Significance and Use of the Periodic Table

Description

This is a dry-lab, hands-on activity where students must arrange elements according to properties.

Go to Top

• Ionization energies, atomic and ionic radii, oxidation states, and electronegativities are four periodic atomic properties of great help in understanding and predicting the chemical properties of elements.
• The chemical and physical properties of the elements depend upon their electronic configurations -- the way that their electrons are arranged. The elements in the periodic table are found to be arranged according to electronic configuration. Knowing the position of an atom in the periodic table allows us to predict its electronic configuration. This, in turn, allows us to predict its most common oxidation states and, based upon knowledge of how other chemicals in the group behave, to predict its chemical reactivity patterns.
• The chart was developed on the basis of similar properties of elements before an understanding of electronic configuration was available. The first activity simulates the historical activity. The second activity takes electronic structure into account.

Go to Top

1. Your instructor will give you a periodic table listing elements and their properties. In this table, the properties of one or two of the elements will be missing.
2. Your task is to predict the properties of the missing elements. The properties of the elements above and below the missing element will provide clues; so will the properties of the elements to the left and right. Your prediction is to be based upon interpolation and extrapolation of these data.
3. Obtain the accepted values for the missing element(s) from a handbook or from the instructor. Compute the percent deviation between the accepted and predicted value.
4. Your instructor will provide you with a table showing the electronic arrangements of atoms. Some of the elements will be missing.
5. Fill in the electronic arrangements for the missing elements. Predict the valance.

Go to Top

Name _____________________________ Class _______

Teacher______________________________

DoChem 033 Significance and Use of the Periodic Table

1. Your instructor will give you a periodic table listing elements and their properties. In this table, the properties of one or two of the elements will be missing.
2. Your task is to predict the properties of the missing elements. The properties of the elements above and below the missing element will provide clues; so will the properties of the elements to the left and right. Your prediction is to be based upon interpolation and extrapolation of these data.
3. Obtain the accepted values for the missing element(s) from a handbook or from the instructor. Compute the percent deviation between the accepted and predicted value.
4. Your instructor will provide you with a table showing the electronic arrangements of atoms. Some of the elements will be missing.
5. Fill in the electronic arrangements for the missing elements. Predict the valance.

Questions

1. State the trend for the following properties of the elements, from left to right, on the Periodic Table, using the words increases or decreases: atomic size (radius); metallic properties; and ionization energy.
2. Prepare a graph of the ionization energies of the elements versus atomic number. Compare the ionization energies of metals with those of nonmetals.
3. Which would you predict to be more reactive, potassium or rubidium? Explain your answer.
4. Is the trend in ionization energies paralleled by a corresponding trend in metallic to nonmetallic characteristics? Illustrate using the elements in the horizontal row beginning with Rb.
5. Identify the most metallic element of Group VA, and explain the basis for your prediction.
6. Identify the most metallic element in the row beginning with potassium, and justify your selection.
7. State the relationship between the group number in the chart and the number of valence or outermost electrons in an atom.
8. Name the element with the outermost electron configuration 4s²4p⁴.
9. Predict the most common oxidation state of gallium, and justify your prediction.
10. Give the family names for the following groups: IA; IIA; VIIA; and 0.

Go to Top

Purpose

• To examine the periodic relationships of the chemical and physical properties of elements.
• To relate the periodicity to electronic structure.
• To predict properties of "unknown" elements by studying the known properties of neighboring elements.

Go to Top

There are no chemicals to use for this activity. Students should bring graph paper.

Go to Top

The teacher may "eliminate" any element by selective photocopying or graphical editing.

Here is a set of figures in addition to the three above, which you may find useful in this activity:

Go to Top

Teacher Preparation: 10 minutes

Class Time: May be used as a homework assignment.

Go to Top

There are no special hazards known for this activity.

Go to Top

Perform this activity in a classroom environment, not in a laboratory.

Go to Top

Save the materials for the next time this activity is used.

Go to Top

Closure Questions:

1. State the trend for the following properties of the elements, from left to right, on the Periodic Table, using the words increases or decreases: atomic size (radius); metallic properties; and ionization energy.
2. Prepare a graph of the ionization energies of the elements versus atomic number. Compare the ionization energies of metals with those of nonmetals.
3. Which would you predict to be more reactive, potassium or rubidium? Explain your answer.
4. Is the trend in ionization energies paralleled by a corresponding trend in metallic to nonmetallic characteristics? Illustrate using the elements in the horizontal row beginning with Rb.
5. Identify the most metallic element of Group VA, and explain the basis for your prediction.
6. Identify the most metallic element in the row beginning with potassium, and justify your selection.
7. State the relationship between the group number in the chart and the number of valence or outermost electrons in an atom.
8. Name the element with the outermost electron configuration 4s²4p⁴.
9. Predict the most common oxidation state of gallium, and justify your prediction.
10. Give the family names for the following groups: IA; IIA; VIIA; and 0.

Answers to Closure Questions:

1. Radii decrease, metallic properties decrease, and ionization potentials increase as one goes from left to right in a row of the Periodic Table.
2. The ionization energies of metals are lower than those of nonmetals. It is easier to remove an electron from a metal atom than from a nonmetal atom.
3. Rubidium is more reactive than potassium because rubidium is a larger atom than potassium and can lose its electron easier.
4. Rubidium has the lowest ionization energy in its row, and is the most metallic element. Xenon has the highest ionization energy in that row, and it is an inert gas. Iodine has the next largest ionization energy in the row, and it is the least metallic, most nonmetallic element in that row.
5. Bismuth is the most metallic element listed in Group VA. It is the last member of the group. It has the lowest ionization energy and the largest radius of the members of that group.
6. Potassium is the most metallic element in its row. It has the lowest ionization energy, and is very reactive.
7. The group number is equal to the number of electrons in the outermost shell.
8. Selenium has outermost electron configuration 4s²4p⁴.
9. The most common oxidation state for gallium would be 3+. Gallium has 3 outermost electrons, which it would lose in the same way as aluminum.
10. Group IA, alkali metals; Group IIA, alkaline earth metals; Group VIIA, halogens; and Group 0, inert gases.

Go to Top

• acid
• base
• metal oxide
• nonmetal oxide
• oxidation
• solubility
• periodic properties

Go to Top