AP ChemistryUnit 3 – Special Reactions

This unit focuses on mathematical tools used in studying chemical reactions. These tools make chemistry easier and more manageable, since they allow you to use the power of math to reduce substances and reactions to the simplest terms possible. Molecules have mass just as objects in our everyday world do. We'll find out how the mass of a molecule is determined, discuss the atomic mass unit, the fundamental unit in atomic, molecular, and formula masses. We'll also consider the mole, the chemist's counting unit for atoms and molecules, and use moles to carry out quantitative chemistry. Explore different ways to express quantities and to relate them with chemical equations. In particular, the skill of predicting products of reactions and writing net ionic equations will be further developed. Our focus will be on acid-base reactions and redox reactions.

Objectives:3.1 Measure and determine the concentration of a solution3.2 Distinguish between various types of acids and bases.3.3 Use titration to determine the neutralization point of a chemical reaction3.4 Identify oxidizers, reducers, and balance redox equations.3.5 Predict the products of redox reactions.

Skills to Master:1. Describe concentration in terms of molarity. 2. Make a solution of a given molarity. 3. Solve problems involving both stoichiometry and solution molarity. 4. Describe the use of gravimetric analysis to determine the amount of a substance

dissolved in a solution. 5. Describe the use of volumetric analysis and titrations to determine the amount of a

substance dissolved in a solution. 6. Discriminate between the equivalence point and the end point of a titration. 7. Use laboratory procedures which are important techniques used to separate

mixtures into their component parts. 8. Use acid-base titration and redox titration to determine amounts in chemical

reactions.9. Gather and/or use experimental measurements to determine the mole ratio of

reactants where the formulas of the products are not known.

Chapter 3 Problem Set: P. 111-118; 11, 13, 17, 19, 60Chapter 4 Problem Set: P. 158-164; 4, 13, 15, 18, 29, 38, 39, 49, 70, 72, 80, 84, 88, 108

Podcast 3.1: Working with SolutionsParts of Solutions

• Solution- • Solute- • Solvent- • Soluble- • Miscible-

Aqueous solutions• Dissolved in water.

• Water is a good solvent because _____________________________• The oxygen atoms have a partial ______________charge.• The hydrogen atoms have a partial _______________charge.• The bond angle is ______º.

Hydration• The process of _____________________________________________ • Ions have charges and attract the opposite charges on the water molecules. • Solubility -

• Usually g/100 mL or M• Varies greatly, but if they do dissolve the ions are separated (hydrated), thus

free to move around• Water can also dissolve non-ionic compounds if they have ____________ bonds.

Electrolytes• Electricity is moving ______________• Only ions that are ________________ can move.• Compounds that conduct an electrical current in an aqueous solution or molten

compound are called ______Solutions are classified three ways.

1. ______________________- completely dissociate (break apart into ions), many ions- Conduct well.

2. ______________________- only a certain fraction dissolves into ions, few ions -Conduct electricity slightly.

3. ______________________- Don’t break apart, no ions- exist as molecules, don’t conduct electricity

**Solubility Rules assist in determining if something will be an electrolyte or notTypes of Electrolyte Solutions

• Acids- form __________ when dissolved.• Strong acids fall apart completely• Memorize the 8 strong acids: (all others are weak)

• Weak acids- don’t dissociate completely.• Bases - form ________ when dissolved.• Strong bases completely break apart (All others are weak bases, including NH3)

Measuring Solutions• Concentration- how much is dissolved.• Molarity =

Calculate the molarity of a solution with 34.6 g of NaCl dissolved in 125 mL of solution.

Concentrations of Solutions: Molarity (M)• Ion Concentration

6.0 M HCl =

6.0 M Na2SO4 =

Molarity can be used as a conversion factor in Dimensional Analysis (T-tables)Example 1: How many moles of HNO3 are in a 2.0 L solution of 0.200 M HNO3?

Example 2: What volume of solution is necessary to provide 2.0 mol of HNO3? We have 0.30 M HNO3.Convert from Molarity to MassExample 1: How many grams of solute are present in 50.0 mL of 0.360 M K2Cr2O7?

