Trends in Periodic Table

Fill-in-the-blank practice flashcards covering periodic trends, shielding effect, effective nuclear charge, atomic and ionic radii, ionization energy, electronegativity, electron affinity, and Group exceptions based on lecture notes.

The ability of electrons in inner orbitals to reduce electrostatic attraction between protons in the nucleus and outer electrons is known as the __________.

shielding effect

Electrons in the same energy level provide __________ shielding effect from each other.

minimal

The equation for effective nuclear charge is $Z_{eff} = Z - \frac{\text{shielding constant}}{\text{atomic number}}$, where $\text{shielding constant}$ is represented by the variable __________.

$$\sigma$$

Atomic radius is defined as __________ the distance between the nuclei of two identical, adjacent atoms.

$$\frac{1}{2}$$

Moving from top to bottom down a group, atomic radius __________ as the main energy level ($n$) increases.

increases

Moving from left to right across a period, atomic radius __________ due to an increase in effective nuclear charge.

decreases

Across a period, transition elements show a __________ decrease in atomic radius because $3d$ electrons shield outer $4s$ electrons more effectively.

slight

Formation of cations involves losing electrons, which leads to less shielding effect, an increased effective nuclear charge, and a __________ in atomic radii.

decrease

Formation of anions involves gaining electrons, which increases shielding effect, decreases effective nuclear charge, and causes an __________ in atomic radii.

increase

In an isoelectronic series, species with more protons have a __________ ionic radius.

smaller

When comparing $K^+$ ($19 - 1 = 18$ electrons) and $S^{2-}$ ($16 + 2 = 18$ electrons), __________ have greater radii than cations.

anions

The energy required to remove electrons from an atom in its gaseous state is called __________.

ionization energy

Removing a second electron requires more energy than removing the first, meaning second ionization energy ($IE_2$) is __________ than first ionization energy ($IE_1$).

greater

Moving from top to bottom down a group, ionization energy __________ because outer electrons are farther from the nucleus.

decreases

Moving from left to right across a period, ionization energy __________ due to increased effective nuclear charge and decreased atomic radius.

increases

Group 3 elements require less ionization energy than Group 2 elements because the electron being removed resides in a valence __________-subshell.

$$p$$

Group 16 elements require less energy to remove an electron than Group 15 elements because Group 16 contains __________ electrons that cause greater electrostatic repulsion.

paired

The ability of an atom to attract shared electrons in a chemical bond is called __________.

electronegativity

Electronegativity does not apply to noble gases because they are already stable and have __________ attraction for shared electrons.

no

The energy change that occurs when an atom or ion in the gaseous state gains an electron is defined as __________.

electron affinity

High electronegativity is characterized by high ionization energy and a highly __________ electron affinity.

negative

Moving from top to bottom down a group, electronegativity __________ because outer electrons are farther from the nucleus.

decreases

Moving from left to right across a period, electronegativity __________ because an increased effective nuclear charge attracts bonding pairs of electrons more strongly.

increases

In Group 13, Gallium has a higher electronegativity than Aluminum because its filled $3d$ subshell electrons are bad at __________ the valence electrons from the nucleus.

shielding

When moving from Indium to Thallium in Group 13, the valence shell $4f$ subshell provides __________ shielding than the $d$ subshell, leading to a high effective nuclear charge and stronger electronegativity.

worst