Net Ionic Reactions

Single & Double Replacement

REPLACEMENT REACTIONS - one element replaces another in a compound - more active metals replace less active metals OR more active nonmetals replace less active nonmetals (see activity chart on other side) A + BC AC + B

KEY: REACTANTS ARE 1 ELEMENT AND 1 COMPOUND - ORDER NOT IMPORTANT - ELEMENT CAN’T BE O2

A. More active metals replace less active metals from their compounds. Hydrogen can act like a metal.

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

Zn + H2SO4 ZnSO4 + H2

B. More active nonmetals replace less active nonmetals from their compounds.

Cl2 + 2 NaI 2 NaCl + I2

C. Very active metals (only the first 5) can replace one (AND ONLY ONE) of the hydrogens from the stable compound water.

2 Na + 2 H2O 2 NaOH + H2 a metal hydroxide is always one of the

products.

IONIC REACTIONS - (also called double replacement reactions) ions, usually in solution, exchange respective cations and anions - these reactions occur when ions are removed from solution by the formation of a precipitate or by the formation of a molecular compound. These reactions can be further divided into the following 3 classifications below:

KEY: REACTANTS ARE 2 COMPOUNDS - UNLESS BOTH ARE OXIDES AB + CD AD + CB

A. Precipitation Reaction - one of the possible products must be a precipitate or no reaction occurs. AgNO3 + HCl AgCl + HNO3 AgCl is a precipitate.

B. Molecular Reaction - the reaction occurs if one of the products is a molecular compound (examples: CO2, H2S, SO2, H2O)H2SO4 + Na2O Na2SO4 + H2O H2O is molecularCaCO3 + 2 HCl CO2 + H2O + CaCl2 CO2 is molecular

C. Acid - Base Reaction - an acid plus a base yields an ionic compound (a salt) and water (a molecular compound).

2 H3PO4 + 3 Ca(OH)2 Ca3(PO4)2 + 6 H2O these are called HCl + NaOH NaCl + H2O neutralization reactions

Directions:For each of the following three reactions, write a balanced equation for the reaction in part (i) and answer the question about the reaction in part (ii). In part (i), coefficients should be in terms of lowest whole numbers. Assume that solutions are aqueous unless otherwise indicated. Represent substances in solutions as ions if the substances are extensively ionized. Omit formulas for any ions or molecules that are unchanged by the reaction. You may use the empty space at the bottom of the next page for scratch work, but only equations that are written in the answer boxes provide will be scored.

Consider the following reaction:

potassium chromate + barium nitrate barium chromate + potassium nitrate

K2CrO4 + Ba(NO3)2 BaCrO4 + 2KNO3

This is an ionic reaction and it occurs in water (like most ionic reactions).

The barium chromate is a precipitate. To indicate what is happening in this reaction we sometimes include the physical state of the reactants and products as follows: (aq) means aqueous or dissolved in water

K2CrO4(aq) + Ba(NO3)2(aq)BaCrO4(s) + 2KNO3(aq)

K2CrO4(aq) + Ba(NO3)2(aq)BaCrO4(s) + 2KNO3(aq)

The above equation is called the molecular equation.

The complete ionic equation as written below:

2K+1(aq) + CrO4-2(aq)+ Ba+2(aq) + 2NO3

-1 (aq) BaCrO4(s) + 2K+1(aq)+ 2NO3

-

1(aq)

Notice that all strong electrolytes are written as ions.

Also notice that there are K+1 and NO3-1 ions on both

sides of the reaction. Since these ions are on both sides of the reaction, they are not actually part of the reaction and are called spectator ions.

The ions that actually participate in the chemical reaction are the following:

CrO4-2 (aq) + Ba+2 (aq) BaCrO4 (s)

This last equation is called the net ionic equation and includes only the participating ions of the reaction.

