2. Review Questions  On your own, complete the 14 review questions  You may use the internet or your notes to complete them.

3. Polyatomic Ions  A polyatomic ion is a charged chemical ion composed of two or more atoms covalently bonded that can be considered to be acting as a single unit.  Ions with a positive charge are called cations  Ions with a negative charge are called anions  These ions are frequently given specific names.  For example: Ammonia is NH 3 and is neutral but the Ammonium ion is NH 4 and has a charge of +1

4. Important Ions

5. Bonds  A chemical bond is an attraction between atoms that allows the formation of chemical substances that contain two or more atoms.  Ionic bonding is a type of chemical bond that involves the electrostatic attraction between oppositely charged ions.  A covalent bond is a chemical bond that involves the sharing of electron pairs between atoms.

6. Bonds  A chemical bond is an attraction between atoms that allows the formation of chemical substances that contain two or more atoms.  Ionic bonding is a type of chemical bond that involves the electrostatic attraction between oppositely charged ions.  A covalent bond is a chemical bond that involves the sharing of electron pairs between atoms.

7. Ionic versus Covalent

8. Naming Compounds  We name ionic and covalent compounds differently.  General rule for ionics: Metal first (cation), Non- metal (anion) second.  Example: Sodium and Chlorine become Sodium chloride  Example: Potassium and the Permanganate ion become Potassium Permanganate.  General rule for covalents: Use prefixes like mono, di, tri, etc.  Example: Carbon and 2 Oxygens become Carbon dioxide.  Example: 4 Sulfurs and 2 Chlorines become Tetrasulfur dichloride

9. Basic Info – Ionic Bonds  An ionic bond involves the transfer of electrons between a non-metal and a metal  These ions have electrostatic attraction.  Example : Sodium (cation – metal) and Chlorine (anion – non-metal)  Sodium has 1 valence electron. Chlorine has 7 valence electrons  Sodium “gives” its electron to chlorine  Compound is Sodium Chloride

10. More examples  Example : Barium and Bromine  Barium has _____ valence electron. Bromine has ____ valence electrons  How many bromines are needed per barium?

11. Trickier example  Example : Aluminum and a Sulfate ion  Aluminum has ______ valence electron(s) and an oxidation # of ___________.  Sulfate has an oxidation # of ______ (look at the list of polyatomic ions!)  How many Aluminums and Sulfates will be needed?  FIND THE LOWEST COMMON MULTIPLE…..

12. Lewis Dot Model  Also called Electron Dot Structures Models  Represent atoms and ions with dots representing the valence electrons  Typically only the s and p orbital = octet rule  Rules:  a) Metals should have dots only.  b) Non-metals have spaces and dots.  c) Period 1 atoms follow the duet rule and Periods 2 & 3 follow the octet rule.  Dot Models for Period 4-7 on are not very accurate because of the d and f orbitals

13. Examples  Phosphorus: Group 1 5 so 5 valence e’s  Non-metal, so would like 3 more e’s  Scandium: Group 3 so 3 valence e’s  Metal, so it would like to get rid of those e’s  Phosphorus would like 3 e’s, Scandium would like to get rid of 3 e’s  Scandium Phosphide

14. Naming Ionic Compounds  Binary – any compound with only 2 different elements.  The non-metal name is changed to end with -ide  Example: Calcium and Chlorine  Step 1. What are the oxidation numbers of Calcium and Chlorine  Step 2. Cross the oxidation numbers  Step 3. Name the compound (METAL then NON- METAL)  Calcium Chloride

15. More Practice  What is the name of the compound: NaFl  Oxidation # Na = ? Oxidation # Fl = ?  What is the name of the compound: FeO 2  Oxidation # Fe= ? Oxidation # O = ?  Determine the formulas:  Oxidation # Al = ? Oxidation # O = ?  Formula: Aluminum Oxide ?  Oxidation # Mg= ? Oxidation # S = ?  Formula: Magnesium Sulfide ?

