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: CHEMICAL REACTIONS Click to add text • Chemical reactions: Reactions that produce new substances • PRODUCT: substance formed during a chemical reaction (right side of arrow) • REACTANT: starting substance(s) in a chemical reaction (left side of arrow) • Law of Conservation of Mass must be satisfied! In this unit you should know… • 1. How to balance chemical equations • 2. Identify the different types of reactions • 3. Be able to predict the products for both single and double replacement reactions • 4. Determine if a reaction will take place using either the activity series of metals or solubility rules • 5. Understand the role of a catalyst in a chemical reaction Evidence of Chemical Reactions Temperature change: endothermic (colder), exothermic (hotter) Color Change Odor Gas Produced (bubbles) Precipitate: formed from 2 liquids Balancing Equations Steps 1) Balance atoms that appear only once on each side. 2) Balance polyatomic ions that appear on both sides as a single unit. 3) Balance hydrogens. 4) Balance oxygens. 5) Never change the subscripts of a compound to balance an equation. Types of Reactions • Synthesis Reaction: • 1. Two or more substances combine to form a single compound. • 2. Usually energy is released (exothermic) • 3. Basic reaction: A + B --> AB Synthesis Reaction Examples: • Element + Oxygen ----> Oxide Compound • Magnesium + Oxygen ---> Magnesium Oxide • Mg + O2 ------> 2 MgO • Metal Oxide + Water ---> Hydroxide Compound • CaO + H2O ---> Ca(OH)2 (base) Decomposition Reactions: • 1. Single compound is broken down into two or more simpler products. • 2. Usually requires energy. • 3. Basic reaction: AB ---> A +B Decomposition Reaction Examples: • Metal Carbonate ----> Metal oxide + carbon dioxide • Ca CO3 ----> CaO + CO2 • Metal Hydroxide ----> metal oxide + water • Ca(OH)2 ---> CaO + H2O • Metal Chlorate ---> metal chloride + oxygen • 2KClO3 ---> 2 KCl + O2 • Oxyacid ---> nonmetal oxide + water • H2SO4 ---> SO3 + H2O SINGLE REPLACEMENT REACTION: • 1. One element replaces a similar element in a compound. • 2. A reactive metal will replace any metal that is less reactive (see pg 288 Activity Series of Metals) • 3. Nonmetal will replace other nonmetals. Activity Series: Single Replacement Reactions Only • One metal will only replace another if it is HIGHER on the activity series • This is because it is a more reactive metal Single Replacement cont. • 4. Basic Reaction: • A + BC ---> AC + B • Y + BX ---> BY + X Single Replacement Examples: • Replacement of a metal in a compound by a more reactive metal • Use activity series to determine if one metal is strong enough to replace the other one. If not, then no reaction will occur • 2Al + 3Fe(NO3)2 ---> 3Fe + 2Al(NO3)3 • Replacement of Halogens. • Cl2 + 2 KBr ---> 2KCl + Br2 • Metal replacing hydrogen in an acid. • Zn + 2HCl ---> ZnCl2 + H2 DOUBLE-REPLACEMENT REACTIONS: • 1.Exchange of positive ions between two compounds. • 2.One compound formed is usually a precipitate, gas, or a molecular compound (often water) • 3. Basic Equation: • AB + CD ---> CB + AD • 4. Use the solubility rules to determine whether or not a reaction will take place Double-Replacement Examples: • Metal oxide + acid ---> water + salt (metal/nonmetal) • MgO + 2 Hcl ---> H2O + MgCl2 • Metal carbonate + acid ---> salt + carbon dioxide +water • CaCO3 + 2 HCl ---> Ca Cl2 + CO2 + H2O • Acids + metal Hydroxide ---> salt + water • HCl + NaOH ----> NaCl + H2O Solubility Rules Overview • List of rules used to determine whether or not a reaction will take place • Remember! In order for a reaction to take place you must produce a gas or a precipitate from 2 liquids. • Solubility rules tell us whether or not a precipitate (solid) is produced • You will often see these letters indicating what state of matter a substance is • • • • Solid (s) Liquid (l) Gas (g) Aqueous (aq) = soluble in water Double Rep. Reactions: Solubility Rules (see handoutdo not have to copy down) • 1. Soluble: All salts containing the ammonium or Group IA ions (Li+, Na+, K+, Rb+, Cs+) • 2. Soluble: All salts containing nitrate (NO3-), acetate (C2H3O2-), and perchlorate (ClO4-) • 3. Soluble: All salts containing Group VIIA ions (Cl-, Br-, I-), except those in Rule 5. • 4. Soluble: All salts containing sulfate (SO4-2). Exceptions are barium sulfate, calcium sulfate, lead II sulfate, and strontium sulfate. Solubility Rules cont. • 5. Insoluble: All salts containing silver ion (Ag+), lead II ions (Pb+2), and mercury I ions (Hg2+2) • 6. Insoluble: All salts containing carbonate, chromates, hydroxides, oxides, phosphates, and sulfides • Exceptions: • Group IIA chromates, except barium chromate are solulbe • Group IIA hydroxides, except magnesium hydroxide, are soluble COMBUSTION REACTIONS: • 1. Oxygen reacting with another substance. • 2. Usually involves hydrocarbons (contain hydrogen & carbon) • 3. Heat is always released. • 4. Basic Equation: CXHY + O2 ---> H2O + CO2 • [x & y represent a ratio of carbon & hydrogen] Combustion Examples: • 4. Complete combustion: • C3H8 + 5O2 ---> 3CO2 + 4H2O • 5. Incomplete combustion: creates carbon monoxide (CO), carbon, & water. [products cannot be predicted] Catalysts • A substance that increases the rate of a chemical reaction by lowering activation energies but is not itself consumed in the reaction. • Example: Enzymes: allow many chemical rxns to occur at a rate that sustains life at normal living temperatures