Page 1 of 12

NAME_________________________________

PERIODIC TRENDS Summer Assignment 2014

You will need internet access to complete this assignment. You can use other resources as needed.

I. Introduction – read carefully!

In the late 1800’s, Dmitrii Mendeleev first proposed a table in which the elements are arranged according

to physical and chemical properties. This table has been modified since that time, but physical and chemical

properties can still be predicted using the periodic table. Trends in various properties can be observed both across

periods and down groups. However, deviations from the trends may occur in a variety of properties. These

variations can usually be explained by electron-electron repulsion, nucleus-electron attraction, or shielding.

Electron configuration and orbital diagrams are helpful tools to help determine the primary effects involved in the

trends as well as deviations from the trend.

In this activity, you will use computer program, the internet and your test to explore the periodic trends of

the following properties: Atomic size, ionic radii, ionization energy, electron affinity, electronegativity and

oxidation numbers. These trends will be explained in terms of the structure of the atoms. Since structure

determines function, a complete understanding of how the structure of the atom affects the function of the atom

will complete our study of atomic structure.

Exact values of atomic size are difficult to determine because atomic wave functions describe probability

distributions and do not have sharp boundaries. Atomic radii can be estimated from data on bond lengths me

measuring the interatomic distance and dividing by two. Even though data from various methods of measurement

can vary, the trends can be studied as long as the atomic radii were all determined by the same method. These

values are usually reported in pm, which is equal to 10-12 m.

The ionization energy is the energy that is required to remove an electron from an atom in the gaseous

phase. You will study sequential ionization energy processes. During the analysis, be cognizant of from which

sublevel and orbital an electron is being removed and whether or not the sublevel is full or half-full. It is not

sufficient to explain a trend by stating the sublevel is “full” or “half-full”. Use forces of attraction and repulsion

along with shielding and a simplified view of effective nuclear charge (Zeff).

The electron affinity is a measure of the energy involved in the process of adding an electron to an atom

in the gaseous phase. The electron affinity is strongly affected by the effective nuclear charge. Be careful

interpreting electron affinity! A high electron affinity in the vocabulary sense indicates the addition is exothermic

and thus has a negative value. A low electron affinity

in the vocabulary sense has a positive value.

II. Models

You will be asked to use a variety of models to help

formulate your answers. The energy level diagram to

the right is also a great reference. Notice how close the

energies of n = 3, 4, 5 are to each other in comparison to

n = 1 & 2.

Page 2 of 12

III. Introductory Terms & Concepts

1. Define the Universal force: Oppositely charged particles _______________ (fill in the blank & finish the

statement)…

2. Describe how these two types of forces are balanced in the ground state arrangement of electrons around

the nucleus

3. Use arrows to diagram these inter-particle forces for the Lithium atom. The helium atom shown on the

left provides a model for you. (What do you think Z stands for?)

4. Define the following terms

a. “SHIELDING” ( see figure for

hints and

http://facultyfp.salisbury.edu/dfrieck/htdocs/212/rev/zeff/shielding.htm )

b. Effective nuclear charge (Wikipedia.org)

Z:

+ 2

─

─

Page 3 of 12

c. Valence electrons

d. Core electrons

5. Write the electron configuration for tin. Underline the core electrons and circle the valence electrons.

6. Calculate the effective nuclear charge for tin.

7. Would shielding have a greater effect across a period or down a group? Explain.

8. Define and write a generalized equation for the following:

a. Ionization energy

b. Electron Affinity

9. Compare and contrast the terms “electron affinity” and “electronegativity”

ELECTRON AFFINITY ELECTRONEGATIVITY

10. As a generalized principle, will an increase in attraction to the nucleus tend to increase or decrease each

of the following? Explain

a. The size of an atom

Page 4 of 12

b. The ionization energy

c. The electronegativity

11. As a generalized principle, will an increased repulsion between electrons tend to increase or decrease each

of the following? Explain

a. The size of an atom

b. The ionization energy

c. The electronegativity

IV. Procedure

Log on to the internet and go to www.webelements.com . Click on the “Periodicity” tab at the top, and

then choose the property you are observing. If you are asked to make comparisons, do the graphs separately and

note the scale on the y-axis. It is not necessary to print or draw the graphs. You can also view the trend

directly on the periodic table. Explore the programs to see what works best for you.

All trends and deviations from trends should be explained in terms of the electron configuration,

shielding, effective nuclear charge, and the structure of the atom. Sample electron configurations and/or orbital

diagrams should be used to support the answer. You may use the noble gas configuration only when instructed.

(You do not need to print the graphs.)

V. Analysis

A. Atomic Radii

12. How are atomic radii measured?

13. Use the program to observe the graph of Radius-metallic. Focus on the atomic radii for groups 2 and 17

(or 7A). Describe the trend in the values down a group and provide an explanation for the trend based

on atomic structure.

Page 5 of 12

14. Observe the atomic radii data for period 4. Describe and explain the periodic trend.

15. Explain why the difference in atomic radii between rubidium and potassium is less than the difference

between sodium and lithium. Refer to the energy level diagram on page one for help

B. Ionic Radii: Go back to the list of periodic trends and look at the graph of Ionic Radius (Pauling).

When you put the cursor on the point for specific elements, the data will be displayed.

