Chemistry: Atoms First 2e PDF Download – OpenStax (BS Chemistry)

Students beginning General Chemistry can download “Chemistry: Atoms First 2e”, an OpenStax textbook by Paul Flowers, Klaus Theopold, Richard Langley, Edward J. Neth, William R. Robinson and contributors, free from the official OpenStax site. It teaches the same general chemistry as OpenStax Chemistry 2e, but in a different order: atoms, electron structure and bonding come first, and stoichiometry waits until Chapter 6.

The 21 chapters run from measurement and atomic theory through bonding, gases, thermodynamics, equilibrium, electrochemistry and kinetics to descriptive inorganic, nuclear and organic chemistry. It suits BS Chemistry, pre-medical and engineering students who prefer to understand atoms and molecules before running reaction calculations, and it works for a two-semester or three-semester course. No prior college chemistry is needed, only comfortable algebra.

Book Overview

Chapter List

Chapter 1: Essential Ideas

Difficulty: Easy · Key topics: chemistry in context, states of matter, classification of matter, physical and chemical properties, SI units, density, significant figures, accuracy and precision, dimensional analysis

This opening chapter sets the vocabulary and habits the whole course relies on. It places chemistry as the central science, outlines the scientific method, and introduces the macroscopic, microscopic and symbolic domains chemists work in. Matter is then classified by state (including plasma) and by composition, separating pure substances from homogeneous and heterogeneous mixtures. Physical and chemical changes and extensive versus intensive properties follow. The chapter closes with SI units, density, significant figures and rounding, accuracy versus precision, and the factor-label method, including temperature conversions between Celsius, Fahrenheit and Kelvin.

Key Points:

• Chemistry is described in three domains: macroscopic (what you observe), microscopic (atoms and molecules) and symbolic (formulas and equations)
• Solids have fixed shape and volume, liquids fixed volume only, gases neither; plasma is a gas containing charged particles
• Homogeneous mixtures have the same composition throughout; heterogeneous mixtures vary from point to point
• Extensive properties (mass, heat) depend on the amount of matter; intensive properties (density, temperature) do not
• Every measurement carries a number, a unit and an uncertainty; density = mass ÷ volume
• Dimensional analysis converts units by multiplying by conversion factors arranged so the unwanted unit cancels

Practice Tip: Convert your own height from centimetres to inches and then to metres using only conversion factors written as fractions. Cancelling the units on paper shows why the method works, and you will use it in every later chapter.

Common Mistake: Copying every digit from the calculator. If the data have three significant figures, the answer cannot honestly have eight.

Important Questions:

• What is the difference between an extensive and an intensive property? An extensive property depends on how much matter is present (mass, volume, heat); an intensive property does not (density, temperature, colour).
• How many significant figures are in 0.00450, and how does that limit a product? Three, because leading zeros do not count but the trailing zero after the decimal point does; a product can have no more significant figures than the least precise factor.

Chapter 2: Atoms, Molecules, and Ions

Difficulty: Easy · Key topics: Dalton’s postulates, laws of definite and multiple proportions, cathode rays, nuclear model, isotopes, atomic mass, molecular and empirical formulas, the mole

The chapter retraces how the atomic picture was built. It starts with Dalton’s postulates and the laws of definite and multiple proportions, then follows the experiments that opened the atom: Thomson’s cathode rays (the electron), Millikan’s measurement of its charge, Rutherford’s gold foil experiment (the nucleus), Chadwick’s neutron and Soddy’s isotopes. Atomic number, mass number, isotopes and average atomic mass are defined next. The last section moves to molecular, empirical and structural formulas, isomers, and the mole with Avogadro’s number and molar mass.

Key Points:

• Law of definite proportions: a pure compound always contains the same elements in the same mass proportions
• Thomson found the electron, Rutherford the nucleus and Chadwick the neutron; Millikan fixed the electron’s charge
• Atomic number (Z) is the number of protons and defines the element; mass number is protons plus neutrons
• Isotopes share Z but differ in neutrons; the periodic-table atomic mass is the abundance-weighted average
• One amu is exactly 1/12 of a carbon-12 atom, about 1.6605 × 10⁻²⁴ g
• One mole is 6.022 × 10²³ entities (Avogadro’s number); molar mass is the mass of one mole in grams

Memory Tip: Remember the three discoveries in order as T-R-C: Thomson (electron), Rutherford (nucleus), Chadwick (neutron).

Common Mistake: Mixing up mass number and atomic mass. The mass number is a whole-number count of protons and neutrons; the atomic mass is a measured average in amu, such as 35.45 for chlorine.

Important Questions:

• How do isotopes of one element differ, and why is a periodic-table atomic mass not a whole number? Isotopes have the same number of protons but different numbers of neutrons; the tabulated value is the average of the naturally occurring isotope masses weighted by abundance.
• How many water molecules are in 18.0 g of water? The molar mass of H₂O is about 18.0 g/mol, so 18.0 g is about 1.00 mol, which is 6.02 × 10²³ molecules.

Chapter 3: Electronic Structure and Periodic Properties of Elements

Difficulty: Hard · Key topics: electromagnetic radiation, photoelectric effect, line spectra, Bohr model, quantum numbers, orbitals, electron configurations, periodic trends, the periodic table, ionic and molecular compounds

Starting from the wave behaviour of light (c = λν) and its particle behaviour (E = hν), the chapter shows why classical physics failed on blackbody radiation, the photoelectric effect and atomic line spectra. Bohr’s quantized orbits for hydrogen come next, followed by the quantum-mechanical model with wavefunctions, orbitals and the four quantum numbers. These feed electron configurations built with the Pauli exclusion principle and Hund’s rule, and then the periodic trends in radius, ionization energy and electron affinity. The chapter ends by organizing the elements in the periodic table and sorting compounds into ionic and molecular types.

