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Chapter 8

Chapter 8 . Reactions in aqueous solutions. Predicting products with aqueous reactants . Driving Force for reactants to form certain products Most common: formation of a solid formation of water formation of gas transfer of electrons. Formation of a precipitate. Vocabulary:

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Chapter 8

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  1. Chapter 8 Reactions in aqueous solutions

  2. Predicting products with aqueous reactants • Driving Force for reactants to form certain products • Most common: formation of a solid • formation of water • formation of gas • transfer of electrons

  3. Formation of a precipitate Vocabulary: Precipitation – formation of a solid in chemical reaction Precipitate - solid formed from chemical reaction Precipitation reaction – reaction in which a solid is formed and separates from solution Ex: K2(CrO4 )(aq) + Ba(No3)2(aq) → K(NO3)(aq) + Ba (CrO4)(s) Notice that 2 aqueous reactants produced a solid and an aqueous product

  4. General rules for solubility • Most nitrates are soluble • Most salts of Na, K, and NH4 are soluble • Most chlorides are soluble • exceptions: AgCl, PbCl2, and Hg2Cl2 • Most sulfates are soluble • exceptions: BaSO4,PbSO4,CaSO4 • ------------------------------------------------------------------------------------------------------------------------------------------------------------ • Most OH compounds are only slightly soluble • exceptions: KOH, Ba(OH)2, Ca(OH)2 are moderately soluble • Most sulfides, carbonates and phosphates are only slightly soluble

  5. How do you know what products will form • You need to know the nature of the reactants • Let’s take Ba(NO3)2(aq) • -the (aq) means that Ba(No3)2 has been dissolved in water • -the formula indicates Ba +2 and NO3-1 ions • **(the ions separate (dissociate) when dissolved in water) • -the formula also indicates 2 NO3 ions form with every 1 Ba ion • -these ions are good electrolytes-conductors of electricity • Let’s take K2CrO4 • -the (aq) means that K2CrO4 also dissolves in water • -the formula also indicates dissociation • - these ions are also good electrolytes • What we have now: • Ba(NO3)2(aq) + K2CrO4 → Products (?)

  6. Predicting the Products • So what are the possibilities? • We know that it’s a double displacement reaction • We also know that K when placed in water will totally dissolve so a product containing K would not form a solid • therefore the products must include the Ba compound forming the precipitate • Knowing the solubility of the ions is important in determining which product will form the precipitate • -soluble solid = solid readily dissolves in water • -insoluble/slightly soluble = only a small amount dissolves in water

  7. Practice • Problem: When an aqueous solution of silver nitrate is added to an aqueous solution of potassium chloride, a white solid forms. Identify this precipitate an write the balance equation. • So what you have is: AgNO3 (aq) + KCl (aq) → white solid • This is a double displacement reaction so we know the products: • AgCl and KNO3 • The question now becomes which one of these is the precipitate • To figure out look at the solubility rules • Rule 1: Most Nitrates are soluble • Rule 2: Most salts containing K are soluble • Rule 3: Most chlorides are soluble with exceptions==AgCl • This means that AgCl MUST be the precipitate • AgNO3(aq) + KCl(aq) → AgCl(s) + KNO3(aq)

  8. Practice • Using the solubility rules predict the products and indicate the states of the products • a) KNO3 (aq) and BaCl2(aq) • b) Na2SO4(aq) and Pb(NO3)2(aq) • KOH(aq) and Fe(NO3)3(aq) • According to the solubility rules nitrates and chlorides are soluble in water SO the ions will stay dissolved in water AND no chemical reaction takes place • According to the solubility rules nitrates are soluble BUT sulfates are only slightly soluble SO Pb(So4) will form the precipitate • According to the solubility rules nitrates are soluble BUT hydroxides are only slightly soluble SO Fe(OH) will form the precipitate

  9. Practice 2 • Predict whether a solid will form when the following pairs of solution are mixed. If so, identify the solid and write a balanced equation for the reaction. • Ba(NO3)2(aq) and NaCl(aq) • Na2S(aq) and Cu(NO3)2(aq) • NH4Cl(aq) and Pb(NO3)2(aq)

  10. Practice 2 answers • No solid forms • Na2S(aq) + Cu((NO3)2(aq) → CuS(s) + 2Na(NO3)(aq) • 2NH4Cl(aq) + Pb(NO3)2(aq) → PbCl2(s) + 2NH4(NO3)(aq)

