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Thermodynamics Class 11 Chemistry NCERT Chapter 5 One Shot | New NCERT CBSE | Full chapter

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Thermodynamics explores energy conversion and its effect on matter, primarily through chemical reactions that involve energy changes during bond formation and breaking. The laws of thermodynamics apply to macroscopic systems, determining reaction feasibility and extent, while emphasizing the interplay between internal energy, enthalpy, and entropy in various processes.

Insights

  • Thermodynamics is the study of heat and energy changes, particularly in chemical processes, where energy is absorbed when bonds break and released when bonds form, emphasizing the significance of energy transformations in reactions.
  • The laws of thermodynamics—First, Second, and Third—apply to large systems made up of many molecules, highlighting how energy changes govern both physical and economic systems, and illustrating the scale and complexity of these interactions.
  • Different types of thermodynamic systems—open, closed, and isolated—are defined by their exchanges of matter and energy, which are crucial for understanding how energy transfers affect internal energy and system behavior.
  • The first law of thermodynamics, expressed as ΔU = Q + W, indicates that the change in internal energy is equal to the heat added to the system plus the work done on the system, underlining the conservation of energy principle.
  • Enthalpy (ΔH) is introduced as a measure of heat transfer at constant pressure, and its relationship with internal energy is crucial for understanding chemical reactions, particularly in the context of formation and dissolution processes.
  • The concept of spontaneity in thermodynamics is linked to changes in entropy (ΔS) and Gibbs free energy (G), where spontaneous processes are characterized by an increase in disorder, as described by the Second Law of Thermodynamics.

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Recent questions

  • What is thermodynamics in simple terms?

    Thermodynamics is the study of heat and energy. It focuses on how energy is converted and transferred, particularly in relation to matter. This field examines the laws governing energy changes, such as how heat affects physical and chemical processes. Thermodynamics is essential for understanding various phenomena, from engines to chemical reactions, as it provides insights into energy efficiency and the feasibility of processes.

  • How do chemical reactions release energy?

    Chemical reactions release energy primarily through the formation of new bonds. When reactants undergo a transformation, bonds in the reactants break, which requires energy input. However, as new bonds form in the products, energy is released. The overall energy change depends on the balance between the energy needed to break bonds and the energy released when new bonds are created. This energy change is crucial for determining whether a reaction is exothermic (releases energy) or endothermic (absorbs energy).

  • What are the laws of thermodynamics?

    The laws of thermodynamics are fundamental principles that describe how energy behaves in physical systems. The First Law states that energy cannot be created or destroyed, only transformed, emphasizing the conservation of energy. The Second Law introduces the concept of entropy, stating that energy transformations increase disorder in a system. The Third Law asserts that as temperature approaches absolute zero, the entropy of a perfect crystal approaches zero. Together, these laws provide a framework for understanding energy interactions in various processes.

  • What is internal energy in thermodynamics?

    Internal energy refers to the total energy contained within a thermodynamic system, encompassing various forms such as kinetic, potential, and thermal energy. It is an extensive property, meaning it depends on the amount of matter present in the system. Changes in internal energy occur due to heat transfer or work done on or by the system. The First Law of Thermodynamics relates changes in internal energy to heat and work, highlighting its significance in analyzing energy transformations during physical and chemical processes.

  • What is the significance of entropy in thermodynamics?

    Entropy is a measure of disorder or randomness in a system and plays a crucial role in thermodynamics. It quantifies the amount of energy in a system that is unavailable to do work. The Second Law of Thermodynamics states that the total entropy of an isolated system can never decrease over time, indicating that natural processes tend to move towards greater disorder. This principle helps predict the spontaneity of reactions and the direction of energy flow, making entropy a key concept in understanding thermodynamic processes and the behavior of systems.