Example 2: If 4.28 g of (NH4)2SO4 is dissolved in enough water to form 300 mL of solution, what is the molarity of the solution?

Example 3: How many milliliters of 0.240 M CuSO4 contain 2.25 g of solute?

Solution PreparationExample 4: How many grams of HCl would be required to make 50.0 mL of a 2.7 M solution?

Example 5: What would the concentration be if you used 27g of CaCl2 to make 500. mL of solution? What is the concentration of each ion?

Dilutions• M1V1 = M2V2

• MconcVconc = MdilVdil

• Shows that the number of moles in each solution is equal

Example 6: A lab requires 500 mL of a 3M HCl solution. How many mL of 12M HCl is needed to make this solution?

Solution Dilution

Example 7: If 25 mL of water is added to 125 mL of a 0.15 M NaOH solution, what will the molarity of the diluted solution be?

ElectrolytesElectrolytes are substances that break up (dissociate or ionize) in water to produce ions. These ions are capable of conducting an electric current. Generally, electrolytes consist of acids, bases, and salts (ionic compounds). Nonelectrolytes are usually covalent compounds, with the exception of acids.

Classify the following compounds as either an electrolyte or a nonelectrolyte.

Compound Electrolyte Nonelectrolyte

1. NaCl2. CH3OH (methyl

alcohol)3. C3H12(OH)3 (glycerol)4. HCl5. C6H12O6 (glucose)6. NaOH7. C2H5OH (ethyl

alcohol)8. CH3COOH (acetic

acid)9. NH4OH10.H2SO4

Molarity (M)

Molarity= molesof soluteliters of solution

Solve the problems below.1. What is the molarity of a solution in which 58 g of NaCl are dissolved in 1.0 L of

solution?

2. What is the molarity of a solution in which 10.0 g of AgNO3 is dissolvedin 500. mL

of solution?

3. How many grams of KNO3 should be used to prepare 2.00 L of a 0.500 M solution?

4. To what volume should 5.0 g of KCl be diluted in order to prepare a 0.25 M

solution?

5. How many grams of CuSO4•5H2O are needed to prepare 100. mL of a 0.10 M

solution?

Molarity by DilutionAcids are usually acquired from chemical supply houses in concentrated form. These acids are diluted to the desired concentration by adding water. Since moles of acid before dilution = moles of acid after dilution, and moles of acid = M x V then…

M 1× V 1=M 2×V 2

Solve the following problems.

1. How much concentrated 18 M sulfuric acid is needed to prepare 250 mL of a 6.0 M

solution?

2. How much concentrated 12 M hydrochloric acid is needed to prepare 100. mL of a

2.0 M solution?

3. To what volume should 25 mL of 15M nitric acid be diluted to prepare a 3.0 M

solution?

4. To how much water should 50 mL of 12 M hydrochloric acid be added to produce a

4.0 M solution?

5. To how much water should 100. mL of 18 M sulfuric acid be added to prepare a 1.5

M solution?

Podcast 3.2: Acid-Base ReactionsAcids

• Acid – substances that ionize in aqueous solutions to form _________, which increases the _________ concentration in solution

• H+ = _________ = hydronium ion• Monoprotic acids (like HCl) will completely ionize• Diprotic acids (H2SO4, H2CO3) ionize in two steps

H2SO4 → HSO4

- → • Triprotic acids (H3PO4, H3PO3) ionize in three similar steps

Bases• Bases – substances that accept or react with _____ ions. They also produce _____

when they dissolve in H2O • Ex. NaOH à

Ca(OH)2 à NH3 + H+ à

• Pay attention to stoichiometry!Acid-Base Reactions

• Neutralization reactions will always produce __________ and a _________ as products. Write the net ionic equation for the following reaction

HCl + NaOH → H2O + NaCl

• ____________ – any ionic compound whose cation comes from a base and whose anion comes from an acid

Titration Demo

Strong Acids and Strong Bases are Strong Electrolytes and Completely Dissociate in Water.

