Chapter 7 of the first-year chemistry course, titled “Thermochemistry,” delves into the fascinating world of energy changes in chemical reactions. In this crucial chapter, students are introduced to the fundamental principles of thermodynamics and how they apply to chemical systems. Thermochemistry explores concepts such as heat, enthalpy, and internal energy, unraveling the mysteries of how energy is transferred during chemical processes.
Students will learn to calculate and interpret energy changes in various reactions, whether they are exothermic (releasing heat) or endothermic (absorbing heat). Additionally, the chapter delves into the concept of calorimetry, providing the tools necessary to measure heat changes experimentally. A solid understanding of thermochemistry is vital for comprehending the behavior of matter in different chemical scenarios, making this chapter a crucial building block in the study of chemistry.
Short Questions Chemistry 1st Year Chapter 7 Thermochemistry
What is the result of a chemical change in terms of energy?
In a chemical change, energy in the form of heat is either evolved or absorbed. This is due to the breaking of bonds in the reactants and the formation of new bonds in the products.
What is the overall energy change in a chemical reaction called?
The overall energy change in a chemical reaction, resulting from the difference between the energy supplied for breaking reactant bonds and that evolved in making product bonds, is known as the heat of reaction.
What is the difference between exothermic and endothermic reactions?
Exothermic reactions release heat into the surroundings, causing the system’s temperature to rise. Endothermic reactions absorb heat from the surroundings, causing the system’s temperature to drop initially.
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What are the units commonly used to express heat changes in thermochemistry?
Heat changes in thermochemistry are commonly expressed in joules (J) and kilojoules (kJ) in the SI system.
What is a spontaneous process in chemistry?
A spontaneous process is a process that takes place on its own without any outside assistance, moves from a non-equilibrium state towards an equilibrium state, and is unidirectional and irreversible.
Provide an example of a spontaneous process.
Water flowing from a higher level to a lower level is an example of a spontaneous process.
What distinguishes a non-spontaneous process from a spontaneous one, and can non-spontaneous processes occur with external energy input?
A non-spontaneous process does not occur on its own and does not occur in nature. However, some non-spontaneous processes can be made to occur by supplying energy to the system from an external source.
How does the concept of free energy help predict whether a reaction will occur spontaneously or not?
The concept of free energy helps predict whether a reaction will occur spontaneously by considering the change in free energy of the system. A negative change in free energy indicates a spontaneous reaction, but a discussion of free energy and its role in predicting spontaneity is beyond the scope of this book.
What is a system in the context of thermochemistry?
A system is any portion of the universe under study, typically materials being tested in a laboratory or considered for analysis in a classroom discussion.
How are surroundings defined in thermochemistry?
Surroundings refer to the remaining portion of the universe that is not part of the system being studied.
What is the boundary in the context of thermodynamics?
The boundary is the real or imaginary surface that separates the system from its surroundings.
Give an example of a system and its surroundings.
One mole of oxygen confined in a cylinder with a piston is a system, while the cylinder, piston, and everything outside the cylinder are the surroundings. Similarly, a cup of water is a system, and the air around it and the table it’s on are the surroundings.
How is the state of a system defined in thermodynamics?
The state of a system refers to its condition, including properties like temperature, volume, and pressure.
Describe the initial state and final state of a system.
The initial state of a system is its condition before any process or change, while the final state is its condition after the change.
What is the formula to calculate the change in temperature (ΔT) of a system?
ΔT = Final temperature – Initial temperature
What is a state function in thermodynamics, and give some examples?
A state function is a macroscopic property of a system that has specific values for both initial and final states and is independent of the path taken to bring about a change. Examples include pressure (P), temperature (T), volume (V), internal energy (E), and enthalpy (H).
How can you calculate the change in volume (ΔV) of a gas, and why is it considered a state function?
ΔV is the difference between the initial volume (V1) and the final volume (V2) of a gas. It is considered a state function because it depends only on the initial and final volumes and is independent of how the volume change was achieved, whether through changes in temperature or pressure.
