Chemical Equations and Reactions

1. Chemical Equationsand Reactions

2. I. Chemical Equations

3. II. Chemical Equations A. The equation must represent facts. B. The equation must contain the correct formulas for the reactants (on the left of the arrow) and the products (on the right of the arrow). C. The law of conservation of mass and energy must be satisfied. Therefore the same number of atoms of each element must appear on each side of a correct chemical equation.

4. Symbol Explanation of symbol + separates 2 or more reactants or products  “yield”, separates reactants from products. ↔ indicates a reversible reaction (s) solid state. Placed after the formula of a substance ↓ Alternative to (s) but used ONLY for a solid PRODUCT, not reactants (l) indicates a liquid reactant or product (aq) indicates an aqueous solution (where some solute has been dissolved in water) (g) indicates a gaseous reactant or product ↑ alternative to (g), but used ONLY for a gaseous PRODUCT Δindicates that heat is supplied to the reaction A formula written above or below the  sign indicates that it is used as a catalyst (something that speeds up the reaction)

5. C. Diatomic and Polyatomic Molecules: Element Formula State Hydrogen H2 gas Nitrogen N2 gas Oxygen O2 gas Fluorine F2 gas Chlorine Cl2 gas Bromine Br2 liquid Iodine I2 solid Sulfur S8 solid Phosphorus P4 solid

6. III. Writing and Balancing Equations

7. Example Write a balanced equation for the following reaction: Na + Cl2 NaCl first write an atom inventory for the total number of atoms of each element on each side of the equation. Na + Cl2 NaCl Reactants Products # Na # Na # Cl # Cl

8. Atom Inventory or Counting Atoms • you must be able to count atoms in order to balance an equation. There are two ways to designate numbers in a formula: • subscripts – small numbers within a formula of a compound. Tells the number of atoms in that compound • MgCl2 – 1 atom of Mg and 2 atoms of chlorine • Sn3N2 – 3 atoms of tin and 2 atoms of nitrogen • Coefficient – the large number in front of the formula of a compound. Tells the number of molecules (in a molecular compound) or formula units (in an ionic compound) or atoms of an element.

9. Remember that atoms cannot be created or destroyed; we must balance an equation using coefficients. Never change a subscript to balance an equation!!

10. Algebraic Method Write the skeleton equation Write a,b,c… on each substance List the atoms present Determine equalities for each atom Assign a value 6 for A Use A to determine other values Reduce, if possible Plug #s in as coefficients

11. Practice __H2O  __H2 + __O2 __Pb(NO3)2 + __Na  __NaNO3 + __Pb __C4H10 + __O2 __CO2 + __H2O

12. Examples Zn + H2O  Zn(OH)2 + H2 C2H6 + O2  CO2 + H2O Na2SO4 + Ba(NO3)2  BaSO4 + NaNO3

13. IV. Types of Chemical Reactions

14. A. Combination or Synthesis • where 2 or more simple substances (elements or compounds) combine to form ONE complex substance • 8Fe + S8 8FeS • 2Sr + O2 2SrO

15. A. Combination or Synthesis Li + P4 N2 + Al  Cl2 + Ca  Na + N2 

16. Special Combination or Synthesis Reactions (Pre-AP Only) the metals that has a variable charge: If one of these metals reacts with fluorine, oxygen, or nitrogen (F, 0, N), these nonmetals will pull the metal to its HIGHEST chargeoroxidation number. Otherwise, when these metals react in a combination reaction, use their LOWEST charge or oxidation number when forming a new compound

17. Practice: Fe + O2 Pb + N2 Sn + S8 Cu + P4  Fe + Br2 Cu + F2

18. B. Decomposition a complex substance (compound) decomposes into 2 or more simple substances. Heat or electricity is usually required. Ex: 2NaCl  2 Na + Cl2 8MgS  8Mg + 2S8

19. Special decomposition reactions to know (Pre-AP only): • 2KClO3 2KCl + 3O2- • all metal chlorates decompose into metal chloride + O2 • CaCO3  CaO + CO2 • metal carbonates decompose into a metal oxide + CO2 • 2KOH  K2O + H2O • metal hydroxides decompose into a metal oxide + H2O

20. C. Combustion Reactions • where oxygen reacts with another substance, usually a hydrocarbon, resulting in the release of energy, usually heat or light. • CH4 + 202 CO2 + 2H20 • Hydrocarbons always produce carbon dioxide and water

21. Common Hydrocarbons

22. Examples C3H8 + O2 C2H2 + O2 Ca + O2

23. D.Single-Replacement • occurs when one element displaces another element in a compound. You must check the “Activity Series of Metals” to see if the “lone” element is active or “strong” enough to displace the element in the compound

24. Activity Series of Metals Li K Ba Ca Na Mg Al Zn Fe Ni Sn Pb H Cu Hg Ag Au

25. Practice: Li + KCl  Sn + ZnCl2 Sn + HCl  Ni + HOH 

26. the halogens F2 Cl2 Br2 I2 Decreasing strength

27. E. Double-Replacement reactions • occur when the cations (positive ions) “switch” places. You do NOT need the “activity series of metals” list in these reactions. When you switch places, be sure to correctly write the formula of the new compound!!!!! • Ex: • 2 NaCl + Mg0  MgCl2 + Na20

28. Practice CuS04 + Al(OH)3 Ca3(P04)2 + ZnCr04

29. Rules for Predicting Double Replacement Reactions: • 1. Predict the products of the double-replacement reaction and indicate the solubility of both of the products • Use the “Solubility Rules” handout (at end of notes) to determine the solubility. • If the compound is soluble that means that it will remain as ions in the solution, if it is insoluble then the compound precipitated out of the reaction (it became the precipitate or solid). • 2. If at least one INSOLUBLE product is formed (which means a precipitate will form) the reaction will occur! • 3. If only SOUBLE products are formed then the reaction will NOT occur (because no precipitate is formed)! • 4. If water is produced the reaction will occur! • 5. If the reaction occurs and one of the compounds formed is soluble then that compound is written as ions and not as a compound.

