The study of the relationship between heat and other forms of energy.
The laws of thermodynamics
Thermodynamics is the study of the relationship between heat and other forms of energy. It is a fundamental field of physics, underpinned by four laws which describe how energy behaves in different systems. The first law states that energy can neither be created nor destroyed; it can only change form. This means that when heat moves from one object to another, the total amount of energy remains constant.
The second law explains why some processes are irreversible: entropy always increases over time as heat flows from hotter objects to cooler ones until equilibrium is reached. Entropy measures the degree of disorder in a system, so this law implies that all natural processes tend towards greater disorder or chaos over time.
The third law states that as temperatures approach absolute zero, entropy tends to a constant value. Essentially, the third law states that if an object reached absolute zero, its atoms would stop moving.,In addition to these three laws, what is commonly known as the zeroth law of thermodynamics states that if two bodies are in equilibrium with a third body, they are also in equilibrium with each other.
The first law of thermodynamics
The first law of thermodynamics states that energy can neither be created nor destroyed, but only changed from one form to another. This is known as the conservation of energy and it means that in an isolated system – without any matter or energy coming in or going out – the total energy must remain constant In other words, when work is done on a system, the total amount of energy remains constant; any increase in one type of energy must be balanced by a decrease in another type.
For example, when a car brakes its kinetic energy is converted into heat due to friction between the brake pads and wheels. Similarly, we can understand metabolic processes in the human body by using the first law of thermodynamics. We obtain energy from the environment by consuming food, and this energy is used to do work, transferred as heat or stored as chemical energy in the body.
The second law of thermodynamics
The second law of thermodynamics states that the entropy, or the degree of disorder in a system, of the universe always increases over time. This means that all natural processes tend towards greater chaos and disorganization as energy is dispersed from hotter to cooler objects. For example, when you leave a cup of coffee on the table it will eventually cool down due to heat transfer from the hot liquid to its surroundings.
Additionally, the second law of thermodynamics states that if a process is irreversible, it must result in an increase in the combined entropy of the system and the environment. A burning campfire is an example of increasing entropy in an irreversible process. As the wood burns, energy and vapors are released, spreading in an expanding cloud and leaving behind nothing but ashes. The remnants of a fire will never be turned back into wood. The foundations for the second law of thermodynamics were laid by the scientist Rudolf Clausius in 1850.
Heat engines and efficiency
Heat engines are devices that convert thermal energy into mechanical work. They operate by transferring heat from a hot place to a cold one, and diverting some of this energy into mechanical energy. The most common heat engines are internal combustion engines. The efficiency of engines such as these is determined by how much useful work they can produce compared to the amount of energy put into them.
The Second Law of Thermodynamics states that it is impossible for any heat engine to reach 100% efficiency due to entropy; some energy will always be lost as waste heat during the process. This means that no matter how advanced an engine may be, there will always be some inefficiency associated with it. Heat engines commonly reach only around 30-50% efficiency, as a result of practical constraints.
Despite this limitation, engineers continue to strive towards greater efficiencies through improved designs and materials – such as turbochargers, which increase air pressure within cylinders for more powerful combustion.
The Carnot cycle
The Carnot cycle is a theoretical thermodynamic cycle that describes the most efficient way to convert heat into work. It consists of four phases: isothermal expansion, adiabatic expansion, isothermal compression and adiabatic compression. During the first two phases, heat is transferred to an ideal gas in the system from a hot reservoir and causes the gas to expand and do work on its surroundings. During the second two phases, heat leaves the system as the gas contracts. This process can be repeated indefinitely for maximum efficiency.
Unfortunately, this idealized model cannot be achieved in reality due to friction and other losses associated with real-world engines. Nevertheless, its principles are still used today in modern engineering designs such as refrigerators which use a reversed analogous cycle to cool their interiors by transferring heat outwards instead of inwards.