What is internal energy in a thermodynamic system?
Internal energy is the sum of kinetic and potential energies of the particles within a system, including translational, rotational, and vibrational energies, as well as attractive forces such as bonds and van der Waal’s forces.
Is the absolute value of internal energy measurable for a system? If not, what can be measured?
The absolute value of internal energy cannot be measured directly, but the change in internal energy (ΔE) during a change in the state of the system can be measured.
What are the two fundamental ways of transferring energy to or from a system?
The two fundamental ways of transferring energy to or from a system are heat and work.
Is heat a state function, and how is it defined?
Heat is not a state function. It is defined as the quantity of energy that flows across the boundary of a system during a change in its state due to a temperature difference between the system and its surroundings.
What is work in thermodynamics, and how is it measured?
Work in thermodynamics is the product of force and distance (W = F x S). It is measured in Joules in SI units.
What is pressure-volume work, and how is it calculated?
Pressure-volume work occurs when a system undergoes expansion or compression. It is calculated as W = -PΔV, where P is the external pressure and ΔV is the change in volume.
According to the first law of thermodynamics, what is the law of conservation of energy?
The first law of thermodynamics, also known as the law of conservation of energy, states that energy cannot be created or destroyed but can be changed from one form to another. The total energy of a system and its surroundings remains constant.
How is the change in internal energy (ΔE) related to heat (q) and work (w)?
The change in internal energy (ΔE) of a system is equal to the sum of heat (q) absorbed by the system and work (w) done by the system: ΔE = q + w.
When does the equation for the change in internal energy simplify to ΔE = qv?
The equation ΔE = qv applies when the volume of the gas is kept constant, and no pressure-volume work is done (ΔV = 0).
What is enthalpy, and how is it represented mathematically?
Enthalpy, represented as H, is a thermodynamic property of a system that accounts for both its internal energy (E) and the product of pressure and volume (PV). Mathematically, it is expressed as H = E + PV.
Is enthalpy a state function, and what are its units of measurement?
Yes, enthalpy is a state function. It is measured in joules (J).
How can you measure the change in enthalpy (ΔH) for a system?
ΔH, the change in enthalpy, can be measured for a change in the state of a system. It is calculated using the formula ΔH = ΔE + PΔV, where ΔE is the change in internal energy, and PΔV represents the work done by the system when it expands against a constant pressure.
Under what conditions is ΔH approximately equal to ΔE?
ΔH is approximately equal to ΔE when there is no significant volume change in the process, as is the case with liquids and solids (ΔV ≈ 0).
What is the significance of ΔH in thermodynamics, and why is it more convenient to work with than ΔE?
ΔH represents the change in enthalpy, which is equal to the heat of reaction at constant pressure. Reactions are often carried out at constant pressure, making ΔH more convenient to work with than ΔE.
What is the standard enthalpy of a reaction (ΔHo)?
The standard enthalpy of a reaction (ΔHo) is the enthalpy change that occurs when a certain number of moles of reactants, as indicated by the balanced chemical equation, react together completely to give the products under standard conditions (25°C and one atmosphere pressure). It is expressed in kJ/mol.
What is the standard enthalpy of formation (ΔH0f)?
The standard enthalpy of formation (ΔH0f) of a compound is the amount of heat absorbed or evolved when one mole of the compound is formed from its elements under standard conditions (25°C and one atmosphere pressure). It is also expressed in kJ/mol.
What is the standard enthalpy of atomization (ΔHoat)?
The standard enthalpy of atomization (ΔHoat) of an element is the amount of heat absorbed when one mole of gaseous atoms is formed from the element under standard conditions (25°C and one atmosphere pressure). It is given in kJ/mol.
What is the standard enthalpy of neutralization (ΔHon)?
