chemistry - silberberg 8e ch.18 - acid-base...
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CHEMISTRY - SILBERBERG 8E
CH.18 - ACID-BASE EQUILIBRIA
CONCEPT: ACID IDENTIFICATION
The most common feature of an acid is that many possess an H+ ion called the _______________________________ .
When it comes to acids there are 2 MAJOR TYPES that exist:
_______________________ are acids where the H+ ion is attached to an electronegative element.
• These types of acids lack the element __________________ and usually possess no __________________ .
• The most common type of these particular acids are the haloacids: _______ , _______ , _______ & _______ .
_______________________ are acids that contain the ________________ , ________________ & ________________.
• They are created by the hydration of nonmetal oxides.
PRACTICE: Which of the following compound(s) cannot be classified as an acid?
a) H2S b) HCN c) H2 d) C6H6 e) All are acids.
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CONCEPT: BINARY ACID STRENGTH
STRONG ACIDS are considered _________________ Electrolytes so they ionize completely in water.
HCl (aq) H2O
H+ (aq) + Cl – (aq)
WEAK ACIDS are considered __________________ Electrolytes so they don’t completely ionize in water.
HF + H2O F – (aq) + H3O+ (aq)
The strength of a BINARY ACID is based on the _________________________ or ________________ of the nonmetal.
• For elements in the same period then look at their __________________ . The ________, the ________ acidic.
• For elements in the same group then look at their __________________ . The ________, the ________ acidic.
BINARY ACID STRENGTH
PRACTICE 1: Which is the weakest acid from the following?
a) H2S b) H2Se c) H2Te d) All would have the same acid strength.
PRACTICE 2: Which of the following acids would be classified as the strongest?
a) CH4 b) NH3 c) H2O d) HF e) PH3
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CONCEPT: OXYACID STRENGTH
The strength of OXYACIDS is based on the number of _____________ or the _____________________ of the nonmetal.
• RULE: If my oxyacid has 2 or More ___________ than ____________ then my oxyacid is a __________ ACID.
HNO3 ___ Oxygens – ___ Hydrogens
C6H5OH ___ Oxygens – ___ Hydrogens
HBrO4 ___ Oxygens – ___ Hydrogens
When comparing the strengths of different oxyacids remember:
• If they have different number of oxygens then the _________ oxygen the ___________ acidic
• If they have the same number of oxygens then the _________ electronegative the nonmetal the ________ acidic.
Electronegativity
H2C2O4 ___ Oxygens – ___ Hydrogens
HSO4 –
___ Oxygens – ___ Hydrogens
3 Exceptions
HIO3 ___ Oxygens – ___ Hydrogens
PRACTICE: Rank the following oxyacids in terms of increasing acidity.
a) HClO3 b) HBrO4 c) HBrO3 d) HClO4
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CONCEPT: BASE STRENGTHS
STRONG BASES are considered _________________ Electrolytes so they ionize completely in water.
NaOH (aq) H2O
Na+ (aq) + OH – (aq)
WEAK BASES are considered __________________ Electrolytes so they don’t completely ionize in water.
NH3 + H2O NH4+ (aq) + OH – (aq)
Bases possess THREE major features: __________________ or __________________ or __________________ .
Group ________:
• Any Group ______ metal when combined with OH –, H –, O2– or NH2 – makes a STRONG BASE.
Group ________:
• Any Group ______ metal, from _____ to _____ , when combined with OH –, H –, O2– or NH2 – makes a STRONG
BASE.
_____________:
• ____________________________________ are considered WEAK BASES.
Ex:
• ____________________________________ are considered WEAK ACIDS.
Ex:
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PRACTICE: ACID & BASE IDENTIFICATION
EXAMPLE: Classify each of the following as a strong acid, weak acid, strong base or weak base.
a) HCHO2 c) H2NNH2
b) (CH3CH2)3NH+ d) HBrO3
PRACTICE 1: Classify each of the following as a strong acid, weak acid, strong base or weak base.
a) KOCH3 b) CH3OH
PRACTICE 2: Classify each of the following as a strong acid, weak acid, strong base or weak base.
a) HOCN b) H5IO6
PRACTICE 3: Classify each of the following as a strong acid, weak acid, strong base or weak base.
a) NaN3 b) SrH2
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CONCEPT: ARRHENIUS ACIDS & BASES
The most general definition for acids and bases was developed by Svante Arrhenius near the end of the 19th century.
• According to him, the _______ cation and the _________ anion are fundamental to the concept of acids and bases.
• His definition however failed to describe acidic and basic behavior in nonaqueous media.