Key Points:

• c = λν for any electromagnetic wave and E = hν for a photon, so light shows both wave and particle behaviour
• Bohr’s model quantizes the hydrogen electron’s orbits; a photon is emitted or absorbed when the electron changes orbit
• The four quantum numbers are n (shell), l (subshell, 0 to n − 1), m_l (orientation) and m_s (spin, ±½)
• Orbitals fill in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p and so on, obeying the Pauli exclusion principle and Hund’s rule
• Atomic radius grows down a group and shrinks across a period; ionization energy mostly does the opposite
• Metals tend to form cations and nonmetals anions; ionic compounds usually pair a metal with a nonmetal, covalent compounds two nonmetals

Memory Tip: Write the subshells in rows (1s; 2s 2p; 3s 3p 3d; 4s 4p 4d 4f) and draw diagonal arrows through them. Following the arrows reproduces the filling order in about ten seconds.

Common Mistake: Removing d electrons first when writing a transition-metal cation. Electrons leave the outermost shell first, so the 4s electrons go before the 3d electrons.

Important Questions:

• State the Pauli exclusion principle and Hund’s rule. No two electrons in an atom can have the same set of four quantum numbers; in degenerate orbitals, electrons occupy separate orbitals with parallel spins before any pairing occurs.
• Why does ionization energy generally increase across a period? The effective nuclear charge rises while electrons are added to the same shell, so electrons are held more tightly; filled and half-filled subshells cause small exceptions.

Chapter 4: Chemical Bonding and Molecular Geometry

Difficulty: Medium · Key topics: ionic bonding, covalent bonding, electronegativity, nomenclature, Lewis structures, octet rule, formal charge, resonance, VSEPR, molecular polarity

This chapter explains how atoms join and what shape the result takes. Ionic bonding is described through the stable electron configurations of cations and anions, and covalent bonding through shared pairs and electronegativity, which decides whether a bond is nonpolar, polar or ionic. A full section teaches naming of ionic compounds, molecular compounds, binary acids and oxyacids. Lewis symbols and structures follow, including octet-rule exceptions, then formal charge and resonance. The chapter ends with VSEPR theory, the difference between electron-pair geometry and molecular structure, and how bond dipoles combine into molecular polarity.

Key Points:

• Representative-metal cations usually reach a noble-gas or completely filled shell configuration; nonmetal anions fill their valence shell
• The electronegativity difference between two atoms decides whether a bond is nonpolar covalent, polar covalent or ionic
• Naming: the metal first, the nonmetal ending in −ide, a Roman numeral for variable charge, and Greek prefixes for molecular compounds
• Most Lewis structures obey the octet rule; exceptions are odd-electron, electron-deficient and hypervalent species
• The preferred Lewis structure has formal charges as close to zero as possible; a resonance hybrid is the average of the resonance forms
• VSEPR places electron regions as far apart as possible; molecular structure describes only the atoms, not the lone pairs

Memory Tip: For any VSEPR problem, first count the regions of electron density around the central atom (a lone pair counts as one region, and so does any bond, single or multiple). Name the electron-pair geometry, then ignore the lone pairs to name the molecular shape.

Common Mistake: Assuming a molecule with polar bonds must be polar. In CO₂ the two polar bonds point in opposite directions and cancel, so the molecule has no dipole moment.

Important Questions:

• What are the electron-pair geometry and molecular structure of H₂O? Four regions of electron density give a tetrahedral electron-pair geometry; two of them are lone pairs, so the molecular structure is bent.
• When does a molecule containing polar bonds have no overall dipole moment? When its shape is symmetric enough for the bond dipoles to cancel, as in linear CO₂ or tetrahedral CCl₄.

Chapter 5: Advanced Theories of Bonding

Difficulty: Hard · Key topics: valence bond theory, sigma and pi bonds, hybrid orbitals, multiple bonds, molecular orbital theory, bond order, paramagnetism, band theory

This chapter picks up where Lewis structures and VSEPR stop working. Valence bond theory describes a bond as overlap of atomic orbitals, with σ bonds along the bond axis and π bonds from sideways overlap, and it adds dipole moments and partial ionic character. Hybrid orbitals (sp through sp³d²) are then used to explain the geometries found in VSEPR, and a short section shows how multiple bonds and resonance fit this picture. The final section builds molecular orbital diagrams for second-period diatomic molecules, bond order, magnetic behaviour and a brief look at band theory.

Key Points:

• A σ bond forms by overlap along the line joining the nuclei; a π bond forms by sideways overlap and has a node along that line
• Hybridization follows the number of electron regions: two give sp, three sp², four sp³, five sp³d, six sp³d²
• Hybrid orbitals either form σ bonds or hold lone pairs; π bonds use unhybridized p or d orbitals
• A double bond is one σ plus one π bond; a triple bond is one σ plus two π bonds
• In-phase combinations of atomic orbitals give bonding MOs; out-of-phase combinations give antibonding MOs
• Bond order = (bonding electrons − antibonding electrons) ÷ 2; unpaired electrons make a species paramagnetic

Practice Tip: Draw the molecular orbital energy diagrams for N₂ and O₂ on paper, fill in the valence electrons, and count bond order and unpaired electrons. The diagram for O₂ shows its magnetism immediately.

Common Mistake: Ignoring lone pairs when assigning hybridization. Lone pairs sit in hybrid orbitals too, so the central atoms in NH₃ and H₂O are sp³ even though they form fewer than four bonds.

Important Questions:

• What is the hybridization of each carbon in ethene (C₂H₄), and why? Each carbon has three regions of electron density (two C–H σ bonds and one C–C σ bond), so it is sp²; the leftover unhybridized p orbitals overlap sideways to form the π bond.
• How does molecular orbital theory explain why O₂ is paramagnetic? The last two valence electrons occupy separate degenerate π* orbitals with parallel spins, so O₂ has two unpaired electrons, something the Lewis structure does not predict.

Chapter 6: Composition of Substances and Solutions

Difficulty: Medium · Key topics: formula mass, percent composition, empirical and molecular formulas, molarity, dilution, mass percentage, ppm and ppb

Once the atomic foundation is in place, this chapter turns to quantities. Formula mass is built from average atomic masses, and the equivalence between amu and g/mol links single particles to laboratory samples. Percent composition is used to work backward from experimental data to an empirical formula and, with the molar mass, to a molecular formula. The second half introduces solutions: solute, solvent, molarity, and the dilution equation, then other concentration units such as mass percentage, volume percentage, mass-volume percentage, parts per million and parts per billion.

