In This Lesson: Valence Electrons and Lewis Dot Structures (Lesson 4 of 4) Today is Tuesday, October (!) 4 th, 2016 Stuff You Need: Periodic Table Pre-Class: You ve probably heard of the special name we give to electrons in the outermost principal quantum number. Do you remember what it is?
Today s Agenda Study Guide for Core Assessment Valence Electrons Lewis Dot Structures Where is this in my book? P. 194 and following
By the end of this lesson You should be able to draw valence electrons as Lewis Dot Structures.
Valence Electrons Valence electrons are the ones available for bonding. Notice where they are:
Da portant Stufz Valence electrons are electrons in the outermost shell (highest energy level). The highest coefficient. IMPORTANT NOTE: d and f sublevel electrons do not figure in bonding because we only look at the highest principal quantum number electrons.
Finding Valence Electrons To find valence electrons, simply perform the usual electron configuration notation. Find the sublevel(s) with the highest principal quantum number. Count the electrons there, ignoring d or f sublevels, if any. Example: 1s 2 2s 2 2p 4 2s 2 2p 4 are the valence electrons (6 total)
Determining Valence Electrons Let s try some. Grab your periodic tables and whiteboards. Tell me the number of valence electrons in the following elements [next slide].
Valence Electrons Element Electron Configuration # of Valence Electrons/Capacity Oxygen (O) 1s 2 2s 2 2p 4 6/8 Hydrogen (H) 1s 1 1/2 Xenon (Xe) Rubidium (Rb) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 8/8 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 1/8 Helium (He) 1s 2 2/2 Boron (B) 1s 2 2s 2 2p 1 3/8 Carbon (C) 1s 2 2s 2 2p 2 4/8 Fluorine (F) 1s 2 2s 2 2p 5 7/8 Aluminum (Al) 1s 2 2s 2 2p 6 3s 2 3p 1 3/8
More Practice Electrons Review Worksheet Do the first page (landscape orientation), but only these columns: Atomic Number Electron Configuration Number of valence electrons
The Octet Rule One other thing Remember when we said that atoms want a full valence electron shell like those super awesome noble gases? Well, atoms want to be like noble gases because a full valence shell makes them more stable than having a partial valence shell. And who doesn t like stability? They do this by adding or dropping electrons.
The Octet Rule What is the electron capacity of a full s sublevel plus a full p sublevel? 8 (s=2, p=6) This idea, of having 8 electrons in the valence shell to be full, is called The Octet Rule. Note: Hydrogen and helium are exceptions. What is capacity for their valence shells? 2, so they only want 2 electrons to be stable.
Group IA (alkali metals) have 1 valence electron (1+)
Group IIA (alkaline earth metals) have 2 valence electrons (2+)
Group IIIA elements have 3 valence electrons (3+)
Group IVA elements have 4 valence electrons (4+)**
Group VA elements have 5 valence electrons (3-)
Group VIA elements have 6 valence electrons (2-)
Group VIIA (halogens) have 7 valence electrons (1-)
Group VIIIA (Noble gases) have 8 valence electrons, except helium, which has only 2 (no charge)
Transition metals ( d block) have 1 or 2 valence electrons (1+ or 2+)
Lanthanides and actinides ( f block) have 1 or 2 valence electrons (1+ or 2+)
About transition metals Transition metals do weird things. Yes, they do have 1 or 2 valence electrons, but they form lots of different ionic charges. The first thing to be aware of is that while full energy levels are the most stable, half-filled sublevels are still mostly stable. To understand this better, let s take a look at Cu, Fe, and Mn.
Copper (Cu) Copper s valence orbital notation: 4s 3d So you can see that copper has a full s sublevel but only an almost full d sublevel. We would expect it to drop it the two 4s electrons, making a charge of 2+. However, because a half-filled sublevel is preferable to this setup, Copper flips one electron up from the 4s sublevel to fill 3d.
Copper (Cu) Then, copper can just drop the s electron. 4s 3d As a result of these two possibilities, copper can have two possible ionic charges: 2+ or 1+.
