Presentation on theme: "Electron Configuration and the Periodic Table"— Presentation transcript:
1Electron Configuration and the Periodic Table Mallard Creek Chemistry - Rines
2Electromagnetic Radiation Property of WavesFrequencyNo. of waves per secondWave LengthDistance between corresponding points in a waveAmplitudeSize of the wave peakWave Nature of Light
3Electromagnetic Radiation Mathematical RelationsC = speed of light = 3.0 x 108 m/sλ (lamda) = wavelength (m)f= frequency (Hz or s-1)This is how we know what color light is emitted!C = λ f
4Frequency is inversely proportional to Wavelength If λ increases f decreasesIf f increases λ decreasesSpeed of the wave is always constant at 3.0 x 108 m/s
5Bohr Model Nucleus: Neutrons and Protons Orbits: Electrons We know both specific energy and location of each electronElectrons orbit the nucleus in certain fixed energy levels (or shells)Energy LevelsNucleus
6Bohr 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
7Bohr’s Model The Photoelectric Effect Atomic Emission Spectra Light releases electronsNot all colors workAtomic Emission SpectraHydrogen gas emitted specific bands of lightBohr’s calculated energies matched the IR, visible, and UV lines for the H atom654321
8Electromagnetic Radiation Photoelectric Effect – There is a minimum frequency to eject the electron
9Electromagnetic Radiation Only explained by “energy packets” of light called a quantumQuantum - minimum amount of energy that can be gained or lost by an atomPhotons are massless particles of light of a certain quantum of energyBased on the frequency and wavelength of the photonPhotoelectric Effect
10Bohr’s Model Excited electrons Energy added to atom – electrons “jump” up energy levelsWhen the atom relaxes - electron “falls” to lower energy levels and emits photonBohr Model of hydrogenReference Sheets!!!!!
11Electromagnetic Radiation Atomic Line SpectraElectrons 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
12Electromagnetic Radiation These observations suggested that electrons must exist in defined energy levelsFirst, the electron absorbs energy and jumps from the ground state to an excited stateNext, the excited electron relaxes to a lower excited state or ground state5 ______4 ______3 ______2 ______1 ______hv5 ______4 ______3 ______2 ______1 ______5 ______4 ______3 ______2 ______1 ______hv
13Electromagnetic Radiation Wave nature could not explain all observations (Plank & Einstein)Photoelectric EffectWhen light strikes a metal electrons are ejectedAtomic Line SpectraWhen elements are heated, they emit a unique set of frequencies of visible and non-visible light.Particle Nature of LightE = hf
14Other Scientists Contributions De BroglieHeisenburgModeled electrons as wavesHeisenberg Uncertainty Principle: states one cannot know the position and energy of an electronElectrons exist in orbital’s of probabilityOrbital - the area in space around the nucleus where there is a 90% probability of finding an electron
15Other Scientists Contributions SchrödingerSchrödinger Wave Equation - mathematical solution of an electron’s energy in an atomQuantum Mechanical Model of the atom – current model of the atom treating electrons as waves.
16Quantum Mechanical Model Nucleus: Neutrons and protonsOrbitals: 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.
