TAKING A CLOSER LOOK AT ELECTRON ORBITALS

Transcription

1 TAKING A CLOSER LOOK AT ELECTRON ORBITALS

2 Mind Catalyst E - Analogy Share-Out Assuming that each pea represents an electron, the bull s eye represents the nucleus, and the area around the bulls eye represents the area around the nucleus, what happens to the probability of finding an electron as the distance from the nucleus increases? Where is there zero probability of finding an electron? What is the overall shape that the spots collectively made on the target sheets? What would this overall shape (in 3-D) correspond to in an atom? Did this shape look familiar to you as you completed the webquest for homework?

3 So, What Did We Cover Last Class? Louis de Broglie proposed that small particles of matter (like electrons) exhibit observable wave-like properties Furthermore, Werner Heisenberg formulated the Uncertainty Principle that states it is impossible for us to know an electron s exact position (where it is) and momentum (where it is going) As a result, we cannot identify specific orbits that electrons travel in as referenced in Bohr s atomic model Instead, Erwin Schrodinger developed a complex math equation based on the Heisenberg Uncertainty Principle and Louis de Broglie s work that described the energy and position of electrons in an atom

4 So, What Did We Cover Last Class? The solutions to the equation tell us: The energy level (n) where the electron resides n = 1, n = 2, n = 3, etc... The regions of space within an atom where an electron is most likely to be found ORBITALS! AND the shape of these orbitals In the webquest completed for homework, you explored these four orbitals and their energies!

5 The s-orbital Spherical shape Seen in all energy levels Lowest in energy Every energy level has one s orbital Each s-orbital can hold up to 2 electrons

6 The p-orbital y-axis z-axis p (x) x-axis p (y) There are three p orbitals: p x, p y and p z All three are dumbbell-shaped Seen in 2 nd energy level and above Can hold up to 2 electrons PER SUBORBITAL (6 electrons total) p (z)

7 The d-orbital Five clover-shaped orbitals Can hold up to 2 electrons per suborbital (10 electrons total) Seen in 3 rd energy level and above

8 The f-orbital Seven equal energy orbitals Shape is not well-defined Each suborbital can hold up to 2 electrons (14 electrons total) Seen in 4 th energy level and above

9 s p d

10

11 Arrangement of Atomic Orbitals The orbitals of an atom are LAYERED!

12 In Summary Each energy level contains orbitals that layer n = 1 made of 1s n = 2 made of 2s and 2p n = 3 made of 3s, 3p, and 3d n = 4 made of 4s, 4p, 4d, and 4f n = 5 and above made of s, p, d, and f-orbitals as well AND each orbitals are made of sub-orbitals s-type orbitals are made of 1 sub-orbital Holds 2 electrons p-type orbitals are made of 3 sub-orbitals Holds 6 electrons TOTAL (2 electrons per sub-orbital) d-type orbitals are made of 5 sub-orbitals Holds 10 electrons TOTAL (2 electrons per sub-orbital) f-type orbitals are made of 7 sub-orbitals Holds 14 electrons TOTAL (2 electrons per sub-orbital)

13 An Introduction to Electron Configurations HOW DO THE ORBITALS FILL UP WITH ELECTRONS?

14 The Big Questions Now that we know how electrons fit into the atomic structure puzzle, we can now answer the questions: How are they arranged into orbitals and suborbitals? How can we communicate the arrangement of atoms in orbitals? What are valence electrons and why are they more important than other electrons?

15 Assigning an Electron s Address Explore Complete the Electron Configurations activity

16 ELECTRON CONFIGURATIONS

17 Predicting Electron Locations Before we begin How do you determine the number of electrons in an atom again? In a NEUTRAL atom, # of protons from periodic table = # of electrons We show the way electrons are arranged in atoms by writing electron configurations The electron configuration of an atom is the complete description of the orbitals occupied by all of its electrons There are rules to follow!

