6: Chemical Composition
- Page ID
- 47420
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)Chapter 6 is concerned with the amounts of substances which participate in chemical reactions, the quantities of heat given off or absorbed when reactions occur, and the volumes of solutions which react exactly with one another. These seemingly unrelated subjects are discussed together because many of the calculations involving them are almost identical in form. The same is true of the density calculations, and of the calculations involving molar mass and the Avogadro constant.
- 6.1: Prelude to Chemical Composition - How Much Sodium?
- This page highlights the importance of understanding composition, illustrating how substances combine and the implications of these interactions. It emphasizes practical applications in evaluating materials like sodium in diets, iron in steel, hydrogen in fuel, and chlorine in freon, all of which support informed decisions in health, industry, and environmental issues.
- 6.2: Counting Nails by the Pound
- This page covers methods for indirectly counting molecules via weighing, emphasizing Avogadro's number and the relationship between mass units and molar mass. It explains how to calculate molar mass, using examples like water, methane, sulfur dioxide, and alum. The page illustrates the process of determining molar mass from atomic masses and provides practical exercises to reinforce these concepts, culminating in the calculation of alum's molar mass at 474 grams per mole.
- 6.3: Counting Atoms by the Gram
- This page provides an overview of the mole concept and Avogadro's number (\(6.02 \times 10^{23}\)), essential in chemistry for quantifying particles. It details methods to convert between the number of particles, moles, and grams using molar mass. It emphasizes the significance of these conversions in chemical calculations and includes example problems to illustrate the practical application of these principles.
- 6.4: Counting Molecules by the Gram
- This page covers key concepts in chemistry related to molecular mass, formula mass, and molar mass, illustrating their importance in laboratory calculations. It details how to calculate molecular mass using ethanol as an example and explains formula mass in ionic compounds like calcium phosphate. Additionally, it provides methods for converting between moles, mass, and particles, using examples of various compounds.
- 6.5: Chemical Formulas as Conversion Factors
- This page explains the use of chemical formulas as conversion factors to relate moles of molecules to moles of atoms, 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.
- 6.6: Mass Percent Composition of Compounds
- This page covers how to determine the percent composition of elements in a compound based on mass data. It explains percent by mass through a practical example involving peanut butter and provides a detailed calculation for a zinc-oxygen compound. The importance of the percent values summing to 100% is highlighted, and an exercise on sulfuric acid helps reinforce the concepts presented.
- 6.7: Mass Percent Composition from a Chemical Formula
- This page explains determining the percent composition of elements in a compound through its chemical formula. It describes calculating each element's mass in one mole, dividing by the molar mass, and multiplying by 100%. An example with dichlorine heptoxide illustrates this calculation, showing the total percent by mass sums to 100%. Additionally, the concept is applied to find the mass of elements in a compound sample, emphasizing the significance of percent composition.
- 6.8: Calculating Empirical Formulas for Compounds
- 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.
- 6.9: Calculating Molecular Formulas for Compounds
- 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.