Lewis Dot Diagrams

1. Lewis Dot Diagrams Mrs. Kay

2. Energy in the atom • Charged particles called electrons surround the nucleus in regions called shells or energy levels

3. Atoms of different elements have different numbers of electrons • Each shell is “full” before electrons move to the next shell.

4. Period number: number of electron shells • Shells will fill in the following order: 2, 8, 8,18 • Group number: number of valence electrons (look at the second digit of the group number)

5. Trends: • Elements of the same group have similar properties because they have the same number of electrons in their outer shell. • The outer shell = valence shell • Electrons that occupy the valence shell are called valence electrons

6. Noble gas configuration • Group 18, the noble gases are the most stable of elements because their valence shell is full with electrons • Less energy required • Stable octet

7. Lewis Dot diagrams • Visual representation of an element and only its valence electrons • Example: sodium, Na has one valence so it has 1 dot representing that electron. • Chlorine has 7 electrons. • Electrons get paired up along 4 sides of the element.

8. Bonding • Ionic bonding: attraction between oppositely charged ions formed when metallic ions (+) transfer electron(s) to nonmetallic ions (-) • Ex: NaCl

9. Not always 1:1 ratio, sometimes need to use subscript to show the number of atoms Ex: CaCl2 The 2 is a subscript, it shows that 2 atoms of chlorine bond with one atom of calcium. The charges need to have a sum of zero

10. Covalent / Molecular bonding • Formed when a pair of electrons is shared by two non-metal atoms. • A single bond is formed when one pair of electrons is shared. • Example: diatomic molecules, like O2, H2, N2, F2, Cl2…

11. Single bonds

12. Polar Covalent Bonding: • Electrons are shared unequally. • They are not ionic, because the electron is not totally removed because there was not enough attraction to totally remove the electron.

13. Polar covalent bond

14. When you have a polar covalent bond, one atom will attract an electron closer than the other. • Called polar molecules because one end is slightly positive and one end is slightly negative

15. Water is Polar! • Water is very common polar compound. • The hydrogen are slightly positive (because almost gave up e-) • The oxygen is slightly negative (because almost took the e-)

16. Naming Compounds: the rules! Mrs. Kay

17. 1st: What are polyatomic ions? • Ions that are made of multiple atoms covalently bonded together. • We treat them like a unit • Example: sulfate, SO4-2so sulfur and oxygen form a negative ion (-2)

18. Given the Formula: Fe2(SO4)3 • Step One: The cation is always the first thing you see in the name, and the anion is always the second thing. In this case, you should recognize that Fe is "iron", and that SO4 is the "sulfate" ion. Generally, if one of these ions has more than one atom in it, you'll need to look it up in a chart. • you need to know eight of the polyatomic ions : hydroxide, nitrate, nitrite, sulfate, sulfite, carbonate, phosphate, ammonium.

19. Step Two: Figure out if you need a Roman numeral in the name.If the cation in the compound you're naming is not a transition metal (group 3-12), then you definitely don't need to use a Roman numeral and the naming is done. If there is, then you need to figure out whether or not the cation can exist in more than one charge. If not, then you don't need a Roman numeral. If so, then move on to Step Three...

20. Step Three: Figure out what the Roman numeral should beBasically, this should be fairly easy. A good rule of thumb is that usually the number of anions you have in the molecule is equal to the charge on the cation, and that the number of cations you have is equal to the number of anions. Using our example, there are three sulfate ions, meaning that iron has a charge of +3. Likewise, since there are two iron atoms, the sulfate has a charge of -2. Since iron has a charge of +3 in this compound, the name in this example is iron (III) sulfate

21. Given the name, find the formula: • Step One: Translate the name into the ionsIn copper (II) fluoride, the cation is the copper (II) ion and the anion is the fluoride ion. Hopefully, you realize that the copper (II) ion is simply Cu2+ and the fluoride ion is F-. If not, then you need to go back and review the rules for naming ions above.

22. Step Two: Put brackets around the ions, but leave the charges on the outside.In this case, the copper (II) ion would be [Cu]2+ and the fluoride ion would be [F]-1. Never change anything in these brackets, ever! • Step Three: Put the ions next to each other.When we do this here, we get [Cu]2+[F]-1

23. Step Four: Cross the charges:The charge on the cation will be equal to the number of anions you have, and the charge on the anion will be equal to the number of cations you have. In our example, you should realize that we have one copper atom (because the charge on fluorine is -1) and two fluoride ions (because the charge on copper is +2). This gives us a formula of: [Cu][F]2 • Step Five: Take the brackets away. The final formula for copper (II) fluoride is then CuF2

24. Naming Covalent compounds: • All covalent compounds have two word names.  The first word typically corresponds to the first element in the formula and the second corresponds to the second element in the formula except that "-ide" is substituted for the end.  As a result, HF is named "hydrogen fluoride", because hydrogen is the first element and fluorine is the second element

25. If there is more than one atom of an element in a molecule, we need to add prefixes to these words to tell us how many are present.  Here are the prefixes you'll need to remember

26. Test what you know here:http://science.widener.edu/svb/tutorial/namingcsn7.html