Free flashcard deck
Thermochemistry and Thermodynamics flashcards
149 cards, Medium, made and shared by a student on FORKSAI.
FORKSAI brings each card back with spaced repetition, right before you would forget it. Free to start, no card needed.
All 149 cards
Card 1
What is defined as the 'system' in thermochemistry?
The system is the specific set of substances and energy (e.g., a chemical reaction) that is being studied or observed.
Card 2
What constitutes the 'surroundings' in thermochemistry?
The surroundings include everything that is not directly involved with the system or reaction being studied.
Card 3
In a chemical reaction, which component is typically considered the 'system'?
In chemistry, the reaction itself (reactants and products, and their associated energy) is the primary system of interest.
Card 4
State the First Law of Thermodynamics.
The First Law of Thermodynamics, also known as the Law of Conservation of Energy, states that the total energy of an isolated system remains constant; energy cannot be created or destroyed, only converted from one form into another.
Card 5
What is the mathematical expression for the First Law of Thermodynamics, and what do its terms represent?
The expression is ΔU = Q - W, where: - ΔU (or ΔE) is the change in internal energy of the system. - Q is the heat transferred to or from the system. - W is the work done by or on the system.
Card 6
State the Second Law of Thermodynamics.
The Second Law of Thermodynamics states that any spontaneous process must lead to an increase in the total entropy of the universe (ΔS_universe > 0).
Card 7
How is the Second Law of Thermodynamics conceptually linked to the spontaneity of a process?
It implies that for a process to be spontaneous, the overall disorder or randomness of the universe (system + surroundings) must increase.
Card 8
Which thermodynamic state function is often used to predict spontaneity based on the Second Law, especially at constant temperature and pressure?
Gibbs Free Energy (ΔG) is used to predict spontaneity, as it combines enthalpy and entropy changes of the system to reflect the overall entropy change of the universe.
Card 9
Define a thermodynamic state function.
A thermodynamic state function is a property of a system that depends only on the current state of the system, not on the path taken to reach that state. Examples include internal energy, enthalpy, entropy, and Gibbs free energy.
Card 10
What is Internal Energy (ΔU or ΔE)?
Internal energy is the total energy of a system, comprising the potential vibrational energy of atoms and the kinetic energy of molecules within the system.
Card 11
What are the standard units for Internal Energy (ΔU or ΔE)?
The standard units for internal energy are Joules per mole (J/mol).
Card 12
Define Enthalpy (H).
Enthalpy is a thermodynamic state function that represents the total heat content of a system. It is the sum of the internal energy (E) and the product of its pressure (P) and volume (V).
Card 13
What is the mathematical definition of Enthalpy (H)?
The mathematical definition of enthalpy is H = E + PV.
Card 14
Under what specific conditions does the change in enthalpy (ΔH) equal the heat transferred (q)?
The change in enthalpy (ΔH) equals the heat transferred (q) when the process occurs at constant pressure and constant volume (i.e., ΔH = q_p).
Card 15
How does the sign of ΔH indicate whether a reaction is endothermic or exothermic?
If ΔH > 0, the reaction is endothermic (absorbs heat). If ΔH < 0, the reaction is exothermic (releases heat).
Card 16
What is the formula for calculating the change in enthalpy for a reaction (ΔrH) from the enthalpies of products and reactants?
ΔrH = ΔH_products - ΔH_reactants (assuming standard conditions and using standard enthalpies of formation).
Card 17
What are the standard units for Enthalpy (ΔH)?
The standard units for enthalpy are Joules per mole (J/mol).
Card 18
Define Entropy (ΔS).
Entropy is a state function that describes a system's tendency to disperse energy and matter, leading to greater randomness or disorder.
Card 19
What is the mathematical definition of entropy change (ΔS) in terms of heat and temperature for a reversible process?
ΔS = q_rev / T, where q_rev is the heat transferred in a reversible process and T is the absolute temperature.
Card 20
What are the standard units for Entropy (ΔS)?
The standard units for entropy are Joules per Kelvin (J/K).
Card 21
Define Gibbs Free Energy (ΔG).
Gibbs Free Energy is a thermodynamic potential that measures the 'useful' or process-initiating work obtainable from an isothermal, isobaric thermodynamic system. It indicates the spontaneity of a process.
Card 22
What does Gibbs Free Energy represent in a closed system at constant temperature and pressure?
It represents the maximum amount of free energy that could be extracted from a closed system to do useful work at constant temperature and pressure.
Card 23
How does the sign of ΔG indicate if a reaction is spontaneous (exergonic) or non-spontaneous (endergonic)?
If ΔG 0, the reaction is non-spontaneous (endergonic). If ΔG = 0, the system is at equilibrium.
Card 24
What is the fundamental equation for Gibbs Free Energy change (ΔrG) in terms of enthalpy and entropy?
ΔrG = ΔH_system - TΔrS.
Card 25
What are the standard units for Gibbs Free Energy (ΔG)?
The standard units for Gibbs Free Energy are Joules per mole (J/mol).
