1 / 33

Electrons in Atoms: Electron Configuration

Electrons in Atoms: Electron Configuration. Chemistry EQs: What is the relationship between matter and energy? How does the behavior of electrons affect the chemistry of atoms ?.

kasie
Download Presentation

Electrons in Atoms: Electron Configuration

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Electrons in Atoms: Electron Configuration Chemistry EQs: What is the relationship between matter and energy? How does the behavior of electrons affect the chemistry of atoms?

  2. SC3 Students will use the modern atomic theory to explain the characteristics of atoms.b. Use the orbital configuration of neutral atoms to explain its effect on the atom’s chemical properties.f. Relate light emission and the movement of electrons to element identification. GPS

  3. Aufbau Principle • Pauli Exclusion Principle • Hund’s Rule • Electron Configuration • Valance Electron • Energy levels • Lewis Structure • Ground state • Excited state • Orbitals • Quantum Vocabulary

  4. Excited Electrons

  5. Bohr’s Model: electrons orbit the nucleus; only orbits in certain energies are permitted. • Ground State- lowest E level • Excited State- Higher than ground state. • The e- are raised to the next level, then release light when they return to ground state • Must have enough E to raise to next level or won’t happen. Bohr Model of the Atom:

  6. Bohr Model

  7. Bohr Model of the atom

  8. DeBroglie- • Quantum Mechanics- light behaves as wave & particles • Visible objects (baseball) have  too small to see, need very small object to detect  • Heisenburg Uncertainty Principle- Can’t know the position & speed of electron at the same time Electrons

  9. E- do not fall towards nucleus • E- reside in electron clouds called orbitals Bohr Model  Quantum Model

  10. Energy levels- region around nucleus where e- likely to be found (electron density is high) • Quantum- amount of energy for e- to jump levels • Continuous- ramp, no units • Quantitized- fixed levels, fixed units Electrons

  11. Schrodinger- estimates the probability of e- to be in certain area; electrons are like a fuzzy cloud, but more dense= more likely to find e- 90% of the time in a particular location Orbitals-Wave functions with corresponding densities (shape and energy) **orbital is NOT the same as Bohr’s orbit

  12. Row=Period=Energy Level => horizontal Column= Group/ Family=> vertical • Elements in the same family have the same # of valence electrons & share similar chemical properties. Valence electrons= e- in to last energy level. • Valence e- correspond w/ group # (does not include transition elements): • Group 1A= 1 valence e- • Group 2A= 2 valence e- • Group 3A (13) = 3 valence e- • Group 4A (14) = 4 valence e- • Group 5A (15) = 5 valence e- • Group 6A (16) = 6 valence e- • Group 7A (17) = 7 valence e- • Group 8A (18) = 8 valence e-  FULL SET; STABLE Electron configuration

  13. Tells the arrangement of electrons around the nucleus of an atom • Written in ground state, which is the lowest energy & most stable arrangement • e- arrange from lowest to highest E level Electron Configuration

  14. Orbital: the 3-D space around the nucleus that describes an electrons probable location. • Energy Level (n): indicate the relative sizes & energies of atomic orbitals. • As n increases, the orbitals become larger, and the e- spends more time farther from the nucleus. • n = major energy levels; n = 1-7; correspond w/ the 7 rows on the P.T. • Sublevel: energy levels contained w/in a energy level; s, p, d, f Electron configuration.

  15. Electron Configuration

  16. s and p Orbitals What shape are the orbitals?

  17. Tells the arrangement of electrons around the nucleus of an atom • Written in ground state, which is the lowest energy & most stable arrangement • Follows 3 rules: • Aufbau Principle • Pauli Exclusion Principle • Hund’s Rule Electron Configuration

  18. Each electron will occupy the lowest available energy level 1st, then higher energy levels. Aufbau Principle

  19. A maximum of 2 electrons with opposite spins can fit in an orbital (No more than 2 e- can occupy orbitals). • e- in the same orbital must have opposite spin (repulsion); • Show each orbital w/ its own box • One is spinning clockwise & the other is counter clockwise, Show this with one arrow going up & one pointing down. NOT Pauli Exclusion Principle   

  20.      • Single electrons with the same spin must occupy each equal energy level before additional electrons with opposite spins can be added • e- enter orbitals singularly, then pair up • Example: when filling the p sublevel with 4e-, each box gets 1 before doubling up one box NOT Hund’s Rule

  21. Using the PT: • The principal quantum number , n = period. • There are 4 blocks: (s, p, d, f) • Noble gases:full s & p level making them inert (stable). • Alkali Metals- s1 • Alkaline Earth Metals- s2 • Transition Elements- d1-d10 • Inner Transition Elements- f1-f14 Electron Configuration-

  22. Electron Fill Sequence

  23. Electron Configuration

  24. 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p Follow the yellow brick road Electron Sequence Model

  25. Examples: • F 1s22s22p5 • Cl 1s22s22p63s23p5 • Al 1s22s22p63s23p1 • Br 1s22s22p63s23p64s23d104p5 Electron Configuration

  26.      1s 2s 2p 3s 3p 4s 3d 4p          1s 2s 2p 3s 3p 4s 3d 4p        1s 2s 2p 3s 3p 4s 3d 4p                   1s 2s 2p 3s 3p 4s 3d 4p Orbital Diagrams”Uses boxes to represent orbitals

  27. Simplified version of writing e- configurations • Noble gas is placed in brackets [ ] Noble Gas Notation

  28. When doing configurations for large numbers of electrons, we can take a short cut using noble gases. (yay!) • Example, lets try Sulfur: (16 electrons) • The noble gas that comes before Sulfur is: Neon • Noble gas is placed in brackets [ ] • Place noble gas in bracket to represent the e- configuration up to that noble gas. • Write the rest of the e- config. for that element. • So we could shortcut by writing: [Ne] 3s2 3p4 Now you try: Manganese and Strontium Noble Gas Shortcut

  29. Electron Configuration and Orbital Notation

  30. e- in the outer most energy level that determines chemical properties • Lithium = 1s2 2s1 • Bromine = 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 • Aluminum = [Ne] 3s2 3p1 • Family Number • VE are the electrons available to form chemical bonds w/ other elements. Valence Electrons

  31. Lewis Dot Structure: Electron Dot • Chemical symbol & valence electrons of atom • Valuable in showing how atoms share electrons in covalent bonds • We can draw Lewis structures for every element using valence electrons • Count the # of valence electron, then arrange then around the symbol for the atom one at a time; up to 8 electrons. • Arrange 1/ side around the symbol, then couple up if more than 4 electrons.

  32. Lewis Dot

More Related