An example forming a molecule instead of a precipitate

potassium hydroxide + hydrochloric acid water + potassium sulfate

Molecular equation:

2KOH (aq) + H2SO4(aq) 2H2O (l) + 2KCl (aq)

Complete ionic equation:2 K+1(aq) + 2OH-1(aq) + 2H+1(aq) + 2Cl-(aq)

2H2O(l) + 2K+1(aq) + 2Cl-(aq)Net ionic equation:

2 OH-1(aq) + 2 H+1(aq) 2 H2O(l)

A little lie we told….Not all acids formed in ionic reactions indicate a reaction

…. only WEAK acids

When we write ionic or net ionic reactions, all strong electrolytes are written as ions.

Reactions that produce one of the 7 strong acids may not occur.

Ex: NaCl (aq) + HNO3(aq) NaNO3 (aq) + HCl(aq)

Na+ + Cl- + H+ + NO3- Na+ + NO3

- + H+ + Cl-

No new substances are formed = no reaction

Oxidation-Reduction (Redox) Reactions

A process where electrons are transferred from one substance to another. One substance is oxidized while another substance is reducedOxidation

• loss of electrons• an increase in oxidation number

Reduction- gain of electrons- a decrease in oxidation number

How can you tell when a redox reaction is taking place?

1) Assign oxidation numbers to atoms in substances

2) Compare oxidation numbers before and after reaction to determine if atom has lost or gained electrons

Rules For Assigning Oxidation Numbers - A bookkeeping

system1. The oxidation number of an atom in an element is 0.

Examples: Na, H2, Br2, S8, NeOx. # 0 0 0 0 0

2. The oxidation number of a monatomic ion is the same as its charge.Examples: Na+1, Ca+2, Al+3, Cl-1, O-2

Ox # +1 +2 +3 -1 -2

3. The sum of the oxidation numbers of all atoms in a neutral compound is zero.

4. The sum of the oxidation numbers of all atoms in an ion is equal to the charge on the ion.

5. In compounds, fluorine is always assigned an oxidation number of -1

6. Hydrogen’s oxidation number will be - +1 when bonded to a nonmetal (HCl)- -1 when bonded to a metal (NaH)

Examples:

NaH CaH2 HCl H2SNa—H H—Ca—H H—Cl H—S—H +1 -1 -1 +2 -1 +1 -1 +1 -2 +1

7.Oxygen usually has an oxidation number of -2. Exceptions: in peroxides, oxygen will be -1 and when combined with F, it will be +2 and in O2 it will be 0Examples:

H2O CaO H2O2

H — O — H Ca—O H—O—O—H

O2-2 OF2

[O—O]-2 F—O—F

+1 -2 +1 +2 -2 +1 -1 -1 +1

-1 -1 -1 +2 -1

8. Halogens usually have an oxidation number of -1. Exception is when chlorine, bromine, iodine are combined with oxygen

Examples:

NaCl MgI2 OCl2 HOBr

Na—Cl, I—Mg—I, Cl—O—Cl, H—O—Br

+1 -1 -1 +2 -1 +1 -2 +1 +1 -2 +1

** If none of the above rules help you get started…look for a atom

with a known charge and use that charge as its oxidation number

CdS: Cd-S

+2 -2

More examples

AlH3

Al +3

H -1

S2O3-2

S +2

O -2

Na2Cr2O7

Na +1

O –2

Cr +6

CO2

O -2

C +4

Use algebra to determine oxidation numbers of "difficult" atoms.

