Ionization Energy
Which element has the highest ionization energy on the periodic table?
Helium
Which element is the most electronegative, and which is the least?
Fluorine (Highest); Francium (Lowest).
As you move down a group on the periodic table, does the atomic radius increase or decrease?
Increases
Generally speaking, do metals or non-metals have a more negative (exothermic) electron affinity?
They generally release more energy when gaining an electron to reach stability.
In what corner of the periodic table (top-left, top-right, bottom-left, bottom-right) are the most metallic elements found?
Bottom-left
Explain the "Shielding Effect" and how it contributes to the trend of ionization energy decreasing down a group.
Inner electrons block the nucleus's pull on outer electrons, making them easier to remove.
Why are the Noble Gases (Group 18) generally assigned an electronegativity value of zero or not included in the standard Pauling scale?
They generally do not form bonds or "want" more electrons.
Why does the atomic radius decreases as you move from left to right across a period, despite the increasing number of electrons.
Increased nuclear charge, more protons pull the electron clouds closer to the nucleus.
While Fluorine is the most electronegative, Chlorine actually has a higher (more negative) electron affinity. Why is Fluorine's affinity lower than expected?
Fluorine is so tiny that adding an electron creates significant crowding/repulsion, slightly offsetting the energy gain.
List three physical properties that decrease as "metallic character" decreases across a period.
Luster, Malleability, Conductivity
Why is the first ionization energy of Oxygen slightly lower than that of Nitrogen, even though Oxygen is further to the right?
Oxygen has a paired electron in a p-orbital; the repulsion makes it easier to remove than Nitrogen’s stable, half-filled shell.
Based on the periodic trend, which bond would be more polar: C-H or C-O?
C-O
Arrange the following elements in order of increasing atomic radius: O, Al, K, S
O<S<Al<K
Group 2 elements (like Beryllium and Magnesium) have electron affinities that are actually positive (endothermic). Why is it energetically unfavorable for them to gain an electron?
Their s-subshells are full; a new electron must go into a higher energy p-orbital, which is energetically costly.
Silicon is considered a metalloid. Explain how its position on the periodic table explains its "semi-conductive" properties compared to Aluminum and Phosphorus.
Silicon has enough protons to hold electrons somewhat tightly (like a non-metal) but a large enough radius to allow some movement (like a metal).
An unknown element has a first ionization energy of 496kJ/mol and a second ionization energy of 4560kJ/mol. Which group on the periodic table does this element likely belong to?
Group 1
How does the "Effective Nuclear Charge" influence the increase of electronegativity across a period?
A higher effective nuclear charge exerts a stronger "tug" on shared electrons in a bond.
Which is larger: a neutral Fluorine atom or a Fluride (F-) ion?
F-, adding an electron increases repulsion between electrons, pushing them further apart.
Describe the difference between Electronegativity and Electron Affinity. (Hint: One is about lone atoms, the other is about atoms in bonds).
Electron Affinity is the energy change for a single isolated atom; Electronegativity is a relative scale for atoms in a bond.
How does the trend of metallic character relate to the trend of oxides?
Metallic oxides tend to be basic, while non-metallic oxides tend to be acidic.
Why does the third ionization energy of Magnesium (Mg) see a massive, non-linear jump compared to its first and second ionization energies?
Core electron removal,The first two electrons are valence; the third must be pulled from a stable noble gas core (2p6), requiring massive energy.
Elements like Cesium and Francium have the lowest electronegativity. In terms of atomic structure, why do these atoms "want" to give away electrons rather than attract them?
Their valence electrons are so far from the nucleus that they have very little "grip" on them.
While atomic radius generally decreases across a period, which specific group of elements often shows a slight "rebound" or increase in measured radius due to their full valence shells and lack of bonding?
Noble Gases
Rank these elements from most to least exothermic electron affinity: Li, C, N, F.
F > C > Li > N
Why does metallic character increase as you move down a group, even though the number of protons in the nucleus is increasing significantly?
Low Ionization Energy: Even with more protons, the "shielding" from many inner shells makes it very easy for the atom to lose its outer electrons, which defines metallic behavior.