GCC CHM 151LL: Chemical Reactions: Introduction to Reaction Types © GCC, 2014 page 1 of 12
Chemical Reactions: Introduction to Reaction Types
In this experiment, you will observe (on video and in lab) examples of the six basic types of chemical
reactions which are described below. You will write your observations of the reactions. Based on
those observations and the given word equations, you will write complete, balanced chemical
equations to effectively communicate the chemistry of the reactions.
Introduction to Reaction Types
1. Combination: A + B → AB
In a combination reaction, a new compound can be formed in one of three ways, by combining: a) 2
elements, b) 1 element and 1 binary compound (consisting of 2 elements), or c) 2 binary compounds.
The following are examples of combination reactions:
The rusting of iron: 4Fe (s) + 3O2 (g) → 2Fe2O3 (s)
The formation of one kind of acid rain: SO3 (g) + H2O (l) → H2SO4 (aq)
2. Decomposition: AB → A + B
In a decomposition reaction, a compound is broken down into two or more substances. In general,
decomposition reactions occur when a solid compound is heated. This type of reaction almost always
produces a gas. The following are examples of decomposition reactions:
Heating mercury (II) oxide produces oxygen gas: 2HgO (s) → 2Hg (l) + O2 (g).
Leaving the cap off the carbonated soft drink bottle allows the carbonic acid to release
carbon dioxide: H2CO3 (aq) → H2O (l) + CO2 (g).
3. Single Replacement: A + BC → AC + B
In this type of reaction, a more “active metal” displaces another element in solution. These reactions
can be further classified as a solid metal reacting with a) a metal ion solution, b) an acid solution, or c)
water. An example of each is provided below:
a) When a solid metal reacts with a metal solution, the solid metal’s ions go into solution
while the metal ions originally in solution plate out onto the surface of solid metal
e.g., Zn (s) + CuSO4 (aq) → ZnSO4 (aq) + Cu (s).
b) When a solid metal reacts with an acid, the metal replaces hydrogen in the acid to produce
hydrogen gas while the metal ion goes into solution with the anion from the acid
e.g., Mg (s) + 2HCl (aq) → H2 (g) + MgCl2 (aq).
c) When a solid metal reacts with water, the metal replaces hydrogen in the water to produce
hydrogen gas while the metal ion goes into solution with hydroxide ion
e.g., Ca (s) + 2 H2O (aq) → H2 (g) + Ca(OH)2 (aq).
To predict whether or not a single-replacement reaction will occur, we refer to the Activity Series for
Metals (shown on the next page).
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Activity Series for Metals
Li > K > Ba > Sr > Ca > Na > Mg > Al > Mn > Zn > Fe >
Cd > Co > Ni > Sn > Pb > (H) > Cu > Ag > Hg > Au
If a solid metal is “more active” – i.e., higher on the Activity Series – than the metal ion in solution or
hydrogen for acids, the more active metal will displace the less active ion, so a reaction occurs. The
more active metal goes into solution while the less active ion either plates out for a metal ion or
bubbles out as a gas for hydrogen ion. If the solid metal is less active than the metal ion or hydrogen
in the compound, then no reaction occurs. Only six metals (Li, K, Ba, Sr, Ca, and Na) – called “active
metals” – react directly with water to produce hydrogen gas and a metal hydroxide solution. These
active metals are the first six metals in the Activity Series.
4. Double Replacement/Precipitation: AB + CD → AD + CB
These reactions involve the mixing of two aqueous ionic compounds to produce a precipitate, an
insoluble ionic compound. The products of a double-replacement/
precipitation reaction can be predicted by switching the cations of the two compounds and using the
Solubility Rules (see below) to determine if the compounds produced are soluble or insoluble.
Solubility Rules for Ionic Compounds in Water
The compound is SOLUBLE if it has:
1. Li+, Na+, K+, or NH4
+ ion (ALWAYS!)
