AP Unit 6 Problems

A supplied dissolution experiment in an insulated cup lowers the solution temperature. Identify energy-flow direction for the dissolution and for the solution’s thermal change. Explain the net interaction balance and why a separate melting experiment can absorb heat without a temperature rise.

A molar physical transformation has initial enthalpy 8.0 and final enthalpy 30.0 kJ mol⁻¹ on a common scale. Construct an energy diagram: label both axes, draw and label the initial and final horizontal levels, and add a vertical arrow showing the forward enthalpy change. Give ΔH for the forward and reverse processes. Explain the effect of adding 100 kJ mol⁻¹ to both reference levels.

Two isolated bodies with constant heat capacities C_hot =30.0 J °C⁻¹ at 70.0 °C and C_cool =90.0 J °C⁻¹ at 10.0 °C reach thermal equilibrium without phase/reaction change. Explain net particle-level energy transfer and find T_f.

Read the initial and corrected-final temperatures from these two thermometer scales. Record each value with its unit and a sensible estimated digit. The smallest marked division is1 °C; estimate to0.1 °C, and treat each recorded reading as having a maximum uncertainty±0.1 °C.

Supplied dissolution data:80.0 g final solution, c=4.00 J g⁻¹ °C⁻¹, C_cup=15.0 J °C⁻¹, and0.0300 mol dissolved solute. Assume no remaining external heat loss. Use your recorded readings to calculate q_dissolution and molar ΔH. Explain the bias from neglecting the cup, and give the largest ΔT uncertainty from the two reading uncertainties. These are supplied instruments/data; no experiment is required.

An ideal supplied thermal experiment gives 250 J to each of two 10.0 g nonreacting samples. Sample A has c =0.500 J g⁻¹ °C⁻¹; sample B has c =2.00 J g⁻¹ °C⁻¹. Find their temperature rises. If A’s molar mass is 40.0 g mol⁻¹, find its molar heat capacity and show that the molar equation gives the same rise.

A supplied exothermic calorimetry trial loses heat to room air during the temperature reading. A student uses the solution’s uncorrected rise alone to infer ΔH. Identify the missing heat-transfer term, predict the bias, and describe a method using an existing temperature-time record that can improve the estimate. Explain why q =mcΔT would not by itself account for any simultaneous melting.

Consider this ideal heating sequence: begin with 36.0 g ice at −10.0 °C; melt it at 0 °C, then warm the liquid to 20.0 °C. Use c_ice =2.10, c_water =4.18 J g⁻¹ °C⁻¹, molar mass 18.0 g mol⁻¹, and ΔH_fus =6.01 kJ mol⁻¹. Calculate total heat and explain the plateau at the particle level.

At the specified boiling temperature/pressure, a substance has ΔH_vap =30.0 kJ mol⁻¹. A 0.250 mol vapor sample condenses completely without further cooling. Find q for the sample, compare it with vaporizing 0.500 mol, and state whether temperature changes during the ideal transition.

For 2CO(g) + O₂(g) → 2CO₂(g), ΔH = −566.0 kJ. A supplied sample contains 0.300 mol CO and 0.100 mol O₂.

1. Find the heat as the reaction proceeds to the limiting-reagent maximum. Explain warming if the released energy is transferred to a surrounding liquid.
2. In a separate supplied post-reaction observation, reaction is complete: the products are at 40.0 °C and a surrounding liquid is at 20.0 °C. They can exchange heat through a conductive wall. The combined assembly is insulated from the room; no further chemical reaction, phase change or work occurs, and both parts have positive heat capacities. Sketch or describe how both temperatures change over time, the direction of heat transfer, and the condition reached at the end. Is the final common temperature necessarily 20.0 °C?
3. Would reaching that final temperature establish chemical equilibrium between CO, O₂ and CO₂? Explain.

A balanced reaction is 2A → B with ΔH =+80.0 kJ for the equation as written. If 0.250 mol B is formed, find q and state ΔH for the reversed equation B → 2A.

Estimate C₂H₄(g) + H₂(g) → C₂H₆(g). Ethene has one C=C and four C–H bonds; ethane has one C–C and six C–H bonds. Average D values: C=C 614, H–H 436, C–C 347, C–H 413 kJ mol⁻¹. Show cancellation of unchanged bonds and discuss the estimate’s limitation.

For N₂(g) + 3H₂(g) → 2NH₃(g), use average D(N≡N) =945, D(H–H) =436, D(N–H) =391 kJ mol⁻¹. Estimate ΔH and separate the breaking input from formation release.

For CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l), supplied 298 K ΔH_f° values in kJ mol⁻¹ are CH₄ −74.8, CO₂ −393.5, H₂O(l) −285.8, O₂ 0. Find ΔH°. Recalculate if H₂O(g), −241.8 kJ mol⁻¹, is the product and explain the difference.

Write the standard formation reaction for exactly one mole of C₂H₅OH(l) at 298 K using carbon graphite, H₂(g), and O₂(g). Explain why formation data for O(g) cannot be replaced by zero when computing another reaction.

Find ΔH for 2C(graphite,s) +2H₂(g) → C₂H₄(g) using: C +O₂ → CO₂, −393.5 kJ; H₂ +½O₂ → H₂O(l), −285.8 kJ; C₂H₄ +3O₂ → 2CO₂ +2H₂O(l), −1411.0 kJ. Show operations and cancellation, then justify addition of enthalpies.

An accounting path for A(g) → B(l) has A(g) → B(g), ΔH =+25.0 kJ mol⁻¹, and B(g) → B(l), ΔH =−40.0 kJ mol⁻¹, at compatible stated conditions. Find ΔH for 0.300 mol A converted to B(l). Explain why stopping at B(g) would answer a different question and whether these two steps prove the actual mechanism.