Chapter 2

Chapter 2 - Issues to Address

• Promotion of Bonding: Understanding what factors contribute to atomic bonding.

• Types of Bonds: Identification and categorization of various types of chemical bonds.

• Properties Inferred from Bonding: Analysis of what physical and chemical properties can be determined based on bond types.

Chapter 2 - Atomic Structure

• Definition of Atom: The fundamental building block of matter.

  • Electrons: Mass of $9.11 imes 10^{-31} ext{ kg}$

  • Protons and Neutrons: Subatomic particles that make up the atomic nucleus.

  • Z (Atomic Number): Represents the number of protons in the nucleus of an atom, equal to the number of electrons in a neutral atom.

  • A (Atomic Mass Unit): Defined as $ext{amu} = rac{1}{12} ext{ mass of } ^{12} ext{C}$; 1 amu is approximately $1.67 imes 10^{-27} ext{ kg}$.

• Atomic Weight: Represents the weight of $6.022 imes 10^{23}$ atoms or molecules.

Concept Check 2.1

• **Atomic Weight Non-Integer Explanation:

  1. Atomic weights represent a weighted average of the atomic masses of naturally occurring isotopes.

  2. The mass of an atom is slightly different from the sum of its individual components due to binding energy mass loss.

• Example: Only Carbon-12 has an integer atomic mass by definition.

Chapter 2 - Isotopes of Chromium

• Chromium Isotopes and Atomic Weights:

  • 4.34% of $^{50} ext{Cr}$ with an atomic weight of 49.9460 amu

  • 83.79% of $^{52} ext{Cr}$ with an atomic weight of 51.9405 amu

  • 9.50% of $^{53} ext{Cr}$ with an atomic weight of 52.9407 amu

  • 2.37% of $^{54} ext{Cr}$ with an atomic weight of 53.9389 amu

• Average Atomic Weight Calculation: Confirmed average atomic weight of Cr = 51.9963 amu.

Atomic Structure Continued

• Valence Electrons: Key to determining properties such as chemical, electrical, thermal, and optical behaviors.

Chapter 2 - Electronic Structure

• Nature of Electrons: Electrons exhibit both wavelike and particulate properties, existing in orbitals defined by probability.

• Energy Levels:

  • Quantum Numbers:

  • n (Principal Quantum Number): Defines the energy level or shell (K, L, M, N, O = 1, 2, 3, etc.)

  • l (Subsidiary Quantum Number): Defines subshells (s, p, d, f).

  • ml (Magnetic Quantum Number): Defines energy states within a magnetic field, ranges from $-l$ to $+l$.

  • ms (Spin Quantum Number): Represents the spin of the electron: $+ rac{1}{2}$ or $- rac{1}{2}$.

• Energy States: Electrons occupy discrete energy states and tend to fill the lowest available energy state.

• Electron Configurations: Valence electrons in unfilled energy levels are crucial for chemical bonding, e.g., for Carbon: $1s^2 2s^2 2p^2$.

Chapter 2 - The Periodic Table

• Periodic Trends: Columns in the periodic table represent elements with similar valence structures, affecting their bonding characteristics.

• Electropositive and Electronegative Elements:

  • Electropositive elements readily donate electrons to become positively charged ions.

  • Electronegative elements readily acquire electrons to become negatively charged ions.

Bonding Forces

• Attractive and Repulsive Forces: The force of bonding can be expressed as:

  • $F<em>N = F</em>A + F_R$

  • Where $F<em>N = 0$ at equilibrium $r = r</em>0$.

• Bonding Energy Dynamics:

  • Bonding energy (E$_0$) represents the minimum energy at the most stable configuration (equilibrium), balancing attractive and repulsive interactions.

Properties Derived from Bonding

• Trends in Bonding Energy:

  • Bond length $r$ and energy $(E<em>0)$ are inversely related to melting temperature $(T</em>m)$; higher $(E_0)$ correlates with higher melting temperatures.

• Coefficient of Thermal Expansion (α):

  • Defined as $rac{ΔL}{L<em>0(T</em>2-T<em>1)}$; larger $(E</em>0)$ leads to smaller α values.

Primary Interatomic Bonds

• Types of Chemical Bonds in Solids:

  • Ionic Bonds: Formed between metals and nonmetals involving electron transfer.

  • Covalent Bonds: Formed by sharing of electron pairs; characterized by specific directional bonding.

  • Metallic Bonds: Involve a delocalized cloud of electrons surrounding positively charged ion cores.

Ionic Bonding

• Nature of Ionic Bonds:

  • Ionic bonds occur between positively and negatively charged ions requiring significant differences in electronegativity.

• Example - NaCl:

  • Sodium donates an electron to chlorine resulting in Na+ and Cl- ions, stabilized by Coulombic attraction.

Covalent Bonding

• Characteristics of Covalent Bonds:

  • Form between atoms with similar electronegativities by sharing electrons; bonds are directional, limiting bonding atom number.

• Example: In methane (CH4), carbon shares electrons with four hydrogen atoms.

Ionic-Covalent Mixed Bonding

• Percent Ionic Character Calculation: The ionic character of the bond can be assessed with the equation: ext{% Ionic Character} = | rac{XA - XB}{1}| imes 100

  • Where $XA$ and $XB$ are the electronegativities of element A and B respectively.

Bonding Energy Comparisons

• Types of Bonding and Associated Energies:

  • Ionic bonding: High energy, low thermal expansion.

  • Covalent bonding: Variable energy depending on molecule structure, generally lower density.

  • Metallic bonding: Also high energy with variable properties.

Summary of Bonding Types

• Ceramics: Characterized by ionic and covalent bonding with high bond energy and melting temperatures.

• Metals: Exhibit metallic bonding with variable bond strengths and moderate thermal properties.

• Polymers: Exhibit covalent and weak secondary bonding, resulting in low densities and lower melting temperatures.