30. SOLUBILITY RULES 1. All salts whose cation is in Group 1 or is NH4+ are soluble – no matter what the anion is. 2. All nitrates and nitrites are soluble. 3. All acetates are soluble. 4. All chlorates and perchlorates are soluble – no matter what the cation is. 5. All chlorides are solubleexcept silver, lead (II), mercury 6. All bromides are solubleexcept silver, lead (II), mercury 7. All iodides are solubleexcept silver, lead (II), mercury 8. All flourides are INSOLUBLE. 9. All sulfates are solubleexcept silver, lead (II), mercury, calcium, strontium and barium 10. All sulfides are INSOLUBLEexcept Group 1 and 2. 11. All phosphates, phosphites, carbonates, chromates, and dichromates are INSOLUBLE unless the cation is in Group 1 or is NH4+. 12. All hydroxides are INSOLUBLEexcept Group 1, calcium,. Strontium, barium and ammonium

31. Net Ionic Equations (Pre-AP Only)

32. F. Net Ionic Equations (Pre-AP Only) – shows only the compounds and ions that undergo a chemical change in a double replacement reaction • Example: Na2S + Cd(NO3)2 Na+ + NO3 + CdS(s) • Step 1: Convert the chemical equation to an overall ionic equation. All reactants are shown as ions. For the products, all soluble ionic compounds are shown as dissociated ions and the precipitates are shown as solids. • Na+ + S 2 + Cd 2+ + NO3 Na+ + NO3 + CdS(s)

33. Step 2: All spectator ions (ions that do not take part in a chemical reaction and are found as ions both before and after the reaction) are removed from the equation. • S2 + Cd2+  CdS(s)

34. Examples BaCO3 + CuSO4 BaSO4(s) + CuCO3 (s) K3PO4 + NaOH  no reaction (no ppt) Na2S + Cd(NO3)2 Na+ + NO3 + CdS(s)

35. Types of Reactions Summary Combination (synthesis) A + B  AB Decomposition AB  A + B Combustion CxHy + O2 CO2 + H2O Single Replacement A + BC  AC + B Double Replacement AB + CD  AD + CB

36. Practice Predicting Products l. AlCl3 2. C2H4 + 02 3. Zn + AgNO3 4. H20 

37. Practice Predicting Products 5. Al + P4 6. NaI + MgS  7. Cl2 + NaBr  8. C6H1206 + O2

38. Practice Predicting Products 1. AlCl3 + Na2CO3 2. Ni + MgSO4 3. Cl2 + K  4. C5H12 + 02

39. Practice Predicting Products (Pre-AP) • sodium metal is placed into water • methane gas is burned in the presence of oxygen • potassium bromide solution is mixed with chlorine gas • a solution of aluminum dichromate is mixed with a solution of lithium oxalate

40. V. Oxidation Reduction Reactions (Pre-AP only)

41. What is REDOX? • Oxidation-Reduction (Redox) – involves a transfer of electrons • One specie is losing electrons OIL – Oxidation is losing Mg0 Mg+2 + 2 e-

42. One specie is gaining electrons RIG – Reduction is gaining Mg+2 + 2 e- Mg0

43. The Species that is oxidized is the reducing agent • The Species that is reduced is the oxidizing agent.

44. Mg0 + O20 Mg+2O-2 Ox Red ag Ox ag Red

45. REDOX reactions MUST: • 1 Have a species that is oxidized and one reduced – YOU cannot have one without the other • The number of electrons gained and lost MUST be the SAME • The number of atoms of each element must be the same on both sides of the equation

46. Balancing REDOX Reactions • 1) Assign oxidation numbers to each atom in the equation • 2) Determine the substances oxidized, reduced, oxidizing agent, reducing agent,

47. 3) Write balanced half-reactions for the oxidation and reduction reactions. • Mg0 + O20 Mg+2O-2 Mg0 Mg+2 + 2 e- O20 + 4 e-  2O-2 

48. 4) Multiply each equation so that the number of electrons lost equals the number of electrons gained. • Mg0 + O20 Mg+2O-2 • 2[Mg0 Mg+2 + 2 e-]= 2Mg0 2Mg+2 + 4 e- O20 + 4 e-  2O-2

49. 5) Add the two half-reactions. Place the coefficients into the original equation. • Mg0 + O20 Mg+2O-2 2Mg0 2Mg+2 + 4 e- O20 + 4 e-  2O-2 2Mg0 + O20  2Mg+2O-2

50. 6) Adjust other ions if necessary. Check all atoms for conservation. Check hydrogen’s and oxygen’s last. 2Mg+ O2 2MgO