Determining Limiting Reagents Guided Practice Problem

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Determining Limiting Reagents Guided Practice
Problem

• Part of the SO2 that is introduced into the
  atmosphere ends up being converted to sulfuric
  acid, H2SO4. The net reaction is
• 2SO2(g) O2(g) 2H2O(l) ? 2H2SO4(aq)
• How much H2SO4 can be formed from 5.0 mol of SO2,
  1.0 mol O2, and an unlimited quantity of H2O?

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Determining Limiting Reagents Guided Practice
Problem

• Consider the following reaction
• 2Na3PO4(aq) 3Ba(NO3)2(aq) ? Ba3(PO4)2(s)
  6NaNO3(aq)
• Suppose that a solution containing 3.50 g of
  Na3PO4 is mixed with a solution containing 6.40 g
  of Ba(NO3)2. How many grams of Ba3(PO4)2 can be
  formed? What is the yield, if experimentally,
  only 4.70 g were obtained from the reaction?

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Chapter 4Aqueous Reactions and Solution
Stoichiometry
CHEMISTRY The Central Science 9th Edition
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Solution Composition

• Solutions are homogenous mixtures of two or more
  substances
• Solute present in smallest amount and is the
  substance dissolved in the solvent.
• Solvent present in the greater quantities and is
  used to dissolve the solute.
• Example NaCl dissolved in Water (water Solvent
  and NaCl Solute)
• Change concentration by using different amounts
  of solute and solvent.
• Molarity Moles of solute per liter of solution.

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Concentrations of Solutions

• Formula for Molarity
• The most widely used way of quantifying
  concentration of solutions in chemistry.
  Molarity is generally represented by the symbol M
  and defined as the number of moles of solute
  dissolved in a liter of solution.

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Class Guided Practice Problem

• Calculate the molarity of a solution made by
  dissolving 23.4 g of sodium sulfate, Na2SO4, in
  enough water to form 125 mL of solution.

Class Practice Problem

• Calculate the molarity of a solution made by
  dissolving 5.00 g of NaCl in sufficient water to
  form 0.125 L of solution.

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Class Guided Practice Problem
Determining Mass using Molarity

• Key concept If we know molarity and liters of
  solution, we can calculate moles (and mass) of
  solute.

• How many grams of C6H12O6 are required to make
  100 mL of 0.278 M C6H12O6?

Class Practice Problem

• How many grams of NaCl are required to make a 1 L
  of 0.500 M NaCl?

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Concentration of Diluted Solutions

• Solutions are routinely prepared in stock
  solutions form.
• Example 12 M HCl
• Solution of lower concentrations are prepared by
  adding more solvent (e.g., water), a process
  called dilution.
• We recognize that the number of moles are the
  same in dilute and concentrated solutions.
• Hence, moles solute before dilution moles
  solute after dilution
• So
• MdiluteVdilute MconcentratedVconcentrated
• or
• Mfinal Vfinal MinitialVinitial

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Class Guided Practice Problem

• How much 3.0 M H2SO4 would be required to make
  500 mL of 0.10 M H2SO4?
• How many milliliters of 5.0 M K2Cr2O7 solution
  must be diluted in order to prepare 250 mL of
  0.10 M solution?

Class Practice Problem
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General Properties of Aqueous Solutions

• Electrolytic Properties
• Aqueous solutions, solutions in water, have the
  potential to conduct electricity.
• The ability of the solution to conduct depends on
  the number of ions in solution.
• There are three types of solution
• Strong electrolytes,
• Weak electrolytes, and
• Nonelectrolytes.

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General Properties of Aqueous Solutions
Electrolytic Properties
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General Properties of Aqueous Solutions

• Molecular Compounds in Water
• Molecular compounds in water (e.g., CH3OH) no
  ions are formed.
• If there are no ions in solution, there is
  nothing to transport electric charge.

