds toward products (K >> 1 favors products, K << 1 favors reactants)
• Le Chatelier’s principle predicts that a system at equilibrium, subjected to a stress (added/removed substance, volume/pressure change, or temperature change), shifts in the direction that partially counteracts that stress
• Comparing the reaction quotient Q (calculated the same way as K, but from current, non-equilibrium concentrations) to K predicts which direction a system will shift to reach equilibrium: if Q < K, the reaction shifts forward; if Q > K, it shifts in reverse
• Only changing TEMPERATURE actually changes the value of K itself; adding/removing a substance or changing volume/pressure shifts the equilibrium POSITION (the actual concentrations present) but leaves K unchanged, since K is a true constant only at a fixed temperature

Memory Tip: “Q tells you which way, K tells you how far” — comparing the reaction quotient Q to the equilibrium constant K tells you which direction a system will shift to reach equilibrium, while K’s actual numerical value tells you how far toward products (or reactants) that equilibrium lies.

Common Mistake: Believing that adding a catalyst, or changing pressure/volume/concentration, changes the value of the equilibrium constant K. Only a change in TEMPERATURE actually changes K’s value — every other stress (catalyst, concentration change, volume/pressure change) shifts the equilibrium POSITION to a new set of concentrations, but K itself, evaluated at that same temperature, remains exactly the same.

Important Questions:

• What does it mean for a chemical reaction to be at “dynamic equilibrium”? Dynamic equilibrium means the forward reaction and the reverse reaction are both still actively occurring, but at exactly equal rates, so there is no NET change in the concentrations of reactants or products over time, even though individual molecules continue reacting in both directions. This is fundamentally different from the reaction having simply stopped — equilibrium is a balance of ongoing opposing processes, not a static endpoint.
• How do you use the reaction quotient Q, compared to the equilibrium constant K, to predict which direction a reaction will shift? Calculate Q using the identical mathematical expression as K, but using the system’s current (not necessarily equilibrium) concentrations or partial pressures. If Q is less than K, the reaction has not yet produced enough product relative to equilibrium, so it will shift forward (toward products) to increase Q until Q = K. If Q is greater than K, the reaction has too much product relative to equilibrium, so it will shift in reverse (toward reactants) until Q decreases to equal K.

Chapter 14: Acid-Base Equilibria

Difficulty: Hard · General Chemistry II · Key topics: Bronsted-Lowry acids and bases, pH and pOH, relative strengths of acids and bases, hydrolysis of salts, polyprotic acids, buffers, acid-base titrations

This chapter applies the general equilibrium framework from Chapter 13 to the specific and hugely important case of acid-base chemistry. It defines acids and bases through the Bronsted-Lowry model (proton donors and acceptors) and introduces the pH/pOH scales for quantifying a solution’s acidity. It distinguishes strong acids/bases (complete ionization, Ka or Kb effectively infinite) from weak acids/bases (partial, equilibrium-governed ionization, with a genuine Ka or Kb value), covers how dissolved salts can themselves shift a solution’s pH through hydrolysis, and extends the framework to polyprotic acids that ionize in multiple, distinct steps. The chapter’s practical core is buffers — solutions that resist pH change — and acid-base titrations, the standard experimental technique for determining an unknown acid or base’s concentration.

Key Points:

• In the Bronsted-Lowry model, an acid is a proton (H+) donor and a base is a proton acceptor; every acid-base reaction produces a conjugate acid-base pair on each side of the equation
• pH = −log[H3O+] and pOH = −log[OH-], and at 25°C, pH + pOH always equals 14 for any aqueous solution — a lower pH means a more acidic (higher H3O+ concentration) solution
• A strong acid or base ionizes essentially completely in water (Ka or Kb is very large); a weak acid or base ionizes only partially, reaching a genuine equilibrium governed by its Ka or Kb value, which must be looked up or calculated, not assumed
• A buffer solution (a weak acid plus its conjugate base, or a weak base plus its conjugate acid) resists significant pH change when small amounts of strong acid or base are added, because the buffer’s components can consume the added H3O+ or OH- through their own equilibrium
• In an acid-base titration, the equivalence point (where moles of acid exactly equal moles of base) does NOT always occur at pH 7 — a weak acid titrated with a strong base has a basic equivalence point (pH > 7) because the resulting salt’s conjugate base hydrolyzes water

Practice Tip: Before calculating any acid or base’s pH, first determine whether it’s strong or weak — a strong acid/base’s [H3O+] or [OH-] can be read directly from its given concentration, while a weak acid/base requires setting up an ICE table with its Ka or Kb to find the actual equilibrium concentration.

