⚗️ General Chemistry II

What you’ll learn

1. The Three QuestionsDistinguish the three independent questions — will it go, how fast, how far — that determine whether a balanced equation describes anything that actually happens.A balanced equation is a claim about bookkeeping, not about reality: it says a reaction could happen without saying it will. Diamond is thermodynamically unstable relative to graphite (ΔG° = −2.9 kJ/mol) yet never converts, which proves that favorability and speed are separate questions. General Chemistry II is the study of the three questions — direction, rate, and extent — that decide a reaction's fate.
2. The Energy LedgerExplain enthalpy change as the net result of a bond-breaking/bond-forming ledger, and use Hess's law to obtain ΔH for a reaction that cannot be measured directly.Breaking bonds always costs energy and forming bonds always releases it; a reaction's enthalpy change is just the difference between those two totals, which is why exothermic reactions are the ones that end in stronger bonds. Because enthalpy is a state function — dependent only on start and end points, not on route — Hess's law lets chemists add known reactions together to get the enthalpy of a reaction no one can measure directly.
3. The Arrow That Isn't EnergyDefine entropy as the number of ways a system's energy and particles can be arranged, and explain why spontaneous change tends toward spreading rather than toward disorder.Entropy counts the microscopic arrangements consistent with a state, so change runs toward whatever can be achieved in the overwhelmingly larger number of ways — not because nature prefers mess, but because there is vastly more of it. The second law says the entropy of the universe increases, which is why a spontaneous process can be locally ordering (water freezing) as long as it dumps more entropy into its surroundings than it removes from itself.
4. The VerdictUse ΔG = ΔH − TΔS to predict spontaneity, and explain why temperature can reverse the verdict for reactions whose enthalpy and entropy terms disagree.Gibbs free energy combines the energy vote (ΔH) and the entropy vote (TΔS) into a single criterion: ΔG < 0 means spontaneous. Because the entropy term is multiplied by temperature, the two votes' relative weight shifts with T — which is why some reactions flip direction at a specific temperature, calculable as the point where ΔH = TΔS.
5. Spontaneous Is Not FastExplain activation energy as a barrier that must be climbed before a favorable reaction can proceed, and use it to reconcile thermodynamic favorability with observed rate.Reacting molecules must first pass through a high-energy transition state in which old bonds are partly broken and new ones only partly formed, and the cost of reaching it is the activation energy. Because that barrier is unrelated to the reaction's overall ΔG, a strongly favorable reaction can be immeasurably slow — which is why fuel, and diamond, and you, persist at all.
6. How Fast, ExactlyInterpret a rate law, and explain why reaction orders must be determined experimentally rather than read from the balanced equation.A rate law states how concentration controls speed, and its exponents — the reaction orders — are experimental findings, not the coefficients from the balanced equation. The reason is that most reactions proceed through several elementary steps, and the rate is set by the slowest of them, so the rate law reports the mechanism's bottleneck rather than the overall stoichiometry.
7. Why Heat Works So WellUse collision theory and the distribution of molecular energies to explain why a modest temperature rise produces a large rate increase.Reactions require collisions that are both hard enough to clear the activation barrier and correctly oriented, and only a small tail of molecules at any temperature qualifies. Because raising the temperature shifts the whole energy distribution, it multiplies the size of that tail rather than adding to it — which is why roughly 10 °C can double a rate, and why refrigeration works.
8. A Better RoadExplain how a catalyst accelerates a reaction by providing an alternative mechanism with a lower activation energy, and why it cannot change the reaction's equilibrium position or ΔG.A catalyst does not push a reaction; it opens a different route with a lower barrier, then emerges unchanged to do it again. Because it alters only the path and not the endpoints, it cannot change ΔG or the equilibrium position — it speeds the forward and reverse reactions equally, which is precisely why it is useful and why it is not a source of free lunch.
