Thermochemistry and Calorimetry Concepts
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Questions and Answers

What are the ions present in the reaction involving Ba²⁺ and SO₄²⁻?

The ions present are Ba²⁺, NO₃⁻, Na⁺, and SO₄²⁻.

What is the net reaction when Ba²⁺ reacts with SO₄²⁻?

The net reaction is Ba²⁺ (aq) + SO₄²⁻ (aq) = BaSO₄ (s).

How is the heat transferred calculated in this calorimetry experiment?

The heat transferred, qH2O, is calculated using the formula qH2O = m x Cₛ x ΔT.

What does the negative sign in the enthalpy change indicate in the reaction?

<p>The negative sign indicates that the reaction is exothermic, meaning heat is released.</p> Signup and view all the answers

What defines a bomb calorimeter?

<p>A bomb calorimeter is a rigid steel container used to measure energy changes by the temperature increase of water.</p> Signup and view all the answers

Explain why no work is done in constant-volume calorimetry.

<p>No work is done because the volume of the system remains constant, resulting in ΔV = 0.</p> Signup and view all the answers

How is the change in enthalpy (ΔH) related to energy at constant pressure?

<p>The change in enthalpy (ΔH) is equal to the heat released or absorbed at constant pressure.</p> Signup and view all the answers

What does a positive entropy change (, \Delta S > 0) indicate about a system?

<p>A positive entropy change indicates an increase in disorder within the system.</p> Signup and view all the answers

According to the second law of thermodynamics, what happens to the entropy of the universe?

<p>The entropy of the universe always increases in spontaneous processes.</p> Signup and view all the answers

Explain the significance of the equation ( \Delta S = S_{final} - S_{initial} ).

<p>The equation signifies the change in entropy, where a positive result indicates an increase in disorder and a negative result indicates a decrease.</p> Signup and view all the answers

Identify an example of a process that results in a negative entropy change and explain why.

<p>The process of water freezing into ice results in a negative entropy change because it transitions from a higher entropy state (liquid) to a lower entropy state (solid).</p> Signup and view all the answers

How does the dissolution of NH4Cl in water affect the system's entropy?

<p>The dissolution of NH4Cl in water results in a positive entropy change, indicating an increase in disorder.</p> Signup and view all the answers

What is the standard enthalpy change of formation (ΔHf) for CO(g)?

<p>-110 kJ mol-1</p> Signup and view all the answers

Calculate the standard enthalpy change of formation (ΔHf) for C4H10(g) using the provided equation.

<p>-125 kJ mol-1</p> Signup and view all the answers

How is the standard enthalpy change of combustion (ΔHc) defined?

<p>It is the enthalpy change when one mole of a substance undergoes complete combustion with oxygen under standard conditions.</p> Signup and view all the answers

What is the significance of the enthalpy level diagram in relation to the oxidation of C(graphite) to CO2(g)?

<p>It visually represents the changes in enthalpy as substances transform from reactants to products.</p> Signup and view all the answers

Identify the enthalpy changes involved when converting 4C(graphite) and 5H2(g) to C4H10(g).

<p>The enthalpy changes include the combustion enthalpies of graphite and hydrogen and the formation enthalpy of C4H10(g).</p> Signup and view all the answers

Explain the difference between ΔHf and ΔHc.

<p>ΔHf refers to the formation of a compound from its elements, while ΔHc refers to the combustion of a substance.</p> Signup and view all the answers

What are the arbitrary enthalpy values assigned to elements in the enthalpy level diagram?

<p>The enthalpy values of elements are taken as zero.</p> Signup and view all the answers

How does the enthalpy of formation for CO(g) relate to the combustion of carbon?

<p>The enthalpy of formation for CO(g) is derived from the combustion enthalpy of carbon and related reactions.</p> Signup and view all the answers

Why is it essential to understand the standard enthalpy changes in combustion and formation reactions?

<p>Understanding these changes is crucial for predicting reaction spontaneity and energy requirements.</p> Signup and view all the answers

What is the standard enthalpy change of solution ($ abla H_{soln}$) for NaCl?

