Lab Investigation: Periodic Trends Data Analysis
Investigation Overview
In this investigation, you will analyze real data about atomic properties to discover and verify periodic trends. This activity develops skills tested on the ACT Science section and SAT data analysis questions.
Learning Objectives
- Analyze data tables and graphs showing atomic properties
- Identify patterns and trends in experimental data
- Draw conclusions based on evidence
- Explain the relationship between atomic structure and observed properties
Data Set: Atomic Radii (in picometers)
| Element | Atomic Number | Period | Group | Atomic Radius (pm) |
|---|---|---|---|---|
| Li | 3 | 2 | 1 | 152 |
| Be | 4 | 2 | 2 | 112 |
| B | 5 | 2 | 13 | 85 |
| C | 6 | 2 | 14 | 77 |
| N | 7 | 2 | 15 | 75 |
| O | 8 | 2 | 16 | 73 |
| F | 9 | 2 | 17 | 72 |
| Na | 11 | 3 | 1 | 186 |
| Mg | 12 | 3 | 2 | 160 |
| K | 19 | 4 | 1 | 227 |
Data Set: First Ionization Energy (kJ/mol)
| Element | Atomic Number | First Ionization Energy |
|---|---|---|
| Li | 3 | 520 |
| Be | 4 | 900 |
| B | 5 | 801 |
| C | 6 | 1086 |
| N | 7 | 1402 |
| O | 8 | 1314 |
| F | 9 | 1681 |
| Na | 11 | 496 |
| K | 19 | 419 |
Analysis Questions
Question 1: Looking at Period 2 elements (Li through F), describe the trend in atomic radius as you move from left to right.
Answer: Atomic radius decreases from left to right across Period 2. Li (152 pm) has the largest radius, while F (72 pm) has the smallest.
Question 2: Compare Li, Na, and K (all Group 1 elements). What trend do you observe in atomic radius going down the group?
Answer: Atomic radius increases going down Group 1: Li (152 pm) < Na (186 pm) < K (227 pm). Each element has electrons in a higher energy level.
Question 3: Based on the ionization energy data, which element in Period 2 is hardest to ionize (remove an electron from)? Why?
Answer: Fluorine (1681 kJ/mol) is hardest to ionize. It has the smallest atomic radius and highest effective nuclear charge in Period 2, so electrons are held most tightly.
Question 4: Compare the ionization energies of Li, Na, and K. What trend do you observe? Explain using atomic structure.
Answer: Ionization energy decreases: Li (520) > Na (496) > K (419). As you go down the group, the valence electron is farther from the nucleus and easier to remove.
Question 5: Notice that B has lower ionization energy than Be. Propose an explanation for this exception to the general trend.
Answer: B has its outermost electron in a 2p orbital, while Be's are in 2s. The 2p orbital is slightly higher in energy and farther from the nucleus, making B's electron easier to remove.
Question 6: O has lower ionization energy than N. Based on electron configuration, why might this be?
Answer: N has a half-filled 2p subshell (2p3), which is particularly stable. O has one paired electron in its 2p4 configuration, which experiences electron-electron repulsion, making it easier to remove.
Question 7: If you were to predict the atomic radius of Rb (Period 5, Group 1), would you expect it to be larger or smaller than K? Estimate a value.
Answer: Rb should be larger than K because it is below K in Group 1. Following the pattern (increase of roughly 40-50 pm per period), Rb should be approximately 248-280 pm. (Actual value: 248 pm)
Question 8: Predict the first ionization energy of Rb compared to K. Would it be higher or lower?
Answer: Rb should have lower ionization energy than K because its valence electron is in a higher energy level (5s vs 4s), farther from the nucleus. (Actual: 403 kJ/mol)
Question 9: Based on the trends, which element would you expect to have the highest electronegativity: Na, Mg, or Cl?
Answer: Cl would have the highest electronegativity. Electronegativity increases across a period, and Cl is farthest to the right. Na and Mg are metals with low electronegativity.
Question 10: Using the data and your knowledge of periodic trends, explain why noble gases were not included in the ionization energy table.
Answer: Noble gases have very high ionization energies because they have full valence shells (stable electron configurations). They would show the highest values in each period but don't typically form ions in normal chemistry.
Conclusion Questions
- Summarize the relationship between atomic radius and ionization energy.
- Why do periodic trends exist? What is the underlying cause?
- How does this type of data analysis help prepare you for standardized tests?
Next Steps
- Practice analyzing data tables and graphs from ACT Science passages
- Review any trends that were unclear
- Move on to Common Mistakes to avoid errors on tests