Card 26
Conceptually, how is the expression for Gibbs free energy change (ΔrG) derived from the Second Law of Thermodynamics (ΔS_universe > 0)?
The Second Law states ΔS_universe = ΔS_system + ΔS_surroundings > 0. For a process at constant T and P, ΔS_surroundings = -ΔH_system / T. Substituting this into the Second Law gives ΔS_system - ΔH_system / T > 0. Multiplying by -T (and reversing the inequality) yields -TΔS_system + ΔH_system < 0. Rearranging gives ΔH_system - TΔS_system < 0, which is ΔrG < 0 for a spontaneous process.
Card 27
What does ΔG = 0 signify for a reaction?
ΔG = 0 signifies that the reaction is at equilibrium, meaning there is no net change in the concentrations of reactants and products over time, and no useful work can be extracted.
Card 28
Under what conditions (signs of ΔH and ΔS) is a reaction always spontaneous?
A reaction is always spontaneous (ΔG 0 (increase in entropy).
Card 29
Under what conditions (signs of ΔH and ΔS) is a reaction never spontaneous?
A reaction is never spontaneous (ΔG > 0) when ΔH > 0 (endothermic) and ΔS < 0 (decrease in entropy).
Card 30
Under what conditions (signs of ΔH and ΔS) is spontaneity dependent on temperature (specifically, spontaneous at high temperatures)?
Spontaneity is temperature-dependent (spontaneous at high T) when ΔH > 0 (endothermic) and ΔS > 0 (increase in entropy). The TΔS term becomes larger than ΔH at high T, making ΔG negative.
Card 31
Under what conditions (signs of ΔH and ΔS) is spontaneity dependent on temperature (specifically, spontaneous at low temperatures)?
Spontaneity is temperature-dependent (spontaneous at low T) when ΔH < 0 (exothermic) and ΔS < 0 (decrease in entropy). The TΔS term remains smaller than ΔH at low T, making ΔG negative.
Card 32
Define an endothermic process.
An endothermic process is a reaction or process which requires or absorbs energy (in the form of heat) from its surroundings (ΔH > 0).
Card 33
Define an exothermic process.
An exothermic process is a reaction or process which releases energy (in the form of heat) into its surroundings (ΔH < 0).
Card 34
Define an endergonic process.
An endergonic process is a reaction or process that absorbs energy in the form of work (or free energy) from the surroundings, meaning ΔG > 0 and it is non-spontaneous.
Card 35
Define an exergonic process.
An exergonic process is a reaction or process that releases energy in the form of work (or free energy) into the surroundings, meaning ΔG < 0 and it is spontaneous.
Card 36
State Lavoisier-Laplace's Law.
Lavoisier-Laplace's Law states that the enthalpy change for a reaction in one direction is equal in magnitude but opposite in sign to the enthalpy change for the reverse reaction.
Card 37
Provide an example illustrating Lavoisier-Laplace's Law.
If the formation of water from hydrogen and oxygen (2H₂(g) + O₂(g) → 2H₂O(l)) is exothermic (ΔH 0) with the same magnitude.
Card 38
State Hess's Law.
Hess's Law states that the total enthalpy change for a chemical reaction is independent of the pathway taken or the number of intermediate steps involved, as long as the initial and final conditions are the same.
Card 39
What is the key implication of Hess's Law regarding reaction enthalpy?
The change of reaction enthalpy (ΔrH) is not dependent on the reaction mechanism or the number of reaction steps. It's a state function.
Card 40
How is Hess's Law useful in calculating reaction enthalpies?
Hess's Law allows for the calculation of an overall reaction enthalpy by summing the enthalpy changes of individual steps, even if those steps cannot be measured directly.
Card 41
Can reaction enthalpy depend on the reaction mechanism according to Hess's Law?
No, Hess's Law explicitly states that reaction enthalpy is independent of the reaction mechanism or number of steps.
Card 42
What is the standard pressure defined for a reaction system?
The standard pressure is 1 atmosphere (atm) or 10⁵ Pascals (1 bar).
Card 43
What is the standard temperature for a reaction system in Celsius and Kelvin?
The standard temperature is typically 25°C (298 K), though 0°C (273 K) can be used depending on context.
Card 44
What is the standard concentration/activity for solutes in a reaction system?
The standard concentration for a solute is 1 Molar (1 M).
Card 45
What is the standard state for pure liquids and solids?
For pure liquids and solids, the standard state is the pure substance itself at 1 atm (or 1 bar) pressure and the specified temperature (usually 298 K).
Card 46
What is the value of the reaction quotient (Q) at the standard state?
At the standard state, the reaction quotient (Q) is defined as 1 (Q=1).
Card 47
Why is characterizing a 'standard state' important in thermochemistry?
The standard state provides a common reference point for comparing thermodynamic properties of different reactions, allowing for consistent calculations and predictions.
Card 48
Define standard molar enthalpy of formation (ΔfH).
The standard molar enthalpy of formation (ΔfH°) is the enthalpy change when one mole of a compound is formed from its constituent elements in their standard states under standard conditions (1 atm/bar, 298 K).
