KAP Chemistry Notes: Chemical Reactions PDF
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This document contains notes on chemical reactions, including types, balancing chemical equations, and predicting products and practice questions. It is designed for a secondary school chemistry course.
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Name__________KEY KAP Chemistry Period Notes: Chemical Reactions REVIEW: A chemical reaction is the process by which the atoms of one or more substances are rearranged to form one or more new substances...
Name__________KEY KAP Chemistry Period Notes: Chemical Reactions REVIEW: A chemical reaction is the process by which the atoms of one or more substances are rearranged to form one or more new substances. This is typically shown by a chemical equation. Besides the general indicator of a chemical reaction (new substance(s) is/are formed), other specific signs of a chemical reaction are: 1. a gas is produced. 2. light is produced. 3. a solid precipitate is formed from 2 solutions. 4. temperature change (temp can go up – exothermic or down - endothermic) 5. permanent color change Some elements do not occur as single atoms when uncombined with other elements.They will bond with themselves, forming a molecule. In any chemical reaction, when you see these elements alone, they must be shown with a diatomic formula: Diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2 “I Bring Clay For Our New Hut” OR “HOBrFINCl” The particle for an element is an atom. The particle for a diatomic element is a molecule. The particle of a molecular compound is a molecule. The particle for an ionic compound is a formula unit. I. Describing Chemical Change A. Writing Equations Words can be used to describe chemical reactions, but that can become long and awkward, not to mention less useful. Chemists use chemical equations to describe reactions. In chemical equations, formulas and symbols are used to represent the reaction. *** HENCE THE IMPORTANCE OF KNOWING THE NOMENCLATURE RULES!!!*** 1 Symbols Used in Chemical Equations: Symbol Meaning + combine with, reacts with → yields, forms, produces reversible reaction (s) solid (l) liquid (g) gas (aq) dissolved in water, aqueous solution heat is supplied to the reaction (s) or () precipitate is formed (2 possibilities) (g) or () gas is released (2 possibilities) MnO2 (Ex. MnO2 is the catalyst) catalyst Ex 1: Solid iron reacts with oxygen gas to produce iron (III) oxide (rust). 4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) (We will learn how to balance chemical equations with coefficients in the next section) Ex 2: Solid sodium bicarbonate reacts with hydrochloric acid to produce aqueous sodium chloride, water, and carbon dioxide gas. NaHCO3(s) + HCl(aq) → NaCl(aq) + H2O(l) +CO2(g) Note: water is assumed to be a liquid, H2O(l) (unless it is in a combustion reaction when it is a gas H2O(g)). Acids are always aqueous (aq). Practice: Write equations for each chemical reaction: 1. Solid sulfur burns in oxygen gas to form sulfur dioxide gas. S(s) + O2(g) → SO2(g) 2. Heating solid potassium chlorate in the presence of the catalyst manganese(IV) oxide produces oxygen gas and solid potassium chloride. MnO2 KClO3(s) → O2(g) + KCl(s) (We must balance this equation) MnO2 2 KClO3(s) → 3 O2(g) + 2 KCl(s) 3. A solution of lithium hydroxide reacts with hydrochloric acid to produce aqueous lithium chloride and water. LiOH(aq) + HCl(aq) → LiCl(aq) + H2O(l) 4. Sodium metal reacts with water to form a solution of sodium hydroxide and hydrogen gas. Na(s) + H2O(l) → NaOH(aq) + H2(g) (we must balance this equation) 2 Na(s) + 2 H2O(l) → 2 NaOH(aq) + H2(g) 2 B. Balancing Chemical Equations In order to correctly balance chemical equations, always remember the law of conservation of mass (In any chemical reaction, mass cannot be created or destroyed). An equation that