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Chapter 8

“Chemical Reactions”

ChemistryJudson High School

Mr. Trotts

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Describing Chemical Reactions OBJECTIVES:

–Describe how to write a word equation.

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Describing Chemical Reactions OBJECTIVES:

–Describe how to write a skeleton equation.

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Describing Chemical Reactions OBJECTIVES:

–Describe the steps for writing a balanced chemical equation.

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All chemical reactions… have two parts:

–Reactants - the substances you start with

–Products- the substances you end up with

The reactants turn into the products.Reactants Products

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Reactants

Products

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CHEMICAL REACTION:a reaction in which one or

more substances are changed to new substances.

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REACTANTS:substances involved in a

chemical reaction.

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PRODUCTS:new substances produced in

a chemical reaction.

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The relationship between reactants and products can be

written as:

REACTANTS PRODUCTS

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In a chemical reaction Atoms aren’t created or destroyed. A reaction can be described several ways:

1. In a sentence (every item is a word) Copper reacts with chlorine to form copper (II)

chloride.

2. In a word equation (some symbols used)

Copper + chlorine copper (II) chloride

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Symbols in equations the arrow separates the reactants

from the products–Read as “reacts to form” or yields

The plus sign = “and” (s) after the formula = solid: AgCl(s)

(g) after the formula = gas: CO2(g)

(l) after the formula = liquid: H2O(l)

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Symbols used in equations(aq) after the formula = dissolved in

water, an aqueous solution: NaCl(aq) is a salt water solution

used after a product indicates a gas has been produced: H2↑

used after a product indicates a solid has been produced: PbI2↓

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Symbols used in equations■ indicates a reversible

reaction (more later)■ shows that

heat is supplied to the reaction■ is used to indicate a

catalyst is supplied, in this case, platinum.

heat ,

Pt

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What is a catalyst? A substance that speeds up a

reaction, without being changed or used up by the reaction.

Enzymes are biological or protein catalysts.

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3. The Skeleton EquationUses formulas and symbols to

describe a reaction

–but doesn’t indicate how many; this means they are NOT balanced

All chemical equations are a description that describe reactions.

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Write a skeleton equation for:1. Solid iron (III) sulfide reacts with

gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas.

2. Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water.

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Now, read these:

Fe(s) + O2(g) Fe2O3(s)

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

NO2 (g) N2(g) + O2(g) Pt

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4. Balanced Chemical EquationsAtoms can’t be created or destroyed

in an ordinary reaction:

–All the atoms we start with we must end up with

A balanced equation has the same number of each element on both sides of the equation.

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I. CONSERVATION OF MASS:

In a chemical reaction, matter is not created or

destroyed but is conserved.

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II. IN OTHER WORDS:The starting mass of the

reactants equals the final mass of the

products.

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Rules for balancing:1) Assemble the correct formulas for all the

reactants and products, use + and →2) Count the number of atoms of each type

appearing on both sides3) Balance the elements one at a time by

adding coefficients where needed (the numbers in front) - save balancing the H and O until LAST!

(I prefer to save O until the very last)4) Check to make sure it is balanced.

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Never change a subscript to balance an equation.

– If you change the formula you are describing a different reaction.

– H2O is a different compound than H2O2

Never put a coefficient in the middle of a formula

2NaCl is okay, but Na2Cl is not.

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WRITING EQUATIONS

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III. Chemical reactions can be described with words

such as:

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solid lead (II) nitrate, dissolved in water, plus solid potassium iodide, dissolved in water, produces solid lead

(II) iodide plus potassium nitrate, dissolved in water

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A. All of this information is important to letting a scientist know what the reactants and products

are as well as their physical states.

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B. A shorthand method has been developed to

describe chemical reactions.

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C. This method uses:1. Chemical formulas

(NaCl for sodium chloride)

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2. Coefficients (numbers to indicate how many molecules)

3. Symbols(for physical state,

catalysts, direction of reaction, etc.)

