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Chapter 11 Chemical Reactions Hingham High School Mr. Dan Clune All chemical reactions • Two parts: • Reactants – what you start with • Products- what you end with • Reactants turn into the products. • Reactants Products In a chemical reaction • The way atoms are joined is changed. • Atoms aren’t created or destroyed. In a chemical reaction • Can be described several ways: 1. In a sentence Copper reacts with chlorine to form copper (II) chloride. 2. In a word equation Copper + chlorine copper (II) chloride Symbols in equations-p.323 Cu + Cl2 CuCl2 • Arrow () separates reactants from products • Read “reacts to form” • Plus (+) sign read “and” Symbols used in equations Cu(s) + Cl2(g) CuCl2(s) • • • • (s) = solid (g) = gas (l) = liquid (aq) - dissolved in water, an aqueous solution. Symbols used in equations • ↑ after product, indicates gas produced •same as (g) - H2↑ •after product, indicates solid produced •same as (s) - PbI2↓ Symbols used in equations indicates reversible reaction heat shows that heat is supplied to the reaction Pt is - indicates a catalyst is supplied, in this case, platinum. , What is a catalyst? • Speeds up a reaction • Is NOT changed or used up by the reaction. • Enzymes are biological or protein catalysts. Skeleton Equation • Uses formulas to describe a reaction • doesn’t indicate how many. • NOT balanced Convert this to an equation • Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas. Fe2S3(s) + HCl(g) FeCl3 + H2S(g) Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water, carbon dioxide gas, sodium nitrate dissolved in water. HNO3(aq) H2O(l) + Na2CO3(aq) + CO2(g) + NaNO3(aq) Now, read these: • Fe(s) + O2(g) Fe2O3(s) • Cu(s) + AgNO3(aq) Ag(s) + Cu(NO3)2(aq) • NO2 (g) Pt N2(g) + O2(g) Balancing Chemical Equations Balanced Equation • Atoms can’t be created or destroyed • All the atoms we start with we must end up with • A balanced equation has the same number of each element on both sides of the equation. C + O O O C O • C + O2 CO2 • This equation is already balanced • What if it isn’t? C + O O C O • C + O2 CO • Need one more O in the products. • Can’t change the formula, because it describes what it is (carbon monoxide in this example) C + O O C O C O • Must be used to make another CO • But where did the other C come from? C + C O O C O C O • Must have started with two C • 2 C + O2 2 CO Finding the number of atoms • The subscript in front of an element is the number of atoms of that element / polyatomic. • Ex) CO2 B2(SO4)3 • C= 1 O= 2 B= 2 SO4= 3 • A coefficient in front of the formula multiplies the amount of elements by the coefficient. • Ex) 3CO2 2B2(SO4)3 • C= 3 O= 6 B= 4 SO4= 6 Finding the number of atoms H2 O H= 2 O= 1 2H2O H= 4 O= 2 B(NO3)2 B= 1 NO3= 2 3B(NO3)2 B= 3 NO3= 6 Pb3(PO4)4 Pb= 3 PO4= 4 2Pb3(PO4)4 Pb= 6 PO4= 8 Rules for balancing: . Determine the correct formulas for reactants and products. . Write a skeleton equation. Rules for balancing: 3. Count the number of atoms of each element appearing on both sides of the equation. 4. Balance the elements one at a time by adding coefficients (the numbers in front) -Save H and O until LAST! Rules for balancing: 5. Check to make sure it is balanced. 