In the 1920

In the 1920’s Werner Heisenberg, Louis De Broglie and Erwin Schrödinger developed a different approach: Wave Mechanics or Quantum Mechanics.

With the idea that electrons behave also as waves, they used the mathematics of waves to describe the energy of the electron.

Shrödinger used the physical principles for describing standing waves to treat the electron.

H = Set of mathematical operators

Y =Wave function of the coordinates x, y and z

E = Total energy of the atom (potential and kinetic)

A specific wave function describes an orbital.

Orbital: The 3-D region in the space where the electron is most likely to be found.

The electrons are moving randomly. They don’t follow a specific path. We now talk about the probability of finding the electron in a given region.

The wave function describes the energy and spatial distribution available for an electron.

Heisenberg uncertainty principle: there is a fundamental limitation to just how precise we can know both the position and the momentum of a particle at a given time. DxD(mv) ³ h

Ground state: Lowest energy state for the electrons in an atom

Excited electrons can jump to other orbitals with higher energy when energy is absorbed.

Aufbau principle: As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to the hydrogen like orbitals.

Pauli exclusion principle: Two electrons in the same atom cannot have the same set of four quantum numbers. If they are in the same level of energy, same orbital of the same orientation the spin must be different (opposite).

Hund’s rule: The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the

Pauli exclusion principle in a particular set of degenerate orbitals. In addition, all of the unpaired electrons have parallel spins.

Atom: The smallest piece of an element that still has the properties of the element. (It contains protons, neutrons, electrons)

Molecules: particles made of more than one atom chemically bonded. There are molecules of elements and molecules of compounds.

Molecules of elements are made of the same type of atoms (H2, O2, N2, F2, Br2, Cl2, I2, etc.)

Molecules of compounds are made of different types of atoms covalently bonded (H2O, CO2, C12H22O11)

Formula Units refer to the smallest ratio of ions in ionic compounds (NaCl, CaF2)

Z= atomic number = # of protons

A= mass number = # of protons + neutrons

Don’t confuse mass number A with relative atomic mass Ar

The relative atomic mass Ar is the mass of an atom relative to carbon 12. It has no units. It represents the average mass of all the isotopes. The mass number A is a whole number because it is equal to the number of protons plus the number of neutrons.

#p= Z = # e # n = A- Z

Isotopes: Atoms of the same element with different mass.

They have the same value of Z but diffrent value of A.

It means that they have the same number of protons but different mumber of neutrons. Examples: Cl-35 and Cl-37

They could also be written