2. C2H3O2
–, NO3
–, ClO4
–
3. Cl–, Br–, or I–, except compounds
with Ag+, Pb+2, Cu2+, and Hg2
+2 are
insoluble
4. SO4
2- except compounds with
Ag2SO4, CaSO4, SrSO4, BaSO4,
PbSO4, and Hg2SO4 are insoluble
The compound is INSOLUBLE if it has:
5. CO3
2–, CrO4
2–, PO4
3–, except compounds
with Li+, Na+, K+, NH4
+ are soluble
6. S2–, except compounds with Li+, Na+, K+,
NH4
+, Ca+2, Sr+2, Ba+2 are soluble
7. Hydroxide ion, OH–, except compounds
with Li+, Na+, K+, NH4
+ are soluble
Soluble ionic compounds will dissolve in water, so their physical states are indicated as aqueous, (aq),
while insoluble ionic compounds will not dissolve in water, so their physical states are indicated as
solid, (s). For a precipitation reaction to occur, at least one of the products must be insoluble; if both
products are soluble, then no reaction occurs. The presence of a precipitate is observed in the lab as a
cloudy mixture that results when two solutions are mixed. The following is an example of a doublereplacement/
precipitation reaction: Pb(NO3) 2 (aq) + K2CrO4 (aq) → 2KNO3 (aq) + PbCrO4 (s)
5. Acid-Base Neutralization: HX (aq) + YOH (aq) → H2O (l) + HX (aq)
These reactions occur between an acid and a base. In general, acids are compounds that produce
hydrogen ions (H+), also called protons, when dissolved in water. The chemical formulas for acids are
most often given with the H’s at the beginning, so acids are usually easy to recognize. A few common
acids are hydrochloric acid, HCl(aq), nitric acid, HNO3(aq), and sulfuric acid, H2SO4(aq). Bases are
compounds that produce hydroxide ions (OH–) when dissolved in water. A few common bases are
sodium hydroxide, NaOH, potassium hydroxide, KOH, calcium hydroxide, Ca(OH)2, and barium
hydroxide, Ba(OH)2. Other types of bases contain carbonate ion, CO3
–2, and hydrogen carbonate (or
bicarbonate) ion, HCO3
–.
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The two types of acid-base neutralization reactions involve a) an acid reacting with a base (containing
the hydroxide ion (OH–)) to produce water and a salt (an ionic compound) or b) an acid reacting with a
base containing carbonate (CO3
–2) or hydrogen carbonate ion (HCO3
–) to produce water, carbon
dioxide gas, and a salt. An example of each is provided below:
a) When an acid reacts with a base containing hydroxide ion (OH–) to produce water and a
salt, the hydrogens from the acid combine with the hydroxide from the base to form water while
the salt is formed by combining the cation from the base with the anion from the acid. The
following is an example of this type of reaction: HCl (aq) + NaOH (aq) → H2O (l) + NaCl
(aq).
b) When an acid reacts with a base containing carbonate (CO3
–2) or hydrogen carbonate
ion (HCO3
–) to produce water, carbon dioxide gas, and a salt, the hydrogens from the acid
combine with the carbonate or hydrogen carbonate from the base to form water and carbon
dioxide gas while the salt is formed by combining the cation from the base combining with the
anion from the acid—e.g., HCl (aq) + NaHCO3 (aq) → H2O (l) + CO2 (g) + NaCl (aq).
6. Combustion Reactions: Hydrocarbon (CxHy) + O2 (g) → CO2 (g) + H2O (g)
In a combustion reaction, a hydrocarbon (composed of C and H) or a hydrocarbon derivative
(composed of C, H, and O) is burned in oxygen to produce carbon dioxide gas and steam. One
example is the combustion of methane (natural gas): CH4 (g) + 2 O2 (g) → CO2 (g) + 2 H2O (g)
**Lab Notebook**
Record observations for all of the chemical reactions carried out during the lab in your lab book.
These observations should include:
• observations of the reactant(s) before the reaction
• observations of the reaction mixture during the reaction
• observations of the product(s) after the reaction.
Your observations of a material should contain the color, clarity and state of matter, plus any
useful descriptions of the material (for example, a sample of magnesium might be described as a
smooth, shiny, silver, opaque solid).
Your observations of the reaction in progress should include anything of potential interest, such as
“the color changed from green to blue”, “a pungent odor is present now”, “the test tube is getting
warmer” or “bubbles are forming on the surface of the magnesium”.