The Arrhenius definition states an acid is a compound that increases _______________ when dissolved in a solvent.
The Arrhenius definition states a base is a compound that increases _______________ when dissolved in a solvent.
PRACTICE 1: Which ions are formed from the dissociation of the following compound?
a) Sr(OH)2 (s) Dissolves in H2O
PRACTICE 2: Which ions are formed from the dissociation of the following compound?
a) H2SO4 (l) Dissolves in H2O
PRACTICE 3: Which ions are formed from the dissociation of the following compound?
a) HBO32-
Dissolves in H2O
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CONCEPT: BRONSTED LOWRY ACIDS & BASES
In 1923, Johannes Brønsted and Thomas Lowry developed a new set of definitions for acids and bases.
According to the Bronsted-Lowry definition, acids are considered _____________________________ and bases are
considered _____________________________.
• Unlike Arrhenius acids and bases, they are not limited to aqueous solutions.
• Every Arrhenius acid is a Brønsted-Lowry acid (and likewise for the bases).
• Brønsted-Lowry acids and bases always occur in pairs called _____________________________________ .
EXAMPLE 1: Write the formula of the conjugate base for the following compound:
HSO4 –
EXAMPLE 2: Write the formula of the conjugate acid for the following compound:
V2O52-
PRACTICE 1: Write the formula of the conjugate base for the following compound:
H2Se
PRACTICE 2: Write the formula of the conjugate for the following compound:
NH2NH2
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PRACTICE: BRONSTED LOWRY ACIDS & BASES (CALCULATIONS)
EXAMPLE 1: Identify the acid, base, conjugate acid and conjugate base in the following reactions:
a) HF (aq) + H2O (aq) F – (aq) + H3O+ (aq)
EXAMPLE 2: Identify the acid, base, conjugate acid and conjugate base in the following reactions:
a) CN – (aq) + H2O (aq) HCN (aq) + OH – (aq)
PRACTICE 1: Which of the following is a Bronsted-Lowry acid? a) CH4 b) HCN c) NH3 d) Br2
PRACTICE 2: Determine the chemical equation that would result when carbonate, CO32-, reacts with water.
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CONCEPT: AMPHOTERIC SPECIES
An amphoteric, or _________________________, is a species that can act as a(n) ACID or BASE.
• Water is prime example of an amphoteric species.
Partially dissociated conjugate bases of polyprotic acids are also amphoteric.
• These compounds possess _________________ and a __________________________.
Ex:
PRACTICE: Which of the following species is/are amphoteric?
a) CO32– b) HF c) NH4
+ d) HPO32- e) CH3O –
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CONCEPT: LEWIS…THE FINAL TYPE OF ACID & BASE
In the 1920s, Gilbert Lewis proposed a new set of definitions for acids and bases.
A Lewis acid is a(n) ______________________________________.
• ________ acts as a Lewis acid when connected to an electronegative element: ___ , ___ , ___ , ___ , or ________
• _____________________ charged hydrogen or metals.
• If your central element has _________________ 8 valence electrons.
A Lewis base is a(n) ______________________________________.
• Compounds with _________________________ .
NH3 H2O CH3OH CH3OCH3
• Compounds with a _________________________ .
CN – OH – CH3O – N3–
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PRACTICE: LEWIS….THE FINAL TYPE OF ACID & BASE (CALCULATIONS)
EXAMPLE: Identify each of the compounds in the following chemical equation.
H3CH2C O
CH2H3C
Al
Br
Br Br
H3CH2C O
CH2H3C
Al
Br
Br
Br
PRACTICE 1: Identify the Lewis acids and bases in the following reactions.
a) H+ + OH – H2O
b) Cl – + BCl3 BCl4–
c) SO3 + H2O H2SO4
PRACTICE 2: Identify each of the following compounds as either a Lewis acid, a Lewis base or neither.
a) ZnCl2 b) CN –
c) NH4+ d) Co3+
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CONCEPT: pH and pOH
To deal with incredibly small concentration values of [H+] and [OH-] we can use the pH scale.
• Under normal conditions, the pH scale operates within the range of ______ to ______ .
By taking the – log of [H+] and [OH-] we can find pH and pOH.
pH = − log[H+ ] pOH = − log[OH− ] p = − log
By recognizing the relationship between [H+] and [OH-] with pH and pOH we can create new formula relationships.
pH = − log[H+ ] pOH = − log[OH− ]
A species with a pH greater than 7 is classified as _____________ and the [H+] is ___________________ than the [OH-].
• The ______________ the base then the ______________ the pH.
A species with a pH less than 7 is classified as _______________ and the [H+] is ____________________ than the [OH-].