Strong Acids Strong Bases

Acid-Base Theories: Acids and Bases have been described by chemists for more than 300 years; however, chemists still disagree on the chemical processes that actually take place.Arrhenius Acids and Bases

Brǿnsted-Lowry Acids and Bases:

Lewis Acids and Bases:

Amphoteric Substance• Water • A substance like this is referred to as “________________” because it act as either an

acid or a base• Other amphoteric substances may include metal or metalloid oxides, amino acids

and proteins, and ammoniaTitrations: Lab technique used to determine the ___________________ of a particular solute in solution.

• Standard Solution – • Titrations can be conducted using acid-base, precipitation, or redox reactions.• For a titration to work effectively, stoichiometrically equivalent quantities of the

reactants must be brought together to reach the _____________________ ________________

__________________ are used to determine when the equivalence point has been reached. • Color change = ________ point (as close to equivalence point as possible)

pH Ranges of Acid-Base Indicators

Example Problem 1: What is the molarity of a NaOH solution if 48.0 mL is needed to neutralize 35.0 mL of 0.144 M H2SO4? 2NaOH + H2SO4 → 2H2O + Na2SO4

Example 2: The quantity of Cl- in a water supply is determined by titrating the sample with Ag+. How many grams of Cl- are in a sample of water if 20.2 mL of 0.100 M Ag+ is needed to react with all chloride in the sample?Ag+ + Cl- → AgCl

Paper Chromatography LabIntroduction

The fact that different substances have different solubilities in a given solvent can be used in several ways to effect a separation of substances from mixtures in which they are present. One widely used technique, which depends on solubility differences, is chromatography.

In the chromatographic experiment a mixture is deposited on some solid adsorbing substance, which might consist of a strip of filter paper, a thin layer of silica gel on a piece of glass, some finely divided charcoal packed loosely in a glass tube, or even some microscopic glass beads coated very thinly with a suitable adsorbing substance and contained in a piece of copper tubing. The components of a mixture pass down and are adsorbed on the solid to varying degrees, depending on the nature of the component, the nature of the adsorbent, and the temperature. A solvent is then caused to flow through the adsorbent solid under applied or gravitational pressure or by capillary motion. As the solvent passes the deposited sample, the various components tend, to varying extents, to be dissolved and swept along the solid. The rate at which a component will move along the solid depends on its relative tendency to be dissolved in the solvent and adsorbed on the solid. The net effect is that the components separate from each other and move along as rather diffuse zones as the solvent passes slowly through the solid. With the proper choice of solvent and adsorbant, it is possible to resolve many complex mixtures by this procedure. If necessary, we can usually recover a given component, removing that part of the solid from the system, and eluting the desired component with a suitable solvent.

The name given to a particular kind of chromatography depends upon the manner in which the experiment is conducted. Thus, we have column, thin-layer, paper, and vapor chromatography, all in very common use. Chromatography in its many possible variations offers the chemist one of the best methods, if not the best method, for resolving a mixture into pure substances, regardless of whether that mixture consists of a gas, a volatile liquid, or a group of nonvolatile, relatively unstable, complex organic compounds.

In this experiment, we shall use paper chromatography to resolve a mixture of substances known as acid-base indicators. These materials are typically brilliant in color, with the colors depending on the acidity of a system in which they are present. A sample containing a few micrograms of the indicator is placed near one end of a strip of filter paper. That end of the paper is then immersed vertically in a solvent. As the solvent rises up the paper by capillary action it tends to carry the sample along with it, to a degree that depends on the solubility of the sample in the solvent and its tendency to adsorb on the paper. When the solvent has risen a distance of L centimeters, the solute, now spread into a somewhat diffuse zone or band, will have risen a smaller distance, say D centimeters. It is found that D/L is, for a given substance under specified conditions, a constant independent of the relative amount of that substance or other substances present. D/L is called the Rf value for that substance under the experimental conditions:

R f=DL

=distance solute movesdistance solvent moves

The Rf value is characteristic of the substance in a given chromatography experiment, and can be used to test for the presence of a particular substance in a mixture of substances with different Rf values.

