Electron Configuration and the Periodic Table Mallard Creek Chemistry - Rines

Electromagnetic Radiation Property of Waves Frequency No. of waves per second Wave Length Distance between corresponding points in a wave Amplitude Size of the wave peak Wave Nature of Light

Electromagnetic Radiation Mathematical Relations C = speed of light = 3.0 x 108 m/s λ (lamda) = wavelength (m) f= frequency (Hz or s-1) This is how we know what color light is emitted! C = λ  f

Frequency is inversely proportional to Wavelength If λ increases f decreases If f increases λ decreases Speed of the wave is always constant at 3.0 x 108 m/s

Bohr Model Nucleus: Neutrons and Protons Orbits: Electrons We know both specific energy and location of each electron Electrons orbit the nucleus in certain fixed energy levels (or shells) Energy Levels Nucleus

Bohr Model Bohr’s Atomic Model of Hydrogen Bohr - electrons exist in energy levels AND defined orbits around the nucleus. Each orbit corresponds to a different energy level. The further out the orbit, the higher the energy level

Bohr’s Model The Photoelectric Effect Atomic Emission Spectra Light releases electrons Not all colors work Atomic Emission Spectra Hydrogen gas emitted specific bands of light Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom 6 5 4 3 2 1

Electromagnetic Radiation Photoelectric Effect – There is a minimum frequency to eject the electron

Electromagnetic Radiation Only explained by “energy packets” of light called a quantum Quantum - minimum amount of energy that can be gained or lost by an atom Photons are massless particles of light of a certain quantum of energy Based on the frequency and wavelength of the photon Photoelectric Effect

Bohr’s Model Excited electrons Energy added to atom – electrons “jump” up energy levels When the atom relaxes - electron “falls” to lower energy levels and emits photon Bohr Model of hydrogen Reference Sheets!!!!!

Electromagnetic Radiation Atomic Line Spectra Electrons in an atom add energy to go to an “excited state”. When they relax back to the ground state, they emit energy in specific energy quanta

Electromagnetic Radiation These observations suggested that electrons must exist in defined energy levels First, the electron absorbs energy and jumps from the ground state to an excited state Next, the excited electron relaxes to a lower excited state or ground state 5 ______ 4 ______ 3 ______ 2 ______ 1 ______ hv 5 ______ 4 ______ 3 ______ 2 ______ 1 ______ 5 ______ 4 ______ 3 ______ 2 ______ 1 ______ hv

Electromagnetic Radiation Wave nature could not explain all observations (Plank & Einstein) Photoelectric Effect When light strikes a metal electrons are ejected Atomic Line Spectra When elements are heated, they emit a unique set of frequencies of visible and non-visible light. Particle Nature of Light E = hf

Other Scientists Contributions De Broglie Heisenburg Modeled electrons as waves Heisenberg Uncertainty Principle: states one cannot know the position and energy of an electron Electrons exist in orbital’s of probability Orbital - the area in space around the nucleus where there is a 90% probability of finding an electron

Other Scientists Contributions Schrödinger Schrödinger Wave Equation - mathematical solution of an electron’s energy in an atom Quantum Mechanical Model of the atom – current model of the atom treating electrons as waves.

Quantum Mechanical Model Nucleus: Neutrons and protons Orbitals: region in space surrounding the nucleus where there is a 95% probability of finding an electron. We know either energy or location of each electron.

Solutions to the Wave Equation Wave Equation generates 4 variable solutions n - size l - shape m - orientation s – spin Address of an electron Quantum Numbers

Quantum Numbers n – Primary Quantum Number Describes the size and energy of the orbital n is any positive # n = 1,2,3,4,…. Found on the periodic table Like the “state” you live in

Quantum Numbers l – Orbital Quantum Number l = 0,1,2,3,4,….(n-1) Sub-level of energy Describes the shape of the orbital l = 0,1,2,3,4,….(n-1) “City” you live in l – Orbital Quantum Number n = 3 l = 0,1,2 n = 2 l = 0,1 n = 1 l = 0

Quantum Numbers l – Orbital Quantum Number # level = # sublevels 1st level – 1 sublevel 2nd level – 2 sublevels 4th level = 4 sublevels l – Orbital Quantum Number

Sublevels are named for their shape Quantum Numbers Sublevels are named for their shape s l = 0 Spherical in shape p l = 1 Dumbbell in shape d l = 2 f l = 3 f d s p

Quantum Numbers m – Magnetic Quantum Number Describes the orientation of the orbital in space Also denotes how many orbital's are in each sublevel For each sublevel there are 2l +1 orbital's “Street” you live on m – Magnetic Quantum Number

