Chapter 10 Chemical Reactions

1. Chapter 10 Chemical Reactions Chemical changes and reactions between compounds Chemical formula revisited: types of elements and ratio making up compound empirical formula : simplest whole number ratio of elements in a compound -general description of how to assemble the compound ionic-exists as separate ions -empirical formula is the right description -simple mixture of ionsNa+ and Cl- ions molecular (covalent)-exists as a particular bonded entity -empirical formula tells what it’s made of -does not describe how it is bonded molecular formula : identifies actual numbers of atoms in a molecule-discrete units of atoms -not appropriate for ionic – array of separate ions -moleculesdescribed by molecular formula -you know a formula is molecular if the numbers are not the simplest whole number ratio glucose: C6H12O6 molecular since divisible by 6 empirical – CH2O

2. Making compounds – cannot count atoms need to know about masses of compounds Formula weight (mass) –mass of one formula unit (amu) sum of atomic weights for ALL atoms in chemical formula general term for ionic and molecular Molecular weight (mass) - mass of molecule add up all atomic weights for each atom in molecule specific to molecules only (covalent) use atomic mass since it is an average over all isotopes – naturally abundant species Formula weight of NaCl: Molecular weight of C12H22O12

3. Also want to be able to describe percent amount of substances in compound Identification of substances percent composition: what percent by mass is from each element % composition = mass of the element in compound total formula weight X100% CAN USE EMPIRICAL OR MOLECULAR Find % comp of each element in NaCl: Ex: find % comp of each element in C8H18 empirical easier C4H9 Try: H3BO3 Ba(NO3)2

4. Chemical Equations Describes what it takes for a chemical reaction to occur (A recipe for producing compounds) Which involves changes in compounds (one to another), states of matter (liquid, solid, gas), and exchanges in energy. GENERAL CHEMICAL EQUATION ReactantsProducts To quantitatively show how to make reaction succeed, you have to describe the numbers of atoms involved LAW OF CONSERVATION OF MASS Matter is neither created nor destroyed The products must balance the reactants C + O2 CO2 C = 1 C = 1 O = 2 O = 2 NEW REACTION: synthesis of carbon monoxide C + O2 CO (unbalanced) Unequal – mass of Reactants different than for products 2C + O22CO Now mass of Reactants is equal with mass of Products Balance equations by changing coefficients NOTE: The formula CAN’T be changed

5. Be careful to take inventory of ATOMS in equation! Ex: CH4 + 2O2 CO2 + 2H2O C=1 C=1 O=4 O=4 H=4 H=4 Many times, Balancing equations is a trial & error process Ex: Combustion of Gasoline (Octane) 2C8H18(g) + 25O2(g) 16CO2(g) + 18H2O(g) However, you should be familiar with the rules which describe balanced chemical reactions. 1. Number of Atoms of each element conserved in reactants and products 2. Cannot change formula of reactants or products 3. Can only change coefficients to balance equation Hints to help in balancing equations 1. Balance compounds with biggest numbers of atoms first 2. Treat polyatomics as single units, especially when on both sides of equation 3. Fractional coefficients are useful

6. Can also identify physical states of compounds in reaction (g) – compound is gaseous (l) – compound is liquid (s) – compound is solid (aq) – compound is an aqueous solution (in water) () - temperature change (energy) (or ) - solid precipitates out (or gas escapes) Precipitation of calcium bicarbonate from hard water - changes in state Ca(HCO3)2 (aq) + Na2CO3 (aq) 2NaHCO3 (aq)+ CaCO3 precipitate HYDROCARBONS-combinations of hydrogen and carbon -important energy source - combustion gasoline, acetylene, propane -react with O2 to form H2O and CO2plus ENERGY Example: butane 2C4H10 + 13O2 8CO2 + 10H2O (lighters) + energy Carbohydrates: hydrocarbon with oxygen plants store sun energy as carbs (photosythesis) we eat and respire oxygenenergy + CO2 GLUCOSE : C6H12O6(s) + 6O2(g) 6CO2(g) + 6H2O(g) Enzymes - slowly release energy in plants and animals free burning produces quick energy release and flame