Key Points:

• Formula mass is the sum of the average atomic masses in the formula; formula mass in amu equals molar mass in g/mol numerically
• Percent composition = (mass of element ÷ mass of compound) × 100
• Empirical formula: convert masses to moles, then divide by the smallest to get the simplest whole-number ratio
• Molecular formula = (molar mass ÷ empirical formula mass) × empirical formula
• Molarity (M) = moles of solute ÷ litres of solution; dilution follows M₁V₁ = M₂V₂
• Percentage units and ppm or ppb are preferred where very small concentrations occur, as in environmental and medical work

Practice Tip: For empirical-formula problems, assume a 100 g sample so that each percentage becomes grams directly. Convert to moles, divide by the smallest value, and multiply by 2 or 3 if the ratios end in .5 or .33.

Common Mistake: Using the volume of the solvent instead of the volume of the final solution in a molarity calculation.

Important Questions:

• How many grams of NaCl are needed to prepare 250 mL of a 0.100 M solution? Moles = 0.100 mol/L × 0.250 L = 0.0250 mol; multiplied by the molar mass 58.44 g/mol this gives about 1.46 g.
• A compound contains 40.0% C, 6.7% H and 53.3% O by mass. What is its empirical formula? In 100 g there are about 3.33 mol C, 6.65 mol H and 3.33 mol O; dividing by 3.33 gives CH₂O.

Chapter 7: Stoichiometry of Chemical Reactions

Difficulty: Medium · Key topics: balancing equations, net ionic equations, precipitation, acid-base and redox reactions, mole ratios, limiting reactant, percent yield, titration, gravimetric and combustion analysis

This chapter makes the balanced equation a calculating tool. It begins with writing and balancing equations, including complete and net ionic equations for reactions in water. The classification section sorts reactions into precipitation, acid-base and oxidation-reduction types, with oxidation numbers and the half-reaction method for balancing redox equations. Mass-mole-mass calculations follow, then limiting reactants, theoretical, actual and percent yield, and a short look at atom economy. The chapter closes with quantitative analysis by titration, gravimetric analysis and combustion analysis.

Key Points:

• A balanced equation has the same number of atoms of each element on both sides; coefficients give mole ratios
• A net ionic equation omits spectator ions and shows only the species that actually change
• Precipitation forms an insoluble product, acid-base reactions transfer H⁺, and redox reactions change oxidation numbers
• The limiting reactant is the one consumed first and fixes the theoretical yield
• Percent yield = (actual yield ÷ theoretical yield) × 100
• Titration uses titrant volume, gravimetric analysis uses the mass of a separated product, and combustion analysis finds elemental composition

Practice Tip: Write every stoichiometry problem as a chain: grams of A to moles of A, then moles of B using the coefficients, then grams of B. Label every factor with units so the chain cancels cleanly.

Common Mistake: Using the coefficients as mass ratios. Coefficients give ratios of moles, so grams must be converted to moles first.

Important Questions:

• What is a limiting reactant and how do you find it? It is the reactant that runs out first; compare the moles available with the moles the equation requires, or calculate the product each reactant could make and take the smaller amount.
• A reaction has a theoretical yield of 25.0 g and an actual yield of 20.0 g. What is the percent yield? (20.0 ÷ 25.0) × 100 = 80.0%.

Chapter 8: Gases

Difficulty: Medium · Key topics: gas pressure, Boyle’s, Charles’s and Avogadro’s laws, ideal gas law, gas density, Dalton’s law, Graham’s law, kinetic-molecular theory, van der Waals equation

The chapter starts with pressure, its units and how barometers and manometers measure it. It then connects pressure, volume, temperature and amount of gas through the individual laws (Amontons, Charles, Boyle, Avogadro) and combines them into PV = nRT. Applications follow: gas density and molar mass, Dalton’s law of partial pressures, collecting gas over water, and stoichiometry with gaseous reactants. Effusion and diffusion lead to Graham’s law, the kinetic-molecular theory explains the laws at the molecular level, and the van der Waals equation corrects for non-ideal behaviour.

Key Points:

• Pressure is force per unit area; common units are pascal, atm, torr and bar
• Boyle: P and V are inversely proportional at constant T and n; Charles: V is proportional to T at constant P and n
• Ideal gas law: PV = nRT, with T always in kelvin
• Dalton’s law: the total pressure of a mixture is the sum of the partial pressures of its components
• Graham’s law: effusion rate is inversely proportional to the square root of molar mass
• Real gases deviate at high pressure and low temperature because molecules have volume and attract one another

Practice Tip: Solve one ideal gas problem twice, once with pressure in atm and the matching value of R, and once in kPa with its matching R. Getting the same volume both times confirms your unit handling.

Common Mistake: Substituting temperature in degrees Celsius into a gas law. It must be converted to kelvin first.

Important Questions:

• State Boyle’s law and Charles’s law. Boyle’s law says the volume of a fixed amount of gas is inversely proportional to its pressure at constant temperature; Charles’s law says the volume is directly proportional to the absolute temperature at constant pressure.
• Why do real gases deviate from ideal behaviour at high pressure and low temperature? At those conditions the finite volume of the molecules and the attractive forces between them can no longer be ignored, which is what the van der Waals equation corrects.

Chapter 9: Thermochemistry

Difficulty: Medium · Key topics: energy, heat, specific heat, calorimetry, enthalpy, Hess’s law, enthalpy of formation, bond energies, lattice energy, Born-Haber cycle

This chapter measures and calculates the heat of chemical change. It defines kinetic and potential energy, heat versus temperature, endothermic and exothermic processes, and specific heat. Calorimetry then shows how coffee-cup and bomb calorimeters turn temperature changes into heat. The enthalpy section covers thermochemical equations, standard enthalpies of combustion and formation, and Hess’s law. The last section, placed here instead of in the bonding chapter, deals with the strengths of covalent bonds (bond energies) and ionic bonds (lattice energy and the Born-Haber cycle).