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The reaction described above would look like this:

Pb(NO3)2(aq) + 2KI(aq) PbI2(s) + 2KNO3(aq)

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IV. COEFFICIENTS: are the numbers placed to the left of the formulas

for the reactants and products.

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A. The coefficients represent the number of

units of each substance taking part in a reaction.

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B. In the reaction above, there are:

1. 2 units each of KNO3(aq) and KI(aq)

2. 1 unit each of Pb(NO3)2(aq) and PbI2(s)

Pb(NO3)2(aq) + 2KI(aq) PbI2(s) + 2KNO3(aq)

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V. Symbols are used to indicate what is

happening in the reaction.

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A. The physical state of the reactants

B. Things added to help the reaction take place

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C. An indicator to show which chemicals are the reactants and which are

the products

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D. Some commonly used symbols are in the following

table.

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SYMBOLS USED IN CHEMICAL EQUATIONS

SYMBOL MEANING

produces, yields or forms; placed between the reactants andthe products

+ plus; placed between individual reactants or products in theequation

(s) solid; placed after the formulas in parentheses on the sameline

(l) liquid(g) gas(aq) aqueous, a solid is dissolved in waterheat

the reactants are heated; things added to make the reactionhappen are written over the arrow.

light

the reactants are exposed to light

Zn

Zinc was added as a catalyst. A catalyst increases the rate ofa reaction, but is not consumed in the reaction.

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Reactants ProductsC3H8(g) + 5O2(g) 3CO2(g) + 4H2O(g)

1 unit propane gas;5 units oxygen gas3 units carbon dioxide gas; 4 units water vapor

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VII. REACTION ENERGY -- Chemical reactions either absorb heat

(endothermic) or release heat (exothermic).

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A. Endothermic reactions must have energy added

(usually thermal, sometimes electrical, or

light)for the reaction to take

place.

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1. This is due to more energy being required to break bonds than to make

new bonds.

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2. Some endothermic reactions cause the

reaction container to feel cool to the touch.

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b. Exothermic reactions have some form of energy released by the reaction (usually it is thermal or

light).

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1. This is due to less energy being required to break bonds than to form

new bonds.

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2. Exothermic reactions cause the reaction

container to feel warm or even hot to the touch.

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BALANCING EQUATIONS

THERE ARE 5 STEPS TO BALANCING EQUATIONS.

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Step 1: Remember these two rules:

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Each capital letter begins a new element

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Coefficients(numbers that multiply the

whole formula)can only be placed in front

of the chemical formula

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Step 2: Determine each type of element present

on both sides of the equation

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Write the equation.

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Make a table under the equation that lists each

type of element.

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List the elements in the center of the table

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Example:

MgO + Br2 MgBr2 + O2 Mg Br O

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Step 3: Determine the number of each type of

element on both sides of the equation.

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Example:

MgO + Br2 MgBr2 + O2

1 Mg 12 Br 21 O 2

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Step 4: Determine which elements are not balanced.

In the example, the unbalanced element is

oxygen.

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Step 5: to balance the number of atoms for the

element,Put a coefficient in front of the formula with the lower number of atoms for that

element to produce an equal number of atoms.

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Remember:The coefficient multiplies all of the elements in the chemical

formula.

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Start balancing by trying to make the numbers of elements in all the other formulas equal

with the most complicated chemical formula (the formula with the most elements and

subscripts).

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Always balance Hydrogen and Oxygen last.

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Water can always be added to the equation

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Example: Place a coefficient of 2 in front of

the MgO.Then multiply the number

of atoms in the table by

the coefficient.

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2 MgO + Br2 MgBr2 + O2

2 1 Mg 12 Br 2

2 1 O 2The coefficient changed the number ofMg also. Now the Mg has to be balanced.

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Go to the other side of the equation and place

a coefficient of 2 in front of the MgBr2.

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2 MgO + Br2 2 Mg Br2 + O2

2 1 Mg 1 22 Br 2 4

2 1 O 2

The coefficient changed the number of Bralso. Now the Br has to be balanced.

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Place a coefficient of 2 in front of the Br2.

