While still analyzing the graph, make predictions about the trend of ionization energy between atoms in the same group on the periodic table. For example, do atoms with larger atomic numbers have greater ionization energy than atoms with smaller atomic numbers in the same group? The LD mac nemst nijl Fo< appens Kje 0nt A’ ener9 | Navigate back to the main page and reset the data using the reset symbol. Now choose two elements that are in the same group. How do their ionization energy values compare? Does this data support your prediction for part h? si99kJlmgl Na 495, ejimo| but 9rcete 7oniza+ Lithwvm hat Smalle astomic Tnis 015 Ptuve [ Pred: cti0n Based on what you have learned and without the assistance of the periodic trends simulation, organize the following lists of atoms from lowest ionization energy to highest ionization energy: Na, Al, Ac, Al, Ar, F, Br, Cl, I, C, F, O, Si, Mg, He, Rb, Mg, S, O. Using the simulation, check your predicted answers to see if you are correct. Reflecting on what you have learned about both atomic radius and ionization energy at this point, which of the following statements best describe these trends? Atoms that have large atomic radii also have large values of ionization energy. Atoms that have small atomic radii will have large values of ionization energy. Explain your choice referencing both the atomic model and subatomic particles.
The Correct Answer and Explanation is:
To predict the trend of ionization energy among elements in the same group, let’s first understand the behavior of ionization energy across the periodic table. Ionization energy is the amount of energy required to remove an electron from a gaseous atom.
Trend in Ionization Energy Within a Group:
- General Trend: In a group (vertical column) of the periodic table, ionization energy decreases as you go down. This is because, as the atomic number increases, the number of electron shells increases, meaning the outermost electrons are farther away from the nucleus and are shielded more by inner electrons. As a result, these outer electrons experience a weaker force of attraction from the nucleus, making them easier to remove (lower ionization energy). So, for example, in the alkali metals group, sodium (Na) has a lower ionization energy than lithium (Li), even though sodium has a larger atomic number. This trend is consistent across other groups as well.
Prediction:
- As you go down a group in the periodic table, the ionization energy will decrease because of the increasing distance between the nucleus and the outer electrons and greater electron shielding.
- This is why sodium (Na), which is in the same group as lithium (Li), has a lower ionization energy compared to lithium.
Comparison of Ionization Energies (Example from Group 1):
- Sodium (Na): Ionization energy = 495 kJ/mol
- Lithium (Li): Ionization energy = 520 kJ/mol
Prediction:
Lithium has a higher ionization energy than sodium, which aligns with the general trend of decreasing ionization energy as you go down a group.
Organizing Atoms by Ionization Energy:
Let’s organize the atoms from lowest to highest ionization energy, considering the periodic trends:
- Rb (lowest ionization energy – alkali metal, farthest from the nucleus, largest atomic radius)
- Na (alkali metal, similar to Rb but higher)
- Mg (alkaline earth metal, slightly smaller atomic radius)
- Al (Group 13, with slightly higher ionization energy than alkali and alkaline earth metals)
- Si (Group 14)
- C (Group 14, smaller atomic radius)
- O (Group 16, high ionization energy)
- F (Group 17, very high ionization energy)
- Ar (Group 18 noble gas, highest among the elements in the same period)
So, Rb has the lowest ionization energy and Ar has the highest in this list.
Explanation of the Atomic Model and Trends:
- Atomic Radius: Atoms with large atomic radii (e.g., alkali metals like Rb and Na) have low ionization energies. The outermost electron is farther from the nucleus and experiences more shielding, making it easier to remove.
- Small Atomic Radius: Atoms with small atomic radii (e.g., noble gases like He, Ar) have high ionization energies. The outermost electron is closer to the nucleus and experiences a stronger force of attraction, requiring more energy to remove it.
Thus, atoms with small atomic radii tend to have high ionization energies, because the force of attraction between the nucleus and the valence electrons is stronger, making it more difficult to remove electrons.