The first part of the experiment will involve the determination of the Rf values for five common acid-base indicators. These substances have colors that will allow you to establish the positions of their bands at the conclusion of the experiment. When you have found the Rf for each substance by studying it by itself, you will use these Rf values to analyze an unknown mixture.

Experimental Procedure1. Take a clean dry beaker and a clean dry test tube to the stockroom and obtain six paper strips and a

sample of your unknown. Handle the strips carefully; whenever you need to work with them, handle them by their edges, because their surfaces can very easily be contaminated by your fingers.

2. Place the strips on a clean dry sheet of paper and make a pencil mark about ¾ inch from one end of each strip.

3. Put two or three drops of the following indicators into separate, clean, dry, small test tubes:Bromothymol blue PhenolphthaleinCongo red Phenol redMethyl orangeFor an applicator use a fine capillary tube, which will be furnished to you by your instructor. Test the application procedure by dipping the applicator into one of the colored solutions and touching it momentarily to a round test piece of filter paper. The liquid from the applicator should form a spot no larger than ¼ inch in diameter. Test the procedure several times.

4. Clean the applicator by dipping it in a few mL of acetone and blowing air through it to dry it. Dip it in one of the indicator solutions and put a ¼” spot on the pencil line on one of the strips. Label the strip at point X so that you can tell which indicator you applied. Clean the applicator and repeat the procedure on the other strips for each of the other indicators and the unknown. Each strip will look like that in figure 1 on completion.

5. Get 50 mL of eluting solvent from the supply in the hood. This solution is made by adjusting the pH of a convenient organic alcohol (ethanol, butanol or hexanol) with concentrated NH3 solution to about pH 10. Pour this solution into three dry 250 mL Erlenmeyer flasks. These will serve as developing chambers in the experiment. Stopper each flask promptly to avoid evaporation of the NH3. The final setup in the flask will resemble that in Figure 2.

6. When you are sure the sample spots are dry, place two of the strips opposite each other on the side of a cork, and place the cork and the strips in the flask so that the ends of the strips, but not the sample spots, are in the solvent. In the same way, place two strips in the second flask and two in the third.

7. Let solvent rise on the strips for 45 to 60 minutes, or until the solvent front has moved at least 3 ½ inches above the pencil lines. Remove the strips from the beakers and put them on the sheet of paper used earlier. Draw a pencil line along the solvent front on each of the strips and let the strips dry for several minutes.

8. When the strips have dried (you may need to hold them in the warm air over a piece of asbestos screen held over a small Bunsen flame), hold each strip over the open mouth of a bottle of concentrated NH3, and note the color of the band associated with each indicator when it is in the alkaline vapor.

9. Repeat the procedure, holding the strips over an open bottle of concentrated HCl, recording the color of each indicator in the presence of an acid. When you are sure that you know the position of the band to be associated with each indicator; measure the distance from the center of the band to the point where the indicator was applied. For each indicator, also measure the distance from the solvent front to the point of application. Calculate Rf values for each indicator.

10. On the strip containing the unknown, measure the Rf values and the colors for each band in the presence of both basic and acidic vapors, and identify those indicators which are in the unknown. Some of the indicators, particularly phenolphthalein, have colors that fade very rapidly, so you should be careful to check for the presence of each indicator separately. If any bands overlap, the different indicator colors in the presence of acidic and basic vapors should enable you to make a positive identification.

Data and Calculations:Distance solvent moved (cm)

Distance sample moved (cm)

Rf Color in NH3 vapor

Color in HCl vapor

Bromthymol

blue

Congo Red

Methyl Orange

Phenol Red

Phenolphthalein

Unknown #1

Unknown #2

Unknown #3

Unknown #4

Unknown #5

Unknown #6

Composition of unknown(s) ____________________________________________

Unknown # __________________________

Bronsted-Lowry Acids and Bases

According to the Bronsted-Lowry Theory, an acid is a proton (H+) donor, and a base is a proton acceptor.

