Chapter 4. Chemical equation and reaction stoichiometry

Objectives: Write balanced chemical equation to

describe chemical reactions From balanced equations, calculate

the moles of reactants and products involved in each of the reactions

Determine which reactant is the limiting reactant in reactions

Objectives

Compare the amount of substance actually formed in a reaction (actual yield) with the predicted amount (theoretical yield) and determine the percent yield

Work with sequential reactions Use the terminology of solutions –

solute, solvent, concentration Calculate concentrations of solutions

when they are diluted

Chemical equations Chemical reactions always involve

changing one or more substances into one or more different substances

In other words, chemical reactions rearrange atoms or ions to form other substances

Chemical equations are used to describe chemical reactions

Chemical equations show:

(1) the substances that react, reactants

(2) the substances formed, products

(3) the relative amounts of the substances involved.

CH4 + 2O2 CO2 + 2H2O

heat

Special conditions required for some reactions are indicated by notation over the arrow

Balance chemical equations

A balanced chemical equation must always include the same number of each kind of atom on both sides of the equation.

Write equations with smallest possible whole-number coefficients

In reactants and In products (Ex.)

Examples

C2H6O + O2 CO2 + H2O

Carbon appear in only one compound on each side, the same is true for hydrogen

Al + HCl AlCl3 + H2

Exercice

Balance the following chemical reactions:

(a) P4 + Cl2 PCl5(b) RbOH + SO2 Rb2SO3 + H2O

(c) P4O10 + Ca(OH)2 Ca3(PO4)2 + H2O

Calculations based on chemical equations

Number of molecules (a balanced chemical equation may be interpreted on a molecular basis)

Number of moles formed (Avogadro’s number)

CH4 + 2O2 CO2 + 2 H2O

1 mol 2 mol1 mol2 mol

6.023 x 1023

Examples

How many O2 molecules react with 30 CH4 molecules according to the preceding equation?

How many moles of water could be produced by the reaction of 5 mol of methane with excess oxygen (more than sufficient)

Mass of a reactant requiredExample

What mass of oxygen (in g) is required to react completely with 24.0 g of CH4

What mass of CH4, in grams, is required to react with 96.0 g of O2

Mass of a product formedExample

Calculate the mass of CO2, in grams, that can be produced by burning 6.00 mol of CH4 in excess O2

The limiting reactant concept

The calculations were based on the reactant that was used up first, called the limiting reactant

Example: what mass of CO2 could be formed by the reaction of 16.0 g of CH4 with 48.0 g of O2

Example

A fuel mixture used in the early days of rocketry is composed of two liquids, hydrazine (N2H4) and N2O4, which ignite on contact to form nitrogen gas and water vapor. How many grams of nitrogen gas form when 1.00X 102 g of N2H4 and 2.00X 102 g of N2O4 are mixed?

Percent yields from chemical reactions Theoretical yield (from a chemical

reaction) is the yield calculated by assuming that the reaction goes to completion

In practice we often do not obtain as much product from a reaction as is theoritically possible Many reactions do not go to completion In some cases, reactants form undesired

products (by-products) In some cases, separation of the desired

products is so difficult that not all of the product formed is successfully isolated

2,4 D (2,4-Dichlorophenoxyacetic acid)

2,4,5 T

Actual yield: is the amount of a specified pure product actually obtained from a given raction

Percent yield = actual yield of product

theoritical yield of productX 100 %

Exercise

A 15.6-g sample of C6H6 is mixed with excess HNO3. We can isolate 18.0 g of C6H5NO2. What is the percent yield of C6H5NO2 in this reaction

Sequential reactions Often more than one reaction is

required to change starting materials into the desired product. These are called sequential reactions.

H3PO4 can be prepared in two-step process P4 + 5O2 P4O10

P4O10 + 6H2O H3PO4

Example

In the above reaction, we allow 272 g of phosphorous to react with excess oxygen, which forms tetraphosphorous decoxide, P4O10, in 89.5% yield. In the second step reaction, a 98.6% yield of H3PO4 is obtained. What mass of H3PO4 is obtained?

