9: Stoichiometry and the Mole
- Page ID
- 568675
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)- 9.1: Avogadro's Number
- This page explains Avogadro's number, \(6.02 \times 10^{23}\), which quantifies the number of representative particles in a mole, allowing chemists to measure atoms and molecules. It discusses the mole as the SI unit for substance amount, with representative particles being atoms for elements and molecules for compounds. Additionally, it mentions National Mole Day, celebrating this key chemistry concept.
- 9.2: Conversions Between Moles and Particles
- This page explains conversion methods between moles, atoms, and molecules, emphasizing the convenience of moles for simplifying calculations. It provides examples on converting carbon atoms to moles and determining hydrogen atoms in water and sulfuric acid. The importance of knowing chemical formulas for accurate calculations is highlighted, accompanied by step-by-step calculation processes. The document concludes with review questions on the discussed concepts.
- 9.3: Molar Mass
- This page explains how to calculate the amount of a substance needed for a solution based on its molar mass, using carbon dioxide and calcium nitrate as examples. It highlights the relationship between moles and grams and underscores the importance of molar mass for accurate measurements in solution preparation. The page also includes review questions to help reinforce understanding of molar mass and its calculations.
- 9.4: Conversions Between Moles and Mass
- This page emphasizes the link between moles and mass using molar mass as a conversion factor, providing practical examples such as calculating the mass needed for specific moles of calcium chloride. The content stresses the significance of precision in calculations and significant figures in laboratory work, concluding with a review of mole and mass conversions.
- 9.5: Conversions Between Mass and Number of Particles
- This page details the conversion between mass and the number of particles using moles, which involves changing grams to moles and then moles to particles. An example with chlorine gas (\(\ce{Cl2}\)) is provided to illustrate this process. Finally, the page includes a review section with questions to enhance comprehension.
- 9.6: Chemical Formulas as Conversion Factors - "Big Particle/Little Particle"
- This page explains the use of chemical formulas as conversion factors to relate moles of molecules (the big particle) to moles of atoms (the little particle), demonstrating constant atomic ratios in compounds like water and ethanol. It introduces the mole concept with practical examples and exercises, reinforcing stoichiometric relationships in chemical calculations through specific numerical values for different substances.
- 9.7: Stoichiometry
- This page introduces stoichiometry, emphasizing its role in using balanced chemical equations to relate reactants and products. It draws parallels to recipes, demonstrating how to convert quantities using ratios from balanced equations. Through examples like pancake ingredient calculations, it highlights the importance of balanced equations in predicting reaction outcomes by facilitating conversions between moles and molecules, underscoring the need for precision in stoichiometric relationships.
- 9.8: Mole-to-Mole Conversion - The Heart of Stoichiometry!
- This page highlights the significance of balancing chemical equations by focusing on moles, demonstrating that coefficients indicate the mole ratios of substances in reactions. It explains that balanced equations can be used to derive conversion factors for calculations, and provides examples for interpreting equations and conducting mole-to-mole conversions.
- 9.9: Mole-Mass and Mass-Mass Calculations
- This page covers the calculation of masses and moles in balanced chemical equations, emphasizing mole-mass conversions. It details the steps for converting grams to moles and vice versa, highlighting the necessity of balancing equations in moles. Additionally, mass-mass calculations are introduced, illustrating a systematic approach that involves converting mass to moles, applying mole ratios, and converting back to mass.
- 9.10: Yields
- This page covers theoretical yield, actual yield, and percent yield in chemical reactions, explaining that theoretical yield is based on theoretical calculations, while actual yield reflects what is produced in practice. Percent yield measures efficiency by comparing these two.
- 9.11: Limiting Reagents
- This page explores the limiting reagent (reactant) in chemical reactions, highlighting its role in restricting product formation when one reactant is fully consumed. It details how to identify the limiting reagent using mole-mass calculations and compares product yields from different reactants. The page also outlines how to compute the mass of the product produced and the remaining quantities of reactants after reactions.
- 9.12: Mass Percent Composition from a Chemical Formula
- This page details the process of calculating the percent composition of elements in a compound using its chemical formula, including an example with dichlorine heptoxide. It describes how to determine the mass of each element per mole and convert it into a percentage of the compound's molar mass. The page also illustrates using percent composition to find the mass of an element in a sample and includes a practice exercise with barium fluoride.
- 9.13: Calculating Empirical Formulas
- This page explains empirical formulas, detailing how to determine them from percent composition via elemental analysis. It presents a methodical approach including sample size assumption, percentage conversion to grams, mole calculation, and obtaining whole-number ratios. An example of iron(III) oxide (\(Fe_2O_3\)) is given, along with an exercise on mercury and chlorine to enhance understanding.
- 9.14: Calculating Molecular Formulas
- This page explains the difference between empirical and molecular formulas, highlighting that molecular formulas provide the actual number of atoms while empirical formulas denote the simplest ratio. It outlines how to derive molecular formulas from percent composition and molar mass, using examples such as glucose and sucrose.