CH.3 - STOICHIOMETRY OF FORMULAS &...

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CHEMISTRY - SILBERBERG 8E

CH.3 - STOICHIOMETRY OF FORMULAS & EQUATIONS

CONCEPT: MASS CONVERSIONS

The _______________ is the chemical unit for the amount of a substance.

One mole (1 mol) contains 6.022 x 1023 entities, which is known as _________________________________________.

• Entities means ______________________ , ______________________ or ______________________.

• We use ______________________ when dealing with a single, individual element.

• We use _________________ or _________________when dealing with more than one element or a compound.

6.022 x 1023 atoms of Fe is equal to 1 mole of Fe and has a mass of 55.85 amu

Atoms Moles Grams

EXAMPLE: Determine the mass (in grams) found in 7.28 x 1028 nitrogen atoms.

6.022 x 1023 molecules of H2O is equal to 1 mole of H2O and has a mass of 18.016 amu

Molecules Moles Grams

EXAMPLE: Determine how many molecules of carbon dioxide, CO2, are found in 75.0 g CO2.



CONCEPT: MASS CONVERSIONS (PRACTICE)

PRACTICE 1: If the density of water is 1.00 g/mL at 25oC calculate the number of water molecules found in 1.50 x 103 µL

of water.

PRACTICE 2: Calculate the number of oxygen atoms found in 783.9 g CuSO4 · 5 H2O.

PRACTICE 3: The density of the sun is 1.41 g/cm3 and its volume is 1.41 x 1027 m3. How many hydrogen molecules are in

the sun if we assume all the mass is hydrogen gas?

PRACTICE 4 (CHALLENGE): A cylindrical copper wire is used for the fences of a house. The copper wire has a diameter

of 0.0750 in. How many copper atoms are found in 5.160 cm piece? The density of copper is 8.96 g/cm3. ( V = π · r2 · h ).



CONCEPT: EMPIRICAL FORMULA

The empirical formula is also known as the __________________________.

• It represents the _______________________________ ratio of moles

of each element in the compound.

The molecular formula is also known as the __________________________.

• It represents the _______________________________ ratio of moles of each element in the compound.

EXAMPLE: What is the empirical formula of dimethylhydrazine, C2H8N2, a colorless liquid used as a rocket fuel?

EXAMPLE: Elemental analysis of a sample of an ionic compound showed 2.82 g of Na, 4.35 g of Cl, and 7.83 g of O. What

is the empirical formula and name of the compound?

EXAMPLE: After a workout session, lactic acid (M = 90.08 g/mol) forms in muscle tissue and is responsible for muscle

soreness. Elemental analysis shows that this compound contains 40% C, 6.7% H and 53.3% O. Determine the molecular

formula.

C6H12O6 =



CONCEPT: COMBUSTION ANALYSIS

Under a combustion reaction a compound made of _______________ or _______________ reacts with ______________ .

• The products formed will be ________________ and ________________ .

EXAMPLE: A 0.2500 g sample contains carbon, hydrogen and oxygen and undergoes complete combustion to produce

0.3664 g of CO2 and 0.1500 g of H2O. What is the empirical formula of the compound?



CONCEPT: CHEMICAL COMPOSITION

The mass per mole of an element is called its _____________________________ (M).

The mass per mole of a compound is called its ____________________________ (M).

• They both have the units of _______________.

1. Elements. To find mass of an element just look up its atomic mass in the periodic table.

EXAMPLE: What is the total mass of each of the following elements?

a. Sodium b. Gold c. Mercury 2. Compounds. The mass of a compound is the sum of the individual masses of the elements in the chemical formula.

EXAMPLE: What is the total mass of each of the following compounds?

a. N2O5 b. C12H22O11 c. (NH4)3PO4



CONCEPT: MASS PERCENT

Mass or weight percent is the percentage of a given element in a compound.

EXAMPLE: What is the percentage of carbon in sodium hydrogen carbonate, NaHCO3?

EXAMPLE: A sample of toothpaste contains tin (II), SnF2. Analysis of a 5.25 g sample contains 8.77 x 10-3 g of F. What is

the percentage of tin (II) fluoride in the sample?

PRACTICE: Hemoglobin contains 0.33% iron and has a molecular weight of 68 kg. How many iron atoms are in each

molecule of hemoglobin?

Mass Percent (%) =

MassComponent)(TotalMass)(

•100



CONCEPT: BALANCING CHEMICAL REACTIONS

When balancing an equation always make sure the ________ and ________ of atoms on both sides of the arrow are equal.

EXAMPLE: Write balanced equations for each of the following by inserting the correct coefficients in the blanks:

a. ____ Al (s) + ____ Cl2 (g) ____ AlCl3 (s)

b. ____ Ba3(PO4)2 (s) + ____ KOH (aq) ____ K3PO4 (aq) + ____ Ba(OH)2 (aq)

c. ____ C4H10 (aq) + ____ O2 (g) ____ CO2 (g) + ____ H2O (l)



CONCEPT: GROUP REACTIVITY

The central principle of Organic Chemistry is based on the _________________________________. • The reactivity of an organic compound is all based on which type is present.

C

CC

C

CC

CO

C

C

CC

CC

CH3

H3C

H3C

OH

Tetrahydrocannabinol (THC)

CH2CH2CH2CH2CH3

Alkane

Alkene

Alkyne

Alcohol

Amine

Aldehyde

Ketone

Carboxylic Acid

Ester



CONCEPT: GROUP REACTIVITY (PRACTICE)

EXAMPLE: In each of the following molecules, identify the type(s) of functional groups present.

a. b.

O

OH

PRACTICE: In each of the following molecules, identify the type(s) of functional groups present.

a. b.

O

O



CONCEPT: STOICHIOMETRIC REACTIONS

2 H2 (g) + 1 O2 (g) 2 H2O (g)

In the above equation the numbers that are in bold are called _______________________.

• They tell us the number of ______________ of each compound that reacts.

• This numerical relationship between compounds in a balanced equation is called __________________________.

STOICHIOMETRIC CHART

Before we get into solving stoichiometric reactions lets work out a plan of attack.

Entities means ______________________ , ______________________ or ______________________.

Entities of Given

Grams of Given

Entities of Unknown

Grams of Unknown

Moles of Given Moles of Unknown

Use this chart when given a chemical equation with the ____________ quantity of a compound or element and asked to find

the ____________ quantity of another compound or element.

EXAMPLE: How many grams of H2O are produced when 12.3 g H2 reacts?

2 H2 (g) + 1 O2 (g) 2 H2O (g)



PRACTICE: STOICHIOMETRIC REACTIONS

EXAMPLE 1: The oxidation of chromium solid is represented by the following equation:

4 Cr (s) + 3 O2 (g) 2 Cr2O3 (s)

a. How many moles of chromium (III) oxide are produced when 34.69 g Cr reacts with excess oxygen gas?

b. How many grams of O2 were needed to produce 4.28 x 103 molecules Cr2O3?

EXAMPLE 2: If the density of ethanol, CH3CH2OH, is 0.789 g/mL, how many milliliters of ethanol are needed to produce 4.8

g of H2O in the following reaction?

CH3CH2OH (l) + 3 O2 2 CO2 (g) + 3 H2O (l)

PRACTICE: Dinitrogen monoxide gas decomposes to form nitrogen gas and oxygen gas. How many molecules of oxygen

are formed when 8.00 g of dinitrogen monoxide decomposes?



CONCEPT: LIMITING REAGENT

In a chemical reaction the reactant that is consumed when a reaction occurs and determines the maximum amount of

product formed is called the _________________________________.

• The amount of product it forms is called the _________________________ yield.

The reactant that remains after the completion of the chemical reaction is called the ____________________ reactant.

EXAMPLE: Chromium (III) oxide reacts with hydrogen sulfide (H2S) gas to form chromium (III) sulfide and water:

Cr2O3 (s) + 3 H2S (g) Cr2S3 (s) + 3 H2O (l) [Balanced]

a. What is the mass of chromium (III) sulfide formed when 14.20 g Cr2O3 reacts with 12.80 g H2S?

b. Identify the limiting reactant, excess reactant and theoretical yield.

c. What mass of excess reactant remains?



CONCEPT: PERCENT YIELD

The percent yield of a reaction is used to determine how effective the chemist was in creating their desired products.

• A high percent yield would signify that the reaction is _______________________________________.

EXAMPLE 1: A scientist performs an experiment in the laboratory and obtains 13.27 g Cr2S3. If his calculations on scratch

paper give him a theoretical yield of 18.23 g what is the percent yield?

PRACTICE: Consider the following balanced chemical reaction:

2 C6H6 (l) + 15 O2 (g) 12 CO2 (g) + 6 H2O (l)

a. If a 2.6 g sample of C6H6 reacted with excess O2 to produce 1.25 g of water, what is the percent yield of water?

b. If the above reaction only went to 75% completion, how many moles of CO2 would be produced if 1.57 x 10-5 molecules of C6H6 were reacted with excess oxygen?



CONCEPT: MOLARITY

Molarity (M) can serve as the connection between the interconversion of ____________ to ____________ and vice versa.

For example, a 5.8 M NaCl solution really means __________________________ per __________________________.

Molarity =MolesSolute)(LitersSolution)(

A typical mixture consists of a smaller amount of one substance, the ________________, dissolved in a larger amount of

another substance, the __________________. Together they form a __________________.



PRACTICE: MOLARITY

EXAMPLE 1: 2.64 grams of an unknown compound was dissolved in water to yield

150 mL of solution. The concentration of the solution was 0.075 M. What was the

molecular weight of the substance?