The standard enthalpy of neutralization (ΔHon) is the amount of heat evolved when one mole of hydrogen ions (H+) from an acid reacts with one mole of hydroxide ions (OH-) from a base to form one mole of water. For strong acids and bases, this value is approximately -57.4 kJ/mol.
What is the definition of the standard enthalpy of combustion (ΔH°c)?
The standard enthalpy of combustion of a substance is the amount of heat evolved when one mole of the substance is completely burnt in excess of oxygen under standard conditions. It is denoted by ΔH°c.
What is the standard enthalpy of combustion for ethanol (C2H5OH)?
The standard enthalpy of combustion of ethanol (C2H5OH) is -1368 kJ mol⁻¹.
Define the enthalpy of solution (ΔH°sol.).
The standard enthalpy of a solution is the amount of heat absorbed or evolved when one mole of a substance is dissolved in so much solvent that further dilution results in no detectable heat change.
What is the enthalpy of solution (ΔH°sol.) for ammonium chloride?
The enthalpy of solution (ΔH°sol.) of ammonium chloride is +16.2 kJ mol⁻¹, indicating an endothermic process.
Explain how to determine the enthalpy change of a reaction using a glass calorimeter.
In a glass calorimeter, the enthalpy change of a reaction can be determined by measuring the temperature change (ΔT) before and after the reaction. The heat evolved or absorbed during the reaction can be calculated using the formula q = m x s x ΔT, where q is the heat, m is the mass of reactants, s is the specific heat of the reaction mixture, and ΔT is the change in temperature.
Why can’t ΔH be measured directly for some compounds like tetrachloromethane (CCl4)?
ΔH cannot be measured directly for some compounds like CCl4 because they cannot be prepared directly by combining their constituent elements, making it difficult to determine their standard enthalpies of formation.
How does Hess’s law of constant heat summation apply to the determination of enthalpy changes?
Hess’s law states that if a chemical change occurs through several different routes, the overall energy change is the same, provided the initial and final conditions are the same. This principle allows chemists to indirectly calculate enthalpy changes by adding or subtracting the enthalpies of other related reactions.
How can you calculate the enthalpy change for the formation of CO(g) using Hess’s law?
By knowing the enthalpy of combustion for graphite to CO2 and the enthalpy of combustion of CO to CO2, you can determine the enthalpy of formation for CO using Hess’s law. In this case, the enthalpy change for the formation of CO(g) is -110.0 kJ/mol.
What is the specific application of Hess’s law in the Born-Haber cycle?
Hess’s law finds its best application in the Born-Haber cycle, which is used to indirectly calculate lattice energies of ionic compounds. The Born-Haber cycle allows for the determination of lattice energies by considering various stages of ionization and atomization.
How is the lattice energy of sodium chloride (NaCl) determined using the Born-Haber cycle?
The lattice energy of NaCl is determined by considering the atomization and ionization energies of sodium and chlorine and their conversion into gaseous ions. The lattice energy can be calculated using the Born-Haber cycle as the difference between the standard enthalpy of formation of NaCl and the total energy change involved in the ionization and atomization processes. In the case of NaCl, the lattice energy is found to be -787 kJ/mol.
MCQ’s Chemistry 1st Year Chapter 7 Thermochemistry
Long Questions Chemistry 1st Year Chapter 7 Thermochemistry
Question: Differentiate between the following:
(i) Internal energy and enthalpy
Internal Energy
Internal energy (U) is a state function that represents the total energy contained within a system.
It includes the kinetic energy (energy due to motion) and potential energy (energy due to position) of particles within the system.
Internal energy is independent of pressure and volume changes; it depends solely on the state of the system.
Enthalpy
Enthalpy (H) is also a state function that accounts for the total heat content of a system.
Enthalpy includes internal energy (U) and the product of pressure and volume (PV).
Enthalpy is particularly useful in studying processes at constant pressure, as enthalpy change (ΔH) directly relates to heat exchange at constant pressure.
(ii) Internal energy change and enthalpy change
Internal Energy Change (ΔU)
ΔU represents the change in internal energy of a system during a process.
