Key Thermodynamics Principles

Why This Matters

Thermodynamics is the backbone of chemistry—it tells you why reactions happen, not just what happens. Every time you're asked to predict whether a reaction will occur spontaneously, calculate heat flow, or explain why ice melts at room temperature, you're applying thermodynamic principles. These concepts connect directly to equilibrium, electrochemistry, and kinetics, making them foundational for everything that follows in Honors Chemistry.

You're being tested on your ability to distinguish between enthalpy and entropy, predict spontaneity using Gibbs Free Energy, and apply conservation of energy in calculations. The exam loves to ask you to compare exothermic versus endothermic processes or determine when temperature affects spontaneity. Don't just memorize definitions—know what concept each principle illustrates and how they work together to explain real chemical behavior.

---

The Fundamental Laws: Rules Energy Must Follow

These three laws establish the boundaries of what's possible in any energy transformation. They're the "constitution" of thermodynamics—everything else builds on them.

First Law of Thermodynamics

• Energy is conserved—it cannot be created or destroyed, only converted between forms like heat, work, and chemical potential energy
• Internal energy change ($\Delta U$) equals heat added ($q$) minus work done by the system ($w$): $\Delta U = q - w$
• Foundation for all calorimetry—this law is why you can track energy flow by measuring temperature changes

Second Law of Thermodynamics

• Entropy of the universe always increases—spontaneous processes move toward greater disorder in the overall system plus surroundings
• No process is 100% efficient—some energy is always "lost" as heat that increases entropy elsewhere
• Explains directionality—this is why heat flows from hot to cold and why you can't unscramble an egg

Third Law of Thermodynamics

• Entropy approaches zero at absolute zero (0 K) for a perfect crystalline substance
• Absolute zero is unattainable—you can get infinitely close but never reach it in finite steps
• Provides the reference point for calculating absolute entropy values ($S°$) that appear in data tables

---

State Functions: Properties That Don't Care About the Path

State functions depend only on where you start and where you end—not how you got there. Think of it like elevation: whether you hike or take a helicopter, the altitude change is the same.

Enthalpy

• Total heat content ($H$) of a system at constant pressure—most lab reactions occur at atmospheric pressure, making this your go-to energy measure
• $\Delta H$ determines heat flow—negative means exothermic (releases heat), positive means endothermic (absorbs heat)
• State function advantage—you can calculate $\Delta H$ for any reaction using tabulated values, even if you can't measure it directly

Entropy

• Measure of disorder ($S$)—more possible arrangements of particles means higher entropy
• Increases with temperature, phase changes, and more particles—gases have more entropy than liquids, which have more than solids
• $\Delta S$ predicts disorder change—positive $\Delta S$ favors spontaneity; negative $\Delta S$ works against it

Gibbs Free Energy

• The spontaneity predictor—combines enthalpy and entropy into one value: $\Delta G = \Delta H - T\Delta S$
• Negative $\Delta G$ = spontaneous; positive $\Delta G$ = non-spontaneous; zero $\Delta G$ = equilibrium
• Temperature is the tiebreaker—when $\Delta H$ and $\Delta S$ have the same sign, temperature determines which factor "wins"

---

Energy Flow in Reactions: Tracking Heat

These concepts help you quantify and categorize the energy changes that accompany chemical processes. Mastering these is essential for any calculation involving heat.

