AP Chemistry

Unit 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 solution

3.2 Distinguish between various types of acids and bases.

3.3 Use titration to determine the neutralization point of a chemical reaction

3.4 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, 60

Chapter 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 Solutions

Parts 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 not

Types 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 Mass

Example 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 Preparation

Example 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?

Electrolytes

Electrolytes 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.  NaCl
2.  CH3OH (methyl alcohol)
3.  C3H12(OH)3 (glycerol)
4.  HCl
5.  C6H12O6 (glucose)
6.  NaOH
7.  C2H5OH (ethyl alcohol)
8.  CH3COOH (acetic acid)
9.  NH4OH
10.  H2SO4

Molarity (M)

Molarity= moles of 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 Dilution

Acids 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…

M1×V1=M2×V2

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 Reactions

Acids

•  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 ammonia

Titrations: 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 Lab

Introduction

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:

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 Procedure

1.  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: