Balance the following representation of the chemical reaction involved when an airbag deploys.

NaN3 Na + N2

EXAMPLE 6.1 Balancing Equations

SolutionSodium atoms are balanced, but the nitrogen atoms are not. For this sort of problem, we will use the concept of the least common multiple. There are three nitrogen atoms on the left (reactants side) and two on the right (products side). The least common multiple of 2 and 3 is 6. Therefore, we need three N2 and two NaN3:

2 NaN3 Na + 3 N2 (not balanced)

We now have two sodium atoms on the left. We can get two on the right by placing the coefficient 2 in front of Na.

2 NaN3 2 Na + 3 N2 (balanced)

Checking, we count two Na atoms and six N atoms on each side. The equation is balanced.

2 x 3 = 6 N atoms 3 x 2 = 6 N atoms

2 Na atoms 1 Na atoms

The reaction between hydrogen and nitrogen to give ammonia, called the Haber process, is typically the first step in the industrial production of nitrogen fertilizers, represented as

H2 + N2 NH3

Balance the equation.

Exercise 6.1A

EXAMPLE 6.1 Balancing Equations continued

Iron ores like Fe2O3 are smelted by reaction with carbon to produce metallic iron and carbon dioxide, represented as

Fe2O3 + C CO3 + Fe

Balance the equation.

Exercise 6.1B

When a fuel such as methane is burned in sufficient air, the products are carbon dioxide and water, represented as

CH4 + O2 CO2 + H2O (not balanced)

Balance the equation.

EXAMPLE 6.2 Balancing Equations

SolutionIn this equation, oxygen appears in two different products and by itself in 02; we leave the oxygen for last and balance the other two elements first. Carbon is already balanced, with one atom on each side of the equation. For hydrogen, the least common multiple of 2 and 4 is 4, and so we place the coefficient 2 in front of H2O to balance hydrogen. Now we have four hydrogen atoms on each side.

CH4 + O2 CO2 + 2 H2O (not balanced)

Now for the oxygen. There are four oxygen atoms on the right. If we place a 2 in front of O2 on the left, the oxygen atoms balance.

CH4 + 2 O2 CO2 + 2 H2O (balanced)

The equation now has one C atom, four H atoms, and four O atoms on each side; it is balanced.

Butane is a common fuel. Its combustion is represented as

C4H10 + O2 CO2 + H2O

Balance the equation.

Exercise 6.2A

EXAMPLE 6.2 Balancing Equations continued

Does it take more oxygen (per molecule) to burn butane (see Exercise 6.2A) than it does to burn methane? How much more CO2 is produced?

Exercise 6.2B

What volume of oxygen is required to burn 0.556 L of propane if both gases are measured at the same temperature and pressure?

C3H8(g) + 5 O2(g) 3 CO2 (g) + 4 H2O(g)

Using the equation in Example 6.3, calculate the volume of CO2(g) produced when 0.492 L of propane is burned if the two gases are compared at the same temperature and pressure.

Exercise 6.3A

If 10.0 L each of propane and oxygen are combined at the same temperature and pressure, which gas will be left over after reaction? What volume of that gas will remain?

Exercise 6.3B

EXAMPLE 6.3 Volume Relationships of Gases

SolutionThe coefficients in the equation indicate that each volume of C3H8(g) requires 5 volumes of O2(g). Thus, we use 5 L O2(g)/1 L C3H8(g) as the ratio to find the volume of oxygen required.

? L O2(g) = O.556 L C3 H8(g) x = 2.78 L O2(g)5 L O2(g)

1 L C3 H8(g)

Calculate (a) the molecular mass of nitrogen dioxide (NO2) an amber colored gas that is a constituent of smog, and (b) the formula mass of ammonium sulfate [(NH4)2SO4] a fertilizer commonly used by home gardeners.

EXAMPLE 6.4 Calculating Molecular Masses

Solutiona. We start with the molecular formula: NO2. Then, to determine the molecular mass, we need only to

add the atomic mass of nitrogen to twice the atomic mass of oxygen.