17Solutions to the Wave Equation Wave Equation generates 4 variable solutionsn - sizel - shapem - orientations – spinAddress of an electronQuantum Numbers
18Quantum Numbers n – Primary Quantum Number Describes the size and energy of the orbitaln is any positive #n = 1,2,3,4,….Found on the periodic tableLike the “state” you live in
19Quantum Numbers l – Orbital Quantum Number l = 0,1,2,3,4,….(n-1) Sub-level of energyDescribes the shape of the orbitall = 0,1,2,3,4,….(n-1)“City” you live inl – Orbital Quantum Numbern = 3l = 0,1,2n = 2l = 0,1n = 1l = 0
20Quantum Numbers l – Orbital Quantum Number # level = # sublevels 1st level – 1 sublevel2nd level – 2 sublevels4th level = 4 sublevelsl – Orbital Quantum Number
21Sublevels are named for their shape Quantum NumbersSublevels are named for their shapesl = 0Spherical in shapepl = 1Dumbbell in shapedl = 2fl = 3fdsp
22Quantum Numbers m – Magnetic Quantum Number Describes the orientation of the orbital in spaceAlso denotes how many orbital's are in each sublevelFor each sublevel there are l +1 orbital's“Street” you live onm – Magnetic Quantum Number
23Can only be one s orbital Quantum NumbersLook at Orbital's as Quantum Numbersl = 1 m = -1, 0, +1For each p sublevel there are 3 possible orientations, so three 3 orbital'sl = 0 m = 0Can only be one s orbital
24Orbital Designations 3d 3 2 -2,-1,0,+1,+2 5 10 3p 1 -1,0,+1 6 3s 2p 2s M2l+1No. of OrbitalNo. of Electron3d32-2,-1,0,+1,+25103p1-1,0,+163s2p2s1s
25Orbital Rules n n2 2n2 4 s, p, d, f 16 32 3 s, p, d 9 18 2 s, p 8 1 s Energy LevelPossible sub-levelsNumber of Sub-levelsnNo. of Orbitalsn2No. of Electrons2n24s, p, d, f16323s, p, d9182s, p81s
26ReflectionHow 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.
32Periodic 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
33Orbital Energy Diagram Sub-level (l)Increasing Energyd ______ ______ ______ ______ ______p ______ ______ ______3 s ______2 s ______1 s ______Orbitals (m)Level (n)An energy diagram for the first 3 main energy levels
34Orbital Energy Diagram and Electron Configuration Increasing Energyp ______ ______ ______3 s ______2 s ______1 s ______1s22s22px22py22pz21s22s22p6Electron Configuration NotationElectron SpinAn energy diagram for Neon
35Orbital Notation 1s22s22p4 electron configuration! Orbital Notation shows each orbitalO (atomic number 8)____ ____ ____ ____ ____ ____1s s px 2py 2pz 3s1s22s22p4 electron configuration!
36Orbital Notation ! ____ ____ ____ ____ ____ ____ 1s 2s 2px 2py 2pz 3s Orbital Notation shows each orbitalO (atomic number 8)____ ____ ____ ____ ____ ____1s s px 2py 2pz 3s!
37Orbital Notation ___ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p Write the orbital notation for SS (atomic number 16)___ __ __ __ __ __ __ __ __1s 2s p s p1s22s22p63s23p4How many unpaired electrons does sulfur have?2 unpaired electrons!
38Orbital Notation ___ ____ ____ ____ ____ ____ ____ ____ ____ Write the orbital notation for SS (atomic number 16)___ ____ ____ ____ ____ ____ ____ ____ ____1s 2s p s pHow many unpaired electrons does sulfur have?
39Valence Electrons Valence Electrons As (atomic number 33) 1s22s22p63s23p64s23d104p3The electrons in the outermost energy level.s and p electrons in last shell5 valence electrons
40Valence Electrons Valence Electrons As (atomic number 33) The electrons in the outermost energy level.s and p electrons in last shell
42Noble Gas Configuration Example - GermaniumX X X X X X X X X X X X X[Ar]4s23d104p2
43Electron Configuration Let’s PracticeP (atomic number 15)1s22s22p63s23p3Ca (atomic number 20)1s22s22p63s23p64s2As (atomic number 33)1s22s22p63s23p64s23d104p3W (atomic number 74)1s22s22p63s23p64s23d104p65s24d105p66s24f145d4Noble Gas Configuration [Ne] 3s23p3 [Ar] 4s2 [Ar] 4s23d104p3 [Xe] 6s24f145d4