Card 49
What is the ΔfH° value for an element in its standard state?
The standard molar enthalpy of formation (ΔfH°) for an element in its most stable form under standard conditions is defined as zero (e.g., ΔfH°(O₂(g)) = 0, ΔfH°(C(graphite)) = 0).
Card 50
How are ΔfH° values typically obtained or used?
ΔfH° values are typically read from predetermined tables and are used to calculate the enthalpy change of any reaction.
Card 51
What is the formula for calculating the standard enthalpy of a reaction (ΔrH°) using standard molar enthalpies of formation?
ΔrH° = ΣnΔfH°(products) - ΣmΔfH°(reactants), where n and m are the stoichiometric coefficients.
Card 52
Concept/Formula: Explain how to use standard molar enthalpies of formation (ΔfH°) to calculate the standard enthalpy of a reaction (ΔrH°).
The standard enthalpy of reaction (ΔrH°) can be calculated by subtracting the sum of the standard molar enthalpies of formation of the reactants (each multiplied by its stoichiometric coefficient) from the sum of the standard molar enthalpies of formation of the products (each multiplied by its stoichiometric coefficient). Elements in their standard states have ΔfH° = 0.
Card 53
Problem: Calculate ΔrH° for the combustion of methane: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l). Given: ΔfH°(CH₄) = -74.8 kJ/mol, ΔfH°(CO₂) = -393.5 kJ/mol, ΔfH°(H₂O) = -285.8 kJ/mol, ΔfH°(O₂) = 0 kJ/mol.
Step-by-step solution: 1. Identify products and reactants: Products: CO₂(g), H₂O(l); Reactants: CH₄(g), O₂(g). 2. Apply the formula: ΔrH° = [1 * ΔfH°(CO₂) + 2 * ΔfH°(H₂O)] - [1 * ΔfH°(CH₄) + 2 * ΔfH°(O₂)]. 3. Substitute values: ΔrH° = [1 * (-393.5 kJ/mol) + 2 * (-285.8 kJ/mol)] - [1 * (-74.8 kJ/mol) + 2 * (0 kJ/mol)]. 4. Calculate: ΔrH° = [-393.5 - 571.6] - [-74.8 + 0] ΔrH° = -965.1 - (-74.8) ΔrH° = -965.1 + 74.8 ΔrH° = -890.3 kJ/mol.
Card 54
Why is knowing ΔfH° useful for calculating the enthalpy of *any* reaction?
Because ΔfH° values are tabulated for many compounds, they serve as building blocks to calculate the enthalpy change for virtually any reaction using Hess's Law, without needing to measure each reaction directly.
Card 55
List some common reactions/processes for which knowing the standard enthalpy change is particularly useful.
It is useful to know ΔH values for common reactions/processes such as combustion, dissolution, hydration, neutralization, and hydrolysis, as well as covalent bond breakage and formation.
Card 56
Why is it particularly useful to know ΔH values for compounds such as H₂O and CO₂?
H₂O and CO₂ are common products in many reactions (especially combustion and biological processes), so knowing their formation enthalpies helps calculate the energy released or absorbed in these widespread reactions.
Card 57
What does the enthalpy of combustion tell us?
The enthalpy of combustion tells us the amount of heat released when one mole of a substance undergoes complete combustion with oxygen under standard conditions, indicating its energy content.
Card 58
What does the enthalpy of dissolution tell us?
The enthalpy of dissolution tells us the heat change when one mole of a substance dissolves in a solvent, indicating whether the dissolution process is endothermic or exothermic.
Card 59
How does the comparison of the reaction quotient (Q) and the equilibrium constant (K) predict the spontaneity of a reaction?
If Q K, the reaction will proceed spontaneously in the reverse direction (ΔG > 0 for forward). If Q = K, the reaction is at equilibrium (ΔG = 0).
Card 60
What is the central mathematical expression that describes the dependence of ΔrG on the reaction quotient (Q) and the equilibrium constant (K)?
The central expression is ΔrG = ΔrG° + RTlnQ.
Card 61
What do the terms R and T represent in the equation ΔrG = ΔrG° + RTlnQ?
R is the ideal gas constant (8.314 J/mol·K), and T is the absolute temperature in Kelvin.
Card 62
How does the sign of ΔrG change when Q < K?
When Q < K, ln(Q/K) is negative, making ΔrG negative. The reaction is spontaneous in the forward direction.
Card 63
How does the sign of ΔrG change when Q > K?
When Q > K, ln(Q/K) is positive, making ΔrG positive. The reaction is non-spontaneous in the forward direction (spontaneous in reverse).
Card 64
What does it mean when Q = K for a reaction?
When Q = K, the reaction is at equilibrium, and ΔrG = 0. There is no net change in the system.
Card 65
Concept/Formula: Explain how to calculate ΔrG using the standard Gibbs free energy change (ΔrG°), the gas constant (R), temperature (T), and the reaction quotient (Q).
The actual Gibbs free energy change (ΔrG) under non-standard conditions can be calculated using the equation ΔrG = ΔrG° + RTlnQ. This allows for the prediction of spontaneity under specific reactant and product concentrations.
