Determining Limiting Reagents Guided Practice Problem - PowerPoint PPT Presentation

About This Presentation
Title:

Determining Limiting Reagents Guided Practice Problem

Description:

Determining Limiting Reagents Guided Practice Problem Part of the SO2 that is introduced into the atmosphere ends up being converted to sulfuric acid, H2SO4. – PowerPoint PPT presentation

Number of Views:123
Avg rating:3.0/5.0
Slides: 38
Provided by: Simo2187
Category:

less

Transcript and Presenter's Notes

Title: Determining Limiting Reagents Guided Practice Problem


1
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?

2
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?

3
Chapter 4Aqueous Reactions and Solution
Stoichiometry
CHEMISTRY The Central Science 9th Edition
4
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.

5
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.

6
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.

7
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?

8
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

9
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
10
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.

11
General Properties of Aqueous Solutions
Electrolytic Properties
12
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.

13
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)

14
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

15
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.

16
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.

17
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.

18
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

19
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)

20
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

21
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.

22
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.

23
Acid-Base Reactions
Strong and Weak Electrolyte Summary
24
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.

25
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)

26
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.

27
Oxidation-Reduction Reactions
Oxidation and Reduction
28
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.

29
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).

30
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.

31
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.

32
(No Transcript)
33
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.

34
Solution Stoichiometry and Chemical Analysis
35
Solution Stoichiometry and Chemical Analysis
  • Titrations

36
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.

37
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.
Write a Comment
User Comments (0)
About PowerShow.com