CCE478 - Terminal Effects: Thermodynamics and Chemistry Notes

Units and Conversions

• Mass, m (g) · Moles, n (mol) · Molar mass, M (g/mol)
  • m = n × M
  • n = m / M
  • M = m / n
• Converting Between Mass and Moles (summary formulas):
  • $m = n\,M$
  • $n = \dfrac{m}{M}$
  • $M = \dfrac{m}{n}$

SI Units and Prefixes

• International System of Units (SI) basics
• SI Base Units:
  • Length: meter, symbol $m$
  • Mass: kilogram, symbol $kg$
  • Time: second, symbol $s$
  • Temperature: kelvin, symbol $K$
  • Amount of substance: mole, symbol $mol$
• Secondary units (examples):
  • Newton, symbol $N$
  • Pascal, symbol $Pa$
  • Joule, symbol $J$
• SI prefixes and their factors (examples):
  • tera (T): $10^{12}$
  • giga (G): $10^{9}$
  • mega (M): $10^{6}$
  • kilo (k): $10^{3}$
  • hecto (h): $10^{2}$
  • deka (da): $10^{1}$
  • deci (d): $10^{-1}$
  • centi (c): $10^{-2}$
  • milli (m): $10^{-3}$
  • micro (µ): $10^{-6}$
  • nano (n): $10^{-9}$
  • pico (p): $10^{-12}$
• SI Derived Units (examples):
  • energy: joule, symbol $J$ (also expressed as $N\cdot m$)
  • force: newton, symbol $N$
  • pressure: pascal, symbol $Pa$
  • power: watt, symbol $W$
  • electric charge: coulomb, symbol $C$
  • electric potential: volt, symbol $V$
  • electric resistance: ohm, symbol $\Omega$
• Special note: Unit degree Celsius is equal in magnitude to unit kelvin.

Temperature and Temperature Scales

• Temperature notions:
  • The centigrade (Celsius) scale was created by Anders Celsius in 1742; widely used.
  • The Kelvin scale is based off Celsius with its 0 point at absolute zero.
• Temperature scale conversions (from the slide):
  • $373.15\,\text{K} = 100^{\circ}\text{C} = 212^{\circ}\text{F}$
  • $298.15\,\text{K} = 25^{\circ}\text{C} = 77^{\circ}\text{F}$
  • $310.15\,\text{K} = 37^{\circ}\text{C} = 98^{\circ}\text{F}$
  • $273.15\,\text{K} = 0^{\circ}\text{C} = 32^{\circ}\text{F}$
  • $283.15\,\text{K} = 10^{\circ}\text{C} = 50^{\circ}\text{F}$
  • Absolute zero: $-273.15^{\circ}\text{C} = 0\,\text{K} = -459.58^{\circ}\text{F}$
  • Room temperature: about $20^{\circ}\text{C} \approx 68^{\circ}\text{F}$
• On the nano-scale, temperature measures the average speed of molecules; absolute zero corresponds to zero average velocity.

Forms of Energy

• Forms listed: Heat/Thermal, Potential Energy, Kinetic Energy, Radiant/Light, Electrical, Sound
• Energy units: Joules (J) as the standard; other common units include ft-lb, BTU, kWh, Calories, electron-volts, etc.
• The slide notes that energy can be stored in and transferred between these forms; the dimension is that of energy.

Dimensions and Units (SI)

• Base quantities and symbols:
  • Length: meter, $m$
  • Mass: kilogram, $kg$
  • Time: second, $s$
  • Temperature: kelvin, $K$
  • Amount of substance: mole, $mol$
  • Luminous intensity: candela, $cd$
• SI Derived Units (examples): $N\,(m)$, $Pa$, $J\,(N\cdot m)$, $s^2$, etc. (illustrative from the slide)
• Prefixes recap (as above).

Chemical Reactions

• Reactant: a substance that enters into a chemical reaction and is transformed into products (starting material).
• Product: a substance produced in the course of a chemical reaction (end result).
• Chemical equations express reactions: reactants → products.

Combustion Reactions

• Combustion: a reactant (fuel) combines with oxygen to produce simpler products and heat.
• Example: $2 \;\mathrm{C<em>4H</em>{10}} + 13\; \mathrm{O<em>2} \rightarrow 8\; \mathrm{CO</em>2} + 10\; \mathrm{H_2O} + \text{heat}$

Balanced Chemical Equations

• A chemical equation shows reactants and products; a balanced equation shows the correct stoichiometric ratios.
• Examples:
  • $\mathrm{H<em>2} + \mathrm{O</em>2} \rightarrow \mathrm{H_2O}$
  • $\mathrm{H<em>2} + \tfrac{1}{2}\mathrm{O</em>2} \rightarrow \mathrm{H_2O}$
  • $2\mathrm{H<em>2} + \mathrm{O</em>2} \rightarrow 2\mathrm{H_2O}$

Balancing Chemical Equations

• Coefficients are placed in front of reactant/product formulas to balance.
• Use the smallest possible whole numbers for coefficients (if possible).
• You cannot change the molecular formulas of the reactants or products to balance.