7. Types of Chemical Reactions Oxidation-Reduction (REDOX) Reaction - GENERAL -electrons transferred from one atom to another -one element oxidized (loses e-) and the other is reduced (gains e-) <balance of charge-losses cancel increases> -oxygen often involved, but F, Cl & other nonmetals do the same thing OXIDIZING AGENT - takes electrons from other substances - takes e- away causing substances to be oxidized - O2 in food and fuels gains e- to form octet - Cl & F in bleach kill bacteria - pools REDUCING AGENT - provides e- to substance being reduced - carbon: gives e- to form octet - REDUCTION OF IRON ORE [iron(III) oxide] 2Fe2O3(s)+ 3C(s) 4Fe(s) + 3CO2 Note reverse crossover To get charge for name BUT reactions are normally defined in terms of the effect on reactants and products: reaction type 3 types of redox eqs.: combine, decompose, replace Ion exchange - not redox since no change in charge ions just replaced

8. Reaction Types Combination (Synthesis) Reaction two or more substances combine to form one X + Y XY lower energy oxidation of metals- rust, burning burning of Mg metal 2Mg(s) + O2(g) 2MgO(s) rusting of iron 4Fe(s) + 3O2(g) 2Fe2O3(s) burning non-metals C(s) + O2(g) CO2(g) Decomposition Reaction a compound is broken down into simpler compounds or constituent elements XY X + Y requires energy HYDROLYSIS - decompose water with electricity - hydrogen fuel cells 2H2O(l) 2H2(g) + O2(g) decomposition of mercury(II) oxide 2HgO(s) 2Hg(s) + O2(g) D electricity

9. Replacement Reaction One atom or polyatomic ion in a compound is replaced XY + Z XZ + Y or XY + A AY + X (-) part replaced (+) part replaced chemical activity - tendency of an element to give up electrons more chemically active if unable to keep electrons MORE ACTIVE METALS GIVE UP ELECTRONS TO LESS ACTIVE ACTIVITY SERIES - Fig. 10.12 shows most active elements on top a metal will replace any metal it is above has a larger chemical activity 2Al(s) +3CuCl2(aq) 2AlCl3(aq) + 3Cu(s) Al more chem active - above Cu - gives e- to Cu+2 - Cu precipitates, Al goes into solution if less chem active, e- would stay where they are (no reaction) Alkali and Alkaline metals- very active often replace other elements 2Na(s) + H2O(l) 2NaOH + H2 H takes e- from Na makes Na+ and H2

10. Ion Exchange Reaction Each ion of one compound is replaced by the ions of a second compound Ions mix together to form - a precipitate (insoluble) - a gas - or water removes ions from the solution of ions AX + BY AY + BX one of products must leave solution (precipitate, gas or water) Example: 3Ca(OH)2(aq) + Al2(SO4)3(aq) 3CaSO4(aq) +2Al(OH)3 Application: Dissolved ions form insoluble aluminum hydroxide “net” for water treatment- mixture traps suspended impurity particles Solubility Tables (Appendix B) if the products are soluble --- no reaction if one product is insoluble ---reaction occurs precipitate or insoluble gas Examples: identify the reaction type 2Al(s) + Fe2O3(s) Al2O3(s) + 2Fe(s) 2Ag(s) + S(g) Ag2S(s) NaCl(aq) + AgNO3(aq) NaNO3(aq) + AgCl 2KClO3(s) 2KCl(s) + 3O2

11. Summary of Chemical Equations Tells how compounds combine together to form new substances what atoms are present before and after (inventory) 2H2(g) + O2 2H2O atomic description: 4 atoms H and 2 atoms O form 4 atoms H and 2 atoms O MOLECULAR description: 2 molecules H2 and 1 molecule O2 forms 2 molecules of H2O Also describes masses (formula weight): reactant mass = 2x2.016amu + 32 amu = 36 amu H2 O2 product mass = 2x18amu =36 amu H2O LAW OF CONSERVATION OF MASS NOTE: 2g of H combines with 8 g of O always the same mass ratio remember amu based on weight of C-12 exactly 12 amu! relationship between mass and number historically misunderstood by Dalton mis-measured mass of H and O to form H20 he thought water was HO Led to failure of Daltons atomic theory