It is a measure of the heat exchanged with the surroundings and the work done on or by the system.
ΔU = q (heat added to the system) – w (work done by the system).
Enthalpy Change (ΔH)
ΔH represents the change in enthalpy of a system during a process, often at constant pressure.
It accounts for both heat exchanged and work done, but it simplifies calculations for systems under constant pressure conditions.
ΔH = ΔU + PΔV, where P is the constant pressure and ΔV is the change in volume.
(iii) Exothermic and endothermic reactions
Exothermic Reaction
An exothermic reaction releases heat energy to the surroundings.
The system loses internal energy, resulting in a negative ΔH (ΔH < 0).
Examples include combustion reactions, where fuels burn to produce heat.
Endothermic Reaction
An endothermic reaction absorbs heat energy from the surroundings.
The system gains internal energy, resulting in a positive ΔH (ΔH > 0).
Examples include the process of photosynthesis in plants, where energy from the sun is absorbed to convert carbon dioxide and water into glucose.
Question: What are spontaneous and non-spontaneous processes. Give examples.
Answer: Spontaneous Processes
A spontaneous process is a physical or chemical change that occurs naturally without any external intervention or influence. These processes tend to proceed in a particular direction without requiring additional energy input. In spontaneous processes, the entropy of the system generally increases, and they tend to move towards a state of higher disorder or randomness.
Examples of spontaneous processes
Ice melting at room temperature: When you take a block of ice out of the freezer and leave it at room temperature, it naturally melts into liquid water. This process is spontaneous because it leads to an increase in entropy.
Dissolving of salt in water: When you add table salt (sodium chloride) to water, it dissolves spontaneously, forming a solution. This process is driven by the tendency of particles to mix and disperse, increasing entropy.
Chemical reactions like rusting: The corrosion of iron (rusting) is a spontaneous chemical process that occurs in the presence of oxygen and moisture. Iron gradually turns into iron oxide (rust) over time.
Non-Spontaneous Processes
A non-spontaneous process is a physical or chemical change that does not occur naturally under given conditions and requires external energy input to proceed. These processes often lead to a decrease in entropy or a move toward a more ordered state.
Examples of non-spontaneous processes
Water freezing at room temperature: To freeze water at room temperature, you need to remove heat energy from it. Water does not naturally freeze under these conditions.
Separating a gas mixture into its components: Separating a mixture of gases into its individual components, such as separating oxygen and nitrogen from air, requires energy input in processes like fractional distillation or cryogenic separation.
Charging a battery: The process of charging a battery is non-spontaneous as it requires an external energy source (e.g., electrical energy) to force the chemical reactions in the battery to proceed in the opposite direction.
It’s important to note that whether a process is spontaneous or non-spontaneous depends on the specific conditions and constraints (e.g., temperature, pressure, concentration) in which it occurs. Some processes that are non-spontaneous under certain conditions may become spontaneous under different conditions.
Question: Explain that burning of a candle is a spontaneous process.
Answer: The burning of a candle is a spontaneous process because it occurs naturally without requiring external energy input and is driven by the fundamental principles of thermodynamics. To understand why the burning of a candle is spontaneous, let’s consider the key factors involved:
Release of Energy: When a candle burns, it undergoes a chemical reaction with oxygen (combustion). During this process, chemical bonds in the candle’s wax (composed of hydrocarbons) are broken and new bonds are formed with oxygen molecules. This chemical reaction releases energy in the form of heat and light, producing a flame. The release of energy is a characteristic of spontaneous processes, as they tend to move towards a state of lower potential energy, releasing energy to the surroundings.
Increase in Entropy: One of the fundamental principles of thermodynamics is the second law, which states that in any spontaneous process, the total entropy (a measure of disorder or randomness) of the universe increases. When a candle burns, it leads to an increase in entropy. This is because the solid wax and gaseous oxygen and combustion products (such as carbon dioxide and water vapor) have a higher degree of disorder than the candle wax alone. The random motion of molecules in the gaseous products contributes to a higher entropy.