Key Points:

• Heat is energy transferred between objects at different temperatures; the SI unit of energy is the joule
• Heat absorbed or released: q = m × c × ΔT, where c is the specific heat
• ΔH is the heat change at constant pressure; reversing an equation flips its sign and scaling the equation scales ΔH
• Hess’s law: if a process is the sum of steps, its enthalpy change is the sum of the step enthalpies
• ΔH° = Σ ΔH°_f(products) − Σ ΔH°_f(reactants)
• Bond dissociation energy measures covalent bond strength; lattice energy measures ionic bond strength and rises with higher charge and shorter ion separation

Practice Tip: In calorimetry problems, write the energy balance first: heat lost by the hot object equals heat gained by the cold one, so q_metal = −q_water. The algebra then follows without sign confusion.

Common Mistake: Forgetting that ΔH changes sign when a reaction is reversed and scales with the coefficients when the equation is multiplied.

Important Questions:

• State Hess’s law. If a reaction can be written as the sum of several steps, the enthalpy change of the overall reaction equals the sum of the enthalpy changes of the individual steps.
• How much heat is needed to warm 100 g of water from 20 °C to 50 °C (c = 4.184 J/g·°C)? q = 100 g × 4.184 J/g·°C × 30 °C ≈ 1.26 × 10⁴ J.

Chapter 10: Liquids and Solids

Difficulty: Medium · Key topics: intermolecular forces, hydrogen bonding, viscosity and surface tension, phase transitions, heating curves, phase diagrams, crystalline solids, unit cells, X-ray diffraction

Using the kinetic-molecular ideas from the gas chapter, this chapter explains condensed matter through the attractions between particles. London dispersion forces, dipole-dipole attractions and hydrogen bonding are compared, and then used to explain viscosity, surface tension and capillary rise. Phase transitions are treated with vapor pressure, boiling points, enthalpies of vaporization and fusion, and heating curves. Phase diagrams introduce the triple point, the critical point and supercritical fluids. The solid-state sections classify ionic, metallic, covalent network and molecular solids, then describe unit cells, closest packing, holes in ionic lattices and X-ray diffraction.

Key Points:

• London forces exist in all matter and grow with molecular size; dipole-dipole forces act between polar molecules
• Hydrogen bonding is a strong dipole-dipole attraction that occurs when H is bonded to F, O or N
• Melting, vaporization and sublimation are endothermic; freezing, condensation and deposition are exothermic
• The triple point is where solid, liquid and gas coexist; above the critical point no distinct liquid phase exists
• Crystalline solids fall into four types: ionic, metallic, covalent network and molecular
• Small cations usually occupy tetrahedral holes and larger cations octahedral holes in a close-packed anion array

Memory Tip: When ranking boiling points of molecules of similar size, ask three questions in order: is there an H on F, O or N (hydrogen bonding), is the molecule polar (dipole-dipole), and how big is it (dispersion)?

Common Mistake: Treating a hydrogen bond as an ordinary chemical bond inside the molecule. It is an attraction between molecules, far weaker than the covalent bonds it connects.

Important Questions:

• Why does water boil at a much higher temperature than H₂S? Water molecules form hydrogen bonds with each other, which are much stronger than the dipole-dipole and dispersion forces in H₂S, so more energy is needed to separate them.
• What are the triple point and the critical point on a phase diagram? The triple point is the one temperature and pressure at which solid, liquid and gas are in equilibrium; the critical point is the end of the liquid-gas curve, beyond which liquid and gas cannot be distinguished.

Chapter 11: Solutions and Colloids

Difficulty: Medium · Key topics: dissolution, electrolytes, solubility, Henry’s law, molality, colligative properties, osmotic pressure, van’t Hoff factor, colloids

This chapter explains why substances dissolve and what a solution does to the properties of its solvent. It looks at the energy changes of dissolution, then at electrolytes, where ion-dipole attractions pull ionic compounds apart in water. The solubility section covers saturated and supersaturated solutions, miscibility, the effect of temperature, and Henry’s law for gases. The longest section covers colligative properties: molality and mole fraction, vapor pressure lowering, boiling point elevation, freezing point depression, osmosis, and the van’t Hoff factor for electrolytes. A final section describes colloids and how they differ from true solutions.

Key Points:

• Strong electrolytes ionize completely, weak electrolytes only slightly, and nonelectrolytes not at all
• Henry’s law: the concentration of a dissolved gas is proportional to its partial pressure above the solution
• Gas solubility falls as temperature rises; for most solids it rises
• Molality = moles of solute ÷ kilograms of solvent
• Colligative properties depend only on the concentration of solute particles; the van’t Hoff factor i corrects for ionization
• Colloid particles stay dispersed and do not settle under gravity, and they are often electrically charged

Practice Tip: Compare the predicted freezing points of 1.00 m glucose and 1.00 m NaCl in water. Seeing that NaCl lowers the freezing point roughly twice as much makes the idea of counting particles concrete.

Common Mistake: Assuming equal molal solutions of different solutes change a property equally. Colligative effects depend on the number of dissolved particles, so an electrolyte produces a larger effect than a nonelectrolyte.

Important Questions:

• What does a colligative property depend on? Only on the concentration of solute particles, not on their identity; vapor pressure lowering, boiling point elevation, freezing point depression and osmotic pressure are all examples.
• Why do lakes hold less dissolved oxygen in summer? The solubility of gases in water decreases as temperature increases, so warm water holds less oxygen.

Chapter 12: Thermodynamics

Difficulty: Hard · Key topics: spontaneity, dispersal of energy and matter, entropy, microstates, second and third laws, standard entropy change, Gibbs free energy, coupled reactions

This chapter asks which processes happen on their own. It separates spontaneous from nonspontaneous processes and shows that spontaneity is linked to how matter and energy are dispersed. Entropy is introduced through microstates (Boltzmann’s statistical view) and through heat divided by temperature, with rules for predicting the sign of the entropy change. The second law (the entropy of the universe increases in a spontaneous process) and the third law (zero entropy for a perfect crystal at 0 K) let you calculate standard entropy changes. The final section defines Gibbs free energy, ways to calculate ΔG, and coupled reactions.

