Atomic Structure and the Periodic Table

Basics of Periodic Table Each box on the table represents an element. In each box… an element symbol the element’s atomic number the element’s average atomic mass Elements arranged in order of increasing atomic number.

Isotope symbol (neutral) Mass number (#p + #n) Atomic number (#P) Isotope name

Isotope symbol (charged)

Average Atomic Mass The masses found on the periodic table are called average atomic masses. They represent the weighted average of all the isotopes found in a sample of the element Isotopes are atoms of the same element with different numbers of neutrons

Example The atomic masses of the two stable isotopes of boron, boron-10 (19.78%) and boron-11 (80.22 %), are amu and amu, respectively. Calculate the average atomic mass of boron.

Remember… Electron configuration…ie. Shorthand Should be able to do shorthand w/o diagonal rule Aufbau principle – electrons fill energy levels and sublevels in order of increasing energy Pauli Exclusion principle – no two electrons can have the same set of four quantum numbers (which means no two electrons can be in the same place at the same time Hund’s rule – when adding electrons to sublevels with more than one orbital, each orbital gets its own electron first before pairing

s orbitals

p orbitals

d orbitals

f orbitals

Quantum Numbers Just as a point on an xy-graph needs a set of two coordinates, each electron has a unique set of four coordinates. These four coordinates represent shell (energy level), subshell (sublevel), orbital, and spin direction of the electron.

Principal Quantum number Represented by n Corresponds to the rows of the periodic table Therefore n = 1, 2, 3, and so on Tells the size of the electron cloud

2 nd Quantum number Represented by l Called the angular momentum quantum number Describes the shape of the orbital l can have the values from 0 to n-1 0 = s sublevel 1 = p sublevel 2 = d sublevel 3 = f sublevel

3 rd Quantum Number Called the magnetic quantum number Describes the orientation in space of the orbital Whether the path of the electron lies on the x, y, or z axis Represented by m l m l can have values from –l to +l if l = 2, then m l = -2, -1, 0, +1, +2

4 th Quantum Number Corresponds to the spin of an electron Represented by m s Clockwise represented by +1/2 Counterclockwise represented by -1/2

Therefore Mg (3, 0, 0, -1/2) Bi (6, 1, +1, +1/2) Co (3, 2, -1, -1/2) Cf (5, 3, -1, -1/2)

Diamagnetism/Paramagnetism Diamagnetic elements have all of their electrons spin paired. Which means they have complete sublevels. Are not affected by a magnetic field Paramagnetic elements do not have all of their electrons spin paired. Strongly affected by a magnetic field

Ground State vs Excited State In a ground state atom, all electrons are in the lowest available sublevels. For an atom in the excited state, one or more electrons have absorbed enough energy to jump to higher energy levels. As soon as possible, those excited electrons will release the energy in the form of a photon, possibly as colored light.

Shorthand for ions Ca 2+ 1s 2 2s 2 2p 6 3s 2 3p 6 K + 1s 2 2s 2 2p 6 3s 2 3p 6 Cl - 1s 2 2s 2 2p 6 3s 2 3p 6 S 2- 1s 2 2s 2 2p 6 3s 2 3p 6 P 3- 1s 2 2s 2 2p 6 3s 2 3p 6 All of these ions have the same configuration as argon and are isoelectronic.