• The ______________ the acid then the ______________ the pH.
A species with a pH equal to 7 is classified as ______________ and the [H+] is _____________________ than the [OH-].
By using – log with the equilibrium expression for water a relationship between pH and pOH can be created.
pH+ pOH =14
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PRACTICE: pH and pOH (CALCULATIONS 1)
EXAMPLE: What is the hydroxide ion and hydronium ion concentration of an aqueous solution that has a pH equal to 6.12?
PRACTICE 1: Which of the following solutions will have the lowest concentration of hydronium ions?
a) 0.100 moles C6H5NH2
b) 0.100 moles Be(OH)2
c) 0.100 moles SrH2
d) 0.100 moles (CH3)2NH
PRACTICE 2: Which of the following statements about aqueous solutions is/are true?
a) For an basic solution the concentration of H3O+ is greater than the concentration of OH –.
b) The pH of a neutral aqueous solution is 7.00 at all temperatures.
c) An acidic solution under normal conditions has a pH value less than 7.00.
d) If the concentration of H3O+ decreases then the concentration of OH – will also decrease.
e) The pH of aqueous solutions is less than 7.
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PRACTICE: pH and pOH (CALCULATIONS 2)
EXAMPLE: A solution is prepared by dissolving 0.235 mol Sr(OH)2 in water to produce a solution with a volume of 750 mL.
a) What is the [OH-]?
b) What is the [H+]?
PRACTICE: What is the Kw of pure water at 20.0°C, if the pH is 7.083?
a) 8.26 × 10-8 b) 6.82 × 10-15 c) 7.23 × 10-14 d) 1.00 × 10-14
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CONCEPT: AUTO IONIZATION OF WATER
Water can react with itself in a reaction called self–ionization where ______________ and ______________ are produced.
H2O (l) + H2O (l)
This reaction is usually written more simply as:
H2O (l) + H2O (l)
The equilibrium equation for water is called the ________________________ (KW) for water and is given by the following:
KW = [H+ ][OH− ]
At 25oC, KW = ___________________, but remember KW, like all other constants K, is temperature dependent.
• Increasing the temperature will ______________ KW.
Constant
0oC
10 oC
50 oC
100 oC
KW
1.14 x 10-14
2.93 x 10-14
5.476 x 10-14
5.13 x 10-13
EXAMPLE: Determine the concentration of hydronium ions for a neutral solution at 25oC and at 50oC.
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CONCEPT: CALCULATING pH and pOH OF STRONG SPECIES
STRONG ACIDS & BASES are considered _________________ Electrolytes so they ionize completely in water.
HCl (aq) H2O
H+ (aq) + Cl – (aq)
NaOH (aq) H2O
Na+ (aq) + OH – (aq)
EXAMPLE 1: Calculate the pH of a 0.0782 M solution of CaH2.
EXAMPLE 2: Calculate the pH of a 0.000550 M HBr solution to the correct number of significant figures.
a) 3.3
b) 3.26
c) 3.260
d) 3.2596
e) All are correct
PRACTICE: Calculate the pH of 50.00 mL of 4.3 x 10-7 M H2SO4.
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CONCEPT: CALCULATING pH and pOH OF WEAK SPECIES
WEAK ACIDS & BASES are considered __________________ Electrolytes so they don’t completely ionize in water.
HF + H2O F – (aq) + H3O+ (aq)
NH3 + H2O NH4+ (aq) + OH – (aq)
EXAMPLE 1: Pryridine, an organic molecule, is a very common weak base.
C5H5N (aq) + H2O (l) C5H5NH+ (aq) + OH- (g)
Assume you have a 0.0225 M aqueous solution of pyridine, C5H5N, determine its pH. The Kb value for the
compound is 1.5 x 10-9.
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PRACTICE: CALCULATING pH and pOH OF WEAK SPECIES (CALCULATIONS 1)
EXAMPLE: An unknown weak base has an initial concentration of 0.750 M with a pH of 8.03. Calculate its equilibrium base constant.
PRACTICE: Determine the pH of a solution made by dissolving 6.1 g of sodium cyanide, NaCN, in enough water to make a
500.0 mL of solution. (MW of NaCN = 49.01 gmol
). The Ka value of HCN is 4.9 x 10-10.
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CONCEPT: ACID & BASE CONSTANTS As you might already realize, there are relatively few strong acids. The great majority of acids are weak acids.
Consider a weak monoprotic acid, HA, and its ionization in water:
HA (aq) + H2O (l) A – (aq) + H3O+ (aq)
The equilibrium expression for this ionization would be:
Ka =
ProductsReac tan ts
=
Where Ka represents the _________________________________________ and it measures the strength of weak acids.
