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Learning Objectives
- Perform and observe the results of a variety of single and double replacement reactions.
- Become familiar with some of the observable signs of these reactions
- Identify the products formed in each of these reactions.
- Write balanced chemical equations for each of the double replacement reactions studied.
Background
During a chemical reaction both the form and composition of matter are changed. Old substances are converted to new substances, which have unique physical and chemical properties of their own. Some of the observable signs that a chemical reaction has occurred include:
- Metallic deposit
- Appearance of bubbles
- Temperature change occurs
- Color change occurs
- Precipitate (cloudy, tiny particles) appears
Note that there are other observable signs for chemical reactions, but these are most likely to be seen in this lab
Single Replacement Reaction
In Part A of this lab we will examine single replacement reactions. This is one type of oxidation-reduction reaction, or redox reaction, because it occurs via a transfer of electrons.
Single replacement reactions have the general form: \( \ce{A + BC -> B + AC} \)
Here, A is an element and BC is usually an aqueous ionic compound or an acid (consisting of B+ and C- aqueous ions). Element A replaces element B in the compound BC; this results in the formation of a new element B and a new ionic compound or acid, AC. If the new element B is a metal, it will appear as a metallic deposit. If it is a gas, it will appear as bubbles.
Example \(\PageIndex{1}\)
magnesium metal + aqueous aluminum chloride
Since Mg is more active than Al, a single replacement reaction will occur. The predicted products are aluminum metal and aqueous magnesium chloride
Reaction Equation
\[ \ce{3 Mg + 2 AlCl3->2 Al + 3 MgCl2} \]
Example \(\PageIndex{2}\)
Aluminum Metal + Aqueous Magnesium Chloride
Since Al is not more active than Mg, a replacement will NOT occur.
Reaction Equation
\[ \ce{Al + MgCl2 -> } \text{NO REACTION} \]
Double Replacement Reactions
All double replacement reactions have the general form:
\[ \ce{AB + CD -> AD + CB} \]
Reactions that can be classified as double replacements include precipitation reactions, neutralization reactions and gas forming reactions.
Precipitation Reactions
Here AB and CD are usually aqueous ionic compounds (or acids) consisting of aqueous ions (A+ and B- , C+ and D- ). When a double replacement reaction occurs, the cations and anions switch partners, resulting in the formation of two new ionic compounds AD and CB, one of which is in the solid state. This solid product is an insoluble ionic compound called a precipitate. To determine whether a product ionic compound will be soluble or insoluble, consult the Solubility Rules provided at the end of the Background section. Note that if both of the predicted products are soluble, a precipitation reaction will not occur.
Example \(\PageIndex{3}\)
aqueous lead (II) nitrate + aqueous sodium chloride
The predicted products are lead (II) chloride (insoluble) and sodium nitrate (soluble). Since one of the predicted products is insoluble, a precipitation reaction will occur.
Reaction Equation
\[ \ce{Pb(NO3)2(aq) + 2 NaCl (aq) ->2NaNO3 (aq) + PbCl2(s)} \]
Neutralization Reactions
Here AB is an acid (consisting of H+ and X- aqueous ions) and BC is a base (consisting of M+ and OH- ions). When a double replacement reaction occurs, the cations and anions switch partners, resulting in the formation of water and a new ionic compound (or salt), which is usually soluble. Neutralization reactions are exothermic, and are generally accompanied by a noticeable release of heat.
Example \(\PageIndex{4}\)
sulfuric acid + aqueous lithium hydroxide
The predicted products are water and lithium sulfate.
Reaction Equation
\[ \ce{H2SO4(aq) + 2 LiOH(aq) ->Li2SO4(aq) + H2O(l)} \]
Gas Forming Reactions
In these reactions one of the products (AD or CB) after the double replacement is in the gaseous state, such as hydrogen sulfide (H2S) or ammonia (NH3). One of the products could also be carbonic acid (H2CO3) or sulfurous acid (H2SO3). Both carbonic acid and sulfurous acid are unstable and will decompose to form carbon dioxide and sulfur dioxide gases, respectively:
\[ \ce{H2CO3(aq) ->H2O(l) + CO2(g)} \label{Carbonic Acid} \]
\[ \ce{H2SO3 (aq) ->H2O (l) + SO2 (g)} \label{Sulfurous Acid} \]
Example \(\PageIndex{5}\)
nitric acid + aqueous sodium bicarbonate
The predicted products are carbonic acid and sodium nitrate. However carbonic acid decomposed to carbon dioxide and water.
