Electron Configurations

1. Electron Configurations

2. Q3. What is the diagonal effect? • Atoms on the diagonal down and to the right have similar properties, due to similar ionization energies and atomic radii • Atoms down and to the left have similar properties due to similar ionization energies and atomic radii • Atoms on the diagonal down and to the right have similar properties. due to similar electron configurations. • Atoms down and to the left have similar properties. Due to similar electron configurations. • Q4. Which statement is true? • Electronegativity is the ability of an atom to attract an electron. • Electron affinity is the ability of an atom to attract an electron. • Electron affinity is the ability of an atom to attract electron density in a bond • Electron affinity and electronegativity are synonomous. • A and C are both true Q1: Rank the following in order of increasing radius: Li, C, F C<F<Li F<C<Li Li<C<F Li<F<C Not enough information. Q2: Rank the following in order of increasing radius: Cu+, Cu2+, O2- Cu+< Cu2+<O2- Cu2+<Cu+<O2- O2-<Cu+<Cu2+ O2-<Cu2+<Cu+ Cu2+<O2-<Cu+

3. Electronegativity Effective Nuclear Charge Electronegativity

4. Explain why atomic radius is the reverse of each of the other trends. • The trends are based on the effective nuclear charge (horizontal) and shielding/distance from the nucleus (vertical). • Since effective nuclear charge increases as you go across the periodic table, • the amount that it will attract an electron (electron affinity) increases, • the amount of energy it takes to pull off an electron (ionization energy) increases • and the amount of electron density the nucleus will try to attract in a bond (electronegativity) will all increase. • However this tighter control over its own, and other atom’s electrons also mean that it will hold its own electrons closer. Making the radius smaller as you go across the periodic table. • Since shielding and distance from the nucleus increase as you go down the periodic table • the amount that it will attract an electron (electron affinity) decreases, • the amount of energy it takes to pull off an electron (ionization energy) decreases • and the amount of electron density the nucleus will try to attract in a bond (electronegativity) will all decrease. • However this looser control over its own, and other atom’s electrons also mean that it won’t hold its own electrons as tightly. Making the radius larger as you go down the periodic table.

5. Ionization Energy C: [He]2s22p2 F • State the p-block exceptions for ionization energy and electron affinity. (remember this is homework feel free to look it up if need be): • Carbon/Nitrogen/Oxygen and Silicon/Sulfur/Phosphorous and Germanium/Arsenic/Selenium. For Example: • for ionization energy C<O<N • for electron affinity N<C<O • b) Explain why ionization energy and electron affinity have these exceptions, while atomic radius and electronegativity do not. • Ionization energy and electron affinity both are based on adding or removing electrons so stable electron configurations change them. • Atomic radius and electronegativity are not electron configuration dependent. N N: [He]2s22p3 Cl O O: [He]2s22p4 P C Be Mg S B Al Electron Affinity