Using a calculator, we need only write down the final answer, 46.0 u. That is, we have no need to record the numbers 14.0 and 32.0.

b. We must make certain that all the atoms in the formula unit are accounted for, which means paying particular attention to all the subscripts and parentheses in the formula. The “(NH4)2” means that both the “N” and the “H4” must be multiplied by 2—that is, the formula indicates a total of two N atoms and eight H atoms. Combining the atomic masses, we have

2 x atomic mass of N = 2 x 14.0 u = 28.0 u8 x atomic mass of H = 8 x 1.01 u = 8.08 u1 x atomic mass of S = 1 x 32.0 u = 32.0 u4 x atomic mass of O = 4 x 16.0 u = 64.0 uFormula mass of (NH4)2SO4 = 132.1 u

1 x atomic mass of N = 1 x 14.0 u = 14.0 u2 x atomic mass of O = 2 x 16.0 u = 32.0 u

Formula mass of NO2 = 46.0 u

EXAMPLE 6.4 Calculating Molecular Masses continued

Calculate the formula mass of (a) sodium azide (NaN3) used in automobile airbags, and (b) phosphoric acid (H3PO4).

Exercise 6.4A

Calculate the formula mass of (a) para-dichlorobenzene (C6H4Cl2) used as a moth repellent, and (b) calcium dihydrogen phosphate [Ca(H2PO4)2], used as a mineral supplement in foods.

Exercise 6.4B

How many grams of N2 are in 0.400 moles N2?

Calculate the mass, in grams, of (a) 0.0728 mol silicon, (b) 55.5 mol H2O, and (c) 0.0728 mol Ca(H2PO4)2.

Exercise 6.5

EXAMPLE 6.5 Mole-to-Mass Conversions

SolutionThe molecular mass of N2 is 2 x 14.0 u = 28.0 u. The molar mass of N2 is therefore 28.0 g/mol. Using the molar mass as a conversion factor (red), we have

? g N2 = 0.400 mol N2 x = 11.2 g N2

28.0 g N2

1 mol N2

Calculate the number of moles of Na in a 62.5-g sample of sodium metal.

Calculate the amount, in moles, of (a) 3.71 g Fe, (b) 165 g butatne, C4H10, and (c) 0.100 mol Mg(NO3)2.

Exercise 6.6

EXAMPLE 6.6 Mass-to-Mole Conversions

SolutionThe molar mass of Na is 23.0 g/mol. To convert from a mass in grams to an amount in moles, we must use the inverse of the molar mass as a conversion factor (1 mol Na/23.0 g Na) to get the proper cancellation of units. When we start with grams, we must have grams in the denominator of our conversion factor (red).

? mol Na = 62.5 g Na x = 2.72 mol Na1 mol Na

23.0 g Na

Calculate the density of (a) nitrogen gas and (b) methane (CH4) gas, both at STP.

Calculate the density of He at STP.

Exercise 6.7A

EXAMPLE 6.7 Density of a Gas at STP

Solutiona. The molar ass of N2 gas is 28.0 g/mol. We multiply by the conversion factor 1 mol N2 = 22.4 L,

arranged to cancel units of moles.

b. The molar mass of CH4 gas is (1 x 12.0) g/mol + (4 x 1.01) g/mol = 16.0 g/mol. Again we use the conversion factor 1 mol CH4 = 22.4 L.

x = 1.25 g/L28.0 g N2

1 mol N2

1 mol N2

22.4 L N2

x = 0.714 g/L16.0 g CH4

1 mol CH4

1 mol CH4

22.4 L CH4

Estimate the density of air at STP (assume 78% N2 and 22% O2) and compare this value to the value of He you calculated in Exercise 6.7A.

Exercise 6.7B

The density of diethyl ether vapor is 3.30 g/L. Calculate the molar mass of diethyl ether.

The density of an unknown gas at STP is 2.30 g/L. Calculate its molar mass.

Exercise 6.8A

EXAMPLE 6.8 Molar Mass from Gas Densities

SolutionThis time we are given the density in g/L. We want molar mass, which has units of g/mol. Once again we have the factor 22.4 L = 1 mol diethyl ether. Clearly, we must cancel the unit of liters and obtain the unit of moles in the denominator.

x = 73.9 g/mol3.30 g

1 L

22.4 L

1 mol

An unknown gaseous compound contains only hydrogen and carbon and its density is 1.34 g/mol at STP. What is the formula for the compound?

Exercise 6.8B

Nitrogen monoxide (nitric oxide), an air pollutant discharged by internal combustion engines, combines with oxygen to form nitrogen dioxide, a yellowish-brown gas that irritates the respiratory system and eyes. The equation for this reaction is

2 NO + O2 2 NO2

State the molecular, molar, and mass relationships indicated by the equation.

