Chapter 2: The Structure of the Atom and the Periodic Table

Chapter 2: The Structure of the Atom and the Periodic Table

Atomic Theory

• Historical Background:

  • The nuclear model for atomic theory is credited to Ernest Rutherford.

  • The modern atomic theory utilizes the Bohr atom model, sometimes referred to as the solar system model.

• Definition of the Atom:

  • An atom is defined as the smallest unit of an element that retains its chemical properties and cannot be decomposed into simpler substances through chemical means.

• Composition of Atoms:

  • Atoms are composed of three types of subatomic particles:

  • Protons: Positively charged particles.

  • Neutrons: Neutral particles with no charge.

  • Electrons: Negatively charged particles.

  • The nucleus of an atom contains protons and neutrons.

Emission Spectrum of Atoms

• Each atom has a unique emission spectrum.

• Example for Hydrogen gas:

  • Notable wavelengths in the spectrum:

  • 434 nm (Violet)

  • 486 nm (Blue-green)

  • 656 nm (Red)

• Wavelength (nm) from various sources is visualized, showcasing peaks at specific wavelengths within particular colors.

Isotopic Symbols and Atomic Statistics

Changing Protons

• Mass Number (A): Represents the sum of protons and neutrons in an atom.

• Atomic Number (Z): Represents the number of protons, which determines the identity of the element.

  • Example: For atoms with 26 protons and 29 neutrons or 5 protons and 6 neutrons.

Changing Neutrons

• Isotopes: Are defined as atoms with the same number of protons but differing in their number of neutrons.

• Isotopes of an element share the same chemical properties.

  • Example Isotopes:

  • Hydrogen (1 proton)

  • Deuterium (1 proton, 1 neutron)

  • Tritium (1 proton, 2 neutrons)

Changing Electrons

• In a neutral atom, the number of electrons equals the number of protons.

• Cation: Positively charged atom where the number of protons exceeds the number of electrons (p > e-).

• Anion: Negatively charged atom where the number of electrons exceeds the number of protons (p < e-).

  • Example for isotopic symbols:

  • An example includes 0^{14}_{7} with 7 protons, 7 neutrons, and 10 electrons (ionized state 3-).

  • Another example includes 0^{41}_{19} with 19 protons, 22 neutrons, and 18 electrons (ionized state 19+).

Electronic Structure

• Protons and neutrons are located within the nucleus, which accounts for almost all the atom's mass.

• Electrons exist in orbital regions (orbitals) of specific energies described by quantum numbers.

Quantum Numbers

• Principal Quantum Number (n): Describes the size of the orbital (also known as the shell), and identifies the row of the periodic table.

• Angular Momentum Quantum Number (l): Describes the shape of the orbital (also known as the subshell).

  • Values of l indicate various subshells:

  • l = 0: s-subshell (1 orbital)

  • l = 1: p-subshell (3 orbitals)

  • l = 2: d-subshell (5 orbitals)

  • l = 3: f-subshell (7 orbitals)

Electron Configurations

• Electron Configuration is defined as the actual arrangement of electrons in atomic orbitals.

• Aufbau Principle: States electrons fill the lowest energy orbitals first.

• Order of Energy Levels:

  • The order in which electron energy levels fill as follows:

  • 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s

• Hund’s Rule: States that orbitals in the same subshell must fill singly (each orbital gets one electron) before any orbital can have a second electron.

• Pauli’s Exclusion Principle: States that in an orbital containing two electrons, they must have opposite spins (only applicable in orbital diagrams).

• Examples of Electron Configurations:

  • Example for Lithium (Li):

  • Electron configuration: Li

  • $1s^2 2s^1$

  • Example for Carbon (C):

  • Electron configuration: C

  • $1s^2 2s^2 2p^2$

Shorthand Electron Configurations

• Shorthand configurations utilize the previous noble gas to represent the core electrons.