Label the Bronsted-Lowry acids and bases in the following reactions and show the direction of proton transfer. H+

H+

Example: H2O + Cl - à OH - + HCl

Acid Base Base Acid

1. H2O + H2O

2. H2SO4 + OH_

3. HSO4 - + H2O

4. OH - + H3O +

5. NH3 + H2O

H +

Example: HCl + OH - à Cl - + H2O

The HCl acts as an acid, the OH- as a base. This reaction is reversible in that the H2O can give back the proton to the Cl-.

Conjugate Acid-Base Pairs

It was shown that after an acid has given up its proton, it is capable of getting back that proton and activing as a base. Conjugate base is what is left after an acid gives up a proton. The stronger the acid, the weaker the conjugate base. The weaker the acid, the stronger the conjugate base.

Fill in the blanks in the table below.

ACID BASE EQUATION

1. H2SO4 HSO4 - H2SO4

2. H3PO4

3. F -

4. NO3-

5. H2PO4 -

6. H2O

7. SO4 2-

8. HPO4 2-

9. NH4 +

10

. H2O

11. Which is a stronger base, HSO4 – or H2PO4

- ? Why?

12. Which is a weaker base, Cl – or NO2 - ? Why?

Podcast 3.3: Acid-Base Reaction StoichiometryFinding the Molarity of an AcidExample 1: 20.0mL of 0.25M NaOH is titrated with 23.2mL of HC2H3O2 to the phenophtalien endpoint. What is the concentration of the acetic acid?

Example 2: 10.0mL of 0.50M NaOH is titrated with 23.2mL of H3PO4 to the phenophtalien endpoint. What is the concentration of the phosphoric acid?

Calculating the Molar Mass of an AcidExample 3: 0.523 grams of an unknown monoprotic acid is titrated to the phenphtalien endpoint with 22.5mL of 0.103M NaOH. What is the molar mass of the acid?

Acid-Base TitrationTo determine the concentration of an acid (or base), we can react it with a base (or acid) of known concentration until it is completely neutralized. This point of exact neutralization, known as the endpoint, is noted by the change in color of the indicator. We use the following equation:

Solve the problems below.

1. A 25.0 mL sample of HCl was titrated to the endpoint with 15.0 mL of 2.0 M NaOH. What was the concentration of the HCl?Balanced equation:

2. A 10.0 mL sample of H2SO4 was exactly neutralized by 13.5 mL of 1.0 M KOH. What is the molarity of the H 2SO4?Balanced equation:

3. How much 1.5 M NaOH is necessary to exactly neutralize 20.0 mL of 2.5 M H3PO4?Balanced equation:

4. How much of 0.5 M HNO3 is necessary to titrate 25.0 mL of 0.05 M Ca(OH)2 solution to the endpoint?Balanced equation:

5. What is the molarity of a NaOH solution if 15.0 mL is exactly neutralized by 7.5 mL of a 0.02 M HC2H3O2 solution?Balanced equation:

MA x VA = Mb x VB where M = molarity V = volume

Podcast 3.4 Redox Reactions: Oxidizers, Reducers, and Assigning Oxidation NumbersOxidation and Reduction

• __________________ – the combination of an element with oxygen to produce oxides.Ex: C(s) + O2(g) → CO2(g)

CH4(g) + 2O2(g) → CO2(g) + 2H2O (g)• __________________ – the loss of oxygen from a compound.

Ex: 2 Fe2O3(s) + 3C(s) → 4Fe(s) + 3CO2(g)Redox Reactions

• Both processes of oxidation and reduction occur simultaneously. You can’t have one without the other.