Common Mistake: Assuming every acid-base titration’s equivalence point occurs at exactly pH 7. This is only true for a strong acid titrated with a strong base; titrating a WEAK acid with a strong base produces a basic equivalence point (pH > 7), because the weak acid’s conjugate base (now present in solution) hydrolyzes water and produces excess OH-, and the reverse is true for a weak base titrated with a strong acid.

Important Questions:

• According to the Bronsted-Lowry definition, what is an acid and what is a base, and what is a conjugate acid-base pair? In the Bronsted-Lowry model, an acid is a substance that donates a proton (H+) in a reaction, and a base is a substance that accepts a proton. When an acid donates its proton, what remains is called its conjugate base; when a base accepts a proton, what results is called its conjugate acid. Every acid-base reaction therefore produces two conjugate acid-base pairs, one on each side of the equation.
• Why does a buffer solution resist significant changes in pH when small amounts of strong acid or base are added? A buffer contains both a weak acid and its conjugate base (or a weak base and its conjugate acid) in appreciable amounts, in equilibrium with each other. When a small amount of strong acid is added, the buffer’s conjugate base component reacts with and consumes most of the added H3O+; when strong base is added, the buffer’s weak acid component reacts with and consumes most of the added OH-. Because both added H3O+ and OH- are largely neutralized by the buffer’s own components rather than accumulating freely in solution, the overall pH changes only slightly instead of shifting dramatically.

Chapter 15: Equilibria of Other Reaction Classes

Difficulty: Hard · General Chemistry II · Key topics: precipitation and dissolution, Lewis acids and bases, coupled equilibria

This chapter extends the equilibrium framework beyond acid-base reactions to two more important reaction classes. It covers precipitation and dissolution equilibria for sparingly soluble ionic compounds, introducing the solubility product constant Ksp and using it to predict whether a precipitate will form (comparing the ion product to Ksp) and to calculate molar solubility. It then broadens the definition of acids and bases further with the Lewis model (an acid accepts an electron pair, a base donates one), which is more general than Bronsted-Lowry and explains reactions, like metal ion complexation, that don’t involve proton transfer at all. The chapter closes with coupled equilibria — systems where two or more equilibria interact and influence each other simultaneously, such as how solution pH affects the solubility of a salt whose anion is itself a weak base.

Key Points:

• The solubility product constant Ksp is the equilibrium constant for a sparingly soluble ionic compound’s dissolution; comparing the ion product (Q, calculated from current ion concentrations) to Ksp predicts whether a precipitate will form (Q > Ksp) or the solution remains unsaturated (Q < Ksp)
• Molar solubility (moles of compound that dissolve per liter) can be calculated from Ksp and vice versa, but the relationship between them depends on the compound’s specific ion ratio (stoichiometry), so it is NOT simply the square root of Ksp for every compound
• The Lewis acid-base model is broader than Bronsted-Lowry: a Lewis acid accepts an electron pair and a Lewis base donates one, which explains acid-base-like behavior (such as metal cation complexation with ligands) that doesn’t involve any proton transfer at all
• The common-ion effect reduces a sparingly soluble salt’s solubility when a solution already contains one of the salt’s own ions from another source, shifting the dissolution equilibrium back toward the solid (undissolved) form
• Coupled equilibria occur when two equilibrium systems share a common species and therefore influence each other — for example, a salt’s solubility can be pH-dependent if its anion is a weak base that reacts with H3O+ in a separate, linked acid-base equilibrium

Memory Tip: Every Lewis acid is an electron-pair ACCEPTOR and every Lewis base is an electron-pair DONOR — the opposite direction from how you might intuitively guess, so it helps to remember that a Lewis base “gives” electrons the same way a Bronsted-Lowry base effectively “takes” a proton by offering a lone pair to bond it.

Common Mistake: Assuming a compound’s molar solubility can always be found by simply taking the square root of its Ksp value. That shortcut only works for a salt with a 1:1 cation-to-anion ratio (like AgCl); a salt with a different ion ratio (like Ca(OH)2, which produces one Ca2+ and two OH- ions) requires setting up the correct Ksp expression with the appropriate stoichiometric exponents before solving for molar solubility.