gives the same number of each type of atom on the left and right sides of the equation is called a balanced equation. Numbers called coefficients are placed in front of the formulas. When no coefficient is written, it is assumed to be 1. The coefficient tells you the number of atoms, molecules or Formula units of the substance used or produced in the reaction. The coefficient also tells you the number of *** MOLES *** of the element orcompound used/produced in the reaction. Interpreting a balanced equation in two ways: 4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) 1) 4 atoms of solid iron react with 3 molecules of oxygen gas to produce 2 formula units of solid iron(III) oxide. 2) 4 moles of solid iron react with 3 moles of oxygen gas to produce 2 moles of solid iron(III) oxide. Tips for Balancing Chemical Reactions Determine the correct formulas for all the reactants and products. Write formulas for reactants on the left and formulas for products on the right with → in between. Separate multiple reactants and products with a + sign. Count the number of atoms of each element in the reactants and products. Balance the elements one at a time by using coefficients. - it’s easier to begin with elements that appear only once. - never balance by changing subscripts within a formula. - once you determine correct formulas, they should never be changed. - many times, you can treat polyatomics as one unit (if they appear on both sides of the equation) Check the numbers of each atom or polyatomic ion to be sure the equation is balanced. Finally, make sure all the coefficients are in the lowest possible whole number ratio. Reduce if necessary. ALL coefficients must reduce if you do this! Ex 1: Write a balanced equation for the reaction of copper metal and an aqueous solution of silver nitrate to form aqueous copper(II) nitrate and silver metal. Reactants → Products Before balancing: Cu(s) + AgNO3(aq) → Cu(NO3)2(aq) + Ag(s) Atom inventory: Cu 1 1 Ag 1 1 NO3 1 2 After balancing: Cu(s) + 2 AgNO3(aq) → Cu(NO3)2(aq) + 2 Ag(s) Ex 2: Balance the following reaction: Before balancing: Al(s) + O2(g) → Al2O3(s) Atom inventory: Al 1 2 O 2 3 After balancing: 4 Al(s) + 3 O2(g) → 2 Al2O3(s) 3 Practice: Write balanced chemical equations: 1. Potassium metal and water react to form aqueous potassium hydroxide and hydrogen gas. 2 K(s) + 2 H2O(l) → 2 KOH(aq) + H2(g) 2. Aqueous calcium hydroxide and sulfuric acid react to form aqueous calcium sulfate and water. (Hint: balance sulfate as 1 unit; balance hydroxide as 1 unit by writing water as HOH instead of H2O) Ca(OH)2(aq) + H2SO4(aq) → CaSO4(aq) + 2 H2O(l) 3. Balance the following equations: a. 2 SO2 + O2 → 2 SO3 b. Fe2O3 + 3 H2 → 2 Fe + 3 H2O c. 4 P + 5 O2 → P4O10 d. 2 Al + N2 → 2 AlN e. * 2 C2H6 + 7 O2 → 4 CO2 + 6 H2O f. * C3H8 + 5 O2 → 3 CO2 + 4 H2O * Trick for balancing combustion reactions like (e) and (f): 1. Balance H first by adding the correct coefficient in front of H2O. The coefficient must be even, if odd, double it and rebalance H. 2. Balance C second. 3. Balance O last. 4 II. Types of Chemical Reactions A. Classifying Chemical Reactions There are 5 basic types of chemical reactions: 1. Synthesis (S): two or more substances combine to create a more complex substance. General Form: A + X → AX Guideline: 1 PRODUCT (COMPOUND) Ex: 4Na(s) + O2 (g) → 2Na2O(s) 2. Decomposition (D): a complex substance is broken down into two or more simpler substances. General Form: AX → A + X Guideline: 1 REACTANT (COMPOUND) Ex: 2KCl(s) → 2K(s) + Cl2 (g) 3. Single replacement (SR): one element takes the place of another in a compound. Guideline: ELEMENT + COMPOUND(aq) → ELEMENT + COMPOUND(aq) 2 types: Type 1: Cation (metal) replacement Type 2: Anion (halogen) replacement General Form: AX + D → DX + A General form: AX + Z → AZ + X Ex: MgBr2(aq) + 2Na(s) → 2NaBr(aq)+ Mg(s) Ex: 2NaCl(aq)+ F2(g) → 