TYPES of CHEMICAL REACTIONS (60 Points)

TYPES OF CHEMICAL REACTIONS (60 points)

Objectives:

• Perform and make observations of different types of chemical reactions
• Qualitatively identify products of these chemical reactions
• Write balanced chemical equations for these chemical reactions

Background:

There are an infinite number of chemical reactions. Chemist have divided these into broad classifications based on certain criteria. The most important classifications are; synthesis, combustion, decomposition, single replacement, double replacement, neutralization and redox. Note that some reactions will fall into more than one classification. For example, all single replacement reactions are also redox reactions.

Reactions are described with chemical equations. A chemical equation is the symbolic representation of the chemicals involved in the reaction. These chemical equations can be used to describe both physical processes and chemical reactions. All chemicalequations consist of reactants, the starting materials, products, the ending materials, the state of matter that the materials are in and a reaction arrow representing that a reaction has occurred.

reactants products

The states of matter are symbolized by subscripts which follow each chemical. The most common states of matter are solid, liquid, gas and aqueous. The first three should be self-explanatory the final, aqueous, occurs when a substance is dissolved in water. Gasses and ions are commonly found in an aqueous state. The symbols are as follows; solid (s), liquid (l), gas (g) and aqueous (aq).

Here are two examples of chemical equations:

H2O(s)  H2O(l)

CH4(g) + O2(g) CO2(g) + H2O(g)

The first is the physical process of ice melting the second is the combustion of methane gas. There is one other aspect necessary in writing a proper chemical equation, the equation must have the same number of each type atom on each side of the reaction arrow. This is the Law of Conservation of Matter, in other words you are not allowed to create nor destroy matter.The first of the two above equations is fine, two hydrogens on each side and one oxygen on each side. The second equation needs to be balanced. This can be accomplished by placing a 2 in front of both the O2(g) and the H2O(g).

CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g)

On each side we find one carbon, four hydrogens and four oxygens. This equation obeys the conservation of matter and is said to be balanced.

Occasionally you may see something written above or below the arrow. This is normally a catalyst, platinum and nickel are a common metal catalysts. Another possibility is that the something is an environmental necessity, for example, heat or ultra violet light.

2 CO(g) + O2(g) 2 CO2(g)

The above chemical equation occurs in an automobiles catalytic converter, converting carbon monoxide into carbon dioxide in the presence of a platinum catalyst. The platinum is the reason the catalytic converter is expensive.

Synthesis reactions, two or more reactants combine to form one product. The general equation is:

A + B  C

A specific example is the tarnishing of a silver tea set:

Ag(s)+ S(s) Ag2S(s)

Decomposition reactions, one reactant breaks to form two or more products. The general equation is:

A B +C

A specific example is the electrolysis of water, electricity is causes this reaction:

2 H2O(l) 2 H2(g) + O2(g)

Single replacement reactions, one chemical replace another in a compound. The general equation is:

A + BX  B + AX

A specific example is zinc metal replacing iron in the iron(III) oxide compound, ships use this reaction to keep their hulls from rusting.

3 Zn(s) + Fe2O3(s)  2 Fe(s) + 3 ZnO(s)

The reactivity of a substance depends on its ability to gain or lose electrons. It is possible to arrange the elements into a series based upon their reactivity. Such a list is called an activity series.

Activity Series for Metals and Non-metals

Single Replacement Reactions

Name Symbol
Metals
Lithium Li+
Sodium Na+
Potassium K+
Rubidium Rb+
Barium Ba+2
Strontium Sr+2
Calcium Ca+2 / Decreasing activity
↓
reacts with water and acids
Magnesium Mg+2
Aluminum Al+3
Manganese Mn
Zinc Zn+2
Chromium Cr+3
Iron Fe
Cadmium Cd+2
Cobalt Co+2
Nickel Ni+2
Tin Sn
Lead Pb / reacts with acids replacing hydrogen
Hydrogen H2
Antimony Sb
Bismuth Bi
Copper Cu
Mercury Hg
Silver Ag+1
Platinum Pt
Gold Au / fairly unreactive
Non-metals
Fluorine F2
Chlorine Cl2
Bromine Br2
Iodine I2

While there are many types of single replacement reactions, this lab will focus on two primary types. In one type, a more active metal, or halogen, replaces a less active metal or halogen from solution. An example of this is:

Zn(s) + CuSO4 (aq)  Cu (s) + ZnSO4(aq)

The other type of single replacement reaction involves the replacement of hydrogen from an acid by a metal. Consider the reaction between zinc and hydrochloric acid:

Zn(s) + 2 HCl (aq)  H2(g) + ZnCl2(aq)

The zinc metal is active enough to replace the hydrogen from the acid. Bubbles of hydrogen gas can be seen rising to the surface, and the piece of zinc is consumed. On the other hand, if the less reactive metal, copper, is placed into a hydrochloric acid solution, no reaction will take place.

