Chapter 14 Acids and Bases
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Transcript of Chapter 14 Acids and Bases
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Chapter 14 Acids and Bases
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Chapter 14Section 1 – Properties of Acids and BasesSection 2 – Acid Base TheoriesSection 3 – Acid Base Reactions
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14.1 Properties of Acids and BasesList five general properties of aqueous acids
and bases.
Name common binary acids and oxyacids, given their chemical formulas.
List five acids commonly used in industry and the laboratory, and give two properties of each.
Define acid and base according to Arrhenius’s theory of ionization.
Explain the differences between strong and weak acids and bases.
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Properties of: Acids Bases
1. Sour taste2. Conducts electricity3. Turns litmus paper
red4. Reacts with bases
to produce salts and water
5. Reacts with some metals and releases hydrogen gas
1. Bitter taste2. Feels slippery3. Conducts
electric current4. Turns litmus
paper blue5. Reacts with
acids to produce salts and water
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Binary Acids
Contains only two different elementsHydrogen & an electronegative, nonmetal
Nomenclature:hydro - _________ - ic acid
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Diatomic Nomenclature
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OxyacidContains hydrogen, oxygen, and a third
element(hydrogen with a polyatomic ion)
Nomenclature:
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Acid Names
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Oxyacids
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Common Industrial AcidsSulfuric Acid
Sulfuric acid is the most commonly produced industrial chemical in the world.
Nitric Acid
Phosphoric AcidHydrochloric Acid
Conc. HCl is commonly referred to as muriatic acid.
Acetic AcidPure acetic acid is a clear, colorless, and
pungent-smelling liquid known as glacial acetic acid.
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Arrhenius Acids and BasesArrhenius Acids:
Increases concentration of H+ ions in solution
Arrhenius Bases:Increases concentration of OH- ions in
solution
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Arrhenius Acid Base Video
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Acid/Base StrengthStrong acid:
Ionizes completely in solution and is an electrolyteHigher the KA, the greater the strength as an acid
K reveals a greater extent of ionization
Example: HCl, HClO4, HNO3
Weak acid:Releases few hydrogen ions in solution
Hydronium ions, anions and dissolved acid molecules present
Examples: HCN, Organic acids – HC2H3O2
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Dissociation Constants
Strong vs. Weak BaseStrong bases ionizes completely in solution and
is a strong electrolyte
KB = dissociation constant of a baseHigher the KB , the greater the strength of a
base
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Aqueous Acids
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s aq + aq2H O –NaOH( ) Na ( ) OH ( )
aq + l aq + aq–3 2 4NH ( ) H O( ) NH ( ) OH ( )
Base StrengthStrong bases:
Ionic compounds containing metal cation and hydroxide ion (OH-)Dissociates in water
Weak bases:Molecular compounds do not follow Arrhenius
definition: Ammonia (NH3)Produces hydroxide ions when it reacts with water
molecules
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Base Strength
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Acidic solution has greater [H3O+] Basic solution has greater [OH–]
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14.2 Acid Base TheoriesDefine and recognize Brønsted-Lowry
acids and bases.
Define a Lewis acid and a Lewis base.
Name compounds that are acids under the Lewis definition but are not acids under the Brønsted-Lowry definition.
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+ –3 4HCl NH NH Cl
l + aq aq + aq–2 3 4H O( ) NH ( ) NH ( ) OH ( )
Bronsted-Lowry Acid
Bronsted-Lowry Acid:Proton (H+) donor
Hydrogen chloride acts as a Bronsted-Lowry acid when it reacts with ammonia.
Water can also act as a Bronsted-Lowry acid
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+ –3 4HCl NH NH Cl
Bronsted-Lowry Base
Bronsted-Lowry Base:Proton acceptor
Ammonia accepts a proton from hydrochloric acid.