6. Make sure the coefficients are in the lowest possible ratio. Don’t you ever… • Never change a subscript to balance an equation. • H2O is a different compound than H2O2 • Never put a coefficient in the middle of a formula X • 2 NaCl is okay, Na2Cl is not. Example 2 H2 + O2 2 H2O R P 4 2 H 2 4 2 O 1 2 2 H2 + O2 2 H2O 2 AgNO3 + Cu Cu(NO3)2 + 2Ag R P 21 1 2 1 Ag Cu NO3 12 1 2 3Mg + N2 Mg3N2 R P 3 1 Mg 2 N 3 2 4 P + 5 O2 P4O10 R 41 10 2 P P 4 O 10 2 Na + 2H2O H2 + 2 NaOH R 21 4 2 2 1 P Na H O 12 3 4 1 2 CH4 + 2 O2 CO2 + 2 H2O R P 1 4 42 C H O 1 2 4 3 4 Section 11.2 Types of Chemical Reactions • OBJECTIVES: • Identify a reaction as combination, decomposition, single-replacement, doublereplacement, or combustion Types of Reactions • 5 major types. • predict the products • predict whether or not they will happen at all • How? We recognize them by their reactants #1 - Combination Reactions • Combine - put together • 2 or more substances combine to make one compound. A + B • Ca +O2 CaO • SO3 + H2O H2SO4 #1 - Combination Reactions • We can predict the products if they are two elements. 3 Mg + N2 Mg3N 2 Write and balance • Ca + Cl2 CaCl2 • 2 Fe + O2 iron (II) oxide 2 FeO • 4 Al + 3O2 2 Al2O3 • Remember that the first step is to write the correct formulas • Then balance by using coefficients only #1 – Combination Reactions • Additional Notes: a) Some nonmetal oxides react with H2O - produces acid: SO2 + H2O H2SO3 (how “acid rain” forms) b) Some metallic oxides react with H2O - produces base: CaO + H2O Ca(OH)2 #2 - Decomposition Reactions • Decompose = fall apart • One reactant falls apart into two or more elements or compounds. AB A + B #2 - Decomposition Reactions •NaCl Na + Cl2 •CaCO3 CaO + CO2 •Note that energy is usually required to decompose electricity #2 - Decomposition Reactions • Binary compounds (made of 2 elements) falls apart into its elements • 2H2O • 2 HgO electricity 2 H2 O + 2 2 Hg + O2 #2 - Decomposition Reactions • If the compound has more than two elements, one product must be given. • The other product will be from the missing pieces • NiCO3 CO2 + Ni • H2CO3(aq) CO2 + ? H2 #3 - Single Replacement Reactions • One element replaces another (new dance partner) • Reactants are an element & cpd • Products will be a different element and different cmpd • Li + KCl K + LiCl (Cations switched) • F2 +2 LiCl 2LiF + Cl2 (Anions switched) #3 Single Replacement Reactions • Metals replace other metals (and H) • 3 K + AlN K3N + Al • Zn + 2 HCl ZnCl2 + H2 • Think of water as: HOH • Metals replace first H, then combines w/ hydroxide (OH). • 2 Na + 2 H2O 2 NaOH + H2 #3 Single Replacement Reactions • Sometimes, the reaction will not happen: Some chemicals are more “active” than others • More active replaces less active The “Activity Series” of Metals Higher • Lithium activity • Potassium • Calcium • Sodium • Magnesium • Aluminum • Zinc • Chromium • Iron • Nickel • Lead • Hydrogen • Bismuth • Copper • Mercury • Silver Lower • Platinum activity • Gold If the lone metal is above the paired metal, replacement will occur. Ex) Li + NaCl Na + LiCl If the lone metal is below the paired metal, replacement will not occur. Ex) Na + LiCl Na + LiCl The “Activity Series” of Halogens Higher Activity Fluorine Chlorine Bromine Iodine If the lone halogen is above the paired halogen, replacement will occur. Lower Activity 2NaCl(s) + F2(g) 2NaF(s) + Cl2(g) MgCl2(s) + Br2(g) No Reaction! #3 Single Replacement Reactions Practice: Higher • Lithium activity • Potassium • Calcium • Sodium • Magnesium • Aluminum • Zinc • Chromium • Iron • Nickel • Lead • Hydrogen • Bismuth • Copper • Mercury • Silver Lower • Platinum activity • Gold • Fe + CuSO4 FeSO4 + Cu • Pb + KCl No reaction • 2 Al + 6 HCl 2 AlCl3 + 3 H2 #4 - Double Replacement Reactions • Two things replace each other. • Reactants must be two ionic compounds. • NaOH + FeCl3 • positive ions change place • NaOH + FeCl3 Fe+3 OH- + Na+1 Cl-1 = NaOH + FeCl3 Fe(OH)3 + NaCl #4 - Double Replacement Reactions • Have certain “driving forces”, or reasons • only happens if one product: a) doesn’t dissolve in water & forms solid (a “precipitate”), or b) is gas that bubbles out, or c) is molecular compound (usually water) Complete and balance: • assume all of the following reactions actually take place: CaCl2 + 2 NaOH Ca(OH)2 + 2 NaCl CuCl2 + K2S CuS + 2 KCl 3 KOH + Fe(NO3)3 3 KNO3 + Fe(OH)3 K2SO4 + BaF2 2 KF + BaSO4 How to recognize which type? • Look at the reactants: El + El = Combination Cpd = Decomposition El + Cpd = Single replacement Cpd + Cpd = Double replacement Practice Examples: • • • • • • • 2 H2 + O2 2H2O 2 H2O 2 H2 + O2 Zn + H2SO4 H2 + ZnSO4 2 HgO 2 Hg + O2 2 KBr + Cl2 2 KCl + Br2 AgNO3 + NaCl AgCl + NaNO3 Mg(OH)2 + H2SO3 MgSO3 + 2 H2O #5 – Combustion Reactions • Combustion means “add oxygen” • Normally, a cpd composed of only C, H, (sometimes O) is reacted with oxygen – called “burning” • Complete combustion, products are CO2 and H2O Combustion Reaction Examples: • C4H12 + 7 O2 4 CO2 + 6 H2O • C6H12O6 + 9 O2 6 CO + 6 H O 2 2 • C8H8 + 10O2 8 CO + 4 H O 2 2 SUMMARY: An equation... • • • • Describes a rxn Must be balanced only balance by changing coefficients special symbols to indicate physical state, catalyst or energy required, etc. Reactions • 5 major types • We can tell what type they are by looking at reactants • Single Replacement happens based on the Activity Series Section 11.3 p. 342 Reactions in Aqueous Solution NiCl2 Co(NO3)2 K2Cr2O7 K2CrO4 CuSO4 KMnO4 Net Ionic Equations • Many reactions occur in water, or aqueous solution • When dissolved in water, many ionic cpds “dissociate”, or separate, into cations & anions • Now write ionic equation Net Ionic Equations • Example (needs to be a double replacement reaction) AgNO3 + NaCl AgCl + NaNO3 1. this is the full balanced equation 2. next, write it as ionic equation by splitting the cpds into their ions: Ag1+ + NO31- + Na1+ + Cl1- AgCl(s) + Na1+ + NO31Solids do not split up. Net Ionic Equations 3. Crossing out ions that did not change (called spectator ions) Ag1+ + Cl1- AgCl (s) This is the net ionic equation Let’s talk about precipitates before we do some other examples Predicting the Precipitate • Insoluble salt is a precipitate • i.e. a solid • General solubility rules are found: a) Table 11.3, p. 344 in textbook Solubility Rules BaCl2 + 2 AgNO3 → Ba(NO3)2 + 2 AgCl (s) 1. Break up into ions Ba+2 + 2Cl- + 2Ag+ + 2NO3- Ba+2 + 2NO3- + 2AgCl (s) 2. Now write the net ionic equation. 2Ag+ + 2NO3- 2AgCl (s) 2 NaCl + Ba(NO3)2 → 2 NaNO3 + BaCl2 1. Break up into ions 2Na+ + 2Cl- + Ba+2 + 2NO3- 2Na+ + 2NO3- + Ba+2 + 2Cl- 2. Now write the net ionic equation. Nothing changed, so no net ionic equation. PbCl2 (s) + Li2O PbO + 2 LiCl 1. Break up into ions PbCl2 (s) + 2Li+ + O-2 Pb+2 + O-2 + 2Li+ + 2Cl2. Now write the net ionic equation. PbCl2 (s) Pb+2 + 2Cl-