9. Reactions That Stop ShortExplain dynamic equilibrium as the point where forward and reverse rates become equal, and interpret the magnitude of K as a statement about the position of that balance.Reactions are two-way, so as products accumulate the reverse reaction accelerates until it matches the forward reaction and net change stops — with reactants still present. Equilibrium is dynamic rather than finished, and the equilibrium constant K reports where the balance sits: large K means product-favored, small K means the reaction barely proceeds.
10. Pushing Back on EquilibriumApply Le Chatelier's principle to predict how concentration, pressure and temperature changes shift an equilibrium, and explain why only temperature changes the value of K.A system at equilibrium responds to an imposed stress by shifting in the direction that partly relieves it, which lets chemists move a balance point without changing the chemistry. Concentration and pressure shift the position while leaving K fixed; temperature is unique in actually changing K, because heat is effectively a participant in the reaction.
11. The CompromiseExplain the Haber–Bosch process's operating conditions as a deliberate compromise between thermodynamics, kinetics and equilibrium, and evaluate the trade-offs each condition represents.Ammonia synthesis is exothermic and reduces the number of gas moles, so equilibrium wants it cold and compressed while kinetics wants it hot — a direct conflict that no single condition resolves. Industry settles on roughly 400–450 °C and 200 atm with an iron catalyst, converting only ~15% per pass and recycling the rest, which turns a mediocre equilibrium into a process that fixes nitrogen at global scale.
12. How Strong Is a Weak AcidExplain acid strength as an equilibrium position quantified by Ka, and distinguish acid strength from concentration.Strong acids ionize essentially completely while weak acids reach an equilibrium in which most molecules stay intact, and Ka reports where that balance sits. Strength and concentration are independent: a concentrated weak acid can have a higher pH than a dilute strong one, because pH counts the H⁺ actually released rather than the acid present.
13. Chemistry That Pushes BackExplain how a buffer resists pH change using a weak acid/conjugate base pair, and why buffering is only possible with a weak acid.A buffer holds a reservoir of both a weak acid and its conjugate base, so added acid is absorbed by the base and added base is neutralized by the acid — Le Chatelier's principle deployed on purpose. It works only for weak acids, because a strong acid keeps no undissociated reservoir to draw on, and it is what holds blood at pH 7.4 against a continuous metabolic acid load.
14. Making a Reaction Do WorkExplain how separating a redox reaction's half-reactions converts chemical free energy into electrical work, and connect cell voltage to ΔG.A redox reaction moves electrons from one substance to another, and physically separating the two halves forces those electrons through a wire, where they can do work instead of releasing heat. Cell voltage is free energy per unit charge (ΔG = −nFE), which makes a voltmeter a direct instrument for reading the thermodynamics from the first half of this course.

Questions this course answers

Grounded in trusted sources

• OpenStax — Chemistry 2e (2019), chs. 5, 12, 13, 14, 16, 17
• Brown, LeMay & Bursten — Chemistry: The Central Science, 14th ed. (2018)
• Chemistry LibreTexts — The Haber Process (ΔH = −92.4 kJ/mol; 400–450 °C; 200 atm; ~15% conversion per pass, ~98% with recycling)
• Vaclav Smil — Enriching the Earth: Fritz Haber, Carl Bosch and the Transformation of World Food Production (2001)
• Wikipedia — 'pH' (measurement table; blood 7.34–7.45; gastric acid 1.0–1.5; ocean surface pH ~8.15 → ~8.05, 1950–2020)
• Wikipedia — 'Daniell cell' (E° = 1.10 V; Zn²⁺/Zn = −0.76 V, Cu²⁺/Cu = +0.34 V)
• Lumen Learning — Acid and Base Ionization Constants (acetic acid Ka = 1.8 × 10⁻⁵; 0.10 M is ~1.3% ionized)
• Standard Gibbs energy of formation, C(diamond) = +2.9 kJ/mol relative to graphite — CBSE/Vedantu thermodynamic tables

Every Wunder lesson is built from real, reputable sources — never invented.

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