Key Points:

• A spontaneous process proceeds without continual outside energy input; a nonspontaneous process requires it
• Entropy S counts the ways a system can be arranged (microstates): S = k ln W
• Entropy order: S_solid < S_liquid < S_gas; it also increases on heating and when solutions form
• Second law: ΔS_univ > 0 for a spontaneous process; third law: a perfect crystal at 0 K has zero entropy
• Gibbs free energy uses only system quantities: ΔG = ΔH − TΔS
• ΔG < 0 means spontaneous, ΔG > 0 nonspontaneous, and ΔG = 0 means equilibrium

Memory Tip: Think of ΔG = ΔH − TΔS as a tug of war. A negative ΔH and a positive ΔS both push ΔG negative, so if both are favourable the process is spontaneous at every temperature.

Common Mistake: Treating spontaneous as meaning fast. Spontaneity only says a process is thermodynamically favourable; it says nothing about the rate, which is the subject of kinetics.

Important Questions:

• State the second law of thermodynamics. For any spontaneous process the entropy of the universe increases (ΔS_univ > 0); if it is zero the system is at equilibrium.
• Predict the sign of ΔS for 2SO₂(g) + O₂(g) → 2SO₃(g). Three moles of gas become two moles of gas, so there are fewer ways to arrange the system and ΔS is negative.

Chapter 13: Fundamental Equilibrium Concepts

Difficulty: Hard · Key topics: reversible reactions, dynamic equilibrium, reaction quotient, equilibrium constant, Le Châtelier’s principle, ICE tables, free energy and equilibrium

Here the book turns from “will it happen” to “how far does it go”. The chapter defines reversible reactions and dynamic equilibrium, then the reaction quotient Q and the equilibrium constant K, with K_c and K_P, homogeneous and heterogeneous equilibria, and coupled equilibria. Le Châtelier’s principle predicts how changes in concentration, temperature, volume and pressure shift an equilibrium. The longest section teaches equilibrium calculations with ICE tables, including the algebra-simplifying assumption, and returns to free energy: how ΔG depends on temperature and on Q, and how ΔG° relates to K.

Key Points:

• At equilibrium the forward and reverse rates are equal and amounts stay constant, but the reactions have not stopped
• Q uses current concentrations and K uses equilibrium concentrations; if Q < K the reaction shifts toward products
• Pure solids and pure liquids are left out of the equilibrium expression
• Le Châtelier: a system disturbed from equilibrium shifts in the direction that counters the disturbance
• A catalyst speeds up both directions equally and does not change K or the equilibrium mixture
• ΔG = ΔG° + RT ln Q, so at equilibrium ΔG = 0 and ΔG° = −RT ln K

Practice Tip: Write the full ICE table (initial, change, equilibrium) for every equilibrium problem before touching the algebra, and finish by substituting your answers back into the K expression to check them.

Common Mistake: Using the “x is small” shortcut without checking afterwards that x really is negligible compared with the starting concentration.

Important Questions:

• What does Q < K tell you about a reaction mixture? The mixture has too much reactant relative to equilibrium, so the net reaction proceeds toward products until Q equals K.
• What effect does a catalyst have on an equilibrium? It lets equilibrium be reached faster but does not change the equilibrium constant or the final composition.

Chapter 14: Acid-Base Equilibria

Difficulty: Hard · Key topics: Brønsted-Lowry acids and bases, conjugate pairs, K_w, pH and pOH, K_a and K_b, salt hydrolysis, polyprotic acids, buffers, titration curves

This chapter applies equilibrium ideas to acids and bases in water. It starts with the Brønsted-Lowry definition, conjugate acid-base pairs, amphiprotic species and the autoionization of water (K_w). pH and pOH follow, then the ionization constants K_a and K_b, percent ionization, and how molecular structure affects acid strength. Hydrolysis explains why many salt solutions are acidic or basic, polyprotic acids ionize in steps, and buffers resist pH change, with the Henderson-Hasselbalch equation. The chapter ends with titration curves and the choice of acid-base indicators.

Key Points:

• A Brønsted-Lowry acid donates a proton and a base accepts one; each acid has a conjugate base and each base a conjugate acid
• K_w = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C, so pH + pOH = 14
• For a conjugate pair K_a × K_b = K_w; the stronger the acid, the weaker its conjugate base
• Salts of a weak base’s cation give acidic solutions; salts of a weak acid’s anion give basic solutions
• Each successive ionization of a polyprotic acid has a smaller K_a
• Buffers: pH = pK_a + log([base] ÷ [acid]); a buffer fails when its capacity is exceeded

Memory Tip: For salts, the stronger parent wins. A strong acid with a weak base gives an acidic salt, a strong base with a weak acid gives a basic salt, and two strong parents give a neutral salt.

Common Mistake: Reading pH as a linear scale. Each pH unit is a factor of ten in hydronium concentration, so pH 3 is ten times more acidic than pH 4.

Important Questions:

• What is the pH of 0.010 M HCl? HCl is a strong acid, so [H₃O⁺] = 0.010 M and pH = −log(0.010) = 2.00.
• What is a buffer and how does it work? A buffer is a solution containing appreciable amounts of a weak conjugate acid-base pair; the base neutralizes added acid and the acid neutralizes added base, so the pH changes only slightly.

Chapter 15: Equilibria of Other Reaction Classes

Difficulty: Hard · Key topics: solubility product, molar solubility, precipitation, common ion effect, Lewis acids and bases, complex ions, formation constants, coupled equilibria

This chapter extends the equilibrium toolkit beyond acids and bases. The first section introduces the solubility product K_sp, shows how to move between K_sp and molar solubility, predicts when a precipitate forms by comparing Q with K_sp, and explains the common ion effect, with examples from medical imaging and wastewater treatment. The Lewis acid-base section generalizes acids and bases to electron-pair acceptors and donors and applies it to complex ions and their formation constants. The last section combines several equilibria at once, such as dissolution with acid-base or complex-formation reactions.

Key Points:

• K_sp is the equilibrium constant for dissolution of a slightly soluble ionic solid, written with ion concentrations only
• Precipitation begins when Q exceeds K_sp and continues until Q equals K_sp
• The common ion effect lowers the solubility of a salt when one of its ions is already present in solution
• A Lewis acid accepts an electron pair; a Lewis base donates one
• A complex ion is a central metal ion bonded to ligands; its formation constant K_f is large when the complex is stable
• Coupled equilibria explain why some salts dissolve better in acid or in a solution containing ligands

Practice Tip: For every solubility problem, write the dissolution equation and the K_sp expression first, then set up an ICE table using the molar solubility s. Doing it in this order prevents most stoichiometry slips.