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+ –3 4HCl NH NH Cl
acid base
Bronsted-Lowry Acid Base Reactions
Protons are transferred from one reactant (the acid) to another (the base)
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aq + l aq + aq–2 3HF( ) H O( ) F ( ) H O ( )
acid conjugate base
Conjugate Acid – BaseConjugate Base:
The species that remains after a Bronsted-Lowry acid has given up a proton
Conjugate Acid:The species that remains after a Bronsted-
Lowry base has accepted a proton
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aq + l aq + aq–2 3HF( ) H O( ) F ( ) H O ( )
Conjugate Acid Base PairsMatch up the acid-base pairs
(proton donor-acceptor pairs)
acid1 base2 conjugate
base1 conjugate
acid2
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g + l aq + aq–2 3HCl( ) H O( ) H O ( ) Cl ( )
strong acid base acid weak base
Strength of Acid Base PairsThe stronger the acid, the weaker the
conjugate baseThe stronger the base, the weaker the
conjugate acid
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aq + l aq + aq–4 2 3 4HClO ( ) H O( ) H O ( ) ClO ( )
aq + l aq + aq–3 2 3 3CH COOH( ) H O( ) H O ( ) CH COO ( )
stronger acid stronger base weaker acid weaker base
weaker acid weaker base stronger acid stronger base
Proton transfer favors the production of the weaker acid and base.
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Acid Base Strength
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aq + l aq + aq–2 4 2 3 4H SO ( ) H O( ) H O ( ) HSO ( )
g + l aq aq–3 2 4NH ( ) H O( ) NH ( ) OH ( )
acid1 base2 acid2 base1
base1 acid2 acid1 base2
AmphotericAny species that can react as either an acid
or a baseExample: water
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Amphoteric Water Video
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Other Amphoteric CompoundsCovalently bonded –OH group in an acid is
referred to as a hydroxyl groupMolecular compounds with hydroxyl groups
can be acidic or amphotericThe behavior of the compound is affected
by the number of oxygen atoms bonded to the atom connected to the –OH group
*The more oxygen’s in a polyatomic formula, the greater the strength of polyatomic as an acid
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Oxyacids of Chlorine
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Brønsted-Lowry Acid Base Video
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g + l) aq + aq–2 3HCl( ) H O( H O ( ) Cl ( )
Monoprotic AcidsCan donate only one proton (hydrogen ion)
per moleculeOne ionization step
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Monoprotic and Diprotic Acids
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l + l aq + aq–2 4 2 3 4H SO ( ) H O( ) H O ( ) HSO ( )
aq + l aq + l– 2–4 2 3 4HSO ( ) H O( ) H O ( ) SO ( )
Polyprotic AcidsDonates more than one proton per moleculesMultiple ionization steps
Diprotic – donates 2 protons Ex:Triprotic – donates 3 protons Ex:
Sulfuric acid solutions contain H3O+, HSO4-, SO4
- ions
1.
2.
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aq + aq aq aq3 3 4H ( ) : NH ( ) [H — NH ] ( ) or [NH ] ( )
Lewis AcidLewis acid:
Atom, ion, or molecule that ACCEPTS an ELECTRON PAIR to form a covalent bond
A proton (hydrogen ion) is a Lewis acid
Lewis base:Atom, ion, or molecule that DONATES an
ELECTRON PAIR to form a covalent bond
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aq + aq aq 3 3 3 3 2Ag ( ) 2 : NH ( ) [H N — Ag — NH ] ( ) or [Ag(NH ) ]
Lewis AcidA lewis acid might not include hydrogenSilver as a lewis acid:
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Lewis Acid Base Video
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Acid and Base Definitions
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Acid Base Definitions Video
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14.3 Acid Base ReactionsDescribe a conjugate acid, a conjugate
base, and an amphoteric compound.
Explain the process of neutralization.
Define acid rain, give examples of compounds that can cause acid rain, and describe effects of acid rain.
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2aq + aq aq lHCl( ) NaOH( ) NaCl( ) H O( )
Neutralization ReactionsWhat does it mean to neutralize
something?
Neutralization reactions: Hydronium and hydroxide ions react to form
water The left over cation and anion in solution
produce a salt (ionic compound)
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Neutralization Reactions
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Neutralization Reaction Video
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g + l aq3 2 2 4SO ( ) H O( ) H SO ( )
Acid RainNO, NO2, CO2, SO2, and SO3 gases from
industrial processes can dissolve in atmospheric water to produce acidic solutions.
Very acidic rain is known as acid rain.Acid rain can erode statues and affect ecosystems.
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Chapter 15 Acid Base Titration and pH
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Chapter 15Section 1 – Aqueous Solutions and the
Concept of pHSection 2 – Determining pH and Titrations
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15.1 Aqueous Solutions and pHDescribe the self-ionization of water.
