Why Understanding Enthalpy Matters in Chemistry
Enthalpy is defined as the amount of internal energy and the output of a thermodynamic system's pressure and volume. Enthalpy is an energy-like property or state function that has energy dimensions (and is thus calculated in joules or erg units). The enthalpy H is equivalent to the sum of the internal energy E and the pressure P multiplied with volume V of the system i.e., H = E + PV, respectively.
Under the law of conservation of energy, the shift in internal energy is equal to the heat transmitted to the device, minus the work performed by it. If a change in volume at constant pressure is the only work performed, the change in enthalpy is exactly equal to the heat transferred to the device. The amount of energy is called the enthalpy (or latent heat of vaporization) and is expressed in units of joules per mole when energy needs to be applied to a substance to shift its phase from a liquid to a gas.
Enthalpy Change
The name given to the amount of heat evolved or consumed in a reaction conducted at constant pressure is Enthalpy transition. The symbol of Enthalpy H is referred to as "delta H". At constant pressure, the equation for the change in internal energy, ∆U = q + w can be written as:
∆U = qP – p∆V
Where qP represents the heat absorbed by the system at constant pressure and – p∆V is the expansion work done due to the heat absorbed by the system. The above equation can be written in the terms of initial and final states of the system which is defined below:
UF – UI = qP –p(VF – VI)
Or qP = (UF + pVF) – (UI + pVI)
Enthalpy (H) can be written as H= U + PV. Putting the value in the above equation, we obtained:
qP = HF – HI = ∆H
Hence, change in enthalpy ∆H = qP, referred to as the heat consumed at a constant pressure by the system. At constant pressure, we can also write,
∆H = ∆U + p∆V
Some Key Points
The heat from the device is lost to the surrounding atmosphere during exothermic reactions. ∆H is negative for such reactions. During endothermic reactions, heat is absorbed from the atmosphere by the system. ∆H is positive for such reactions.
Enthalpy of Reactions:
Energy change (U) is equal to the amount of heat produced and the work carried out. Pressure-volume work is called work performed by an expanding gas (or just PV work). For instance, consider a gas-producing reaction, such as dissolving a piece of copper in concentrated nitric acid.
Cu(s)+ 4HNO3(aq) → Cu(NO3)2 (aq) + 2H2O(l) + 2NO2(g)
The quantity of PV work performed by multiplying the external pressure P by the volume change induced by the piston movement (almost V) is found. At constant external pressure, (here, atmospheric pressure),
W = −PΔV
The negative sign associated with PV work performed means that when the volume increases, the device loses energy. The work performed by the system is negative if the volume increases at constant pressure (V> 0), implying that a system has lost energy by performing work on its surroundings. Conversely, the work performed by the system is positive if the volume decreases (almost V<0), which implies that the environment has worked on the system, thereby increasing its energy.
The internal energy U of a system is the sum of all its components' kinetic energy and potential energy. It is the inner energy shift that generates heat plus function. Chemists typically use a related thermodynamic quantity called enthalpy (H) to calculate the energy changes that occur in chemical reactions. Systems’ enthalpy is defined as the sum of their internal energy U plus the product of their pressure P and volume V:
H=U+PV
Since all state functions are internal energy, strain and volume, enthalpy is also a state function. We can therefore characterize a shift in enthalpy ('H) accordingly.
ΔH=Hfinal −Hinitial
If at constant pressure (i.e. for a given P, ΔP=0) a chemical shift occurs, the change in enthalpy ( ΔH) is
ΔH=Δ(U+PV)
=ΔU + ΔPV
=ΔU + PΔV
Substituting q+w for ΔU (First Law of Thermodynamics) and −w for PΔV we obtain
ΔH=ΔU+PΔV
=qp+w−w
=qp
The p subscript is used here to emphasize that this equation is only valid for a constant pressure phase. It is observed that the shift in enthalpy, the H of the system, is equal to the heat obtained or lost at constant pressure.
ΔH=Hfinal−Hinitial
=qp
System in Thermodynamics
A thermodynamic system is a part of matter with a defined boundary on which we concentrate our attention. The system boundary might be fixed or flexible, and it can be real or fictitious. The 3 types of systems in Thermodynamics are-
Isolated System - A system that is separated from its surroundings is unable to exchange both energy and mass. The cosmos is thought to be a self-contained system.
Open System - Both mass and energy can be moved between the system and its surroundings in an open system.
Closed System - The transmission of energy happens across the closed system's border, but the transfer of mass does not.
Different Branches of Thermodynamics
Thermodynamics has 4 major branches, they are
Classical Thermodynamics - The behaviour of matter is investigated using a macroscopic approach in Classical Thermodynamics. Individuals consider units such as temperature and pressure, which aids in the calculation of other properties and the prediction of the characteristics of the matter conducting the process.
Statistical Thermodynamics - Every molecule is in the limelight in Statistical Thermodynamics, which means that the properties of each molecule and how they interact are taken into account to characterise the behaviour of a group of molecules.
Chemical Thermodynamics - The study of how heat & work interacts with each other in a given chemical reactions and state transitions is known as Chemical Thermodynamics.
Equilibrium Thermodynamics - Equilibrium Thermodynamics refers to study of energy & matter transitions as they approach towards equilibrium.
Tips to study Thermodynamics
Thermodynamics is one of the most important chapters in chemistry as aspects of it (Carnot Engine), also appears again in Physics in later chapters.
