Chapter 3 - Atoms, Electrons, and Periodic Table

Chapter 3.1: The Nuclear Atomic Model

• John Dalton’s atomic theory: atoms are tiny, solid, and indestructible spheres
  • No longer applies
  • Atoms can be divisible
• Dmitri Mendeleev’s periodic table: listed known elements increasing atomic mass
  • Similar chemical properties were listed in the same column
  • The theory has undergone many changes
• Joseph John Thomson: discovered the existence of a negatively charged particle (known today as an electron)
  • His model of the atom contained electrons, but the entire sphere carried a uniform positive charge
• Ernest Rutherford: discovered that the atom is mainly made up of empty space but has a concentrated positive charge in the center
  • This is also known as the discovery of the atomic nucleus
  • Nuclear model: diagram of an atom that pictures electrons in motion around the atomic nucleus
  • Also known as the planetary model
• The Discovery of neutron help explain why the atomic nucleus did not fly apart or what the total mass of an atom was
• Niels Bohr: developed a similar atomic model to Rutherford but helped explained that electrons had certain restrictions:
  • Atom has only specific energy levels, called stationary states
  • While in a stationary state atoms do not emit energy
  • Atom can change stationary states when it emits or absorbs specific quantities of energy that is equal t the difference in energy between two stationary states
• Planck suggested that matter, at the atomic level, can absorb or emit only discrete quantities of energy; Quantum of energy: these specific quantities
• Photons: particle-like energy, (ex. light)
• Emission spectra: when atoms in an excited state emit photons when they fall to lower energy levels
• Absorption spectrum: results when atoms absorb photons of certain wavelengths and are excited from lower to higher energy levels

Chapter 3.2: The Quantum Mechanical Model of the Atom

• Quantum mechanical model of the atom: a model of the atom that described atoms as having wave-like properties
• Uncertainty principle: impossible to know the position and momentum of an object beyond a certain measure of precision
• Orbitals: wave functions that give information about an electron’s energy and location within an atom
• Ground state: electron-density probabilities for the lowest energy level
• Quantum numbers: Principle quantum number (n): a positive whole number that describes energy level of atomic orbital and relative size Larger n value means larger energy level and electron is farther away from the nucleus Second quantum number/Orbital-shape quantum number (l): describes the orbital shape, a positive integer from 0 to (n-1); refers to to the energy sublevels within each principal energy level;l=0: sl=1: p l=2: dl=3: fTo find energy sublevel/type of orbital, you combine principle quantum number with the letter of orbital shape (s,p,d, or f) Third Quantum number/Magnetic quantum number (m1): values ranging from -1 to +1 and indicates the orientation of orbital in the same around the nucleus

Chapter 3.3: Electron Configurations and Periodic Trends

• Fourth Quantum number/Spin quantum number (ms): specifies the direction of where the electron is spinning
  • Can either be +½ or -½
• Pauli exclusion principle: only two electrons of opposite spin could occupy an orbital
  • No two electrons in an atom have the same four quantum numbers
• Electrons do not actually occupy orbitals, instead, calculate probability densities for electrons
• Electron configuration: shorthand notation of number and arrangement of electrons
  • Atom’s chemical property associated with ground-state electron configuration
• Aufbau principle: build up the ground state electronic structure for each atom
• Orbital diagram: uses a box for each orbital in each given principal energy level
  • Empty box: orbital with no electrons
  • Arrow up/down: orbital with an electron that spins in that direction
  • Two arrows mean filled orbital
• Hund’s rule: no two electrons can have the same set of quantum numbers
  • The orbital diagram must have at least one before adding a second electron to the orbital
• Some elements have different electron configurations from predicted
  • Cr, Cu have electron configurations that result in more stable electrons, so 4s has 1 electron filled instead of 2
• Atomic radius: measurable property between nuclei of bonded
  • Molecules atomic radius is half the distance between nuclei of identical atoms that bonded together
  • As n increases, higher probability of finding electrons far away from the nucleus since atomic volume is larger
  • Effective nuclear charge: the net force of attraction between electrons and the nucleus
  • Down group: inner electrons increase which increases shielding effect
  • Across period: as you go across decrease in atomic size results in effective nuclear charge increasing
• Ionization energy: the energy needed to remove one electron from ground state gaseous atom
  • Decreases as you go down a group due to the atomic radius (increase in distance between nucleus and valence electrons)
  • Increases across a period due to the atomic radius (effective nuclear charge)

Electron affinity: change in energy that occurs when an electron is added to a gaseous atom