Define pH, and give the pH of a neutral solution at 25°C.
Explain and use the pH scale.
Given [H3O+] or [OH−], find pH.
Given pH, find [H3O+] or [OH−].
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l + l aq + aq–2 2 3H O( ) H O( ) H O ( ) OH ( )
Self Ionization of WaterTwo water molecules produce a hydronium ion
and hydroxide ion by proton transfer
In water at 25°C, [H3O+] = 1.0 ×10−7 M and [OH−] = 1.0 × 10−7 M
The ionization constant of water, Kw
Kw = [H3O+][OH−]
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At 25OC
Kw = [H3O+][OH−] = (1.0 × 10−7)(1.0 × 10−7) = 1.0 × 10−14
Kw = 1.0 x 10-14
Kw increases as temperature increases
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Ion Concentration[H3O+] = [OH−]
neutral
[H3O+] > [OH−]
acidic
[H3O+] > 1.0 × 10−7 M
[OH−] > [H3O+]
basic
[OH−] > 1.0 × 10−7 M
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s aq + aq2H O –NaOH( ) Na ( ) OH ( )
-14 -14
-123 – -2
1.0 10 1.0 10[H O ] 1.0 10 M
[OH ] 1.0 10
Calculating ConcentrationStrong acids and bases are considered
completely ionized or dissociated in aqueous solutions.
1 mol 1 mol 1 mol1.0 × 10−2 M NaOH therefore, [OH−] = 1.0
× 10−2 M[H3O+] is calculated using Kw
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-14 -14
– -10-4
3
1.0 10 1.0 10[OH ] 5.0 10 M
[H O ] 2.0 10
Example Problem 1Given: [HCl] = 2.0 × 10−4 M
[H3O+] = ______________Unknown: [OH-] = ?
Kw = [H3O+][OH−] = 1.0 × 10−14
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pHDefinition of the pH of a solution: negative of the
common logarithm of the hydronium ion concentration, [H3O+].
pH = −log [H3O+]
Example: a neutral solution has a [H3O+] = 1×10−7
pH = −log [H3O+] = −log(1 × 10−7) = −(−7.0) = 7.0
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pH Values as Specified [H3O+]
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The pH Scale
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pOHThe pOH of a solution is defined as the negative of
the common logarithm of the hydroxide ion concentration, [OH−].
pOH = −log [OH–]
pH + pOH = 14.0
Example: a neutral solution has a [OH–] = 1×10−7
the pH of this solution is?
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Calculating [H3O+] from pH
Finding the [H3O+] from pH requires taking the antilog of the negative pH
[H3O+] = antilog (-pH)
You can find the [OH−] by also taking the antilog of the negative pOH.
[OH-] = antilog (-pOH)
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The Circle of pH
pH
pOH
[ H3O+]
[ OH-]
-log [H3O+]
antilog (-pH)
antilog (-pOH)
-log [OH-]
[ H3O+] [ OH-] = 1.0x10-14pH + pOH = 14
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pOH Video
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pH Values of Some Common Materials
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Approximate pH Range of Common Materials
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Comparing pH and pOH Video
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pH of Weak Acids and BasesThe pH of solutions of weak acids and
weak bases must be measured experimentally.
The [H3O+] and [OH−] can then be calculated from the measured pH values.
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Significant FiguresThere must be as many significant figures to
the right of the decimal as there are in the number whose logarithm was found.
Example: [H3O+] = 1 × 10−7
one significant figure
pH = 7.0
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15.2 Determining pH and TitrationsDescribe how an acid-base indicator
functions.
Explain how to carry out an acid-base titration.
Calculate the molarity of a solution from titration data.
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IndicatorsAcid-base indicators: compounds whose
colors are sensitive to pH.
The pH range over which an indicator changes color is called its transition interval.
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pH MeterspH meter determines the pH of a
solution by measuring the voltage between the two electrodes that are placed in the solution.
The voltage changes as the hydronium ion concentration in the solution changes.
Measures pH more precisely than indicators
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Color Ranges of Indicators
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Color Ranges of Indicators
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Color Ranges of Indicators
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Antacids Video with Methyl Orange
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TitrationNeutralization occurs when hydronium
ions and hydroxide ions are supplied in equal numbers by reactants.