2NaF(aq)+ Cl2(g) 4. Double replacement (DR): ions from two aqueous ionic compounds switch places. Guideline: COMPOUND(aq) + COMPOUND(aq) → COMPOUND + COMPOUND General Form: AX + DZ → DX + AZ Ex: LiCl(aq) + AgNO3(aq) → AgCl(s) + LiNO3(aq) 5. Combustion (C): a substance combines with oxygen and burns, releasing a large amount of energy in the form of light and/or heat. We will only focus on hydrocarbon combustion in this class. General Form: (hydrocarbon combustion) CXHY + O2 (g) → CO2 (g) + H2O(g) (Hydrocarbon) Ex: CH4(g) + 2O2(g) → CO2 (g) + 2H2O(g) Practice: Balance the following equations and identify the type of reaction: TYPE 1. Pb(NO3)2 + K2CrO4 → PbCrO4 2 KNO3 ______DR___________ 2. 2 C3H6 + 9 O2 → 6 CO2 + 6 H2O ______C____________ 3. 2 K + ZnCl2 → 2 KCl + Zn _____SR (metal)____ 4. 4 Li + O2 → 2 Li2O ______S____________ 5. _____ MgCO3 → MgO + CO2 (BALANCED!) ______D___________ 6. ____ Cl2 + 2 KI → 2 KCl + I2 _____SR (halogen)__ 5 B. Predicting Products of Chemical Reactions 1. Synthesis, Decomposition and Combustion To predict the products of a chemical reaction, you must first recognize the type ofreaction: Reactants Probable type element + element Synthesis (S) 1 compound Decomposition (D) element + compound(aq) Single replacement (SR) compound(aq) + compound(aq) Double replacement (DR) Hydrocarbon (CXHY) + oxygen Combustion (C) You then need to be familiar with the mechanism of that reaction so you can predict the product(s). Ex 1: Solid magnesium reacts with solid sulfur → ? Type of reaction: synthesis Write and balance: Mg(s) + S (s) → MgS(s) Ex 2: Solid copper(I) oxide ⎯⎯ → Type of reaction: decomposition Write and balance: 2 Cu2O(s) ⎯⎯ → 4 Cu(s) + O2 (g) Ex 3: Methane (CH4) gas reacts with oxygen gas ⎯⎯ →? Type of reaction: combustion Write and balance: CH4(g) + 2 O2(g) ⎯⎯ → CO2 (g) + 2 H2O(g) Practice: Classify by reaction type and write a fully balanced equation by predicting products: TYPE __S___ 1. Sodium metal + chlorine gas → 2 Na(s) + Cl2(g) + → 2 NaCl(s) __C___ 2. Propane (C3H8) gas + oxygen gas → C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g) __D___ 3. Sulfur dioxide gas → SO2(g) → S(s) + O2(g) 6 2. Detailed Instructions: Single Replacement Reactions Activity Series of Metals and H2 One element will replace another element in a compound IF it is more reactive. Otherwise, no reaction occurs. a) Activity Series of Metals and H2 DECREASING REACTIVITY Hydrogen can also participate in single replacement reactions with metals. (See H2 in activity series). Whether one metal will displace another metal from a compound can be determined by the relative reactivity of the two metals. The activity series of metals lists metals in order of decreasing reactivity. (Meaning metals at the bottom are less reactive than those at the top). In a single replacement reaction, a metal will replace any metal that is below it in the activity series. Ex 1: ___Mg(s) + Zn(NO3)2(aq) Mg(NO3)2(aq)+ Zn(s) Ex 2: ___Mg(s) + LiNO3(aq) No reaction (NR) b) Activity Series of Halogens Activity Series of Halogens A non-metal can also replace another non-metal. (On periodic table) This is usually limited to halogens, and they have their own activity series. The activity of halogens decreases as you move down group 17 on the periodic table. Ex 1: 2KI(aq) + Br2(l) → 2KBr(aq) + I2(s) Ex 2: CaF2(aq) + I2(s) → No reaction (NR) Practice – Single replacement: Decide if the following single replacement reactions will occur, then predict the products and balance the equation. If no reaction occurs, write NO REACTION (NR). Cation replacement (metal or hydrogen replacement) 1. Cu(s) + KNO3(aq) → NO REACTION 2. 