Important Questions:

• What does it mean if a solution’s ion product (Q) is greater than a compound’s Ksp, and what will happen? If the calculated ion product Q exceeds the compound’s Ksp, the solution is supersaturated with respect to that compound — it contains more dissolved ions than the equilibrium (saturation) condition allows. The system will respond by precipitating solid compound out of solution, which removes dissolved ions and continues until Q decreases back down to equal Ksp, restoring equilibrium.
• How does the Lewis acid-base definition differ from and extend beyond the Bronsted-Lowry definition? The Bronsted-Lowry definition limits acids and bases to proton donors and acceptors specifically. The Lewis definition is broader: a Lewis acid is any species that accepts a pair of electrons, and a Lewis base is any species that donates a pair of electrons, regardless of whether a proton is transferred at all. This lets the Lewis model explain reactions like a metal cation (a Lewis acid) bonding with ligands (Lewis bases) to form a complex ion, which is acid-base-like behavior that the Bronsted-Lowry model, requiring proton transfer, cannot describe.

Chapter 16: Thermodynamics

Difficulty: Hard · General Chemistry II · Key topics: spontaneity, entropy, the second and third laws of thermodynamics, free energy

This chapter answers a question thermochemistry (Chapter 5) couldn’t: not just how much heat a reaction releases or absorbs, but whether the reaction will happen spontaneously at all under given conditions. It introduces entropy (S) as a measure of a system’s disorder or number of accessible microstates, and the Second Law of Thermodynamics, which states that the total entropy of the universe increases for any spontaneous process. The Third Law establishes an absolute reference point (zero entropy for a perfect crystal at absolute zero), enabling actual entropy calculations. The chapter’s culmination is Gibbs free energy (G), which combines enthalpy and entropy into a single quantity: a negative ΔG means a process is spontaneous under the given conditions, unifying the enthalpy and entropy factors into one predictive criterion.

Key Points:

• Entropy (S) measures a system’s disorder, or more precisely, the number of equivalent ways (microstates) its particles and energy can be arranged — higher entropy means more possible microstates
• The Second Law of Thermodynamics states that for any spontaneous process, the total entropy of the universe (system plus surroundings) increases; a process that would decrease the universe’s total entropy is never spontaneous
• The Third Law of Thermodynamics establishes that a perfect crystal at absolute zero (0 K) has exactly zero entropy, providing the reference point that lets absolute entropy values be tabulated and calculated for real substances
• Gibbs free energy combines enthalpy and entropy into a single spontaneity criterion: ΔG = ΔH − TΔS; a negative ΔG means the process is spontaneous as written, at that temperature
• A process can be spontaneous due to a favorable enthalpy term (exothermic, ΔH negative), a favorable entropy term (ΔS positive), or a combination of both — and because ΔG depends on temperature (through the TΔS term), a process can switch from nonspontaneous to spontaneous, or vice versa, purely by changing the temperature

Practice Tip: To predict how temperature affects a reaction’s spontaneity, check the SIGNS of ΔH and ΔS separately in ΔG = ΔH − TΔS: if they have the same sign, spontaneity depends on temperature (favorable at high T if ΔS is positive, favorable at low T if ΔH is negative); if they have opposite signs, the reaction is spontaneous at all temperatures or nonspontaneous at all temperatures.

Common Mistake: Assuming an exothermic reaction (negative ΔH) is automatically spontaneous. Spontaneity depends on the COMBINATION of enthalpy and entropy through ΔG = ΔH − TΔS, not on ΔH alone — an exothermic reaction with a sufficiently negative ΔS (large entropy decrease) can still have a positive ΔG and be nonspontaneous, especially at high temperature where the TΔS term dominates.

Important Questions:

• What does the Second Law of Thermodynamics state, and how does it differ from simply saying “a system’s entropy always increases”? The Second Law states that for any spontaneous process, the TOTAL entropy of the universe (the system plus its surroundings) increases — not necessarily the system’s entropy alone. A system’s own entropy can decrease during a spontaneous process (such as water freezing, which is spontaneous below 0°C despite the system becoming more ordered), as long as the surroundings’ entropy increases by a larger amount, keeping the universe’s total entropy change positive overall.
• How does Gibbs free energy (ΔG) combine enthalpy and entropy to predict whether a process is spontaneous? Gibbs free energy is calculated as ΔG = ΔH − TΔS, combining the enthalpy change (heat released or absorbed) and the entropy change (weighted by absolute temperature) into one quantity. A negative ΔG indicates the process is spontaneous as written under those conditions; a positive ΔG indicates it is nonspontaneous (though the reverse process would be spontaneous); this single criterion correctly accounts for cases where enthalpy and entropy favor opposite outcomes, which neither factor alone could resolve.