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Summary

00:00

Understanding Thermodynamics and Energy Dynamics

  • Thermodynamics studies heat dynamics, focusing on energy conversion and its effects on matter, particularly in chemical processes, termed chemical thermodynamics.
  • Chemical reactions involve bond formation and breaking, consuming energy when bonds break and releasing energy when bonds form, highlighting energy changes in reactions.
  • The laws of thermodynamics—First, Second, and Third—apply to macroscopic systems, which consist of large numbers of molecules, emphasizing energy changes in these systems.
  • Macroscopic systems can be compared to economies, where national economic discussions differ from individual financial situations, illustrating the scale of thermodynamic systems.
  • Thermodynamics can determine the feasibility of reactions and the extent to which they occur, but it does not provide information on reaction speed or pathways.
  • Three types of thermodynamic systems are defined: open systems (exchange of matter and energy), closed systems (exchange of energy only), and isolated systems (no exchange of matter or energy).
  • Thermodynamic processes include isothermal (constant temperature), isobaric (constant pressure), isochoric (constant volume), and adiabatic (no heat transfer), each with specific characteristics.
  • The state of a system is defined by properties such as pressure, volume, and temperature, known as state variables or state functions, which depend on the system's state.
  • Path functions, like work and heat transfer, depend on the route taken to achieve a state, illustrating that different paths can result in different energy transfers.
  • Internal energy represents the total energy within a thermodynamic system, encompassing various forms such as potential, mechanical, and electrical energy.

16:39

Understanding Internal Energy and Thermodynamics

  • Extensive properties depend on the quantity of matter, while intensive properties do not; internal energy is an extensive property influenced by the amount of matter present.
  • The absolute value of internal energy cannot be determined, but changes in internal energy can be measured when matter enters or leaves a system or when heat is exchanged.
  • Internal energy changes occur when heat is absorbed (positive change) or released (negative change), denoted as ΔU for change in internal energy.
  • An isothermal process occurs at constant temperature, resulting in no change in internal energy (ΔU = 0) since there is no temperature difference to drive heat transfer.
  • In an adiabatic system, no heat transfer occurs, but work can be done on or by the system, leading to changes in internal energy through mechanical or electrical work.
  • Mechanical work can be performed by stirring a liquid, while electrical work can be done using an immersion rod to heat water, both increasing internal energy.
  • The change in internal energy (ΔU) is the same regardless of whether mechanical or electrical work is done, as demonstrated by Joule's experiments on adiabatic work.
  • The first law of thermodynamics states that the change in internal energy (ΔU) equals the heat transferred (Q) plus the work done (W), expressed as ΔU = Q + W.
  • In an isolated system, no heat or work transfer occurs, resulting in constant internal energy, illustrating the law of conservation of energy.
  • The relationship between state functions and path functions is highlighted, where ΔU is a state function, while Q and W are path-dependent, emphasizing the importance of understanding energy transfer processes.

32:26

Understanding Work in Gas Compression and Expansion

  • The change in volume is calculated as the final volume minus the initial volume, indicating a decrease when the piston moves inward, resulting in a negative delta V.
  • Pressure is defined as force per unit area, and the force on the piston can be expressed as pressure multiplied by area (F = P × A).
  • Work done on the system by moving the piston is calculated as force multiplied by the distance covered (W = F × l), which can also be expressed as W = P × ΔV.
  • When the volume decreases, delta V is negative, indicating that work done on the system must be positive, leading to the equation W = -P × ΔV.
  • During gas compression, work is done on the system, while during expansion, the system does work, resulting in different signs for work done based on the volume change.
  • The work done during a reversible process is calculated using integration, expressed as W = -∫(P dV) from V1 to V2, where V1 is the initial volume and V2 is the final volume.
  • In irreversible processes, work is done in one shot, while reversible processes involve gradual changes, allowing for the possibility of reversing the process.
  • The area under the pressure-volume curve represents the work done, with irreversible processes showing a larger area due to instantaneous changes compared to gradual changes in reversible processes.
  • Free expansion of gas in a vacuum results in zero work done, as external pressure is zero, leading to the conclusion that ΔU = Q in such scenarios.
  • At constant volume, the change in internal energy equals the heat transfer, establishing a critical relationship in thermodynamics for processes occurring under these conditions.