Favorable Reaction Conditions: The burning of a candle is favored under typical room conditions. There is an ample supply of oxygen in the air, and the heat generated by the initial ignition sustains the combustion reaction. As long as there is a continuous supply of oxygen and fuel (wax), the candle continues to burn spontaneously.
Natural Driving Force: The burning of a candle is a result of the natural tendency of chemical systems to reach a state of lower potential energy. Hydrocarbon molecules in the wax have a higher potential energy due to the bonds between carbon and hydrogen atoms. Combustion allows these molecules to release energy by forming more stable bonds with oxygen.
Question: Is it true that a non-spontaneous process never happens in the universe? Explain it.
Answer: A non-spontaneous process can indeed happen in the universe, but it requires the input of external energy to drive the process. It’s important to understand that whether a process is spontaneous or non-spontaneous depends on the specific conditions and constraints present in a given system. Here’s an explanation:
Spontaneous vs. Non-Spontaneous Processes
Spontaneous processes: These are processes that occur naturally without external intervention. They tend to move in a particular direction without requiring additional energy input. Spontaneous processes often lead to an increase in entropy (disorder) and a decrease in the potential energy of the system.
Non-spontaneous processes: These are processes that do not occur naturally under given conditions and require external energy input to proceed. Non-spontaneous processes often lead to a decrease in entropy or a move toward a more ordered state.
External Energy Input
Non-spontaneous processes can be made to happen by supplying the necessary external energy. This energy input can come from various sources, such as mechanical work, electrical energy, heat, or other forms of energy.
For example, freezing water at room temperature is a non-spontaneous process because it requires removing heat energy from the water. However, if you place the water in a freezer (which extracts heat), the process becomes feasible and occurs as a result of the external energy input provided by the freezer.
Energy Conservation
The fundamental principle of energy conservation (the first law of thermodynamics) states that energy cannot be created or destroyed; it can only be converted from one form to another.
When a non-spontaneous process occurs, the external energy input compensates for the energy required to drive the process against the natural thermodynamic tendencies.
Question: What is the first law of thermodynamics. How does it explain that (i) qv=ΔE (ii) qp =ΔH
Answer: The first law of thermodynamics, also known as the law of energy conservation, states that energy cannot be created or destroyed in an isolated system. Instead, it can only change forms or be transferred between the system and its surroundings. This fundamental principle can be expressed mathematically as:
ΔE = q – w
Where:
ΔE represents the change in the internal energy of the system.
q represents the heat added to the system.
w represents the work done by the system on its surroundings.
Now, let’s see how the first law of thermodynamics explains the given expressions:
(i) qv = ΔE
In this equation, qv represents the heat added to a system at constant volume (v).
ΔE represents the change in the internal energy of the system.
The first law states that the change in internal energy (ΔE) of a system is equal to the heat added to the system (q) minus the work done by the system (w).
When the system is at constant volume (v), no work is done (w = 0), so the equation simplifies to ΔE = qv. This means that the change in internal energy is equal to the heat added when no work is done.
(ii) qp = ΔH
In this equation, qp represents the heat added to a system at constant pressure (p).
ΔH represents the change in enthalpy of the system.
Enthalpy (H) is defined as H = E + PV, where E is the internal energy, P is the pressure, and V is the volume.
The first law, ΔE = q – w, can be rearranged as ΔE = q + w (since work done by the system is negative). If we consider work done at constant pressure (w = -PΔV), then ΔE = q – PΔV.
Now, we introduce enthalpy, H = E + PV, and find the change in enthalpy, ΔH = ΔE + Δ(PV). Substituting the previous equation into this one, we get ΔH = q – PΔV + Δ(PV).
At constant pressure (p), Δ(PV) simplifies to PΔV, and the equation becomes ΔH = q – PΔV + PΔV.