Procedure:
Safety and waste disposal directions are listed with each procedure.
General Directions:
1. Carry out the reactions using the approximate quantities of reagents indicated. Unless otherwise
stated, use test tubes. To estimate 2 mL, measure 2 mL of water in a graduated cylinder and pour it
into a test tube. Save this test tube for comparison.
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2. When combining solutions in a test tube, tap the tube a few times or use the Vortex mixer to ensure
that the solutions have mixed completely.
3. To heat a solid in a test tube, position the test tube holder near the top of the test tube, and hold the
test tube in a slanted position so that the opening of the test tube is pointed away from people.
Keep the bottom of the test tube in the hottest part of the burner, but continuously move it back and
forth over the flame to avoid “hot spots” (overheating one part of the test tube).
4. There are different concentrations of the HCl and NaOH used in this laboratory session. Check
labels carefully for the proper chemical and concentration!
A. Acids and Bases.
CAUTION! NaOH and HCl can damage skin, eyes and clothing on contact. Rinse off any spills
immediately with plenty of water for 10 minutes. In the event of a spill in the laboratory, notify
your instructor immediately.
Place one piece each of red litmus paper and blue litmus paper on a watch glass, leaving a 1-inch
space between them. Place a drop of 0.1M HCl(aq) on each piece of litmus paper using a stirring rod
and record your observations. Then place a drop of 0.1M NaOH(aq) on each piece of litmus paper and
record your observations. Place a drop of deionized water on each piece of litmus paper and record
your observations.
Red litmus paper Blue litmus paper
Before reaction
Reaction with 0.1M HCl
Reaction with 0.1M NaOH
Reaction with H2O
B. Combination Reactions
1. Heat a piece of copper wire strongly in the Bunsen burner flame (using crucible tongs) until a
change in appearance is noted. Record any changes in the appearance of the copper wire in your
lab report. Place the cooled wire in the regular trash.
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CAUTION: Do not look directly at the Mg ribbon as it burns, or you may damage your eyes.
1. Hold a strip of magnesium ribbon in the burner flame (using crucible tongs).
2. Scrape the ash away from any unreacted Mg metal and place only the ash in a watch glass. Add a
few drops of distilled H2O. Carefully crush and stir the ash/water mixture with a stirring rod.
Place one drop of the solution on blue litmus paper and another drop on red litmus paper.
3. Dispose of the wet ash and any unreacted Mg in the waste jar in the hood. Rinse off the pieces of
litmus paper with water, then dispose of them in the regular trash.
Copper Metal Magnesium metal
Before heating
During heating
After heating
Red litmus Blue litmus
Magnesium ash solution
Make sure you conclude whether the ash is acidic or basic.
C. Decomposition Reactions
1. Place approximately half a spatula full (roughly pea-sized) of copper(II) carbonate in a dry test
tube. If you do not have a clean, dry test tube, ask your instructor for one. Do not try to dry
a test tube during the laboratory period. Observe the color of the sample. Using a test tube
clamp, heat the test tube over a Bunsen burner until you notice a color change (approximately 30
seconds – 1 minute). Be sure to constantly move the test tube to avoid overheating the glassware!
Cool the test tube in an empty beaker. Record the color of the solid sample after heating. When
cool, dispose of the contents in the waste jar in the hood.
Copper (II) carbonate
Before heating
During heating
After heating
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D. Single-Replacement Reactions
CAUTION: AgNO3 will stain skin and clothes!
1. Place a piece of copper wire in a test tube with enough 1M AgNO3 to cover it. Allow the test tube
to stand for 5-10 minutes. Note changes in the appearance of both the wire and the solution.
Dispose of the contents of the test tube in the waste jar in the hood.
Copper Metal AgNO3 solution
Before reaction
During reaction
After reaction
CAUTION: 3 M HCl(aq) can damage skin and clothing on contact. Rinse any spills on skin
immediately with plenty of water for 10 minutes. Neutralize all spills on the lab bench with water or
NaHCO3 solution, and rinse your hands thoroughly.