• Thus they are called Oxidation-Reduction Reactions, or Redox Reactions• So what does that mean for reactions not involving oxygen? Today, chemists use a

broad definition that deals with the transfer of electrons• Oxidation = __________ of electrons• Reduction = __________ of electrons• LEO say ________• Substances that lose e- are ___________• Substances that gain e- are ___________• Substances that are oxidized are also called the ________________________

(Substance doing the reducing)• Substances that are reduced are also called the ________________________ (Substance

doing the oxidizing)Example 1: Mg + S → MgS

For this reaction we can write two ½ -reactions (or ionization equations)

Mg

S

Oxidation Numbers: used to help keep track of electrons in redox reactionsIn any reaction:

• An INCREASE in oxidation # of a substance indicates _________________• A DECREASE in oxidation # of a substance indicates __________________• If there are no changes in any oxidation #’s then you do NOT have a REDOX

reactionOxidation Rules – Apply IN ORDER

1. Atoms in elemental form will have an oxidation number of zero (Diatomics included.)

2. Monatomic ions will have an oxidation # equal to it’s charge.3. Oxygen has an oxidation # of -2, unless it is a peroxide, then it will be -1.4. Hydrogen has an oxidation # of +1 when it is bonded to nonmetals and -1 when

bonded to metals.5. Halogens are usually -1, unless they are combined with oxygen, then they will be

positive.6. The sum of the oxidation #’s will equal zero for a compound or will equal the value

of the charge for an ion.

Example 1: Determine the oxidation number of each element in Na2SO4. • Is the substance elemental? • Is the substance ionic? • If the substance is ionic, are there any monoatomic ions present? • Which elements have specific rules?

• Which element(s) do(es) not have rules? • Let S = oxidation number of sulfur. • Solve for S using algebra.

Example 2: Determine the oxidation number of each element in K2C2O4.• Is the substance elemental? • Is the substance ionic? • If the substance is ionic, are there any monoatomic ions present? • Which elements have specific rules? • Which element(s) do(es) not have rules?

Solve for C using algebra.

Practice Problems: Determine the oxidation number of each element in the following compounds:

1) Ba(NO3)2

2) NF3

3) (NH4)2SO4

Assigning Oxidation NumbersAssign oxidation numbers to all of the elements in each of the compounds or ions below.

1. HCl 11.H2 SO3

2. KNO3 12.H2SO4

3. OH- 13.BaO2

4. Mg3N2 14.KMnO4

5. KClO3 15.LiH

6. Al(NO3)3 16.MnO2

7. S8 17.OF2

8. H2O2 18.SO3

9. PbO2 19.NH3

10.NaHSO4 20.Na

Oxidation and Reduction Round Table – Get your Oxidation and Reduction Assignment from your teacher.• Pass ONE paper around the table clockwise.• Each team member uses a different color to write the

answer when the paper comes to them. • When the paper comes to you, solve the sample problem

while explaining your reasoning OUT LOUD to your team– The person to your right will be your coach– The person to your left will be your accuracy checker– The person at the diagonal is your encourager

Podcast 3.5: Balancing Redox ReactionsBalancing Redox Reactions in ACIDIC Conditions

Ensure that the electrons lost equal the electrons gained. Utilizes the ½ reactions

Follow these steps:1. Divide equation into ½ reactions, one for oxidation and one for reduction2. Balance each ½ reaction

a. Balance elements other than H and O firstb. Balance O by adding waterc. Balance H by adding H+

d. Balance charge by adding electrons3. Multiply reactions by a whole number so e-’s are equal4. Add two ½ reactions together and cancel out like terms.5. Double check your answer to ensure atoms are balanced.

Example Reaction: MnO4- + C2O4

2- → Mn2+ + CO2

1) MnO4- →

C2O42- →

2a &2b)

2c)

2d)

3)

4)

5) Double Check atoms and charges

Balance the Following Redox Reaction: An acidic solution of potassium dichromate is added to a solution of iron (II) nitrate.

Balancing Redox in BASIC Solution1. Write the ½ - reactions 2. Balance everything except H and O3. Balance O by adding H2O4. Balance H by adding H+

5. Add OH- to both sides to cancel the H+

6. Cancel out any extra water and OH-

7. Balance Charge with electrons (multiply so electrons cancel8. Add the ½ - reactions

Example: A solution of potassium permanganate is mixed with an alkaline solution of sodium sulfite.

