Chapter 1-7: Introduction to Atomic Properties and Periodic Trends

Flashcards covering key vocabulary related to atomic structure, periodic trends, ionization energy, and electron affinity, based on lecture notes.

Van der Waals Radius

A way of defining the radius of an atom or an ion by the boundary drawn based on its outermost electrons.

Valence Shell

The outermost electron shell of an atom.

Periodic Trends

Systematic variations in properties of elements across the periodic table, explained by the location of valence electrons and the number of protons.

Effective Nuclear Charge (Z_effective)

The net positive charge experienced by an electron from the nucleus, which attracts the valence electrons and brings them closer.

Cation

An ion formed when a neutral atom loses one or more electrons, resulting in a net positive charge and typically a smaller size due to the loss of a valence shell.

Anion

An ion formed when a neutral atom gains one or more electrons, resulting in a net negative charge and typically a larger size due to electron-electron repulsion.

Atomic Size Trend (Left to Right)

Atomic size decreases across a row on the periodic table because the increasing number of protons leads to a higher effective nuclear charge, pulling electrons closer.

Atomic Size Trend (Top to Bottom)

Atomic size increases down a group on the periodic table because valence electrons are in higher n-shells, meaning they are further away from the nucleus.

Ionization Energy (IE)

The energy required to completely remove an electron from an atom.

First Ionization Energy (IE1)

The energy required to remove the first electron from an atom, usually the outermost valence electron.

Ionization Energy Trend (Left to Right)

Ionization energy generally increases across a row because a higher effective nuclear charge makes it harder to remove electrons.

Ionization Energy Trend (Top to Bottom)

Ionization energy generally decreases down a group because valence electrons are further from the nucleus, experiencing a weaker effective nuclear charge and are easier to remove.

Core Electrons

Electrons in inner shells that are much closer to the nucleus and are significantly harder to remove than valence electrons, leading to a dramatic jump in successive ionization energies.

Electron Affinity (EA)

The energy change that occurs when an electron is added to a neutral atom to form an anion; a more negative value indicates greater stability for the anion.

Electron Affinity Trend (Left to Right)

Electron affinity generally increases (becomes more negative) across a row because more protons lead to a greater effective nuclear charge, better attracting an additional electron.

Electron Affinity Trend (Top to Bottom)

Electron affinity generally decreases (becomes less negative) down a group because the valence shell is further from the nucleus, resulting in weaker attraction for an added electron.

Main Group Elements (Charge Predictability)

Elements where the number of valence electrons dictates predictable ionic charges (e.g., Group 1 is +1, Group 2 is +2) because taking away electrons leads to a full outer shell.

Transition Metals (Charge Predictability)

Elements with unpredictable ionic charges due to the involvement of d orbitals in their electron configurations.