Chapter 3 Stoichiometry

Chapter 3Stoichiometry Calculation with Chemical Formulas and Equations

OutlinesChemical EquationsPatterns of Chemical ReactivityAtomic and Molecular WeightsThe MoleEmpirical Formula from AnalysisQuantitative Information from Balanced EquationsLimiting Reactants

Chemical EquationsA chemical equation is written to represent a chemical reaction. The following example illustrates the information that is conveyed by a chemical equation. Methane is the principal component of natural gas. Methane burns in the presence of oxygen to produce carbon dioxide and water. In this reaction methane and oxygen are the reactants, and carbon dioxide and water are the products. The chemical equation corresponding to this reaction is:

Chemical EquationsAlthough this equation is a qualitative statement of fact, "Methane and oxygen react to form carbon dioxide and water," it is not quantitatively correct. The law of conservation of mass has not been obeyed. To correct this, we must balance the equation by changing coefficients where appropriate.

Chemical EquationsCoefficients are the numbers that appear to the left of the chemical symbols. When no number appears to the left of a chemical symbol, as is the case for each reactant and product in this equation, the coefficient is presumed to be one. This type of equation can be balanced by inspection. The first disparity to be corrected is the number of hydrogen atoms. In the unbalanced equation there are four hydrogen atoms to the left of the arrow and only two to the right. To fix this, we change the coefficient of water, on the right, to a 2.CH4 + 2O2 CO2 + 2H2O

Chemical EquationsThe ability to balance equations proficiently is vitally important. Some tips:Adjust the coefficients for compounds first and elements last. For now, balance equations using only whole numbers.

Patterns of Chemical ReactivityElements in the same group in the periodic table tend to have very similar chemical properties. Sodium and potassium, for example, each react with water in the same way to produce the corresponding metal hydroxide and hydrogen gas.

Patterns of Chemical ReactivityIn fact, all of the members of group 1A (alkali metals) react with water in the same way. The equation corresponding to this general reaction is

where M can represent any group 1A metal.

Patterns of Chemical ReactivitySodium and Potassium in Water

Patterns of Chemical ReactivityBecause there are so many different chemical reactions, it is helpful to organize them by recognizing that many reactions fall into one of a small number of categories. For instance, 2NaN3 2Na + 3N2 CaCO3 CaO + CO2 2H2O2 2H2O + O2 appear to be very different equations.

Patterns of Chemical ReactivityThe reactions they represent, though, all fall into the category of decomposition reactions. Essentially the opposites of decomposition, combination reactions are those in which two or more reactants react to form a single product.

Patterns of Chemical Reactivity

Patterns of Chemical ReactivityOther categories of common reactions include combustion reactions and oxidation-reduction reactions. Combustion reactions of hydrocarbons and related compounds constitute an enormous number of reactions, many of which are very useful. Combustion of such compounds involves combination with oxygen to produce carbon dioxide and water. We have already seen an example of a combustion reaction in Section 3.1, namely the combustion of methane.

Patterns of Chemical ReactivityOther substances can undergo combustion, too. For example, hydrogen combines with oxygen to produce only water.

Patterns of Chemical ReactivityElemental sulfur combines with oxygen to produce sulfur dioxide.

Patterns of Chemical ReactivityAirbags

Atomic and Molecular WeightsThe atomic mass scale gives the mass of each element relative to the mass of 12C. The amu is defined by assigning a mass of 12 amu as the mass of a 12C atom. 1 amu = 1.66054 x 10-24 g1 g = 6.02214 x 1023 amuThe atomic mass unit scale allows us to use a chemical formula to determine the percentage composition by mass of a compound.

Atomic and Molecular WeightsFor instance, each carbon dioxide, CO2, molecule consists of one carbon atom and two oxygen atoms. Oxygen has a mass of 15.9949 amu (an atom of 16O has a mass of 15.9949/12 that of an atom of 12C), the percentage composition of CO2 is 27.28 percent carbon by mass and (2 x 15.9949amu)/(12amu + 2 x 15.9949amu) x 100% = 72.72 percent oxygen by mass.

Atomic and Molecular WeightsAverage atomic massMost elements occur in nature as mixtures of isotopesThe atomic mass of an element is determined using its various isotopes and their relative abundancesFormula and Molecular weightThe formula weight is the sum of the atomic weights of each atom in its chemical formula If he chemical formula of a substance is its molecular formula, the formula weight is also called the molecular weight

Atomic and Molecular WeightsPercentage composition from formulasMolecular weight: sum of atomic weights of each atom in chem. formulaChem. formula percentage compositionEx: Percentage composition of C12H22O11: MW of C12H22O11: 342 g/mol

Atomic and Molecular Weights

The MoleJust as a baker uses "dozen" to mean twelve, a chemist uses "mole" (abbreviated mol) to mean 6.022 x 1023. Further, just as the word dozen can apply to any collection of twelve objects, the word mole can apply to any collection of 6.022 x 1023 objects, whether they be atoms, molecules, or ions. Using the atomic mass unit scale, we can determine the mass of a mole of water molecules.

