1. Chapter 5 Electrons in Atoms

2. Bohr In 1913 Bohr published a theory about the structure of the atom based on an earlier theory of Rutherford's. Rutherford had shown that the atom consisted of a positively charged nucleus, with negatively charged electrons in orbit around it

3. Bohr’s “Quantum Mechanical Model” of the atom Bohr expanded upon this theory by proposing that electrons travel only in certain successively larger orbits He suggested that the outer orbits could hold more electrons than the inner ones these outer orbits determine the atom's chemical properties

4. The Periodic Table: organizes elements by atomic number and… Groups/families: elements have the same physical and chemical properties. Rows/periods: elements have the same number of electron shells. Organizing atoms in the periodic table

5. Practice Question 1 1.Name another element that would have similar chemical properties to chlorine. 2.Name an atom that is in the same period as chlorine.

6. Electrons All atoms have an equal number of protons and electrons –Atoms are electrically neutral Atoms have no charge Symbol: Ne An equal number of positive protons and negative electrons results in zero charge

7. Practice Question 2 How many electrons orbit: –A magnesium atom? –A sulfur atom? –A hydrogen atom?

8. 1 valence e - 4 valence e - Valence electrons are: responsible for chemical behavior of atom used for chemical bonding located in the outer electron shell

9. study question 5 Total number of electrons Number of valence electrons nitrogen phosphorus

10. There are 7 possible energy levels where an electron can be found. Energy levels are represented by the periods (horizontal rows) on the periodic table The total number of electrons that can fit in an energy level can be found using the equation: 2n 2 (where n = an energy level 1-7)

11. Within each energy level, there are sublevels –s sublevel- consist of e - from groups 1 and 2 –p sublevel- consist of e - from groups 13-18 –d sublevel- consists of e - from groups 3-12 –f sublevel- consists of e - from the lanthanide and actinide series

12. s p d (n-1) f (n-2) 12345671234567 6767

13. Orbitals Each sublevel can be broken down into orbitals: –s sublevel: has 1 orbital –p sublevel: has 3 orbitals –d sublevel: has 5 orbitals –f sublevel: has 7 orbitals Each orbital can only hold a maximum of 2 electrons.

14. Orbitals Complete the chart in your notes as we discuss this. The first level (n=1) has an s orbital. It has only 1. There are no other orbitals in the first energy level. We call this orbital the 1s orbital. Energy Level Sub- levels Total OrbitalsTotal Electrons Total Electrons per Level n = 1s1 (1s orbital)22 n = 2spsp 1 (2s orbital) 3 (2p orbitals) 2626 8 n = 3spdspd 1 (3s orbital) 3 (3p orbitals) 5 (3d orbitals) 2 6 10 18 n = 4spdfspdf 1 (4s orbital) 3 (4p orbitals) 5 (4d orbitals) 7 (4f orbitals) 2 6 10 14 32

15. Where are these Orbitals? http://www.biosulf.org/1/images/periodictable.png 1s 2s 3s 4s 5s 6s 7s 3d 7p 6p 5p 4p 3p 2p 5f 4f 6d 5d 4d

16. Electron configurations Electron configurations are similar to postal “zipcodes”. –They represent a general area where an electron can be found. Examples Hydrogen has 1 electron: 1s 1 He has 2 electrons: 1s 2 Li has 3 electrons: 1s 2 2s 1 Be has 4 electrons: 1s 2 2s 2 B has 5 electrons: 1s 2 2s 2 2p 1

17. Electron Configurations

18. Rules for Electon Configurations 3 rules govern electron configurations. 3 rules govern electron configurations. – Aufbau Principle – Pauli Exclusion Principle – Hund’s Rule

19. Each line represents an orbital. 1 (s), 3 (p), 5 (d), 7 (f) Low Energy High Energy

20. The Pauli Exclusion Principle states that an atomic orbital may have up to 2 electrons and then it is full. The Pauli Exclusion Principle states that an atomic orbital may have up to 2 electrons and then it is full. Wolfgang Pauli, yet another German Nobel Prize winner

21. Don’t pair up the 2p electrons until all 3 orbitals are half full.

22. ElementConfigurationElementConfiguration H Z=11s 1 He Z=21s 2 Li Z=31s 2 2s 1 Be Z=41s 2 2s 2 B Z=51s 2 2s 2 2p 1 C Z=61s 2 2s 2 2p 2 N Z=71s 2 2s 2 2p 3 O Z=81s 2 2s 2 2p 4 F Z=91s 2 2s 2 2p 5 Ne Z=101s 2 2s 2 2p 6 (2p is now full) Na Z=111s 2 2s 2 2p 6 3s 1 Cl Z=171s 2 2s 2 2p 6 3s 2 3p 5 K Z=191s 2 2s 2 2p 6 3s 2 3p 6 4s 1 Sc Z=211s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 Fe Z=261s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 Br Z=351s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 Note that all the numbers in the electron configuration add up to the atomic number for that element. Ex: for Ne (Z=10), 2+2+6 = 10

24. Electron Configurations ElementConfiguration H Z=11s 1 Li Z=31s 2 2s 1 Na Z=111s 2 2s 2 2p 6 3s 1 K Z=191s 2 2s 2 2p 6 3s 2 3p 6 4s 1 This similar configuration causes them to behave the same chemically. It’s for that reason they are in the same family or group on the periodic table. Each group will have the same ending configuration, in this case something that ends in s 1. This similar configuration causes them to behave the same chemically. It’s for that reason they are in the same family or group on the periodic table. Each group will have the same ending configuration, in this case something that ends in s 1.