2. Place a small piece of zinc metal in a test tube containing 2 mL of 3 M HCl, and record your
observations. Dispose of the contents of the test tube in the waste jar in the hood.
Zinc Metal HCl solution
Before reaction
During reaction
After reaction
E. Double Replacement/Precipitation Reactions
CAUTION: AgNO3 will stain skin and clothing! Pb containing compounds are toxic and should not
be ingested. HCl, HNO3 and NaOH are corrosive and can cause chemical burns and damage
clothing. Any hazardous chemical spilled on skin must be rinsed off with plenty of water for 10-
15 minutes. If any spills occur in the laboratory, notify your instructor immediately.
Use the 0.10 M solutions of each of these reagents (except Na2SO4 which is 1M, HNO3 which is 3M
and AgNO3 which is 1M). You will obtain solutions of AgNO3(aq), NaCl(aq), Ba(NO3)2(aq),
HNO3(aq) and Pb(NO3)2(aq). To EACH of them you will add solutions of NaNO3(aq), NaCl(aq),
Na2SO4(aq), NaOH(aq), KI(aq), and saturated Na2CO3(aq).
First, record observations of each solution before the solutions are mixed.
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You will mix the pairs of chemicals and observe the reactions between them, watching in particular for
the appearance of a precipitate. All observations should be recorded in your data table in your
notebook. If no change occurs, write “NR” for “No reaction”. If a precipitate appears or if the
solution changes in any other way, record your observations of the change.
Here is an example data table for different reagents. Your data table will be much larger (consider
using the “landscape” orientation of the notebook, because this table will be wider than it is tall).
KNO3(aq) KI(aq) KOH(aq)
CuCl2(aq) NR A brown, opaque solid
is present in a clear,
dark purple solution.
A translucent, blue gel-like
ppt formed immediately.
Pb(NO3)2(aq) NR An opaque yellow ppt
formed in the clear
colorless solution.
The solution turned cloudy
white. Slowly a white
opaque ppt settled in the
clear, colorless solution.
KCl(aq) NR NR NR
Procedure
NOTE: All waste for this part of the experiment should be poured into the labeled waste
containers in the hood and the test tubes rinsed with a minimum amount of water, which
should also be placed into the waste container. DO NOT dispose of any solutions or solids
down the drain.
1. Wash your well-plate thoroughly with soap and water, then rinse it completely with deionized
water. A dirty well-plate can give incorrect results.
2. Place 5 drops of each aqueous solution in the correct wells based on the table you constructed for
your observations.
3. For each of the following combinations, mix 10 drops of each solution in a clean test tube, so any
reactions that take place can be observed on a larger scale:
• 0.10 M Ba(NO3)2(aq) and 0.10 M NaOH(aq)
• 3 M HNO3(aq) and 3 M NaOH(aq)
• 3 M HNO3(aq) and Saturated Na2CO3(aq)
4. Have your lab instructor sign off on your Double Replacement/Precipitation reaction observation table.
5. In the discussion section of your laboratory notebook, write a balanced chemical equation and clearly
identify the solid product for any precipitation reactions that you observe.
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F. Combustion Reaction
1. Place about 10-15 drops of 2-propanol (isopropyl alcohol, C3H7OH) in a small evaporating dish.
2. Ignite a wooden splint in the Bunsen burner and use the wooden splint to light the alcohol.
2-propanol
Before ignition
During combustion
After combustion
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Chemical Reactions: Introduction to Reaction Types: Lab Report
Name: ________________________
Partner(s): _____________________
Section Number: ________________
Word Equations and Balanced Chemical Equations
Translate each of the following word equations into a balanced chemical reactions by writing the
correct chemical formulas (including physical states) for the reactants and products. Make sure
to balance each equation.