16. Polyatomic Ions  Polyatomic ions are exceptions to the “ide” rule  Sodium hydroxide  What is hydroxide? What is its oxidation #?  What is sodium’s oxidation number?  Potassium permanganate  What is permanganate? What is its oxidation #?  What is potassium’s oxidation number?  Don’t change the way the polyatomic ion is written!

17. Writing the formula 1) Write the symbol for the first element… 2) Place the Oxidation # on the top of the element (VI) = 6 3) Name the polyatomic ion - sulfite – NOTICE MY PARENTHESIS 4) Put the two (metal first) side by side and crossover the charges 5) Question? What do we do with the 2 and the 6? Reduce….. Manganese (VI) Sulfite

18. More practice 1) Write the name for the first symbol… 2) Decide whether or not it needs parenthesis in the middle…all except Group 1+2, Al, Zn, Ag need ( ) 3) Write the name of the polyatomic ion – (B 4 O 7 ) 2- 4) To find the number in the ( ), we need to recognize the ratio of Chromium atoms to Tetraborate molecules 5) Since Tetraborate has ALWAYS got a 2- charge we need to… 6) As you can see, the 1 to 2 ratio became a 2 to 4 ratio and the 4 crosses back up top to give Cr a 4+ Oxidation # 7.) Report your final answer Example: Write the formal name of Cr(B 4 O 7 ) 2

19. Facts about Ionics  Ionic compounds form crystals  They have high melting and boiling points  Ionic compounds are hard and brittle  They conduct electricity when they are dissolved in water  These are called electrolytes or aqueous solutions  They don’t conduct electricity when they are solid

20. Facts about Ionics  Compounds having water attached are called hydrates or hydrous  Compounds that have no water are called anhydrous.  The chemistry storeroom has many compounds that absorb water over time. These are called hygroscopic Special compounds for industrial use that absorb water are called humectants.

21. Naming Hydrates  We name the ionic compound the same way but we add _______hydrate, where the blank is the number of waters.  1 = mono, 2 = di, 3 = tri, 4 = tetra, 5 = penta, 6 = hexa, 7 = hepta, 8 = octa, 9 = nona, 10 = deca  What is Ba(OH) 2 8H 2 O?

22. Acids  Acids are compounds that donate protons (Hydrogen ions) when in solution.  We will learn more about acids later.  Bases are compounds that accept protons  They often have an hydroxide ion in the compound  Most acids have a Hydrogen cation in as the first element in their chemical formula.  Safety: When diluting acids, chemists will always add the concentrated acid to water.  It makes the solution more dilute and less likely to release heat!

23. Rule 1  Acids with Hydrogen as the cation and a non- metal as the anion are named by adding HYDRO prefix to the front and an –ic suffix to the end.  Example : Hydrogen and Fluorine:  HF, called Hydrofluoric Acid  On your own:  Hydrosulfuric Acid  Hydroiodic Acid

24. Rule 2  Acids with Hydrogen as the cation and a polyatomic “ate” group as the anion are named by replacing the –ate suffix with an –ic suffix and adding the term acid at the end.  Example : Hydrogen and Carbonate  Formula for Carbonic Acid: H 2 (CO 3 )  On your own :  Sulfuric Acid  Nitric Acid  Oxalic Acid

25. Rule 3  Acids with Hydrogen as the cation and a polyatomic “ite” group as the anion are named by replacing the –ite suffix with an -ous suffix and adding the term acid at the end.  Example : Hydrogen and Carbonite: write the formula and oxidation number for carbonite:  Formula for Carbonous Acid: H 2 (CO 2 )  On your own :  Sulfurous Acid  Nitrous Acid

26. Rule 4 & 5  Rule 4 & 5 are for super-oxygen rich polyatomic ions called per______ates and for oxygen- depleted polyatomic ions called hypo______ites.

27. Electronegativity  Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons.  Scale goes from 0 to 4.0  Fluorine is the most electronegative atom  Francium is the least.

28. Differences in Electronegativity  When atoms prefer to bond when there is a large difference in electronegativity.  If the difference is greater than 1.8, then the bond is ionic and the electrons are transferred.  If the difference is less, then the bond is covalent and the electrons are shared.  The more equal the electronegativity values are, the more equally the electrons are shared.