Atomic and ionic radii, ionization energy, electron affinity

Atomic and Ionic Radii

1. The atomic radius is the distance from the nucleus to the outermost shell of an atom.
2. Atomic radius decreases across a period due to increased nuclear charge, which pulls electrons closer to the nucleus.
3. Atomic radius increases down a group because of the addition of new electron shells.
4. Cations (positively charged ions) have smaller radii than their parent atoms due to the loss of electrons and reduced electron-electron repulsion.
5. Anions (negatively charged ions) have larger radii than their parent atoms because of increased electron repulsion.
6. The size of atoms and ions influences chemical reactivity and bonding properties.
7. The ionic radius follows the same periodic trend as atomic radius, depending on the ion's charge and size.

Ionization Energy

1. Ionization energy is the energy required to remove an electron from a neutral atom in its gaseous state.
2. It is expressed in units of kilojoules per mole (kJ/mol).
3. Ionization energy increases across a period due to higher nuclear charge, which makes it harder to remove an electron.
4. Ionization energy decreases down a group as the atomic size increases, making it easier to remove outer electrons.
5. Successive ionization energies (1st, 2nd, 3rd, etc.) increase because removing additional electrons requires overcoming a greater attraction to the nucleus.
6. Elements with a stable noble gas configuration have very high ionization energies.
7. Alkali metals (Group 1) have the lowest ionization energy in their respective periods, making them highly reactive.
8. Noble gases (Group 18) have the highest ionization energy due to their full valence shells.

Electron Affinity

1. Electron affinity is the energy change that occurs when an electron is added to a neutral atom in its gaseous state.
2. A positive electron affinity indicates that energy is released when an electron is added, making the process exothermic.
3. A negative electron affinity suggests that energy is absorbed, making the process endothermic.
4. Electron affinity generally increases across a period due to increased nuclear charge, which attracts additional electrons more strongly.
5. Electron affinity decreases down a group because of the increased distance between the nucleus and the incoming electron.
6. Halogens (Group 17) have the highest electron affinities because they are one electron away from a stable noble gas configuration.
7. Noble gases have very low or negligible electron affinities because their valence shells are full.
8. Electron affinity values are influenced by atomic size, nuclear charge, and electronic configuration.

Key Comparisons and Trends

1. Atomic radius: Decreases across a period, increases down a group.
2. Ionic radius: Smaller for cations, larger for anions compared to their parent atoms.
3. Ionization energy: Increases across a period, decreases down a group.
4. Electron affinity: Increases across a period, decreases down a group.
5. These trends help explain the reactivity and bonding behavior of elements.

Applications of Periodic Trends

1. Understanding periodic trends aids in predicting chemical reactions and the formation of compounds.
2. Explains why metals tend to lose electrons and non-metals tend to gain electrons.
3. Helps in determining the stability of ions and the strength of chemical bonds.
4. Plays a critical role in the design of new materials and pharmaceutical compounds.

Key Points

1. Trends in atomic and ionic radii explain the reactivity of alkali metals and halogens.
2. Ionization energy is highest for noble gases and lowest for alkali metals.
3. Elements with high electron affinity, like fluorine, form strong negative ions.
4. Periodic trends are essential for understanding the properties of elements and their compounds.