Unit. 5 Electron Configuration

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

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 3

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

C = λ  f 5

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

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

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

E = h f Photoelectric Effect 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 Photoelectric Effect E = h f H = Plank’s constant As frequency increases energy increases

Electrons in an atom add energy to go to an “excited state”. 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

5 ______ 4 ______ 3 ______ 2 ______ 1 ______ 5 ______ 4 ______ 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

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

Transitions of the visible spectrum There are three main series of transition Paschen Series- IR Drops into 3rd Level Balmer Series - Visible Drops into 2nd Level Lyman Series - UV Drops into 1st Level

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

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.

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

l – Orbital Quantum Number 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

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 Spherical in shape p l = 1 Dumbbell in shape d l = 2 f l = 3 f d s p

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 m = 0, ±1, ±2, ±3, ±l “Street” you live on m – Magnetic Quantum Number

Can only be one s orbital 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 Designation n l 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

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

How is the Bohr model different from the model of Dalton? 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.

Quantum…. What….. ?

Aufbau Principal Lowest energy orbital available fills first “Lazy Tenant Rule” 27

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

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

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

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

An energy diagram for Neon 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

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

____ ____ ____ ____ ____ ____ ____ ____ ____ 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!

Bellringer What is the electron configuration for As ? A. Use the long form B. Use the noble gas notation Log onto SFHS website – navigate to Staff / Curtis and select the Gizmo link

Valence Electrons As (atomic number 33) 1s22s22p63s23p64s23d104p3 The electrons in the outermost energy level. s and p electrons in last shell 5 valence electrons 37

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

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

Let’s Practice P (atomic number 15) Ca (atomic number 20) 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

Noble Gas Configuration Your Turn N (atomic number 7) [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

Full energy level Full sublevel Half full sublevel

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 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

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

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

X 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

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 • • • •