Double replacement reactions, chemicals in each of two compounds switch compounds. The general equation is:

AB + CD  AD + CB

A specific example is allows for the removal of toxic barium from a water source by adding a compound containing sulfate.

BaCl2(aq) + Na2SO4(aq)  BaSO4(s) + 2 NaCl(aq)

Certain substances are not very soluble in water. Frequently such substances are generated in a reaction vessel by the addition of various other substances that are themselves soluble. For example, silver chloride is not soluble in water. If an aqueous solution of silver nitrate (very soluble) is mixed with an aqueous solution of sodium chloride (very soluble), the combination of silver ions from one solution and the chloride ions from the other solution produces silver chloride, which then forms a precipitate that settles to the to the bottom of the container. The solution that remains above the precipitate of silver chloride is a solution of sodium nitrate. The silver ions and the sodium ions have switched partners:

AgNO3 (aq)+ NaCl (aq)  AgCl (s)+ NaNO3(aq)

The silver ion and the sodium ion have replaced each other in this process; this sort of reaction is referred to as a double replacement reaction.

To clarify what is really happening in such reactions, it is often more instructive to write the reaction in its net ionic form. In a net ionic equation for a precipitation reaction, only the ions involved in actually forming the precipitate are shown; the other ions contained in the original reagents used are called spectator ions and are still present in the solution after the precipitate has been formed. For example, the net ionic reaction for the previous reaction is:

Ag+(aq) + Cl-(aq)  AgCl (s)

The net ionic reaction is especially instructive because it implies that any solution containing silver ions should react with any solution containing chloride ions since it is these ionic species that are really reacting. For example, if dilute hydrochloric acid, HCl, were added to a solution of silver nitrate, a precipitate would be expected to form. The reactions of silver nitrate with sodium chloride and of silver nitrate and hydrochloric acid are exactly the same when the net ionic reaction is considered.

Neutralization reactions occur when the hydrogen ion from an aqueous acid combines with the hydroxide ion from an aqueous base to produce water and an aqueous salt. For example,

HNO3 (aq) + NaOH (aq)  H2O (l) + NaNO3(aq)

The net ionic equation for neutralization reactions is the same and is typical of the reaction between acids and bases in aqueous solution:

H+(aq) + OH-(aq)  H2O (l)

Predicting Products

The products of a chemical reaction may often be predicted by applying known facts about common reaction types. While there are hundreds of different ”kinds” of chemical reactions, only four general types of reactions will be considered; single displacement, double displacement, decomposition, and synthesis. The following sections give examples of these general types.

Synthesis

In a synthesis reaction two or more simple substances (compounds and/or elements) are combined to form one new and more complex substance. Here the general form is

Element + element  compound

or

compound + compound  compound

A + XAX

The following are some general types of synthesis reactions.

1. Combination of elements.

Fe(s) + S(1)  FeS(s)

2Na(s) + Cl2(g)  2NaCl(s)

2. Combination of an acid anhydride with water to give an acid.

SO2(g) + H2O(1)  H2SO3(aq)

N2O3(g) + H2O(1)  2HNO2(aq)

CO2(g) + H2O(1) H2CO3(aq)

P2O5(s) + 3H2O(l) 2H3PO4(aq)

3. Combination of a basic anhydride or a metallic oxide with water to form a base.

Na2O(s) + H2O(1)  2NaOH(aq)

CaO(s) + H2O(1)  Ca(OH)2(aq)

BaO(s) + H2O(1)  Ba(OH)2(aq)

4. Combination of the metal of a basic oxide with the nonmetal of an acidic oxide to form a salt.

CO2(g) + Na2O(s) Na2CO3(s)

P2O5(s) + 3BaO(s)  Ba3(PO4)2(s)

SO2(g) + MgO(s)  MgSO3(s)

Decomposition

When energy in the form of heat, electricity, light, or mechanical shock is supplied, a compound may decompose to form simpler compounds and/or elements. The general form for this type of reaction is

compound  two or more substances

AX  A + X

The following are some general types of decomposition reactions.

