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Periodic Properties of the Halogens

Reminder: Lab practical in the next two labs!

Lab practicals will take place during labs 5 and 6. For more information, see Instructions for Pipetting/Dilution for Lab Practical in the lab manual appendix.

Goals

  • To explore periodic trends among the halogens (Group 17 (VIIA) elements).
  • To design and use qualitative identification techniques.
  • Improve observation skills

Background

Periodic Trends

The Periodic Table of the Elements is so named because of the periodic variation of many important elemental properties along a row (Period) or down a column (Group or Family) of the table. For example, atomic size decreases from left to right along a Period as well as from bottom to top within a Group. Ionization energy, electronegativity, and electron affinity also vary in similar manners (see the sections in your textbook).

Oxidation-Reduction Reactions

The effects of these periodic variations in properties are often expressed in the relative chemical reactivities of the members of a Period or Group. This experiment is designed to allow you to determine the relative reactivities of three (Cl, Br, and I) of the Group 17 elements (i.e the halogens) in reduction - oxidation reactions (referred to as redox reactions). Redox reactions involve the exchange of electrons. In the example below, 'X' represents a halogen (e.g. F, Cl, Br, or I) and 'Y¯' denotes a halide ion. A reaction of halogen X2 with halide Y¯ would be as follows: 

X2  +  2  Y¯   →    2  X¯  +  Y2

Redox reactions are best understood when written in terms of two half reactions.

A Redox Reaction can be divided into two half reactions:

A half reaction  that

With a reactant that undergoes

That reactant is defined as

Loses Electron(s)

Oxidation

Reducing agent

Gains Electron(s)

Reduction

Oxidizing Agent

Viewed as a half reaction the reduction of X2 would be:

X2  +  2 e¯   →   2  X¯               a gain of electrons

And the oxidation half reaction of Y¯ would be:

2  Y¯   →     Y2  +  2 e¯             a loss of electrons

In this reaction, X2 is being reduced; it is called the oxidizing agent because it is taking electrons away from Y¯, causing Y¯ to be oxidized.  Conversely, Y¯ is being oxidized; it is called the reducing agent because it is causing X2 to be reduced.

Upon successful completion of the exercise, you will be able to arrange these halogens in order of their relative reactivity as oxidizing agents. Likewise, you will be able to arrange the halides in order of their relative reactivity as reducing agents.

Balancing a Redox Reaction using the Half-reaction Method

Reaction to be balanced:           MnO4-1(aq) + I-(aq) → Mn2+(aq) + I2(aq) (in acidic solution)

  1. First identify the oxidation and reduction state of the atoms involved.

    2. The oxidation state of oxygen in a compound is typically -2. In addition, the oxidation states of atoms in a molecule must add up to the charge of the molecule. Therefore, the oxidation state of Mn in MnO4-1 is +7.

    The oxidation states of I and Mn in I-, Mn2+, and I2 are -1, +2, and 0, respectively.

  2. Next, determine which species are being oxidized and which are being reduced based on the oxidation states for all atoms. Then separate the equation into two half reactions. Include electrons in your reactions.

    Oxidation half reaction: 2I-(aq) →I2(aq) + 2e-

    Reduction half reaction: MnO4-1(aq) + 5e- → Mn2+(aq)

  3. If there are unblanaced oxygen atoms in your half reactions, add H2O(l) and H+(aq) (because we are working with acidic solution) to balance.

    Oxidation half reaction: 2I-(aq) →I2(aq) + 2e-

    Reduction half reaction: 8H+(aq) + MnO4-1(aq) + 5e- → Mn2+(aq) + 4H2O(l)

  4. Before the half reactions can be added together, the half reactions may need to be modified to ensure that the overall reaction will consume the same number of electrons that are being produced. In this case, we have two electrons produced in the oxidation half reaction and five consumed in the reduction half reaction. The least common multiple is 10. So we multiply each half-reactant by appropriate coefficients so that ten electrons are exchanged.

    Oxidation half reaction: 10I-(aq) → 5I2(aq) + 10e-

    Reduction half reaction: 16H+(aq) + 2MnO4-1(aq) + 10e- → 2Mn2+(aq) + 8H2O(l)

  5. Finally, combine the two half reaction to give the overall reaction.  Verify that the numbers of atoms and the charges are balanced.