Example: Aluminum metal reacts with oxygen.
aluminum metal + oxygen gas →
4 Al (s) + 3 O2 (g) → 2 Al2O3 (s)
(Rewrite Entire Balanced Reaction)
A. Combination Reactions
1. Copper reacts with oxygen.
copper metal + oxygen gas →
2. Magnesium metal reacts with oxygen.
magnesium metal + oxygen gas →
3. Magnesium oxide (ash) reacts with water.
magnesium oxide + water →
B. Decomposition Reactions
1. Water decomposes.
water →
2. Copper(II) carbonate decomposes.
copper(II) carbonate →
Δ
Δ
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12
C. Single-Replacement Reactions
1. Copper reacts with silver nitrate to form silver metal and copper(II) nitrate.
copper metal + silver nitrate →
2. Zinc metal reacts with hydrochloric acid to produce zinc chloride and hydrogen.
zinc metal + hydrochloric acid →
D. Double Replacement (precipitation) and Acid Base Reactions
Refer to your data table for the following selected sets of reactants and fill in the following blanks and
beaker drawings. If there is no net ionic reaction because all the ions are spectators still complete the
molecular reaction, the ionic reaction, and the beaker drawings, then put NR only for the net ionic
reaction. An example NOT from this experiment is presented first.
Example: calcium acetate and ammonium sulfate Reaction type: ____precipitation________
Molecular: Ca(CH3COO)2(aq) + (NH4)2SO4(aq) → CaSO4(s) + 2 NH4CH3COO(aq)________________________________
Ionic: Ca2+(aq) + 2 CH3COO-(aq) + 2 NH4
+(aq) + SO4
2-(aq) → CaSO4(s) + 2 CH3COO-(aq) + 2 NH4
+(aq)____________
Net Ionic: Ca2+(aq) + SO4
2-(aq) → CaSO4(s)______________________________________________________________
+
→
1. lead (II) nitrate and potassium iodide Reaction type: _______________
Molecular: ________________________________________________________________________
Ionic: ____________________________________________________________________________
Net Ionic: ________________________________________________________________________
+
→
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2. nitric acid and sodium hydroxide Reaction type: _______________
Molecular: ________________________________________________________________________
Ionic: ____________________________________________________________________________
Net Ionic: ________________________________________________________________________
+
→
3. barium nitrate and sodium sulfate Reaction type: _______________
Molecular: ________________________________________________________________________
Ionic: ____________________________________________________________________________
Net Ionic: ________________________________________________________________________
+
→
F. Combustion Reactions
1. Isopropyl alcohol (C3H8O) undergoes combustion. Write the balanced chemical equation including
physical states:
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Balancing and Categorizing Chemical Equations:
Balance each of the 12 chemical equations given below, and identify each as one of the six types listed
below.
Combination reaction (C)
Decomposition reaction (D)
Single-Replacement reaction (SR)
Double-Replacement/Precipitation reaction (DR)
Acid-Base Neutralization reaction (N)
Combustion reaction (B)
TYPE
_____ 1. _____ Al (s) + _____ NiCl2 (aq) → _____ Ni (s) + _____ AlCl3 (aq)
_____ 2. _____ Ba(OH)2 (s) + ____ FeCl3 (aq) → ____ BaCl2 (aq) + ____ Fe(OH)3 (s)
_____ 3. _____ C4H10 (l) + _____ O2 (g) → _____ CO2 (g) + _____ H2O (g)
_____ 4. _____ KClO3 (l) → _____ KCl (l) + _____ O2 (g)
_____ 5. _____ Al (s) + _____ I2 (s) → _____ AlI3 (s)
_____ 6. _____ H2SO4 (aq) + _____ Mg(OH)2 (s) →_____ H2O (l) + _____ MgSO4 (aq)
_____ 7. _____ CH3OH (l) + _____ O2 (g) → _____ CO2 (g) + _____ H2O (g)
_____ 8. _____ Ca (s) + _____ O2 (g) → _____ CaO (s)
_____ 9. _____ Mg (s) + _____ CO2 (g) → _____ MgO (s) + _____ C (s)
____ 10. _____ Na3PO4 (aq) + ____ MgCl2 (aq) → ____ Mg3(PO4)2 (s) + ___ NaCl (aq)
____ 11. _____ HgO (s) → _____ Hg (l) + _____ O2 (g)
____ 12. _____ H3PO4 (aq) + ____ NaOH (aq) → _____ H2O (l) + ____ Na3PO4 (aq)
Δ
Δ
Δ
Δ
