Key Concepts of Atomic Structure- The Nuclear Atom

Chapter 1: Introduction

• Development of Atomic Structure Understanding

  • Ancient Greek philosopher Democritus coined the term "atom," proposing matter consists of indivisible particles.

  • John Dalton supported this theory in the early 1800s through the law of multiple proportions, showing fixed mass percentages in compounds.

• Early Models of Atoms

  • J.J. Thompson (1897) discovered electrons and proposed the plum pudding model, in which electrons are embedded in a positive "soup."

  • Ernest Rutherford (1908-1913) challenged this model, discovering that atoms have a small, dense nucleus, concluding that atoms are mostly empty space.

Chapter 2: Protons and Electrons

• Refinement of Atomic Models

  • Niels Bohr (1913) suggested electrons orbit the nucleus at specific distances and energies.

  • Erwin Schrödinger (1926) introduced the wave model, describing electron distribution as a cloud rather than fixed orbits.

• Structure of Atoms

  • Atoms consist of protons and neutrons in the nucleus, with electrons in surrounding shells.

  • Protons (+1 charge) and electrons (-1 charge) are charged particles, while neutrons are neutral.

  • The mass of neutrons is slightly larger than protons, while electrons are significantly lighter meaning, we don’t consider their mass in normal calculations.

• Forces in Atoms

  • Electrostatic attraction keeps electrons near the nucleus, while the strong nuclear force counters electrostatic repulsion between protons, ensuring nucleus stability.

Chapter 3: Number of Protons

• Atomic Number and Mass Number

  • Atomic number (number of protons) defines the element; mass number is the sum of protons and neutrons.

  • Neutrons help stabilize the nucleus without increasing repulsion among protons.

• Mass Measurement

  • Atomic mass units (AMU) are used for measuring atomic mass, with 1 AMU defined as 1/12th the mass of a carbon atom.

  • The mass of an atom is not equal to its mass number due to mass-energy conversion in stable nuclei.

• Ions

  • Atoms are electrically neutral with equal protons and electrons; ions have unequal numbers, resulting in positive (cations) or negative (anions) charges.

Chapter 4: Atomic Mass (RAM)

• Isotopes

  • Atoms of the same element can have different mass numbers due to varying neutron counts; these are called isotopes.

  • Isotopes are represented with a letter symbol, subscript (atomic number), and superscript (mass number).

• Relative Atomic Mass (RAM)

  • RAM is an average of atomic masses of isotopes in a sample, reflecting their relative abundances.

  • Calculation of RAM involves weighted averages based on isotopic masses and their abundances.

• Example Calculations

  • For isotopes with equal abundance, RAM is the average of their masses.

  • For chlorine (75% Cl-35 and 25% Cl-37), RAM = (35 * 75 + 37 * 25) / 100 = 35.5.

  • For magnesium (79% Mg-24, 10% Mg-25, 11% Mg-26), RAM = (24 * 79 + 25 * 10 + 26 * 11) / 100 = 24.3.

Chapter 5: Conclusion

• Periodic Table Representation

  • Elements are represented by symbols with atomic numbers and RAM displayed.

• Mass Spectrometry

  • Mass spectrometers measure isotopic masses and abundances by ionizing atoms and analyzing their mass-to-charge ratios.

  • Ions are accelerated and deflected in a magnetic field, with the degree of deflection indicating their m/z ratios, displayed as peaks in