Iron (Fe) Iron does something kinda similar: 4s For iron, the first thing it can do is drop both s electrons. 2+ ion. Or, it could drop both s electrons and one d. 3+ ion. 3d
Fun Fact: Iron (Fe) The fact that iron has four unpaired electrons in its d sublevel is the reason iron is/can be magnetized at room temperature (along with cobalt and nickel and others). Having four electrons spinning all in the same direction makes for easy magnetic field induction, and it s called having an orbital magnetic moment. Note that there are other factors at play here, one of which is the sea of electrons concept you ll learn next unit.
Manganese (Mn) Now for Manganese: 4s Manganese can drop both s electrons. 2+ ion. Or, it can drop all 7 electrons. 7+ ion. 3d Or about five other possibilities!
Multivalent Elements These, and other metals, are multivalent they have several different configurations of their valence electrons. Therefore, they form multiple charges. Here s a present for you a periodic table with a listing of multivalent metals: Periodic Table Polyatomic Ions and Multivalent Elements Only
About Group IVA Group IVA also does some weird things. Carbon, for example, would like to either gain or lose 4 electrons. But how many does it have total? 6 So gaining/losing 4 is kinda hard for such a small atom. Even Si is too small. So C and Si share electrons instead of losing or gaining.
About Group IVA However, Ge, Sn, and Pb are all big enough to ionize. Their outer electrons are very far away, and what s 4/82 to lead? So, Ge tends to lose all four valence electrons. 4+ ion. Sn and Pb either lose all four 4+ ion. or just lose the p sublevel electrons. 2+ ion. So Sn and Pb have two different possible oxidation states (ionic charges).
About Group IVA In addition to Group IVA, other nearby large elements do the same sort of thing: Antimony (Sb) 3+ or 5+ 3+ = p dropped; 5+ = s and p dropped. Bismuth (Bi) 3+ or 5+ 3+ = p dropped; 5+ = s and p dropped. Thallium (Tl) 1+ or 3+ 1+ = p dropped; 3+ = s and p dropped. Polonium (Po) 2+ or 5+ 2+ = one from p dropped, one from s dropped; 5+ = all of p dropped and one from s.
More Practice Electrons Review Worksheet Finish the first page.
Lewis Dot Notation Lewis Dot Notations for Period 2 elements. Lithium Beryllium Boron Carbon Li Be B C Nitrogen Oxygen Fluorine Neon N O F Ne
Creating Lewis Dot Structures Step 1: Determine the number of valence electrons. Step 2: Draw them around the element abbreviation one-by-one. Step 3: Check your answer. Make sure you only have electron pairs if you already have four electron singles.
Lewis Dot Practice Cl Se Al K Si Ca
Practice Electrons Review Worksheet Try the reverse side. Tough ones: Zn, Ag, Fe (TRY THEM!)
Summary Valence Electrons Electrons in the outermost energy level (highest n number) just s and p sublevels. Electrons can be shown via Lewis Dot Diagrams. Octet Rule Atoms react to get eight electrons in their outermost shell. Except H and He. That makes em stable.
Closure Part 1 Draw the dot structure for the element Bromine: Draw the dot structure for the element Thallium: Draw the dot structure for the element Selenium: Draw the dot structure for the element Magnesium:
Closure Part 2 Draw the dot structure for the element Potassium: Draw the dot structure for the element Helium: Draw the dot structure for the element Aluminum: Draw the dot structure for the element Hydrogen:
Closure Part 3 If you have an atom on the right side of the table, let s say Chlorine, how many electrons does it need to get to 8? 1 From where might it get that electron? A cation. And in which group would that cation be? Alkali metals (Group I), because they each have one electron they d like to give away. And what would be a possible donor element? Na, Li, K, Rb, Cs, et cetera. They make salts like in our flame tests!
Closure Part 4 Electron Configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 1 Shorthand Electron Configuration: [Kr] 5s 2 4d 10 5p 1 Valence Electron Configuration: 5s 2 5p 1 Orbital Notation: [arrows] Dot Notation: Uh dots.