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Electron Configuration and the Periodic Table Mallard Creek Chemistry - Rines.

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1 Electron Configuration and the Periodic Table Mallard Creek Chemistry - Rines

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

3 Electromagnetic Radiation  Mathematical Relations C = speed of light = 3.0 x 10 8 m/s C = speed of light = 3.0 x 10 8 m/s λ (lamda) = wavelength (m) λ (lamda) = wavelength (m) f= frequency (Hz or s -1 ) f= frequency (Hz or s -1 ) This is how we know what color light is emitted! This is how we know what color light is emitted! C = λ  f

4 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 10 8 m/s

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

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

7 Bohr’s Model  The Photoelectric Effect  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

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

9 Electromagnetic Radiation Photoelectric Effect 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  Based on the frequency and wavelength of the photon

10 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!!!!!

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

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

13 Electromagnetic Radiation Particle Nature of Light Particle Nature of Light  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. E = hf

14 Other Scientists Contributions De Broglie De Broglie Heisenburg 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

15 Other Scientists Contributions Schrödinger 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.

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

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

18 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

19 Quantum Numbers  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 n = 3 l = 0,1,2 n = 2 l = 0,1 n = 1 l = 0

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

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

22 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

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

24 Orbital Designations Orbital Designation nlM2l+1 No. of Orbital No. of Electron 3d32-2,-1,0,+1, p31-1,0,+136 3s p21-1,0,+136 2s s10012

25 Orbital Rules Energy Level Possible sub-levels Number of Sub-levels n No. of Orbitals n 2 No. of Electrons 2n 2 4 s, p, d, f s, p, d s, p 248 1s112

26 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 2 nd level? 4 th level.

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

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

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

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

31 Periodic Table & Electron Configuration

32 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

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

34 Orbital Energy Diagram and Electron Configuration p ______ ______ ______ 3 s ______ p ______ ______ ______ 2 s ______ 1 s ______ An energy diagram for Neon Increasing Energy Electron Spin 1s 2 2s 2 2p x 2 2p y 2 2p z 2 2p 6 1s 2 2s 2 Electron Configuration Notation

35 Orbital Notation  Orbital Notation shows each orbital  O (atomic number 8) ____ ____ ____ ____ ____ ____ ____ ____ ____ ____ ____ ____ 1s 2s 2p x 2p y 2p z 3s 1s 2s 2p x 2p y 2p z 3s  1s 2 2s 2 2p 4 electron configuration!

36 Orbital Notation  Orbital Notation shows each orbital  O (atomic number 8) ____ ____ ____ ____ ____ ____ ____ ____ ____ ____ ____ ____ 1s 2s 2p x 2p y 2p z 3s 1s 2s 2p x 2p y 2p z 3s !!!!

37 Orbital Notation  Write the orbital notation for S  S (atomic number 16) ___ __ __ __ __ __ __ __ __ ___ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 1s 2s 2p 3s 3p  1s 2 2s 2 2p 6 3s 2 3p 4  How many unpaired electrons does sulfur have? 2 unpaired electrons!

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

39 Valence Electrons  Valence Electrons  As (atomic number 33)  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3  The electrons in the outermost energy level.  s and p electrons in last shell  5 valence electrons

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

41 zShorthand Configuration S 16e - Valence Electrons Core Electrons S16e - [Ne] 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 4 Valence Electrons Longhand Configuration Longhand Configuration

42 [Ar] 4s 2 3d 10 4p 2 Noble Gas Configuration Example - Germanium Example - Germanium X X X X X X X X X X X X X

43 Electron Configuration Let’s Practice  P (atomic number 15)  1s 2 2s 2 2p 6 3s 2 3p 3  Ca (atomic number 20)  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2  As (atomic number 33)  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3  W (atomic number 74)  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 4 Noble Gas Configuration [Ne] 3s 2 3p 3 [Ar] 4s 2 [Ar] 4s 2 3d 10 4p 3 [Xe] 6s 2 4f 14 5d 4

44 Electron Configuration Your Turn  N (atomic number 7)  1s 2 2s 2 2p 3  Na (atomic number 11)  1s 2 2s 2 2p 6 3s 1  Sb (atomic number 51)  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 3  Cr (atomic number 24)  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 Noble Gas Configuration [He] 2s 2 2p 3 [Ne] 3s 1 [Kr]5s 2 4d 10 5p 3 [Ar] 4s 2 3d 4

45 Stability  Full energy level  Full sublevel  Half full sublevel

46 Exceptions  Copper  Expect: [Ar] 4s 2 3d 9  Actual: [Ar] 4s 1 3d 10  Silver  Expect: [Kr] 5s 2 4d 9  Actual: [Kr] 5s 1 4d 10  Chromium  Expect: [Ar] 4s 2 3d 4  Actual: [Ar] 4s 1 3d 5  Molybdenum  Expect: [Kr] 5s 2 4d 4  Actual: [Kr] 5s 1 4d 5 Exceptions are explained, but not predicted! Atoms are more stable with half full sublevel

47 Stability Atoms create stability by losing, gaining or sharing electrons to obtain a full octet Atoms create stability by losing, gaining or sharing electrons to obtain a full octet Isoelectronic with noble gases Isoelectronic with noble gases Atoms take electron configuration of the closest noble gas

48 Stability  Na (atomic number 11)  1s 2 2s 2 2p 6 3s 1  1s 2 2s 2 2p 6 = [Ne] Na 1 Valence electron Metal = Loses Ne

49 Try Some  P -3 (atomic number 15)  1s 2 2s 2 2p 6 3s 2 3p 6  Ca +2 (atomic number 20)  1s 2 2s 2 2p 6 3s 2 3p 6  Zn +2 (atomic number 30)  1s 2 2s 2 2p 6 3s 2 3p 6 3d 10  Lost valence electrons (s and p) Full Octet

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

51 Lewis Structures  Shows valence electrons only!  s & p electrons 1. Write noble gas configuration for the element 2. Place valence electrons around element symbol in order X s electrons p electrons

52 Try Some Write the Lewis structures for: Write the Lewis structures for: Oxygen (O) Oxygen (O) –[He] 2s 2 2p 4 Iron (Fe) Iron (Fe) –[Ar] 4s 2 3d 6 Bromine (Br) Bromine (Br) –[Ar] 4s 2 3d 10 4p 5 O Fe Br Valence electrons

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

54 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

55 Periodic Table 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 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 He broke the pattern of increasing atomic weight to keep similar reacting elements together

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

57 Periodic Table Organization Metalloids Metals Non-Metals

58 Periodic Table Organization Representative Elements Transition Metals Inner Transition Metals

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

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

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

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

63 Ionization Energy Noble gases have the highest first Ionization Energy

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

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

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


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