H3O+(aq) + OH−(aq) 2H2O(l)Titration: the controlled addition and
measurement of the amount of a solution of known concentration required to react completely with a measured amount of a solution of unknown concentration.
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Titration Pointsequivalence point: point at which the
two solutions used in a titration are present in chemically equivalent amounts
end point: point in a titration at which an indicator changes color
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Which indicator do I choose?pH less than 7
Indicators that change color at pH lower than 7 are used to determine the equivalence point of strong-acid/weak-base titrations.
strong-acid/weak-base titration = acidic.pH at 7
Indicators that undergo transition at about pH 7 are used to determine the equivalence point of strong-acid/strong base titrations.
strong acids/strong bases = salt solution with a pH of 7.
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Which indicator do I choose?pH greater than 7
Indicators that change color at pH higher than 7 are used to determine the equivalence point of weak-acid/strong-base titrations.
weak-acid/strong-base = basic
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Titration Curve Strong Acid and a Strong BaseEquivalence Point:
pH at 7
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Titration Curve Weak Acid and a Strong BaseEquivalence Point:
pH higher than 7
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Titration Curve Strong Acid and a Weak BaseEquivalence Point:
pH less than 7
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Titration Problems:* Can be used to determine concentration of
unknown solution or volume of added standard
1. Start with the balanced equation for the neutralization reaction
2. Make amount of acid and base chemically equivalent to each other (1 to 1 mol ratio).
3. Determine the molarity of the unknown solution.
Equation: M1V1 = M2V2
1: starting solution2: added standard
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Molarity and Titrationstandard solution: solution that contains the
precisely known concentration of a solute
primary standard: highly purified solid compound used to check the concentration of the known solution
The standard solution can be used to determine the molarity of another solution by titration.
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Performing a Titration – Set up
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Performing a Titration – Set up Acid
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Performing a Titration – Starting Amount
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Performing a Titration – Set up Base
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Performing a Titration - Titrating
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Performing a Titration – End Point
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1 mol 1 mol 1 mol 1 mol
Molarity and TitrationDetermine the molarity of an acidic solution, 10
mL HCl, by titration
1. Titrate acid with a standard base solution20.00 mL of 5.0 × 10−3 M NaOH was titrated
2. Write the balanced neutralization reaction equation.
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
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-3-45.0 10 mol NaOH 1 L
20 mL 1.0 10 mol NaOH used1 L 1000 mL
-4-21.0 10 mol HCl 1000 mL
1.0 10 M HCl10.0 mL 1 L
Molarity and Titration4. Calculate the number of moles of NaOH used
in the titration. 20.0 mL of 5.0 × 10−3 M NaOH is needed to reach
the end point
5. mol of HCl = mol NaOH = 1.0 × 10−4 mol
6. Calculate the molarity of the HCl solution
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Example ProblemIn a titration, 27.4 mL of 0.0154 M
Ba(OH)2 is added to a 20.0 mL sample of HCl solution of unknown concentration until the equivalence point is reached. What is the molarity of the acid solution?
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Ba(OH)2 + 2HCl BaCl2 + 2H2O1 mol 2 mol 1 mol 2
mol
Example Problem SolutionGiven: 27.4 mL of 0.0154 M Ba(OH)2
Unknown: ? M HCl of 20.0 mLSolution:
Write balanced equation:
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2
2 2
mol Ba(OH) 1 LmL of Ba(OH) solution mol Ba(OH)
1 L 1000 mL
22
-42
0.0154 mol Ba(OH)24.7 mL of Ba(OH) solution
1 L1 L
4.22 10 mol Ba(OH)1000 mL
1. Calculate Moles of Given
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–42
2
–4
2 mol HCl4.22 10 mol of Ba(OH)
1 mol Ba(OH)
8.44 10 mol HCl
2. Write a mole ratio: moles of base used to moles of acid produced
22
2 mol HClmol of Ba(OH) in known solution mol HCl
mol Ba(OH)
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amount of solute in unknown solution (mol) 1000 mL
volume of unknown solution (mL) 1 L
molarity of unknown solution
3. Calculate Unknown Molarity
-2
-48.44 10 mol HCl 1000 mL
20.0 m4.22 10
L 1M l
LHC