49:54

Thermodynamics of Internal Energy and Enthalpy

  • Internal energy is influenced by particle motion; temperature changes affect this motion, thus altering internal energy levels in substances.
  • Isothermal processes maintain constant temperature, resulting in zero change in internal energy (ΔU = 0) during such transformations.
  • The first law of thermodynamics states ΔU = Q + W, where Q is heat added and W is work done; in isothermal conditions, ΔU remains zero.
  • For isothermal irreversible processes, heat absorbed (Q) equals negative work done (W), expressed as Q = -W, with W calculated as P_ext * ΔV.
  • In free expansion (isothermal into a vacuum), external pressure (P_ext) is zero, leading to no work done (W = 0) and thus Q also equals zero.
  • For isothermal expansion against a constant external pressure of 1 atmosphere, work done is calculated as W = -P_ext * (V_final - V_initial), with V_initial = 2 liters and V_final = 10 liters.
  • In reversible isothermal expansion, work done (W) is calculated using the formula W = 2.303 * nRT * log(V_final/V_initial), where n = number of moles and R = gas constant.
  • Enthalpy (ΔH) is introduced for processes at constant pressure, where ΔH = ΔU + PΔV, highlighting the relationship between internal energy and enthalpy.
  • For gases, the difference in moles of gaseous products and reactants (Δn) is crucial in calculating changes in enthalpy and internal energy during reactions.
  • Extensive properties, like internal energy and enthalpy, depend on the quantity of matter, while intensive properties, such as pressure and temperature, do not.

01:07:29

Physical Properties of Matter and Energy

  • The size of a bottle affects its volume; increasing size increases volume, while decreasing size reduces it, emphasizing the relationship between size and volume in physical properties.
  • Temperature influences the properties of matter; as temperature increases, the matter expands, while pressure remains constant, indicating that temperature is a critical factor in physical changes.
  • Heat capacity is defined as the amount of heat required to raise the temperature of a given mass by one degree, represented mathematically as q = c * ΔT, where c is heat capacity.
  • Specific heat capacity refers to the heat needed to raise the temperature of one unit mass of a substance by one degree, expressed as q = m * s * ΔT, where m is mass and s is specific heat.
  • The relationship between heat capacity and specific heat capacity is established as heat capacity equals mass multiplied by specific heat capacity (C = m * s).
  • At constant volume, heat transfer equals the change in internal energy (ΔU = qV), while at constant pressure, it equals the change in enthalpy (ΔA = qP).
  • A calorimeter measures changes in internal energy and enthalpy, based on the principles of heat transfer at constant volume and pressure, respectively.
  • Reaction enthalpy is the change in enthalpy during a chemical reaction, calculated as the sum of the enthalpy of products minus the sum of the enthalpy of reactants.
  • Standard enthalpy refers to the enthalpy of substances in their standard states at 298 Kelvin and 1 bar pressure, denoted with a superscript circle.
  • Enthalpy changes during phase transformations, such as fusion, vaporization, and sublimation, are defined for one mole of substance at standard states, with positive values indicating heat absorption.

01:25:02

Understanding Enthalpy and Thermochemical Equations

  • Moles of gaseous reactants are zero, leading to a delta n value of zero; the ideal gas constant (R) is 8.3 J/(mol·K) at 298 K.
  • Convert energy units from joules to kilojoules by multiplying by 10, resulting in 41.53 kJ/mol for internal energy of vaporization.
  • Standard enthalpy of formation is defined as the enthalpy change when one mole of a compound forms from its most stable elements in standard states.
  • Reference states for elements are their most stable forms: H2 gas for hydrogen and O2 gas for oxygen, ensuring reactants and products are pure at 1 bar pressure.
  • Standard molar enthalpy of formation is denoted as ΔfH, while reaction enthalpy is denoted as ΔrH, with formation requiring elements, not compounds, as reactants.
  • Example: Calcium carbonate (CaCO3) forms from calcium (Ca), carbon (C), and oxygen (O), not from other compounds like CaO2, which does not qualify for formation enthalpy.
  • Enthalpy of formation applies to one mole of a compound, while reaction enthalpy can involve multiple moles, as seen in hydrogen bromide (HBr) formation.
  • The enthalpy of formation for elements in their reference state is conventionally zero, e.g., solid iron (Fe) and gaseous hydrogen (H2) both have ΔfH = 0.
  • A thermochemical equation is a balanced chemical equation that includes the reaction enthalpy value, expressed in kJ/mol, and must be balanced in terms of moles, not molecules.
  • Hess's Law states that the total enthalpy change for a reaction is the sum of the enthalpy changes for each step, applicable when reactions occur at the same temperature.

