7: Chemical Reactions - Energy, Rates, and Equilibrium
- Page ID
- 86229
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)- 7.1: Energy and Chemical Bonds
- This page defines energy as the capacity to supply heat or perform work, distinguishing between potential and kinetic energy. It highlights energy's significance in daily life, particularly in chemical processes, fuel combustion, and industrial applications. The law of conservation of energy is emphasized, stating that energy changes form but is not created or destroyed. Additionally, the page briefly contrasts energy changes in chemical and nuclear reactions.
- 7.2: Heat Changes during Chemical Reactions
- This page covers bond dissociation energy and its relevance in chemical reactions, explaining how bond formation lowers energy and releases heat (exothermic) while bond breaking requires energy (endothermic). It introduces enthalpy change (ΔH) and its calculation using bond energies of reactants and products.
- 7.3: Exothermic and Endothermic Reactions
- This page covers endothermic and exothermic reactions, detailing their energy dynamics, where endothermic reactions absorb energy (positive \(\Delta H\)) and exothermic reactions release energy (negative \(\Delta H\)). It includes examples, energy diagrams, a summary table of key distinctions, and the impact on surrounding temperature. The page concludes with review questions to enhance comprehension of energy's importance in chemical bonds and reactions.
- 7.4: Why Do Chemical Reactions Occur? Free Energy
- This page covers spontaneous reactions, emphasizing their characteristics, such as a decrease in enthalpy and an increase in entropy. It introduces the Gibbs free energy (\(\Delta G\)) concept, indicating that a negative \(\Delta G\) denotes a spontaneous reaction, while a positive value signals nonspontaneity.
- 7.5: How Do Chemical Reactions Occur? Reaction Rates
- This page covers chemical kinetics, focusing on reaction rates, defined as changes in concentration over time, and activation energy, the minimum energy required for reactions. Effective collisions between reactant particles are necessary for reactions to occur. Potential energy diagrams illustrate energy changes and activation energy barriers during reactions.
- 7.6: Effects of Temperature, Concentration, and Catalysts on Reaction Rates
- This page explores factors influencing chemical reaction rates, focusing on concentration, temperature, and catalysts. It explains that higher concentrations or temperatures increase collision frequency and effectiveness, speeding up reactions. Catalysts reduce activation energy to accelerate reactions without being consumed, illustrated by manganese (IV) oxide in hydrogen peroxide decomposition.
- 7.7: Reversible Reactions and Chemical Equilibrium
- This page discusses chemical equilibrium as a dynamic state where forward and reverse reactions occur at equal rates, exemplified by the hydrogen and iodine reaction forming hydrogen iodide. It emphasizes that at equilibrium, reactant and product concentrations remain constant, and reactions persist in a balanced manner, symbolized by a double arrow. Key insights focus on the continuous nature of reactions despite achieving equilibrium.
- 7.8: Equilibrium Equations and Equilibrium Constants
- This page provides an overview of equilibrium constants (Keq and KP) in chemical reactions. It explains the law of mass action, detailing how to derive the Keq expression from product and reactant concentrations. The relationship between Keq and KP is explored, highlighting how to calculate KP from Keq using temperature and Δn. The page also addresses heterogeneous equilibria, emphasizing that pure solids and liquids are excluded from Keq calculations.
- 7.9: Le Chatelier’s Principle- The Effect of Changing Conditions on Equilibria
- This page explains Le Chatelier's principle, which states that a system at equilibrium will adjust to counteract changes in concentration, pressure, or temperature. It discusses how these stresses affect equilibrium shifts in reactions, particularly in the synthesis of ammonia and calcium carbonate. The effects of adding or removing substances, along with temperature and pressure adjustments, are covered, as well as the role of catalysts.