1.1 Atomic Structure: The Nucleus

Atoms are the building blocks of matter, consisting of a nucleus surrounded by electrons. The nucleus contains protons and neutrons, while electrons orbit in shells. Understanding atomic structure is crucial for grasping chemical properties and reactions.

Isotopes are atoms with the same number of protons but different neutron counts. This affects their mass and stability. Atomic weight, calculated from isotope masses and abundances, is essential for stoichiometry and determining molar mass in chemical calculations.

Atomic Structure and the Nucleus

Structure of atoms

• Atoms fundamental units of matter made up of a dense, positively charged nucleus surrounded by negatively charged electrons in shells or orbitals
• Nucleus contains protons (positive charge, mass ~1 atomic mass unit (amu)) and neutrons (no charge, mass ~1 amu)
• Electrons orbit nucleus in shells or orbitals have a negative charge and negligible mass compared to protons and neutrons
  • Number of electrons in an atom determines its chemical properties (reactivity, bonding)
  • Electrons arranged in energy levels (shells) and sublevels (orbitals) around the nucleus (Bohr model, quantum mechanical model)
• Atoms are electrically neutral with equal numbers of protons and electrons
  • Ions formed when atoms gain or lose electrons (cations are positively charged, anions are negatively charged)

Isotopes and atomic properties

• Isotopes are atoms of the same element with different numbers of neutrons have the same number of protons and electrons but vary in the number of neutrons (e.g., carbon-12, carbon-13, carbon-14)
• Atomic number (Z) is the number of protons in an atom's nucleus defines the identity of an element (e.g., all carbon atoms have 6 protons)
  • All isotopes of an element have the same atomic number
• Mass number (A) is the sum of the number of protons and neutrons in an atom's nucleus differs among isotopes of the same element (e.g., carbon-12 has 6 protons and 6 neutrons, carbon-13 has 6 protons and 7 neutrons)
  • Represented as a superscript to the left of the element symbol ($^{12}C$, $^{13}C$)
• Isotopes have similar chemical properties but different physical properties (mass, radioactivity)
  • Some isotopes are stable (non-radioactive) while others are unstable (radioactive) and undergo decay (emit particles and energy over time)

Nuclear properties and stability

• Nuclear force is the strong attractive force between nucleons (protons and neutrons) that overcomes electrostatic repulsion between protons
• Nuclear stability is determined by the ratio of neutrons to protons in the nucleus
  • Stable nuclei have a specific range of neutron-to-proton ratios
• Binding energy is the energy required to break apart a nucleus into its constituent nucleons
  • Higher binding energy per nucleon generally indicates greater nuclear stability
• Radioactivity is the spontaneous emission of particles or energy from unstable nuclei
  • Radioactive decay occurs when nuclei have an unfavorable neutron-to-proton ratio

Calculation of atomic weight

• Atomic weight is the weighted average of the masses of an element's naturally occurring isotopes takes into account the relative abundance of each isotope
• To calculate atomic weight:

1. Multiply the mass of each isotope by its fractional abundance (expressed as a decimal)
2. Sum the products of mass and fractional abundance for all isotopes

• Example calculation for chlorine with two naturally occurring isotopes, $^{35}Cl$ (75.77%) and $^{37}Cl$ (24.23%):
  • Atomic weight of Cl = $(34.97 \text{ amu} \times 0.7577) + (36.97 \text{ amu} \times 0.2423) = 35.45 \text{ amu}$
• Atomic weights are used in stoichiometry calculations and for determining the molar mass of compounds
  • Molar mass is the sum of the atomic weights of all atoms in a compound (g/mol)