16. Compare the atomic radii and the ionic radii on the same graph for each of the groups 2 and 17 (or 7A).

ELEMENT ATOMIC RADII IONIC RADII

Potassium

Strontium

Sulfur

Bromine

a) Which is larger for cation formation: atomic radii or ionic radii (circle one) Explain.

b) Which is larger for anion formation: atomic radii or ionic radii (circle one) Explain.

c) Fill out the following chart. Justify your claim in 16a regarding cation size using atomic structure.

Particle Electron configuration Outer electron

config.

Sr

Sr2+

Justification:

Page 6 of 12

d) Fill out the following chart. Justify your claim in 16b regarding anion size using atomic structure.

Particle Orbital diagram for n=4 electrons

Br [Ar] ____ ____ ____ ____ ____ ____ ____ ____ ____

4s 3d 4p

Br─

[Ar] ____ ____ ____ ____ ____ ____ ____ ____ ____

4s 3d 4p

Justification:

e) Write the complete electron configurations for the following ions. Determine the number of

subatomic particles in each ion. Rank the ions from smallest (1) to largest (5). Justify your

prediction.

ION ELECTRON CONFIGURATION # of

Protons

# of

Electrons

RANK

P3─ 1s22s22p63s23p6

S2─

Cl─

K+

Ca2+

JUSTIFICATION

Page 7 of 12

C. Ionization Energies

17. Examine the graph of the 1st ionization energy for periods 2 and 3.

a)

Describe and explain the period trend.

b) Fill in the orbital diagrams below for Mg and Al. Note the 1st ionization energies for these two

elements do not follow the trend. Explain this anomaly (a deviation from the common rule or trend)

observed between Mg and Al in period three.

Mg: [Ne] ____

3s

Al: [Ne] ____ ____ ____ ____

3s 3p

c) Fill in the orbital diagrams below for N and O and explain the anomaly observed between N and O

in period two

N: [He] ____ ____ ____ ____

2s 2p

O: [He] ____ ____ ____ ____

2s 2p

18. Examine the 1st ionization energy for all elements below.

Page 8 of 12

Look

specifically at the

elements in groups 2 and 16. Describe and explain in terms of the universal force the trend down a group. Are

there exceptions?

19. Look up the 2nd ionization energy for the following representative elements in groups 1, 2, and 13. Fill in

the orbital diagrams and circle the electron that would be removed during the second ionization step.

Why is there such a large difference between the first and second ionization energies in group 1 compared

to groups 2 and 13?

Page 9 of 12

GROUP 1st

ionization

energy

2nd

ionization

energy

ORBITAL diagram for a representative element. Circle the 2nd

e─ that is removed in each

1

Na

____ ____ ____ ____ ____ ____ ____ ____ ____

1s 2s 2p 3s 3p

2

Mg

____ ____ ____ ____ ____ ____ ____ ____ ____

1s 2s 2p 3s 3p

13

Al

____ ____ ____ ____ ____ ____ ____ ____ ____

1s 2s 2p 3s 3p

Explanation of difference:

20. Examine the table of successive ionization energies in the table below.

Determine whether a valence electron is being removed or a core electron. Where does the largest

increase in ionization energy occur? Explain using atomic structural principles.

21. Write reaction equations for the 1st, 2nd, 3rd ionization for arsenic. Based on the trend observed for

aluminum, where would you predict the major increase in ionization energy to occur for arsenic? Why?

IONIZATION IONIZATION REACTION

1ST

As + IE1 As+ + e-

2ND

3RD

QUESTION ANSWER:

Page 10 of 12

D. Oxidation Numbers

22. Define the term “oxidation number” or “oxidation state”. (These are synonyms so only one term needs to

be defined.)

On the last page is a periodic table with oxidation numbers listed. Use it to answer these questions.

23. Explain the common oxidation numbers listed below using electron configurations. Write the electron

configuration and CIRCLE the electrons that would be removed for cations. For anion formation,

CIRCLE the orbital/sublevel that would change. Write the electron configurations for the ions formed.

Element ION NOBLE GAS ELECTRON

CONFIGURATION FOR ATOM

NOBLE GAS ELECTRON

CONFIGURATION FOR ION

K +1

Mn +2

+3

+7

Element ION NOBLE GAS ELECTRON

CONFIGURATION FOR ATOM

NOBLE GAS ELECTRON

CONFIGURATION FOR ION

As -3

+3

+5

Br -1

+1

+5

E. Electron affinity & Electronegativity

NOTE: Electron affinity is a measure of the actual energy involved in the process. A negative value is

exothermic and a positive value is endothermic. Since exothermic processes are FAVORABLE, as the VALUE

for the electron affinity becomes more NEGATIVE, the element is said to have a HIGHER affinity in the

vocabulary sense. Seems a little flip-flopped! When in doubt – look for fluorine – the element that has a strong

affinity for electrons. Because this can get confusing, we will study the trends in ELECTRONEGATIVITY.

Electronegativity is a central concept in the discussion of bonding.

1. Graph the electronegativity values for period 3. Describe and explain the general trend.

2. Graph the electronegativity values for group 1. Describe and explain the general trend.

Page 11 of 12

Page 12 of 12