The MoleConverting this to grams gives us a number of more convenient magnitude.

The MoleIt is not a coincidence that we ended up with the same number in grams as the original number of amus. Just as one molecule of water has a mass of 18.02 amu, one mole of water molecules has a mass of 18.02 grams. Likewise, the mass of a molecule of CO2 is 44.01 amu, and the mass of a mole of CO2 molecules is 44.01 grams. Half of that mass would contain half as many particles. T

The Mole

Empirical Formula from AnalysisChemical substances such as water, carbon dioxide, hydrogen peroxide, and sodium chloride have formulas that describe the relative amounts of their constituent elements. The formula for water, H2O, indicates that a water molecule contains two hydrogen atoms and one oxygen atom.

Empirical Formula from AnalysisLikewise, the formula for carbon dioxide indicates that each CO2 molecule contains one carbon and two oxygen atoms. These two formulas are molecular formulas. They tell exactly how many of each type of atom make up a molecule. Hydrogen peroxide's molecular formula is H2O2. Each hydrogen peroxide molecule is made up of two hydrogen atoms and two oxygen atoms.

Empirical Formula from AnalysisThere is another way to express the formula for hydrogen peroxide, though. The empirical formula tells us the relative number of atoms of each type in that molecule. Hydrogen peroxide, then, with a ratio of hydrogen atoms to oxygen atoms of 1:1, has an empirical formula of HO. Obviously, an empirical formula does not contain as much information as a molecular formula.

Empirical Formula from AnalysisMolecular substances have both molecular formulas and empirical formulas. In many cases, as with water and carbon dioxide, they are the same. Ionic compounds do not exist as molecules and, as a result, do not have molecular formulas. Rather, ionic substances such as sodium chloride and magnesium chloride have only empirical formulas, NaCl and MgCl2, respectively.

Empirical Formula from AnalysisJust as you were able to use a formula to calculate percent composition of a compound, you can use percent composition to determine an empirical formula. Determine the empirical formula of a compound that is 81.32 percent carbon, 5.12 percent hydrogen, and 13.56 percent oxygen by mass.

Empirical Formula from AnalysisAssume total mass = 100 gC = 81.32 g H = 5.12 g O = 13.56 gC:H:O = 81.32/12 : 5.12/1 : 13.56/16C:H:O = 6.7767 : 5.12 : 0.8475 = 8:6:1Empirical formula of a compound: C8H6O

Empirical Formula from AnalysisGeneral procedure for determining empirical formulas:

Mass % elementsAssume 100 g sampleGrams of each elementUse atomic weightsMoles of each elementCalculate mol ratioEmpirical formulaGiven:Find:

Quantitative Information from Balanced EquationsIn addition to being both a qualitative and quantitative statement of fact, a balanced chemical equation is an algebraic equality in which the arrow is the equal sign. This means that we can manipulate a balanced chemical equation in much the same way as we can manipulate any other algebraic equality. For instance, we can multiply one of the reactants by a numberas long as we multiply everything else in the equation by that same number.

Quantitative Information from Balanced EquationsThe quantitative statement of fact is, "Two molecules of hydrogen react with one molecule of oxygen to produce two molecules of water." Or, "Two moles of hydrogen react with one mole of oxygen to produce two moles of water." Suppose we wanted to know how many moles of water would be produced by the reaction of four moles of hydrogen? We can simply multiply the amount of hydrogen by 2as long as we multiply everything else by 2 as well.

Quantitative Information from Balanced EquationsThe result tells us that four moles of hydrogen would react with two moles of oxygen and would produce four moles of water. Balanced equations can be used to make predictions about amounts of reactants consumed and/or products produced by a chemical reaction.

Quantitative Information from Balanced EquationsConsider the combustion of propane. How many grams of carbon dioxide will be produced by the combustion of 25.0 grams of propane?

Sample ExerciseHow many grams of water are produced in the combustion of 1.00 g of glucoseSolution:Convert grams of C6H12O6 mol of C6H12O6 Convert moles of C6H12O6 moles of H2OConvert moles of H2O grams of H2O

Sample Exercise

Limiting ReactantsWhen chemists carry out reactions, they seldom do so by using stoichiometric amounts of reactants. In general, one reactant will be consumed completely and will limit the amount of product that can be formed.

Limiting ReactantsLimitting reactants

Limiting ReactantsThe amount of product that can be produced if the limiting reactant is consumed completely is the theoretical yield. In practice, the theoretical yield is seldom achieved. Under real world conditions a variety of factors, including experimental error and efficiency of the reaction, make it impossible to produce the theoretical amount.

Limiting ReactantsThe smaller amount that is actually produced is called the actual yield. The percent yield is calculated using the actual (measured) and theoretical (calculated) yields and is a measure of the efficiency of the overall process.

Limiting Reactants

Limiting ReactantsLimiting Reagents

Thank You

Documents

Chapter 3 Stoichiometry