Exercise 6.9Hydrogen sulfide, a gas that smells like rotten eggs, burns in air to produce sulfur dioxide and water according to the equation

2 H2S + 3 O2 2 SO2 + H2OState the molecular, molar, and mass relationships indicated by this equation.

SolutionThe molecular and molar relationships can be obtained directly from the equation; no calculation is necessary. The mass relationship requires a little calculation: Molecular: Two molecules of NO react with one molecule of O2 to form two molecules of NO2.Molar: 2 mol of NO reacts with 1 mol of O2 to form 2 mol of NO2.Mass: 60.0 g of NO (2 mol NO x 30.0 g/mol) reacts with 32.0 g (1 mol O2 x 32.0 g/mol) of O2 to form 92.0 g (2 mol NO2 x 46.0 g/mol) of NO2.

CONCEPTUAL EXAMPLE 6.9 Molecular, Molar, and Mass Relationships

When 0.105 mol of propane is burned in a plentiful supply of oxygen, how many moles of oxygen is consumed?

C3H8 + 5 O2 3 CO2 + 4 H2O

EXAMPLE 6.10 Molar Relationships

For the combustion of propane in Example 6.10, (a) How many moles of carbon dioxide is formed when 0.529 mol of C3H8 is burned? (b) How many moles of water is produced when 76.2 mol of C3H8 is burned? (c) How many moles of carbon dioxide is produced when 1.010 mol of O2 is consumed?

Exercise 6.10

SolutionThe equation tells us that 5 mol O2 is required to burn 1 mol C3H8. We can write

1 mol C3H8 5 mol O2

where we use the symbol to mean “is stoichiometrically equivalent to.” From this relationship we can construct conversion factors to relate moles of oxygen to moles of propane. The possible conversion factors are

Which one do we use? Only if we multiply the given quantity (0.105 mol C3H8) by the factor on the right do we get an answer with the asked-for units (moles of oxygen).

? mol O2 = 0.105 mol C3H8 x = 0.525 mol O2

1 mol C3H8

5 mol O2

5 mol O2

1 mol C3H8

and

5 mol O2

1 mol C3H8

Calculate the mass of oxygen needed to react with 10.0 g of carbon in the reaction that forms carbon dioxide.

EXAMPLE 6.11 Mass Relationships

SolutionStep 1 The balanced equation is

C + O2 CO2

Step 2 The molar masses are 2 x 16.0 = 32.0 g/mol for O2 and 12.0 g/mol for C.

Step 3 We convert the mass of the given substance, carbon, to an amount in moles.

? mol C = 10.0 g C x = 0.833 mol C

Step 4 We use coefficients from the balanced equation equation to establish the stoichiometric factor (red) that relates the amount of oxygen to that of carbon.

0.833 mol C x = 0.833 mol O2

Step 5 We convert from moles of oxygen to grams of oxygen.

0.833 mol O2 x = 26.7 g O2

1 mol O2

12.0 g C

1 mol O2

1 mol C

32.0 g O2

1 mol O2

Calculate the mass of oxygen (O2) needed to react with 0.334 g of nitrogen (N2) in the reaction that forms nitrogen dioxide.

Exercise 6.11A

EXAMPLE 6.11 Mass Relationships continued

Calculate the mass of carbon dioxide formed by burning 775 g of each of (a) methane (CH4) and (b) butane (C4H10).

Exercise 6.11B

We can also combine the five steps into a single setup. Note that the units in the denominators of the conversion factors are chosen so that each cancels the unit in the numerator of the preceding term.

(The slightly different answers are due to rounding in the intermediate steps.)

1 mol C

12.0 g C10.0 g C x x x = 26.6 g O2

1 mol O2

1 mol C

32.0 g O2

1 mol O2

We start here This converts g C to mol C

This releasemol C to mol O2

This convertsmol O2 to g O2

The answer: the number and the unit

The decomposition of sodium azide (NaN3) produces sodium metal and nitrogen gas. The gas is used to inflate automobile airbags. What mass of nitrogen, in grams, can be made from 60.0 g of sodium azide?

EXAMPLE 6.12 Mass Relationships

Solution1. We start by writing and balancing the chemical equation, which shows that 2 mol NaN3 produces 2 mol

Na and 3 mol N2.

2 NaN3 2 Na + 3 N2

2. The molar mass of NaN3 is 23.0 g/mol + (3 x 14.0) g/mol = 65.0 g/mol and the molar mass of N2 is (2 x 14.0) g/mol = 28.0 g/mol.