12. Gay-Lusaac and Avogadro gases combine in whole number ratio (at constant temperature and pressure-STP) 2H2(g) + O2 2H2O 2 volumes of H2 for each volume O2 gives one volume of water vapor LAW OF COMBINING VOLUMES equal volumes of gas at STP have equal numbers of molecules FIXED formula for water: (twice as much H2) H2O SHOWED H and O diatomic (H2&O2) since two volumes of water produced New interpretation of chemical equation: coefficients inchemical equation give the volumesof gas BUT we need to know how to get numbers of molecules, and atoms by working with mass mass is easier to determine (measure) for high yield reactions (chemical co. pharmacy, etc)

13. How do masses and numbers of atoms relate? Each element has a averaged atomic weight based on the mixtures of isotopes (natural abundance) each weight a comparison to carbon-12: Study weight mass relationship for C-12 the number of atoms in 12.0 g of C-12 is 6.02x1023experimentally counted Avogadro’s number - 6.02x1023 SI unit for counting numbers Defines the mole - an amount of a substance that contains Avogadro’s number 6.02x1023 H2O molecules - 1 mole of water - 1 mole of O atoms - 2 moles of H atoms - 3 moles total of atoms All element weights based on a comparison to carbon-12! remember: mass of one mole C-12=12 g numerically equal to the atomic weight Weight of one mole of any element is equivalent to its atomic weight (in grams)

14. Define weights based on large amounts of atoms gram-atomic weight - the mass (in grams) of one mole of an element -numerically equal to atomic wt. (in grams) -represents 6.02x1023 atoms - atomic weight since a mixture of isotopes - one gram-atomic weight of any element contains the same number of atoms 6.02x1023 gram-formula weight - the mass (in grams) of one mole of the compound -represents 6.02x1023 formula units -numerically equal to formula wt. (in grams) - applies to ionic or covalent gram-molecular weight - the mass (in grams) of one mole of the molecules (covalent only) -gram-formula weight of a molecular compound -represents 6.02x1023 molecules -numerically equal to formula wt. (in grams) Can now go back and forth between number and mass like conversion factors

15. Conversions between mass and number Gram weights can be thought of as conversion factors: grams per mole Example: NaCl- formula weight = 58.44 amu gram-formula wt.= 58.44 grams 58.44 g for every mole conversion identity 58.44 g = 1 mole NaCl 24.7 g of NaCl is how many moles? 24.7 g =0.423 moles Example: 0.773 moles of C2H4 has what mass? gram formula wt.=2x12.01g + 4x1.008g = 28 g (per mole C2H4) Convert: 0.773 moles = 21.64 g of C2H4 How many molecules? 0.773 moles = 4.65x1023 C2H4 molecules How many C atoms: 2 C atoms per mole 0.773 moles =1.546 moles 1 mole 58.44 g 28 g 1 mole 6.02x1023 molecules 1 mole 2 C atoms 1 mole C2H4

16. Suggested problems for Chapter 10: • Parallel Exercises - # 1, 2, 3, 5 • Additional problems: • Convert the following to numbers of moles • 6.7 g of iron 6.7g = 0.12 moles Fe • 45.3 g of Fe2(SO4)3 45.3 g=0.113 moles Fe2(SO4)3 c) 8.4x1024molecules of C6H12O6 8.4x1024molecules = 13.95 moles d) 9.4x1023 Na+ ions of C6H12O6 2. Convert the following into masses • 0.55 moles of silver =59.34 g • 4.7x1022 molecules of CH4 =1.26 g • 2 moles of AgCl - Answer: 286.8 g • 5.6x1024 Cl- ions - Answer: 330.23 g 3. In two moles of BaF2: a) how many moles of F are there? Answer: four b) how many Ba+2 ions are there? Answer: two c) what is the mass of F? Answer: 76 g d) what is the mass percent of F? (Ch. 9) Answer: 21.7% 1 mole Fe 55.8 g 1 mole Fe2(SO4)3 399.9 g 1 mole C6H12O6 6.02x1023 molecules 107.9 g 1 mole silver 16.16 g in a mole 6.02x1023 molecules