Referring to a periodic table, arrange the following atoms in order of increasing atomic radius, ionization energy and electronegativity: P; Si, N?
The Correct Answer and Explanation is:
To arrange the atoms phosphorus (P), silicon (Si), and nitrogen (N) in terms of increasing atomic radius, ionization energy, and electronegativity, we need to understand periodic trends.
1. Atomic Radius:
The atomic radius is the distance from the nucleus to the outermost electron. As you move across a period (from left to right), the atomic radius decreases because the nuclear charge increases, pulling the electrons closer to the nucleus. As you move down a group (from top to bottom), the atomic radius increases because additional electron shells are added, making the atom larger.
Order of increasing atomic radius:
- N < P < Si
Explanation:
- Nitrogen (N) is in period 2, group 15, and has a smaller atomic radius compared to phosphorus and silicon due to having fewer electron shells.
- Phosphorus (P) is in period 3, group 15, and is larger than nitrogen because it has an additional electron shell.
- Silicon (Si), although also in period 3, is in group 14 and has a larger atomic radius than phosphorus, as it has fewer protons, leading to a weaker attraction between the nucleus and the outermost electrons.
2. Ionization Energy:
Ionization energy refers to the energy required to remove an electron from an atom. Ionization energy increases across a period because the nuclear charge increases, making it harder to remove electrons. It decreases down a group because the outer electrons are farther from the nucleus and more shielded.
Order of increasing ionization energy:
- Si < P < N
Explanation:
- Silicon (Si) has the lowest ionization energy because it is the farthest to the left in the periodic table, so its outer electrons are farther from the nucleus and are more easily removed.
- Phosphorus (P) has a higher ionization energy because it is further to the right, with a stronger pull from the nucleus on its outer electrons.
- Nitrogen (N) has the highest ionization energy because it is the furthest right, with a stronger nuclear charge and fewer electron shells, making it more difficult to remove an electron.
3. Electronegativity:
Electronegativity is the tendency of an atom to attract electrons in a bond. It increases across a period and decreases down a group.
Order of increasing electronegativity:
- Si < P < N
Explanation:
- Silicon (Si) has the lowest electronegativity because it is in group 14, which generally has lower electronegativity values.
- Phosphorus (P) has a higher electronegativity than silicon but is still lower than nitrogen, as it is in group 15.
- Nitrogen (N) has the highest electronegativity because it is in group 15, period 2, and tends to attract electrons more strongly than both silicon and phosphorus.
In summary:
- Atomic radius: N < P < Si
- Ionization energy: Si < P < N
- Electronegativity: Si < P < N