The PΔV terms cancel out, leaving us with ΔH = q at constant pressure (qp = ΔH). This means that, under constant pressure conditions, the change in enthalpy is equal to the heat added to the system.
Question: How will you differentiate between ΔE and ΔH? Is it true that ΔH and ΔE have the same values for the reactions taking place in the solution state.
Answer: ΔE and ΔH are related thermodynamic quantities, but they represent different aspects of a chemical or physical process:
ΔE (Change in Internal Energy):
ΔE represents the change in the internal energy of a system during a chemical or physical process.
It takes into account the total energy stored within the system, including both kinetic and potential energy of particles.
ΔE is generally used to describe processes that occur at constant volume (qv) or in cases where work done by or on the system is negligible.
ΔH (Change in Enthalpy):
ΔH represents the change in enthalpy of a system during a chemical or physical process.
Enthalpy (H) is defined as H = E + PV, where E is the internal energy, P is pressure, and V is volume.
ΔH takes into account both the internal energy change (ΔE) and the work done against or by the surroundings due to changes in volume.
ΔH is typically used to describe processes that occur at constant pressure (qp).
Regarding whether ΔH and ΔE have the same values for reactions taking place in the solution state, it’s essential to understand that these values can be different. The key factor that determines the difference between ΔH and ΔE is the presence or absence of work done on or by the system during the process.
In many cases, when reactions occur in the solution state at constant pressure (such as reactions in open containers), ΔH and ΔE may indeed have approximately the same values because any work done against or by the surroundings due to volume changes is minimal. However, there can be exceptions, especially in cases where changes in volume are significant or where other forms of work are involved.
In summary, while ΔE and ΔH are related thermodynamic quantities, they represent different aspects of energy changes during a process. Whether they have the same values for reactions in the solution state depends on the specific conditions and constraints of the process. For processes at constant pressure and in open containers, ΔH and ΔE may often be very close in value.
Question: What is the difference between heat and temperature? Write a mathematical relationship between these two parameters.
Answer: Heat and temperature are related but distinct concepts in thermodynamics:
Heat
Heat (q) is a form of energy transfer between two objects or systems due to a temperature difference.
It represents the energy that flows from a higher-temperature object or system to a lower-temperature one.
Heat is measured in units of energy, such as joules (J) or calories (cal).
It is a scalar quantity and does not have direction.
Temperature:
Temperature (T) is a measure of the average kinetic energy of the particles (atoms or molecules) in a substance or system.
It provides information about the hotness or coldness of an object or system.
Temperature is measured in units such as degrees Celsius (°C), Kelvin (K), or degrees Fahrenheit (°F).
Temperature is a scalar quantity, and it does not involve the transfer of energy.
Mathematical Relationship between Heat and Temperature
The mathematical relationship between heat and temperature is described by the following equation:
q = mcΔT
Where:
q represents the heat energy transferred (in joules or calories).
m is the mass of the substance involved (in kilograms or grams).
c is the specific heat capacity of the substance (in J/g°C or cal/g°C).
ΔT represents the change in temperature (in degrees Celsius or Kelvin).
This equation shows that the amount of heat transferred (q) is directly proportional to:
The mass of the substance (m): More massive objects require more heat to change their temperature.
The specific heat capacity (c) of the substance: Different substances require different amounts of heat to change their temperature. Specific heat capacity is a measure of how much heat energy is needed to raise the temperature of a given mass of a substance by a certain amount.
The change in temperature (ΔT): Larger temperature differences require more heat to achieve the same change in temperature.
Question: How do you measure the heat of combustion of a substance by bomb calorimeter.
Answer: Measuring the heat of combustion of a substance using a bomb calorimeter is a common laboratory technique in thermodynamics and chemistry. The bomb calorimeter is a device designed to measure the heat released (or absorbed) during a combustion reaction with high precision. Here’s a step-by-step guide on how to measure the heat of combustion using a bomb calorimeter:
Materials and Equipment Needed
- Bomb calorimeter (consisting of a bomb, ignition system, calorimeter chamber, and thermometer).