Card 66
Problem: Given ΔrG° = +10 kJ/mol for a reaction at 298 K, calculate ΔrG if the reaction quotient Q = 0.1. (R = 8.314 J/mol·K)
Step-by-step solution: 1. Convert ΔrG° to Joules: ΔrG° = +10,000 J/mol. 2. Apply the formula: ΔrG = ΔrG° + RTlnQ. 3. Substitute values: ΔrG = 10,000 J/mol + (8.314 J/mol·K)(298 K)ln(0.1). 4. Calculate ln(0.1) ≈ -2.3026. 5. Calculate RTlnQ: (8.314)(298)(-2.3026) ≈ -5708 J/mol. 6. Calculate ΔrG: ΔrG = 10,000 J/mol - 5708 J/mol = +4292 J/mol or +4.29 kJ/mol. (Since ΔrG is positive, the reaction is non-spontaneous under these conditions.)
Card 67
What is the mathematical equation relating the standard Gibbs free energy change (ΔrG°) to the equilibrium constant (K) of a reaction?
ΔrG° = -RTlnK.
Card 68
How is the equation ΔrG° = -RTlnK derived from the general ΔrG equation?
At equilibrium, ΔrG = 0 and Q = K. Substituting these into the general equation ΔrG = ΔrG° + RTlnQ gives 0 = ΔrG° + RTlnK, which rearranges to ΔrG° = -RTlnK.
Card 69
Why is the expression ΔrG° = -RTlnK valid for a reaction occurring at the standard state, even if it's an initial or non-equilibrium state?
The ΔrG° term itself *defines* the standard Gibbs free energy change, which is a fixed value for a given reaction at a specific temperature. It's the hypothetical change if reactants in their standard states converted to products in their standard states. The relationship to K connects this standard change to the equilibrium state, which is reached regardless of the initial state.
Card 70
Concept/Formula: Explain how to calculate the equilibrium constant (K) from the standard Gibbs free energy change (ΔrG°).
The equilibrium constant (K) can be calculated from the standard Gibbs free energy change (ΔrG°) using the equation K = e^(-ΔrG°/RT). A more negative ΔrG° indicates a larger K, favoring products at equilibrium.
Card 71
Problem: Calculate the equilibrium constant (K) for a reaction at 298 K if ΔrG° = -30 kJ/mol. (R = 8.314 J/mol·K)
Step-by-step solution: 1. Convert ΔrG° to Joules: ΔrG° = -30,000 J/mol. 2. Apply the formula: K = e^(-ΔrG°/RT). 3. Calculate the exponent: -ΔrG°/RT = -(-30,000 J/mol) / (8.314 J/mol·K * 298 K) = 30,000 / 2477.572 ≈ 12.108. 4. Calculate K: K = e^(12.108) ≈ 1.81 x 10⁵. (A large K value indicates that products are heavily favored at equilibrium.)
Card 72
Specify the standard state conditions for a biological system (biochemical reactions).
The biochemical standard state (ΔG'°) is typically defined as 1 bar (10⁵ Pa) pressure, 25°C (298 K), and pH 7.0 (where [H⁺] = 10⁻⁷ M).
Card 73
Define a 'high energy compound' in biochemistry.
A high energy compound is a biochemical molecule that possesses a high potential for group transfer (e.g., phosphoryl group) and releases a significant amount of free energy (more than approximately -25 kJ/mol) upon hydrolytic cleavage of a specific bond.
Card 74
What is the typical ΔrG threshold for hydrolytic cleavage to classify a compound as 'high energy'?
Compounds are typically classified as 'high energy' if they release more than ΔrG > (-)25 kJ/mol during hydrolytic cleavage.
Card 75
Provide an example of a high energy compound.
Adenosine triphosphate (ATP) is a prime example of a high energy compound.
Card 76
What characteristics contribute to a compound being 'high energy'?
High energy compounds often have bonds (e.g., phosphoanhydride bonds in ATP) that, upon hydrolysis, relieve electrostatic repulsion, stabilize products through resonance, and allow for greater solvation.
Card 77
Define coupled biochemical reactions.
Coupled biochemical reactions are sequential or consecutive reactions that share a common intermediate substance, allowing an energetically unfavorable reaction to be driven by an energetically favorable one.
Card 78
How do coupled reactions overcome thermodynamically unfavorable steps?
A thermodynamically unfavorable (endergonic, ΔG > 0) reaction can be driven by coupling it to a highly favorable (exergonic, ΔG < 0) reaction, such that the overall net ΔG for the coupled process is negative.
Card 79
What is the role of a common intermediate substance in coupled reactions?
The common intermediate links the two reactions, allowing the free energy released from the favorable reaction to be effectively transferred and utilized by the unfavorable reaction.
Card 80
Concept/Formula: Explain how to calculate the overall ΔrG for coupled reactions.
When two reactions are coupled, their Gibbs free energy changes are additive. The overall ΔrG for the coupled reaction is the sum of the ΔrG values of the individual reactions: ΔrG_overall = ΔrG₁ + ΔrG₂.
