Predicting Products of Chemical RXNS Make sure you have a whiteboard and dry erase marker

Recap: Chemical Equations Coefficient: large number before the element/compound. Indicates # of the element/compound used in rxn Subscript: small number next to element symbol Indicates the number of each element present in the compound/molecule. States of Matter: Solid (s), Liquid (l), Gas (g), Aqueous (aq) New:  = reversible reaction = catalyst (in this case, Pt is the catalyst) or = heat is added

Recap: Balancing List all the element (Single Cap. Letter) on the reactant side of the equation. Count # of atoms on reactant(left of arrow) & product(right of arrow) side for EACH element. Use coefficient to balance the equation Multiply: coefficient x subscript to change # of atoms.

Lets review a few of the Problems: __FeCl3 + __Be3(PO4)2  __BeCl2 + __FePO4 __AgNO3 + __LiOH  __AgOH + __LiNO3 __Mg + __Mn2O3  __MgO + __Mn

Packet Practice

Recap: Rxn Types Synthesis: A + B  AB Decomposition: AB  A + B 2+ reactants  1 product Decomposition: AB  A + B 1 reactant  2+ products Single Replacement: A + BX  B + AX element + compound  element + Compound Double Replacement: AX + BY  AY + BX compound + compound  compound + compound Combustion: AB + O2  H2O + CO2 or Oxide O2 on REACTANT SIDE to produce CO2 & H2O or an oxide

1. Synthesis (redox) Two or more elements or compounds that combine to form a single substance either ionic or covalent. General equation: A + B  AB EX. Iron and chlorine combine to produce Iron chloride. 2Fe (s) + Cl2 (g)  2FeCl3 (s) Neutral Neutral Iron +3 oxidized, Chlorine -1 reduced Iron is the reducing agent, Chlorine is the oxidizing agent

Synthesis Special Cases Non-metal oxide + water  acid Use oxidation numbers to find the oxidation number of the non-metal. Use the oxidation number of the non-metal to determine the oxyanion of the acid. Must balance charges from the ions when you write the formula for the acid. Example: SO2 + H2O  H2SO3 Non-metal oxide + metal oxide  salt Use the oxidation number of the non-metal to determine the oxyanion of the salt. The cation is the metal ion. Must balance charges from the ions when you write the formula for the salt. Examples: CO2 + metal oxide  metal carbonate SO2 + metal oxide  metal sulfite SO3 + metal oxide  metal sulfate metal oxide + water  base The metal is the cation of the base. The anion is OH-1 (definition of a base). Must balance charges from the ions when you write the formula for the base. Example: Al2O3 + H2O  Al(OH)3

Practice 1. CaO(s) + H2O(l)  Ca(OH)2(aq) 2. H2(g) + Cl2(g)  2HCl(g) 3. CO2(g) + H2O(l)  4. K2O(s) + CO2(g)  Ca(OH)2(aq) 2HCl(g) H2CO3(aq) K2CO3(s)

2. Decomposition A single compound is broken down into two or more substances. Opposite of synthesis. General equation: AB A + B EX. Hydrogen peroxide decomposes into water and oxygen. 2H2O2 (aq)  O2 (g) + 2H2O (l)

Decomposition Special cases acid  Non-metal oxide + water salt  Non-metal oxide + metal oxide Uses the oxidation steps from the combination reactions in reverse. Find the oxidation number of the non-metal in the oxyanion. Use the oxidation number of the non-metal to determine the formula of the non-metal oxide. Must balance charges from the ions when you write the formula for the metal oxide. The metal cation is the metal from the salt. Metal halate  metal halide + oxygen gas Metal carbonate  metal oxide + carbon dioxide Metal peroxide  metal oxide + oxygen gas  

Practice 1. BaCO3  BaO + CO2 2. Ni(ClO3)2  NiCl2 + 3O2 3. Ag2O  4. H2SO4  BaO + CO2 NiCl2 + 3O2 2 4Ag + O2 H2O + SO3

