1. UNIT 4: Formulas and Equations (Review Book Topic 2) How can we distinguish between quantitative and qualitative information? What are the different types of formulas? How can we know that elements form a compound and in what proportions? How can we write formulas? How can we name compounds? What are the different parts of a chemical equation? What are the differences between endothermic and exothermic reactions? How can chemical equations demonstrate the Law of Conservation? What are the different type of chemical reactions? How can we predict products of a chemical reactions? How can we determine an unknown reactant, product, or mass in a chemical equations?

2. AIM: How can we distinguish between quantitative and qualitative information?  Chemical Symbols: Each element has a unique one-, two-, or three- letter symbol  The first letter is always capitalized (Table S, PT)  Almost all symbols are written without a subscripts as monatomic  Diatomic Molecules: elements that exist in nature as two identical atoms covalently bonded into a diatomic molecule  Br 2, I 2, N 2, Cl 2, H 2, O 2, F 2

3. AIM: How can we distinguish between quantitative and qualitative information?  Chemical Formulas: compounds are composed of elements chemically combined in fixed ratios  Formulas use chemical symbols and number to show both qualitative and quantitative information about a substance  Qualitative: information that relates to things that cannot be counted or measured-  What elements are in the compound  Quantitative: information that deals with things that can either be counted or measured  The number of atoms of each element in the compound (subscript or coefficient can give us this info)

4. AIM: How can we distinguish between quantitative and qualitative information? Example: Determine the quantitative and qualitative information in the following examples: 1. CaCO 3 2. Zn 3 (PO 4 ) 2

5. AIM: What are the different types of formulas  Empirical: simplest ratio  Ionic formulas are always empirical formulas  Molecular: actual ratio  covalently bonded substances form molecules, in some cases the empirical represents both empirical and molecular –H 2 O-  In other cases molecular formula is a multiple of the empirical formula C 6 H 12 O 6 – six times the empirical CH 2 O

6. AIM: How can we know what elements form a compound and in what proportions?  Atoms and compound are electrically neutral – equal numbers of negative (electrons) and positive (protons)  Ions – can be positive or negative  Ionic charge: indicated by a superscript positive – lost electrons, negative – gained electrons Cl -, Al +3  Polyatomic Ion: group of atoms covalently bonded together, possessing a charge – Table E

7. AIM: How can we know what elements form a compound and in what proportions?  Forming a compound: many ways – one way:  By attraction of oppositely charged ions  Monatomic or polyatomic ions attract each other in a ratio that produce a neutral compound  Coefficients: written in front of a formula, applies to entire formula, multiple coefficient and subscript to determine the number of each type of element

8. AIM: How can we know what elements form a compound and in what proportions?  Hydrates  Compounds that contain definite amounts of water molecules  Ex: BaCl 2 ·2H 2 O – Barium chloride traps 2 water molecules

9. AIM: AIM: How can we write formulas?  Compounds must be electrically neutral  For many elements, the oxidation state is equal to the charge found in the top right corner of each element box  1:1  Na + & Cl - yields NaCl  Mg 2+ & S 2- yields MgS  Not 1:1 – Mg 2+ & Cl -  Write the charge of one ion as the subscript of the other without the sign (# only)  Thus MgCl 2 ; that is 1 Mg with a 2+ & 2 Cl with 1- each yields (2+) + 2(1-) = 0

10. AIM: AIM: How can we write formulas? Examples: Write the formulas for the following: a. Al +3 and Br – b. Ba +2 and CO 3 -2 c. Cu 2+ and CO 3 -2 d. Pb +2 and Cl – e. Pb +4 and CrO 4 -2

11. AIM: How can we name compounds?  Binary Ionic (Metal–positive & Nonmetal–negative)  1 st element (metal) – retains its name  2 nd element (nonmetal) – change ending to “ide”  Ex: KCl is Potassium chloride  Other Ionic (Contains Polyatomic Ions)  Same as Binary except all polyatomic ions retain their name  Examples:  KNO 3 is __________________  NH 4 Cl is _____________________  NH 4 NO 3 is___________________ Potassium nitrate Ammonium chloride ammonium nitrate

12. AIM: How can we name compounds? Examples: Write the names for the following formulas: a. Mg(SO 4 ) b. Na(OH) c. Ca(OH) 2 d. Li 3 (PO 4 ) e. (NH 4 )Cl

13. AIM: How can we name compounds?  Covalent (2 Nonmetals)  Need prefixes to tell reader how many of each  Exception – if only one 1 st element, don’t use mono  Examples  NO is _____________________  N 2 O 4 is _____________________  Stock System (multiple oxidation states)  mostly the touchable metals; Iron, Tin, Copper...  a Roman numeral proceeding the metal tells the reader the oxidation number  Examples  FeCl 2 is _________________  FeCl 3 is _________________ Nitrogen monoxide Dinitrogen tetroxide Iron (II) chloride Iron (III) chloride

14. AIM: AIM: What are the different parts of a chemical equation?  Reactant(s) yield Product(s)  Reactant are on the left of the arrow (yield sign) and products are to the right of the arrow  State of matter is indicated by the letter inside the parenthesis  Example: C (s) + O 2 (g) → CO 2 (g)  Identify the reactants and products in the equation above…….

15. AIM: What are the differences between endothermic and exothermic reactions?  Endothermic  Heat is required for a reaction to occur, thus, energy is found on the reactant side  Example: H 2 O (s) + energy → H 2 O (l)  Exothermic  Heat is produced in a reaction, thus energy is found on the product side  Example: H 2 O (g) → H 2 O (l) + energy

16. AIM: AIM: How can chemical equations demonstrate the Law of Conservation? - BALANCING  The Law of Conservation of Mass & Charge must be upheld  Example:  __ H 2 (g) + __ O 2 (g) → __ H 2 O (g)  Remember – It’s coefficient x subscript to find the # of atoms 2 2 1

17. AIM: AIM: How can chemical equations demonstrate the Law of Conservation? - BALANCING  Nothing can be created or destroyed – law of conservation of mass  Count up the atoms on both side and fill in any missing elements or compounds  Same thing for missing mass  Practice: If 103.0g of potassium chlorate are decomposed to form 62.7g of potassium chloride and oxygen gas according to the equation 2KClO 3  2KCl + 3O 2 how many grams of oxygen are formed?

18. AIM: What are the different types of chemical reactions?  Synthesis  2 or more reactants form 1 product  A + B → AB  Decomposition  1 reactant breaks down into 2 or more products  AB → A + B  Single Replacement  1 element replaces another  A (element) + BX (compound) → B (element) + AX (compound)  Double Replacement  2 elements/polyatomic ions replace two others  AB (compound) + CD (compound) → AD (compound) + CB (compound)

19. AIM: AIM: How can we predict products of a chemical reactions?  Single Replacement Reactions:  If the individual metal is above the metal that is in the compound a reaction will occur  Double Replacement Reactions:  If a solid is formed (Table F)  If a gas is formed  If a molecular substance such as water is formed

20. Using TABLE F 1. Cross out the first element or compound 2. Look for second element or compound on Table F check to see if it is paired with an exception 3. Determine solubility Ex) NH 4 Cl  MgOH Soluble Insoluble

21. TABLE F AgNO 3 (aq) + NaCl(aq)  AgCl(_____) + NaNO 3 (____) s aq

22. AIM: How can we determine an unknown reactant, product, or mass in a chemical equations?  Nothing can be created or destroyed – law of conservation of mass  Count up the atoms on both side and fill in any missing elements or compounds  Same thing for missing mass