Card 81
Problem: Reaction 1: A → B, ΔrG₁ = +20 kJ/mol. Reaction 2: B → C, ΔrG₂ = -35 kJ/mol. Calculate the overall ΔrG for the coupled reaction A → C.
Step-by-step solution: 1. Identify the individual ΔrG values: ΔrG₁ = +20 kJ/mol, ΔrG₂ = -35 kJ/mol. 2. Apply the formula for coupled reactions: ΔrG_overall = ΔrG₁ + ΔrG₂. 3. Substitute values: ΔrG_overall = (+20 kJ/mol) + (-35 kJ/mol). 4. Calculate: ΔrG_overall = -15 kJ/mol. (The overall coupled reaction is spontaneous.)
Card 82
Explain the role of ATP in metabolic reactions.
ATP (adenosine triphosphate) serves as the primary energy currency of the cell, providing the energy required for various metabolic processes by undergoing hydrolysis to release free energy.
Card 83
How does ATP act as an energy currency in the cell?
ATP stores chemical energy in its phosphoanhydride bonds. When these bonds are broken (hydrolyzed), a significant amount of free energy is released, which can be used to power cellular work.
Card 84
What makes the phosphodiester bonds (specifically phosphoanhydride bonds) in ATP 'high energy'?
The phosphoanhydride bonds in ATP are 'high energy' due to associated electronegative charges between the phosphate groups that repel them, and the hydrolysis products are more stable (e.g., due to resonance stabilization and increased solvation).
Card 85
What process releases energy from ATP?
The cleavage of the phosphodiester (phosphoanhydride) bonds, typically through hydrolysis, releases energy from ATP.
Card 86
How does ATP facilitate thermodynamically unfavorable reactions?
ATP facilitates unfavorable reactions through 'ATP coupling,' where the energy released from ATP hydrolysis (a thermodynamically favorable reaction) is used to drive a thermodynamically unfavorable reaction.
Card 87
What is 'ATP coupling'?
ATP coupling is the process where the hydrolysis of ATP (a thermodynamically favorable, exergonic reaction with a negative ΔG) is linked to a thermodynamically unfavorable reaction (an endergonic reaction that cannot proceed without energy input), making the overall coupled reaction spontaneous.
Card 88
In ATP coupling, is the hydrolysis of ATP a thermodynamically favorable or unfavorable reaction?
The hydrolysis of ATP is a thermodynamically favorable (exergonic) reaction, meaning it has a negative ΔG and releases energy spontaneously.
Card 89
What is the typical range for ΔrG of ATP hydrolysis under physiological conditions?
Under physiological conditions, the ΔrG for ATP hydrolysis to ADP + Pi is typically around -30.5 kJ/mol, making it a highly exergonic reaction.
Card 90
Predict the spontaneity of a process if ΔH 0.
Always spontaneous (exergonic) at all temperatures, as ΔG = ΔH - TΔS will always be negative.
Card 91
Predict the spontaneity of a process if ΔH > 0 and ΔS < 0.
Never spontaneous (endergonic) at any temperature, as ΔG = ΔH - TΔS will always be positive.
Card 92
Predict the spontaneity of a process if ΔH < 0 and ΔS < 0.
Spontaneous at low temperatures, non-spontaneous at high temperatures. ΔG will be negative when the magnitude of ΔH is greater than TΔS.
Card 93
Predict the spontaneity of a process if ΔH > 0 and ΔS > 0.
Spontaneous at high temperatures, non-spontaneous at low temperatures. ΔG will be negative when the magnitude of TΔS is greater than ΔH.
Card 94
Why is the equilibrium constant (K) a useful value to know for a reaction?
K provides information about the relative amounts of reactants and products at equilibrium, indicating the extent to which a reaction proceeds in the forward or reverse direction.
Card 95
What is the significance of a very large equilibrium constant (K >> 1)?
A very large K indicates that at equilibrium, the reaction strongly favors the formation of products, meaning the reaction proceeds almost to completion.
Card 96
What is the significance of a very small equilibrium constant (K << 1)?
A very small K indicates that at equilibrium, the reaction strongly favors the reactants, meaning very little product is formed.
Card 97
What is the relationship between the standard Gibbs free energy change (ΔrG°) and the extent of a reaction?
A more negative ΔrG° corresponds to a larger equilibrium constant (K) and a greater extent of reaction towards products at equilibrium.
Card 98
How does temperature affect the spontaneity of a reaction if both ΔH and ΔS have the same sign?
If ΔH and ΔS have the same sign, temperature will determine spontaneity. For example, if both are positive, higher temperatures make it spontaneous, while if both are negative, lower temperatures make it spontaneous.
Card 99
What is the difference between ΔG and ΔG°?
ΔG is the Gibbs free energy change under any given conditions, while ΔG° is the standard Gibbs free energy change under standard state conditions (1 atm/bar, 298 K, 1 M concentration).
Card 100
In the context of the Second Law of Thermodynamics, what does 'increase in entropy in the universe' specifically refer to?