Chapter 17: Electrochemistry

Difficulty: Hard · General Chemistry II · Key topics: review of redox chemistry, galvanic cells, electrode and cell potentials, potential/free energy/equilibrium, batteries and fuel cells, corrosion, electrolysis

This chapter connects chemistry directly to electricity, by studying oxidation-reduction (redox) reactions where electrons transfer from one species to another. It reviews assigning oxidation states and balancing redox equations, then introduces galvanic (voltaic) cells, which harness a spontaneous redox reaction’s electron transfer to generate usable electrical current by physically separating the oxidation and reduction half-reactions. The chapter covers standard electrode potentials and how to combine them to predict a cell’s overall voltage and whether a given redox reaction is spontaneous, then connects cell potential mathematically to Gibbs free energy and the equilibrium constant — three different ways of describing the identical underlying spontaneity. It closes with practical applications: batteries, fuel cells, corrosion (unwanted spontaneous oxidation), and electrolysis, which uses external electrical energy to force an otherwise nonspontaneous redox reaction to occur.

Key Points:

• A galvanic (voltaic) cell generates electrical current from a SPONTANEOUS redox reaction by physically separating the oxidation half-reaction (at the anode) and the reduction half-reaction (at the cathode), forcing electrons to flow through an external wire between them
• A standard cell potential (E°cell) is calculated from the two half-reactions’ standard reduction potentials; a positive E°cell indicates the overall redox reaction is spontaneous under standard conditions
• Cell potential, Gibbs free energy, and the equilibrium constant are three connected descriptions of the same underlying spontaneity: ΔG° = −nFE°cell and ΔG° = −RT ln K, linking all three quantities together mathematically
• Electrolysis uses an external power source to force electrons through a cell in the reverse of their spontaneous direction, driving an otherwise NONSPONTANEOUS redox reaction — the electrical energy input compensates for the reaction’s unfavorable (positive) ΔG
• Corrosion (such as iron rusting) is itself a spontaneous, unwanted electrochemical process, and understanding galvanic cells explains both why it happens and how sacrificial anodes (a more easily oxidized metal) can be used to protect a structure from corroding instead

Memory Tip: “An Ox, Red Cat” — oxidation happens at the Anode, and reduction happens at the Cathode, in BOTH galvanic and electrolytic cells. What changes between the two cell types is only whether the overall reaction is spontaneous (galvanic, generating current) or driven by an external power source (electrolytic, consuming current).

Common Mistake: Assuming a galvanic cell and an electrolytic cell always have their anode and cathode assigned the same electrical charge sign. In a galvanic cell, the anode is negative and cathode positive; in an ELECTROLYTIC cell, an external power source reverses the charge signs (anode becomes positive, cathode negative) even though oxidation still occurs at the anode and reduction still occurs at the cathode in both cases — the electrode NAMES track the reaction type, not a fixed charge sign.

Important Questions:

• What is the difference between a galvanic cell and an electrolytic cell? A galvanic (voltaic) cell harnesses a SPONTANEOUS redox reaction, using the reaction’s own favorable free energy change to generate usable electrical current that flows through an external circuit. An electrolytic cell does the opposite: it uses an external electrical power source to force electrons through a NONSPONTANEOUS redox reaction, driving it forward despite its unfavorable free energy change — essentially running a galvanic cell’s chemistry in reverse by supplying energy from outside.
• How are cell potential, Gibbs free energy, and the equilibrium constant related to each other? These three quantities are mathematically linked descriptions of the same underlying reaction spontaneity: ΔG° = −nFE°cell connects free energy to cell potential (where n is moles of electrons transferred and F is Faraday’s constant), and ΔG° = −RT ln K connects free energy to the equilibrium constant. A reaction with a positive standard cell potential has a negative ΔG° (spontaneous) and a large equilibrium constant K (strongly favors products) — all three facts describing the identical reaction from different but consistent angles.