Systems in Thermodynamics

• Define a system: the portion of the universe under study.
• Surroundings: everything outside the system.
• System types:
  • Isolated: no heat or matter exchange with surroundings
  • Closed: exchange of heat allowed, but no matter exchange
  • Open: exchange of heat and matter allowed
• The First Law applies to the system and its surroundings, not just the system alone.
• The transcript shows a condensed energy balance relation for a process, illustrating the partition between system and surroundings.

First Law of Thermodynamics

• Core idea: energy exists in multiple forms, but total energy of a system and its surroundings is conserved.
• Axioms described in the transcript:
  • Axiom 1: Internal energy, $U$, is an intrinsic property of a system, related to measurable coordinates.
  • Axiom 2: The total energy of the system plus surroundings is conserved.
• General form (standard thermodynamics): $\Delta U = Q - W$ where $Q$ is heat added to the system and $W$ is work done by the system.
• Emphasis in the notes is on the role of internal energy and energy exchange with surroundings.

Internal Energy

• Internal energy arises from molecular motion and interactions:
  • Translational, rotational, and vibrational kinetic energy
  • Potential energy from intermolecular bonds
  • Electronic-nuclear interactions and bond energies that hold atoms/molecules together
• Internal energy is called internal to distinguish it from kinetic and potential energies with a particular frame of reference.
• There is no concise standalone thermodynamic definition of internal energy; only changes in U are used in thermodynamics.
• A linked reference is provided to Crash Course Chemistry for further explanation.

Enthalpy

• Definition: $H \equiv U + P V$
• Work for a mechanically reversible, closed-system process: $W = -\int P\,dV$
• Combining with the energy balance leads to the use of enthalpy as a convenient property, $H = U + P V$.
• Constant-volume processes: $\Delta U = Q$ (since $W = 0$).
• Constant-pressure processes: $-V = \dfrac{d}{d t}(U + P V) = \dfrac{dQ}{d t}$; equivalently, at constant pressure, $\Delta H = Q$.
• Implication: Enthalpy plays a role at constant pressure analogous to internal energy at constant volume.

Enthalpy and Constant-Pressure Processes (Key Insight)

• For constant pressure: $\Delta H = Q$ (heat added at constant pressure).
• This forms the basis for calorimetry measurements at atmospheric pressure and similar constant-pressure analyses.

Work (Pressure-Volume) and Its Significance

• Categories of work include:
  • Mechanical work
  • Pressure-Volume (PV) work
  • Rotational work
  • Spring work
  • Non-mechanical work: electric field, electrical polarization, magnetic work, gravitational work
• PV work formula (classic): $W = -\int P\,dV$
• PV work describes the work associated with the displacement of a piston in a cylinder, i.e., compression or expansion of a gas.
• Conditions for PV work to be valid:
  • The system is infinitesimally displaced from a state of internal equilibrium
  • The system is infinitesimally displaced from a state of mechanical equilibrium with its surroundings

Closed Systems (Recap)

• In a closed system, mass is constant (no transfer of matter across the boundary).
• All energy exchange with surroundings occurs as heat or work.
• The total energy change of the surroundings equals the net energy transferred to or from it as heat and work.

Summary: Energy Interconversions and Boundaries

• The First Law emphasizes energy conservation across system + surroundings.
• The distinction between constant-volume and constant-pressure processes links Q, W, U, and H in practical ways:
  • Constant volume: $\Delta U = Q$ (no PV work)
  • Constant pressure: $\Delta H = Q$ (heat at constant pressure equals enthalpy change)

Real-World Relevance and Connections

• These concepts underpin energy release in combustion, safety analyses, and lethality assessments where heat transfer, work, and energy storage determine outcomes.
• The material connects foundational thermodynamics to practical calculations in chemical reactions, engineering systems, and safety assessments.

References in the Transcript

• Crash Course Chemistry – Enthalpy (video references provided in the transcript):
  • https://www.youtube.com/watch?v=SV7U4yAXL5I&list=PL8dPuuaLjX tPHzzYuWy6fYEaX9mQQ8oGr&index=19
• Additional practice problems and examples are indicated in the transcript.

Questions?