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2 MgO + 2 Br2 2 Mg Br2 + O2

2 1 Mg 1 2 4 2 Br 2 4

2 1 O 2

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The equation is now balanced with 2-Mg, 4-

Br, and 2-O.

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OTHER HINTS

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The diatomic molecules(Br I N Cl H O F)

are always written as two-atom molecules

(Br2, I2, N2, Cl2, H2, O2, F2).

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A coefficient can be changed when needed to

get every element to balance.

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Many times there can be an even amount of atoms on one side and an odd amount of atoms on the

other side.

]Find a common multiple.

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Example:

N2 + H2 NH3

2 N 1 2 H 3There are 2 hydrogen on the left side.There are 3 hydrogen on the right side.

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1. Determine the common multiple between

2 and 3 (it is 6).

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2. Multiply the H2 by 3 and the NH3 by 2.

N2 + 3 H2 2 NH3

2 N 1 26 2 H 3 6

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The equation is now balanced with 2-N and

6-H.

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Sometimes hydrogen and oxygen can cause

problems. Two or more sources of hydrogen or

oxygen on one side of the equation can occur.

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Example:

ZnS + O2 ZnO + SO2

The product side has oxygen from 2 sources.

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1. Make the odd source of oxygen even by

multiplying ZnO by 2. ZnS + O2 2 ZnO + SO2

This 1 Zn 1 2unbalances 1 S 1the Zinc. 2 O 3 4 (2-O’s

from 2ZnO plus2-O’s from SO2)

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2. Place a 2 in front of the ZnS.

2 ZnS + O2 2 ZnO + SO2

2 1 Zn 1 2This 2 1 S 1unbalances 2 O 3 4the sulfur.

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3. Place a 2 in front of the SO2.

2 ZnS + O2 2 ZnO + 2 SO2

2 1 Zn 1 2 2 1 S 1 2This 2 O 3 4 6unbalances (2 O’s from 2ZnO plusthe oxygen more. 4 O’s from 2SO2)

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4. Multiply the O2 by 3 to balance the equation.2 ZnS + 3 O2 2 ZnO + 2 SO2

2 1 Zn 1 2 2 1 S 1 2 6 2 O 3 4 6

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The equation is now balanced with 2-Zn, 2-S,

and 6-O on each side.

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A polyatomic ion that is not broken into parts by the reaction (it stays the same on both sides) is

treated like an element.

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Example: The equation,

H3PO4 + Ca(OH)2 Ca3(PO4)2 +

H2O,

has two polyatomic ions, phosphate (PO4) and

hydroxide (OH).

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The phosphate has remained intact.

The hydroxide was broken up to form water.

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1. Make a table and list all of the elements.

H3PO4 + Ca(OH)2 Ca3(PO4)2 + H2OList the PO4 1 PO4 2as if it is an 1 Ca 3element. 5 H 2 List the H

2 O 1 and Oseparately and last.

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Note that the oxygen in the phosphate ion is

not included when the number of oxygen

atoms is listed.

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2. The most complicated formula is the Ca3(PO4)2. Make the other formulas balance with the

Ca3(PO4)2.

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Start by balancing the calcium, multiply

Ca(OH)2 by 3. H3PO4 + 3Ca(OH)2 Ca3(PO4)2 + H2OThe PO4 1 PO4 2is now 3 1 Ca 3unbalanced. 9 5 H 2

6 2 O 1Do not worry about H and O yet.

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3. Multiply the H3PO4 by 2. 2H3PO4 + 3Ca(OH)2 Ca3(PO4)2 + H2O

2 1 PO4 2 3 1 Ca 3

12 9 5 H 2 6 2 O 1This leaves 12 H’s and 6 O’s. The ratio of

H to O is 2:1. This means that the reaction made

6 molecules of water.

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4. Finish balancing the equation by multiplying

the water by 6.2H3PO4+3Ca(OH)2 Ca3(PO4)2 + 6 H2O

2 1 PO4 2 3 1 Ca 3

12 9 5 H 2 6 2 O 1 6

The equation is now balanced with 2-PO4, 3-Ca, 12-H and 6-O.