3. We convert the mass of the given substance, sodium azide, to an amount in moles.

60.0 g NaN3 x = 0.923 mol NaN3

4. We use coefficients from the balanced equation to establish the stoichiometric factor that relates the amount of nitrogen gas to that of sodium azide.

0.923 mol NaN3 x = 1.38 mol N2

1 mol NaN3

65.0 g NaN3

3 mol N2

2 mol NaN3

Ammonia reacts with phosphoric acid (H3PO4) to form ammonium phosphate [(NH4)3PO4]. What mass in grams of ammonia is needed to react completely with 74.8 g of phosphoric acid?

Exercise 6.12A

EXAMPLE 6.12 Mass Relationships continued

a. The decomposition of potassium chlorate (KCIO3) produces potassium chloride (KCI) and O2 gas. What mass in grams of oxygen can be made from 2.47 g of potassium chlorate?

b. Phosphorus reacts with oxygen to form tetraphosphorus decoxide. The equation is

P4 + O2 P4O10 (not balanced)

What mass in grams of tetraphosphorus decoxide can be made from 3.50 g of phosphorus?

Exercise 6.12B

5. We convert from moles of nitrogen gas to grams of nitrogen gas.

1.38 mol N2 x = 38.6 g N2

As is usually the case, all the steps just outlined can be combined into a single setup.

60.0 g NaN3 x x x = 38.8 g NaN3

Notice that the units of the numerator in one stoichiometric factor are the units in the denominator of the next stoichiometric factor. In this way, the correct cancellation of units occurs and the units of the final numerator are the units of your answer.

28.0 g N2

1 mol N2

1 mol NaN3

65.0 g NaN3

3 mol NaN3

2 mol NaN3

28.0 g N2

1 mol N2

A gas is enclosed in a cylinder fitted with a piston. The volume of the gas is 2.00 L at 0.524 atm. The piston is moved to increase the gas pressure to 5.15 atm. Which of the following is a reasonable value for the volume of the gas at the greater pressure?

0.20 L 0.40 L 1.00 L 16.0 L

SolutionThe pressure increase from 0.524 atm to 5.15 atm is almost tenfold. The volume should drop to about one-tenth of the initial value. We estimate a volume of 0.20 L. (The calculated value is 0.203 L.)

A gas is enclosed in a 10.2-L tank at 1208 mmHg. (The mmHg is a pressure unit; 760 mmHg = 1 atm.) Which of the following is a reasonable value for the pressure when the gas is transferred to a 30.0-L tank?

300 mmHg 400 mmHg 3,600 mmHg 12,000 mmHg

Exercise 6.13

EXAMPLE 6.13 Boyle's Law: Pressure-Volume Relationships

A cylinder of oxygen has a volume of 2.25 L. The pressure of the gas is 1470 pounds per square inch (psi) at 20 °C. What volume will the oxygen occupy at standard atmospheric pressure (14.7 psi) assuming no temperature change?

EXAMPLE 6.14 Boyle’s Law: Pressure-Volume Relationships

SolutionWe find it helpful to first separate the initial from the final condition.

Initial Final Change

P1 = 1470 psi P2 = 14.7 psi

V1 = 2.25 L V2 = ?

Then use the equation V1P1 = V2P2 and solve for the desired volume or pressure. In this case, we solve for V2.

V2 =

V2 = = 225 L

Because the final pressure (14.7 psi) is less than the initial pressure (1470 psi), we expect the final volume (225 L) to be larger than the original volume (2.25 L), and we see that it is.

The pressure goes down, therefore, the volume goes up.

2.25 L x 1470 psi14.7 psi

V1P1

P2

EXAMPLE 6.14 Boyle’s Law: Pressure-Volume Relationshipscontinued

A sample of air occupies 73.3 mL at 98.7 atm and 0 ºC. What volume will the air occupy at 4.02 atm and 0 ºC?

Exercise 6.14A

A sample of helium occupies 535 mL at 988 mmHg and 25 °C. If the sample is transferred to a 1.05-L flask at 25 °C, what will be the gas pressure in the flask?

Exercise 6.14B

A balloon indoors, where the temperature is 27 °C, has a volume of 2.00 L. What would its volume be (a) in a hot room where the temperature is 47 °C, and (b) outdoors, where the temperature is –23 ºC? (Assume no change in pressure in either case.)