Common Mistake: Ranking the solubility of two salts by comparing their K_sp values directly. This only works when the formulas have the same ion ratio; otherwise compare calculated molar solubilities.

Important Questions:

• What does it mean when Q is greater than K_sp? The solution is supersaturated with respect to the solid, so a precipitate forms until the ion concentrations fall to the point where Q equals K_sp.
• Why is a fluoride-containing solid such as CaF₂ more soluble in acidic solution? Fluoride is a weak base; H⁺ converts F⁻ to HF, which removes fluoride from the solution and pulls the dissolution equilibrium forward.

Chapter 16: Electrochemistry

Difficulty: Medium · Key topics: redox review, half-reactions, galvanic cells, standard electrode potentials, Nernst equation, batteries, fuel cells, corrosion, electrolysis

Electrochemistry links redox chemistry to energy and equilibrium. The chapter reviews oxidation numbers and balancing redox equations in acidic and basic solution, then builds galvanic cells with separate half-cells, a salt bridge and cell notation. Standard electrode potentials, measured against the standard hydrogen electrode, let you compute cell potentials and predict spontaneity. The relationships between E°, ΔG° and K and the Nernst equation for nonstandard conditions follow. Applied sections cover batteries and fuel cells, corrosion and its prevention, and electrolysis with its stoichiometry of current and time.

Key Points:

• Oxidation (loss of electrons) occurs at the anode and reduction (gain of electrons) at the cathode
• A salt bridge lets ions move between half-cells to keep each side electrically neutral
• The standard hydrogen electrode is defined as exactly 0 V; E°_cell = E°_cathode − E°_anode
• ΔG° = −nFE°, and a positive E° means a spontaneous cell reaction
• The Nernst equation gives the cell potential when concentrations are not standard
• Electrolysis forces a nonspontaneous reaction with an external source; moles of electrons = current × time ÷ Faraday constant

Memory Tip: “An Ox and a Red Cat”: oxidation happens at the anode, and reduction happens at the cathode. This holds for both galvanic and electrolytic cells.

Common Mistake: Multiplying a half-reaction’s potential when the half-reaction is multiplied to balance electrons. Potential is an intensive property, so the E° value stays the same.

Important Questions:

• How are the anode and cathode identified in a galvanic cell? The anode is where oxidation occurs and is the source of electrons in the external circuit; the cathode is where reduction occurs and receives the electrons.
• How is the standard cell potential related to free energy? ΔG° = −nFE°, so a positive standard cell potential corresponds to a negative ΔG° and a spontaneous reaction.

Chapter 17: Kinetics

Difficulty: Medium · Key topics: reaction rates, factors affecting rates, rate laws, reaction order, integrated rate laws, half-life, collision theory, Arrhenius equation, mechanisms, catalysis

Kinetics asks how fast reactions happen. The chapter defines reaction rate as the change in amount per unit time and relates the rates of different species through stoichiometry. It lists the factors that change rates: the nature of the reactants, surface area, temperature, concentration and catalysts. Rate laws and reaction order are determined from experimental data, and integrated rate laws give concentration versus time and half-lives. Collision theory and the Arrhenius equation explain the effect of temperature and activation energy. The chapter ends with reaction mechanisms, the rate-determining step, and homogeneous, heterogeneous and enzyme catalysis.

Key Points:

• Rate laws are found by experiment and cannot be read from the balanced equation
• The overall reaction order is the sum of the orders for each reactant; the units of k depend on the order
• First order: ln[A] falls linearly with time and t_1/2 = 0.693 ÷ k, independent of concentration
• Second order: 1/[A] rises linearly with time; zero order: [A] falls linearly
• Collision theory: reaction needs collisions with enough energy and proper orientation; the Arrhenius equation links k, E_a and temperature
• The slowest elementary step is rate-determining; a catalyst lowers activation energy but does not change equilibrium

Practice Tip: Given concentration versus time data, plot [A], ln[A] and 1/[A] against time. The plot that gives a straight line tells you the reaction order, and its slope gives k.

Common Mistake: Taking the exponents in the rate law from the coefficients of the balanced equation. Only elementary steps work that way.

Important Questions:

• How is the order of a reaction determined? Experimentally, for example by comparing initial rates when concentrations are changed one at a time, or by seeing which plot of concentration against time is linear.
• A first-order reaction has k = 0.0231 min⁻¹. What is its half-life? t_1/2 = 0.693 ÷ 0.0231 min⁻¹ ≈ 30.0 minutes.

Chapter 18: Representative Metals, Metalloids, and Nonmetals

Difficulty: Medium · Key topics: periodicity, alkali and alkaline earth metals, metal preparation, metalloids, nonmetals, hydrogen, carbonates, nitrogen, phosphorus, oxygen, sulfur, halogens, noble gases

This is the book’s descriptive-chemistry tour of the s- and p-block elements. It begins with periodic trends in the properties and reactions of the representative metals, then explains how reactive metals are isolated by electrolysis (sodium, potassium, aluminum) or chemical reduction (magnesium, zinc, tin). Sections on the six metalloids and on the general structure of nonmetals follow. The rest of the chapter works through hydrogen, carbonates, nitrogen, phosphorus, oxygen, sulfur, the halogens and the noble gases, covering occurrence, preparation, typical compounds and uses.

Key Points:

• Alkali metals form 1+ ions and react vigorously with water to give hydrogen and a basic hydroxide solution; group 2 metals form 2+ ions and are less reactive
• Sodium, potassium and aluminum are made by electrolysis; magnesium, zinc and tin mainly by chemical reduction
• The metalloids (B, Si, Ge, As, Sb, Te) are semiconductors with properties between metals and nonmetals
• Most nonmetal oxides are acid anhydrides, and most metal oxides are base anhydrides
• The very strong N≡N triple bond makes nitrogen gas unreactive, so nitrogen fixation takes a lot of energy
• Xenon forms fluorides, showing the noble gases are not completely inert (Bartlett, 1962)

Memory Tip: Two sentences explain a large part of this chapter: metal oxides are basic (they react with acids), and nonmetal oxides are acidic (they react with water to form acids).

