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Writing Electron Configurations

Writing Electron Configurations. Ok...let’s simplify this. . Every atom has a nucleus. In that nucleus we have protons (positive charge) and neutrons (no charge) Surrounding the nucleus we have electrons.

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Writing Electron Configurations

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  1. Writing Electron Configurations

  2. Ok...let’s simplify this. • Every atom has a nucleus. • In that nucleus we have protons (positive charge) and neutrons (no charge) • Surrounding the nucleus we have electrons. • Electrons have a negative charge, however they do not “fall” into the nucleus (despite being attracted to the positive proton), why?

  3. Ok...let’s simplify this • Because electrons exist in “orbitals” • An orbital is the region of space in which you will find an electron. (For our purposes) • This is where it gets confusing... • There are different levels of orbitals. • We represent these different levels of orbitals with the letter “n”.

  4. So if an electron is in the 1st orbital surrounding the nucleus it is n=1. • If an electron is in the 3rd orbital surrounding the nucleus it is n=3. • Easy enough, right?

  5. There’s more. • We also have the letter “l” this represents what type of orbital the electron is in. • The different types of orbitals are as follows... • S, p, d, f, and g • These letters refer to the shape of the orbital.

  6. So now, we have two pieces of information to help us determine where an electron is. • We know the number of orbital it is in and we know the shape of the orbital.

  7. Review: Quantum Numbers • Principle Quantum Number (n) represents the energy orbital number (how far the orbital is from the nucleus). • Orbital-Shape Quantum Number Orbital (l) represents the shape of the orbital sublevel (s is round, p is like 2 balloons, 2 d orbitals put together) • Magnetic Quantum Number represents the direction of the orbital sublevel (l)

  8. n can be any number from 1 on. • l can be any number from 0 to (n-1) • ml is –l to +l

  9. Ok...back to new stuff • The Fourth Quantum Number (ms) is used to represent the direction of the spin on the electron. • There are only two possible values for this number, + ½ or – ½ . • Electrons spin like a top.

  10. Electron Spin

  11. Pauli Exclusion Principle • We have a new principle to learn! • The Pauli Exclusion Principle states that no two electrons in any atom will have the same 4 quantum numbers. • For example you would not find two electrons with n=4, l=3, ml=-3 and ms= - ½ • You could find n=4, l=3, ml=-3, ms= - ½ and n=4, l=3, ml= -3, ms= + ½

  12. Example

  13. Electron Configurations

  14. An electron configuration can be written for each atom. • That is, we can write out where we can find each electron in an atom. • And remember, no two electrons will have the same set of 4 quantum numbers in the same atom.

  15. Generally when we write electron configurations we write the configuration of the atom in it’s ground state. • Ground state means that the electron is neutral and is not an isotope.

  16. Example • Hydrogen, in it’s ground state, has one electron (atomic number 1, mass of 1). • The electron configuration for Hydrogen is 1s1. • Remember, n=1, l=0 and the small, raised 1 symbolizes that there is one electron in the s sublevel.

  17. Helium is He 1s2 symbolizing there are 2 electrons in the s sublevel. • Lithium is 1s22s1 symbolizing that there are 2 electrons in the 1 s sublevel and 1 electron in the 2 s level. • Fluorine is written as 1s22s22p5 • Why do you think that there are only 2 electrons in the s sublevel, but there can be 5 electrons in the p sublevel?

  18. Remember the Pauli Exclusion Principle? • What did the Pauli Exclusion Principle tell us? • That no two electrons can have the same 4 quantum numbers in an atom. • Let’s write out the possible values of the 4 quantum numbers in the s sublevel. • How many electrons do you think is possible in the p sublevel? • How about the d sublevel?

  19. There is a maximum of 2 electrons in the s sublevel. • There is a maximum of 6 electrons in the d sublevel. • There is a maximum of 9 electrons in the p sublevel.

  20. Aufbau Principle • Each electron occupies the lowest energy orbital available – For example, all of the “2” orbits must be filled before electrons can go into the 3s orbital • Atom’s are “built-up”

  21. Do you think that the direction of the arrow symbolizes anything? • Why do you think that in the Carbon electron configuration the 2 p level has two electrons in the same direction?

  22. We can tell the electron configuration of an element just by looking at the periodic table! • The “long form periodic table” shows us which sublevel contains the atom’s valence electrons.

  23. A helpful guide to what goes where

  24. Let’s try it. • How would we write the configuration for: • Hydrogen • Helium • Boron • Carbon • Nitrogen

  25. Noble Gas Notation • The last column of the periodic table is called the Noble Gases. • The Noble Gases have full valence electron shells. • So we when we out electron configurations for our elements we can shorten it like this...

  26. Shorten Electron Configuration Notation • Ne – 1s2 2s2 2p6 • Na – 1s 2 2s2 2p6 3s1 – [Ne]3s1

  27. Try it out! • In your notebook write out the electron configuration for the first 12 elements. • We have already done some!

  28. Homework • Read pages 146-148 • Start working on your assignment.

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