Podcast 3.6: Types of Redox ReactionsSimple Redox1. Hydrogen Displacement

Ca(s) + 2H2O(l) --> Ca(OH)2(s) + H2

2. Metal DisplacementZn(s) + CuSO4(aq) ---> ZnSO4(aq) + Cu(s)

3. Halogen DisplacementCl2(g) + KBr(aq) ----> 2KCl(aq) + Br2(l)

4. CombustionCH4(g) + 2O2(g) ---> CO2(g) + 2H2O(g)

Complex Redox

Disportionation: one substance both oxidizes and reducesCl2(g) + 2OH-(aq) ------> OCl-(aq) + Cl-(aq) + H2O(l)

Reactions involving oxoanions such as Cr2O72-

14H+(aq) + Cr2O72- + 6 Fe2+ ---> Cr3+ + 7 H2O + 6 Fe3+

Redox Reaction Prediction

Important Oxidizers Formed MnO4

- (acid solution) Mn+2

MnO4- (basic solution) MnO2

MnO2 (acid solution) Mn+2

Cr2O7-2 (acid) Cr+3

CrO4-2 Cr+3

HNO3, conc NO2

HNO3, dilute NOH2SO4, hot conc SO2

Metallic Ions (Sn4+) Metallous Ions (Sn2+)Free Halogens (Br2) Halide ions (Br-)HClO4 Cl-

H2O2 H2OHalates (IO3

-) Halogens (I2)

Important Reducers Formed in ReactionHalide Ions (Cl-) Halogens (Cl2)Free Metals (Fe) Metal Ions (Fe3+)Metalous Ions (Cu1+) Metallic ions (Cu2+)Nitrite Ions Nitrate IonsSulfite Ions SO4

-2

Free Halogens (dil, basic, sol) Hypohalite ionsFree Halogens (conc, basic sol) Halate ionsC2O4

2- CO2

H2O2 O2

Redox reactions involve the transfer of electrons. The oxidation numbers of at least two elements must change. Single replacement, some combination and some decomposition reactions are redox reactions.To predict the products of a redox reaction, look at the reagents given to see if there is both an oxidizing agent and a reducing agent. When a problem mentions an acidic or basic solution, it is probably is redox.

Predicting Products of Redox Reactions1. Solid copper is added to a dilute nitric acid solution

2. Ethanol has completely burned in air.

3. Sodium metal is added to water.

4. Hydrogen peroxide is added to a solution of iron(II)sulfate.

Memorize This!!!

Put it on Flash Cards

Know it by the Unit 3 Test

5. Al + MnO4- MnO2 + Al(OH)4

-

Podcast 3.7 Redox Stoiochiometry50.0mL of 0.10M KMnO4 is titrated to the endpoint with 20.0 mL of FeSO4. The solution is acidified.

Balance the equation.

What is the concentration of the Fe2+ [Fe2+]?

More Handy Flashcards to Make1. Brown Gas= Nitrogen Dioxide (NO2)2. Black ppt. = Silver Sulfide (Ag2S)3. Green Sol.= Nickel 2+4. Yellow to reddish orange sol.= Iron 2+ or 3+5. Blue or Green sol.= Copper 1+ or 2+6. Yellow sol.=Chromate7. Purple sol.= Permangenate (MnO4-)8. Red flame= Strontium 2+9. Yellow (Orange) flame= Sodium 1+10.Deep red (crimson) flame= Lithium 1+11.Bright yellow ppt. = Lead Iodide (PbI2)12.White ppt.= Silver Chloride (AgCl)13.Pink sol.= Cobalt 3+14.Orange sol.= Dichromate15.Pink sol.=Manganese 2+16.Blue-green flame= Copper 2+17.Red-orange flame=Calcium 2+18.Green flame= Barium 2+19.Violet flame= Potassium 1+20.Dark blue sol.= Cu(NH3)42+

AP Chemistry Lab Report Rubric

Oxidation-Reduction Titration

Name: _______________________________________

Table of ContentsPoints Earned

Points Possible

Includes the title, page numbers, and date of experiment 1Title

Capitalized appropriately, relates to the experiments, underlined at the top of the lab report 1

