Chemical reactions Laboratory task

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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).

GCC CHM 151LL: Chemical Reactions: Introduction to Reaction Types © GCC, 2014 page 2 of 12

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 double- replacement/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–.

GCC CHM 151LL: Chemical Reactions: Introduction to Reaction Types © GCC, 2014 page 3 of 12

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.

GCC CHM 151LL: Chemical Reactions: Introduction to Reaction Types © GCC, 2014 page 4 of 12

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.

GCC CHM 151LL: Chemical Reactions: Introduction to Reaction Types © GCC, 2014 page 5 of 12

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

GCC CHM 151LL: Chemical Reactions: Introduction to Reaction Types © GCC, 2014 page 6 of 12

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.

GCC CHM 151LL: Chemical Reactions: Introduction to Reaction Types © GCC, 2014 page 7 of 12

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.

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

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 →

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: ________________________________________________________________________

→

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:

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)