Questions and Answers
If ΔE is positive, energy can be written as a product in the equation for the reaction.
True
Which response is false regarding the reaction CH4(g) + 2O2(g) → CO2(g) + 2H2O?
What is ΔE for the combustion of a 1.00-g sample of hexane, C6H14?
−5.91 × 10^3 kJ/mol
What is ΔE for the combustion of a 0.900-g sample of toluene, C7H8?
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What is the amount of work done when a system is compressed from 50.0 L to 5.0 L at a constant pressure of 10.0 atm?
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What work is done in the conversion of 1.00 mole of Ni to Ni(CO)4 at 75°C?
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What is ΔE for the conversion of 2.00 moles of NO2 to N2O4 at 125°C?
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Which statement is false regarding thermodynamic processes?
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Which one of the following statements is false?
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Which of the following statements regarding spontaneous changes is false?
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What is the change in entropy of the reaction N2(g) + 3H2(g) → 2NH3(g)?
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Which of the following statements about free energy is false?
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Which statement is incorrect?
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Which of the following statements about the first law of thermodynamics and energy is false?
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Which of the following statements is a correct interpretation of the First Law of Thermodynamics?
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Which one of the following thermodynamic quantities is not a state function?
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Which term is not correctly matched?
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The enthalpy change, ΔH, of a process is defined as:
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Which statement regarding enthalpy change is incorrect?
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How much heat was absorbed by the calorimeter in an exothermic reaction that liberates 7.58 kJ of heat in a coffee cup calorimeter containing 157 grams of solution?
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Calculate the heat capacity of the calorimeter (in J/°C) if 4.168 kJ of heat is added to a calorimeter containing 75.40 g of water, which increases in temperature from 24.58°C to 35.82°C.
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Calculate the molar heat of neutralization for the reaction between 50.0 mL solution of 1.2 M HCl and 50.0 mL of 1.3 M NaOH resulting in a final temperature of 29.8°C.
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How much heat is absorbed in the complete reaction of 3.00 grams of SiO2 with excess carbon?
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How much heat is released when 75 g of octane is burned completely?
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What is the molar heat of neutralization, ΔH, for the reaction of 5.5 grams of HCl with excess Ba(OH)2 that releases 8300 J of heat?
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How much heat is released when the burning of 80.3 g of SiH4 at constant pressure gives off heat?
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What is the ΔH of the reaction for the roasting of 48.7 g of ZnS at constant pressure that gives off heat?
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Which of the following statements is incorrect regarding the thermochemical standard state?
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Which of the following substances is not in its standard state?
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For which of the following substances does ΔHf^0 = 0?
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Which of the following substances is not correctly matched with its molar heat of formation?
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Calculate the amount of heat released in the complete combustion of 8.17 grams of Al to form Al2O3(s).
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How much heat energy is liberated when 11.0 grams of manganese is converted to Mn2O3 at standard state conditions?
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When 32.1 g of H2 reacts with excess silicon to form SiH4(g) at standard conditions, how much heat is absorbed?
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Which of the following is not a formation reaction?
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From the data provided, calculate ΔH0 at 25°C for this reaction: 4HCl(g) + O2(g) → 2Cl2(g) + 2H2O(g).
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Given the following at 25°C: 1/2N2(g) + O2(g) → NO2(g) ΔH0 = 33.2 kJ. Calculate the ΔH0 for 2NO2(g) → N2O4(g).
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Given the enthalpy changes for the following reactions, calculate for CO(g): C (graphite) + O2(g) → CO2(g) = −393.5 kJ, CO(g) + 1/2O2 → CO2(g) ΔH0 = −283.0 kJ.
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Given the standard heats of formation for the following compounds, calculate for the reaction: CH4(g) + H2O(g) → CH3OH + H2(g).
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Calculate ΔH0 at 25°C for the reaction 2ZnS(s) + 3O2(g) → 2ZnO(s) + 2SO2(g).
