Flashcards for:
Understanding Periodic Trends in Chemistry

The atomic radius decreases across a period because the increasing number of protons increases the effective nuclear charge, pulling electrons closer to the nucleus.

First ionization energy is the energy required to remove the first electron from an atom.

The atomic radius increases down a group due to the addition of new energy levels (or shells), which increase the size of the atom.

Ionization energy increases across a period due to higher effective nuclear charge and smaller atomic radius, which increases the force needed to remove an electron.

Ionization energy decreases down a group because the larger atomic radius increases the distance between nucleus and outer electrons, reducing the force and energy required to remove an electron.

Electron affinity is the amount of energy released when an atom gains an electron.

Fluorine has the highest electronegativity, with a value of 4.0 on the Pauling scale.

Electron affinity increases across a period as atoms become more stable upon gaining electrons, releasing more energy, especially in elements like fluorine.

Electronegativity increases across a period because the higher effective nuclear charge leads to stronger attraction of electrons.

Electronegativity is undefined for noble gases because their electron configuration is full and they do not need to attract more electrons.

The atomic radius increases down a group due to the addition of new electron shells leading to a larger distance from the nucleus.

Second ionization energy is the energy needed to remove the second electron from an atom.

An increase in effective nuclear charge across a period results in greater electronegativity.

Electron affinity decreases down a group because a larger atomic radius leads to less stable negative ions, causing the reactions to be more endothermic.

The distance from the center of the nucleus to the outermost electron shell.