- Oxygen supply (typically provided by a high-pressure oxygen cylinder).
- Sample of the substance whose heat of combustion you want to measure.
- Weighing balance.
- Ignition system (usually an electric ignition system).
- Thermometer and temperature recording device (digital thermometer or calorimeter thermometer).
- Water and a calorimeter vessel.
- Heat insulation (e.g., insulating materials or a calorimeter jacket).
Procedure
Calibrate the Calorimeter
Fill the calorimeter vessel with a known quantity of water (measured in grams or milliliters) and record its initial temperature (T_initial).
Install the thermometer in the calorimeter chamber and connect it to a temperature recording device.
Prepare the Sample
Weigh the sample of the substance whose heat of combustion you want to measure accurately. The mass should be sufficient to produce a measurable temperature change in the calorimeter.
Record the mass of the sample (m_sample).
Assemble the Bomb:
Load the weighed sample into the bomb calorimeter. If necessary, add any required reagents (e.g., an ignition source) for the combustion reaction.
Seal the bomb tightly to ensure that no oxygen escapes during the reaction.
Charge the Bomb with Oxygen
Connect the bomb to the oxygen supply source and fill it with oxygen at high pressure. Ensure that the bomb is pressurized to the specified level.
Ignite the Sample
Use the ignition system to initiate the combustion reaction inside the bomb. This typically involves passing an electric current through a resistance wire, which ignites the sample.
Monitor the Temperature Change
As the combustion reaction proceeds, it releases heat. This heat is transferred to the water in the calorimeter vessel.
Monitor and record the maximum temperature reached by the water in the calorimeter chamber (T_final) as the reaction progresses.
Calculate the Heat of Combustion
Calculate the heat of combustion (ΔH_combustion) using the calorimeter formula:
ΔH_combustion = (m_water * C_water * ΔT) / n_sample
Where:
m_water is the mass of water in the calorimeter vessel (in grams or milliliters).
C_water is the specific heat capacity of water (usually taken as 4.18 J/g°C or 1 cal/g°C).
ΔT is the temperature change (T_final – T_initial) in degrees Celsius.
n_sample is the number of moles of the sample used in the combustion reaction.
Calculate the Heat of Combustion per Mole
Calculate the heat of combustion per mole of the substance by dividing ΔH_combustion by the number of moles of the sample (n_sample).
Correct for Any Sources of Error
Consider any corrections needed for heat loss to the surroundings or the calorimeter itself.
Report the Result
Express the heat of combustion per mole in units such as joules per mole (J/mol) or calories per mole (cal/mol).
Note: The calorimeter constant, which relates the heat capacity of the calorimeter to the temperature change, is often determined in a calibration step and used in the calculations. Additionally, it’s crucial to follow safety precautions when working with high-pressure oxygen and ignition sources to ensure safe experimentation.
By following this procedure and accurately measuring the temperature change, mass of the sample, and other relevant parameters, you can determine the heat of combustion of the substance with a bomb calorimeter.
Question: Define heat of neutralization. When a dilute solution of a strong acid is neutralized by a dilute solution of a strong base, the heat of neutralization is found to be nearly the same in all the cases. How do you account for this?
Answer: The heat of neutralization, also known as the enthalpy change of neutralization, refers to the heat energy released or absorbed when an acid and a base react to form water and a salt in a neutralization reaction. It’s a fundamental concept in thermodynamics and can be represented by the following general chemical equation:
H⁺(aq) + OH⁻(aq) → H₂O(l) + energy
The key points regarding the heat of neutralization are as follows:
Magnitude of Heat of Neutralization: The heat of neutralization is usually quite consistent and similar for strong acid-strong base reactions. In such cases, it’s typically close to a standard value of approximately -57.3 kJ/mol or -13.7 kcal/mol. This means that for every mole of water formed in the neutralization reaction, approximately 57.3 kJ of heat energy is released.