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Calculate ΔH0 for the following reaction at 25.0°C: Fe3O4(s) + CO(g) → 3FeO(s) + CO2(g).
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Calculate the standard enthalpy change for the reaction C(graphite) + 4HNO3 → CO2(g) + 4NO2(g) + 2H2O.
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Calculate the standard enthalpy change for the reaction 12NH3(g) + 21O2(g) → 8HNO3 + 4NO(g) + 14H2O(g).
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Evaluate ΔH0 for the reaction SiO2(s) + 4HF(aq) → SiF4(g) + 2H2O.
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Use the provided data to calculate for benzene, C6H6, at 25°C and 1 atm.
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Given the following data, calculate for HCN(g) at 25°C: 2NH3(g) + 3O2(g) + 2CH4(g) → 2HCN(g) + 6H2O.
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Given that ΔH0 for the oxidation of sucrose, C12H22O11(s), is −5648 kJ per mole at 25°C, evaluate for sucrose.
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At 25°C for CO(g), given that ΔH0 for the reaction 2CH4(g) + O2(g) + 4Cl2(g) → 8HCl(g) + 2CO(g) is −809.9 kJ, calculate for CO(g).
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How much heat is evolved in the formation of 35.0 grams of Fe2O3(s) at 25°C and 1.00 atm pressure?
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How much heat is released when 6.38 grams of Ag(s) reacts according to the equation at standard state conditions?
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How much heat is released or absorbed during the reaction of 10.0 grams of SiO2 with excess hydrofluoric acid?
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How much heat would be released or absorbed if 575 g of H2 is produced from the reaction?
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How much heat would be released if 12.0 g of methane, CH4, was completely burned in oxygen to form carbon dioxide and water at standard state conditions?
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Calculate the standard heat of vaporization for tin(IV) chloride, SnCl4, in kJ per mole.
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Which of the following techniques cannot be used to calculate ΔHrxn?
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The heat of reaction of one of the following reactions is the average bond energy for the N-H bond in NH3. Which one?
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Evaluate ΔH0 for the following reaction from the given bond energies: 2HBr(g) → H2(g) + Br2(g).
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Estimate the enthalpy change for the reaction given the average bond energies below: CH4(g) + 2Cl2(g) → CH2Cl2(g) + 2HCl(g).
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Estimate the heat of reaction at 298 K for the reaction given: Br2(g) + 3F2(g) → 2BrF3(g).
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Calculate the average N−H bond energy in NH3(g).
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Calculate the average S−F bond energy in SF6.
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Given: H−H bond energy = 435 kJ, Cl−Cl bond energy = 243 kJ, and the standard heat of formation of HCl(g) is −92 kJ/mol, calculate the H−Cl bond energy.
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Calculate the average bond energy in kJ per mol of bonds for the C−H bond.
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The heat of gaseous acetylene, H−C≡C−H, is 227 kJ/mol. What is the C≡C bond energy?
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Which of the following statements about internal energy, E, is false?
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Consider the following reaction at constant pressure. Which response is true? N2(g) + O2(g) → 2NO(g)
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Which statement concerning sign conventions for ΔE = q + w is false?
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Study Notes
Thermodynamics and Energy
- Energy is defined as the capacity to perform work or transfer heat.
- Kinetic energy refers to the energy associated with motion, while potential energy is linked to an object’s position or composition.
- Processes absorbing energy from surroundings are termed endothermic, while those releasing energy are exothermic.
- The First Law of Thermodynamics states that energy cannot be created or destroyed, only transformed.
Enthalpy (ΔH)
- Enthalpy change (ΔH) quantifies heat transferred during a process at constant pressure.
- Enthalpy is a state function, meaning its value depends only on the current state, not how it arrived there.
- The specific heat capacity of a solution can be used to calculate heat absorbed or released based on temperature changes.
Calorimetry
- When substances in a calorimeter react, the heat absorbed or released can be measured using changes in temperature and the specific heat formula.
- The heat capacity of a calorimeter helps in determining the total heat exchange during a reaction.