Reaction Equation
\[ \ce{HNO3 (aq) + NaHCO3 (aq) ->NaNO3(aq) + H2CO3 (aq)} \]
Decomposes
Final equation:
\[ \ce{HNO3 (aq) + NaHCO3 (aq) ->NaNO3 (aq) + H2O(l) + CO2(g)} \]
Writing Equations for Reactions
Write chemical formulas for each reactant and place a yield arrow (\(\ce{->}\)) after the last reactant.
Identify the reaction type – precipitation, neutralization or gas forming. If you determine that a reaction will occur, write the correct formula(s) of the products after the arrow.
If you determine that a reaction will not occur, write “no reaction” after the arrow. Balance the equation (to ensure mass conservation). Be sure to include the physical states of all reactants and products in your final equation.
Review: Chemical Formulas and Names of Ionic Compounds
Ionic compounds are formed when positive ions (cations) and negative ions (anions) are attracted to each other via strong electrostatic forces. This attraction is called ionic bond. The basic rules for writing the chemical formulas of ionic compounds:
- Determine the formulas and charges on the cation and anion involved in the compound.
- Combine the ions in a ratio that results in the formation of a neutral ionic compound.
- The charge of all the cations must equal the total charge of all the anions in the compound. The number of each element present in the compound becomes a subscripts in the chemical formula.
Example \(\PageIndex{6}\)
Write the formula for iron (III) chloride
The basic rules for writing names of ionic compounds:
- Both the cation and anion must be named.
- Always name the cation first, then the anion.
Solution
First identify the cation and the anion in this compound – iron (III) – Fe cation and chloride Cl - 3+ anion. For a neutral compound, one Fe is needed for every 3 Cl - . The formula of the compound is FeCl3.
Example \(\PageIndex{7}\)
Name the ionic compound Al(NO3)3. First identity the cation and anion in this compound.
Solution
Cation = Al+3: aluminum
Anion = NO3- : nitrate
The name of this compound is aluminum nitrate.
Experimental Procedure
Safety
- Be especially cautious when using 6M HCl, 3M H2SO4 and 6M NaOH as they can burn your skin.
- Be aware that skin discoloration will result from contact with AgNO3. If you feel any tingling sensations or see any color changes on your skin, flush with water immediately for a minimum of 15 minutes. Inform your instructor of any chemical contact as soon as possible.
- WEAR SAFETY GOGGLES, CLOSED-TOE SHOES
Materials and Equipment
Solids: sodium bicarbonate, copper, zinc, magnesium
Solutions: 6 M sodium hydroxide, 3 M sulfuric acid, 6 M hydrochloric acid, 1 M sodium chloride; all other solutions are 0.1 M and include sodium chloride, iron (III) chloride, ammonium hydroxide, sodium carbonate, cobalt (II) chloride, sodium phosphate, copper (II) sulfate, potassium nitrate, barium chloride, silver nitrate, copper (II) sulfate, zinc nitrate, aluminum sulfate. Equipment: 9 small test tubes, test tube rack, large beaker, 2 small beakers.
Part A: Single Replacement Reactions
Use the medium sized test tubes; place them in a large beaker. Use clean test tubes that have been rinsed with distilled water. The test tubes do not have to be dry. Use 3-mL quantities of all solutions. A good estimate is to measure out 3 mL of water in a test tube and use it as a comparison.
Place one piece of metal in the test tube first, and then add the solution. The metal should be completely immersed in the solution used. If results are not obtained immediately, give the reaction some time. Some reactions take longer than others.
Perform the following reactions, and record your observations for each on the data sheet. All waste is to be disposed of in the container in the hood!
- Zinc metal + hydrochloric acid
- Copper metal + aqueous silver nitrate
- Copper metal + aqueous zinc nitrate
- Magnesium metal + hydrochloric acid
PartB: Double Replacement Reactions
Use approximately 3-mL quantities of all solutions. A good estimate is to measure out 3 mL of water in a test tube and use it as a comparison.
Perform the following reactions, and record your observations for each on the data sheet. If results are not obtained immediately, give the reaction some time. Some reactions take longer than others. All waste is to be disposed of in the container in the hood!
- Aqueous sodium chloride + aqueous silver nitrate
- Aqueous sodium phosphate + aqueous copper (II) sulfate
- Hydrochloric acid + solid sodium bicarbonate (just a small scoop)
- Aqueous barium chloride + sodium sulfate
- Hydrochloric acid + aqueous sodium hydroxide
- Aqueous sodium carbonate + aqueous cobalt (II) chloride
- Aqueous sodium chloride + aqueous potassium nitrate
- Aqueous iron (III) chloride + aqueous ammonium hydroxide
When finished, complete the data sheet by writing the balanced equation for each reaction.
