Chemistry Chapter 12

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Chemistry Chapter 12
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• stoichiometry.
• that portion of chemistry dealing with numerical
  relationships in chemical reactions the
  calculation of quantities of substances involved
  in chemical equations

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• A balanced chemical equation can be interpreted
  in terms of different quantities, including
  numbers of atoms, molecules, or moles mass and
  volume.

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• Mass and atoms are conserved in every chemical
  reaction.

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• Number of Atoms
• At the atomic level, a balanced equation
  indicates that the number and type of each atom
  that makes up each reactant also makes up each
  product.

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• Number of Molecules
• The balanced equation indicates that one
  molÂecule of nitrogen reacts with three molecules
  of hydrogen. Nitrogen and hydrogen will always
  react to form ammonia in a 132 ratio of
  molecules.

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• Moles
• A balanced chemical equation also tells you the
  number of moles of reactants and products. The
  coefficients of a balanced chemical equation
  indicate the relative numbers of moles of
  reactants and products in a chemical reaction

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• Mass
• A balanced chemical equation obeys the law of
  conservation of mass. This law states that mass
  can be neither created nor destroyed in an
  ordinary chemical or physical process.

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• Volume
• If you assume standard temperature and pressure,
  the equation also tells you about the volumes of
  gases. Recall that 1 mol of any gas at STP
  occupies a volume of 22.4 L.

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• However, molecules, formula units, moles, and
  volumes are not necessarily conservedalthough
  they may be. Consider, for example, the formation
  of hydrogen iodide,

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• (5) Key Concept How is a balanced equation
  similar to a recipe?Hint
• (6) Key Concept How do chemists use balanced
  equations?Hint
• (7) Key Concept Chemical reactions can be
  described in terms of what quantities?Hint
• (8) Key Concept What quantities are always
  conserved in chemical reactions?Hint

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• (9)Interpret the given equation in terms of
  relative numbers of representative particles,
  numbers of moles, and masses of reactants and
  products.
• 2K (s) 2H2O (l) ? 2KOH (aq) H2 (g)

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• (10)Balance this equation C2H5OH(l) O2 (g) ?
  CO2 (g) H2O(g). Show that the balanced equation
  obeys the law of conservation of mass.

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• mole ratio
• a conversion factor derived from the coefficients
  of a balanced chemical equation interpreted in
  terms of moles

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• In chemical calculations, mole ratios are used to
  convert between moles of reactant and moles of
  product, between moles of reactants, or between
  moles of products.

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• When multiplying and dividing measurements, the
  rounded answer can have no more significant
  figures than the least number of significant
  figures in any measurement in the calculation.

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• When adding and subtracting measurements, the
  answer can have no more decimal places than the
  least number of decimal places in any measurement
  in the problem. The difference of 8.78 cm - 2.2
  cm 6.58 cm is rounded to 6.6 cm (one decimal
  place).

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• In a typical stoichiometric problem, the given
  quantity is first converted to moles. Then the
  mole ratio from the balanced equation is used to
  calculate the number of moles of the wanted
  substance. Finally, the moles are converted to
  any other unit of measurement related to the unit
  mole, as the problem requires.

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• Recall from Chapter 10 that the mole can be
  related to other quantities as well. For example,
  1 mol 6.02 10 23 representative particles,
  and 1 mol of a gas 22.4 L at STP. These two
  relationships provide four more conversion
  factors that you can use in stoichiometric
  calculations.

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• (21) Key Concept How are mole ratios used in
  chemical calculations?Hint
• (22) Key Concept Outline the sequence of steps
  needed to solve a typical stoichiometric
  problem.Hint
• (23)Write the 12 mole ratios that can be derived
  from the equation for the combustion of isopropyl
  alcohol.
• 2C3H7OH( l ) 9O2 (g) ? 6CO2 (g) 8H2O(g)

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• 24)The combustion of acetylene gas is represented
  by this equation
• 2C2H2 (g) 5O2 (g) ? 4CO2 (g) 2H2O(g)
• How many grams of CO2 and grams of H2O are
  produced when 52.0 g C2H2 burns in oxygen?

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• In a chemical reaction, an insufficient quantity
  of any of the reactants will limit the amount of
  product that forms.

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• When one molecule (mole) of N2 reacts with three
  molecules (moles) of H2, two molecules (moles) of
  NH3 are produced. What would happen if two
  molecules (moles) of N2 reacted with three
  molecules (moles) of H2? Would more than two
  molecules (moles) of NH3 be formed?

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• Before the reaction takes place, nitrogen and
  hydrogen are present in a 23 molecule (mole)
  ratio. The reaction takes place according to the
  balanced equation. One molecule (mole) of N2
  reacts with three molecules (moles) of H2 to
  produce two molecules (moles) of NH3. At this
  point, all the hydrogen has been used up, and the
  reaction stops. One molecule (mole) of unreacted
  nitrogen is left in addition to the two molecules
  (moles) of NH3 that have been produced by the
  reaction.

