Chemistry Unit 3 - Electron Config.
The Periodic Table and Atomic Structure
Introduction to the Periodic Table
Historical Development
Dimitri Mendeleev (1869): Developed the first periodic table, organizing elements by atomic mass and noting periodic trends in their properties.
Henry Moseley: Revised Mendeleev's table by arranging elements according to their atomic number, which more accurately reflected the observed periodic relationships in physical and chemical properties.
Periodic Table Organization
Groups (Columns): Vertical columns on the periodic table.
Elements within the same group share similar valence electron configurations, meaning they have the same number of valence electrons.
Key Groups:
Group 1: Alkali metals (reactive)
Group 2: Alkaline earth metals (reactive)
Groups 3-12: Transition metals
Group 15: Nitrogen family
Group 16: Oxygen family (reactive)
Group 17: Halogens (very reactive)
Group 18: Noble gases (inert/non-reactive)
The group number often corresponds to the number of valence electrons (with exceptions in the transition metals).
Periods (Rows): Horizontal rows on the periodic table.
The period number corresponds to the principal quantum number (energy level) of the valence electrons.
Element Classifications
Metals:
Good conductors of heat and electricity.
Shiny luster.
Primarily solid at room temperature (mercury is the exception).
Ductile (can be drawn into wires).
Malleable (can be hammered into sheets).
Nonmetals:
Poor conductors of heat and electricity.
Many exist as gases naturally.
Solid forms are typically brittle.
Metalloids:
Exhibit properties of both metals and nonmetals.
Atomic Models
Bohr Model
Electron Path: Electrons orbit the nucleus in fixed, circular paths (energy levels).
Limitation: Primarily accurate only for hydrogen (one electron); fails to describe heavier elements accurately.
Wave Mechanical Model (Quantum Model)
Electron Path: Electrons do not follow fixed orbits. Instead, they occupy orbitals, which are three-dimensional regions of space where there is a high probability of finding an electron.
Uncertainty: This model aligns with Heisenberg's Uncertainty Principle, stating that one cannot simultaneously know both the exact position and momentum of a particle.
Electron Configuration
Principles of Electron Configuration
Electron Configuration: The arrangement of electrons within an atom's orbitals.
Aufbau Principle: Electrons fill the lowest energy levels and orbitals first before occupying higher energy levels.
Pauli Exclusion Principle: Each orbital can hold a maximum of two electrons, and these two electrons must have opposite spins (e.g., one clockwise, one counterclockwise).
Hund's Rule: Within a sublevel (like pp, dd, or ff), each orbital must contain at least one electron before any orbital can contain a second electron. All singly occupied orbitals within a sublevel have the same electron spin.
Representing Electron Configuration
Full Electron Configuration: Shows the distribution of all electrons in an atom.
Example: Neon (Ne) has the electron configuration 1s22s22p61s22s22p6.
Noble Gas Configuration (Shorthand): Uses the preceding noble gas in the periodic table to represent the core electrons, followed by the configuration of the valence electrons.
Example: Sodium (Na) has an electron configuration of 1s22s22p63s11s22s22p63s1. Since Neon (Ne) has the configuration 1s22s22p61s22s22p6, the noble gas configuration for Sodium is [Ne]3s1[Ne]3s1.
Example: Tin (Sn) has the electron configuration 1s22s22p63s23p64s23d104p65s24d105p21s22s22p63s23p64s23d104p65s24d105p2. The preceding noble gas is Krypton (Kr), which has the configuration 1s22s22p63s23p64s23d104p61s22s22p63s23p64s23d104p6. Therefore, the noble gas configuration for Tin is [Kr]5s24d105p2[Kr]5s24d105p2.
Valence and Core Electrons
Definitions
Valence Electrons: Electrons in the outermost principal energy level (highest energy level) of an atom. These are the electrons involved in chemical bonding. The highest principal energy level is indicated by the largest principal quantum number (coefficient) in the electron configuration.
Example: For Sodium (Na) with configuration 1s22s22p63s11s22s22p63s1, the highest principal energy level is 3, and there is 1 electron in that level (3s13s1). Thus, Sodium has 1 valence electron.
Example: For Tin (Sn) with configuration 1s22s22p63s23p64s23d104p65s24d105p21s22s22p63s23p64s23d104p65s24d105p2, the highest principal energy level is 5, and there are 4 electrons in that level (5s25s2 and 5p25p2). Thus, Tin has 4 valence electrons.
Core Electrons: The inner electrons that are not in the outermost principal energy level. They are not involved in chemical bonding.
Example: For Sodium (Na) $1s22s22p6∣3s11s22s22p6∣3s1$, the core electrons are 1s22s22p61s22s22p6.
Example: For Tin (Sn) $1s22s22p63s23p64s23d104p6∣5s24d105p21s22s22p63s23p64s23d104p6∣5s24d105p2$, the core electrons are 1s22s22p63s23p64s23d104p61s22s22p63s23p64s23d104p6.
Achieving Stability: Ionic Charges
The Octet Rule
Atoms strive to achieve a full valence shell, which typically means having 8 valence electrons (like noble gases), for stability.
Atoms will gain or lose valence electrons to achieve this stable configuration, whichever requires less energy.
Predicting Ionic Charges
The number of valence electrons an atom has dictates how it will achieve a stable electron configuration and, therefore, its likely ionic charge.
General Trends by Group (for main group elements):
Group 1 (1 valence electron): Tends to lose 1 electron to form a +1 ion.
Group 2 (2 valence electrons): Tends to lose 2 electrons to form a +2 ion.
Group 13 (3 valence electrons): Tends to lose 3 electrons to form a +3 ion.
Group 14 (4 valence electrons): Can lose 4 electrons (+4) or gain 4 electrons (-4).
Group 15 (5 valence electrons): Tends to gain 3 electrons to form a -3 ion.
Group 16 (6 valence electrons): Tends to gain 2 electrons to form a -2 ion.
Group 17 (7 valence electrons): Tends to gain 1 electron to form a -1 ion.
Group 18 (8 valence electrons): Already stable (inert), so typically form no ions (0 charge).
Isoelectronic Species
Definition
Isoelectronic: Atoms or ions that have the same electron configuration.
Example
Consider Oxygen (O).
Oxygen's electron configuration is 1s22s22p41s22s22p4. It has 8 electrons.
To achieve a stable octet, Oxygen tends to gain 2 electrons, forming the oxide ion, O2−O2−.
The electron configuration of O2−O2− is 1s22s22p61s22s22p6. This configuration has 10 electrons.
The noble gas Neon (Ne) has the electron configuration 1s22s22p61s22s22p6.
Therefore, the O2−O2− ion and a Neon atom are isoelectronic because they share the same electron configuration (1s22s22p61s22s22p6).