1. If some acids are heated, they decompose to form water and an acidic oxide.

H2SO3(aq)  SO2(g) + H2O(1)

H2CO3(aq) CO2(g) + H2O(1)

2. When some metallic hydroxides are heated, they decompose to form a metallic oxide and water.

Ca(OH)2(s)  CaO(s) + H2O(g)

2Fe(OH)3(s)  Fe2O3(s) + 3H2O(g)

3. Some metallic carbonates decompose to form a metallic oxide and carbon dioxide when heated.

Li2CO3(s)  Li2O(s) + CO2(g)

CaCO3(s)  CaO(s) + CO2(g)

4. Metallic chlorates decompose to form metallic chlorides and oxygen when heated.

2KC103(s)  2KC1(s) +302(g)

Ni(C103)2(s)  NiC12,(s) + 302(g)

5. Most metallic oxides are stable, but a few decompose when heated.

2HgO(s)  2Hg(1) + O2,(g)

2Ag2O(s)  4Ag(s) + O2(g)

6. Some compounds cannot be decomposed by heat, but can be decomposed into their elements by electricity.

2H2O(1)  2H2(g) + O2(g)

2NaC1(1) 2Na(1) + Cl2,(g)

MgCl2(1)  Mg(1) + Cl2(g)

2H2O(1)  2H2

Single Displacement

One metallic element displaces another metallic element in a compound, or a nonmetallic element displaces another nonmetallic element in a compound. A single displacement has the general form

element + compound element + compound

A+BX AX+B

X+BY BX+Y

The following are some general types of single displacement reactions.

1. An active metal will displace the metallic ion in a compound of a less active metal.

Fe(s) + Cu(NO3)2(aq)  Fe(NO3)2(aq) + Cu(s)

Pb(s) + 2AgC2H3O2(aq)  Pb(C2H3O2)2,(aq) + 2Ag(s)

1. Some active metals such as sodium and calcium will react with water to give a metallic hydroxide and hydrogen gas.

2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)

Ca(s) + 2H2O(1)  Ca(OH)2(aq) + H2(g)

1. Active metals such as zinc, iron, and aluminum will displace the hydrogen in acids to give a salt and hydrogen gas.

Zn(s) + 2HC1(aq)  ZnCl2(aq) + H2(g)

Fe(s) + H2SO4(aq)  FeSO4(aq) + H2(g)

4. Halogens (which are active nonmetals) will displace less active halogens.

Cl2(g) + 2NaBr(aq)  2NaC1(aq) + Br2(aq)

Br2(g) + 2KI(aq)  2KBr(aq) + I2(g)

For example, use the Activity Chart to either predict the products or determine that a reaction does not take place:

(a) magnesium is added to a solution of iron (III) chloride:

First write the formulas and physical states of the reactants:

Mg (s) + FeCl3 (aq)

Then predict the products based on the Activity Chart:

Mg (s) + FeCl3 (aq)  MgCl2 + Fe

Balance the equation:

3 Mg (s) + 2 FeCl3 (aq) 3 MgCl2 + 2 Fe

Insert the physical states:

3 Mg (s) + 2 FeCl3 (aq)  3 MgCl2(aq) + 2 Fe (s)

Double Displacement

The positive and negative ions of two compounds are interchanged. The form of these reactions is easy to recognize,

compound + compound  compound + compound

AX + BY  AY + BX

The following are some general types of double displacement reactions.

1. A reaction between an acid and a base yields a salt and water. Such a reaction is a neutralization reaction.

2KOH(aq) + H2SO4(aq)  K2SO4 (aq) +H2O(l)

Ca(OH)2(aq) + 2HNO3(aq)  Ca(NO3)2(aq) + 2H2O(1)

2. Reaction of a salt with an acid forms a salt of the acid and a second acid which is volatile.

2KNO3(aq) + H2SO4(aq)  K2SO4(aq) + 2HNO3(g)

FeS(c) + 2HC1(aq)  FeCI2(aq) + H2S(g)

2NaC1(aq) + H2SO4(aq)  Na2SO4(aq) + 2HCl(g)

2a. This same reaction of a salt with an acid or base may yield a compound which can be decomposed.

