Chapter 8: Periodic properties of the elements

Electron configuration Electrons occupy orbitals following a set of rules The number of electrons determine the eventual electronic configuration Orbital diagrams can be used to visually represent how electrons occupy orbitals Orbital diagram Electron configuration

Pauli exclusion principle Quantum numbers can indentify each individual electrons Pauli exclusion principle – no two electrons can have the same four quantum numbers Electron configuration Quantum Numbers Orbital diagram

Aufbau principle Aufbau principle - When placing more electrons, fill the lowest energy orbitals first Lower l value, lower energy s < p < d < f Lower n value, lower energy Exceptions once you get to d and f orbitals 4s  3d

Diagonal rule

Start from top left (hydrogen) Follow order of increasing atomic number 6s  then 4f  then 5d 7s  then 5f  then 6d

Hund’s rule Hund’s rule – when placing more electrons, fill in empty orbitals first with parallel spins, then pair up half filled orbitals with opposite spin electrons Occurs with degenerate (equal energy) orbitals (ex: 3px, 3py, 3pz)

Practice: Electron configuration for Vanadium atomic number = 23 Number of electrons = 23 Use electrons to fill lower energy orbitals in correct order (refer to chart) Fill in vacant orbitals first before pairing two electrons up on degenerate orbitals

Noble Gas Abbreviation The electron configuration of the noble gas that precedes the element in question is represented by the noble gas’ bracketed symbol He = 1s2 C = 1s22s22p2 = [He] 2s22p2 A quicker way to write out electron configurations Electron configuration of Xe, xenon 1s22s22p63s23p64s23d104p65s24d105p6 Electron configuration of W, tungsten 1s22s22p63s23p64s23d104p65s24d105p66s24f145d4 or [Xe]66s24f145d4

d-block and f-block exceptions Sometimes transition metals, lanthanides, and actinides can deviate from the expected rules Chromium expected valence shell – 4s23d4 Chromium observed valence shell – 4s13d5 Sometimes it saves energy to half fill or completely fill the lower sublevel, than to have it in the 4s or 5s orbital

Valence Electrons Electrons in the outermost (valence) shell The shell with the highest n value C = 1s22s22p2 --- 2 is highest number count up electrons in all the valence orbitals 2 electrons in the 2s orbital and 2 electrons in the 2p orbital 2 + 2 = 4 valence electrons P = 1s22s22p63s23p3 ------- 3 is highest number 2 + 3 = 5 valence electrons Core electron – the rest of the electrons aside from the valence electrons The inner electrons, like a core from an apple

Electron configuration for ions When writing electron configurations for ions, start from the neutral atoms configuration, then remove or add the appropriate amount of electrons dependant on the charge of the ion When removing electrons, pull out of the highest principle quantum number orbitals first V = [Ar] 4s23d3 V2+ = [Ar] 4s03d3 = [Ar] 3d3 electrons dependant on the charge of the ion

Coulomb’s law: implications Electron configuration rules derived from interactions of the electrons and the nucleus Coulomb’s law – potential energy of two charged particles dependant on magnitude of charges (q1, q2) and the distance between them (r) If charges are the same ( both positive or both negative) the energy is positive Since like charges repel each other, the closer they are the more potential energy they have If charges are opposite (one negative and one positive) the energy will be negative Opposites attract, the closer they are, the more negative the potential energy Larger magnitude of q will increase the potential energy +3 charge will have stronger attraction than a +1 charge

Shielding The nucleus of an atom is positively charged and will attract negatively charged electrons In multielectron atoms electrons are repelled by the other electrons present Shielding – electrons preventing other electrons from experiencing full effects of nuclear charge Effective nuclear charge (Zeff) – The approximate charge an electron will experience For Li+ 3 protons, 3+ 2 electrons, 2- 1s2 An incoming electron outside the “shielded” region experiences 3 – 2 = +1 charge Zeff = +1

Penetration The actual nuclear charge Z is weakened by the charge shielded by screening electrons S, giving the effective nuclear charge Zeff If an electron were to “penetrate” inside of the shielded region, it would then experience the full nuclear charge This is penetration

Orbital ordering explained Radial distribution functions for orbitals solved from the schrödinger equation When comparing orbitals with the same n value, the 2s has higher penetration than the 2p orbital Lower l value = higher penetration

4s has greater penetration than the 3d orbital, enough be a lower energy orbital in comparison

Size of an atom Atomic radius – the radius of an atom (spherical in shape) Non bonding atomic radius or van der Waals radius – determined upon freezing an element into the solid state, and measuring the distance between atom centers using the density Bonding atomic radius or covalent radius Nonmetals: half the distance between two of the atoms bonded together Metals: One-half the distance between two of the atoms next to each other in a crystal of the metal

Periodic trend for atomic radius As we move down a column/family, the radius increases As we move to the right across a period, the radius decreases Transition metals are a slight expection, generally staying constant throughout

Zeff trend Effective nuclear charge will increase going to the right across a period Core electrons shield better than valence electrons Going to the right increases proton number and actual charge Z, but the S increases at a lower rate, resulting in a higher Zeff going to the right

Magnetic properties of atoms and ions Unpaired electrons generate a magnetic field due to its spin Atoms or ions that have unpaired electrons are called paramagnetic More unpaired electrons make a more paramagnetic atom or ion Using the electron configuration you can determine the presence of unpaired electrons If all electrons are paired, then the atom or ion is diamagnetic diamagnetic paramagnetic

Ionic radii When an atom loses or gains electrons, it changes the size of the atom Upon losing electrons, the atom gets smaller Cations are smaller than the neutral atom Upon gaining electrons, the atom gets larger Anions are larger than the neutral atom Presence of more electrons causes electron repulsion, that increases repulsion

Isoelectronic ionic trends Isoelectronic – to have the same number of electrons, also being the same net electron configuration For isoelectronic ions, the higher the number of protons, the greater the nuclear charge which pulls the electrons in, the smaller the size

Ionization energy (IE) IE – the energy required to remove an electron from an atom or ion in the gaseous state Removing the first electron is the first ionization energy (IE1) Removing another electron is the second ionization energy (IE2) So on and so on Depending on the electron being removed, and where it is located, this relates to the energy that it will take to remove it

IE1 trends

IE1 trends

IE2,3,4…n trends

Electron Affinities (EA) EA – how easily an atom or ion will accept an electron in the gaseous state The EA is the energy associated with the gaining of an electron Values are typically negative since an atom releases heat when it gains an electron As EA becomes more negative, it becomes more favorable for an atom or ion to gain an electron

Metallic character

The end of Chapter 8