Thermodynamics also has a lot of weightage in competitive exams like JEE, etc. To get good marks in Thermodynamics, the fundamentals of the chapter, i.e. understanding of different types of systems and energy, etc must be crucial as it sets the base for the complex topic that builds on it. The students should start by going through the NCERT chapter once or twice and then solve all the NCERT Exercises. They can find solutions to these Exercises at Vedantu's official website. If the student has some doubts or needs revision of any topic of Thermodynamics, they can check out Vdantu's Youtube Channel. Here they can find several video lectures on Thermodynamics and other topics.
Teachers also conduct live sessions where questions and solutions are discussed to help the students who don't have the best means to a teacher. Solving questions is really important for understanding Thermodynamics as it has many concepts which relate to formulas and conditions. Students can find the list of important questions of thermodynamics and other topics at Vedantu;'s official website. Previous year's questions are also one of the most suggested ways to study and prepare for exams as it helps to break down the questioning pattern and help the student to explore different types of questions.
This also helps them to create an exam-like situation that gives them a real test with time limits. They can find the papers and solutions at Vedantu's official website. Students should definitely utilize these FREE resources to clear their concepts and their way to get good marks. Vedantu is trying to bring the best out of every child who wants to achieve something or has a zeal to do hard work. A little hard work with some guidance can help any child to achieve the sky and make their parents and teachers proud.
FAQs on Enthalpy Definition and Its Mathematical Derivation
1. What is enthalpy and how is its mathematical expression derived?
Enthalpy, denoted by the symbol H, is a thermodynamic property of a system that represents its total heat content. It is mathematically defined as the sum of the system's internal energy (U) and the product of its pressure (P) and volume (V). The derivation is as follows:
The fundamental equation is: H = U + PV.
According to the first law of thermodynamics, the change in internal energy (ΔU) is given by ΔU = q + w, where q is heat and w is work done.
For a process at constant pressure, the work done (w) by the system is -PΔV. Substituting this, we get ΔU = q_p - PΔV, where q_p is heat at constant pressure.
The change in enthalpy (ΔH) is ΔH = ΔU + PΔV. Substituting the expression for ΔU gives ΔH = (q_p - PΔV) + PΔV.
This simplifies to ΔH = q_p, meaning the change in enthalpy is equal to the heat absorbed or released by the system at constant pressure.
2. What is the key difference between enthalpy (H) and internal energy (U)?
The key difference lies in the conditions under which they measure heat change. Internal energy (U) represents the total energy contained within a system (kinetic and potential), and its change (ΔU) is equal to the heat transferred at constant volume (q_v). In contrast, enthalpy (H) also includes the energy associated with the system's pressure and volume (the 'PV' term). Its change (ΔH) is equal to the heat transferred at constant pressure (q_p). Most chemical reactions in a lab occur at constant atmospheric pressure, making enthalpy the more practical measure.
3. Why is enthalpy (H) a more convenient state function for chemists than internal energy (U)?
Enthalpy is more convenient for chemists because most chemical reactions are carried out in open containers, such as beakers or flasks, which are subject to constant atmospheric pressure. Under these common laboratory conditions, the heat change measured is directly equal to the change in enthalpy (ΔH). Measuring heat change at constant volume (which corresponds to ΔU) would require using a sealed, rigid container called a bomb calorimeter, which is less practical for routine experiments. Therefore, enthalpy provides a direct and easily measurable way to track heat flow in typical chemical processes.
4. How is the sign of an enthalpy change (ΔH) used to classify a reaction?
The sign of the enthalpy change (ΔH) indicates whether a reaction releases or absorbs heat. This is used for classification as follows:
A negative ΔH (ΔH < 0) signifies an exothermic reaction. In this case, the system releases heat into the surroundings, and the enthalpy of the products is lower than the enthalpy of the reactants.
A positive ΔH (ΔH > 0) signifies an endothermic reaction. Here, the system absorbs heat from the surroundings, and the enthalpy of the products is higher than the enthalpy of the reactants.
5. Is enthalpy an extensive or an intensive property? Explain why.
Enthalpy is an extensive property. This means its value depends on the amount of substance present in the system. For example, the enthalpy change for burning 2 moles of methane (CH₄) is exactly double the enthalpy change for burning 1 mole of methane. If the amount of the substance is doubled, the total heat content (enthalpy) also doubles. In contrast, an intensive property, like temperature or density, does not change with the amount of substance.
6. What are some practical examples of processes involving enthalpy change?
Enthalpy changes are observed in many everyday processes:
Exothermic (Negative ΔH): The combustion of fuels like petrol in a car engine, the burning of wood in a fireplace, and the chemical reaction in a hand warmer all release heat into the surroundings.
Endothermic (Positive ΔH): An instant cold pack becoming cold is a classic example. When the inner pouch is broken, a salt like ammonium nitrate dissolves in water, absorbing heat from the surroundings and feeling cold. Photosynthesis is another endothermic process, where plants absorb energy from sunlight.
7. Can enthalpy be described simply as the 'total heat' of a system?
No, this is a common misconception. A system does not 'contain' heat; heat (q) is a form of energy transfer. Enthalpy (H) is a state function representing the system's capacity to release heat at constant pressure. It is defined as H = U + PV. The more accurate statement is that the change in enthalpy (ΔH) is equal to the heat transferred (q_p) during a process that occurs at constant pressure. So, while related, enthalpy itself is a property, not the heat it contains.
