Tuesday 12 th June 2007

Periodic Table Lab

Chemical Concepts Demonstrated: Flame tests are used to identify the presence of a relatively small number of metal ions in a compound. Not all metal ions give flame colours. For Group 1 compounds, flame tests are usually by far the easiest way of identifying which metal you have got. For other metals, there are usually other easy methods which are more reliable - but the flame test can give a useful hint as to where to look.

What are the limitations of this test?
The value of the flame test is limited by interference from other brighter colors and by ambiguities where certain different metals cause the same flame color. Sodium, in particular, is present in most compounds and will color the flame. Sometimes a colored glass is used to filter out light from one metal. Cobalt glass is often used to filter out the yellow of sodium.

Experimental:

Clean a platinum or nichrome (a nickel-chromium alloy) wire by dipping it into concentrated hydrochloric acid or nitric acid and then holding it in a hot Bunsen flame. Repeat this until the wire doesn't produce any colour in the flame.

When the wire is clean, moisten it again with some of the acid and then dip it into a small amount of the solid you are testing so that some sticks to the wire. Place the wire back in the flame again.

If the flame colour is weak, it is often worthwhile to dip the wire back in the acid again and put it back into the flame as if you were cleaning it. You often get a very short but intense flash of colour by doing that.

Chemicals/Materials:

1. Chloride salts of Li, Na, K, Rb, Cs, Ca, Ba, Cu, Pb, Fe (II) and Fe(III) Sr (nitrate salt).
2. Glass rods with loops of Pt wire.
3. Bunsen burner/clicker.
4. Concentrated nitric acid or hydrochloric acid.

Observations:

flame colour
Li	red
Na	strong persistent orange
K	lilac (pink)
Rb	red (reddish-violet)
Cs	blue? violet? (see below)
Ca	orange-red
Sr	red
Ba	pale green
Cu	blue-green (often with white flashes)
Pb	greyish-white

It should be noted that sodium is present as an impurity in many if not most metal salts. Because sodium imparts an especially intense color to a flame, flashes of the sodium may be observed in nearly all solutions tested.

Scientific Concepts:

When an element is burned, the electrons will be excited. Then as these electrons fall back from one energy level to another, they will emit photons of light. These photons will have different colors depending on the element and its discrete energy levels. That is, different wavelengths of light (colors) will be emitted when the electrons of different elements go down the step(s) between their energy level(s). Each element will have its own set of steps, therefore each will have its own color or set of colors

The Periodic Table - Not Just A Work Of Art!

Relative Reactivity of Metals and the Activity Series

Objective

The goals of this experiment are:

• to observe the reactions of several metals with cold water, hot water, acids and then other metal ions.
• to prepare an activity series of the metals based on the observations from the above reactions.

Introduction

What is there to know about the periodic table? Why is it important? Why does it appear in nearly every science lecture room and labs? Is it just a portrait of an aspect of chemistry or does it serve a useful purpose? Why is the name periodic appropriate? Why is the table arranged in such a way? What are the important features of the table? Does it give order to the 109 known elements?

By the end of thiese experiments, you will be able to answer all the above questions.

First, we are going to look at the relative reactivity of some of the elements.

Relative Reactivity of Metals and the Activity Series

A superficial glance at the Periodic Table will reveal that all known elements are listed by their chemical symbols. An in depth glance at the Periodic Table yields information on the mass of an atom of the element in atomic mass units (amu) for the molar mass of a mole (6.02 x 10²³) of atoms in grams below the chemical symbol for each element. Above the chemical symbol for each element, there is a second number listed, the atomic number, which gives the number of protons (positively charged particles in the nucleus), or the number of electrons (negatively charged outside the nucleus) for a neutral atom.

Mendeleev arranged the elements in the Periodic Table in order of increasing atomic number in horizontal rows of such length that elements with similar properties recur periodically; that is to say, they fall directly beneath each other in the Table. The elements in a given vertical column are referred to as a family or group. The physical and chemical properties of the elements in a given family change gradually as one goes from one element in the column to the next. By observing the trends in properties, the elements can be arranged in the order in which they appear in the Periodic Table.

We will be looking closely at the alkaline metals (Group 1) and the metals in general as well as the halogens (Group 7).