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• limiting reagent
• any reactant that is used up first in a chemical
  reaction it determines the amount of product
  that can be formed in the reaction

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• excess reagent
• a reagent present in a quantity that is more than
  sufficient to react with a limiting reagent any
  reactant that remains after the limiting reagent
  is used up in a chemical reaction

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• Percent Yield

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• theoretical yield
• the amount of product that could form during a
  reaction calculated from a balanced chemical
  equation it represents the maximum amount of
  product that could be formed from a given amount
  of reactant

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• actual yield
• the amount of product that forms when a reaction
  is carried out in the laboratory

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• percent yield
• the ratio of the actual yield to the theoretical
  yield for a chemical reaction expressed as a
  percentage a measure of the efficiency of a
  reaction

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• The percent yield is a measure of the efficiency
  of a reaction carried out in the laboratory.

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• A percent yield should not normally be larger
  than 100. Many factors cause percent yields to
  be less than 100. Reactions do not always go to
  completion when this occurs, less than the
  calculated amount of product is formed. Impure
  reactants and competing side reactions may cause
  unwanted products to form. Actual yield can also
  be lower than the theoretical yield due to a loss
  of product during filtration or in transferring
  between containers. Moreover, if reactants or
  products have not been carefully measured, a
  percent yield of 100 is unlikely.

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• Figure 12.13 Sodium hydrogen carbonate (Na HCO3)
  will decompose when heated. The mass of NaHCO3 ,
  the reactant, is measured. The reactant is
  heated. The mass of one of the products, sodium
  carbonate (Na2 CO3), the actual yield, is
  measured. The percent yield is calculated once
  the actual yield is determined. Predicting What
  are the other products of this reaction?

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• (33) Key Concept In a chemical reaction, how does
  an insufficient quantity of a reactant affect the
  amount of product formed?Hint
• (34) Key Concept How can you gauge the efficiency
  of a reaction carried out in the laboratory?Hint

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• (35)What is the percent yield if 4.65 g of copper
  is produced when 1.87 g of aluminum reacts with
  an excess of copper(II) sulfate?
• 2Al(s) 3CuSO4 (aq) ? Al2 (SO4) 3 (aq) 3Cu(s)

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• 12.1 The Arithmetic of Equations
• A balanced chemical equation provides the same
  kind of quantitative information that a recipe
  does.Hint

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• Chemists use balanced chemical equations as a
  basis to calculate how much reactant is needed or
  product is formed in a reaction.Hint

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• A balanced chemical equation can be interpreted
  in terms of different quantities, including
  numbers of atoms, molecules, or moles mass and
  volume.Hint
• Mass and atoms are conserved in every chemical
  reaction.Hint

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• 12.2 Chemical Calculations
• In chemical calculations, mole ratios are used to
  convert between moles of reactant and moles of
  product, between moles of reactants, or between
  moles of products.Hint

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• In a typical stoichiometric problem, the given
  quantity is first converted to moles. Then the
  mole ratio from the balanced equation is used to
  calculate the moles of the wanted substance.
  Finally, the moles are converted to any other
  unit of measurement related to the unit mole.Hint

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• 12.3 Limiting Reagent and Percent Yield
• In a chemical reaction, an insufficient quantity
  of any of the reactants will limit the amount of
  product that forms.Hint
• The percent yield is a measure of the efficiency
  of a reaction performed in the laboratory.Hint

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• 12.1 The Arithmetic of Equations
• (36)Interpret each chemical equation in terms of
  interacting particles.
• 2KClO3 (s) ? 2KCl(s) 3O2 (g)
• 4NH3 (g) 6NO(g) ? 5N2 (g) 6H2O(g)
• 4K(s) O2 (g) ? 2K2O(s)

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• (37)Interpret each equation in Problem 36 in
  terms of interacting numbers of moles of
  reactants and products.

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• 38)Calculate and compare the mass of the
  reactants with the mass of the products for each
  equation in Problem 36. Show that each balanced
  equation obeys the law of conservation of mass

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• 12.2 Chemical Calculations
• (39)Explain the term mole ratio in your own
  words. When would you use this term?

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• (40)Carbon disulfide is an important industrial
  solvent. It is prepared by the reaction of coke
  with sulfur dioxide.
• 5C(s) 2SO2 (g) ? CS2 (l) 4CO(g)
• How many moles of CS2 form when 2.7 mol C reacts?

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• How many moles of carbon are needed to react with
  5.44 mol SO2?

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• How many moles of carbon monoxide form at the
  same time that 0.246 mol CS2 forms?

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• How many mol SO2 are required to make 118 mol CS2?

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• 41)Methanol (CH3OH) is used in the production of
  many chemicals. Methanol is made by reacting
  carbon monoxide and hydrogen at high temperature
  and pressure.
• CO(g) 2H2 (g) ? CH3OH(g)

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• How many moles of each reactant are needed to
  produce 3.60 10 2 g CH3OH?

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• Calculate the number of grams of each reactant
  needed to produce 4.00 mol CH3OH.

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• How many grams of hydrogen are necessary to react
  with 2.85 mol CO?

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• 42)The reaction of fluorine with ammonia produces
  dinitrogen tetrafluoride and hydrogen fluoride.
• 5F2 (g) 2NH3 (g) ? N2F4 (g) 6HF(g)

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• If you have 66.6 g NH3, how many grams of F2 are
  required for complete reaction?