• Format: [Noble Gas] + Configuration of last filled row.

  • Example for Strontium (Sr):

  • Complete configuration: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2$

  • Shorthand configuration: $[Kr]5s^2$

• Core Electrons: The electrons found the inner shells are typically not involved in bonding and reactivity.

• Valence Electrons: The electrons in the outermost shell are crucial in determining reactivity and bonding behavior.

Common Ion Charges

• The number of valence electrons influences the formation of common ion charges, which can be predicted from the group position within the periodic table.

• Ion Formation: Atoms can lose or gain electrons to reach the nearest noble gas configuration.

Atomic Mass

• Atomic Mass: Represents the weighted average of all naturally occurring isotopic masses for an element.

• The mass listed on the periodic table is the atomic mass, not the mass number.

• Formula for calculating atomic mass:
  Atomic\,Mass = \sum (Isotopic\,Mass) \times \left(\frac{Natural\,Abundance}{100}\right).

• Example Calculation:

  • For Neon, having three isotopes:

  • Atomic\,Mass = 20.0 * 0.9048 + 21.0 * 0.0027 + 22.0 * 0.0925 = 20.1877 \approx 20.2 amu

Atomic Mass Examples

• Example for calculating the average atomic mass of Chlorine:

  • Natural abundance: 75.53% of 35Cl (mass = 34.969 amu) and 24.47% of 37Cl (mass = 36.966 amu)

  • Calculation yields: Average atomic mass of Chlorine = 35.46 amu.

• Example for calculating average atomic mass of Sulfur:

  • 95.00% of sulfur atoms at mass of 31.972 amu, 0.76% at mass of 32.971 amu, and 4.22% at mass of 33.967 amu.

  • Calculation yields: Average atomic mass of Sulfur = 32.06 amu.

Periodic Table Trends

• Groups: Consists of 18 columns.

  • Main Group/Representative elements:

  • Group 1: Alkali metals

  • Group 2: Alkali Earth metals

  • Group 17: Halogens

  • Group 18: Noble gases

• Periods: There are 7 rows.

Periodic Table Questions

• Example questions to identify elements based on given classifications within periods and groups.

Atomic Radii

• Atomic Radii: Defined as the radius of an atom measured from the center of the nucleus to the outermost shell.

• Atomic radii have specific trends:

  • They decrease across a period.

  • They increase down a group.

  • Comparison of sizes:

  • Cations are smaller than their neutral atoms while anions are larger.

Ionization Energy (IE)

• Ionization Energy: Refers to the energy needed to remove an electron from an atom in the gas phase.

• Trends in ionization energy:

  • Ionization energy decreases down a group.

  • Ionization energy increases across a period.

  • Notably, smaller atoms are more difficult to ionize and noble gases require significantly more energy to ionize.

Metallic Character

• Metallic Character: Represents a measure of properties characteristic to metals.

• Trends in metallic character:

  • Increases down a group and decreases across a period.

  • More metallic elements tend to lose electrons (form cations), whereas less metallic elements tend to gain electrons (form anions).

Atomic Properties Examples

• Evaluate the following sets of atomic properties, determining which atoms possess indicated traits.

  • Determining larger radii, smaller ionization energy, greater metallic character, and ranking by atomic radii.

Conclusion of Chapter 2

• Summation of key learning objectives:

  • Understanding atomic structure.

  • Differentiating between elements, isotopes, and ions.

  • Identifying the numbers of protons, neutrons, and electrons in atoms.

  • Writing both full and shorthand electron configurations for atoms.

  • Drawing orbital box diagrams for atoms.

  • Recognizing and navigating the periodic table layout.

  • Identifying common ion charges using valence electrons.

  • Matching group classifications with the periodic table and identifying elements by position.

  • Calculating atomic mass utilizing isotopic masses and percent abundance.

  • Demonstrating understanding of periodic table trends, including atomic radii, ionization energy, and metallic character.