Chapter 18: Representative Metals, Metalloids, and Nonmetals

Difficulty: Medium · General Chemistry II · Key topics: periodicity, occurrence and preparation of representative metals, structure/properties of metalloids and nonmetals, hydrogen, carbonates, nitrogen, phosphorus, oxygen, sulfur, halogens, noble gases

This descriptive chemistry chapter surveys the main-group (representative) elements region-by-region across the periodic table, applying the periodic trends and bonding theory built up over the previous seventeen chapters to real elements and their compounds. It covers how representative metals are found in nature and industrially extracted/prepared, then works systematically through the nonmetal groups — hydrogen’s unique dual position, carbon and carbonates, nitrogen and phosphorus (both essential to biological molecules), oxygen and sulfur, the halogens (the most reactive nonmetal family), and the largely inert noble gases. Throughout, the chapter connects each element family’s chemical behavior back to its position and electron configuration on the periodic table, rather than presenting facts as an unconnected list.

Key Points:

• Periodicity (the recurring pattern of properties established in Chapter 6) directly predicts and organizes the descriptive chemistry of every main-group element family covered in this chapter, rather than requiring each element’s behavior to be memorized independently
• Hydrogen occupies a unique position on the periodic table — it can behave somewhat like an alkali metal (losing its one electron) or somewhat like a halogen (gaining one electron to complete a duet), and doesn’t fit cleanly into either group
• Metalloids (such as silicon, germanium, and arsenic) have intermediate properties between metals and nonmetals, including semiconductor electrical behavior that makes them essential to modern electronics
• Nitrogen and phosphorus, though in the same periodic group, differ sharply in reactivity — N2’s very strong triple bond makes elemental nitrogen gas relatively unreactive at room temperature, while elemental phosphorus is considerably more reactive
• The halogens (Group 17) are the most reactive nonmetal family, readily gaining one electron to achieve a noble-gas electron configuration, while the noble gases (Group 18) are the LEAST reactive family, already possessing a full valence shell

Memory Tip: Rather than memorizing each element family’s chemistry as a separate, disconnected fact set, always start from that family’s position on the periodic table and its valence electron configuration (from Chapter 6) — group trends in reactivity, bonding preference, and common oxidation states almost always follow directly from where an element sits.

Common Mistake: Assuming elements in the same periodic group always behave nearly identically simply because they share the same group number. While group members do share the same valence electron count (and therefore similar bonding preferences), reactivity and physical properties can still vary substantially within a group — nitrogen gas (N2) is comparatively unreactive due to its strong triple bond, while phosphorus, directly below it in the same group, is considerably more reactive.

Important Questions:

• Why doesn’t hydrogen fit cleanly into either the alkali metal group or the halogen group of the periodic table? Hydrogen has just one electron in its single occupied energy level, which it can either lose (like an alkali metal, becoming H+) or gain one more electron to achieve a filled duet (somewhat like a halogen achieving a filled octet). Because it can display chemical behavior resembling either family depending on the reaction, but shares neither family’s other characteristic properties (like alkali metals’ softness and high reactivity as solids, or halogens’ diatomic gas/liquid states), hydrogen is usually placed in its own unique position rather than firmly belonging to either group.
• What makes metalloids chemically and physically distinct from both metals and nonmetals? Metalloids display properties intermediate between metals and nonmetals — they can have a somewhat metallic luster like metals, yet are typically brittle like nonmetals rather than malleable. Their most technologically important distinguishing feature is semiconductor electrical behavior: unlike metals (which conduct electricity well) or most nonmetals (which are insulators), metalloids like silicon conduct electricity only moderately and in ways that can be precisely controlled, making them the foundation of modern electronic devices.

Chapter 19: Transition Metals and Coordination Chemistry

Difficulty: Hard · General Chemistry II · Key topics: occurrence/preparation/properties of transition metals, coordination chemistry, spectroscopic and magnetic properties of coordination compounds

This chapter turns from main-group elements to the transition metals, whose partially filled d subshells give them chemistry distinctly different from anything covered so far — multiple stable oxidation states, and often vividly colored compounds. It covers where transition metals occur naturally and how they’re extracted and processed industrially, then introduces coordination chemistry — complex ions formed when a central transition metal ion bonds to surrounding ligands (Lewis bases) through coordinate covalent bonds. The chapter explains how to name and determine the geometry of coordination compounds, then closes with the spectroscopic and magnetic properties that make transition metal complexes so visually and practically distinctive: their vivid colors (from d-orbital electron transitions absorbing specific wavelengths of visible light) and their varying magnetic behavior (from unpaired d electrons).