EXAMPLE 2: A solution is prepared by dissolving 0.1408 mol calcium nitrate, Ca(NO3)2, in enough water to make 100.0 mL

of stock solution. If 20.0 mL of this solution is then mix with an additional 90 mL of deionized water, calculate the

concentration of the calcium nitrate solution.

PRACTICE 1: What is the molarity of calcium ions of a 650 mL solution containing 42.7 g of calcium phosphate?

PRACTICE 2: A solution with a final volume of 750.0 mL was prepared by dissolving 30.00 mL of benzene (C6H6, density =

0.8787 gmL

) in dichloromethane. Calculate the molarity of benzene in the solution.

M =MolesSolute)(LitersSolution)(



CONCEPT: NORMALITY

Another measurement for concentration usually encountered is normality (N), which represents the number of equivalents per liter of solution.

N =equivalents of soluteLiters of solution

equivalent (eq) = n ×moles

• An equivalent is the mass of a compound that can either donate or accept an __________ or __________.

An equivalent (eq) and n are both determined by the compound being used (acid or base) and if the compound is undergoing a redox reaction.

Acids

For an acid the number for n is based on the number of _________________________ present.

EXAMPLE 1: Determine the number of equivalents for each of the acids given.

a) 2.5 moles CH3COOH b) 133.4 g H3PO4

Bases

For a base the number for n is based on the number of _________________________ present.

EXAMPLE 2: Determine the number of equivalents for the following base.

a) 50.0 mL of 0.165 M Ca(OH)2

Redox Reactions

For a redox reaction the number for n is based on the number of ________________________ transferred.

EXAMPLE 3: Based on the given redox reaction determine the value for n.

MnO4– (aq) + H+ (aq) Mn2+ (aq) + H2O (l)



PRACTICE: NORMALITY EXAMPLE 1: What is the normality of a solution made by dissolving 325.1 g HNO3 in enough water to create a 750.0 mL solution?

EXAMPLE 2: Determine the equivalent weight of the following compounds.

a) Al(OH)3 b) H2CO3

EXAMPLE 3: What volume, in mL, of 50.0 g H2SO4 is needed to create a 0.300 N H2SO4 solution?

EXAMPLE 4: If a concentrated 3.25 M H3PO4 solution possesses a density of 1.350 g/mL, what is its normality?



CONCEPT: MOLARITY & CHEMICAL REACTIONS

Whenever we are provided given information in a reaction we use ___________________ to find any unknown information.

• In aqueous reactions, this given information is typically in units of __________________ or __________________ .

Entities means ______________________ , ______________________ or ______________________.

Volume of Given Moles of Given Moles of Unknown Volume of Unknown

Entities of Unknown

Grams of Unknown

Use this chart when given a chemical equation with the known quantity in either ________ or ________ of a compound or

element and asked to find the unknown quantity of another compound or element.

EXAMPLE: How many grams of sodium metal are needed to react with 38.74 mL of 0.275 M NaOH?

2 Na (s) + 2 H2O (l) H2 (g) + 2 NaOH (aq)



PRACTICE: MOLARITY & CHEMICAL REACTIONS

PRACTICE 1: How many milliliters of 0.325 M HCl are needed to react with 16.2 g

of magnesium metal?

2 HCl (aq) + Mg (s) MgCl2 + H2 (g)

PRACTICE 2: What is the molarity of a hydrobromic acid solution if it takes 34.12 mL of HBr to completely neutralize 82.56

mL of 0.156 M Ca(OH)2?

2 HBr (aq) + Ca(OH)2 (aq) CaBr2 (aq) + 2 H2O (l)

PRACTICE 3 (CHALLENGE): Iron (III) can be oxidized by an acidic K2Cr2O7 solution according to the net ionic equation:

Cr2O72- + 6 Fe2+ + 14 H+ 2 Cr3+ + 6 Fe3+ + 7 H2O

If it takes 30.0 mL of 0.100 M K2Cr2O7 to titrate a 25 mL Fe2+ solution, what is the molar concentration of Fe2+?

M =MolesSolute)(LitersSolution)(



3. How many molecules of hexane are contained in 55.0 mL of hexane? The density of hexane is 0.6548 g/mL and the molar mass is 86.17 g/mol.



4. How many SO3 ions are contained in 120.0 mg of Na2SO3? The molar mass of Na2SO3 is 126.05 g/mol.



5. What mass of phosphorus pentafluoride, PF5, has the same number of fluorine atoms as 50.0 g of oxygen difluoride, OF2?



6. How many bromide ions are there in 4.50 moles of gallium bromide?



7. How many moles of oxygen atoms are required to combine with 3.05 moles of Pb to create lead (IV) phosphate?



8. How many cations are there in 100.0 g of lithium nitride?



10. Which of the following amounts would contain the least atoms? a) 10.0 g Sr

b) 10.0 g Br

c) 10.0 g Mg

d) 10.0 g Li



11. Which of the following amounts have the most molecules?

a) 15.0 g N2

b) 15.0 g Br2

c) 15.0 g O2

d) 15.0 g I2



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