2 Na(s) + 2 H2O(l) → 2 Na(OH)(aq) + H2(g) (Hint: think of H2O as HOH) Anion replacement (halogen replacement) 4. Cl2(s) + 2 LiBr(aq) → 2 LiCl(aq) + Br2(l) 3. I2(s) + NaF(aq) → NO REACTION Note: the more reactive metal or halogen is always in the compound on the product side 7 3. Detailed Instructions: Double-Replacement Reactions All double-replacement reactions produce either a precipitate (abbreviated ppt, an insoluble solid made from 2 solutions), a gas, or water (H2O(l). We will focus on the formation of a precipitate in this unit. To predict products, simply swap the cations (or anions) in each reactant. The ions “switch partners.” After switching partners, in order to determine if a precipitate forms, consult the solubility rules. (The solubility chart to the right is provided on assessments on your formula chart). If a product is insoluble, a solid(s) precipitate is formed, and a reaction occurs. If both products are soluble(aq) no reaction occurs. Making a solution of an ionic compound. Recall that ionic compounds (also called salts) are all solid crystals at room temperature. When ionic compounds are soluble in water, we can make an aqueous solution by dissolving the solid crystals in water. Ex : Making an aqueous solution of sodium chloride NaCl(s) H→ 2O NaCl(aq) When solid sodium chloride crystals are dissolved in water, the Na+ and Cl⁻ ions dissociate (separate) 8 Examples Predict the products for the following pairs of aqueous solutions (DR). Determine whether a precipitate will occur by consulting the solubility rules above. If a precipitate forms (and therefore a reaction occurs), write a balanced equation, INCLUDING ALL STATE SYMBOLS. (aq) means aqueous/soluble; (s) means solid/insoluble/precipitate. If no reaction occurs, (both products are aqueous) write NR. Ex 1: AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq) (AgCl is the precipitate) Ex 2: LiClO3(aq) + KBr(aq) → NR (both possible products – LiBr and KClO3 are soluble (aqueous) Ex 3: An acid reacts with a base. This is a different type of double replacement reaction called “neutralization”. We will discuss neutralization reactions in detail in a later unit. NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l) There is a reaction because water is a liquid and therefore a new molecule is formed. 9 Practice – Double Replacement: Predict the products and write the balanced equation when the following pairs of aqueous solutions are mixed. If there is a reaction, indicate which product is the precipitate. If no reaction occurs, write NR. 1. ammonium bromide and rubidium acetate → NH4Br(aq) + RbC2H3O2(aq) → RbBr(aq) + NH4C2H3O2(aq) NR 2. sodium carbonate and magnesium iodide → Na2CO3(aq) + MgI2(aq) → 2 NaI(aq) + MgCO3(s) (MgCO3(s) is the ppt) Mixed practice – Predicting Products For each of the following, indicate the type of reaction. If no reaction occurs, write NR. Predict the products and write a balanced equation. Indicate which product is the precipitate (ppt) for DR reactions. TYPE ____NR_____ 1. solutions of rubidium hydroxide and lithium chlorite are mixed → NO REACTION because both products are soluble: LiOH(aq) + RbClO2(aq) ____SR_____ 2. fluorine gas is bubbled through a solution of barium chloride → F2(g) + BaCl2(aq) → BaF2(aq) + Cl2(g) ____S______ 3. potassium metal and solid sulfur are combined → 2 K(s) + S(s) → K2S(s) ____DR_____ 4. a solution of calcium chloride is added to a solution of lead(II) nitrate → CaCl2(aq) + Pb(NO3)2 → Ca(NO3)2(aq) + PbCl2(s) (PbCl2(s) is the ppt) _____C_____ 5. methane (CH4(g)) gas burns with oxygen gas → CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) ____NR_____ 6. hydrogen gas is bubbled through a solution of zinc acetate → NO REACTION because H cannot replace Zn _____D_____ 7. solid lithium oxide is heated → 2 Li2O(s) → 4 Li(s) + O2(g) _____SR_____ 8. iron metal is added to a solution of tin(II) sulfate → (use Fe3+) 2 Fe(s) + 3 SnSO4(aq) → Fe2(SO4)3(aq) + 3 Sn(s) ____NR_____ 9. solid iodine is added to a solution of cobalt(II) bromide → NO REACTION because I cannot replace Br 10 C. REDOX REACTIONS Assigning Oxidation