When looking at weak bases we don’t use Ka, but instead _______, which represents the __________________________.
• The relationship between Ka and Kb can be expressed with the following equation:
In general, the __________________ the Ka the stronger the acid and the __________________ the concentration of H+.
In general, the __________________ the pKa the stronger the acid and the __________________ the concentration of H+.
PRACTICE: If the Kb of NH3 is 1.76 x 10-5, determine the acid dissociation constant of its conjugate acid.
KW =Ka ⋅Kb
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PRACTICE: ACID & BASE CONSTANTS (CALCULATIONS 1)
EXAMPLE 1: Knowing that HF has a higher Ka value than CH3COOH, determine, if possible, in which direction the following equilibrium lies.
HF (aq) + CH3COO – (aq) F – (aq) + CH3COOH (aq)
a) Equilibrium lies to the left.
b) Equilibrium lies to the right.
c) Equilibrium is equal and balanced.
d) Not enough information given.
EXAMPLE 2: What is the equilibrium constant for the following reaction and determine if reactants or products are favored.
HCN (aq) + ClO2 – (aq) CN – (aq) + HClO2 (aq)
The acid dissociation constant of HCN is 4.9 x 10-10 and the acid dissociation of HClO2 is 1.1 x 10-2.
HCN (aq) + H2O (aq) CN – (aq) + H3O+ (aq)
HClO2 (aq) + H2O (aq) ClO2 – (aq) + H3O+ (aq)
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PRACTICE: ACID & BASE CONSTANTS (CALCULATIONS 2)
EXAMPLE: Which of the following solutions will have the lowest pH?
a) 0.25 M HC2F3O2
b) 0.25 M HIO4
c) 0.25 M HC3H5O3
d) 0.25 M H2CO3
e) 0.25 M HSeO4–
PRACTICE 1: Which Bronsted-Lowry base has the greatest concentration of hydroxide ions?
a) C2H8N2 (Kb = 8.3 x 10-5)
b) C5H5N (Kb = 1.7 x 10-9)
c) (CH3)3N (Kb = 1.0 x 10-6)
d) C3H7NH2 (Kb = 3.5 x 10-4)
e) C6H5NH2 (Kb = 3.9 x 10-10)
PRACTICE 2: Which Bronsted-Lowry acid has the weakest conjugate base?
a) HCNO (Ka = 2.0 x 10-4)
b) HF (Ka = 3.5 x 10-4)
c) HN3 (Ka = 2.5 x 10-5)
d) H2CO3 (Ka = 4.3 x 10-7)
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CONCEPT: ACID & BASE NEUTRALIZATION
When an acid neutralizes a base an ionic compound called a _______________ is formed.
• These solutions can be neutral, acidic or basic, based on acid-base properties of the cations and anions formed.
RULES FOR IDENTIFYING YOUR IONS
CATIONS (POSITIVE IONS)
1) Transition Metals: If your transition metal has a charge of +2 or higher it is acidic. If the charge is less than +2 then it is
neutral.
EX:
2) Main-Group Metals: If your main-group metal has a charge of +3 or higher it is acidic. If the charge is less than +3 then
it is neutral.
EX:
3) Positive Amines are acidic.
EX:
ANIONS (NEGATIVE IONS)
1) NEGATIVE ION: If you have a negative ion then add an H+ to it. If you create a weak acid then your negative ion is
basic.
EX:
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PRACTICE: ACID & BASE NEUTRALIZATION 1
EXAMPLE: Determine if each of the following compounds will create an acidic, basic or neutral solution.
a) NaOCl b) PbCl4
PRACTICE 1: Determine if each of the following compounds will create an acidic, basic or neutral solution.
a) LiC2H3O2 b) C6H5NH3Br
PRACTICE 2: Determine if each of the following compounds will create an acidic, basic or neutral solution.
a) Co(HSO4)2 b) Sr(HSO3)2
PRACTICE 3: Determine if each of the following compounds will create an acidic, basic or neutral solution.
a) C3H7NH3F
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PRACTICE: ACID & BASE NEUTRALIZATION 2
EXAMPLE 1: Determine whether each compound will become more soluble in an acidic solution.
a) NaBr b) LiCl c) KIO
EXAMPLE 2: Determine the pH of a 0.50 M NH4Cl solution. The Kb of NH3 is 1.75 x 10-5.
PRACTICE: Determine the pH of a 0.55 M NaCN solution. The Ka of hydrocyanic acid, HCN, is 4.9 x 10-10.