Abstract Purpose clearly stated. Hypothesis states what you are doing, what you predict

will happen, and why you think that will happen. If…Then…Because

10

Independent variable, dependent variables, constants (at least 3), and the control are identified

Procedure summarized appropriately Results provide an accurate description of the outcome of the experiment Application to theory and the real world are discussed knowledgably and

meaningfully.Data

Organized table that shows the data you have collected during the experiment 3Prelab QuestionsPrelab Questions 7Analysis

Statistical Analysis of Class Data includes mean, standard deviation, and check for outliers.Postlab Questions – Calculations show each step in the determination of oxalic acid concentration.

15

Conclusion Summarize results, analyze errors, and discuss further questions/improvements 15

Discussion/Reflection Discuss what you learned from this experiment and how it relates to what we

are learning in class and applications in the real world (your world). 4

Explain how the theoretical concepts we are learning in class directly apply to the lab experience 3

Total Points 59

Net Ionic EquationsNon-Redox Reactions

Determine the type of reaction so you know what sort of products might form.

Look for special key words: excess, concentrated, dilute, equimolar, etc Dissociate all strong acids, strong bases, and strong electrolytes (look out

for carbonic acid in acidic solution).Redox Reactions *

Watch for single-replacement, composition, disportionation, and combustion so you can tell what type of redox reaction it is.

Assign oxidation numbers to each atom. Assume acidic conditions unless otherwise specified. Write the half-reactions and balance them first.

Work together with a partner using the Rally Coach method to write the Net Ionic Equations for the chemical reactions described below.

Pass ONE paper and ONE pencil/pen back and forth between you and your partner.

When the paper comes to you, solve the sample problem while explaining your reasoning OUT LOUD to your team.

While your partner solves the problemo offer genuine praise and encouragement “yes that’s right” or “good, keep

going”o Offer appropriate coaching “Tip-Tip-Tell”

1. Solid potassium chlorate is strongly heated.

2. Solid silver chloride is added to a solution of excess concentrated hydrochloric acid.

3. A solution of ethanoic (acetic) acid is added to a solution of barium hydroxide.

4. Ammonia gas is bubbled into a solution of hydrofluoric acid.

5. Zinc metal is placed in a solution of copper(II) sulfate.

6. Hydrogen phosphide (phosphine) gas is added to boron trichloride gas.

7. A solution of nickel(II) bromide is added to a solution of potassium hydroxide.

8. Hexane is combusted in air.

9. Solid calcium carbonate is strongly heated.

10.A strip of magnesium metal is placed in a solution of iron(II) chloride.

11.Boron trifluoride gas is mixed with ammonia gas.

12.Excess concentrated hydrochloric acid is added to a solution of nickel(II) nitrate.

13.*A solution of tin (II) chloride is added to an acidified solution of potassium permanganate.

14.Propanal is burned in air.

15.A strip of aluminum foil is placed in liquid bromine.

16.Solid copper(II) sulfide is strongly heated in air.

17.Sulfur dioxide gas is bubbled into distilled water.

18.A drop of potassium thiocyanate solution is added to a solution of iron (III) nitrate.

19.A piece of copper wire is placed in a solution of silver nitrate.

20.Solutions of potassium hydroxide and propanoic acid are mixed.

21.An excess of nitric acid solution is added to a solution of tetraamminecopper (II) sulfate.

22.Solid potassium oxide is added to water.

23.*A solution of iron (II) chloride is added to an acidified solution of sodium dichromate.

24.Chlorine gas is bubbled through a solution of potassium bromide.

25.Solutions of strontium nitrate and sodium sulfate are mixed.

26. Powdered magnesium carbonate is heated strongly.

27.An excess of sodium hydroxide solution is added to a solution of magnesium nitrate.

28.Solid lithium hydroxide is added to water.

29.Solutions of ammonia and hydrofluoric acid are mixed.

30.A piece of aluminum metal is added to a solution of silver nitrate.