  11. Describing reactions in aqueous solution • 3 Types • Molecular Equation – shows overall reactions • may or may not show actual states of R and P • Ex: K2CrO4(aq) + Ba(NO3)2→BaCrO4(s) + 2KNO3(aq) • 2) Complete Ionic Equation - shows all R and P that are strong electrolytes • Ex: 2K +(aq) + CrO4(aq)- + Ba+(aq) +2NO3(aq)-→BaCRO4(s) + 2K+(aq) + 2NO3-(aq) • 3) Net Ionic Equation – shows all components that undergo a change • Spectator ions (ions that do not participate directly the reaction) are not included • Ex: Ba+2(aq) + CrO4-2(aq) →BaCrO4(s)

  12. Classwork/Homework • Page 253 1-6 • Page 271-272 1-16

  13. Acids and Bases • Bases: substance that produces OH- ions when dissolved in water • (proton acceptor) • tastes bitter • strong bases-good electrolytes • weak bases- poor electrolytes • litmus paper turns blue • phenolphthalein is pink • emulsifies fats and oils • Ex: NaOH, KOH, Ca(OH)2 • Acids: substance that produces H+ ions when dissolved in water (proton donor) • taste sour • good electrolytes • litmus paper turns red • phenolphthalein is colorless • reacts with metals to form H↑ and metal • corrosive • Ex: HCl, H2SO4,HNO3

  14. What occurs when acids and bases react • Salts and water is produced • Consider a Neutralization Reaction because of the stability of water • Individual properties of the acid and bases are lost to the salt • Ex: HCl(aq) + Na(OH)(aq) →H2O(l) +NaCl(aq) • Because HCl, NaOH and NaCl can completely dissolve in water then the • Complete ionic equation is: • H+(aq) +Cl-(aq) +Na+(aq) + OH-(aq) →H20(l) + Na+(aq) + Cl-(aq) • **** Spectator Ions include: Cl- and Na+ • Net Ionic equation is: H+(aq) + OH-(aq) → H2O(l)

  15. Practice • Nitric Acid id a strong acid. Write the molecular, complete ionic and net ionic equations for the reaction of aqueous acid and aqueous potassium hydroxide. • Molecular equation: • HNO3(aq) + KOH(aq) →H2O(l) + KNO3(aq) • Complete Ionic equation: • H+(aq) + NO3-(aq) + K+(aq) + OH-(aq) →H2O(l) + K+(aq) +NO3-(aq) • Net Ionic Equation: • H+(aq) + OH-(aq) →H2O(l) • NOTE: The products will always include H2O and a SALT that may be soluble or insoluble.

  16. Reactions of metals and nonmetals • Considered Oxidation-Reduction Reactions (Redox)= any chemical change in which the elements undergo a change in oxidation number. • This involves the transfer of electrons when forming compounds • Ex: Na + Cl→NaCl • Na metal has a charge of O; Cl nonmetal has a charge of O • However, when they combine they have charges • Na+1 and Cl-1 • 3)Because the Na loses an electron to the chlorine and gets a (+) charge it is said to be • “oxidized”. Na0→Na+ +1e- • 4)Because the Cl gains an electron from the Na and gains a (-) charge it is said to be • “reduced”. Cl0 + 1e-→Cl-

  17. Rules for assigning oxidation numbers • All uncombined elements have an oxidation number of zero. • A monoatomic ion has an oxidation number equal to its charge. • Fluorine has an oxidation number of -1 in all compounds. • Oxygen has an oxidation number of -2 in all compounds. • Exception: Peroxide H2O2 Oxygen Oxidation number is -1. • Hydrogen has an oxidation number of +1. • Exception: when in combination with metals it is -1 (NaH) • The sum of the oxidation number of all atoms in a neutral compound is • equal to zero.

  18. Practice • For each reaction, show how electrons are gained or lost. Label the reduced and oxidized element. • 2 Na(s) + Br2(l) →2NaBr(s) • 2Ca(s) + O2(g) →2CaO(s) • Solution: • Oxidized: Na → Na+ + 1e- Reduced: Br + 1e- →Br- • Oxidized: Ca→Ca+2 + 2e- Reduced: O + 2e- →O-2

  19. Classwork/Homework • Cw: Lab: Precipitation Reactions—Complete 1 & 2 • CW/HwReadPages 254-262 and • Complete section review : Page 262 1-6

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