Question

directions: use your notes to answer the following questions.

1. rank the following elements by increasing atomic radius: carbon, aluminum, oxygen, potassium.
2. rank the following elements by increasing electronegativity: sulfur, oxygen, neon, aluminum.
3. why does fluorine have a higher ionization energy than iodine?
4. why do elements in the same family generally have similar properties?
5. indicate whether the following properties increase or decrease from left to right across the periodic table.

a. atomic radius (excluding noble gases)
b. first ionization energy
c. electronegativity

1. what trend in atomic radius occurs across the periodic table? what causes this trend?
2. what trend in ionization energy occurs across a period on the periodic table? what causes this trend?
3. circle the atom in each pair that has the largest radius.

a. al or b
b. na or al
c. s or o
d. o or f
e. br or cl
f. mg or ca

Explanation:

1. Atomic Radius Ranking:

• Atomic radius generally increases down a group and decreases across a period.
• Oxygen (O) is in period 2, carbon (C) is also in period 2 but to the left of O (atomic radius decreases left - to - right in a period).
• Aluminum (Al) is in period 3, and potassium (K) is in period 4.

1. Electronegativity Ranking:

• Noble gases (like neon, Ne) have very low electronegativity as they have a full valence shell and are not likely to gain electrons.
• Electronegativity increases across a period. Aluminum (Al) is a metal with low electronegativity, sulfur (S) is more electronegative than Al, and oxygen (O) is more electronegative than S.

1. Ionization Energy of F vs I:

• Ionization energy is the energy required to remove an electron.
• Fluorine (F) has a smaller atomic radius. The electrons are closer to the nucleus and more strongly held. Iodine (I) has a larger atomic radius, so the outermost electrons are less strongly held.

1. Similar Properties in a Family:

• Elements in the same family (group) have the same number of valence electrons. Valence electrons determine an element's chemical properties.

1. Periodic Table Trends:

• Atomic Radius: Excluding noble gases, as we move from left to right across a period, the number of protons increases, pulling the electrons closer (decrease in atomic radius).
• First Ionization Energy: Increases from left to right as it becomes harder to remove an electron (due to increased nuclear charge and smaller atomic radius).
• Electronegativity: Increases from left to right as elements have a greater tendency to attract electrons.

1. Atomic Radius Across Periodic Table:

• Atomic radius decreases from left to right across a period. Caused by increasing nuclear charge (number of protons) which pulls the electrons closer.

1. Ionization Energy Across a Period:

• Ionization energy increases from left to right. Due to smaller atomic radius (electrons closer to nucleus, more energy required to remove them) and increased nuclear charge.

1. Largest Radius in Pairs:

• a. Al or B: Al is in period 3, B in period 2. Larger period → larger radius (Al).
• b. Na or Al: Na is to the left of Al in period 3. Left - side elements in a period have larger radius (Na).
• c. S or O: S is in period 3, O in period 2 (S).
• d. O or F: O is to the left of F in period 2 (O).
• e. Br or Cl: Br is in period 4, Cl in period 3 (Br).
• f. Mg or Ca: Ca is in period 4, Mg in period 3 (Ca).

Answer:

1. \(O < C < Al < K\)
2. \(Ne < Al < S < O\)
3. Fluorine has a smaller atomic radius (electrons closer to nucleus, more strongly held) compared to iodine.
4. Same number of valence electrons.

5.

• a. Decrease
• b. Increase
• c. Increase

1. Decreases from left to right. Caused by increasing nuclear charge pulling electrons closer.
2. Increases from left to right. Due to smaller atomic radius and increased nuclear charge.

8.

• a. \(Al\)
• b. \(Na\)
• c. \(S\)
• d. \(O\)
• e. \(Br\)
• f. \(Ca\)