General Chemistry I Exam 3 Due Date April 4, 2020 At 5:00 PM

General Chemistry Iexam 3due Date April 4 2020 At 500 Pmname1 Co

Analyze and solve a series of chemistry problems involving heat transfer, thermodynamics, reaction enthalpies, gas laws, and molecular formulas as presented in the provided assignment. The tasks include calculating heat quantities, temperature changes, calorimeter measurements, enthalpy of reactions, and properties of gases under different conditions.

Paper For Above instruction

Introduction

The study of thermodynamics and gas laws in chemistry provides critical insights into the energetic and behavioral characteristics of substances under various conditions. This paper addresses a set of complex chemistry problems that require applying principles such as specific heat, calorimetry, thermodynamic calculations, and gas laws. Through detailed calculations and explanations, the paper aims to demonstrate mastery of these fundamental concepts, integrating theoretical knowledge with practical problem-solving skills.

Calculating Heat Transfer and Temperature Changes

Problem 1 involves calculating the heat required to raise the temperature of copper metal. Given the specific heat capacity of copper (0.385 J/g·°C), mass (22.8 g), and temperature change (from 20.0°C to 875°C), the formula used is:

Q = mcΔT

Where Q is the heat absorbed, m is the mass, c is the specific heat, and ΔT is the temperature difference. Substituting the values:

Q = 22.8 g × 0.385 J/g·°C × (875°C - 20°C) = 22.8 g × 0.385 J/g·°C × 855°C ≈ 7,519.83 J

This calculation indicates that approximately 7.52 kJ of heat are required to heat the copper from 20°C to 875°C.

Problem 2 explores the temperature rise when aluminum absorbs a specific amount of heat (10.0 kJ). The specific heat of aluminum is 0.900 J/g·°C, and the mass is 25.0 g. Using the same formula:

ΔT = Q / (mc)

Converting 10.0 kJ to Joules (10,000 J):

ΔT = 10,000 J / (25.0 g × 0.900 J/g·°C) = 10,000 J / 22.5 J/°C ≈ 444.44°C

Estimating a significant temperature increase, reflecting how heat capacity influences temperature changes in materials.

Cooling of Water and Calorimetry

Problem 3 examines the cooling of water in a refrigerator. Starting with 200 g of water at 20°C, and losing 11.7 kJ as it cools, the new temperature is determined by:

ΔT = Q / (mc)

Where Q = 11,700 J, m = 200 g, c = 4.184 J/g·°C:

ΔT = 11,700 J / (200 g × 4.184 J/g·°C) ≈ 13.96°C

The new temperature is:

T_final = T_initial - ΔT = 20°C - 13.96°C ≈ 6.04°C

The water cools to approximately 6.04°C, illustrating principles of heat transfer and thermodynamics.

Calorimeter Heat Capacity and Combustion Reactions

Problem 4 calculates the heat capacity of a bomb calorimeter based on the combustion of benzoic acid. Given the mass of benzoic acid (0.7521 g), its heat of combustion (-26.42 kJ/g), temperature rise (3.60°C), and the mass of water (1,000 g), the total heat released is:

Q = mass × heat of combustion = 0.7521 g × -26.42 kJ/g ≈ -19.87 kJ

The heat capacity (C) of the calorimeter is obtained via:

C = Q / ΔT = 19,870 J / 3.60°C ≈ 5,519 J/°C

This value reflects the calorimeter's capacity to absorb heat during combustion, excluding the water's contribution.

Endothermic and Thermodynamic Processes

Problem 5 asks identifying an endothermic process. Endothermic reactions absorb heat from surroundings, characterized by positive enthalpy change. Among the options, process B (H2O(g) → H2O(l)) involves condensation, which releases heat; thus, it is exothermic. The endothermic process is indicated as the one involving heat absorption: in this case, process A (O2(g) + 2H2(g) → 2H2O(g)) is the formation of water vapor from gases, which requires energy input, making it endothermic.

Bomb Calorimetry and Enthalpy Calculations

Problem 6 involves a combustion of magnesium in a bomb calorimeter. With a mass of 0.1326 g and a total heat capacity of 5,760 J/°C, and a temperature increase of 0.570°C, the enthalpy change per mole is calculated as follows:

Q = C × ΔT = 5,760 J/°C × 0.570°C ≈ 3,283.2 J

Number of moles of Mg:

n = 0.1326 g / 24.305 g/mol ≈ 0.00546 mol

Enthalpy of combustion per mole:

ΔH = Q / n ≈ 3,283.2 J / 0.00546 mol ≈ -601,241 J/mol or -601.24 kJ/mol

Enthalpy of Formation through Hess's Law

Problem 7 asks to find the standard enthalpy of formation of ethanol. Using known enthalpies of combustion of ethanol, CO2, and H2O, and applying Hess's Law, the enthalpy of formation of ethanol is derived:

ΔHf (C2H5OH) = ΔH°combustion - [2 × ΔH°f (CO2) + 3 × ΔH°f (H2O)]

Consolidated calculations estimate ΔHf of ethanol as approximately -278 kJ/mol, aligning with standard literature values.

Organic Compound Formation and Reaction Enthalpies

Problems 8 and 9 analyze the enthalpies of glycine and methanol formation. Utilizing combustion data and Hess's Law, the enthalpy of formation for glycine and methanol are computed, with values close to literature data: approximately -523 kJ/mol for glycine and -201 kJ/mol for methanol, respectively, verifying their stability and energetic profiles in biological and chemical processes.

Gas Laws and Molar Calculations

Problems 11 and 12 employ ideal gas law principles to calculate final volume and temperature changes under constant pressure. For example, in problem 11, the final volume when nitrogen gas is heated from 20°C to 220°C is calculated via:

V2 = V1 × (T2 / T1)

With V1 = 32.4 L, T1 = 293 K, T2 = 493 K:

V2 ≈ 32.4 L × (493 / 293) ≈ 54.58 L

This demonstrates how temperature affects gas volumes according to Charles's Law.

Other Gas Calculations and Molecular Formulas

Problems 13-19 further explore gas behavior under changes in pressure, temperature, and volume, as well as molecular formulas and reaction stoichiometry. Calculations reveal the relationships between variables, enabling accurate predictions of gas behavior in chemical systems.

Reaction Stoichiometry and Gas Production

Problem 19 estimates the amount of MnO2 needed to produce a certain volume of chlorine gas under specified conditions, applying ideal gas law and molar volume at STP, emphasizing the importance of reaction stoichiometry and molar relationships.

Conclusion

The problems addressed demonstrate the application of fundamental principles of thermodynamics, calorimetry, and gas laws in various chemical contexts. Mastery of these calculations not only deepens understanding of energetic and molecular behaviors but also prepares students for practical and theoretical challenges in chemistry and related fields. Accurate computation, integrating concepts such as heat transfer, reaction enthalpies, and gas behaviors, showcases the importance of these principles in understanding real-world chemical phenomena and laboratory practices.

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