    Balanced redox reaction: 16H+(aq) + 10I-(aq) + 2MnO4-1(aq) → 5I2(aq) + 2Mn2+(aq) + 8H2O(l)

Qualitative Analysis

A chemical reaction may be analyzed quantitatively or qualitatively. In quantitative analysis, the amounts of reactants and products are carefully measured and those data are used to make precise interpretations about the reaction. In qualitative analysis, changes in a reaction mixture are observed and that information is used to determine whether or not a reaction ocurred. Qualitative changes include: the formation of a precipitate, the evolution of a gas, the release of energy as heat, the dissolution of a solid, and changes in the color or changes in clarity (clear vs. cloudy) of a solution. In this experiment, we will be observing color and clarity changes accompanying redox reactions, i.e. we will be doing qualitative analysis.

Spectator Ions and Net Ionic Equations

Many ionic compounds dissolve in water to form solvated ions. With rare exceptions, salts containing Group 1 elements are water soluble. For example, NaCl dissolves to form Na+(aq) and Cl¯(aq), where (aq) means dissolved in aqueous (water) solution. In addition to atomic ions, many molecular ions exist that remain as a molecular unit when dissolved, e.g. potassium permanganate, KMnO4, dissolves to form K+(aq) and MnO4¯(aq). When we write a reaction, it can be written in terms of ionic compounds:

NaOH + HCl   →  H2O + NaCl

or in terms of solvated ions:

Na+(aq) + OH¯(aq) + H+(aq) + Cl¯(aq)   →  H2O + Na+(aq) + Cl¯(aq)

You may notice that there are some species in the above equation that do not participate in the reaction. They show up identically on both sides of the equation. They are known as spectator ions. When a reaction is written in terms of solvated ions, the spectator ions should be removed, leaving the net ionic equation:

OH¯(aq) + H+(aq) →  H2O

This reaction is not a redox reaction, so we do not have to worry about balancing half reactions. In your results for this lab, you will have to balance half reactions and then write the resulting net ionic equations.

Synopsis of the Experiment

In Part I of this experiment, you will develop a qualitative identification scheme which you can use to identify the presence of each halogen/halide used in the lab.

In Part II, you will mix each halide with each halogen to see which halogens are capable of oxidizing which halides.  Using your identification scheme, you will be able to identify products if they form.  In addition, you will combine each halogen with solid copper to see whether it can oxidize the copper.  Using your results, you will be able to draw conclusions about the relative oxidizing strengths of the halogens.

In Part III, you will mix each halide with CuSO4 and with KMnO4.  Using your qualitative identification scheme, you will be able to identify products if the halides are successful as reducing agents.  You will be able to draw conclusions about the relative reducing strengths of the halides.

Preparation

Reading Assignment:

  • Description of experiment—see below.
  • The Group 7A Elements—Zumdahl, sec. 19.7, pp. 914–921.
  • Periodic Properties—Zumdahl, sec. 12.15, pp. 560–568.
  • Review Oxidation-Reduction Reactions—Zumdahl, sec. 4.10–4.11, pp. 115–129.

Questions:

    Fill out the prelab worksheet that can be found at the end of the Experiment section.

Experiment

To print instructions, select the portion that you with to print, choose File/Print, and choose "selection" to prevent printing the entire document.

Safety

**Aqueous chlorine and bromine may produce vapors that cause eye irritation. Make sure to keep the caps on these solutions when they are not in use. Please rinse your eyes thoroughly with water if you experience the slightest eye irritation in lab.

**Copper sulfate solution is a skin and eye irritant. Please rinse any affected area with copious amounts of water.
  • Wear safety glasses!
  • Do not pour solutions down the drain! Collect all halogen/halide wastes in the appropriately labeled disposal containers provided.

Materials and Equipment

  • Test tubes (13 mm)
  • Mineral oil
  • Aqueous solutions of Cl2, Br2, and I2
  • 1 M aqueous solutions of NaCl, NaBr, and NaI
  • 1 M CuSO4
  • 0.01 M KMnO4 in 1 M H2SO4
  • Cu turnings
  • Parafilm

Instructions

Students will work in groups of 2 unless otherwise specified by your instructor.