Exercise

SiC is an important ceramic material that is made by allowing sand SiO2 to react with powdered carbon at high temperature. Carbon monoxide is also formed. When 100.0 kg of sand is processed, 51.4 kg of SiC is recovered. What is the percent yield of SiC from this process?

Concentrations of solutions A solution is a homogeneous mixture, at

the molecular level, of two or more substances.

Solute: The dispersed (dissolved) phase Solvent: The dispersing medium The solutions used in the laboratory are

usually liquids, and the solvent is often water. These are called aqueous solutions

Percent by mass

mass of solute

mass of solution

x 100 %percent solute =

or

mass of solutex 100 %percent =

mass of solventmass of solute +

Molarity (M)

The number of moles of solute per liter of solution

number of moles of solute

number of liters of solution

molarity =

Dilution of solutions

Recall the definition of molarity, we have

volume (in L) x molarity = number of moles of solute

V1M1 = V2M2

Example

Isotonic saline is a 0.15 M aqueous solution of NaCI that simulates the total concentration of ions found in many cellular fluids. Its uses range from a cleansing rinse for contact lenses to a washing medium for red blood cells. How would you prepare 0.80 L of isotonic saline from a 6.0 M stock solution?

Using solutions in chemical reactions Usually we carry out a reaction in a

solution, therefore we must calculate the amounts of solutions that we need.

Example (amount of solute)Calculate (a) the number of moles of

H2SO4 and (b) the number of grams of H2SO4 in 500 mL of 0.324 M H2SO4 solution

Exercise (Solution stoichiometry) and (Volume of solution required)

Calculate the volume in liters and in milliliters of a 0.324 M solution of H2SO4 required to react completely with 2.792 g of Na2CO3.

Find the volume in liters and in milliliters of a 0.505 M NaOH solution required to react with 40.0 mL of 0.505 M H2SO4 solution

Exercise

What is the maximum mass of Ni(OH)2 that could be prepared by mixing two solutions that contain 25.9 g of NiCl2 and 10.0 g of NaOH, respectively?

Chapter 5. Some types of chemical reactions Objectives:

Describe the periodic table and some of the relationships that it summarizes

Recognize and descibe nonelectrolytes, strong and weak electrolytes.

Recognize and classify acids, bases, and salts

Assign oxidation number to elements, when they are free, in compounds, or in ions

Objectves (continue) Name and write formulas for

common binary and ternary inorganic compounds

Recognize oxidation-reduction reactions and identify which species are oxidized, reduced, oxidizing agents, and reducing agents

Recognize and describe classes of reactions

The periodic table: metals, nonmetals, and metalloids Atomic weight Atomic number of an element is

the number of protons in the nucleus of its atoms

Elements are arranged in the periodic table in order of increasing atomic number.

The properties of the elements are periodic functions of their atomic number

Metallic character decrease

Increase

Transition metals

Noble gases

The vertical colums are referred to as groups or families

The horizontal rows are called periods

Elements in a group have similar chemical and physical properties, and those within a period have properties that change progressively across the table.

Names of some common groups The Group IA elements (except H)

are referred to as alkaline metals The Group IIA elements are called

alkaline earth metals The Group VIIA elements are called

halogens (“salt formers”) The Group VIIIA elements are

called noble (rare) gases

Some physical properties of metals and nonmetals

Metals High EC that

decreases with increasing temperature

High thermal conductivity

Metallic gray or silver luster

Almost all are solids Malleable Ductile

Nonmetals Poor electrical

conductivity (except C in graphite)

Good heat insulator No metallic luster Solids, liquids, or

gases Brittle in solid state Nonductile

Some chemical properties of metals and nonmetals

Metals Outer shells contain

few electrons (usually 3 or fewer)

Form cations by losing electrons

Form ionic compounds with nonmetals

Solid state characterized by metallic bonding

Nonmetals Outer shells contain

4 or more electrons Form anions by

gaining elecrtons Form ionic

compounds with metals, and molecular (covalent) other compounds with nonmetals

Covalently bonded molecules

Metalloids (semi-metals) show some properties that are characteristic of both metals and nonmetals

Example: B, Si, Ge, As, Te Many of the metalliods, such as Si, Ge, and

Sb act as semiconductors (for electronic curcuits). Semiconductor are insulators at lower temperatures but become conductors at higher temperatures.