Atomic Structures: Electron Configuration

Spectra

Light through a prism – each different wavelength pass through the prism at a different angle 

  • Each color is blended together, not separate – a continuous spectrum 

    • Contains every wavelength and every frequency of light waves from red to violet

    • Also contains every possible amount of energy from high energy violet waves to lower energy red waves

  • Emission line spectrum: discontinuous; an atom, element, or molecule in an excited state returns to a lower energy state

    • Energy produced by the excited atom is quantized

      • Quantized: can have only certain values within a given range

    • Each element has a different line spectrum 

    • Each atom of an element absorbs and emits the same frequencies of electromagnetic radiation.

    • These observations led to the Bohr model

Quantum Mechanics

  • Quantum: discrete quantity of energy that can either be lost or gained by an electron 

  • Quantum Mechanics: study of motion of objects that are atomic or subatomic in size and thus demonstrate wave-particle duality

    • motion of particles can only be affected as they gain or lose energy in discrete amounts called quanta

  • wave-particle duality: electron is both a particle and a wave

  • electron cloud: location of electrons in quantum mechanical model of the atom

  • has variable densities: high density where the electron is most likely to be found and low density where the electron is least likely to be. 

  • orbital: three-dimensional region of space that indicated where there is a high probability of finding an electron

Heisenberg Uncertainty Principle

  • states that it is impossible to determine simultaneously both the position and the velocity of a particle

  • cannot specify accurately the location of an electron — can only say that there is a probability that it exists within this uncertain volume of space

Quantum Mechanical Atomic Model

  • Bohr model explained hydrogen spectrum but not other atoms with more than one electron

  • made assumptions not from experimental conclusions

  • exact path of the electron was restricted to very well-defined circular orbits → later proved wrong, that electrons do not travel around the nucleus in simple circular orbits

  1. electron treated mathematically as a wave (not a particle like in Bohr model) (electron has properties of both a wave and particle)

  2. energy is described in terms of probability of locating the electron in a region of space outside nucleus  

Quantum Numbers

  • describes energy level of an electron using four quantum numbers

  • no two electrons have same set of four quantum numbers

  • Pauli Exclusion Principle: states that no more than two electrons can occupy the same orbital in an atom. The two electrons must also have opposing spins

  1. principal quantum number or principal energy level (n)

  2. sublevel or second quantum number (l)

  3. orbital or third quantum number (m)

  4. electron spin (ms )

1. principal quantum number or principal energy level (n )

  1. describes the most probable distance of the electrons from the nucleus 

  2. has whole number values — 1, 2, 3, 4, … called principal energy levels

  3. similar to Bohr’s K, L, M, N shells

  4. maximum number of electrons in any principal energy level n is 2n2

  5. 1 = 2e, 2 = 8e, 3 = 18e, …

2. sublevel or second quantum number (l )

  • Orbital Angular Momentum Quantum Number

  • given number you can find shape of atomic orbital

  • l = 0 — s 

  • l = 1 — p

  • l = 2 — d

  • l = 3 — f

3. Orbital or magnetic quantum number (m )

  • orbital is region of space around a nucleus in which the probability of finding an electron is high

  • specifies the orientation in space of an orbital of a given energy (n) and shape (l)

  • there are 2l+1 orbitals in each sublevel

  • within a given sublevel in free atom, the energies of orbitals are equal

  • each sublevel has a specific number of orbitals

  • 1 orbital in s sublevel —   0

  • 3 orbitals in p sublevel —   -1, 0, 1

  • 5 orbitals in d sublevel —   -2, -1, 0, 1, 2

  • 7 orbitals in f sublevel —   -3, -2, -1, 0, 1, 2, 3

  • each sublevel and its orbitals are related to the shape of the region that an electron may occupy

    • s is spherical 

    • p is dumbell

    • d is figure eight with a donut wrapped around it 

    • f  is three dimensional clover

  • note: orbital does not represent the path of an electron; it represents a region in which the electron can be found 90% of the time

4. Electron spin (ms)

  • electron acts like it’s spinning on an axis 

  • negatively charged object spins it creates magnetic field

    • counter-clockwise — north pole “up”

    • clockwise - north pole “down”

  • electrons can have the same three principal quantum numbers — MUST HAVE OPPOSITE SPINS

Electron Configuration

  • Electrons in the same orbital are called paired electrons

  • Electron orbitals having the same energy levels are called degenerate orbitals

  • electrons will normally occupy the lowest energy orbital available

Hund’s rule

  • predicts the way electrons will fill the orbitals within a given sublevel

  1. No orbital in a sublevel may contain two electrons unless all of the orbitals in that sublevel contain at least one electron → orbitals fill one electron at a time

  2. All unpaired electrons on a given orbital have the same spin

  • degenerate orbitals are filled evenly with the same spin before electrons are filled into higher energy levels.

  • Note: called orbital diagrams or orbital notation 

  • What do you notice about the electron structure of elements in the same vertical column 

    • that elements in the same family have similar electron structures leading to similar properties 

      • ie. Lithium and sodium both alkali metal family

Transition Elements

  • Third principal energy level holds 18 electrons – 3s (2), 3p (6), 3d (10)

  • 3d is higher in energy than 4s – therefore 4s is filled before filling 3d

  • Elements that complete inner, or lower, energy levels before completing outer, or higher, energy levels are called transitions elements

Aufbau Principle

  • Aufbau principle: states that an electron occupies orbitals in order from lowest energy to highest

  • You can use what you know about principal energy levels electron capacities to figure out the electron configuration of elements when using the information on the Periodic Table

  • These numbers are only for elements in ground state - you can figure out if an atom is in an excited state by comparing a given notation to the one on the periodic table

  • ie oxygen: 2-6  → 1s2 2s2 2p4

  • ie calcium: 2-8-8-2  → 1s2 2s2 2p6 3s2 3p6 4s2 

What do you notice about Noble Gases?

  • Their outermost principal energy has their full complement of electrons

  • No room for electrons from other atoms to enter these levels → do not readily form chemical bonds with other elements (AKA inert gases)

Abbreviated Notation Using Noble Gases

  • These filled energy levels also allow noble gases to be used as shortcuts for writing electron notations of other elements

  • You take the element and work backward to the closest noble gas, put that symbol in brackets and then write the rest of the notation of the element leaving out anything that matches the notation of the noble gas

    • ie. Mg → [Ne] 3s2

    • ie. K → [Ar] 4s1