Intro to Chemistry Unit 6 Review: Electronic Structure and Periodic Properties

Key Concepts and Terminology

• Atom consists of a positively charged nucleus surrounded by negatively charged electrons
• Protons positively charged particles located in the nucleus
• Neutrons neutral particles located in the nucleus
• Electrons negatively charged particles that orbit the nucleus in shells or orbitals
• Atomic number ($Z$) represents the number of protons in an atom
• Mass number ($A$) represents the total number of protons and neutrons in an atom
• Isotopes atoms of the same element with different numbers of neutrons
• Electron configuration arrangement of electrons in an atom's orbitals

Atomic Structure Basics

• Atoms are the fundamental building blocks of matter
• Atoms are composed of protons, neutrons, and electrons
• Protons and neutrons form the dense nucleus at the center of the atom
  • Protons have a positive charge ($+1$)
  • Neutrons have no charge ($0$)
• Electrons orbit the nucleus in shells or orbitals
  • Electrons have a negative charge ($-1$)
• The number of protons in an atom determines its element
• Atoms of the same element can have different numbers of neutrons (isotopes)
• Atoms are electrically neutral when the number of protons equals the number of electrons

Electron Configuration

• Electron configuration describes the arrangement of electrons in an atom's orbitals
• Electrons fill orbitals in a specific order based on energy levels
  • Aufbau principle electrons fill orbitals from lowest to highest energy
  • Pauli exclusion principle no two electrons in an atom can have the same set of four quantum numbers
  • Hund's rule electrons fill orbitals of the same energy singly before pairing
• Electron configurations are written using the format: $1s^2 2s^2 2p^6 3s^2 3p^6$
  • The number represents the principal quantum number ($n$)
  • The letter represents the orbital type ($s$, $p$, $d$, or $f$)
  • The superscript represents the number of electrons in that orbital
• Noble gas notation can be used to abbreviate electron configurations (e.g., $[Ne] 3s^2 3p^1$)

Quantum Numbers and Orbitals

• Quantum numbers describe the properties of electrons in an atom
• Four quantum numbers are used to describe an electron's state:
  1. Principal quantum number ($n$) represents the energy level or shell
  2. Angular momentum quantum number ($l$) represents the subshell or orbital type
  3. Magnetic quantum number ($m_l$) represents the orientation of the orbital in space
  4. Spin quantum number ($m_s$) represents the spin of the electron ($+\frac{1}{2}$ or $-\frac{1}{2}$)
• Orbitals are regions in space where electrons are likely to be found
• Orbital shapes correspond to the angular momentum quantum number ($l$)
  • $s$ orbitals are spherical
  • $p$ orbitals are dumbbell-shaped
  • $d$ orbitals have various shapes (cloverleaf, double dumbbell, etc.)
  • $f$ orbitals have complex shapes

Periodic Table Organization

• The periodic table organizes elements based on their atomic number and electron configuration
• Elements are arranged in rows (periods) and columns (groups)
  • Periods represent the principal quantum number ($n$) of the valence electrons
  • Groups represent the number of valence electrons and similar chemical properties
• The periodic table is divided into four blocks based on the type of orbital being filled
  • $s$-block (groups 1 and 2)
  • $p$-block (groups 13-18)
  • $d$-block (groups 3-12)
  • $f$-block (lanthanides and actinides)
• Valence electrons are the electrons in the outermost shell of an atom
  • Valence electrons determine an element's chemical properties and reactivity

Periodic Trends and Properties

• Periodic trends describe the regular patterns in properties across the periodic table
• Atomic radius generally decreases from left to right and increases from top to bottom
  • Atomic radius is the distance from the nucleus to the outermost shell of electrons
• Ionization energy generally increases from left to right and decreases from top to bottom
  • Ionization energy is the energy required to remove an electron from an atom
• Electron affinity generally increases from left to right and decreases from top to bottom
  • Electron affinity is the energy released when an atom gains an electron
• Electronegativity generally increases from left to right and decreases from top to bottom
  • Electronegativity is the ability of an atom to attract electrons in a chemical bond
• Metallic character generally decreases from left to right and increases from top to bottom
  • Metallic character refers to the tendency of an element to exhibit metallic properties (e.g., conductivity, malleability)

Applications in Chemistry

• Understanding electronic structure and periodic properties is essential for predicting chemical behavior
• Electron configuration helps predict the formation of chemical bonds
  • Atoms tend to gain, lose, or share electrons to achieve a stable noble gas configuration
• Periodic trends influence the reactivity and properties of elements
  • Electronegativity differences determine the type of chemical bond formed (ionic, covalent, or metallic)
  • Ionization energy and electron affinity affect the formation of ions and the stability of compounds
• Knowledge of electronic structure is crucial for understanding spectroscopy and analytical techniques
  • Atomic emission and absorption spectra arise from electron transitions between energy levels
• Periodic properties are used in the design and synthesis of new materials
  • Selecting elements with specific properties (e.g., conductivity, catalytic activity) based on their position in the periodic table

Practice Problems and Examples

• Write the electron configuration for the following elements:
  • Carbon (C)
  • Chlorine (Cl)
  • Iron (Fe)
• Identify the element based on the given electron configuration:
  • $1s^2 2s^2 2p^4$
  • $[Ar] 4s^2 3d^{10} 4p^5$
• Compare the atomic radii of the following pairs of elements:
  • Sodium (Na) and Potassium (K)
  • Nitrogen (N) and Phosphorus (P)
• Predict the type of bond formed between the following pairs of elements:
  • Lithium (Li) and Fluorine (F)
  • Carbon (C) and Oxygen (O)
• Calculate the number of protons, neutrons, and electrons in an atom of $^{238}U$
• Explain why the first ionization energy of Oxygen (O) is higher than that of Nitrogen (N)