Example: H2SO4

H is +1 * 2 = +2 O is -2 * 4 = -8

2 + x + -8 = 0 S is +6

Example: ClO4-1

O is -2 * 4 = -8 -8 + x = -1

Cl is +7

Example: NH4+1

H is +1* 4 = +4

4 + x = 1

N is -3

Tricky Ones:FeSO4

The SO4 part has to have a charge of -2O is -2S is +6

To make a neutral compound Fe must be +2

Fe2(SO4)3

The SO4 part has to have a charge of -2O is -2S is +6There are 3 sulfate ions for a total negative charge of -6We need a total positive charge of +6Each Fe must be +3

2 Fe2O3 (s) + 3 C (s) 4 Fe (s) + 3 CO2 (g)

+3 -2 0 0 +4 –2

Fe is reduced, going from +3 to 0 and C is oxidized, going from 0 to +4, the O undergoes no change

Reducing agent

Causes reduction

Loses electrons

Undergoes oxidation

Oxidation number of atom increases

Oxidizing agent

Causes oxidation

Gains electrons

Undergoes reduction

Oxidation number of atom decreases

Assign oxidation numbers, indicate what is oxidized and reduced, indicate what is the

oxidizing agent and reducing agent

1) Ca (s) + 2 H+1 Ca+2 (aq) + H2 (g) 0 +1 +2 0

Ca is oxidized – increasing from 0 to +2H+1 is reduced – decreasing from +1 to 0 Ca is the reducing agentH+1 is the oxidizing agent

2) 2 Fe+2 (aq) + Cl2 (aq) 2 Fe+3 (aq) + 2 Cl-1 (aq)

+2 0 +3 -1

Fe+2 is oxidized – increasing from +2 to +3

Cl2 is reduced – decreasing from 0 to -1

Fe+2 is the reducing agent

Cl2 is the oxidizing agent

In general, metals act as reducing agents (are oxidized) and nonmetals act a oxidizing agents (are reduced).

 

Replacement Reactions

Are redox reactions!

(So are many of the reactions we studied in chapter 3!)

Example: It is possible for metals to be oxidized in the presence of a salt (in solution):

Fe(s) + Ni(NO3)2(aq) Fe(NO3)2(aq) + Ni(s)

OR

Fe(s) + Ni+2(aq) Fe+2(aq) + Ni(s)

In this reaction iron has been oxidized to Fe2+ while the Ni2+ has been reduced to Ni.

The Activity Series

We can list metals in order of decreasing ease of oxidation.

This list is the activity series.

The metals at the top of the activity series are called active metals and are easily oxidized.

The metals at the bottom of the activity series are called noble metals and NOT easily oxidized.

A metal in the activity series can only be oxidized by a metal ion below it.

(Higher will replace lower)

If we place Cu into a solution of Ag+ ions, then Cu2+ ions can be formed because Cu is above Ag in the activity series:

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

or

Cu(s) + 2Ag+(aq) Cu2+(aq) + 2Ag(s)

Decompositon of acids:F. Acids (hydrogen combined with some

negative ion) decompose into water and nonmetallic oxides.

H2CO3 H2O + CO2

2 H3PO4 P2O5 + 3 H2O

*This is NOT a redox reaction – the oxidation number on the nonmetal must stay constant. Use this knowledge to predict the formula for the nonmetal oxide.

Acid Decomposition

H2SO3 SxOy + H2O

Hydrogen: +1

Oxygen: -2

Sulfur: +4

H2SO3 SO2 + H2O

Sulfur: +4

Since oxygen usually

Has an oxidation number

Of -2. The formula must be

SO2.

Composition Reactions:E. Oxides of nonmetals combine with water to form

compounds called acids. These acids will usually be made from an “ate” ion.

P2O5 + 3 H2O 2 H3PO4

SO3 + H2O H2SO4

F. Oxides of metals combine with oxides of nonmetals to yield ionic compounds. These compounds usually contain an “ate” ion.

Na2O + SO3 Na2SO4

*These are NOT redox reactions…oxidation numbers stay constant. Use that knowledge to predict the product!

SO2 + H2O H2SOx

O: -2 O: -2

S: +4 S: +4

H: +1

In order for the oxidation numbers to add to zero, the formula must be H2SO3

P4O10 + CaO Ca3(POx)2

O: -2 O: -2

P: +5 P: +5

H: +1

In order for the oxidation numbers to add to zero, the formula must be Ca3(PO4)2