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General Properties of Aqueous Solutions

• Ionic Compounds in Water
• Ions dissociate in water (NaCl).
• In solution, each ion is surrounded by water
  molecules.
• Transport of ions through solution causes flow of
  current.
• Other substances that are not ionic compound
  dissociate in water to form ions.
• For example (HCl)

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General Properties of Aqueous Solutions

• Strong and Weak Electrolytes
• Strong electrolytes completely dissociate in
  solution.
• For example
• Weak electrolytes produce a small concentration
  of ions when they dissolve.
• These ions exist in equilibrium with the
  unionized substance.
• For example

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General Properties of Aqueous Solutions

• Some General Terms
• Acids - substances that able to ionize in
  solution to form hydrogen ion (H) and increase
  the concentration of H in the solution.
• For example, HCl dissociate in water to form H
  and Cl- ions.
• Bases - are substances that can react with or
  accept H ions.
• For example, OH- will accept H from HCl forming
  H2O.
• Salts - are ionic compounds that can be formed by
  replacing one or more of the hydrogen ions of an
  acid by a different positive ion.
• For example, NaCl instead of HCl.

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General Properties of Aqueous Solutions

• Identifying Strong and Weak Electrolytes
• Most salts are strong electrolytes (NaCl, CaCO3.
• Most acids are weak electrolytes. However, HCl,
  HBr, HI, HNO3, H2SO4, HClO3, and HClO4 are strong
  acids.
• The common strong bases are the hydroxides,
  Ca(OH)2, of the alkali metals and the heavy
  alkaline earth metals.
• Most other substances are nonelectrolytes.

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Precipitation Reactions

• Exchange (Metathesis) Reactions
• When two solutions are mixed and a solid is
  formed, the solid is called a precipitate.
• Metathesis reactions involve swapping ions in
  solution
• AX BY ? AY BX.
• HCl NaOH ? NaCl H2O
• Metathesis reactions will lead to a change in
  solution if one of three things occurs
• an insoluble solid is formed (precipitate),
• formation of either a soluble weak or
  nonelectrolytes,
• an insoluble gas is formed.

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Class Practice Problem

• Write a balanced equation for the reaction
  between phosphoric acid, H3PO4, and potassium
  hydroxide, KOH.

• H3PO4 3KOH ? 3H2O K3PO4
• AX BY ? AY BX

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Precipitation Reactions

• Writing Reaction Equations
• Ionic equation used to highlight reaction
  between ions.
• Molecular equation all species listed as
  molecules
• HCl(aq) NaOH(aq) ? H2O(l) NaCl(aq)
• Complete ionic equation lists all ions
• H(aq) Cl-(aq) Na(aq) OH-(aq) ? H2O(l)
  Na(aq) Cl-(aq)
• Net ionic equation lists only unique ions
• H(aq) OH-(aq) ? H2O(l)

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Class Guided Practice Problem

• Write the net ionic equation for the reactions
  that occur when solutions of KOH and Co(NO3)2 are
  mixed.

• 2OH- Co2 ? Co(OH)2

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Acid-Base Reactions

• Acids with one acidic proton are called
  monoprotic (e.g., HCl).
• Acids with two acidic protons are called diprotic
  (e.g., H2SO4).
• Acids with many acidic protons are called
  polyprotic.

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Acid-Base Reactions

• Identifying Strong and Weak Electrolytes
• Water soluble and ionic strong electrolyte
  (probably).
• Water soluble and not ionic, but is a strong acid
  (or base) strong electrolyte.
• Water soluble and not ionic, and is a weak acid
  or base weak electrolyte.
• Otherwise, the compound is probably a
  nonelectrolyte.

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Acid-Base Reactions
Strong and Weak Electrolyte Summary
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Acid-Base Reactions

• Neutralization Reactions and Salts
• Neutralization occurs when a solution of an acid
  and a base are mixed
• HCl(aq) NaOH(aq) ? H2O(l) NaCl(aq)
• Notice we form a salt (NaCl) and water.
• Salt ionic compound whose cation comes from a
  base and anion from an acid.
• Neutralization between acid and metal hydroxide
  produces water and a salt.

