Ions, Electron Configurations, and Periodic Trends: Study Guide for GOB Chemistry

Ions and Ion Formation

Definition and Types of Ions

Atoms form ions by gaining or losing electrons to achieve a stable electron configuration, often resembling that of noble gases. This process is fundamental in understanding chemical bonding and reactivity.

• Cation: A positively charged ion formed when an atom loses one or more electrons (typically metals).

• Anion: A negatively charged ion formed when an atom gains one or more electrons (typically non-metals).

• Isoelectronic: Species that have the same number of electrons and thus the same electron configuration.

Example: Sodium (Na) loses one electron to form Na+, which is isoelectronic with neon (Ne).

Ion Formation Diagrams

Visual representations show how atoms lose or gain electrons to form cations or anions. For example, boron (B) can lose electrons to form B+, while oxygen (O) can gain electrons to form O2−.

Practice Problems

• Determine the number of protons, neutrons, and electrons for given isotopes and ions.

• Identify pairs of species that are both isoelectronic and isotopic.

Table: Isotope and Ion Data

Symbol	Protons	Neutrons	Electrons	Mass Number	Net Charge
Co2+	27	32	25	59	2+
Se	34	46	36	80	2−
Sn	50	70	48	120	2+
Xe	54	78	54	132	0

Octet Rule and Electron Configurations

Octet Rule

Main group elements tend to gain, lose, or share electrons to achieve a full outer shell of eight electrons (octet), leading to greater stability.

• Metals lose electrons to achieve the electron configuration of the nearest noble gas.

• Non-metals gain electrons to achieve a noble gas configuration.

Example: Magnesium (Mg, atomic number 12) loses two electrons to form Mg2+ with a filled outer shell.

Electron Configurations for Ions

Electron configuration describes the arrangement of electrons in an atom or ion. For ions, electrons are added (anions) or removed (cations) from the outermost shell.

• For cations: Remove electrons from the highest principal quantum number (n).

• For anions: Add electrons to the next available orbital.

Example: The electron configuration for Na+ is $1s^2 2s^2 2p^6$.

Steps for Writing Electron Configurations

1. Write the electron configuration for the neutral atom.

2. Add or remove electrons according to the ion's charge.

Example: For N3−, start with nitrogen's configuration and add three electrons: $1s^2 2s^2 2p^6$.

Main Group and Transition Elements

Main Group Elements

Main group elements (Groups 1A–8A) have predictable charges based on their group number. The atomic number equals the number of protons in the nucleus.

• Group 1A: +1 charge (e.g., Na+)

• Group 2A: +2 charge (e.g., Mg2+)

• Group 7A: −1 charge (e.g., Cl−)

Exceptions: Lead (Pb) and Tin (Sn) can have variable charges. Heavy metals like Bismuth (Bi), Polonium (Po), and elements with atomic numbers 114–118 also have variable charges.

Transition Metals (Type II Metals)

Transition metals (Groups 3B–2B) can form multiple positive charges due to their electron arrangements. Their charges are less predictable and often require specification (e.g., Fe2+ vs. Fe3+).

Example: Cadmium (Cd) typically forms a +2 ion: Cd2+.

Periodic Trends: Ionization Energy and Electron Affinity

Ionization Energy (IE)

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. It is a key periodic trend that increases across a period (left to right) and decreases down a group (top to bottom).

• High IE: Electrons are not easily lost (e.g., noble gases).

• Low IE: Electrons are easily lost (e.g., alkali metals).

Equation: $\text{Ionization Energy} = N(g) \rightarrow N^+(g) + e^-$

Example: Fluorine (F) has a higher ionization energy than sodium (Na).

Electron Affinity (EA)

Electron affinity is the energy released when an electron is added to a gaseous atom or ion. It generally becomes more negative (more energy released) across a period and less negative down a group.

• High EA: Atoms readily accept electrons (e.g., halogens).

• Low EA: Atoms do not readily accept electrons (e.g., noble gases).

Equation: $C(g) + e^- \rightarrow C^-(g) + \text{Energy}$

Example: Chlorine (Cl) releases more energy upon gaining an electron than sodium (Na).

Summary Table: Periodic Trends

Trend	Across a Period (→)	Down a Group (↓)
Ionization Energy	Increases	Decreases
Electron Affinity	Becomes more negative	Becomes less negative

Practice and Application

• Identify which elements or ions are isoelectronic or isotopic.

• Predict the charge of ions formed by main group and transition elements.

• Rank elements by ionization energy and electron affinity.

• Write electron configurations for ions.

Example: Arrange F, Li, C, O, Be from greatest to least tendency to accept an electron: F > O > C > Be > Li.

Additional info: These notes are based on standard GOB Chemistry curriculum and include inferred explanations and tables for clarity and completeness.