18 Rule #1 Aufbau Principle Electrons are added one at a time to the lowest energy orbitals, or subshells, available until all the electrons of the atom have been accounted for aufbau is German for build up or construct 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f

19 Energies of Electron Orbitals

20 Rule #2 Pauli s Exclusion Principle An orbital can hold only two electrons In other words, no two electrons can ever be in the same place at the same time The electrons MUST HAVE OPPOSITE SPINS Electrons are associated with spin, either one way or the other like a top These spins are called spin up and spin down So, maximum number of electrons held in each orbital is as follows: 2 for s 6 for p (2 x 3 p-orbitals) 10 for d (2 x 5 d-orbitals) 14 for f (2 x 7 f-orbitals)

21 Rule #3 Hund s Rule Electrons must fill a subshell such that each orbital has a spin up electron before they are paired with spin down electrons More energetically favorable A bus analogy: If you enter a bus and don t know anyone on it, you will pick a seat that is completely empty rather than one that already has a person in it i.e. Electrons are unfriendly!

22 In Summary Electrons fill in order from lowest to highest energy Ground floor first then up The Pauli exclusion principle holds An orbital can hold only two electrons Two electrons in the same orbital must have opposite signs (spins) You must know how many electrons can be held by each orbital Hund s rule applies The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons for a set of orbitals By convention, all unpaired electrons are represented as having parallel spins with the spin up

23 Applying the Electron Rules Lucky for you, you don t have to memorize the order in which electrons fill the orbitals (subshells) You can just use the periodic table! The PT follows the Aufbau principle notice (n-1) d-orbitals are filled after ns and before np orbitals AND (n-2) f-orbitals are filled after ns! Each element square represents ONE electron in that particular orbital Start with H and move through the table in order until the desired element is reached!

24 Illustrating Electron Configurations There are several ways to represent electron configurations: Full electron configurations Condensed (noble gas) electron configurations Orbital diagrams

25 Full Electron Configurations In most cases, it is sufficient to write a list of all of the occupied subshells and indicate the number of electrons in each subshell with a superscript. H 1s 1 C 1s 2 2s 2 2p 2 Ar 1s 2 2s 2 2p 6 3s 2 3p 6

26 Practice! #7 and #10 on page

27 There s a Quicker Way Introducing Condensed (Noble Gas) Notation Noble gases have full valence shells, so you can condense electron configurations to eliminate the electrons already accounted for by the closest noble gas It cuts down on a lot of writing, and that s a good thing! To use noble gas notation: Write the symbol for the preceding noble gas [in brackets] to represent all of the electrons in its electron configuration Add the rest of the electrons at the end Example -the full configuration for As-(Arsenic) is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 Notice, the part in red is the same as Argon s configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 The noble gas configuration will start with the gas in the row before it [Ar] 4s 2 3d 10 4p 3

28 Practice! #13 and #17 on page

29 Orbital Diagrams An orbital box diagram goes one step further by also illustrating the spins of the electrons This notation uses boxes to represent orbitals One arrow ( ) represents 1 e - 2 arrows ( ) represent 2 e - Same rules as full electron configurations apply!

30 Example Phosphorus Full Electron Configuration 1s 2 2s 2 2p 2 2p 2 2p 2 3s 2 3p 1 3p 1 3p 1 Phosphorus Orbital Diagram 1s 2s 2p 3s 3p The orbital box diagram indicates that the three electrons in the 3p subshell all have parallel (unpaired) spins

31 Order of Orbital Filling 5d 5f 5p 7p 3d 5s 4f 7s 3p 4d 3s 4s 4p 6s 6p 6d 2s 1s 2p Electron Configuration for Platinum (Element 78) 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 8

32 What About Valence Electrons? When considering the principal electron shells (n = 1,2,3, ), there are two types of electrons: Core Electrons: electrons in the filled inner shell(s) of an atom Valence Electrons: electrons in the unfilled outer shell of an atom All elements in the same group have similar chemical properties because they have the same number of valence electrons in their outer shell!

33 What Orbitals Correspond to Valence Electrons Location? For elements in the first three periods: The core electrons are those in the preceding noble gas configuration The additional electrons in the outer shell are the valence electrons Example: Full Configuration - B 1s 2 2s 2 2p 1 Noble Gas Config - B [He] 2s 2 2p 1 Core: 1s 2 Valence: 2s 2 2p 1 (Shell with n = 1) (Shell with n = 2)

34 Location of Valence Electrons For elements in the fourth period and below in groups 3A 7A, the filled d subshells are also part of the core, even though they are not included in the noble gas configuration Full Config - Se 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 4 Noble Gas Config - Se [Ar] 4s 2 3d 10 4p 4 Core: Valence: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 4 ONLY ELECTRONS IN HIGHEST NUMBERED s- AND p- ORBITALS ARE VALENCE ELECTRONS!