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Unit 4b: Electron Configuration Mapping the Electron

Unit 4b: Electron Configuration Mapping the Electron. Goal: Be able to describe how the following are related 1. Complete ChemQuest 11 2. Write a sentence that demonstrates how energy level , sublevel and orbital are related. Goal: Be able to describe how the following are related.

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Unit 4b: Electron Configuration Mapping the Electron

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  1. Unit 4b: Electron ConfigurationMapping the Electron

  2. Goal: Be able to describe how the following are related 1. Complete ChemQuest 11 2. Write a sentence that demonstrates how energy level, sublevel and orbital are related.

  3. Goal: Be able to describe how the following are related Energy levels can be divided into sublevels that are made up of orbitals.

  4. Mapping out electron configuration • Energy Level • Ring on a Bohr model • Quantized: a certain amount of energy (stable energy amount) • Sublevel • Shape of region where electrons are found • (sphere, dumbbell, clover, indescribable) • Four types: s, p, d and f • Each orbital can hold 2 electrons • S has 1 orbital; maximum electrons = 2 sphere • P has 3 orbitals; maximum electrons = 6 dumbbell • D has 5 orbitals; maximum electrons = 10 clover • F has 7 orbitals; maximum electrons = 14 indescribable • Orbitals • make up each sublevel • regions where electrons are located • maximum of 2 electrons in an orbital

  5. Goal: Be able to describe how the following are related

  6. Keeping in mind the last slide… What if we represented the orbital diagrams with the nucleus on the x-axis?

  7. - - N S Why can two negative electrons exist in same orbital and not repel?Spin Quantum Number, ms North South Electron aligned against magnetic field, ms = - ½ Electron aligned with magnetic field, ms = + ½ The electron behaves as if it were spinning about an axis through its center. This electron spin generates a magnetic field, the direction of which depends on the direction of the spin. Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 208

  8. Electron Cloud  Orbital Shape

  9. Orbitals and their Shape S orbital – Sphere P orbital – Dumbbell D orbital – Clover F orbital – Indescribable (flower)

  10. Energy Level = row in a theatre Sublevels = sections in a row Orbitals = seats in a section Mapping out electron configuration Nucleus

  11. Orbitals and Energy Levels Starts at energy level *Every orbital can hold up to 2 electrons

  12. Filling Rules for Electron Orbitals Aufbau Principle: • Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for. • (Fill lowest energy level to highest energy level) *Aufbau is German for “building up” Hund’s Rule: • Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results. • (Electrons spread out in a sublevel before pairing) Pauli Exclusion Principle: • An orbital can hold a maximum of two electrons. • To occupy the same orbital, two electrons must spin in opposite directions. • (Maximum of 2 electrons with opposite spin per orbital)

  13. Orbital Diagrams • Orbital diagrams are very similar to electron configurations • They show the placement of electrons in an atom • Boxes or lines are used to represent orbitals S has 1 orbital P has 3 orbitals D has 5 orbitals F has 7 orbitals • The energy level and sublevel are written under the boxes (example 1s or 2p) • Electrons are represented by half arrows • ONLY 2 arrows per box pointing opposite directions

  14. Three rules are used to build the Orbital Diagrams: • Aufbau principle: Electrons occupy orbitals of lower energy levels first • “Electrons love the nucleus” • Hund’s Rule: In a set of orbitals, the electrons will fill the orbitals in a way that would give the maximum number of parallel spins (maximum number of unpaired electrons). • Analogy: Students could fill each seat of a school bus, one person at a time, before doubling up • Pauli Exclusion Principle: No two electrons can have the same set of quantum numbers • An orbital can hold only two electrons and they must have opposite spin • One arrow points up, the other points down

  15. Orbital Diagram Example • Draw the orbital diagram for Nitrogen: • Write the configuration for nitrogen • Draw the boxes you need and label them 3. Fill in arrows following Aufbau’s, Hund’s, and Pauli’s rules Why are the arrows in 2p in separate boxes? 1s 2s 2p

  16. Electron Configurations Orbital Filling Element 1s 2s 2px 2py 2pz 3s Configuration Electron H He Li C N O F Ne Na 1s1 1s2 1s22s1 1s22s22p2 1s22s22p3 1s22s22p4 1s22s22p5 1s22s22p6 1s22s22p63s1

  17. Practice • Draw the orbital diagram for Boron • Draw the orbital diagram for Bromine (you may use short hand just don’t for get to include the noble gas before the element in brackets [ ]) • Draw the orbital diagram for Titanium

  18. Definition of Electron Configuration • Is a form of notation showing how electrons are distributed among atomic orbitals and energy levels. • The format for writing electron configurations includes a series of [Number][letter] [superscript number] 1s2

  19. Standard Notation of Fluorine Number of electrons in the sub level 2,2,5 1s22s22p5 Principle Energy Level Numbers 1, 2, 2 Subshell

  20. Blocks in the Periodic Table s starts at 1, p starts at 2, d starts at 3, and f starts at 4

  21. Writing an Electron Configuration- Selenium • Find the element you are looking for on the periodic table • Always start your configuration at hydrogen

  22. Selenium Continued 3. Write the energy number and letter of the subshell, then as a superscript write the number of electrons held in that subshell 1s22s22p63s23p64s23d104p4

  23. Mapping out electron configuration • Electrons are located in energy levels or shells around the nucleus of the atom • Principle Energy Level (n)– Energy level where the electron(s) is(are) found (“n” is the row of the PT) Examples: n= 1, 2, 3, 4, 5, 6, 7 • As “n” increases, electrons: • are farther from the nucleus • Have higher amounts of energy • Electrons want to be as close to the nucleus as they can so they can expend the LEAST amount of energy.

