1. Ionic Compounds and Acid Nomenclature

2. The force that holds two elements together. - Bonds form to seek the lowest energy state and to meet the maximum number of valence electrons. Chemical Bonds

3. Lewis Dot Compounds

4. Ionic Compounds Ions= atoms or groups of atoms that have a positive or negative charge. The difference in charge holds the two atoms together Negative= gaining an electron (-) Positive= losing or donating an electron (+) 1. Have high melting and boiling points 2. Forms Solids 3. Made from a metal and non-metal.

5. Ionic Compounds If an atom is normally neutral (same # of Protons and Electrons) And It donates one electron away it is no longer neutral but positive. So the Sodium becomes Positive known as a Cation Chlorine becomes Negative known as a Anion Cation wrote as Na ⁺ Anion wrote as Cl¯ Cation=Metals and Positive Anions=Non-Metals and Negative

6. MONOATOMIC IONS- ONLY HAVE ONE ATOM AND HAVE A SPECIFIC (+) OR (-) CHARGE. OXIDATION NUMBER OR IONIC CHARGE- CHARGE OF A MONOATOMIC ION AND IS EQUAL TO THE NUMBER OF ELECTRONS TRANSFERRED FROM OR TO AN ATOM. Ions

7. Oxidation Numbers Oxidation numbers are written as superscripts to indicate charge but are not included in the formula. Ex: Li 1+ or O 2-

8. NAMING MONOATOMIC IONS Cations Name the element and add the word Ion. Example: Na ⁺= Sodium Ion Anions Drop the ending off the element name and adding “ide” Example: O²¯= Oxide

9. Naming Transition Metal Ions Transition Metals have more than one Ionic charge. Example: Iron can form two Cations Fe ²⁺ or Fe³⁺ Fe ²⁺= Iron (II) ion Fe³⁺= Iron (III) ion

10. Practice Lithium= Fluorine= Oxygen= Carbon= Barium= Cesium= Au 2+= Gold (III) Ion =

11. Binary Ionic Compounds Binary Ionic Compound- made of two ions/elements; Naming Binary Ionic Compounds The cation keeps the normal name from periodic table and the Anion receives an –ide ending. Example: NaCl= Sodium Chloride KI= Potassium Iodide

12. Writing a formula for Ionic Compounds Binary Ionic Compounds = cation + anion 1. Determine the oxidation number of each cation and anion. 2. The net charge of the compound must equal “0” Example: Sodium Chloride Na ⁺ and Cl¯ +1 and -1 =0 Formula = NaCl

13. Naming Simple Binary Ionic Compounds Example 2 Calcium Chloride Ca ²⁺ and Cl¯ 2+ and -1 = +1 Add one more Cl to = 0 Formula= CaCl₂

14. Naming Simple Binary Ionic Compounds Practice!!! Potassium Iodide= Calcium Phosphide= Magnesium Chloride= Magnesium Oxide= Li ₂O= KBr= Metals Always Placed Before Non-Metals!!!!!!

15. Polyatomic Ions Polyatomic Ions-Tightly bound group of atoms that behave as a unit and carry a charge. Ex. SO ₃²¯, ClO₂⁺

16. Names for Polyatomic Ions Most of the names end with –ite or –ate but not all. The –ite ending indicates one less oxygen than the -ate ending. But –ite and –ate does not determine the actual number of oxygen atoms.

18. Cations and Polyatomic Ions When adding Cations (metals) to polyatomic ions add the basic element name to the Polyatomic Ion. Example: CaSO ₄ = Calcium Sulfate Writing Formulas The charge for the whole compound must equal “0”. Example: Magnesium Chlorate

19. Writing Formulas with Cations and Polyatomic Ions The charge for the whole compound must equal “0”. Example: Magnesium Chlorate Magnesium= +2 and Chlorate = -1 We must have 2 chlorate for the compound to equal Zero. Formula= Mg(ClO ₃)₂ When adding an extra Polyatomic Ion Use Parenthesis and add desired amount behind.

20. Ionic Compounds with Transition Metals Naming 1. Recognize the metal is a transition metal. FeCl ₂ Fe = Iron a transition metal unknown charge 2. Use the Non-metal to determine the charge of the transition metal. Chlorine charge = -1 x 2 = -2 3. The charge of the whole compound must equal zero. Iron must have a +2 Charge Name of FeCl₂ = Iron (II) Chloride

21. Naming Acid Rules Anion Ends In –ide Hydro- stem-ic Acid Example with Chlorine = Hydrochloric Acid Anion Ends In –ite Stem of Anion –ous Acid Example with Sulfite= Sulfurous Acid Anion Ends In –ate Stem of Anion-ic Acid Example with Nitrate= Nitric Acid

22. Acids Acids- compounds that produce hydrogen ions when dissolved in water Acids to Know Hydrochloric Acid=HCl Sulfuric Acid= H ₂SO₄ Nitric Acid= HNO₃ Acetic Acid= HC₂H₃O₂ Phosphoric Acid= H₃PO₄ Carbonic Acid= H₂CO₃

23. END OF Unit 4!!!

24. Naming Covalent/Molecular Compounds Covalent Compounds are made of two non-metals. Pre-fixes are used to indicate the number of atoms in each compound.

25. Naming Covalent/Molecular Compounds Rules 1. The first element name is given followed by the second with an “ide” ending. 2. The first element gets a prefix if more than one. 3. The second element always gets a prefix. Examples NO=Nitrogen Monoxide N ₂O= Dinitrogen Monoxide N₂O₄=Dinitrogen TetraOxide

26. Naming Covalent/Molecular Compounds Formula Copper(II) Sulfate = CuS Roman Numeral = Charge of your Metal Use the Roman Numeral to figure out the charge on your non-metal. Copper (II) equals + 2 Charge The total charge on the whole compound must still equal 0.

27. Molecular Compounds Molecular Compounds or when two atoms of different elements create bonds. Examples: CO ₂, H₂O, PCL₃ Characteristics: 1. Low melting points and Low boiling points 2. Form liquids or gases. 3. Made from two or more non-metals

28. MOLECULE-MADE UP OF TWO OR MORE ATOMS OF THE SAME ELEMENT. EXAMPLES: 0 ₂ (OXYGEN YOU BREATHE) O₃ (OZONE) Molecules and Bonds

29. Ionic Bonds Vs. Covalent Bonds Ionic BondsCovalent Bonds Bonded together by the attraction of opposite charges (+ and -) Bonded together by the sharing of electrons Electrons are donated/transferred and taken giving one atom a (+) charge and the other a (-) charge. Electrons are shared between both atoms charges remain the same. Polar MoleculeNon-Polar Molecule Made of a metal and non-metalMade up of two non-metals Weaker BondsStronger Bonds