44Electron Configuration Noble GasConfiguration[He] 2s22p3[Ne] 3s1[Kr]5s24d105p3[Ar] 4s23d4Your TurnN (atomic number 7)1s22s22p3Na (atomic number 11)1s22s22p63s1Sb (atomic number 51)1s22s22p63s23p64s23d104p65s24d105p3Cr (atomic number 24)1s22s22p63s23p64s23d4
45StabilityFull energy levelFull sublevelHalf full sublevel
46Exceptions Exceptions are explained, but not predicted! CopperExpect: [Ar] 4s2 3d9Actual: [Ar] 4s1 3d10SilverExpect: [Kr] 5s2 4d9Actual: [Kr] 5s1 4d10ChromiumExpect: [Ar] 4s2 3d4Actual: [Ar] 4s1 3d5MolybdenumExpect: [Kr] 5s2 4d4Actual: [Kr] 5s1 4d5Exceptions are explained, but not predicted!Atoms are more stable with half full sublevel
47Atoms take electron configuration of the closest noble gas StabilityAtoms create stability by losing, gaining or sharing electrons to obtain a full octetIsoelectronic with noble gases+1+2+3+4-3-2-1Atoms take electron configuration of the closest noble gas
48Stability Na (atomic number 11) 1s22s22p63s1 1s22s22p6 = [Ne] 1 Valence electronMetal = LosesNeNa
49Try Some Full Octet P-3 (atomic number 15) Ca+2 (atomic number 20) 1s22s22p63s23p6Ca+2 (atomic number 20)Zn+2 (atomic number 30)1s22s22p63s23p63d10Lost valence electrons (s and p)Full Octet
50Try Some P-3 (atomic number 15) Ca+2 (atomic number 20) Zn+2 (atomic number 30)
51X Lewis Structures 6 3 4 1 s electrons 7 2 5 8 p electrons Shows valence electrons only!s & p electronsWrite noble gas configuration for the elementPlace valence electrons around element symbol in orderX6341s electronsp electrons7258
52Try Some O Fe Br Valence electrons Write the Lewis structures for: Oxygen (O)[He] 2s2 2p4Iron (Fe)[Ar] 4s2 3d6Bromine (Br)[Ar] 4s2 3d10 4p5•••O•••Valence electronsFe•••••Br••••
53Try SomeWrite the Electron Configuration & Lewis structures for:Oxygen (O)Iron (Fe)Bromine (Br)OFeBr
54What Do I Need to Know? How the periodic table is arranged Be able to identify subcategories of the periodic tableHow the elements within a group are similarHow the elements within a period are similarBe able to compare and contrast the electronegativities, ionization energies, and radii of metals and non-metals
55Periodic Table What He Did Some Problems Dmitri Mendeleev – Father of the Periodic TableWhat He DidPut elements in rows by increasing atomic weightPut elements in columns by similar propertiesSome ProblemsHe 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
56Arranged by Atomic # Columns = Groups MosleyArranged by Atomic # Columns = GroupsRows = Periods
59Metals and Non-metals Metals Non-metals Shiny Malleable Ductile (pulled into wires)Conduct heat and electricityLow specific heatHigh melting pointsSolidsLose electronsDullBrittlePoor conductorsLow melting/boiling pointsVaried propertiesVaried phases
60Atomic Radius Atomic Radius = ½ the distance between adjacent nuclei Increases towards Francium
61Ionic Radius K Cl- K+ Cl Cations Anions Positive Ion Metals Lose electronsRadius gets smaller!Negative IonNon-metalsGain electronsRadius gets larger!KCl-K+Cl
62Ionization Energy Energy required to remove an electron from an atom Why are there peaks in this trend?
63Noble gases have the highest first Ionization Energy
64Electronegativity Increases Pull of electrons in a covalent bond“Attraction” of atoms towards an electronFluorine is “the man”Electronegativity Increases
65Periodic Trends Nuclear Charge increases Atomic radius decreases Ionization energy increasesElectronegativity increasesIonization energy decreasesElectronegativity decreasesOrbital Size increasesAtomic radius increases
66What Do I Need To Know? How are electrons arranged in an atom The two natures of electromagnetic radiation: Particles vs. WavesHow to use the periodic table to list the configuration or orbital diagramWhat quantum numbers are and how they are related to electron configuration.How the periodic table is arrangedThe basic periodic trends