Key Points:

• Transition metals characteristically exhibit MULTIPLE stable oxidation states (unlike most main-group metals, which typically have one dominant oxidation state), because electrons can be removed from both the outer s subshell and the underlying d subshell relatively easily
• A coordination compound consists of a central transition metal ion (or atom) surrounded by ligands — Lewis bases that donate an electron pair to form a coordinate covalent bond to the metal
• The coordination number (how many ligand donor atoms bond directly to the central metal) determines a complex’s geometry — commonly 4 (tetrahedral or square planar) or 6 (octahedral)
• Transition metal complexes are often vividly colored because their partially filled d orbitals split into different energy levels when ligands bond to the metal, and electrons absorb specific wavelengths of visible light to jump between these split d-orbital levels
• A complex’s magnetic behavior (paramagnetic if it has unpaired d electrons, diamagnetic if all d electrons are paired) depends on both the metal’s specific d-electron count and the ligands’ field strength, which determines exactly how the d orbitals split and how electrons fill them

Memory Tip: Transition metal complexes are colored specifically because of electron transitions BETWEEN split d-orbital energy levels, not because of the bulk metal’s own inherent color — the identical metal ion can produce dramatically different colors depending entirely on which ligands are bonded to it, since different ligands split the d orbitals by different amounts.

Common Mistake: Assuming a transition metal ion has one single, fixed oxidation state the way many main-group metals do. Most transition metals commonly form compounds in several different oxidation states (iron commonly forms both Fe2+ and Fe3+ compounds, for example) — always check which specific oxidation state a compound’s name or formula indicates, rather than assuming a single default value for that metal.

Important Questions:

• What is a coordination compound, and what role do ligands play in it? A coordination compound consists of a central metal ion or atom (usually a transition metal) bonded to a surrounding group of ligands. Ligands are Lewis bases — molecules or ions with an available lone pair of electrons — that each donate an electron pair to form a coordinate covalent bond directly to the central metal, together determining the complex’s coordination number and geometry.
• Why are many transition metal coordination compounds vividly colored? When ligands bond to a transition metal’s central ion, they cause the metal’s normally equal-energy (degenerate) d orbitals to split into two or more different energy levels, with the exact splitting pattern and magnitude depending on the specific ligands and geometry involved. Electrons in the lower-energy split d orbitals can absorb a specific wavelength of visible light to jump up to the higher-energy split d orbitals; because only certain wavelengths are absorbed, the wavelengths that pass through or reflect back are what the human eye perceives as the complex’s characteristic color.

Chapter 20: Organic Chemistry

Difficulty: Medium · General Chemistry II · Key topics: hydrocarbons, alcohols and ethers, aldehydes/ketones/carboxylic acids/esters, amines and amides

This chapter introduces organic chemistry — the chemistry of carbon-based compounds that make up all known living things — building from the previous nineteen chapters’ bonding and structure principles applied specifically to carbon’s unique bonding versatility. It starts with hydrocarbons (compounds containing only carbon and hydrogen): alkanes, alkenes, alkynes, and aromatic compounds, along with IUPAC naming conventions and the concept of isomerism (identical molecular formula, different structure). It then works through the major carbon-based functional groups in roughly increasing oxidation-state order: alcohols and ethers, then aldehydes, ketones, carboxylic acids, and esters, and finally nitrogen-containing amines and amides — each functional group defined by a specific characteristic group of atoms that determines the compound’s chemical behavior regardless of the rest of the molecule.

Key Points:

• Carbon’s unique ability to form four strong covalent bonds, including stable chains and rings with itself, is what makes the vast diversity of organic compounds possible — no other element forms nearly as many distinct stable compounds
• Isomers share the identical molecular formula but have different structural arrangements of their atoms, and can have substantially different physical and chemical properties despite that identical formula
• A functional group is a specific, recognizable arrangement of atoms within a larger molecule (like -OH for alcohols, or -COOH for carboxylic acids) that determines the compound’s characteristic chemical reactivity, largely independent of the rest of the molecule’s structure
• Hydrocarbons are classified by their carbon-carbon bonding: alkanes have only single bonds, alkenes contain at least one double bond, alkynes contain at least one triple bond, and aromatic compounds contain a benzene-type ring with delocalized pi electrons
• Organic functional groups follow a rough increasing-oxidation trend from alcohols through aldehydes/ketones to carboxylic acids, and many common reactions (like an alcohol oxidizing to an aldehyde, and then further to a carboxylic acid) move a molecule along this trend

Practice Tip: When identifying an unfamiliar organic molecule’s likely chemical behavior, first locate and identify its functional group(s) — the specific atoms and bonding pattern present (like -OH, C=O, or -COOH) — rather than trying to reason from the molecule’s overall size or the total number of carbons it contains.