Heat of Neutralization
- The molar heat of neutralization refers to the heat released when an acid reacts with a base.
- Typical values are calculated using the amount of heat exchanged and the stoichiometry of the reaction.
Standard State and Formation Reactions
- Standard state refers to a substance's most stable form at one atmosphere and specified temperature (usually 298 K).
- The heats of formation (ΔHf) relate to the energy changes when one mole of a compound forms from its elements in their standard states.
Bond Energies
- Bond energy is a measure of the strength of the bonds in a compound; breaking bonds requires energy, while forming bonds releases energy.
- Overall reaction enthalpies can be calculated using average bond energies.
Specific Calculations
- Calculating ΔH for reactions involves using enthalpies of formation or measuring heat changes through calorimetry.
- Average bond energy calculations can estimate the strength of various bonds within a molecule.
Work in Thermodynamics
- Work can be done on or by a system, generally related to volume changes under pressure.
- The work done during thermal processes can be quantified through formulas linking pressure, volume, and changes therein.
Application of Thermodynamic Principles
- Careful application of thermodynamic laws enables the prediction of energy changes and the feasibility of chemical reactions, often vital in industrial and laboratory contexts.
Remember
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Each reaction can liberate or absorb significant amounts of energy, influencing reaction drive and practicality.
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Understanding these principles lays the groundwork for further studies in chemistry and related fields.### Thermodynamic Changes and Reactions
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Formation of Nickel Carbonyl:
- Reaction: Ni(s) + 4CO(g) → Ni(CO)4(g)
- ΔH value: +8.68 × 10^3 J
-
NO2 to N2O4 Reaction:
- Reaction: 2NO2(g) → N2O4(g)
- Work done calculated: +3,300 J at 125°C
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False Statements about Energy:
- Bomb calorimeters do not measure ΔH directly.
- In a typical chemical reaction, work done (w) is not usually larger than heat (q).
- For constant pressure processes, work is zero if gaseous moles do not change.
Changes in Gibbs Free Energy (ΔG)
-
Standard Energy Changes:
- For H2(g) + O2(g) → H2O(g): ΔE per mole = -240 kJ
- For SiO2(s) + 4HF(aq) → SiF4(g) + 2H2O: ΔE0 = -5.02 kJ/mol
- For reaction of NH3 and O2: ΔE0 = -3,503 kJ/mol
Spontaneity and Entropy
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Spontaneous Reactions:
- Spontaneity is favored by heat release and increased dispersal of matter.
- All exothermic reactions are not necessarily spontaneous.
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Entropy Changes:
- Positive ΔS indicates increased dispersal of matter.
- Process involving changes like rain forming leads to negative ΔS.
Key Concepts in Thermodynamics
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Third Law of Thermodynamics:
- Absolute entropy (S) is zero only at 0 K; it is positive at T > 0 K.
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Heat and Work:
- ΔH = 0 and ΔE = 0 at the melting point in phase changes.
- Free energy change (ΔG) can be calculated using heat and entropy changes.
Chemical Disequilibrium
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Endothermic and Exothermic Processes:
- Reaction entropy can be negative even with exothermic processes.
- Gibbs Free Energy changes are essential in determining thermodynamic favorability.
Reactions and Calculations
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Gibbs Free Energy Calculation:
- ΔG0 for reactions like P4O10 + 6H2O → 4H3PO4 = -363.7 kJ.
- ΔG can be influenced by both enthalpy (ΔH) and entropy (ΔS).
Entropy of Solutions
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Absolute Entropy:
- 1 M NaCl at 50°C possesses the highest absolute entropy among given states.
Conclusion
- Energy changes, entropy, and thermodynamic laws interact closely throughout chemical processes and reactions. The ability to calculate Gibbs Free Energy changes and recognize spontaneous reactions is pivotal in chemistry.
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Test your knowledge with these flashcards focused on concepts from Chapter 7 of Physics. Explore energy types, laws of thermodynamics, and other related definitions. Determine which statements about energy and processes are correct or incorrect.