EXAMPLE 6.15 Charles’s Law: Temperature-Volume Relationships

SolutionFirst, and most important, convert all temperatures to the Kelvin scale:

T(K) = t(°C) + 273

The initial temperature (T1) in each case is (27 + 273) = 300 K and the final temperatures are (a) (47 + 273) = 320 K and (b) (–23 + 273) = 250 K.a. We start by separating the initial from the final condition.

Initial Final Changet1 = 27 °C t2 = 47 °C

T1 = 300 K T2 = 320 K

V1 = 2.00 L V2 = ?

Solving the equation

=

V1

T1

V2

T2

EXAMPLE 6.15 Charles’s Law: Temperature-Volume Relationships continued

for V2, we have

V2 =

V2 = = 2.13 L

As expected, because the temperature increases, the volume must also increase.

b. We have the same initial conditions as in (a), but different final conditions.

Initial Final Change

t1 = 27 °C t2 = -23 °C

T1 = 300 K T2 = 250 K

V1 = 2.00 L V2 = ?

Again using Charles’s law, we solve the equation for V2:

V2 =

V2 = = 1.67 L

As we expected, the volume decreased because the temperature decreased.

V1T2

T1

V1T2

T1

2.00 L x 250 K

300 K

2.00 L x 320 K

300 K

a. A sample of oxygen gas occupies a volume of 2.10 L at 25 °C. What volume will this sample occupy at 150 °C? (Assume no change in pressure.)

b. A sample of hydrogen occupies 692 L at 602 °C. If the pressure is held constant, what volume will the gas occupy after being cooled to 23 °C?

Exercise 6.15A

At what Celsius temperature will the initial volume of oxygen in Exercise 6.15A occupy 0.750 L? (Assume no change in pressure.)

Exercise 6.15B

EXAMPLE 6.15 Charles’s Law: Temperature-Volume Relationships continued

EXAMPLE 6.16 Ideal Gas Law

Use the ideal gas law to calculate (a) the volume occupied by 1.00 mol of nitrogen gas at 244 K and 1.00 atm pressure, and (b) the pressure exerted by 0.500 mol of oxygen in a 15.0-L container at 303 K.

Determine (a) the pressure exerted by 0.0330 mol of oxygen in an 18.0-L container at 313 K, and (b) the volume occupied by 0.200 mol of nitrogen gas at 298 K and 0.980 atm.

Exercise 6.16A

a. We start by solving the ideal gas equation for V.

V =

V = x x 244 K = 20.0 L

b. Here we solve the ideal gas equation for P.

P =

P = x x 303 K = 0.829 atm

nRT

P

nRT

V

1.00 mol

1.00 atm

0.0821 L • atm

mol • K

Solution

0.500 mol

15.0 L

0.0821 L • atm

mol • K

Determine the volume of nitrogen gas produced from the decomposition of 130 g sodium azide (about the amount in a typical automobile airbag) at 25 °C and 1 atm.

Exercise 6.16B

EXAMPLE 6.16 Ideal Gas Law continued

Calculate the molarity of a solution made by dissolving 3.50 mol of NaCl in enough water to produce 2.00 L of solution.

Calculate the molarity of a solution that has 0.0500 mol of NH3 in 5.7 L of solution.

Exercise 6.17A

Molarity (M) = = = 1.75 M NaCl

We read 1.75 M NaCl as “1.75 molar NaCl.”

Solutionmoles of solute

liters of solution

3.50 mol NaCl

2.00 L solution

Calculate the molarity of a solution made by dissolving 0.750 mol of H3PO4 in enough water to produce 775 mL of solution.

Exercise 6.17B

EXAMPLE 6.17 Solution Concentration: Molarity

What is the molarity of a solution in which 333 g of potassium hydrogen carbonate is dissolved in enough water to make 10.0 L of solution?

Calculate the molarity of each of the following solutions.a. 18.0 mol of H2SO4 in 2.00 L of solutionb. 3.00 mol of KI in 2.39 L of solutionc. 0.206 mol of HF in 752 mL of solution (HF is used for etching glass.)

Exercise 6.18

EXAMPLE 6.18 Solution Concentration: Molarity

First, we must convert grams of KHCO3 to moles of KHCO3

333 g KHCO3 x = 3.33 mol KHCO3

Now use this value as the numerator in the defining equation for molarity. The solution volume, 10.0 L, is the denominator.

Solution

1 mol KHCO3

100.1 g KHCO3

3.33 mol KHCO3

10.0 L solution Molarity = = 0.333 M KHCO3

How many grams of NaCl is required to prepare 0.500 L of typical over-the-counter saline solution(about 0.15 M NaCl)?