Entropy
Entropy is a measure of the amount of energy in a system that is unavailable for doing work. It can be thought of as the disorder or randomness within a system, and it tends to increase over time. This means that any process which involves an increase in entropy will also involve an increase in disorder. For example, when ice melts into liquid water, its molecules become more disordered and spread out; this results in an increase in entropy and thus less energy available to do work.
The Second Law of Thermodynamics states that the entropy of the universe always increases over time, meaning that all processes tend towards greater disorder and lower efficiency. This law has far-reaching implications for our universe: without it, stars would not burn fuel efficiently enough to sustain life on Earth! In fact, some scientists believe that the increasing entropy of our universe may eventually lead to its eventual heat death – at which point no free thermodynamic energy remains in the universe.
Gibbs free energy
Gibbs free energy is a thermodynamic quantity that describes the maximum amount of work which can be done by a closed system. It can also determine whether a reaction will occur spontaneously or not. It is calculated using enthalpy and entropy, two other thermodynamic quantities. Enthalpy measures the total energy of a system, including both heat and work; it can be worked out by adding the internal energy of a system to the product of its temperature and pressure.
Entropy measures the amount of energy in a system that is unavailable for doing work. Gibbs free energy is calculated by deducting the product of entropy and temperature from enthalpy. This gives us an indication of how much useful work can be extracted from any given process – if this value is negative then the reaction will occur spontaneously; if positive then it won’t.
Phase changes and phase diagrams
The most common states of matter are solid, liquid and gas. Solids have a fixed shape and volume, liquids take the shape of their container but maintain a constant volume, while gases fill their containers completely and expand or contract to match changes in pressure. Common phase changes between these states include melting (solid to liquid), vaporization (liquid to gas) and sublimation (solid to gas).
Phase diagrams are graphical representations of how temperature and pressure affect the state of a substance. They show which phases will be present at different temperatures and pressures, as well as the boundaries between them. For example, water’s phase diagram shows that it can exist as ice below 0°C at atmospheric pressure; above this point it melts into liquid water until 100°C when it boils off into steam. By plotting points on such diagrams we can gain insight into how substances behave under different conditions – for instance why some materials become brittle when cooled too quickly or why certain foods spoil more easily in warmer climates!
Ideal gases
An ideal gas is a theoretical concept used to describe the behavior of gases under certain conditions. It assumes that all molecules in the gas have no interactions with each other, allowing them to move freely, randomly, and independently. This simplifies calculations when considering thermodynamic processes such as expansion or compression, making it easier to predict how a real-world system will behave.
At standard temperature and pressure (STP), many common gases can be considered ideal for the purposes of scientists, including nitrogen, oxygen, hydrogen and carbon dioxide. These gases obey the Ideal Gas Law which states that PV = nRT where P is pressure, V is volume, n is number of moles of gas present, R is the ideal gas constant, and T is temperature in Kelvin. The Ideal Gas Law also explains why some substances become lighter when they are heated: because their molecules spread out further due to increased thermal energy causing them to take up more space per unit mass.
Real gases and the van der Waals equation
Real gases differ from ideal gases in that they have molecules which take up a finite volume and have attractive forces between them. At high pressures, the volume of most real gases is higher than the ideal gas law predicts. At sufficiently low temperatures, attractive forces cause the molecules to stick together more easily, leading to gases condensing into liquids.
The van der Waals equation takes these effects into account and provides a better description of real gas behavior than the Ideal Gas Law. It includes two additional terms: one for molecular size and another for intermolecular attraction.
These terms become increasingly important as pressure or temperature increases – for example when carbon dioxide is compressed at room temperature it behaves very differently than predicted by the Ideal Gas Law!
Interestingly enough, this equation was first proposed by Dutch physicist Johannes Diderik van der Waals back in 1873 – over 150 years ago! Since then it has been used extensively to model real-world systems such as combustion engines and refrigeration cycles. Today it remains an invaluable tool for understanding thermodynamics on both macroscopic and microscopic scales.