PROCEDURE

I. Activity Series

Part 1. Reactions of Metals with Water

1.  Place 5 mL H₂O in each of four clean tubes and label them as follows:

A.	^Mg
B.	^Cu
C.	^Zn
D.	^Ca

2. Use sandpaper or steel wool to remove the oxide from the surfaces of Mg, Cu, and Zn.
3. Place several small pieces of Mg, Cu, and Zn in the correctly labeled test tube prepared above. Place two or three (not more!) pieces of Ca turnings in the test tube labeled "Ca".
4. Watch for evidence of reaction by noting evolution of gas bubbles and any change in the color or size of the metal. Record your observations and write net ionic equations for each reaction.

Note: Trapped air bubbles on the metal surfaces are not indicative of a reaction.

CAUTION: H₂ is FLAMMABLE!

CAUTION: Residual calcium should be discarded in a special container designated by your instructor.

Metal	Observations
Mg
Cu
Zn
Ca

Part 2. Reactions of Metals with HCl

CAUTION: The reaction of Ca with HCl is not studied. Residual calcium should be discarded in a special container designated by your instructor.

1. Decant the water from each test tube used in the procedure above and leave the pieces of metal that remain unreacted in each test tube.
2. Place the test tubes in a test tube rack/holder.
3. Add 2 mL of 3 M HCl solution to each test tube.

CAUTION: Some of the test tubes may become very hot. Leave them in the rack/holder while you are making observations.

4. Observe relative rate of H₂ gas evolution for up to 10 minutes and record your observations on your report form.
5. Based on the observations in the previous steps, list the elements that react in 3M HCl in order of increasing strength as reducing agents and write net ionic equations for all reactions.

Metal	Observations
Mg
Cu
Zn

Part 3. Reactions of Metals with Other Metal Ions

1. Place a clean 1 inch-square of metal foil (sheet) of each of these metals Cu, Zn and Pb on a flat surface.
2. Clean the metal surfaces by sanding them with fine sandpaper or steel wool.
3. Place one or two drops in spots of each of these solutions in a clockwise order on the metal surfaces:

A.	0.5 M Ag^+
B.	0.5 M Cu^2+
C.	0.5 M Zn^2+
D.	0.5 M Pb^2+

4. NOTE: Do not test a cation of a metal on a square of the same metal such as Cu²⁺ ion and Cu metal.

5. Watch for color changes in each spot as evidence of reaction. If you are not sure whether the reaction has occurred, rinse the plate with water. A distinct spot of a different color on the surface is good evidence for the reaction.
6. Write net ionic equations for each reaction . Arrange Ag, Cu, Pb and Zn in order of their increasing strength as reducing agents. If a metal A reacts with a cation of another metal B, metal A is a stronger reducing agent, more reactive than metal B.
7. Rinse and dry each square of metal and return it to the correct beaker on the reagent shelf for other students to use.

	Zn	Cu	Pb
Ag^+
Cu^2+		Do not test
Zn^2+	Do not test
Pb^2+			Do not test

II. Ionic Equations

1. Clean a 24 -well plastic plate. Fill in 1/4 of the wells (about 20 drops) with the following solutions according to the table:

	1	2	3	4	5
A	0.1 M Mg(NO3)2	0.1 M ZnSO4	0.1 M Pb(NO3)2	0.1 M AgNO3	0.1 M Cu(NO3)2
B	0.1 M Mg(NO3)2	0.1 M ZnSO4	0.1 M Pb(NO3)2	0.1 M AgNO3	0.1 M Cu(NO3)2
C	0.1 M Mg(NO3)2	0.1 M ZnSO4	0.1 M Pb(NO3)2	0.1 M AgNO3	0.1 M Cu(NO3)2

2. Add the same amount (20 drops) of 0.1 M NaOH in row A. Add 20 drops of 6M NaOH to the solutions in row B. Add 20 drops of 0.1 M NaI to the solutions in row C.
3. Mix the solutions with a glass rod of a toothpick. Rinse the glass rod or get a new toothpick when changing wells. Record your observations and write ionic equations.
4. Fill out your observations in the following table:

	Mg(NO3)2	ZnSO4	Pb(NO3)2	AgNO3	Cu(NO3)2
0.1 M NaOH
6 M NaOH
0.1 M NaI