Common Mistake: Assuming two compounds with the identical molecular formula must have identical properties. Isomers share the same molecular formula but differ in how their atoms are actually arranged/bonded, and this structural difference can produce dramatically different physical properties (boiling point, solubility) and chemical reactivity — molecular formula alone never fully determines a compound’s identity or behavior.

Important Questions:

• What is a functional group, and why is it useful for predicting an organic molecule’s chemical behavior? A functional group is a specific, recognizable arrangement of atoms within a larger organic molecule (such as the -OH group defining an alcohol, or the -COOH group defining a carboxylic acid) that reacts in characteristic, predictable ways. It’s useful because a molecule’s chemical reactivity is determined largely by its functional group(s) rather than by the rest of its carbon skeleton, so recognizing the functional group present lets you predict how an unfamiliar molecule will likely react, based on that functional group’s known general behavior.
• What are isomers, and why can two isomers have very different properties despite sharing the same molecular formula? Isomers are two or more compounds that share the identical molecular formula (the same numbers of each type of atom) but differ in how those atoms are actually connected or arranged in space. Because a molecule’s physical and chemical properties depend heavily on its specific structure — which atoms are bonded to which, and in what three-dimensional arrangement — not just on which atoms are present, isomers can have substantially different boiling points, reactivities, and even biological effects, despite having identical molecular formulas.

Chapter 21: Nuclear Chemistry

Difficulty: Medium · General Chemistry II · Key topics: nuclear structure and stability, nuclear equations, radioactive decay, transmutation and nuclear energy, uses of radioisotopes, biological effects of radiation

This closing chapter shifts focus from the electronic structure that has governed every previous chapter’s chemistry to the nucleus itself, where a fundamentally different set of rules applies. It covers what makes a nucleus stable or unstable (the neutron-to-proton ratio and binding energy), how to balance nuclear equations (conserving mass number and atomic number rather than simply conserving atoms as in ordinary chemical equations), and the major modes of radioactive decay — alpha, beta, gamma, and others — each with distinct penetrating power and biological hazard. The chapter covers half-life as the standard way to quantify decay rate, nuclear transmutation and fission/fusion as sources of nuclear energy, and closes with practical applications (medical imaging and treatment via radioisotopes) alongside the biological effects and safety considerations of ionizing radiation exposure.

Key Points:

• Nuclear reactions involve changes to the nucleus itself (proton/neutron composition), unlike ordinary chemical reactions, which involve only electron rearrangement while the nucleus remains completely unchanged
• Balancing a nuclear equation requires conserving BOTH mass number (the sum of protons and neutrons, superscript) and atomic number (the number of protons, subscript) on both sides — not conserving individual atoms of specific elements the way ordinary chemical equations do
• Radioactive decay follows first-order kinetics, characterized by a fixed half-life — the time for exactly half of a radioactive sample to decay, which is completely independent of the sample’s initial amount or its temperature/chemical environment
• Alpha particles are the least penetrating but most damaging if ingested/inhaled; beta particles penetrate further; gamma rays are the most penetrating and require dense shielding (like thick lead or concrete) — a radioisotope’s decay mode largely determines its handling and safety precautions
• Nuclear fission (splitting a heavy nucleus) and nuclear fusion (combining light nuclei) both release enormous amounts of energy relative to chemical reactions, because nuclear binding energy differences are vastly larger than chemical bond energy differences

Memory Tip: Nuclear half-life is completely constant and unaffected by temperature, pressure, or chemical environment — unlike a chemical reaction’s rate (Chapter 12), which speeds up with heat or a catalyst, a radioisotope decays at exactly the same statistical rate no matter what chemical compound it’s part of or what conditions it’s kept under.

Common Mistake: Trying to balance a nuclear equation the way an ordinary chemical equation is balanced, by conserving the identity/number of each specific element. Nuclear equations must instead conserve total mass number and total atomic number across the reaction, since the elements themselves actually transmute (change identity) during radioactive decay — the parent and daughter nuclides are frequently different elements entirely.

