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11: Chemical Reactions

Last updated

Mar 21, 2025
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- 10.13: Determining Molecular Formulas
- 11.1: Word Equations

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11.1: Word Equations
This page explores the evolution of recipe documentation from handwritten books to online formats, drawing parallels with writing chemical equations. It emphasizes the transformation of reactants into products in chemical reactions, indicated by a yield sign rather than an equal sign. The text provides examples such as tarnishing silver and burning methane gas, noting that while word equations describe reactions qualitively, they lack quantitative data.
11.2: Chemical Equations
This page draws a parallel between shrimp gumbo and writing chemical equations, highlighting how chemical equations depict reactions with reactants and products in their formulas. It emphasizes the skeleton equation's role in showing formulas without relative amounts and underscores the importance of knowing the physical states of substances. Additionally, it outlines various symbols used in chemical equations for states of matter, catalysts, and heating indicators.
11.3: Balancing Equations
This page explains how to balance chemical equations, focusing on the law of conservation of mass, which requires equal numbers of atoms on both sides of a reaction. It offers guidelines for writing and balancing equations, emphasizing the importance of correct formulas, atom counting, and using the lowest coefficient ratios. An example of balancing the reaction between lead(II) nitrate and sodium chloride is provided to illustrate the process.
11.4: Combination Reactions
This page explains the enhanced function of a wheel rim when combined with a tire for a safe ride. It also discusses combination reactions, where multiple substances unite to form new ones, such as elements reacting with oxygen to create oxides. The text provides examples of balancing equations for these reactions, including potassium oxide formation and the reaction of calcium oxide with water.
11.5: Decomposition Reactions
This page discusses Antoine Lavoisier's contributions to modern chemistry, focusing on decomposition reactions. It defines a decomposition reaction as the breakdown of a compound into simpler substances, needing energy input. Examples include the transformation of mercury(II) oxide into mercury and oxygen, and calcium carbonate into calcium oxide and carbon dioxide. It also mentions that various compounds can decompose through methods like heat or electric current.
11.6: Combustion Reactions
This page provides an overview of combustion reactions, emphasizing their need for oxygen and energy release. It discusses examples like roasting marshmallows and the combustion of hydrocarbons, specifically propane and ethanol, producing carbon dioxide and water. The Hindenburg disaster serves as a historical example of combustion with hydrogen. Additionally, review questions are included to enhance understanding of the topic.
11.7: Single Replacement Reactions
This page discusses the tarnishing of silver as a chemical reaction between silver and hydrogen sulfide, resulting in silver sulfide formation. It reviews single-replacement reactions, where elements substitute one another in compounds, with examples like magnesium displacing copper, zinc reacting with hydrochloric acid, and chlorine replacing bromine. The content also highlights the activity series, which ranks the reactivity of metals and halogens.
11.8: Activity Series
This page explains the reactivity differences between sodium and silver, highlighting that sodium reacts violently with water while silver does not. It introduces the "activity series," ranking metals by their reactivity, which helps predict single-replacement reactions. An example is provided showing that nickel can replace lead but not iron in reactions. The page emphasizes the role of the activity series in predicting how metals interact with water and acids.
11.9: Double Replacement Reactions
This page discusses barter as an analogy for double-replacement chemical reactions, highlighting the exchange of ions between compounds to form new substances. These reactions typically occur in aqueous solutions and can produce solid precipitates, gases, or molecular compounds. Examples include the reaction between potassium iodide and lead (II) nitrate, resulting in a yellow precipitate, and the interaction between sodium sulfide and hydrochloric acid producing hydrogen sulfide gas.

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