It refers to the sum of the entropy change of the system (ΔS_system) and the entropy change of the surroundings (ΔS_surroundings). For a spontaneous process, ΔS_universe = ΔS_system + ΔS_surroundings > 0.
Card 101
How can an endothermic reaction (ΔH > 0) still be spontaneous?
An endothermic reaction can be spontaneous if the increase in entropy (ΔS > 0) is sufficiently large to overcome the unfavorable enthalpy change, especially at high temperatures (where -TΔS term becomes more negative than ΔH is positive).
Card 102
How can a reaction that decreases entropy (ΔS < 0) still be spontaneous?
A reaction that decreases entropy can still be spontaneous if it is highly exothermic (ΔH < 0), meaning the release of heat to the surroundings significantly increases the entropy of the surroundings, making ΔS_universe positive, especially at low temperatures.
Card 103
What is the relationship between the magnitude of ΔH and ΔS and the 'difficulty' of a reaction?
Large positive ΔH (highly endothermic) or large negative ΔS (large decrease in disorder) makes a reaction less favorable (more 'difficult' to be spontaneous), while large negative ΔH (highly exothermic) or large positive ΔS (large increase in disorder) makes it more favorable.
Card 104
What is a 'phosphoryl group transfer' and why is it important in metabolism?
A phosphoryl group transfer is the transfer of a phosphate group from one molecule (like ATP) to another. It's crucial in metabolism for activating molecules, driving unfavorable reactions, and regulating enzyme activity by changing the recipient molecule's energy state or conformation.
Card 105
Give an example of how ATP's phosphoryl group transfer activates a molecule for a subsequent reaction.
In glycolysis, ATP transfers a phosphate group to glucose, forming glucose-6-phosphate. This phosphorylation makes glucose more reactive and traps it within the cell, committing it to further metabolic steps.
Card 106
How does the hydrolysis of ATP to ADP + Pi compare energetically to the hydrolysis of ATP to AMP + PPi?
Hydrolysis of ATP to AMP + PPi (pyrophosphate) releases even more free energy than hydrolysis to ADP + Pi. The subsequent hydrolysis of PPi to 2Pi further drives the overall reaction, effectively releasing ~61 kJ/mol (compared to ~30.5 kJ/mol for ATP to ADP).
Card 107
Why is the exact ΔG of ATP hydrolysis variable in a living cell?
The exact ΔG (not ΔG°) of ATP hydrolysis in a living cell is variable because it depends on the actual cellular concentrations of ATP, ADP, and Pi, as well as pH and Mg²⁺ concentration, which are rarely at standard state values.
Card 108
What is the typical physiological concentration of ATP relative to ADP in a healthy cell, and what is its significance?
A healthy cell typically maintains a high ATP:ADP ratio. This high ratio keeps the actual cellular ΔG of ATP hydrolysis more negative than ΔG° (-30.5 kJ/mol), providing a strong driving force for coupled reactions.
Card 109
How does the principle of coupled reactions apply to the synthesis of macromolecules like proteins or nucleic acids?
The synthesis of macromolecules is generally endergonic. These processes are driven by coupling them with the hydrolysis of high-energy compounds, primarily ATP (or GTP), which provides the necessary free energy.
Card 110
What is the role of an enzyme in a coupled reaction?
Enzymes catalyze both the exergonic and endergonic steps of coupled reactions. They ensure that the energy released from the favorable reaction is efficiently funneled to the unfavorable reaction, often by forming a temporary enzyme-substrate-ATP complex.
Card 111
Can a non-spontaneous reaction occur without coupling?
No, a non-spontaneous reaction (ΔG > 0) cannot occur on its own without an input of energy. It requires coupling to an exergonic process or an external energy source.
Card 112
What is meant by a 'reaction quotient' (Q)?
The reaction quotient (Q) is a measure of the relative amounts of products and reactants present in a reaction at any given time. It is calculated in the same way as the equilibrium constant (K) but for non-equilibrium conditions.
Card 113
How does the value of Q relate to the current state of a reaction relative to equilibrium?
If Q K, the reaction proceeds in reverse to reach equilibrium. If Q = K, the reaction is at equilibrium.
Card 114
When calculating ΔrG, why is it important to use the absolute temperature (Kelvin) instead of Celsius?
The thermodynamic equations involving temperature (e.g., ΔG = ΔH - TΔS, ΔG° = -RTlnK) are derived using fundamental thermodynamic principles that require temperature to be expressed on an absolute scale (Kelvin) to avoid mathematical issues (e.g., division by zero or negative temperatures).
Card 115
What is the significance of the gas constant (R) in thermodynamic equations?
The gas constant (R) acts as a proportionality constant, converting energy units (Joules) to temperature and amount of substance units (Kelvin and moles) in equations relating free energy, temperature, and concentration/pressure terms.
Card 116
How does the Second Law of Thermodynamics relate to the concept of 'irreversibility' in natural processes?
