Chemical Reaction and Equations - Davis School District · Chemical Reaction and Equations ......

Chemical Reaction

and Equations

Evidence of a Reaction

Chemical Equations

Balancing Chemical

Equations

Writing and Balancing Chemical

Reactions and Equations

• 1. List four observations that suggest that a

chemical reaction has taken place.

• 2. List three requirements for a correctly

written chemical equation.

• 3. Write a word equation and a formula

equation for a given chemical reaction.

• 4. Write a chemical equation by balancing a

formula equation by inspection.

Evidence of a Chemical Reaction

1. The release of

light

or

heat .

More Light and Heat Evidence

Evidence of a Chemical Reaction

2. The production of

a gas

(bubbles) .

More Gas Production Evidence

O2

2H2 O2 O2 2

H2

H2O

Zn

H2 HCl

Evidence of a Chemical Reaction

3. The formation of

a precipitate

(solid)

when two aqueous

solutions are

mixed together.

Pb(NO3)2

KI

PbI2

When gases form a Precipitate

NH4Cl(s)

Evidence of a Chemical Reaction

4. A color change is often evidence of a

chemical reaction.

PbI2

Characteristics of Chemical

Equations

1. The equation represents known facts.

2. The equation must contain correct formulas!

Reactant + Reactant Product + Product

3. The equation must obey The Law of Conservation of Mass.

Mass of Reactants = Mass of Products

Word Equations

A chemical equation in which the reactants and products are represented by words (names).

Methane + Oxygen Carbon dioxide + Water

The means yields or produces

The + means “and”

Formula Equation

• Reactants and products are described by using their symbols and/or formulas.

CH4 + O2 CO2 + H2O

• A formula equation does not obey “The Law of Conservation of Mass.”

Chemical Equations

1. Use symbols and formulas to represent the reactants and products.

2. Adjust the coefficients in front of each reactant or product to identify the number of each reactant or product.

CH4 + 2 O2 CO2 + 2 H2O

Ones are not written but they are still there!

Chemical Equations

3. Use special symbols in parenthesis to indicate

the state of the reactants or products.

(s) = solid

(l) = liquid

(g) = gas

(aq) = aqueous (dissolved in water)

Chemical Equations

4. Use symbols above the yield arrow to

represent catalysts or special conditions.

heat Δ atm pressure MnO2

Reversible Reactions and Equations

Some chemical reactions occur in both directions.

N2 (g) + 3 H2 (g) 2 NH3 (g)

Balanced Chemical Equations

The coefficients represent the relative

amounts of reactants and products.

1 CH4 + 2 O2 1 CO2 + 2 H2O

1 molecule 2 molecules 1 molecule 2 molecules

1 mole 2 moles 1 mole 2 moles

The Importance of a Balanced

Chemical Equation cont. Coefficients can be used to determine the relative

masses of reactants and products

1 CH4 + 2 O2 1 CO2 + 2 H2O

1 mole 2 moles 1 mole 2 moles

1(16.05g) + 2(32.00g) = 1(44.01g) + 2(18.02g)

80.05 g = 80.05 g

Law of Conservation of Mass

A Chemical Reaction

Evidence = ?

CH4 2 O2

Reactants

Light from a Big Bang!

CH4 2 O2 CO2 2 H2O

ENERGY

Reactants Products

Balancing Chemical Equations

1.Write the Word (Name) Equation.

Methane + Oxygen Carbon dioxide + Water

2. Write the Formula Equation:

CH4 + O2 CO2 + H2O

Balancing Chemical Equations

3. Change the Coefficients as needed.

4. NEVER change a subscript !

5. Balance different types of atoms one

at a time.

9. Elements

Use the symbol from the periodic table for an

element.

Na Au S C Sn P

The following elements always appear as

diatomic molecules in chemical equations.

H2 N2 O2 F2 Cl2 Br2 I2

Balancing Chemical Equations

6. Balance elements that only appear once

on each side of the equation first.

7. Balance polyatomic ions that appear on

both sides of the equation as one unit.

8. Balance “H” and “O” last.

Checking Your Work

Count atoms to be sure that the equation

is balanced.

Counting = Coefficient x subscript

Example One

• Sodium metal (solid) combines with

chlorine gas to produce solid sodium

chloride.

Example One

• Sodium metal (solid) combines with chlorine gas

to produce solid sodium chloride.

Na (s) + Cl2 (g) NaCl (s)

• Remember that chlorine is diatomic!

Example One

• Sodium metal (Solid) combines with

chlorine gas to produce solid sodium

chloride.

Na (s) + Cl2 (g) NaCl (s)

2 Na (s) + Cl2 (g) 2 NaCl (s)

Example Two

• When copper metal (solid) reacts with

aqueous silver nitrate, the products formed

are aqueous copper (II) nitrate and silver

metal (solid).

Example Two

• When copper metal (solid) reacts with

aqueous silver nitrate, the products formed

are aqueous copper (II) nitrate and silver

metal (solid).

Cu (s) + AgNO3 (aq) Cu(NO3)2 (aq) + Ag(s)

Cu

Ag

N

O

1 1

1 1

1 2

2

2

6 6

Example Two

• When copper metal (solid) reacts with

aqueous silver nitrate, the products formed

are aqueous copper (II) nitrate and silver

metal (solid).

Cu (s) + AgNO3 (aq) Cu(NO3)2 (aq) + Ag(s)

Cu (s) + 2 AgNO3 (aq) Cu(NO3)2 (aq) + 2 Ag(s)

Example Three

• Solid Iron (III) oxide and carbon monoxide

gas react to form solid iron and carbon

dioxide gas.

Example Three

• Solid Iron (III) oxide and carbon monoxide

gas react to form solid iron and carbon

dioxide gas.

Fe2O3 (s) + CO(g) Fe(s) + CO2(g)

Fe

C

O

Example Three

• Solid Iron (III) oxide and carbon monoxide

gas react to form solid iron and carbon

dioxide gas.

Fe2O3 (s) + CO(g) Fe(s) + CO2(g)

Fe2O3 (s) + 3 CO(g) 2 Fe(s) + 3 CO2(g)

THE END