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Practice Balancing Examples

1. _AgNO3 + _Cu _Cu(NO3)2 + _Ag

2. _Mg + _N2 _Mg3N2

3. _P + _O2 _P4O10

4. _Na + _H2O _H2 + _NaOH

5. _CH4 + _O2 _CO2 + _H2O

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Types of Chemical Reactions OBJECTIVES:

–Describe the five general types of reactions.

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Types of Chemical Reactions

OBJECTIVES:

–Predict the products of the five general types of reactions.

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Types of Reactions There are millions of reactions. We can’t remember them all, but luckily

they will fall into several categories. We will learn 5 major types. Will be able to predict the products. For some, we will be able to predict

whether or not they will happen at all. We recognize them by their reactants

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#1 - Combination Reactions or Synthesis

Reactions Combine = put together 2 substances combine to make one

compound. Ca +O2 CaO SO3 + H2O H2SO4

We can predict the products if the reactants are two elements.

Mg + N2

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Complete and balance: Ca + Cl2 Fe + O2 (assume iron (II) oxide is product)

Al + O2 Remember that the first step is to write

the correct formulas – you can still change the subscripts at this point, but not later!

Then balance by using the coefficients only

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#1 - CombinationNote:

a) Some nonmetal oxides react with water to produce an acid:

SO2 + H2O H2SO3

b) Some metallic oxides react with water to produce a base:

CaO + H2O Ca(OH)2

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#2 - Decomposition Reactionsdecompose = fall apart

one reactant breaks apart into two or more elements or compounds.

NaCl Na + Cl2

CaCO3 CaO + CO2

Note that energy (heat, sunlight, electricity, etc.) is usually required

electricity

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#2 - Decomposition ReactionsCan predict the products if it is a

binary compound-Made up of only two elements

–breaks apart into its elements:H2OHgO

electricity

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#2 - Decomposition Reactions If the compound has more than

two elements you must be given one of the products

–The other product will be from the missing pieces

NiCO3 CO2 + ___

H2CO3(aq) CO2 + ___

heat

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#3 - Single ReplacementOne element replaces anotherReactants must be an element and

a compound.Products will be a different element

and a different compound.Na + KCl K + NaClF2 + LiCl LiF + Cl2

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#3 Single Replacement Metals replace other metals (and they

can also replace hydrogen) K + AlN Zn + HCl Think of water as: HOH

–Metals replace one of the H, and then combine with the hydroxide.

Na + HOH

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#3 Single Replacement We can even tell whether or not a single

replacement reaction will happen:–Some chemicals are more “active”

than others–More active replaces less active

There is a list on page 333 - called the Activity Series of Metals

Higher on the list replaces lower.

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The Activity Series of the Metals Lithium

Potassium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Nickel Lead HydrogenHydrogen Bismuth Copper Mercury Silver Platinum Gold

1) Metals can replace other metals provided that they are above the metal that they are trying to replace.

2) Metals above hydrogen can replace hydrogen in acids.

3) Metals from sodium upward can replace hydrogen in water.

Higher activity

Lower activity

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#3 Single Replacement Practice:

6. Fe + CuSO4

7. Pb + KCl

8. Al + HCl

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#4 - Double Replacement Two things replace each other.

–Reactants must be two ionic compounds.

–Usually in aqueous solution NaOH + FeCl3

–The positive ions change place. NaOH + FeCl3 Fe+3 OH- + Na+1 Cl-1

NaOH + FeCl3 Fe(OH)3 + NaCl

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The Activity Series of the Halogens

Fluorine Chlorine Bromine Iodine

Halogens can replace other halogens in compounds, provided that they are above the halogen that they are trying to replace.

2NaCl(s) + F2(g) 2NaF(s) + Cl2(g)

MgCl2(s) + Br2(g) ???No Reaction

???