3. Single replacement One element replaces a second element in a compound. -the more reactive metal will replace the least reactive metal. -the more reactive nonmetal will replace the least reactive nonmetal. General equation: A + BX  AX + B (metal) Q2 + RS  RQ + S2 (nonmetal) EX. 2AgNO3 (aq) + Cu(s)  2Ag(s) + Cu(NO3)2 (aq) Silver nitrate reacts with copper to produce silver and copper

Activity Series A chart of metals listed in order of declining relative reactivity. The top metals are more reactive than the metals on the bottom. The top metals will always replace metals below it in a chemical reaction

Practice 1. Au(s) + AgNO3(aq) → No Reaction 2. Fe(s) + Cu(NO3)2(aq) → 3. Ca(s) + H2O(l) → 4. Cl2(g) + KI(s)  No Reaction Cu(s) + Fe(NO3)2(aq) CaO(s) + H2(g) I2(g) + KCl(s)

4. Double replacement Involves an exchange of positive ions between two reacting compounds. Reactants are two ionic compounds One of the following applies: One product precipitates from solution or One product is a gas and bubbles out of solution or One product is a molecular compound such as water. Use Solubility chart to determine precipitate! AX + BY  A Y + B X Ex. AgNO3 (aq) + NaCl (aq)  AgCl (s) + NaNO3 (aq)

Double Replacement Special cases: Acid-Base neutralization reactions: Acid + Base  Salt + water Acid: Cation = H+ Base: Anion = OH-1 Salt: Anything that can be made from an acid and a base that is not water. Gas forming reactions: Acid + metal carbonate  salt + water + carbon dioxide *Double replacement reaction that produces H2CO3 - H2CO3 breaks down as soon as it forms. Follows the rules for metal carbonate decomposition. Acid + metal sulfate  salt + water + sulfur dioxide -Same type of problem as the acid + metal carbonate.

Practice 1. Ca(C2H3O2)2 + Na2CO3  2. NaCl + AgNO3  3. HNO3 + Na2CO3  4. HBr + KOH  CaCO3(s) +2NaC2H3O2(aq) AgCl(s) + NaNO3(aq) 2NaNO3 + H2CO3 H2O(l) + CO2(g) H2O(l) + KBr(aq)

5. Combustion An element or compound reacts with oxygen, often producing energy as heat or light. CxHyOz + O2  CO2 + H2O Some synthesis (combination) reaction fall under this category. Metal reacting to form metal oxides General form…you get the oxide of every element you burn (except oxygen). Complete combustion is excess oxygen…gives you CO2 as the oxide of carbon. Incomplete combustion is limited oxygen…gives you CO as the oxide of carbon.

Practice 1. hexane(C6H14) burns in air. 2. ethene (C2H4) burns in air. 3. benzene (C6H6) burns in air. 4. Magnesium ribbon burns in air. 2C6H14 + 19O2  12CO2 + 14H2O C2H4 + 3O2  2CO2 + 2H2O 2C6H6 + 15O2  12CO2 + 6H2O 2Mg + O2  2MgO

Net Ionic Equations Net ionic equations show only the species actually involved in the reaction. As you first learn to write net ionic equations, you will write three different equations for each reaction.

Do not ionize solids, pure liquids, or gases. Steps for Net Ionic Equations Write the complete molecular equation. (This is the type of equation that you are used to writing.) Write the complete ionic equation. To do this, you must ionize everything that is soluble and ionized in solution. Everything else is left together. Do not ionize solids, pure liquids, or gases. Write the net ionic equation. To do this, cancel out all ions that are not participating in the reaction (spectator ions) and rewrite the equation.

Ex. Sodium chloride + silver nitrate NaCl(aq) + AgNO3(aq)  Na+ + Cl- + Ag+ + NO3-  Na+ + NO3- + AgCl(s) Cl- + Ag+  AgCl NaNO3(aq) + AgCl(s) Ex. HCl + Ba(OH)2  1. 2HCl(aq) + Ba(OH)2(aq)  2. 2H+ + 2Cl- + Ba2+ + 2OH-  2H2O + Ba2+ + 2Cl- 3. 2H+ + 2OH-  2H2O H+ + OH-  H2O (reduce) 2H2O(l) + BaCl2(aq)