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Acid-Base Reactions

• Acid-Base Reactions with Gas Formation
• Sulfide and carbonate ions can react with H in a
  similar way to OH-.
• 2HCl(aq) Na2S(aq) ? H2S(g) 2NaCl(aq)
• 2H(aq) S2-(aq) ? H2S(g)
• HCl(aq) NaHCO3(aq) ? NaCl(aq) H2O(l) CO2(g)

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Oxidation-Reduction Reactions

• Oxidation and Reduction
• When a metal undergoes corrosion it loses
  electrons to form cations
• Ca(s) 2H(aq) ? Ca2(aq) H2(g)
• Oxidized atom, molecule, or ion becomes more
  positively charged.
• Oxidation is the loss of electrons.
• Reduced atom, molecule, or ion becomes less
  positively charged.
• Reduction is the gain of electrons.

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Oxidation-Reduction Reactions
Oxidation and Reduction
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Oxidation-Reduction Reactions

• Oxidation Numbers
• Oxidation number for an ion the charge on the
  ion.
• Oxidation number for an atom the hypothetical
  charge that atom would have if it was an ion.
• Oxidation numbers are assigned by a series of
  rules
• If the atom is in its elemental form, the
  oxidation number is zero. E.g., Cl2, H2, P4.
• For a monoatomic ion, the charge on the ion is
  the oxidation state.

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Oxidation-Reduction Reactions

• Oxidation Numbers
• Nonmetal usually have negative oxidation numbers
• Oxidation number of O is usually 2. The
  peroxide ion, O22-, has oxygen with an oxidation
  number of 1.
• Oxidation number of H is 1 when bonded to
  nonmetals and 1 when bonded to metals.
• The oxidation number of F is 1.
• The sum of the oxidation numbers for the atom is
  the charge on the molecule (zero for a neutral
  molecule).

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Oxidation-Reduction Reactions

• Oxidation of Metals by Acids and Salts
• Metals are oxidized by acids to form salts
• Mg(s) 2HCl(aq) ? MgCl2(aq) H2(g)
• During the reaction, 2H(aq) is reduced to H2(g).
• Metals can also be oxidized by other salts
• Fe(s) Ni2(aq) ? Fe2(aq) Ni(s)
• Notice that the Fe is oxidized to Fe2 and the
  Ni2 is reduced to Ni.

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Oxidation-Reduction Reactions

• Activity Series
• Some metals are easily oxidized whereas others
  are not.
• Activity series a list of metals arranged in
  decreasing ease of oxidation.
• The higher the metal on the activity series, the
  more active that metal.
• Any metal can be oxidized by the ions of elements
  below it.

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Solution Stoichiometry and Chemical Analysis

• There are two different types of units
• laboratory units (macroscopic units measure in
  lab)
• chemical units (microscopic units relate to
  moles).
• Always convert the laboratory units into chemical
  units first.
• Grams are converted to moles using molar mass.
• Volume or molarity are converted into moles using
  M mol/L.
• Use the stoichiometric coefficients to move
  between reactants and product.

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Solution Stoichiometry and Chemical Analysis
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Solution Stoichiometry and Chemical Analysis

• Titrations

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Solution Stoichiometry and Chemical Analysis

• Titrations
• Suppose we know the molarity of a NaOH solution
  and we want to find the molarity of an HCl
  solution.
• We know
• molarity of NaOH, volume of HCl.
• What do we want?
• Molarity of HCl.
• What do we do?
• Take a known volume of the HCl solution, measure
  the mL of NaOH required to react completely with
  the HCl.

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Solution Stoichiometry and Chemical Analysis

• Titrations
• What do we get?
• Volume of NaOH. We know molarity of the NaOH,
  we can calculate moles of NaOH.
• Next step?
• We also know HCl NaOH ? NaCl H2O. Therefore,
  we know moles of HCl.
• Can we finish?
• Knowing mol(HCl) and volume of HCl (20.0 mL
  above), we can calculate the molarity.