Numbers A number assigned to an element, based on the distribution of electrons. The same element can have very different properties in different oxidation states. Rule Example – Oxidation # 1 The oxidation number of any uncombined element is 0. Na(s) = 0 O2(g) = 0 2 The ox. # of a monatomic ion equals the charge of the ion. Cl⁻ = –1 Na+ = +1 3 In compounds, Group 1 & 2 elements and Al have ox. #s of Ox # of Na in Na2SO4 = +1 +1,+2, +3, respectively. (same as the charge of their ion). Ox. # of Ca in CaCO3 = +2 Ox # of Al in AlCl3 = +3 4 The more electronegative element in a binary molecular Ox. # of O in NO = –2 compound is assigned the # equal to the charge it would have if it (O is more electronegative were an ion. than N) Ox # of Cl in CCl4 = –1 (Cl is more electronegative than C) 5 The ox. # of fluorine in a compound is always –1. Ox. # of F in LiF = –1 Ox # of F in CaF2 = –1 6 The sum of ox. # of all atoms in a neutral compound is 0. Ox. # of C in CaCO3 = +4 1. 1We know ox # of Ca = +2 2. We know ox # of each O = –2 –2 x 3 Oxygens =–6 3.Therefore ox # of C = +4 Sum is 0 = 0 7 The sum of the ox. # of all atoms in a polyatomic ion equals Ox. # of S in SO42⁻ = +6 the charge of the ion. 1. We know ox # of each O = –2 –2 x 4 oxygens = –8 2. Therefore ox # of S = +6 sum is charge of SO42⁻ = –2 8 Oxygen has an ox. # of –2 Ox. # of O in NO2 = –2 Exceptions: when combined with F, it is +2. Ox. # of O in OF2 = +2 in a peroxide like H2O2 or Na2O2, it is –1. Ox. # of O in Na2O2 = –1 9 The ox. # of H in most compounds is +1 Ox. # of H in H2O = +1 Exception: when it is combined with a metal, it is –1. Ox. # of H in LiH = –1 Practice Assign oxidation numbers to the underlined element. When assigning oxidation numbers, we assign the oxidation for 1 atom of the element. Example: The oxidation # of each F in SrF2 = –1 1. Ca 0 5. Al2O3 +3 9. NaH +1 2. H2 0 6. FeCl3 +3 10. PH3 –3 3. MgCl2 +2 7. PCl5 +5 11. PO33⁻ +3 4. K2S –2 8. N2O4 +4 12. CO32⁻ +4 11 Oxidation is a reaction in which the atoms or ions of an element experience an increase in oxidation state. An increase in oxidation state means that the oxidation number becomes more positive. Therefore, the species has lost electrons. 0 +1 Ex: Na → Na+ + e⁻ Reduction is a reaction in which the atoms or ions of an element experience a decrease in oxidation state. A decrease in oxidation state means that the oxidation number becomes more negative. Therefore, the species has gained electrons. 0 ⁻1 Ex: Cl2 + 2e⁻ → 2Cl⁻ You can use “OIL RIG”: Oxidation Is Loss (of electrons) Reduction Is Gain (of electrons) Since oxidation is the losing of electrons and reduction is the gaining of electrons, they have to occur simultaneously, and the number of electrons lost in oxidation must equal the number of electrons gained in reduction. **Any chemical reaction in which at least 2 elements undergo changes in oxidation number is an oxidation-reduction reaction, or redox reaction.** Practice: Determine whether the following elements have been oxidized or reduced and label the reaction type. Example 1: 4Fe + 3O2 → 2Fe2O3 TYPE of Reaction: __Synthesis_____ Element Ox.# Reactants side Ox.# Products side Lose/Gain e- Oxidized/Reduced Fe 0 +3 lose Oxidized O 0 ⁻2 Gain Reduced Example 2: Mg + 2HCl → MgCl2 + H2 TYPE of Reaction: _____SR_____ Element Ox.# Reactants side Ox.# Products side Lose/Gain e- Oxidized/Reduced Mg 0 +2 Lose Oxidized H +1 0 Gain Reduced Cl ⁻1 ⁻1 Neither Neither Example 3: LiCl + NaOH → NaCl + LiOH TYPE of Reaction: ____DR______ Element Ox.# Reactants side Ox.# Products side Lose/Gain e- Oxidized/Reduced Li +1 +1 Neither Neither Cl ⁻1 ⁻1 Neither Neither Na +1 +1 Neither Neither O ⁻2 ⁻2 Neither Neither H +1 +1 Neither Neither Summarizing tip: All of the 5 basic reaction types (with a few exceptions) we have learned are redox reactions except for one. What is it? DOUBLE REPLACEMENT 12