What mass in grams of potassium hydroxide is required to prepare 2.00 L of 6.00 M KOH?

Exercise 6.19A

EXAMPLE 6.19 Solution Preparation: Molarity

What mass in grams of potassium hydroxide is required to prepare 100.0 mL of 1.00 M KOH?

Exercise 6.19B

SolutionFirst we use the molarity as a conversion factor to calculate moles of NaCl.

Then we use the molar mass to calculate the grams of NaCl.

0.500 L solution x = 0.075 mol NaCl0.15 mol NaCl

1 L solution

0.75 mol NaCl x = 44 g NaCl58.4 g NaCl

1 mol NaCl

Concentrated hydrochloric acid has a concentration of 12.0 M HCl. How many milliliters of this solution would one need to get 0.425 mol of HCl?

What volume in milliliters of 15.0 M aqueous ammonia (NH3) solution do you need to get 0.445 mol of NH3?

Exercise 6.20A

EXAMPLE 6.20 Moles from Molarity and Volume

What mass in grams of HNO3 is in 500 mL of rain that has a concentration of 2.0 x 10–5 M HNO3?

Exercise 6.20B

Solution

We would need 0.0354 L (35.4 mL) of the solution to have 0.425 mol. Remember that molarity is moles per liter of solution, not per liter of solvent.

Liters of HCl solution = =

= = 0.0354 L

moles of solute

molarity

0.425 mol HCl

12.0 M HCl

0.425 mol HCl

12.0 mol HCl/L

Two-stroke engines use a mixture of 120 mL of oil dissolved in enough gasoline to make of 4.0 liters of fuel. What is the percent by volume of oil in this mixture?

What is the volume percent of ethanol in a solution that has 58.0 mL water in 625 mL of an ethanol–water solution?

Exercise 6.21A

EXAMPLE 6.21 Percent by Volume

Assume that the volumes are additive, and determine the volume percent toluene (C6H5CH3) in a solution made by mixing 40.0 mL of toluene with 75.0 mL of benzene (C6H6).

Exercise 6.21B

Solution

Percent by volumes = x 100% = 3.0%120 mL oil

4000 mL solution

Describe how to make 775 mL of vinegar (about a 5.0% by volume solution of acetic acid in water).

Describe how to prepare 450 mL of an aqueous solution that is 70.0% isopropyl alcohol by volume.

Exercise 6.22A

EXAMPLE 6.22 Solution Preparation: Percent by Volume

Describe how you would prepare exactly 2.00 L of an aqueous solution that is 9.77% acetic acid by volume.

Exercise 6.22B

Solution

Volume of solute =Percent by volume x volume of solution

100%

We begin by rearranging the equation for percent by volume to solve for volume of solute.

Substituting, we have

Take 39 mL of acetic acid and add enough water to make 775 mL of solution. Notice that we don’t simply add 775 mL of water, because the final volume of solution must be 775 mL.

= = 39 mL5.0% x 775 mL

100%

What is the percent by mass of a solution of 25.5 g of NaCl dissolved in 425 g (425 mL) of water?

Hydrogen peroxide solutions for home use are 3.0% by mass solutions of H2O2 in water. What is the percent by mass of a solution of 9.40 g of H2O2 dissolved in 335 g (335 mL) of water?

Exercise 6.23A

EXAMPLE 6.23 Percent by Mass

Sodium hydroxide (NaOH, lye) is used to make soap and is very soluble in water. What is the percent by mass of a solution that contains 1.00 kg of NaOH dissolved in 950 mL of water?

Exercise 6.23B

Use these values in the above percent-by-mass equation:Solution

Percent by mass = x 100% = 5.66% NaCl25.5 g NaCl

(25.5 + 425) g solution

Describe how to make 430 g of an aqueous solution that is 4.85% by mass NaNO3.

Describe how you would prepare 125 g of an aqueous solution that is 4.50% glucose by mass.

Exercise 6.24A

EXAMPLE 6.24 Solution Preparation: Percent by Mass

Describe how you would prepare 1750 g of isotonic saline, a commonly used intravenous (IV) solution that is 0.89% sodium chloride by mass.

Exercise 6.24B

We begin by rearranging the equation for percent by mass to solve for mass of solute.

Substituting, we have

Take 20.9 g of NaNO3 and add enough water to make 430 g of solution.

Solution

Mass of solute =percent by mass x mass of solution

100 %

= = 20.9 g4.85% x 430 g

100 %