  24. Mapping out electron configuration Each energy level has a maximum capacity for electrons. Maximum in each level = 2n2; where n= 1, 2, 3, 4, 5, 6, 7 • Example: for n =1: max = 2(1)2 electrons • You solve for: n = 2 n = 3 n = 4 = 2 electrons

  25. Mapping out electron configuration • Each energy level contains sublevels or subshells • Sublevels– subdivisions of “n” (sections of seats within a row) • The number of sublevels in each energy level is equal to n. Therefore: n = 1 has one sublevel n = 2 has two sublevels n = 3 has three sublevels n = 4 has four sublevels n = 5 has five sublevels but the fifth one is not used in ground state elements.

  26. Mapping out electron configuration • The names for the various sublevels ares, p, d, f. • The first sublevel in any level is always s; the second is p; third is d; fourth is f. Therefore: • if n = 3 & it has 3 sublevels, their names ares p d • if n = 2 & it has 2 sublevels, their names are…? • if n = 1 & it has 1 sublevel, its name is…? • if n = 4 & it has 4 sublevels, their names are…?

  27. Practice Time • Write the electron configuration of the following elements • Beryllium • Cadmium • Bromine • Uranium • Iron

  28. Electron Configuration with Ions • When doing electron configuration with ions, write the configuration and then add (if it is an anion) or subtract (if it is a cation) the charge from the number of electrons (superscript) • Example: O-2 • Oxygen: 1s22s22p4 • Oxygen ion: 1s22s22p6 Add two electrons

  29. Practice Time • Write the electron configuration of the following ions • Ca+2 • P-3 • Mg+2 • F-1

  30. Noble Gas Notation(AKA shorthand notation) • Use the last noble gas that is located in the periodic table right before the element. • Write the symbol of the noble gas in brackets. • Write the remaining configuration after the brackets. • Ex: • Fluorine: • [He] 2s2 2p5

  31. Practice • Write the noble gas configuration for the following elements: • Chlorine • Mercury • Lanthanum (La) • Argon • K+ • F-

  32. Electrons and Energy Levels • Electrons in their energy levels are considered to be in their ground state • If electrons are given energy, they can jump up in energy levels • We call this the excited state • When they fall back to their energy level, they release a photon of light • A photon of light is a tiny particle of light • The color of the light corresponds to the amount of energy the electron released when it goes from its excited state to its ground state

  33. Energy absorbed Energy released as a photon

  34. Energy Levels • Electrons can only move to certain energy levels • Every element releases a photon or photons with unique amounts of energies • This can be used to identify elements • ex: Was used to identify elements in the sun

  35. Low energy High energy Low Frequency High Frequency Electromagnetic Spectrum- the range of wavelengths or frequencies over which electromagnetic radiation extends. X-Rays Radio waves Microwaves Ultra-violet Gamma Rays Infrared . Visible Light (ROY G BIV) Long Wavelength Short Wavelength

  36. Quantum Numbers • 4 Quantum Numbers • The 4 quantum numbers describe the exact location of an electron • It is like the “address” of an electron • Quantum numbers can also be written as ranges to describe a set of electrons

  37. 1st Quantum Number • Principal Quantum Number (n) – Describes the size and energy level of an orbital • The values of n are integers ≥ 1 • The greater the n… the greater the Energy… the greater the distance from nucleus & the less stable

  38. 2ndQuantum Number • Angular Momentum Quantum Number (l) - Defines the subshell or the shape of the orbital • Allowed values of l are integers ranging from 0 to n-1 • s subshell l = 0 • p subshell l = 1 • d subshell l = 2 • f subshell l = 3

  39. 3rd Quantum number • Magnetic Quantum Number (ml) – describes the 3D orientation of an orbital • Values of ml range from - l to + l • Each box in the orbital diagram indicates a different ml value • Example: l value for p is 1 so ml range from -1 to 1

  40. 4th Quantum number • Spin Quantum Number (ms) – describes the electron’s magnetic spin direction • Values for ms are + ½ (spin up/ clockwise) or – ½ (spin down/ counterclockwise) • The direction of the arrow in the orbital diagram determines the ms value

  41. Quantum Numbers - Practice • What are the 4 quantum numbers of: • The last electron in P • The 5th electron in Na • The 9th electron in Cl • The 5th electron in O • The last electron in Cu • The last electron in Fe

  42. Summary for electrons… • Draw all of the following for H • Lewis Dot Diagram—simplified notation showing only valence electrons • Electron Configuration—the shorthand (in numbers and letters) that represent the locations of electrons in the atom • Orbital Diagram— a visual representation of electron configuration • Quantum Numbers- The exact location of an electron or of valance electrons

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