5.1 Energy Basics

Energy Types and Changes

Energy is what drives every physical and chemical process. It shows up in different forms, and it constantly converts from one form to another. Understanding these forms and conversions is the foundation of thermochemistry.

Types and changes of energy

Energy comes in two broad categories:

• Kinetic energy is the energy of motion. It depends on an object's mass and velocity. A moving car and a flowing river both have kinetic energy.
• Potential energy is stored energy based on an object's position or configuration. There are several types:
  • Gravitational potential energy depends on an object's height and mass (a book sitting on a shelf)
  • Chemical potential energy is stored within chemical bonds (the energy in food or fuel)
  • Elastic potential energy is stored in stretched or compressed objects (a stretched spring)

The Law of Conservation of Energy states that energy cannot be created or destroyed, only converted from one form to another. This is a big deal in chemistry. A rollercoaster is a classic example: at the top of a hill, energy is mostly gravitational potential. As it drops, that converts to kinetic energy. At the bottom, kinetic energy is at its maximum.

Chemical processes involve energy changes tied to breaking and forming bonds:

• Exothermic reactions release energy to the surroundings (combustion, for example, gives off heat)
• Endothermic reactions absorb energy from the surroundings (like an ice pack getting cold when activated)

Heat vs thermal energy vs temperature

These three terms get mixed up constantly, but they mean different things:

• Thermal energy is the total kinetic energy of all the particles in a substance due to their random motion. It depends on both the number of particles and their average kinetic energy. A bathtub of warm water has more thermal energy than a cup of boiling water, even though the cup is hotter, because the bathtub has far more particles.
• Temperature measures the average kinetic energy of particles in a substance. It reflects how intense the thermal energy is, not how much total energy there is. It's measured in Kelvin (K), Celsius (℃), or Fahrenheit (℉).
• Heat is the transfer of thermal energy between substances due to a temperature difference. Heat always flows from higher temperature to lower temperature. It's measured in joules (J) or calories (cal).

Specific heat and heat capacity

Different substances heat up at different rates. That's where specific heat comes in.

Specific heat capacity ($c$) is the amount of heat required to raise the temperature of 1 gram of a substance by 1°C (or 1 K). It's measured in J/(g·°C) and varies by substance. Water has a notably high specific heat capacity of 4.18 J/(g·°C), which is why it heats up and cools down slowly compared to metals. This property makes water an effective coolant and is a major reason coastal climates are more moderate than inland ones.

Heat capacity ($C$) is the amount of heat required to raise the temperature of an entire object or sample by 1°C. It's measured in J/°C and depends on both the mass and the specific heat of the substance:

$C = mc$

So a large pot of water has a much greater heat capacity than a small cup, even though they have the same specific heat capacity.

Calculations for heat transfer

Heat transfer is calculated using this formula:

$Q = mc\Delta T$

where:

• $Q$ = heat transferred (J)
• $m$ = mass (g)
• $c$ = specific heat capacity (J/(g·°C))
• $\Delta T$ = temperature change (°C or K)

Steps to solve a heat transfer problem:

1. Identify the mass ($m$), specific heat capacity ($c$), and temperature change ($\Delta T$) of the substance
2. Plug the values into $Q = mc\Delta T$
3. Solve for $Q$

Example: How much heat is needed to raise the temperature of 500 g of water from 20°C to 80°C? (Specific heat of water = 4.18 J/(g·°C))

• $Q = mc\Delta T$
• $Q = (500 \text{ g})(4.18 \text{ J/(g·°C)})(80°C - 20°C)$
• $Q = (500)(4.18)(60)$
• $Q = 125{,}400 \text{ J} = 125.4 \text{ kJ}$

A positive $Q$ value means the substance absorbed heat. A negative $Q$ means it released heat. If a problem asks you to find $\Delta T$ or $m$ instead, just rearrange the same formula.

Energy, Work, and Power

A few more terms you'll see in thermochemistry:

• Work is the transfer of energy when a force moves an object in the direction of that force.
• Power is the rate at which work is done or energy is transferred, measured in watts (W), where 1 W = 1 J/s.
• Thermodynamics is the broader study of heat, temperature, and their relationship to energy and work. Thermochemistry is one branch of it.
• Entropy is a measure of disorder or randomness in a system. Natural processes tend to increase entropy over time. You'll explore this more in later units, but for now, just know that energy transformations aren't perfectly efficient because some energy always disperses as heat, increasing entropy.