Chapter 6

Chemical Reactions:An Introduction

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Chemical Reactions

• Reactions involve chemical changes in matter that result in new substances.

• Reactions involve rearrangement and exchange of atoms to produce new molecules.

• Reactants Products

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Evidence of Chemical Reactions

• A chemical change occurs when new substances are made.

• Visual clues (permanent):– Color change, precipitate formation, gas bubbles,

flames, heat release, cooling, light

• Other clues:– New odor, permanent new state

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Evidence of Chemical Reactions (cont.)

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Chemical Equations

• Shorthand way of describing a reaction

• Provides information about the reaction:– Formulas of reactants and products– States of reactants and products– Relative numbers of reactant and product

molecules that are required– Can be used to determine weights of reactants

used and of products that can be made

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Conservation of Mass

• Matter cannot be created or destroyed.

• In a chemical reaction, all the atoms present at the beginning are still present at the end.

• Therefore, the total mass cannot change.

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Combustion of Methane

• Methane gas burns to produce carbon dioxide gas and liquid water– Whenever something burns, it combines with O2(g).

CH4(g) + O2(g) CO2(g) + H2O(l)

H

HC

H

HOO+

O

O

C + OH H

1 C + 4 H + 2 O 1 C + 2 O + 2 H + O1 C + 2 H + 3 O

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Combustion of Methane Balanced

• To show a reaction obeys the Law of Conservation of Mass, it must be balanced.

CH4(g) + 2 O2(g) CO2(g) + 2 H2O(l)

H

HC

H

H

OO

+

O

O

C +

OH H

OO

+O

H H

+

1 C + 4 H + 4 O 1 C + 4 H + 4 O

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Writing Equations

• Use proper formulas for each reactant and product.• Proper equation should be balanced.

– Obey Law of Conservation of Mass.– All elements on reactants side also on product side.– Equal numbers of atoms of each element on reactant

side as on product side

• Balanced equations show the relationship between the relative numbers of molecules of reactants and products. – Can be used to determine mass relationships

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Symbols Used in Equations

• Symbols used after chemical formula to indicate state:– (g) = gas; (l) = liquid; (s) = solid– (aq) = aqueous, dissolved in water– e. g. NH3(aq) indicates ammonia dissolved in

water

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Sample – Recognizing Reactants and Products

• When magnesium metal burns in air it produces a white, powdery compound, magnesium oxide.

– Burning in air means reacting with O2

– Metals are solids, except for Hg, which is liquid.

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Recognizing Reactants and Products (cont.)

• Write the equation in words– Identify the state of each chemical

magnesium(s) + oxygen(g) magnesium oxide(s)

• Write the equation in formulas– Identify diatomic elements– Identify polyatomic ions– Determine formulas

Mg(s) + O2(g) MgO(s) (unbalanced)

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Balancing by Inspection

• Count atoms of each element– Polyatomic ions may be counted as one “element” if

they do not change in the reaction.

Al + FeSO4 Al2(SO4)3 + Fe

1 SO4 3

– If an element appears in more than one compound on the same side, count each element separately and add.

CO + O2 CO2

1 + 2 O 2

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Balancing by Inspection (cont.)

• Pick an element to balance.– Avoid elements from 1b

• Find least common multiple (LCM) and factors needed to make both sides equal.

• Use factors as coefficients in equation.– If already a coefficient, then multiply by new

factor• Recount and repeat until balanced.

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Example #1

• When magnesium metal burns in air it produces a white, powdery compound, magnesium oxide.

– Burning in air means reacting with O2

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Example #1 (cont.)

• Write the equation in words.– Identify the state of each chemical

magnesium(s) + oxygen(g) magnesium oxide(s)

• Write the equation in formulas.– Identify diatomic elements– Identify polyatomic ions– Determine formulas

Mg(s) + O2(g) MgO(s) (unbalanced)

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Example #1 (cont.)

• Count the number of atoms of on each side– Count polyatomic groups as one “element” if on

both sides– Split count of element if in more than one

compound on one side

Mg(s) + O2(g) MgO(s)

1 Mg 1

2 O 1

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Example #1 (cont.)

• Pick an element to balance– Avoid element in multiple compounds

• Find least common multiple of both sides & multiply each side by factor so it equals LCM

Mg(s) + O2(g) MgO(s)

1 Mg 1

1 x 2 O 1 x 2

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Example #1 (cont.)

• Use factors as coefficients in front of compound containing the element

– If coefficient is already there, multiply them together

Mg(s) + O2(g) 2 MgO(s)

1 Mg 1 x 2

1 x 2 O 1 x 2

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Example #1 (cont.)

• Recount

Mg(s) + O2(g) 2 MgO(s) 1 Mg 2

2 O 2• Repeat

2 Mg(s) + O2(g) 2 MgO(s) 2 x 1 Mg 2

2 O 2

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Example #2

• Under appropriate conditions, at 1000°C ammonia gas reacts with oxygen gas to produce gaseous nitrogen monoxide and gaseous water.

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Example #2 (cont.)

• Write the equation in words.– Identify the state of each chemical

ammonia(g) + oxygen(g) nitrogen monoxide(g) + water(g)

• Write the equation in formulas.– Identify diatomic elements– Identify polyatomic ions– Determine formulas

NH3(g) + O2(g) NO(g) + H2O(g)

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Example #2 (cont.)

• Count the number of atoms of on each side.

– Count polyatomic groups as one “element” if on both sides

– Split count of element if in more than one compound on one side

NH3(g) + O2(g) NO(g) + H2O(g)

1 N 13 H 2

2 O 1 + 1

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Example #2 (cont.)

• Pick an element to balance– Avoid elements in multiple compounds

• Find least common multiple of both sides & multiply each side by factor so it equals LCM

NH3(g) + O2(g) NO(g) + H2O(g)

1 N 12 x 3 H 2 x 3

2 O 1 + 1

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Example #2 (cont.)

• Use factors as coefficients in front of compound containing the element.

2 NH3(g) + O2(g) NO(g) + 3 H2O(g)

1 N 12 x 3 H 2 x 3

2 O 1 + 1

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Example #2 (cont.)

• Recount2 NH3(g) + O2(g) NO(g) + 3 H2O(g)

2 N 16 H 6

2 O 1 + 3

• Repeat2 NH3(g) + O2(g) 2 NO(g) + 3 H2O(g)

2 N 1 x 26 H 6

2 O 1 + 3

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Example #2 (cont.)

• Recount2 NH3(g) + O2(g) 2 NO(g) + 3 H2O(g)

2 N 26 H 6

2 O 2 + 3

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Example #2 (cont.)

• Repeat– When you are forced to attack an element that is in 3 or

more compounds, find where it is uncombined. You can find a factor to make it any amount you want, even if that factor is a fraction.

– We want to make the O on the left equal 5, therefore we will multiply it by 2.5

2 NH3(g) + 2.5 O2(g) 2 NO(g) + 3 H2O(g)

2 N 26 H 6

2.5 x 2 O 2 + 3

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Example #2 (cont.)

• Multiply all the coefficients by a number to eliminate fractions

– x.5 2, x.33 3, x.25 4, x.67 3

2 x [2 NH3(g) + 2.5 O2(g) 2 NO(g) + 3 H2O(g)]

4 NH3(g) + 5 O2(g) 4 NO(g) + 6 H2O(g)

4 N 412 H 12

10 O 10