17: Thermochemistry
- Page ID
- 53868
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)- 17.1: Chemical Potential Energy
- This page discusses gunpowder's composition and explosive nature, its development in the ninth century by the Chinese, and differentiates between potential and kinetic energy. It explains chemical potential energy found in substances such as gasoline and dynamite, highlighting how energy is released through combustion. Additionally, it addresses the properties of dynamite, particularly its instability due to nitroglycerin, and methods to stabilize it to avoid unintended explosions.
- 17.2: Heat
- This page discusses how blacksmiths shape iron by heating it, which makes the metal more workable due to increased atomic movement. It explains that heat energy transfers between objects until thermal equilibrium is reached. Additionally, thermochemistry, a study of energy changes during chemical reactions and state changes, is mentioned. It highlights the relationship between bond formation and breaking, which impacts the total chemical potential energy of a system.
- 17.3: Exothermic and Endothermic Processes
- This page outlines basic thermochemistry principles using a campfire analogy. It explains exothermic and endothermic processes, emphasizing energy conservation during changes. The system represents the matter studied, while the surroundings include everything else. Endothermic processes absorb heat, lowering surrounding temperatures; exothermic processes release heat, raising them.
- 17.4: Heat Capacity and Specific Heat
- This page explains heat capacity and specific heat, emphasizing their effects on temperature changes in objects. It illustrates how mass and chemical composition influence heating rates, using a wading pool versus a larger swimming pool as an example. Water's high specific heat (4.18 J/g°C) requires more energy to warm up compared to other substances like metals, making it an effective coolant and contributing to moderate coastal climates.
- 17.5: Specific Heat Calculations
- This page discusses the role of water in cooling car engines through effective heat absorption. It explains specific heat's influence on temperature changes, detailed by the equation \( q = c_p \times m \times \Delta T \). An example illustrates specific heat calculations for cadmium, showing it is similar to other metals.
- 17.6: Enthalpy
- This page explains enthalpy as the heat content of a system at constant pressure, emphasizing its role in chemical reactions. It highlights how enthalpy is affected by the energy needed to break bonds, the energy released in bond formation, and factors like material quantity and physical states. The page notes that reversing a reaction retains the same enthalpy value with an opposite sign, and discusses how enthalpy change identifies reactions as endothermic or exothermic.
- 17.7: Calorimetry
- This page explains calorimetry, which measures heat transfer in chemical reactions and physical processes using a calorimeter. Originally, food calories were measured with a bomb calorimeter, but are now determined based on protein, carbohydrate, and fat content. Temperature changes in the surrounding liquid reveal if a reaction is exothermic or endothermic, allowing for the calculation of heat transfer using the liquid's mass, specific heat, and temperature change.
- 17.8: Thermochemical Equations
- This page discusses the rising costs of home heating and the importance of choosing the right fuel based on thermochemical data. It highlights the exothermic nature of methane combustion, releasing 890.4 kJ, compared to the endothermic decomposition of calcium carbonate, which absorbs 177.8 kJ. Understanding thermochemical equations and the heat of reaction is vital for energy transfer and efficiency in heating systems.
- 17.9: Stoichiometric Calculations and Enthalpy Changes
- This page discusses growing concerns over manufacturing emissions and the subsequent need for pollutant reduction equipment. Studies focus on measuring product output and energy changes associated with pollution. It highlights the use of stoichiometry in analyzing enthalpy changes, specifically through calculating energy released from the combustion of sulfur dioxide, demonstrating the conversion from mass to moles for accurate assessments.
- 17.10: Heats of Fusion and Solidification
- This page explains the heat transfer process when holding an ice cube, highlighting how heat energy from the hand melts the ice without changing temperature due to the phase change. It covers the concepts of molar heat of fusion and solidification, noting that their values are equal but opposite. Additionally, it mentions calculations to determine heat absorbed or released, exemplified by the melting of 31.6 g of ice, which absorbs around 10.5 kJ.
- 17.11: Heats of Vaporization and Condensation
- This page discusses natural resources for electric power generation, emphasizing renewable energy sources such as geothermal power. It covers the concepts of heat of vaporization and condensation, explaining the energy dynamics during these state changes and highlighting that vaporization demands more energy than fusion.
- 17.12: Multi-Step Problems with Changes of State
- This page explains the energy-intensive process of converting ice at -30°C to steam at 140°C, which involves multiple steps: heating ice, melting it, heating water, vaporizing water, and heating steam. The total energy required for this transformation is calculated to be 133.4 kJ, with the highest energy absorption occurring during the vaporization phase.
- 17.13: Heat of Solution
- This page emphasizes the importance of slowly adding concentrated sulfuric acid to water during dilutions to avoid splattering due to heat release. It explains the concept of heat of solution and provides examples of thermal reactions in hot and cold packs. The molar heat of solution is defined as the heat change per mole of solute dissolved, and sample calculations illustrate how to calculate temperature changes from the heat released when substances like sodium hydroxide dissolve in water.
- 17.14: Heat of Combustion
- This page discusses the use of ethanol in gasoline to enhance fuel efficiency due to its high octane rating, despite potential increases in air pollution. It explains the concept of molar heat of combustion, which measures the energy released during combustion reactions involving carbon and oxygen, producing carbon dioxide and water. The molar heat of combustion for ethanol is noted as 1370 kJ/mol, and bomb calorimetry is mentioned as a method for this measurement.
- 17.15: Hess's Law of Heat Summation
- This page explains the complexities involved in calculating the energy associated with operating an acetylene torch. It highlights the use of Hess's law to determine enthalpy changes indirectly through combustion reactions. By analyzing the combustion of carbon, hydrogen, and acetylene, the heat of formation for acetylene is found to be 228.3 kJ, which signifies that the reaction is endothermic and absorbs heat.
- 17.16: Standard Heat of Formation
- This page discusses the Hope diamond, valued at $350 million, and compares it to graphite, which is much cheaper. It explains that the differences between these two forms of carbon arise from their distinct organizational structures and formation conditions. It also defines standard heat of formation (\( \Delta H^\text{o}_\text{f} \)), describing it as the enthalpy change involved in forming a mole of a compound from its elements at specified standard conditions.
- 17.17: Calculating Heat of Reaction from Heat of Formation
- This page discusses the global sourcing and price control of natural diamonds, highlighting the rise of synthetic diamonds made from carbon for industrial use. It also covers thermodynamic concepts, specifically the standard heat of reaction (\(\Delta H^\text{o}\)), which can be calculated using Hess's law by subtracting the total heats of formation of reactants from products, helping to identify the exothermic or endothermic nature of reactions.