01:40:56

Understanding Enthalpy in Chemical Reactions

  • Enthalpy of bond dissociation refers to the energy required to break a bond in a molecule, such as H2, resulting in the formation of hydrogen atoms.
  • The change in enthalpy during a reaction can also be termed as enthalpy of atomization, which is equivalent to bond dissociation enthalpy for diatomic molecules.
  • In polyatomic molecules like methane (CH4), breaking the four C-H bonds requires different amounts of energy due to the varying stability of the remaining bonds.
  • The concept of mean bond enthalpy is introduced to calculate average energy required to break bonds in polyatomic molecules, as not all bonds require the same energy.
  • Lattice enthalpy is defined as the enthalpy change when an ionic compound is formed from gaseous ions, and it is always a positive value due to energy absorption.
  • The Born-Haber cycle is a method used to indirectly calculate lattice enthalpy, as it cannot be determined experimentally.
  • Lattice enthalpy differs from standard enthalpy of formation, which measures the enthalpy change when a compound forms from its most stable elements.
  • The Born-Haber cycle involves multiple steps, including sublimation, ionization, bond dissociation, and electron gain enthalpy, to calculate the overall lattice enthalpy.
  • Hess's law states that the total enthalpy change for a reaction is equal to the sum of the enthalpy changes for each step in the process.
  • An example calculation shows that if all necessary enthalpy values are provided, lattice enthalpy can be determined, with a sample value of -788 kJ/mol for a specific reaction.

01:59:38

Understanding Enthalpy in Salt Solutions

  • A salt solution is created by dissolving salt in water, and adding more water to this solution results in dilution, which increases the solvent's volume significantly.
  • Enthalpy of solution is defined as the enthalpy change when one mole of a substance dissolves in a solvent, such as salt dissolving in a specific mass of water.
  • The enthalpy change during dissolution is denoted as ΔA, indicating the energy change associated with the process of dissolving a solute in a solvent.
  • When salt dissolves in water, it dissociates into ions (e.g., Na⁺ and Cl⁻), and the enthalpy of this reaction can be calculated using Hess's law.
  • The enthalpy of solution can be expressed mathematically as the sum of lattice enthalpy and hydration enthalpy, which are crucial for understanding dissolution processes.
  • Salts with high lattice enthalpy require significant energy to break apart, making them less soluble in water or other solvents.
  • As more solvent is added, the value of ΔA becomes more negative, indicating an increase in the enthalpy of dilution until it reaches a limiting value.
  • The enthalpy of dilution is defined as the heat withdrawn from the surroundings when additional solvent is added, causing the solution to become hotter.
  • Enthalpy of solution and enthalpy of dilution are distinct concepts, with the latter specifically measuring the change in ΔA due to added solvent.
  • Spontaneous processes in thermodynamics are characterized by a decrease in energy and an increase in disorder, which is measured by entropy, denoted as S.

02:17:53

Thermodynamics and Entropy Explained

  • Entropy change (ΔS) is zero for reversible processes, while for irreversible processes, ΔS is greater than zero, indicating total entropy increases in the universe.
  • The Gibbs function (G) is defined as G = H - TΔS, where a negative ΔG indicates a spontaneous process, while a positive ΔG indicates a non-spontaneous process.
  • The Second Law of Thermodynamics states that the change in entropy (ΔS) is always greater than or equal to zero, reinforcing that spontaneous reactions increase overall disorder.
  • The Third Law of Thermodynamics asserts that the entropy of pure crystalline substances approaches zero as temperature approaches absolute zero (0 Kelvin), applicable only to pure crystals.
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