When preparing to perform an experiment, it is often advantageous to construct a chart or table in your notebook for recording observations and results. This not only guides your work while you are conducting the experiment but also makes review of the results easier after the work is complete. Sample tables for Parts 1 and 2 of this experiment are given. You will need to prepare your own table for recording your results of Part 3.

Before you perform Part 1 - prepare three solutions for use in Part 2

  1. Obtain 3 test tubes.
  2. In each test tubes, place a small piece of copper turning (approximately 2 to 3 cm in length). Fill each about half full with a different aqueous halogen solution. Cover the test tubes with Parafilm and then allow them to sit for at least one hour.

Part 1. Preliminary Study - Making a Qualitative Identification Scheme

Before we can perform reactions, we need a qualitative identification scheme which we can use to test for presence of the halogens and halides that will be reactants and products in our reactions. Begin by observing the provided halogen or halide solution in two different solvents: mineral oil (density of ~0.8g/mL) and water (density of ~1.0 g/mL).  In the elemental form, halogens (X2) are neutral and preferentially dissolve in mineral oil because the molecules are nonpolar.  In water the halogens have a pale color, while in mineral oil the halogens have a darker color (for example, iodine is deep purple).  Halides (X¯) are negatively charged and readily dissolve in water because the ions are polar.  Aqueous halide solutions are colorless.

  1. Prepare a table to record your data in. The table below is given as an example of the type of data table you should create in your notebook to document your observations.

    Click to open an Excel version of this table

    Part 1. Preliminary Study –Qualitative Identification Scheme

    table

  2. Obtain 6 test tubes for Part 1. Fill each to a depth of ~2 cm with a different halogen or halide. Examine and record the colors of the aqueous solutions of each halogen and sodium halide.
  3. Next, add a layer of mineral oil ~1 cm thick to each of the test tubes containing aqueous solutions. Shake vigorously using a small square of parafilm to cover the opening of the test tube. Allow the mixture to settle for a moment.  In your notebook, record your observations, including the color and clarity of each solvent layer. 

Part 1—Questions for thought:

  1. Are the solvents miscible (capable of being mixed)? If not, which solvent forms the top layer and which solvent the lower layer?
  2. What can you conclude about the solubilities of the halogens and halides in mineral oil?  Explain the preference for one solvent or the other.
  3. Do halides or halogens possess a unique color defining their presence? Which solvent accentuates their presence in solution?

Part 2. Halogens as Oxidizing Agents

As stated earlier, we will be studying redox reactions.

When a halogen and a halide are combined, two outcomes can occur: 1) nothing, 2) the reaction above will occur. We will be comining various halogens and halides and looking for indications of the presence of products to determine whether or not a reaction occurred. In this way, the relative oxidizing strength of halogens can be determined.

Another way of testing the strength of halogens is to look for reaction products when they are combined with metals.  If a halogen is introduced to elemental copper, Cu(s), a reaction may occur in which the elemental copper disoolves and the Cu2+(aq) ion forms in the test tube (indicated by blue-green color).

  1. Prepare a Table

Prepare a table to record your data in. The table below is given as an example of the type of data table you should create in your notebook to document your observations.

Click to open an Excel version of this table

Part 2.  Halogens as Oxidizing Agents

table

  1. Form a Hypothesis
  1. Prior to performing reactions in this part of the lab, use the prelab's summary of the trends in the halogen's reactivities to form a hypothesis on whether an exchange of electrons will occur. In the data table below, pencil in a yes if you expect a reaction to occur or a no if no reaction will occur.
  2. Present your hypothesis and predictions to your instructor before proceeding with the experiment.
  1. Proceed with Experimentation
  1. In the 6 test tubes, combine the appropriate halogens and halides in approximately equal volumes. Add ~1 cm of mineral oil to each test tube and mix vigorously. Allow the contents of the test tubes to settle a moment and then record your observations.
  2. Record observations of these six test tubes as well as the three prepared earlier (halogens with Cu). It can be difficult to observe reactions of Cu with halogens, even if you wait an hour, so take a look at the solutions prepared by your instructor several hours prior to lab.
  3. Use the results from the Qualitative Identification Scheme from Part 1 to interpret the results (Y for reaction; N for no reaction). Were your predictions correct? Remember that just because a color appears, that does not mean that a reaction has occurred. A reaction has occurred if you observe disappearance of reactants and/or evidence of products.
Part 2—Questions for thought:
  1. What physical change(s) did you observe which indicated that a reaction had occurred with the copper metal?
  2. Write the balanced net ionic chemical equation for each reaction that occurred.  Do not write an equation if no reaction occurred.  Note which chemical species is being oxidized and which is being reduced.
  3. What statements can you make about the relative oxidizing strengths of the halogens? Support your statements by referring to experimental results. Arrange the halogens in order of increasing strengths as oxidizing agents.