The Second Law implies that all spontaneous processes are irreversible. While energy is conserved (First Law), entropy continuously increases in the universe for spontaneous changes, meaning the system cannot return to its original state without external work and a further increase in the entropy of the universe.
Card 117
What is the difference between 'work' (W) and 'heat' (Q) in the context of the First Law of Thermodynamics?
Work (W) is energy transferred due to a force acting over a distance (e.g., expansion work). Heat (Q) is energy transferred due to a temperature difference. Both are forms of energy transfer, but they represent different mechanisms.
Card 118
Why is enthalpy often more useful than internal energy for describing chemical reactions?
Most chemical reactions occur at constant atmospheric pressure, not constant volume. Under constant pressure, the heat exchanged (q_p) is equal to ΔH, making enthalpy a direct and convenient measure of energy changes in these common conditions.
Card 119
What is the practical implication of a large negative ΔrG for a reaction in an industrial setting?
A large negative ΔrG indicates a highly spontaneous and product-favored reaction. This is desirable industrially as it means high yields can be achieved without continuous energy input, though reaction rate also needs to be considered.
Card 120
What is the practical implication of a large positive ΔrG for a reaction in an industrial setting?
A large positive ΔrG indicates a non-spontaneous, reactant-favored reaction. This means significant energy input or extreme conditions (e.g., high temperature, pressure) would be required to drive the reaction towards products, making it potentially uneconomical.
Card 121
Why is understanding thermochemistry important for drug development?
Thermochemistry helps in understanding drug-receptor binding affinity (ΔG of binding), predicting drug stability (ΔH, ΔS of decomposition), and optimizing drug formulation by understanding solubility and dissolution enthalpies.
Card 122
How does thermochemistry help in understanding climate change?
Thermochemistry is crucial for understanding the energy balance of the Earth, the heat absorption/emission by greenhouse gases, and the energetics of combustion processes that release CO₂, informing climate models and mitigation strategies.
Card 123
What is the connection between bond energy and enthalpy?
Bond energy is the energy required to break a specific bond. In a reaction, the overall enthalpy change (ΔH) is approximately the sum of the energies required to break bonds in reactants minus the sum of the energies released when new bonds form in products.
Card 124
Why is the value of ΔfH° for elements in their standard state defined as zero?
By definition, the formation of an element from itself in its most stable form requires no energy change, providing a consistent reference point for all other enthalpy of formation calculations.
Card 125
How does the concept of 'activity' differ from 'concentration' in thermodynamics, especially for non-ideal solutions?
Concentration is a measure of the amount of substance per unit volume. Activity is a thermodynamic measure of the 'effective concentration' of a species, taking into account non-ideal behavior (intermolecular interactions). For ideal solutions, activity approximates concentration.
Card 126
In the equation ΔrG = ΔH - TΔrS, what does the TΔrS term represent?
The TΔrS term represents the amount of energy that is 'unavailable' to do useful work because it is dispersed as increased disorder (entropy) within the system and surroundings at a given temperature.
Card 127
What is the significance of the 'entropy of the universe' in the Second Law of Thermodynamics?
The 'entropy of the universe' (ΔS_universe) is the ultimate criterion for spontaneity. It states that for any spontaneous process, the total entropy of the system and its surroundings must increase, meaning the universe always tends towards greater disorder.
Card 128
How does the concept of 'free energy' relate to the capacity to do work?
Free energy (Gibbs free energy, ΔG) represents the maximum amount of non-expansion work that can be extracted from a thermodynamically closed system at constant temperature and pressure. A negative ΔG means the system can do work.
Card 129
What happens to the ΔG of a reaction as it approaches equilibrium?
As a reaction approaches equilibrium, the magnitude of ΔG decreases, eventually reaching zero at equilibrium, where no more net work can be done by the system.
Card 130
Why is it important to distinguish between ΔrH and ΔrG when discussing spontaneity?
ΔrH only indicates whether a reaction releases or absorbs heat. ΔrG, however, considers both enthalpy and entropy changes, providing the true measure of spontaneity and the capacity for useful work. An exothermic reaction (negative ΔrH) isn't necessarily spontaneous if the entropy change is unfavorable.
Card 131
What does a negative ΔrG'° value for a biochemical reaction indicate?
A negative ΔrG'° indicates that the biochemical reaction is spontaneous under standard biochemical conditions (pH 7, 25°C, 1 bar, 1 M initial concentrations, except H+ and H2O).
Card 132
What does a positive ΔrG'° value for a biochemical reaction indicate?
A positive ΔrG'° indicates that the biochemical reaction is non-spontaneous under standard biochemical conditions and would require energy input (e.g., through coupling with ATP hydrolysis) to proceed.
Card 133
Why is the pH of 7 specified for the biochemical standard state (ΔG'°)?
Biological systems operate at or near neutral pH. Specifying pH 7 as standard ensures that calculations of free energy changes are relevant to physiological conditions, as [H⁺] concentration significantly affects many biochemical reactions.
Card 134
What is the role of Mg²⁺ ions in ATP hydrolysis in biological systems?
Mg²⁺ ions often bind to ATP, stabilizing its negative charges and influencing its conformation. This binding can affect the actual ΔG of ATP hydrolysis and is often required for enzyme activity in ATP-dependent reactions.
Card 135
How do living organisms maintain a low entropy state locally, despite the Second Law of Thermodynamics?
Living organisms maintain a low entropy state locally (i.e., highly organized structures) by continuously increasing the entropy of their surroundings. They consume ordered nutrients and release heat and disordered waste products, thereby increasing the total entropy of the universe.
Card 136
What is an 'intermediate substance' in the context of coupled reactions?
An intermediate substance is a compound formed in an earlier step of a reaction sequence that is then consumed in a later step. In coupled reactions, the intermediate links the exergonic and endergonic processes.
Card 137
Why is it important to consider the 'actual' ΔG, rather than just ΔG°, in biological systems?
Biological systems rarely operate under standard (1 M) concentrations. The actual cellular concentrations of reactants and products, as well as pH and temperature, can significantly alter the spontaneity and direction of a reaction, which is reflected by ΔG, not just ΔG°.
Card 138
Explain the concept of 'energy landscape' in relation to ΔG.
The energy landscape (or reaction coordinate diagram) plots the free energy of a system as a reaction proceeds. ΔG represents the difference in free energy between reactants and products, indicating the overall favorability of the process.
Card 139
What is the difference between 'spontaneous' in thermodynamics and 'fast' in kinetics?
Thermodynamic spontaneity (ΔG < 0) indicates whether a reaction *can* occur without external energy input. Kinetic speed indicates *how fast* a reaction occurs. A spontaneous reaction can be very slow if it has a high activation energy.
Card 140
How does the concept of enthalpy apply to phase changes (e.g., melting or boiling)?
Phase changes involve specific enthalpy changes (e.g., enthalpy of fusion for melting, enthalpy of vaporization for boiling). These values represent the heat absorbed or released during the phase transition at constant temperature and pressure.
Card 141
What is the significance of the term 'q_rev' in the definition of entropy change (ΔS = q_rev / T)?
q_rev refers to heat transferred in a *reversible* process. This is a theoretical construct for calculating entropy because entropy is a state function. For irreversible processes (all real spontaneous processes), ΔS is still calculated using this definition, even though q for an irreversible process would be different.
Card 142
Why can't work (W) and heat (Q) be considered state functions?
Work and heat are not state functions because their values depend on the specific path taken during a process, not just on the initial and final states of the system.
Card 143
How can the Gibbs-Helmholtz equation (derived from ΔG = ΔH - TΔS) be used to predict the temperature at which a reaction becomes spontaneous?
By setting ΔG = 0, we can find the equilibrium temperature (T_eq = ΔH / ΔS) where the reaction switches from non-spontaneous to spontaneous (or vice-versa). For T > T_eq or T < T_eq, the spontaneity can be predicted based on the signs of ΔH and ΔS.
Card 144
What happens to the internal energy of an ideal gas when it undergoes isothermal expansion (constant temperature)?
For an ideal gas, internal energy (ΔU) depends only on temperature. Therefore, during isothermal expansion, ΔU = 0, meaning any heat absorbed is entirely converted into work done by the gas (Q = W).
Card 145
What is an 'isolated system' in thermodynamics, and how does it relate to the First Law?
An isolated system is one that cannot exchange either energy (heat or work) or matter with its surroundings. For an isolated system, the First Law states that its total internal energy remains constant (ΔU = 0).
Card 146
What is the key difference between an 'open system' and a 'closed system'?
An open system can exchange both matter and energy with its surroundings. A closed system can exchange energy (heat or work) but not matter with its surroundings.
Card 147
Why are standard conditions (e.g., 298 K) chosen for thermodynamic tables, even if reactions often occur at different temperatures?
Standard conditions provide a universal reference point for comparing thermodynamic data. While actual reactions occur at various temperatures, these standard values can be used with equations (like ΔG = ΔH - TΔS) to calculate properties at other temperatures.
Card 148
What is meant by the 'activity' of a pure liquid or solid in thermodynamic calculations?
The activity of a pure liquid or solid in its standard state is defined as 1. This means their concentrations do not change significantly during a reaction, and they do not affect the reaction quotient (Q) or equilibrium constant (K) expressions.
Card 149
How does the concept of entropy explain the mixing of two ideal gases?
When two ideal gases mix, their entropy increases spontaneously (ΔS > 0). This is because the mixed state has more possible microstates (greater disorder) than the separated state, making the mixing process energetically favorable in terms of entropy.
Study Thermochemistry and Thermodynamics the way that sticks
Reading cards is not the same as remembering them. In FORKSAI you answer each card before you see the back, and the ones you miss come back first.
Study these cards freeMore flashcard decks
- Género Bacillus y Clostridium flashcards (100 cards)
- Ψηφιακά Παιχνίδια και Μάθηση flashcards (100 cards)
- Philosophy AS flashcards (100 cards)
- Nervous system flashcards (100 cards)
- ECG, IGEL and Iv flashcards (100 cards)
- Biology - Year 11 flashcards (100 cards)