Higher Activity

Lower Activity

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#4 - Double Replacement Has certain “driving forces”

–Will only happen if one of the products:

a) doesn’t dissolve in water and forms a solid (a “precipitate”), or

b) is a gas that bubbles out, or

c) is a molecular compound (usually water).

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Complete and balance:assume all of the following

reactions actually take place: 9. CaCl2 + NaOH

10. CuCl2 + K2S

11. KOH + Fe(NO3)3

12. (NH4)2SO4 + BaF2

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How to recognize which type Look at the reactants:

E + E = Combination (synthesis)

C = Decomposition

E + C = Single replacement

C + C = Double replacement

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Practice Examples: reaction type and product

13. H2 + O2

14. H2O 15. Zn + H2SO4 16. HgO

17. KBr +Cl2

18. AgNO3 + NaCl

19. Mg(OH)2 + H2SO3

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#5 - Combustion Means “add oxygen” Normally, a compound composed of

only C, H, (and maybe O) is reacted with oxygen – usually called “burning”

If the combustion is complete, the products will be CO2 and H2O.

If the combustion is incomplete, the products will be CO (or possibly just C) and H2O.

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Combustion Examples:

C4H10 + O2 (assume complete)

C4H10 + O2 (incomplete)

C6H12O6 + O2 (complete)

C8H8 +O2 (incomplete)

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SUMMARY: an equation... Describes a reaction Must be balanced in order to follow the

Law of Conservation of Mass Can only be balanced by changing the

coefficients. Has special symbols to indicate

physical state, if a catalyst or energy is required, etc.

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ReactionsCome in 5 major types.We can tell what type they are by

looking at the reactants.Single Replacement happens

based on the Activity SeriesDouble Replacement happens if the

product is a precipitate (insoluble solid), water, or a gas.

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Reactions in Aqueous Solution

OBJECTIVES:

–Describe the information found in a net ionic equation.

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Reactions in Aqueous Solution OBJECTIVES:

–Predict the formation of a precipitate in a double replacement reaction.

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Net Ionic Equations Many reactions occur in water- that is,

in aqueous solution Many ionic compounds “dissociate”, or

separate, into cations and anions when dissolved in water

Now we are ready to write an ionic equation

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Net Ionic Equations Example:

–AgNO3 + NaCl AgCl + NaNO3

1. this is the full equation

2. now write it as an ionic equation

3. can be simplified by eliminating ions not directly involved (spectator ions) = net ionic equation

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Na+

Al3+

S2–

2Ca2+ PO43–

3Cl–

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Review: forming ions Ionic (i.e. salt) refers to +ve ion plus -ve ion Usually this is a metal + non-metal or metal +

polyatomic ion (e.g. NaCl, NaClO3, Li2CO3) Polyatomic ions are listed on page 95 (aq) means aqueous (dissolved in water) For salts (aq) means the salt exists as ions NaCl(aq) is the same as: Na+(aq) + Cl–(aq) Acids form ions: HCl(aq) is H+(aq) + Cl–(aq),

Bases form ions: NaOH(aq) is Na+ + OH–

Q - how is charge determined (+1, -1, +2, etc.)?A - via valences (periodic table or see pg. 95) F, Cl gain one electron, thus forming F–, Cl–

Ca loses two electrons, thus forming Ca2+

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Charge can also be found via the compound E.g. in NaNO3(aq) if you know Na forms Na+, then

NO3 must be NO3– (NaNO3 is neutral)

By knowing the valence of one element you can often determine the other valences

Q - Write the ions that form from Al2(SO4)3(aq)?Step 1 - look at the formula: Al2(SO4)3(aq)Step 2 - determine valences: Al3 (SO4)2

Background: valences and formulas

(Al is 3+ according to the periodic table) Step 3 - write ions: 2Al3+(aq) + 3SO4

2–

(aq) Note that there are 2 aluminums because

Al has a subscript of 2 in the original formula

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Practice with writing ionsQ - Write ions for Na2CO3(aq)

A - 2Na+(aq) + CO32–(aq) (from the PT Na is 1+.