Can only be one s orbital Quantum Numbers Look at Orbital's as Quantum Numbers l = 1 m = -1, 0, +1 For each p sublevel there are 3 possible orientations, so three 3 orbital's l = 0 m = 0 Can only be one s orbital

Orbital Designations 3d 3 2 -2,-1,0,+1,+2 5 10 3p 1 -1,0,+1 6 3s 2p 2s M 2l+1 No. of Orbital No. of Electron 3d 3 2 -2,-1,0,+1,+2 5 10 3p 1 -1,0,+1 6 3s 2p 2s 1s

Orbital Rules n n2 2n2 4 s, p, d, f 16 32 3 s, p, d 9 18 2 s, p 8 1 s Energy Level Possible sub-levels Number of Sub-levels n No. of Orbitals n2 No. of Electrons 2n2 4 s, p, d, f 16 32 3 s, p, d 9 18 2 s, p 8 1 s

Reflection How is the Bohr model different from the earlier models of the atom? Who contributed to the modern model of the atom? How is it different from Bohr’s? Why do atoms give unique atomic line spectra? What are ground and excited states? Is 2d possible? 4f ? 2s ? 6p? 1p? How many total orbital's in the 2nd level? 4th level.

Does Your Head Hurt Yet?? Quantum…. What….. ?

Aufbau Principle Aufbau Principal Lowest energy orbital available fills first “Lazy Tenant Rule”

Pauli’s Exclusion Principle No two electrons have the same quantum #’s Maximum electrons in any orbital is two () Pauli Exclusion Principle

Hund’s Rule Hund’s Rule RIGHT WRONG When filling degenerate orbital's, electrons will fill an empty orbital before pairing up with another electron. Empty room rule RIGHT WRONG

Periodic Table & Electron Configuration

Periodic Table & Electron Configuration Using the periodic table for the filling order of orbitals, by going in atomic number sequence until you use all the needed electrons in the element

Orbital Energy Diagram Sub-level (l) Increasing Energy d ______ ______ ______ ______ ______ p ______ ______ ______ 3 s ______ 2 s ______ 1 s ______ Orbitals (m) Level (n) An energy diagram for the first 3 main energy levels

Orbital Energy Diagram and Electron Configuration Increasing Energy p ______ ______ ______ 3 s ______ 2 s ______ 1 s ______ 1s2 2s2 2px2 2py2 2pz2 1s2 2s2 2p6 Electron Configuration Notation Electron Spin An energy diagram for Neon

Orbital Notation 1s22s22p4 electron configuration! Orbital Notation shows each orbital O (atomic number 8) ____ ____ ____ ____ ____ ____ 1s 2s 2px 2py 2pz 3s 1s22s22p4 electron configuration!

Orbital Notation ! ____ ____ ____ ____ ____ ____ 1s 2s 2px 2py 2pz 3s Orbital Notation shows each orbital O (atomic number 8) ____ ____ ____ ____ ____ ____ 1s 2s 2px 2py 2pz 3s !

Orbital Notation ___ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p Write the orbital notation for S S (atomic number 16) ___ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 1s22s22p63s23p4 How many unpaired electrons does sulfur have? 2 unpaired electrons!

Orbital Notation ___ ____ ____ ____ ____ ____ ____ ____ ____ Write the orbital notation for S S (atomic number 16) ___ ____ ____ ____ ____ ____ ____ ____ ____ 1s 2s 2p 3s 3p How many unpaired electrons does sulfur have?

Valence Electrons Valence Electrons As (atomic number 33) 1s22s22p63s23p64s23d104p3 The electrons in the outermost energy level. s and p electrons in last shell 5 valence electrons

Valence Electrons Valence Electrons As (atomic number 33) The electrons in the outermost energy level. s and p electrons in last shell

Shorthand Configuration Valence Electrons Longhand Configuration S 16e- 1s2 2s2 2p6 3s2 3p4 Core Electrons Valence Electrons Shorthand Configuration S 16e- [Ne] 3s2 3p4

Noble Gas Configuration Example - Germanium X X X X X X X X X X X X X [Ar] 4s2 3d10 4p2

Electron Configuration Let’s Practice P (atomic number 15) 1s22s22p63s23p3 Ca (atomic number 20) 1s22s22p63s23p64s2 As (atomic number 33) 1s22s22p63s23p64s23d104p3 W (atomic number 74) 1s22s22p63s23p64s23d104p65s24d105p66s24f145d4 Noble Gas Configuration [Ne] 3s23p3 [Ar] 4s2 [Ar] 4s23d104p3 [Xe] 6s24f145d4