Common Mistake: Memorizing each preparation as an isolated fact. Group them by method instead: electrolysis for the most reactive metals, chemical reduction for the rest.

Important Questions:

• Why are alkali metals never found as free elements, and how are they produced? They are so reactive that they occur only in compounds, so the free metals (for example sodium and potassium) are prepared by electrolysis.
• Why is nitrogen gas so unreactive? The N≡N triple bond is very strong, so a large amount of energy is needed to break it before nitrogen can form other compounds.

Chapter 19: Transition Metals and Coordination Chemistry

Difficulty: Hard · Key topics: d-block elements, oxidation states, metallurgy, coordination compounds, ligands, geometries, isomerism, crystal field theory, high-spin and low-spin complexes, colour

The chapter first describes transition metals as elements with partially filled d orbitals, comparing their reactivity, common oxidation states, how iron, copper and silver are isolated from their ores, and the behaviour of their halides, oxides, hydroxides and carbonates. It then moves to coordination compounds: ligands, coordination numbers, the common tetrahedral, square planar and octahedral shapes, naming, and geometric and optical isomerism. The final section introduces crystal field theory to explain the colour and magnetism of complexes through splitting of the d orbitals.

Key Points:

• Transition metals have partially filled d orbitals and usually show several stable oxidation states
• Ligands donate electron pairs to the central metal ion through coordinate covalent bonds
• Polydentate ligands bind through more than one donor atom and form chelates
• Common geometries: tetrahedral and square planar (coordination number 4), octahedral (coordination number 6)
• Isomers include cis/trans geometric isomers and optical isomers that are non-superimposable mirror images
• Crystal field splitting (Δ_oct): strong-field ligands give low-spin complexes, weak-field ligands give high-spin complexes

Practice Tip: Sketch the d-orbital splitting diagram for an octahedral d⁶ ion, once with a weak-field and once with a strong-field ligand, and count the unpaired electrons in each. The difference is the whole idea of high-spin versus low-spin.

Common Mistake: Assuming coordination number equals the number of ligands. A polydentate ligand such as ethylenediamine occupies two coordination sites by itself.

Important Questions:

• What are ligands and what is a coordination number? Ligands are molecules or ions that donate electron pairs to a central metal atom or ion; the coordination number is the number of donor atoms attached to that metal.
• How do high-spin and low-spin octahedral complexes differ? In a low-spin complex the t_2g orbitals fill completely before any electron enters the e_g orbitals (large splitting from strong-field ligands); in a high-spin complex the orbitals are singly occupied first (small splitting from weak-field ligands).

Chapter 20: Nuclear Chemistry

Difficulty: Medium · Key topics: nuclear binding energy, nuclear stability, nuclear equations, radioactive decay modes, half-life, radiometric dating, fission and fusion, radioisotope uses, biological effects of radiation

Unlike every earlier chapter, this one changes the nucleus rather than the electrons. It explains nuclear structure, the strong nuclear force, mass defect and binding energy (E = mc²), and why only some nuclides are stable. Nuclear equations are balanced for mass number and charge using alpha, beta, positron and gamma emission and electron capture. The decay section covers decay series, first-order decay kinetics, half-lives and radiocarbon dating. Further sections describe transmutation, fission reactors and fusion, uses of radioisotopes in medicine and industry, and the biological effects, detection and units of radiation.

Key Points:

• Nuclei contain protons and neutrons (nucleons) held together by the short-range strong nuclear force
• Mass defect is the “missing” mass that becomes binding energy; binding energy per nucleon is highest near mass number 56
• Nuclear equations balance both mass number and charge on the two sides
• Radioactive decay is first order and each radioisotope has its own half-life
• Fission splits a heavy nucleus and can sustain a chain reaction; fusion joins light nuclei and releases energy but is not yet commercial
• Alpha particles penetrate least, gamma rays most; shielding, distance and short exposure limit radiation dose

Practice Tip: After writing any nuclear equation, add the mass numbers on each side and then the atomic numbers on each side. Both sums must match, and this check catches almost every error.

Common Mistake: Assuming alpha radiation is harmless because it is the least penetrating. It is stopped by skin, but an alpha emitter that is inhaled or swallowed can be highly damaging to tissue.

Important Questions:

• What are the products of alpha decay of uranium-238? An alpha particle (a helium-4 nucleus) and thorium-234, because the mass number falls by 4 and the atomic number by 2 (92 to 90).
• What is mass defect and why does it matter? A nucleus has slightly less mass than its separate nucleons; the difference is converted to the binding energy that holds the nucleus together, according to E = mc².

Chapter 21: Organic Chemistry

Difficulty: Medium · Key topics: alkanes, alkenes, alkynes, aromatic hydrocarbons, skeletal structures, nomenclature, alcohols, ethers, aldehydes, ketones, carboxylic acids, esters, amines, amides

The book closes with a first look at the chemistry of carbon. The hydrocarbon section covers alkanes, alkenes, alkynes and aromatic compounds, how to read and draw skeletal structures, isomers, and the basic rules for naming alkanes and alkenes. Later sections show how hydrocarbon derivatives are defined by functional groups: the alcohol and ether groups, the carbonyl-based families (aldehydes, ketones, carboxylic acids and esters), and amines and amides, including the condensation of an amine with a carboxylic acid. Short features link these ideas to plastics recycling, carbohydrates, DNA and proteins.

Key Points:

• Alkanes have only single bonds, alkenes at least one C=C, alkynes at least one C≡C, and aromatic hydrocarbons have rings with delocalized π electrons
• A functional group is an atom or group of atoms that largely determines a compound’s properties
• Alcohols contain –OH and ethers contain R–O–R
• The carbonyl group C=O is the key unit of aldehydes (–CHO), ketones, carboxylic acids (–CO₂H) and esters (–CO₂R)
• Amines are basic and contain nitrogen bonded to carbon; amides have nitrogen on one side of a carbonyl
• Isomers have the same molecular formula but different arrangements of atoms

Practice Tip: Draw every isomer of C₄H₁₀ and C₅H₁₂ (two and three respectively) and name each one. Doing this teaches the naming rules faster than reading them.