<p>+4.98 kJ mol-1</p> Signup and view all the answers

What is the bond dissociation energy of the Cl-Cl bond?

<p>242 kJ mol-1</p> Signup and view all the answers

Is the bond dissociation process endothermic or exothermic, and why?

<p>Endothermic, because bond breaking requires energy input.</p> Signup and view all the answers

How does the enthalpy change of solution for LiCl compare to that of NaCl?

<p>LiCl has a $ abla H_{soln}$ of -37.2 kJ mol-1, indicating it releases energy, unlike NaCl.</p> Signup and view all the answers

What does a positive value for $ abla H_{soln}$ indicate about the dissolution of a solute?

<p>It indicates that the dissolution is endothermic.</p> Signup and view all the answers

State the bond dissociation energy of the O=O bond.

<p>498 kJ mol-1</p> Signup and view all the answers

Why are bond dissociation enthalpy values greater than zero?

<p>Because breaking a bond requires an input of energy.</p> Signup and view all the answers

What happens to the enthalpy of NaCl solution when it is diluted?

<p>The enthalpy change remains positive, meaning it continues to absorb heat.</p> Signup and view all the answers

Identify the type of reaction represented by the dissolution of LiCl in water.

<p>Exothermic reaction.</p> Signup and view all the answers

What trend is observed in bond dissociation energies as the size of the halogen increases?

<p>Bond dissociation energies decrease.</p> Signup and view all the answers

Calculate the enthalpy change for the given reaction: CH4(g) + H2O(g) → CO(g) + 3H2(g). What does a positive ΔHreaction indicate?

<p>ΔHreaction = 198 kJ. A positive ΔHreaction indicates that the reaction is endothermic.</p> Signup and view all the answers

What bond enthalpy values are used to calculate the ΔHreaction for iodoethane and water vapor?

<p>The bond enthalpies used are C–I (238 kJ/mol), O–H (463 kJ/mol), C–O (360 kJ/mol), and H–I (299 kJ/mol).</p> Signup and view all the answers

In the reaction 2O3(g) → 3O2(g), what is the calculated ΔHreaction, and what does this value indicate about the reaction?

<p>ΔHreaction = -206 kJ. This negative value indicates that the reaction is exothermic.</p> Signup and view all the answers

Why is exothermicity not the only reason for the spontaneity of a reaction?

<p>Some spontaneous processes are endothermic, meaning they absorb energy without requiring additional heat once initiated.</p> Signup and view all the answers

What is the significance of activation energy (Ea) in spontaneous reactions?

<p>Activation energy is the initial energy needed to start a reaction, even a spontaneous one, before it can proceed naturally.</p> Signup and view all the answers

Describe an example of a spontaneous physical change and a spontaneous chemical change.

<p>An example of a spontaneous physical change is the condensation of steam at 25°C; a chemical change is the burning of wood once ignited.</p> Signup and view all the answers

What is the total bond enthalpy change when breaking bonds in the reaction: N–N, N≡N, 4N–H, and 2F–F?

<p>Total bond enthalpy change for breaking is 163 + 945 + 4<em>390 + 2</em>158 = 2613 kJ.</p> Signup and view all the answers

What are mean bond enthalpies, and why are they used in reaction calculations?

<p>Mean bond enthalpies are average values used to estimate the energy required to break chemical bonds, simplifying complicated calculations.</p> Signup and view all the answers

In terms of spontaneity, explain why the melting of ice is not classified as spontaneous.

<p>Melting of ice requires heat input to transition from solid to liquid and does not occur naturally at temperatures below 0°C.</p> Signup and view all the answers

State the relationship between bond formation and energy changes in a chemical reaction.

<p>The formation of bonds releases energy, while breaking bonds requires energy input.</p> Signup and view all the answers

Study Notes

Introduction to Thermodynamics

  • This is a course on thermodynamics (CHEM 217) at the University of Ghana, Department of Chemistry.

System, Boundary, Surroundings

  • A system is the part of the universe under study.
  • Surroundings comprise the rest of the universe.
  • The real or imaginary surface separating the system from its surroundings is the boundary.