Endothermic and Exothermic Reactions

• Exothermic ($\Delta H < 0$) releases heat to surroundings—combustion, neutralization, and most bond-forming processes
• Endothermic ($\Delta H > 0$) absorbs heat from surroundings—photosynthesis, melting ice, and dissolving ammonium nitrate
• Surroundings temperature tells the story—if the container feels hot, the reaction is exothermic; if cold, endothermic

Heat Capacity

• Amount of heat to raise temperature by 1°C—larger heat capacity means the substance resists temperature change
• $C_p$ vs. $C_v$—heat capacity at constant pressure ($C_p$) is always greater because some energy does expansion work
• Critical for calorimetry—the equation $q = mC\Delta T$ uses specific heat capacity to convert temperature change to energy

Calorimetry

• Experimental measurement of heat—isolates a reaction in an insulated container to track all energy transfer
• Coffee-cup calorimeter measures at constant pressure (gives $\Delta H$); bomb calorimeter measures at constant volume (gives $\Delta U$)
• Assumes no heat loss—$q_{reaction} = -q_{surroundings}$ lets you calculate enthalpy changes from temperature data

---

Calculation Strategies: Finding $\Delta H$ Without Direct Measurement

These tools let you calculate enthalpy changes for reactions you can't easily run in a calorimeter. Hess's Law and formation data are exam favorites.

Hess's Law

• Enthalpy changes are additive—if you can break a reaction into steps, the sum of step $\Delta H$ values equals the overall $\Delta H$
• Manipulate equations strategically—reverse a reaction (flip the sign of $\Delta H$), multiply coefficients (multiply $\Delta H$ by the same factor)
• State function proof—this works precisely because enthalpy doesn't depend on the pathway

Heat of Formation

• $\Delta H_f°$ is the enthalpy change when 1 mole of compound forms from elements in their standard states
• Elements in standard state have $\Delta H_f° = 0$—this is your reference point (e.g., $O_2(g)$, $C(graphite)$, $Fe(s)$)
• Master equation: $\Delta H_{rxn}° = \Sigma \Delta H_f°(products) - \Sigma \Delta H_f°(reactants)$

Heat of Combustion

• $\Delta H_c°$ is the enthalpy change when 1 mole of substance burns completely in $O_2$
• Always negative—combustion reactions release energy (that's why we burn fuels)
• Useful for organic compounds—often easier to measure than formation, and you can work backward to find $\Delta H_f°$

---

Predicting Spontaneity: Will It Happen on Its Own?

Understanding when reactions proceed without outside help is one of the most powerful applications of thermodynamics. This is where everything comes together.

Spontaneous and Non-Spontaneous Processes

• Spontaneous doesn't mean fast—it means thermodynamically favorable; diamond converting to graphite is spontaneous but takes millions of years
• Spontaneous processes increase universe entropy—even if system entropy decreases, surroundings entropy must increase more
• $\Delta G$ is the ultimate test—forget memorizing rules; just calculate $\Delta G = \Delta H - T\Delta S$

Thermochemical Equations

• Include $\Delta H$ with the balanced equation—shows exactly how much heat per mole of reaction as written
• Physical states matter—$H_2O(l)$ vs. $H_2O(g)$ have different enthalpies; always specify
• Coefficients scale the heat—double the coefficients, double the $\Delta H$

State Functions

• Path-independent properties—enthalpy, entropy, internal energy, and Gibbs Free Energy all qualify
• Why this matters for calculations—you can use any convenient path (like Hess's Law steps) to find the change
• Contrast with path functions—heat ($q$) and work ($w$) depend on how the process occurs, so they're not state functions

---

Quick Reference Table

---

Self-Check Questions

1. Which two state functions combine in the Gibbs Free Energy equation, and what does each one measure?

2. A reaction has $\Delta H = +50 \text{ kJ}$ and $\Delta S = +150 \text{ J/K}$. At what temperature (approximately) does this reaction become spontaneous? What thermodynamic principle explains why temperature matters here?

3. Compare and contrast Hess's Law and using heats of formation—when would you choose one method over the other for calculating $\Delta H_{rxn}$?

4. A student claims that all exothermic reactions are spontaneous. Using the Gibbs Free Energy equation, explain why this statement is incorrect and provide a specific example of when an exothermic reaction would be non-spontaneous.

5. How does the Second Law of Thermodynamics explain why a coffee-cup calorimeter must be insulated, and what assumption does this allow you to make in your calculations?