CaCO3(aq) + 2HC1(aq)  CaCl2(aq) + H2CO3(aq) H2CO3(aq) CO2(g) + H2O(1)

or

CaCO3(s) + 2HC1(aq)  CaCl2(aq) + CO 2(g) + H2O(1) K2SO3(aq) + 2HNO3(aq)  2KNO3(aq) + SO2(g) + H2O(1) NH4C1(aq) + NaOH(aq)  NaC1(aq) + NH2(g) + H2O(1)

3. Reaction of some soluble salts produces an insoluble salt and a soluble salt.

AgNO3(aq) + NaC1(aq)  AgC1(s) + NaNO3(aq)

Na2SO4(aq) + Ba(NO3)2(aq)  BaSO4(s) + 2NaNO3(aq)

CuSO4(aq) + Na2S(aq)  CuS(s) + Na2SO4(aq)

Ca(C2H3O2)2(aq) + (NH4)2CO3(aq)  CaCO3(s) + 2NH4C2H3O2(aq)

The solid precipitate can be predicted using the solubility rules/chart:

Solubility Chart

I - Insoluble, S - Soluble, SS - slightly soluble

Ag / Al / Ba / Bi / Ca / Cd / Co / Cr / Cu / Fe / H / Hg / K / Mg / Mn / Na / NH4 / Ni / Pb / Sn / Sr / Zn
Acetate / I / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S
Bromide / I / S / S / S / S / S / S / S / S / S / I / S / S / S / S / S / S / I / S / S / S
Carbonate / I / I / I / I / I / I / I / I / I / I / S / I / S / I / I / S / S / I / I / I / I / I
Chlorate / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S
Chloride / I / S / S / S / S / S / S / S / S / S / S / I / S / S / S / S / S / S / I / S / S / S
Chromate / I / S / I / S / S / S / S / S / S / S / S / I / S / S / S / S / S / S / I / S / I / S
Cyanide / I / S / I / I / S / I / I / I / I / S / S / S / S / S / S / I / I / S / I
Fluoride / S / S / I / S / I / S / S / S / I / I / S / S / I / I / S / S / I / I / S / I / I
Hydroxide / I / I / S / I / SS / I / I / I / I / I / S / I / S / I / I / S / S / I / I / I / SS / I
Iodide / I / S / S / S / S / S / S / S / S / S / S / I / S / S / S / S / S / S / I / S / S / S
Nitrate / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S / S
Oxide / I / I / S / I / I / I / I / I / I / I / S / I / S / I / I / S / SS / I / I / I / S / I
Phosphate / I / I / I / I / I / I / I / I / I / I / S / I / S / I / I / S / S / I / I / I / I / I
Silicate / I / I / I / I / I / I / I / I / I / I / S / I / S / I / I / S / S / I / I / I / I / I
Sulfate / SS / S / I / S / SS / S / S / S / S / S / S / I / S / S / S / S / S / S / I / S / I / S
Sulfide / I / I / I / I / I / I / I / I / I / I / S / I / S / I / I / S / S / I / I / I / I / I
Sulfite / I / I / I / I / I / I / I / I / I / I / S / I / S / I / I / S / S / I / I / I / I / I

General Solubility Trends:

• All compounds of the ammonium ion (NH4+), and of the Alkali metal (Group IA) cations, are soluble.
• All nitrates and actetates are soluble.
• All chlorides, bromides, and iodides are soluble EXCEPT those of silver, lead, and mercury(I).
• All sulfates are soluble EXCEPT those of silver, lead, mercury(I), barium, strontium, and calcium.
• All carbonates, sulfites, and phosphates are insoluble EXCEPT those of ammonium and alkali metal (Group I) cations.
• All hydroxides are insoluble EXCEPT those of ammonium, barium, and alkali metal (Group I) cations.
• All sulfides are insoluble EXCEPT those of ammonium, Alkali metal (Group I) cations, and Alkali earth metal (Group II) cations.
• All oxides are insoluble EXCEPT those of calcium, barium, and Alkali metal (Group I) cations; these soluble ones actually react with the water to form hydroxides.

For example:

Predict the products of a reaction between aqueous rubidium phosphate and aqueous

titanium (IV) nitrate.

First write the formulas of the reactants:

Rb3PO4 (aq) + Ti(NO3)4 (aq)

Then predict the products based on the cations trading places with one another:

Rb3PO4 (aq) + Ti(NO3)4 (aq) Ti3(PO4)4 + RbNO3

Balance the equation:

4 Rb3PO4 (aq) + 3 Ti(NO3)4 (aq)  Ti3(PO4)4 + 12 RbNO3

Finally, based on the solubility rules, insert the physical states of the products. The rules state:

“All sulfates are soluble EXCEPT those of silver, lead, mercury(I), barium, strontium, and calcium.” And “All nitrates and actetates are soluble.” Therefore, the Ti3(PO4)4 compound will be a solid and the RbNO3 will be aqueous.