Important Questions:

• What is nuclear half-life, and what factors does it depend on? Nuclear half-life is the time required for exactly half of a given sample of a radioactive isotope to decay. It is a fixed, characteristic property of each specific radioisotope and is completely independent of the sample’s initial size, temperature, pressure, or the chemical compound the radioactive atoms happen to be part of — unlike ordinary chemical reaction rates, nuclear decay rates cannot be sped up or slowed down by any ordinary physical or chemical means.
• Why must both mass number and atomic number be conserved when balancing a nuclear equation, and what does this reflect physically? Mass number (protons plus neutrons) and atomic number (protons alone) must both balance across a nuclear equation because these are the two quantities nuclear reactions actually conserve at the level of nucleons; unlike ordinary chemical reactions, a nuclear reaction can change WHICH element is present, since it directly changes the number of protons in the nucleus. Conserving mass number and atomic number reflects the physical reality that nuclear reactions can transmute one element into a genuinely different element, something no ordinary chemical reaction can ever do.

Download Chemistry 2e PDF (Free)

This book is free from its official source, OpenStax. Click below to download the complete PDF — a free web-based reader edition (with per-section, linkable pages) is also available on the OpenStax site if you’d rather read online.

How to Study This Book

Chapters 1–10 form General Chemistry I: Essential Ideas through Liquids and Solids — matter, atomic structure, stoichiometry, thermochemistry, bonding, gases, and the condensed phases. This block is largely self-contained and builds the foundation every later chapter assumes.

Chapters 11–17 form the core of General Chemistry II: Solutions, Kinetics, Equilibrium, Acid-Base Equilibria, Equilibria of Other Reaction Classes, Thermodynamics, and Electrochemistry — all six of the equilibrium/thermodynamics chapters (13–17) build directly on the equilibrium concepts introduced in Chapter 13, so treat that chapter as load-bearing for everything after it.

Chapter 6 (Electronic Structure and Periodic Properties) is worth extra time even though it appears early — its electron-configuration and periodic-trend material is assumed without re-explanation in Chapters 7, 8, 18, and 19.

Chapters 7 and 8 (Chemical Bonding and Molecular Geometry, then Advanced Theories of Covalent Bonding) should be read in that order — Chapter 8’s valence bond and molecular orbital theories are explicitly built as extensions of Chapter 7’s Lewis structures and VSEPR, not a fresh start.

Chapters 18–21 (Representative Metals/Metalloids/Nonmetals, Transition Metals and Coordination Chemistry, Organic Chemistry, Nuclear Chemistry) are descriptive-chemistry chapters that some one-year general chemistry courses cover only selectively, or skip in favor of a dedicated organic chemistry course later — check your syllabus before budgeting heavy study time here.

This is a two-semester, 21-chapter general chemistry text — if your course covers only one semester, confirm with your instructor whether it’s Chapters 1–10 or a different split, since some programs divide the material differently than OpenStax’s own chapter numbering suggests.

Used In These Programs

This book is used for the General Chemistry I & II course sequence in: BS Chemistry, Pre-Medical, Engineering, and other STEM degree programs. Browse all Chemistry books or all Chemistry category books.

Who Should Read This

Chemistry 2e is written for a first-year university student taking a two-semester general chemistry sequence — typically a BS Chemistry, Pre-Medical, Engineering, or other STEM major with no prior college-level chemistry background required, though basic algebra and comfort with scientific notation are assumed throughout. Its traditional, example-and-exercise-driven style, with large end-of-section problem sets and a full Chapter Review for every chapter, suits a student who learns best by working through many worked examples before attempting problems independently.

Applicable Universities

This book is useful for students at Pakistani universities offering BS Chemistry, Pre-Medical, Engineering, or other STEM programs, including Punjab University, Virtual University, COMSATS, FAST, UET, NUST, GIKI, and other HEC-recognized institutions, where a two-semester General Chemistry sequence is a standard first-year requirement.

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Related Books

• University Physics Volume 1 – OpenStax (a common concurrent course for STEM students)
• Calculus Volume 1 – Strang & Herman (OpenStax)

Chemistry 2e is this project’s first Chemistry-category book, covering the complete two-semester General Chemistry sequence for BS Chemistry, Pre-Medical, and Engineering students. Browse more Physics books or Mathematics books for the rest of your semester.

Chemistry 2e, by Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson, and contributors. OpenStax, Rice University. Free under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 licence. Access for free at https://openstax.org/details/books/chemistry-2e