There are 2, thus we have 2Na+. There is only one CO3. It must have a 2- charge)

Notice that when ions form from molecules, charge can be separated, but the total charge (and number of each atom) stays constant.

Q - Write ions for Ca3(PO4)2(aq) & Cd(NO3)2(aq)

A - 3Ca2+(aq) + 2PO43–(aq)

A - Cd2+(aq) + 2NO3–(aq)

Q - Write ions for Na2S(aq) and Mg3(BO3)2(aq)

A - 2Na+(aq) + S2–(aq), 3Mg2+(aq)+ 2BO33–(aq)

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Types of chemical equationsEquations can be divided into 3 types

1) Molecular, 2) Ionic, 3) Net ionic Here is a typical molecular equation:Cd(NO3)2(aq) + Na2S(aq) CdS(s) + 2NaNO3(aq) We can write this as an ionic equation (all

compounds that are (aq) are written as ions):Cd2+(aq) + 2NO3

–(aq) + 2Na+(aq) + S2–(aq) CdS(s) + 2Na+(aq) + 2NO3

–(aq) To get the NET ionic equation we cancel out all

terms that appear on both sides:Net: Cd2+(aq) + S2–(aq) CdS(s)

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Equations must be balanced There are two conditions for molecular,

ionic, and net ionic equationsMaterials balance

Both sides of an equation should have the same number of each type of atom

Electrical balanceBoth sides of a reaction should have the same net charge

Q- When NaOH(aq) and MgCl2(aq) are mixed, _______(s) and NaCl(aq) are produced. Write balanced molecular, ionic & net ionic equationsMg(OH)2

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NaOH(aq) + MgCl2(aq) Mg(OH)2(s) + NaCl(aq)

Next, balance the equation

First write the skeleton equation2

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Ionic equation:2Na+(aq) + 2OH-(aq) + Mg2+(aq) + 2Cl-(aq)

Mg(OH)2(s) + 2Na+(aq) + 2Cl-(aq)

LiNO3

Ca3(PO4)2

NaC2H3O2

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Net ionic equation:

2OH-(aq) + Mg2+(aq) Mg(OH)2(s)

Write balanced ionic and net ionic equations:

CuSO4(aq) + BaCl2(aq) CuCl2(aq) + BaSO4(s)

Fe(NO3)3(aq) + LiOH(aq) ______(aq) + Fe(OH)3(s)

Na3PO4(aq) + CaCl2(aq) _________(s) + NaCl(aq)

Na2S(aq) + AgC2H3O2(aq) ________(aq) + Ag2S(s)

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Cu2+(aq) + SO42–(aq) + Ba2+(aq) + 2Cl–(aq)

Cu2+(aq) + 2Cl–(aq) + BaSO4(s)

Net: SO42–(aq) + Ba2+(aq) BaSO4(s)

Fe3+(aq) + 3NO3–(aq) + 3Li+(aq) + 3OH–(aq)

3Li+(aq) + 3NO3–(aq) + Fe(OH)3(s)

Net: Fe3+(aq) + 3OH–(aq) Fe(OH)3(s)

For more lessons, visit www.chalkbored.com

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2Na3PO4(aq) + 3CaCl2(aq) Ca3(PO4)2(s)+ 6NaCl(aq)

6Na+(aq) + 2PO43–(aq) + 3Ca2+(aq) + 6Cl–(aq)

Ca3(PO4)2(s)+ 6Na+(aq) + 6Cl–(aq)

Net: 2PO43–(aq) + 3Ca2+(aq) Ca3(PO4)2(s)

2Na+(aq) + S2–(aq) + 2Ag+(aq) + 2C2H3O2–(aq)

2Na+(aq) + 2C2H3O2–(aq) + Ag2S(s)

Net: S2–(aq) + 2Ag+(aq) Ag2S(s)

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Predicting the Precipitate Insoluble salt = a precipitate General solubility rules are found:

a) Table 8.3, p. 227b) Reference section - page 887

Table A.7 (back of textbook) gives combinations of cations and anions that are soluble or insoluble in water