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• How many grams of NH3 are required to produce
  4.65 g HF?
• How many grams of N2F4 can be produced from 225 g
  F2?

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• 43)What information about a chemical reaction is
  derived from the coefficients in a balanced
  equation?

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• (44)Lithium nitride reacts with water to form
  ammonia and aqueous lithium hydroxide.
• Li3N(s) 3H2O(l) ? NH3 (g) 3LiOH( aq)

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• What mass of water is needed to react with 32.9 g
  Li3N?
• When the above reaction takes place, how many
  molecules of NH3 are produced?

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• Calculate the number of grams of Li3N that must
  be added to an excess of water to produce 15.0 L
  NH3 (at STP).

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• .3 Limiting Reagent and Percent Yield
• (45)What is the significance of the limiting
  reagent in a reaction? What happens to the amount
  of any reagent that is present in an excess?

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• (46)How would you identify a limiting reagent in
  a chemical reaction?

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• (47)In a reaction chamber, 3.0 mol of aluminum is
  mixed with 5.3 mol Cl2 and reacts. The reaction
  is described by the following balanced chemical
  equation.
• 2Al 3Cl2 ? 2AlCl3

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• Identify the limiting reagent for the reaction.
• Calculate the number of moles of product formed.

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• Calculate the number of moles of excess reagent
  remaining after the reaction.

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• (48)Heating an ore of antimony (Sb2S3) in the
  presence of iron gives the element antimony and
  iron(II) sulfide.
• Sb2S3 (s) 3Fe(s) ? 2Sb(s) 3FeS(s)
• When 15.0 g Sb2S3 reacts with an excess of Fe,
  9.84 g Sb is produced. What is the percent yield
  of this reaction?

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• (49)Calcium carbonate reacts with phosphoric acid
  to produce calcium phosphate, carbon dioxide, and
  water.
• 3CaCO3 (s) 2H3PO4 (aq) ? Ca3 (PO4) 2 (aq)
  3CO2 (g) 3H2O(l)

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• How many grams of phosphoric acid react with
  excess calcium carbonate to produce 3.74 g Ca3
  (PO4) 2?

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• Calculate the number of grams of CO2 formed when
  0.773 g H2O is produced.

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• 50)Nitric acid and zinc react to form zinc
  nitrate, ammonium nitrate, and water.
• 4Zn(s) 10HNO3 (aq) ? 4Zn(NO3) 2 (aq) NH4NO3
  (aq) 3H2O(l)

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• How many atoms of zinc react with 1.49 g
  HNO3?Hint
• Calculate the number of grams of zinc that must
  react with an excess of HNO3 to form 29.1 g
  NH4NO3.Hint

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• (51)Hydrazine (N2H4) is used as rocket fuel. It
  reacts with oxygen to form nitrogen and water.
• N2H4 (l) O2 (g) ? N2 (g) 2H2O(g)

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• How many liters of N2 (at STP) form when 1.0 kg
  N2H4 reacts with 1.0 kg O2?Hint

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• How many grams of the excess reagent remain after
  the reaction?Hint

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• 52)When 50.0 g of silicon dioxide is heated with
  an excess of carbon, 32.2 g of silicon carbide is
  produced.
• SiO2 (s) 3C(s) ? SiC(s) 2CO(g)

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• What is the percent yield of this reaction?Hint
• How many grams of CO gas are made?Hint

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• 53)If the reaction below proceeds with a 96.8
  yield, how many kilograms of CaSO4 are formed
  when 5.24 kg SO2 reacts with an excess of CaCO3
  and O2?
• 2CaCO3 (s) 2SO2 (g) O2 (g) ? 2CaSO4 (s)
  2CO2 (g)

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• 54)Ammonium nitrate will decompose explosively at
  high temperatures to form nitrogen, oxygen, and
  water vapor.
• 2NH4NO3 (s) ? 2N2 (g) 4H2O(g) O2 (g)
• What is the total number of liters of gas formed
  when 228 g NH4NO3 is decomposed? (Assume STP.)

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• 55)In an experiment, varying masses of sodium
  metal are reacted with a fixed initial mass of
  chlorine gas. The amounts of sodium used and the
  amounts of sodium chloride formed are shown on
  the following graph.

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• Explain the general shape of the graph.Hint
• Estimate the amount of chlorine gas used in this
  experiment at the point where the curve becomes
  horizontal.Hint

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• 56)The manufacture of compound F requires five
  separate chemical reactions. The initial
  reactant, compound A, is converted to compound B,
  compound B is converted to compound C, and so
  forth. The diagram below summarizes the stepwise
  manufacture of compound F, including the percent
  yield for each step. Provide the missing
  quantities or missing percent yields. Assume that
  the reactant and product in each step react in a
  one-to-one mole ratio.

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• 57)Given a certain quantity of reactant, you
  calculate that a particular reaction should
  produce 55 g of a product. When you perform the
  reaction, you find that you have produced 63 g of
  product. What is your percent yield? What could
  have caused a percent yield greater than 100?