4 Rb3PO4 (aq) + 3 Ti(NO3)4 (aq)  Ti3(PO4)4 (s) + 12 RbNO3 (aq)

Materials

1 M NaCl.1 M Cu(NO3)2

1 M NaOH.1 M Fe(NO3)3

1 M NaNO3.1 M AgNO3

6 M HCl 6 M NH4OH

2 M HCl CuCO3

Mossy zincMagnesium ribbon

Copper woolSteel wool

Equipment

Pipets, thin stemBeakers

Distilled water and wash bottleReaction plate, 48-well plate

BalanceToothpicks

Paper towelsBunsen burner

Crucible tongsMicrospatula

Test tubesTest tube holder

Wood splintsTest tube rack

Evaporating dishMatches

Safety Precautions

Copper(II) nitrate and silver nitrate solutions are slightly toxic by ingestion. Silver nitrate and sodium hydroxide solutions are skin and eye irritants; silver nitrate will also stain skin and clothing. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles. Wash hand thoroughly if you contact any of the above chemicals.

Wash hands thoroughly with soap and water before leaving the lab.

Laboratory Procedure:

For each reaction you perform, make observations and predict the products formed. Below is a section titled results, write the balanced chemical equations for these reactions here. If no observable reaction occurred write no reaction or NR.

This lab will use a flaming or smoldering piece of wood as a qualitative test for either CO2(g) or O2(g). If the fire goes out, you have CO2 being generated, if the fire flames up, you have O2 being generated.

1. With your lab partner, go to a chemical reaction station.

2. Record, in your data table, your observations of the appearance of the reactants before you perform the reaction.

3. Follow the instructions at each station to complete the chemical reaction.

4. Record your observations both during and after the reaction.

5. Clean all equipment and dispose of the chemicals properly.

6. Move to another station and repeat steps (1) – (5).

Station #1:

Magnesium metal and oxygen gas

a. Acquire a piece of magnesium metal approximately 2 cm long, scrub it will steel wool to return its luster.

b. Light the Bunsen burner.

c. Using crucible tongs, hold the magnesium ribbon in the burner flame until it ignites.

d. Once ignited, place it in the evaporating dish to burn.

e. Note and record its appearance.

Station #2:

Hydrogen Chloride and Ammonia

a. In a fume hood, add 5 drops of 6 M HCl, hydrochloric acid, to a test tube.

b. In a fume hood, add 5 drops of 6 M NH4OH, ammonium hydroxide, to a test tube.

c. Remove the test tubes from the hood. Bring them together so the test tube mouths are as close as possible to each other and tilted toward each other.

d. Record any observations.

e. Return the test tubes to the “used test tube beaker” in the fume hood

Station #3:

Copper(II) Carbonate

a.Measure as close to 1.000 gram of CuCO3, copper(II) carbonate onto a piece of paper and place it into a clean, dry test tube.

b.Heat the test tube and contents gently, moving the test tube in and out of the flame from the top of the solid to the bottom of the solid.

c.Using the glowing splint test, check periodically for the presence of an evolving gas as you heat the test tube.

d.After a couple of minutes of gentle heating, increase the flame and heat vigorously for no more than 10 minutes.

e.You should not quit heating until you get a negative test for the presence of this gas.

Allow five minutes of cooling before you remove the test tube and mass it along with its contents.

Station #4:

Silver Nitrate and Copper

a. Put 10 drops of 0.10 M AgNO3, silver nitrate into a test tube.

b. Insert a short piece of copper wool into the test tube.

c. Allow the reaction to take place for at least 30 minutes.

d. Check up on this several times during the 30 minutes and record ALL changes.

Station #5:

HCl and Zinc

a. Add 5ml of 6M HCl to a clean dry test tube.

b. Affix the test tube to a ring stand.

c. Catch a wood splint on fire.

d. Add a small piece of Zn metal to the test tube, wait 10 seconds.

e. Carefully insert the burning splint into the test tube, expect a sound and pop.

f. Make observations of this second test tube.

Station #6:

Copper Nitrate and Zinc

a. Add 10 drops of 0.10 M Cu(NO3)2, copper(II) nitrate to a test tube.

b. Add a small piece of zinc metal into the test tube.

c. Allow the reaction to take place for at least 30 minutes.

d. Check up on this several times during the 30 minutes and record ALL changes.

Station #7: Well-Plate Reactions:

Place the well plate on a paper towel and number the well plate wells as follows:

Number one set of wells 1 – 3, then skip a well and number the next set 4 – 6, then skip a well and number the last set 7 – 9.