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