Common Mistake: Forgetting that every unlabelled line end or corner in a skeletal structure is a carbon atom carrying enough hydrogens to make four bonds.

Important Questions:

• How do alkanes, alkenes, alkynes and aromatic hydrocarbons differ? Alkanes contain only single carbon-carbon bonds, alkenes at least one double bond, alkynes at least one triple bond, and aromatic hydrocarbons contain ring structures with delocalized π electrons.
• How does an aldehyde differ from a ketone? Both contain a carbonyl group; in an aldehyde at least one hydrogen is attached to the carbonyl carbon, while in a ketone the carbonyl carbon is bonded to two carbon groups.

Download Chemistry: Atoms First 2e PDF (Free)

This book is free from its official source, OpenStax. The button below opens the official OpenStax page, where you can download the complete PDF or read every chapter online in the web edition.

How to Study This Book

The first thing to know is that this book is a reordering, not a different course. Chemistry: Atoms First 2e contains the same general chemistry material as Chemistry 2e, arranged so atomic and molecular structure comes first. Chapters 2 to 5 cover atoms, electron configurations, bonding and molecular orbitals before any stoichiometry. Mole calculations and balanced equations only arrive in Chapters 6 and 7, and the book itself says this delay is meant to give you time to settle into chemistry study habits before the heavy calculation chapters.

Read Chapters 1 to 5 in order, and treat Chapter 3 as the foundation for everything after it. Quantum numbers, electron configurations and periodic trends are assumed without re-explanation in the bonding chapters (4 and 5) and in the descriptive chemistry of Chapters 18 and 19. Chapter 5 depends on Chapter 4: hybridization and molecular orbitals are extensions of Lewis structures and VSEPR, not replacements for them.

Chapters 6 and 7 are where the quantitative work starts, and they carry the rest of the book. Mole conversions, molarity, limiting reactants and percent yield are used from Chapter 8 onward. If your course is calculation-heavy, spend most of your practice time here and in Chapter 8 (Gases) before moving on.

The order from Chapter 9 to Chapter 15 is worth following as written. Thermochemistry (9) leads to thermodynamics (12), which comes before equilibrium (13) in this book. Chapter 13 is load-bearing: acid-base equilibria (14) and solubility and complex-ion equilibria (15) both use its ICE-table method, and Chapter 16 uses its link between free energy and K. Chapter 17 (Kinetics) comes late, after electrochemistry. Its calculations rely mainly on the concentration work of Chapters 6 and 7, so some courses teach it earlier; just note that its Catalysis section refers back to equilibrium.

Chapters 10 and 11 (liquids, solids and solutions) are less tightly chained to their neighbours. They build on Chapter 4 for polarity, on Chapter 8 for the kinetic-molecular picture, and Chapter 11 draws on Chapter 6 for concentration units.

Chapters 18 to 21 are descriptive and survey chapters that many one-year courses cover selectively. Chapter 18 in particular is a long catalogue of elements; read it for patterns (basic metal oxides, acidic nonmetal oxides, preparation by electrolysis or reduction) rather than trying to memorize every reaction. Chapter 20 (Nuclear) is relatively self-contained, and Chapter 21 (Organic) is only an introduction and does not replace a full organic chemistry textbook.

The book has more than 1,700 end-of-chapter exercises, about 290 worked examples with Check Your Learning problems, and 13 appendices (A to M) of reference tables: periodic table, essential mathematics, units, constants, thermodynamic data, ionization constants, solubility products, formation constants, standard electrode potentials and half-lives. Appendix B (Essential Mathematics) is worth reading if exponents, logarithms or quadratic equations feel rusty. Odd-numbered answers are provided to students through OpenStax; even-numbered answers are available only to instructors.

Used In These Programs

This book is used for the General Chemistry course in: BS Chemistry, and in Pre-Medical, Engineering and other STEM programs that require general chemistry. Browse all General Chemistry books, all Chemistry books, or the full Chemistry category.

Who Should Read This

This book is written for first-year university students taking general chemistry, and it works for a two-semester or a three-semester sequence. It suits BS Chemistry students, pre-medical students and engineers who want atomic and molecular structure explained before stoichiometry, and anyone who plans to continue into organic or inorganic chemistry, since that path starts from atoms and bonding. No previous college chemistry is assumed, but comfort with algebra, logarithms and scientific notation is. It is not a calculus-based text. If your course follows the traditional order, with measurement and stoichiometry first, read OpenStax Chemistry 2e instead; it covers the same material in a different sequence, and it is a better match for a syllabus that is organized that way. Pick one of the two and follow it, rather than mixing them. If you want a quick revision of the basics before starting, the Chemistry 1st Year (FSC) and Chemistry Class 9 books in Related Books below cover the school-level foundations.

Applicable Universities

This book is useful for students at Pakistani universities offering BS Chemistry, pre-medical, engineering and other science programs, including Punjab University, Virtual University, COMSATS, UET, NUST, GIKI and other HEC-recognized institutions, wherever general chemistry is a first-year requirement.

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Related Books

• Chemistry 1st Year Book (Punjab Board FSC) – the school-level foundation for atoms, the periodic table and bonding
• Chemistry 2nd Year FSC Book (Punjab Board) – a college-level companion for equilibrium, acids and bases, and organic basics
• Chemistry Class 9 English Medium (Punjab Board PCTB) – quick revision if the basics feel unfamiliar
• Calculus Volume 1 – Strang & Herman (OpenStax) – not needed for this chemistry book, but useful if your degree adds calculus alongside general chemistry

Chemistry: Atoms First 2e gives BS Chemistry, pre-medical and engineering students a free, complete general chemistry course in the atoms-first order. Browse more Chemistry books to find the rest of your semester.

Chemistry: Atoms First 2e, by Paul Flowers, Klaus Theopold, Richard Langley, Edward J. Neth, William R. Robinson, and contributors. OpenStax, Rice University, in partnership with the University of Connecticut. Free under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International licence. Access for free at https://openstax.org/details/books/chemistry-atoms-first-2e