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PRACTICE: ACID & BASE NEUTRALIZATION 2
EXAMPLE 1: Determine whether each compound will become more soluble in an acidic solution.
a) NaBr b) LiCl c) KIO
EXAMPLE 2: Determine the pH of a 0.50 M NH4Cl solution. The Kb of NH3 is 1.75 x 10-5.
PRACTICE: Determine the pH of a 0.55 M NaCN solution. The Ka of hydrocyanic acid, HCN, is 4.9 x 10-10.
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CONCEPT: DIPROTIC ACIDS
Diprotic acids and bases are compounds that can donate or accept _______ H+ ion.
For diprotic acids ________ their equations can be illustrated by:
For diprotic bases ________ their equations can be illustrated by:
Based on these equations the relationship between the different forms of diprotic species are:
As a result of these equations for diprotic acids and bases the relationship between Ka and Kb will be:
Ka1 ⋅Kb2 =Kw
Ka2 ⋅Kb1 =Kw
When dealing with diprotic acids:
1) H2A can be treated as a monoprotic acid and we use ________ can be used to find pH.
2) HA – represents the intermediate form and we use ________ can be used to find pH.
3) A2– represents the basic form and we use ________ can be used to find pH.
H2A (aq) + H2O (l) HA– (aq) + H3O+ (aq) Ka1 =ProductsReac tan ts
=
HA– (aq) + H2O (l) A2– (aq) + H3O+ (aq) Ka2 =ProductsReac tan ts
=
A2– (aq) + H2O (aq) HA – (aq) + OH – (aq) Kb1 =ProductsReac tan ts
=
HA – (aq) + H2O (aq) H2A (aq) + OH – (aq) Kb2 =ProductsReac tan ts
=
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PRACTICE: DIPROTIC ACIDS CALCULATIONS 1
EXAMPLE 1: Sulfurous acid, H2SO3, represents a diprotic acid with a Ka1 = 1.6 x 10-2 and Ka2 = 4.6 x 10-5. Calculate the pH and concentrations of H2SO3, HSO3– and SO32– when given 0.250 M H2SO3.
EXAMPLE 2: Determine the pH of 0.115 M Na2S. Hydrosulfuric acid, H2S, contains Ka1 = 1.0 x 10-7 and Ka2 = 9.1 x 10-8.
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CONCEPT: POLYPROTIC ACIDS
Our understanding of diprotic acids and bases can be used to understand polyprotic acids and bases.
For polyprotic acids ________ their equations can be illustrated by:
For polyprotic bases ________ their equations can be illustrated by:
As a result of these equations for polyprotic acids and bases the relationship between Ka and Kb will be:
Ka1 ⋅Kb3 =Kw
Ka2 ⋅Kb2 =Kw
Ka3 ⋅Kb1 =Kw
When dealing with polyprotic acids:
• H3A can be treated as a monoprotic acid and we use ________ can be used to find pH.
• A3– represents the basic form and we use ________ can be used to find pH.
H2A−
[H+ ] ≈ Ka1Ka2[ ]0 +Ka1Kw
Ka1 + [ ]0 HA2−
[H+ ] ≈ Ka2Ka3[ ]0 +Ka2Kw
Ka2 + [ ]0
H3A (aq) + H2O (l) H2A– (aq) + H3O+ (aq) Ka1 =ProductsReac tan ts
=
H2A– (aq) + H2O (l) HA2– (aq) + H3O+ (aq) Ka2 =ProductsReac tan ts
=
HA2– (aq) + H2O (l) A3– (aq) + H3O+ (aq) Ka3 =ProductsReac tan ts
=
A3– (aq) + H2O (l) HA2– (aq) + OH – (aq) Kb1 =ProductsReac tan ts
=
HA2– (aq) + H2O (l) H2A – (aq) + OH – (aq) Kb2 =
ProductsReac tan ts
=
H2A – (aq) + H2O (l) H3A (aq) + OH – (aq) Kb3 =
ProductsReac tan ts
=
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PRACTICE: POLYPROTIC ACIDS CALCULATIONS
EXAMPLE 1: Determine the pH of 0.300 M sodium hydrogen phosphate, Na2HPO4. Phosphoric acid, H3PO4, contains Ka1 = 7.5 x 10-3, Ka2 = 6.2 x 10-8 and Ka3 = 4.2 x 10-13.
EXAMPLE 2: Determine the pH of 0.300 M citric acid, H3C6H5O7 it possesses Ka1 = 7.4 x 10-4, Ka2 = 1.7 x 10-5 and Ka3 =
4.0 x 10-7.
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