Quantum Theory & Periodicity

Atomic Models Rutherford Nucleus with electrons revolving in orbits Couldn’t explain why electrons didn’t crash into nucleus

Bohr’s Model Electrons can only be certain distance from nucleus Each distance corresponds to certain quantity of energy Close to nucleus – lowest energy level Difference in energy between 2 energy levels = quantum of energy Cannot be in between levels Does not give off energy while in a given level

Present Day Electrons located in orbitals Regions with high probability of finding electrons Electrons clouds

Electrons & Light Bohr Electrons can move from low to high energy by absorbing energy Electrons are unstable at high energy level  move to lower energy level by releasing energy Released as light (with specific wavelength) Each move from a level will release light of a different wavelength Ground state – at lowest possible energy Excited state – higher energy Pg 94 - picture of hydrogen emitting light

Electron Configuration Form of notation which shows how the electrons are distributed among various orbitals and energy levels 1s1 = hydrogen 1 = energy level s = sublevel 1 = number of electrons in that sublevel

So if n = 3  sublevels would be s, p, and d (3 total) n = energy level Indicates how many sublevels there are n = 1  1 sublevel n = 2  2 sublevels n = 3  3 sublevels n = 4  4 sublevels Sublevels 1st sublevel = s 2nd = p 3rd = d 4th = f So if n = 3  sublevels would be s, p, and d (3 total)

What is the order of the sublevels?

Orbitals Each type of sublevel holds different number of orbitals Orbitals can hold 2 electrons Pauli Exclusion Principle Sublevel # of orbitals Max # of electrons s 1 2 p 3 6 d 5 10 f 7 14

Aufbau Principle Electrons fill orbitals with the lowest energy first

Orbital Notation (or diagram) Follow Hund’s rule – orbitals are filled with one electron of the same spin before pairing it with the second electron of a different spin

Orbitals & Electron Capacity of First Four Energy Levels Principle Energy Level Type of sublevel Number of orbitals per type Number of orbitals per level (n2) Max # of electrons (2n2) 1 s 2 4 8 p 3 9 18 d 5 16 32 f 7

Exceptions Cr Cu [Ar] 4s2 3d4  [Ar] 4s1 3d5 Half filled and full d orbitals are more stable

Areas of Periodic Table Main group elements Group 1, 2, 13-18 Regular and consistent with same number of valence electrons Group 1- alkali metals Group 2- alkaline earth metals Group 17- halogens Group 18- noble gases

Metals Alkali Alkaline earth React with water to make alkaline solutions (base) K + H20  KOH Stored in oil Not found as pure elements in nature Alkaline earth Found as compounds in nature also More stable  takes more energy to lose 2 electrons

Metals Transition Lanthanides & Actinides Different configurations Placed due to lack of space Named after 1st element in row Actinides are radioactive

Halogens Most reactive nonmetals  only needs to gain 1 electron React with metal to make salts F, Cl  g Br  l I, At  s

Other groups Noble gases Hydrogen Full set of electrons in outer energy level Stable Inert- but able to get Xe to react Hydrogen In class by itself 1 proton, 1 electron Reacts with many elements

Terms to know Electron shielding Nuclear charge Inner electrons pulled into nucleus due to attractive forces protect the electrons farther out Only a factor when going down a group Nuclear charge Protons in nucleus attract the electrons (pull them in) Only a factor when going across a period

Atomic radius Size of atom Hard to determine Going down a group Another principal energy level is filled Electron shielding protect outer electrons from being pulled in So as you move down, radius increases

Atomic radius (cont.) Going across a period Electrons being added to same energy level Pulled in due to attractive forces with protons (nuclear charge increases) So going across, radius decreases

Ionic radius (size of ion) Metals - the ion of a metal is smaller than a neutral atom of that metal forms an ion as it loses electrons from the valence shell, typically emptying the shell of all electrons size of atom is smaller.

Non-metals - the ion of a non-metal is larger than a neutral atom of that non-metal forms an ion as it gains electrons in its valence shell  By filling the valence shell it gets larger because it is experiencing a high electron-electron repulsion in the valence shell.

Ionization Energy Energy required to remove an electron Moving down a group Element contains more electrons More energy levels Electron shielding protects them  takes less energy to remove So as you move down, IE decreases

Ionization energy (cont.) Moving across a period Number of protons and electrons increase Electrons added to the same energy level Nuclear charge increases (pull electrons in) More energy to remove electrons So as you move across, IE increases

Successive Ionization Energies - more energy is required to remove additional electrons IE1 < IE2 < IE3

Electronegativity Ability of atom in compound to attract electrons Number that is derived from various measurable properties Scale from 0 to 4, with a value of 4 representing an element with the greatest attraction for a free electron Higher number, bigger pull on electrons towards itself

Electronegativity (cont.) Going down a group Distance to valence electrons increases Cannot attract them Smaller electronegativity Going across a period Electrons pulled in closely Higher electronegativity