Electron Configuration Noble Gas Configuration [He] 2s22p3 [Ne] 3s1 [Kr]5s24d105p3 [Ar] 4s23d4 Your Turn N (atomic number 7) 1s22s22p3 Na (atomic number 11) 1s22s22p63s1 Sb (atomic number 51) 1s22s22p63s23p64s23d104p65s24d105p3 Cr (atomic number 24) 1s22s22p63s23p64s23d4

Stability Full energy level Full sublevel Half full sublevel

Exceptions Exceptions are explained, but not predicted! Copper Expect: [Ar] 4s2 3d9 Actual: [Ar] 4s1 3d10 Silver Expect: [Kr] 5s2 4d9 Actual: [Kr] 5s1 4d10 Chromium Expect: [Ar] 4s2 3d4 Actual: [Ar] 4s1 3d5 Molybdenum Expect: [Kr] 5s2 4d4 Actual: [Kr] 5s1 4d5 Exceptions are explained, but not predicted! Atoms are more stable with half full sublevel

Atoms take electron configuration of the closest noble gas Stability Atoms create stability by losing, gaining or sharing electrons to obtain a full octet Isoelectronic with noble gases +1 +2 +3 +4 -3 -2 -1 Atoms take electron configuration of the closest noble gas

Stability Na (atomic number 11) 1s22s22p63s1 1s22s22p6 = [Ne] 1 Valence electron Metal = Loses Ne Na

Try Some Full Octet P-3 (atomic number 15) Ca+2 (atomic number 20) 1s22s22p63s23p6 Ca+2 (atomic number 20) Zn+2 (atomic number 30) 1s22s22p63s23p63d10 Lost valence electrons (s and p) Full Octet

Try Some P-3 (atomic number 15) Ca+2 (atomic number 20) Zn+2 (atomic number 30)

X Lewis Structures 6 3 4 1 s electrons 7 2 5 8 p electrons Shows valence electrons only! s & p electrons Write noble gas configuration for the element Place valence electrons around element symbol in order X 6 3 4 1 s electrons p electrons 7 2 5 8

Try Some O Fe Br Valence electrons Write the Lewis structures for: Oxygen (O) [He] 2s2 2p4 Iron (Fe) [Ar] 4s2 3d6 Bromine (Br) [Ar] 4s2 3d10 4p5 • • • O • • • Valence electrons Fe • • • • • Br • • • •

Try Some Write the Electron Configuration & Lewis structures for: Oxygen (O) Iron (Fe) Bromine (Br) O Fe Br

What Do I Need to Know? How the periodic table is arranged Be able to identify subcategories of the periodic table How the elements within a group are similar How the elements within a period are similar Be able to compare and contrast the electronegativities, ionization energies, and radii of metals and non-metals

Periodic Table What He Did Some Problems Dmitri Mendeleev – Father of the Periodic Table What He Did Put elements in rows by increasing atomic weight Put elements in columns by similar properties Some Problems He left blank spaces for what he said were undiscovered elements (he was right!) He broke the pattern of increasing atomic weight to keep similar reacting elements together

Arranged by Atomic # Columns = Groups Mosley Arranged by Atomic # Columns = Groups Rows = Periods

Periodic Table Organization Metalloids Metals Non-Metals

Periodic Table Organization Representative Elements Transition Metals Inner Transition Metals

Metals and Non-metals Metals Non-metals Shiny Malleable Ductile (pulled into wires) Conduct heat and electricity Low specific heat High melting points Solids Lose electrons Dull Brittle Poor conductors Low melting/boiling points Varied properties Varied phases

Atomic Radius Atomic Radius = ½ the distance between adjacent nuclei Increases towards Francium

Ionic Radius K Cl- K+ Cl Cations Anions Positive Ion Metals Lose electrons Radius gets smaller! Negative Ion Non-metals Gain electrons Radius gets larger! K Cl- K+ Cl

Ionization Energy Energy required to remove an electron from an atom Why are there peaks in this trend?

Noble gases have the highest first Ionization Energy

Electronegativity Increases Pull of electrons in a covalent bond “Attraction” of atoms towards an electron Fluorine is “the man” Electronegativity Increases

Periodic Trends Nuclear Charge increases Atomic radius decreases Ionization energy increases Electronegativity increases Ionization energy decreases Electronegativity decreases Orbital Size increases Atomic radius increases

What Do I Need To Know? How are electrons arranged in an atom The two natures of electromagnetic radiation: Particles vs. Waves How to use the periodic table to list the configuration or orbital diagram What quantum numbers are and how they are related to electron configuration. How the periodic table is arranged The basic periodic trends