Chemical reactions describe processes involving chemical change
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1 1.1 Chemical Reactions 1.2 Chemical Equations Chemical reactions describe processes involving chemical change The chemical change involves rearranging matter Converting one or more pure substances into new pure substances Reactants Products Chemical equations are used to describe chemical reactions The chemical symbols for the reactants are shown on the left The chemical symbols for the products are shown on the right An arrow ( ) is used to indicate that reactants are converting to products A plus sign () is used to separate individual id reactants and products 1.3 Chemical Equations 1.4 Chemical Equations hydrogen and oxygen react to form water Chemical equation describing the action: Subscripts? Coefficients? 2 H 2 (g) O 2 (g) 2 H 2 O(l) Letters in parentheses? 2 H 2 (g) O 2 (g) 2 H 2 O(l) The equation accomplishes several things: Adheres to the fundamental law of conservation of matter Atoms are neither created nor destroyed in chemical reactions, they are only rearranged 1
2 1.5 Balancing Chemical Equations All of the matter present in the reactants is also present in the products it s just rearranged Chemical equations = mathematical equations? Some of the same rules apply A balanced equation has the same number and type of atoms in the reactants (on the left) as in the products (on the right) Stoichiometry 1.6 Balancing Chemical Equations For example, in the equation describing the formation of liquid water from hydrogen gas and oxygen gas 2H 2 (g) O 2 (g) 2H 2 O(l) There are four hydrogen atoms on both the left and right sides of the equation There are two oxygen atoms on both the left and right sides of the equation Therefore the equation is balanced 1.7 Balancing Chemical Equations Remember. 2H 2 (g) O(g) 2 2H 2 O(l) The 2 to the left of H 2 (g) and H 2 O(g) refers to the number of molecules present in the balanced equation It is a multiplier for every atom in the molecule l that t follows The subscript 2 in O 2 (g) and NO 2 (g) refers to the number of atoms of this type that are present in each molecule (or ionic compound) 1.9 Balancing Chemical Equations H 2 S(aq) I 2 (aq) HI(aq) S(s) KClO 3 (s) KCl(s) O 2 (g) Ca(OH) 2 (s) H 3 PO 4 (aq) Ca 3 (PO 4 ) 2 (s) H 2 O(l) C 6 H 14 (l) O 2 (g) CO 2 (g) H 2 O(l) 2
3 2.1 Types of Reactions 2.2 Double Replacement Reaction Types Double Replacement Precipitation Acid-Base Combination Decomposition Single Replacement Oxidation-Reduction Gas Forming Double replacement reactions are also called metathesis reactions or partner swapping reactions They have the form AX BY BX AY HF NaOH NaF H 2 O = H = O = Na = F 2.3 Double Replacement 2.4 Mixtures and Solutions Double replacement reactions often take place in water and are of two basic types: Acid base neutralization reactions: HCl(aq) NaOH(aq) NaCl(aq) H 2 O(l) (acid) (base) (salt) (water) H 2 SO 4 (aq) 2 NaOH(aq) Na 2 SO 4 (aq) H 2 O(l) Precipitation reactions (solid forms): Ba(NO 3 ) 2 (aq) Na 2 SO 4 (aq) BaSO 4 (s) 2 NaNO 3 (aq) Here the barium sulfate precipitates out of solution Solute Solvent Solution Aqueous Precipitate 3
4 2.5 Mixtures and Solutions 2.6 Solubility Rules Electrolyte Strong electrolyte Weak electrolyte Non-electrolyte Compounds Containing Solubility Exceptions Group I (Li, Na, K ) and NH 4 Soluble Nitrates Soluble Acetates (CH 3COO ) Soluble Perchlorates (ClO 4 ) Soluble Chlorides, Bromides, Iodides Soluble Ag, Pb 2, Hg Fluorides Soluble Mg 2, Ca 2, Ba 2, Ag, Pb 2 Sulfates Soluble Ca 2, Ba 2, Pb 2, Hg Sulfides Insoluble Group I ions and NH 4 Carbonates Insoluble Group I ions and NH 4 Phosphates Insoluble Group I ions and NH 4 Hydroxides Insoluble Group I ions and NH Precipitation Reactions 2.8 Ionic Equations Precipitation reactions When an ion appears on both sides of a chemical equation it can be canceled out Spectator Ions Writing Equations Molecular Total ionic Net Ionic Equation Net ionic 4
5 2.9 Ionic Equations 2.10 Ionic Equations Mg(ClO 4 ) 2 (aq) K 2 CO 3 (aq) MgCO 3 (s) KClO 4 (aq) t Ba(NO 3 ) 2 (aq) Na 2 SO 4 (aq) BaSO 4 (s) NaNO 3 (aq) KNO 3 (aq) CaCl 2 (aq) 3.1 Acids and Bases Chapter Concentrations of Acids Acid The ph scale Base Strong acid Strong base 5
6 Chapter Concentrations of Acids 3.4 Acid-Base Reactions The ph scale HNO 3 (aq) NaOH(aq) 3.5 Acid-Base Reactions 4.1 Combination Reactions H 2 SO 4 (aq) KOH(aq) Combination reactions have the form A B C Two or more reactants produce a single product 2 H 2 O 2 2 H 2 O = H = O 6
7 4.2 Combination Reactions Other combination reactions: 2 Na(s) S(s) Na 2 S(s) SO 3 (g) H 2 O(l) H 2 SO 4 (aq) 4.3 Decomposition Reactions Further examples: 2 HgO(s) 2 Hg(l) O 2 (g) This reaction was used by Joseph Priestley in the discovery of oxygen in 1774 CaCO 3 (s) CaO(s) CO 2 (g) This reaction is used industrially to produce both lime (CaO) and CO 2 from limestone (CaCO 3 ) 4.4 Decomposition Reactions Decomposition reactions have the form A B C Single reactant breaks down into two or more products 4.5 Single Replacement Single replacement reactions are also called substitution reactions They have the form A BX B AX, where A and B are elements and BX and AX are compounds 2 H 2 O 2 2 H 2 O O 2 = H = O H 2 CuO = H = O H 2 O Cu = Cu 7
8 4.6 Single Replacement 4.7 Gas Forming Reactions Other single replacement reactions 3 C(s) 2 Fe 2 O 3 (s) 4 Fe(s) 3 CO 2 (g) Cu(NO 3 ) 2 (aq) Zn(s) Zn(NO 3 ) 2 (aq) Cu(s) Reactions leading to the formation of a gas Examples Single replacement reactions are also oxidationreduction reactions Oxidation Reduction 8
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Objectives: Identify, define, and explain: combination reaction, synthesis reaction, decomposition reaction, single replacement reaction, double replacement reaction, combustion reaction, rapid oxidation,
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