Reasons for Consistency
Stoichiometry: The heat of neutralization is primarily determined by the stoichiometry of the neutralization reaction, which involves the transfer of one proton (H⁺ ion) from the acid to one hydroxide ion (OH⁻) from the base to form water. This one-to-one stoichiometry results in a consistent energy change for different strong acid-strong base combinations.
Strong Electrolytes: Strong acids and strong bases completely ionize in solution, which means that they exist predominantly as ions. This high degree of ionization ensures that the reactions are highly quantitative and that all the reactants participate in the neutralization.
Negligible Heat Capacity Effects: The heat capacity of the resulting water in the reaction is relatively low compared to the other components of the reaction mixture. Therefore, the heat capacity effects of water are often considered negligible in the calculation of the heat of neutralization.
Enthalpy Change is an Extensive Property: The heat of neutralization is an extensive property, which means it depends on the quantity of reactants involved. Therefore, the heat of neutralization is typically reported per mole of water formed.
In summary, the consistency in the heat of neutralization for dilute solutions of strong acids and strong bases can be attributed to the stoichiometry of the neutralization reaction and the complete ionization of the strong electrolytes. This consistency is a valuable feature in thermodynamic studies, as it allows for the reliable calculation of enthalpy changes in various chemical reactions involving neutralization.
Question: What is lattice energy? How does Born-Haber cycle help to calculate the lattice energy of NaCl?
Answer: Lattice energy is the energy required to completely separate one mole of a solid ionic compound into its individual gaseous ions. It is a measure of the strength of the ionic bonds within the crystal lattice of the compound. Lattice energy is typically expressed in units of energy per mole, such as joules per mole (J/mol) or kilojoules per mole (kJ/mol).
For example, in the case of sodium chloride (NaCl), the lattice energy represents the energy needed to break all the Na⁺ and Cl⁻ ions apart from their arrangement in the solid crystal and convert them into isolated gaseous ions.
The Born-Haber cycle is a thermodynamic cycle used to calculate the lattice energy of an ionic compound, such as NaCl. It involves a series of steps that account for various energy changes associated with the formation of the compound from its constituent elements. Here’s how the Born-Haber cycle helps calculate the lattice energy of NaCl:
Formation of Gaseous Ions
The first step involves the formation of gaseous sodium (Na) and chlorine (Cl) atoms from their respective elements (Na₂(g) → 2Na(g) and Cl₂(g) → 2Cl(g)). This step requires energy equal to the ionization energy of sodium (I(Na)) and the electron affinity of chlorine (EA(Cl)).
ΔH₁ = I(Na) + EA(Cl)
Formation of Solid Sodium Chloride (NaCl)
The next step involves the formation of solid NaCl from its gaseous ions (Na⁺ and Cl⁻). This is the reverse of the lattice energy process, so it releases energy equal to the lattice energy (LE).
ΔH₂ = -LE
Additional Energy Changes
Other energy changes may be involved in the process, such as the sublimation of sodium (Na(s) → Na(g)), which requires energy equal to the sublimation energy of sodium (ΔH_sub).
ΔH₃ = ΔH_sub
Overall Enthalpy Change
The overall enthalpy change for the formation of NaCl can be calculated by summing up the individual energy changes:
ΔH_f = ΔH₁ + ΔH₂ + ΔH₃
To calculate the lattice energy (LE), we rearrange the equation as follows:
LE = -(ΔH_f – ΔH₁ – ΔH₃)
The Born-Haber cycle allows us to indirectly calculate the lattice energy of NaCl based on experimentally determined values for ionization energies, electron affinities, sublimation energies, and the enthalpy of formation of NaCl. The negative sign is used because lattice energy is released when the solid ionic compound forms from its constituent gaseous ions. This cycle is a valuable tool for understanding and calculating the energetics of ionic compounds.