Part 3. Halides as Reducing Reagents

One way of testing the reacting strength of a halide is to look for reaction products when it is combined with an oxidizing agent that has a distinctive aqueous color that identifies its presence. The loss of the oxidizing agent's characteristic aqueous color indicates its disappearance upon reaction to product.   If a halide is mixed with a solution of copper sulfate (CuSO4(aq)) or potassium permanganate (KMnO4(aq)), the halide may act as a reducing agent to form solid copper (Cu0(s)) or Mn2+(aq), respectively.

Now that you are familiar with the general experimental approach and how to distinguish which halogen is present in a reaction, design an experiment to investigate the reaction of each of the aqueous halide solutions with both copper (II) sulfate (CuSO4(aq)) and acidic potassium permanganate (KMnO4(aq)) solution.  If a reaction occurs, the halide will produce its corresponding halogen. Will mineral oil be helpful in identifying products for this reaction? Generate a data table similar to Part 2’s data table to record your results. Be sure to develop an identification scheme for the aqueous oxidizing agents. Do you need a control for Part 3?

Part 3—Questions for thought:

  1. Why were the copper ion and permanganate ions wise choices to serve as the oxidizing agents?
  2. What statements can you make about the relative reducing strengths of the halides? Support your statements by referring to experimental results. Arrange the halides in order of increasing strengths as reducing agents.
  3. Write the balanced net ionic chemical equation for each reaction that occurred (net ionic means that spectator ions [ions that are not involved in the reaction] are not listed).  Do not write an equation if no reaction occurred. Note which chemical species is being oxidized and which is being reduced. The exercise below is designed to help you brush up on the skills you need in order to answer this question.

Prelab & Results

Click here for Prelab worksheet: pdf or Word format

Click here for Results worksheet: pdf or Word format

Hints for balancing reactions:

  1. Two half reactions, one oxidation and one reduction, must be added together to make a redox reaction. See example above.The number of electrons produced in the oxidation reaction must be the same as the number of electrons used in the reduction reaction. See example above.
  2. If a halogen reacts, it becomes the corresponding halide. If a halide reacts, it becomes the corresponding halogen. See half reactions above.
  3. Solid copper has no charge (Cu0(s)). It becomes Cu2+(aq).
  4. Copper sulfate (CuSO4(aq)) separates in solution to become Cu2+(aq) and SO42–(aq). SO42–(aq) is a spectator ion (does not participate in the reaction). The product is solid copper (Cu0(s)).
  5. Potassium permanganate (KMnO4(aq)) separates in solution to become K+(aq) and MnO4–(aq). K+(aq) is a spectator ion. The product is Mn2+(aq). This is a complex reaction. It is worked out in full above.

Question for thought:

  1. Explain the trends you discovered in this experiment based on the position of the halogens/halides in the Periodic Table. Use molecular and ionic size, ionization energy, and electron affinity to support your conclusions. 

For more information:

Group 17 – The Halogens:
http://www.chemsoc.org/viselements/pages/data/intro_groupvii_data.html
Royal Society of Chemistry web site which has a visual elements section and other chem. stuff

Introduction to the Halogens:
http://www.wpbschoolhouse.btinternet.co.uk/page03/The_Halogens.htm
“Doc Brown Chemistry Clinic”  by a chemistry teacher in the UK
very colorful, with some cute graphics (can “cute” and “chemistry” be used on the same page?!)

Redox Balancing Practice Problems:
http://wc.pima.edu/~skolchens/C152OL/Ch21/REDOX.htm
http://www.chemistrycoach.com/Redox.htm (Has solutions too!)

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Created By: Adilia James '07 and Sarah Coutlee '07
Maintained By: Nick Doe
Date Created: July 3, 2006
Last Modified: October 22, 2009
Expiration Date: July 3, 2009