Types of Thermodynamic Systems

  • Isolated system: Cannot transfer matter or energy with the surroundings.
  • Closed system: Can transfer energy (heat, work, radiation) but not matter.
  • Open system: Can transfer both energy and matter.

Homogeneous and Heterogeneous Systems

  • Homogeneous system: Uniform throughout, consisting of a single phase (e.g., pure solid, liquid, or gas).
  • Heterogeneous system: Consists of two or more physically distinct phases (e.g., ice in contact with water).
  • A phase is a homogeneous, physically distinct, and mechanically separable portion of a system.

Intensive and Extensive Properties

  • Intensive properties: Do not depend on the amount of matter (e.g., pressure, temperature, density, concentration).
  • Extensive properties: Depend on the amount of matter (e.g., volume, number of moles, enthalpy, entropy, Gibbs' free energy).

State of a System

  • A thermodynamic system is in a state when all its properties are fixed.
  • Fundamental macroscopic properties determining a system's state: pressure (P), temperature (T), volume (V), mass, and composition.

Thermodynamics

  • Study of energy flow in physical or chemical transformations.
  • First law of thermodynamics: Energy of the universe is constant (conservation of energy).

Thermochemistry

  • Study of heat changes associated with chemical reactions.

Work and Energy

  • Energy is the capacity to do work.
  • Internal energy (U or E) is the total energy stored in a system (cannot be measured absolutely).
  • Work is the transfer of energy to a system by a process equivalent to raising or lowering a weight. (work done on a system is positive, work done by a system is negative).
  • When energy is transferred to a system by doing work, with no other change occurring, ∆U = w.

The First Law of Thermodynamics

  • The total internal energy of an isolated system is constant.
  • In practice, we only measure the change in internal energy.
  • ΔE = Efinal - Einitial = Eproducts - Ereactants .
  • ΔEsystem + ΔEsurroundings = 0

Work

  • Work required to move an object against an opposing force is calculated by multiplying the opposing force by the distance moved against it.
  • Work = opposing force x distance moved
  • Unit is the joule (J), where 1 J = 1 kg.m²/s².
  • Types of work associated with chemical processes are work done by a gas through expansion or compression and motion of a car.

Deriving the Equation for Work

  • Work done by a gas expanding against an external pressure is w = -PΔV.
  • Work done on a gas (compression) is +PΔV.
  • Change in volume (ΔV) is the product of the area of the piston times the distance it moves.

Calculating the work of reversible isothermal expansion of a gas

  • The work done in a reversible isothermal expansion is w = -nRTln(Vfinal/Vinitial).

Examples (various calculations)

  • Various examples are provided demonstrating the application of the equations mentioned.

Enthalpy / Heat

  • Enthalpy (H): Energy transferred due to a temperature difference.
  • Energy flows from a high-temperature region to a low-temperature region.
  • Heat (q): Energy transfer as a result of a temperature difference. Heat transfer into a system is positive. Heat transfer out of a system is negative.
  • When no other changes except heat transfer occur. ΔΕ =q

Measuring Heat; Calorimetry

  • Science of measuring heat using observations of temperature change
  • Calorimeter is a device for determining heat associated with a chemical reaction.

Heat Capacity

  • Heat capacity (C): The energy required to raise the temperature of a substance by one unit.
  • Units: J/°C or J/K
  • Specific heat capacity: Energy required to raise the temperature of one gram by one degree Celsius,
  • Molr heat capacity : Energy required to raise the temperature of one mole of a substance by one degree Celsius
  • Heat capacities of metals differ from that of water.

Constant-Pressure Calorimetry

  • Measuring enthalpy changes of reactions in solution.
  • At constant pressure, the heat of reaction, q, equals enthalpy change, ∆H: qp=∆H.

Constant-Volume Calorimetry

  • Measuring enthalpy changes in reactions that occur in a rigid or fixed-volume container.
  • For reactions occurring at constant volume q = ΔE
  • Bomb calorimeters are devices used to measure these reactions.

Heat Transfers at Constant Pressure

  • Enthalpy (H) is a state function.
  • Change in enthalpy (ΔH) is equal to heat gained or released at constant pressure.
  • Every chemical reaction involves energy changes.

Determining a reaction enthalpy from experimental data

  • Sample problems are provided demonstrating procedures and formulas to determine enthalpy values when certain parameters are given. (using coffee-cup calorimetry)

Standard Enthalpies of Formation

  • Standard reaction enthalpy per mole of formula units of a substance created from its elements in the most stable form.
  • Elements' standard enthalpy of formation is zero.

Combining Reaction Enthalpies: Hess' Law

  • The overall enthalpy change of a reaction is the sum of the enthalpy changes of the steps into which the reaction can be divided.
  • Enthalpy and reaction enthalpy values are independent of the path.

Characteristics of Enthalpy Changes and Hess Law

  • Key characteristics and rules to perform operations using Hess Law are discussed.

Using Hess's Law

  • Procedures are shown demonstrating how to utilize Hess' law to determine enthalpy values of specific reactions.

Bond Dissociation Enthalpy/Energy

  • Enthalpy change for breaking one mole of a bond under standard conditions (gaseous state).
  • Bond enthalpy values differ across environments or compounds.

Bond Enthalpy/Energy

  • Average of bond dissociation enthalpies for a specific bond.
  • Provides estimations for enthalpy changes.

Estimation of B.E. from ΔH and ΔHatm

  • Procedure for estimating bond energy from known enthalpy values using Hess' Law.

Example Problems (various calculations)

  • Various example calculations are given, demonstrating procedures and formulas for determining enthalpy values.

Entropy and Disorder

  • Entropy (S) is a measure of randomness or disorder.
  • Entropy of an isolated system increases in a spontaneous process.
  • Entropy is a state function.
  • Entropy changes (ΔS) are the change in degree of disorder ΔS = Sfinal - Sinitial.

Calculating the change in entropy due to an increase in temperature

  • Formulas are provided for calculating the entropy change due to temperature change in constant volume.

Calculating the change in entropy when an ideal gas expands isothermally

  • How to calculate the entropy change when an ideal gas expands isothermally.

Calculating Entropy Changes When Both Temperature and Volume Change

  • How to determine entropy change when both temperature and volume changes occur.

Standard Molar Entropies

  • Calculate the standard molar entropies for a material.

Calculating the standard reaction entropy

  • Determine the overall reaction entropy.

Global Changes in Entropy

  • How to determine the total entropy change.

The Overall Change in Entropy

  • Calculate the overall change in entropy.

Calculating the Total Entropy Change for the Expansion of an Ideal Gas

  • A procedure for calculating total entropy change when an ideal gas expands.

Assessing Spontaneity

  • A procedure using enthalpy and entropy to assess whether a reaction is spontaneous under standard conditions.

Equilibrium

  • Equilibrium is the state where there's no net change in the composition of the reaction mixture.

Focusing on the System

  • Using Gibbs Free Energy to determine spontaneity for a reaction.

Summary of factors that affect spontaneity

  • Summary of factors, such as enthalpy, entropy, and temperature that affect spontaneity.

Deciding Whether a Process is Spontaneous

  • Procedure for determining the Gibbs free energy change for a reaction and concluding spontaneity based on the ΔG value.

Gibbs Free Energy of Reaction

  • The Gibbs Free energy of reaction is the standard Gibbs free energy of reaction per mole for the formation of a compound formed from its elements in their most stable form. (G°).

Calculating a Standard Gibbs Free Energy of Formation from Enthalpy and Entropy Data

  • How to calculate standard Gibbs free energy of formation.

The Dependence of Free Energy on Pressure

  • How pressure affects the free energy of a chemical reaction (introducing the reaction quotient Q).

Calculating Equilibrium Constant

  • How to calculate the equilibrium constant given a standard Gibbs Free Energy changes.

Criteria for Spontaneity in a Chemical Reaction

  • Summary of spontaneous processes, equilibrium, and Non-spontaneous processes and their associated conditions.

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Description

This quiz tests your understanding of key concepts in thermochemistry and calorimetry, including ionic reactions, enthalpy changes, and the second law of thermodynamics. Questions cover the role of heat transfer, the significance of entropy changes, and the use of bomb calorimeters. Prepare to deepen your grasp of crucial principles in energy chemistry!

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