3.1 Forms of energy and energy transfer

Forms of Energy and Energy Transfer

Energy exists in multiple forms, and thermodynamics is fundamentally about tracking how energy moves and converts between those forms. Understanding these distinctions is essential because the First Law of Thermodynamics rests on one core idea: energy is conserved during every transfer and conversion. This topic lays the groundwork for analyzing any thermodynamic system you'll encounter in this course.

Forms of Energy

Kinetic energy

Kinetic energy is the energy an object has because of its motion. It depends on both mass and velocity:

$KE = \frac{1}{2}mv^2$

Notice that velocity is squared, so doubling an object's speed quadruples its kinetic energy. A 1,000 kg car traveling at 30 m/s has $KE = \frac{1}{2}(1000)(30)^2 = 450{,}000 \text{ J}$ (450 kJ). This squared relationship is why high-speed collisions are so much more destructive than low-speed ones.

Potential energy

Potential energy is stored energy due to an object's position or configuration. The two most common types:

• Gravitational potential energy: $PE = mgh$, where $h$ is height above a reference point. A 5 kg book on a 2 m shelf has $PE = (5)(9.81)(2) = 98.1 \text{ J}$ relative to the floor.
• Elastic potential energy: Energy stored in deformed elastic materials like compressed springs or stretched rubber bands, calculated using $PE = \frac{1}{2}kx^2$, where $k$ is the spring constant and $x$ is the displacement.

Thermal energy

Thermal energy comes from the random motion of particles (atoms and molecules) within a substance. The faster the particles move on average, the higher the temperature and the greater the thermal energy. Thermal energy depends on both the temperature and the amount of substance present. A bathtub of warm water holds more thermal energy than a cup of boiling water, even though the cup is at a higher temperature. Thermal energy is closely related to the internal energy of a system, which you'll use extensively when applying the First Law.

Chemical energy

Chemical energy is stored in the bonds between atoms within molecules. When those bonds break and new ones form during a chemical reaction, energy is either released (exothermic) or absorbed (endothermic). Gasoline stores roughly 34 MJ per liter of chemical energy, which an engine converts into thermal and then mechanical energy. Batteries convert chemical energy into electrical energy through electrochemical reactions.

Electrical energy

Electrical energy is associated with the movement of electric charges. It can be stored in electric fields (as in a capacitor) or transported through conductors via electric current. Power lines carry electrical energy over long distances, and lightning is a dramatic natural example of electrical energy discharge.

Energy Conservation and Conversion

The conservation principle

Energy cannot be created or destroyed. It can only be converted from one form to another. The total energy of an isolated system (one with no energy or mass crossing its boundary) remains constant. This is the most fundamental idea in thermodynamics.

Common energy conversions

• Kinetic → Potential: A roller coaster car climbing a hill slows down as kinetic energy converts to gravitational potential energy.
• Potential → Kinetic: A skydiver in free fall accelerates as gravitational potential energy converts to kinetic energy.
• Chemical → Thermal: Burning fuel in a campfire breaks chemical bonds and releases thermal energy.
• Electrical → Kinetic: An electric motor converts electrical energy into rotational kinetic energy to drive tools or vehicles.

Most real processes involve a chain of conversions. A coal power plant, for example, goes from chemical → thermal → kinetic → electrical.

Efficiency

No real energy conversion is 100% efficient. Some energy is always dissipated as waste heat due to friction, electrical resistance, or other irreversibilities. Efficiency is defined as:

$\eta = \frac{\text{useful energy output}}{\text{total energy input}} \times 100\%$

An incandescent bulb converts only about 5% of electrical energy to light (the rest becomes heat), while an LED achieves roughly 40-50%. Maximizing efficiency is a central goal in engineering design.

Energy Transfer Processes

Energy crosses a system boundary through three mechanisms: work, heat, and mass transfer.

Work

Work is energy transfer that occurs when a force acts through a distance. In its general form:

$W = \int F \cdot ds$

For a gas expanding against a piston, this becomes $W = \int P \, dV$, where $P$ is pressure and $V$ is volume. Work is a path-dependent quantity, meaning the amount of work depends on how the process occurs, not just the start and end states.

Heat

Heat is energy transfer driven by a temperature difference between a system and its surroundings. It occurs through three modes:

• Conduction: Direct molecular contact (a metal pan on a hot burner)
• Convection: Bulk fluid motion carrying energy (warm air rising from a heater)
• Radiation: Electromagnetic waves requiring no medium (sunlight warming the Earth)

For sensible heat transfer (no phase change), the energy transferred is:

$Q = mc\Delta T$

where $m$ is mass, $c$ is specific heat capacity, and $\Delta T$ is the temperature change. Like work, heat is path-dependent.

Mass transfer

In open systems, matter flows across the system boundary and carries energy with it. The energy associated with flowing mass is characterized by its enthalpy ($h$), which accounts for both internal energy and the flow work needed to push the mass into or out of the system. Heat exchangers, turbines, and compressors are all open systems where mass transfer is a critical energy mechanism.

First Law of Thermodynamics Applications

Statement of the First Law

The First Law is simply the conservation of energy applied to a thermodynamic system. For a closed system:

$\Delta U = Q - W$

The change in internal energy ($\Delta U$) equals the net heat added to the system ($Q$) minus the net work done by the system ($W$).

Closed systems

A closed system has no mass crossing its boundaries. Energy can only enter or leave as heat or work. Examples include a sealed piston-cylinder device and a closed refrigeration loop. For these systems, you apply $\Delta U = Q - W$ directly.

Open systems

An open system allows mass to flow in and out. The energy balance must account for heat, work, and the energy carried by mass flow. A steam turbine is a classic open system: high-energy steam enters, transfers energy to the turbine shaft as work, and exits as lower-energy steam.

Problem-solving approach

When tackling First Law problems, follow these steps:

1. Define the system and its boundaries. Draw a sketch and label what's inside versus outside.
2. Classify the system as closed (no mass flow) or open (mass crosses the boundary).
3. Identify all energy transfers crossing the boundary: heat in or out, work in or out, and mass flow (for open systems).
4. Apply the First Law with the correct form of the energy balance.
5. Use consistent sign conventions. A common convention: $Q > 0$ for heat into the system, $W > 0$ for work done by the system. Always state which convention you're using, since textbooks vary on this.