Chapter 3-Acid n Base,Preparation and Dilution
Transcript of Chapter 3-Acid n Base,Preparation and Dilution
CHM 160 (CHAPTER 2)
ACID AND BASES
AND
PREPARATION AND DILUTION OF STOCK SOLUTION
• An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity.
• A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity.
nonelectrolyte weak electrolyte strong electrolyte
ELECTROLYTIC PROPERTIES
METHOD OF DISTINGUISHING BETWEEN ELECTROLYTES AND NONELECTROLYTES
• A pair of inert electrodes (Cu or Pt) is immersed in a beaker of water.
• To light the bulb, electric current must flow from one electrode to the other, thus completing the circuit.
• By adding NaCl (ionic compound), the bulb will glow.
• NaCl breaks up into Na+ and Cl- ions when dissolves in water.
• Na+ are attracted to the negative electrode.
• Cl- are attracted to the positive electrode.
• The movement sets up an electric current that is equivalent to the flow of electrons along a metal wire.
• Strong Electrolyte – 100% dissociation (breaking up of compound into cations and anions
NaCl (s) Na+ (aq) + Cl- (aq)H2O
• Weak Electrolyte – not completely dissociated
CH3COOH CH3COO- (aq) + H+ (aq)
A reversible reaction. The reaction can occur in both directions.
• Hydration is the process in which an ion is surrounded by water molecules arranged in a specific manner.
• Hydration helps to stabilize ions and prevents cations from combining with anions.
H2O
Nonelectrolyte does not conduct electricity?
No cations (+) and anions (-) in solution
C6H12O6 (s) C6H12O6 (aq)H2O
• Have a sour taste. - Vinegar owes its taste to acetic acid. - Citrus fruits contain citric acid.
• React with certain metals to produce hydrogen gas.
• React with carbonates and bicarbonates to produce carbon dioxide gas
• Cause color changes in plant dyes.
2HCl (aq) + Mg (s) MgCl2 (aq) + H2 (g)
2HCl (aq) + CaCO3 (s) CaCl2 (aq) + CO2 (g) + H2O (l)
• Aqueous acid solutions conduct electricity.
PROPERTIES ACIDS
• Have a bitter taste.
• Feel slippery. Many soaps contain bases.
• Cause color changes in plant dyes.
• Aqueous base solutions conduct electricity.
• Examples:
PROPERTIES OF BASES
ROLE OF WATER TO SHOW PROPERTIES ROLE OF WATER TO SHOW PROPERTIES OF ACIDSOF ACIDS
• Anhydrous pure acid (without water) does not show acidic properties.
• In dry form, acids exist as neutral covalent molecules.
• Dry acids do not dissociate to form hydrogen ion (H+).
• When a pure acid is dissolved in water, it will show the properties of acids.
• This is because acids will dissociate in water to form H+ or hydroxonium/hydronium ion, H3O+ which are free to move.
• For example:
i) HCl in liquid methylbenzene (organic solvent) - does not show acidic properties.
ii) HCl in water – show acidic properties
ROLE OF WATER TO SHOW PROPERTIES ROLE OF WATER TO SHOW PROPERTIES OF ALKALIOF ALKALI
• Dry base does not show alkaline properties.
• A base in dry form, contains hydroxide ions (OH-) that are not free to move. Thus, the alkaline properties cannot be shown.
• In the presence of water, bases can dissociate in water to form hydroxide ions, OH-, which are free to move. Thus, alkaline properties are shown.
• For example:
i) ammonia in tetrachlomethane (organic solvent) – do not show alkaline properties
ii) ammonia in water – show alkaline properties
DEFINITION OF ACID AND BASE
Arrhenius Brønsted-Lowry
Lewis
Arrhenius acid is a substance that produces H+ (hydrogen ion) or hydronium ion (H3O+) in water
Arrhenius base is a substance that produces OH- in water
DEFINITION OF ACID AND BASE BY ARRHENIUS
Examples of acid:
CO2 (g) + H2O (l) H2CO3 (aq)
H2CO3 (aq) + H2O(l) H3O+ (aq) + HCO3- (aq)
nonmetal oxides + H2O acid
Examples of bases:
NaOH (s) Na+ (aq) + OH- (aq)
N2H4 (aq) + H2O N2H5+ (aq) + OH- (aq)
metal oxides + H2O bases
Na2O (s) + H2O (l) 2NaOH (aq)
* Limited only to aqueous solutions
• A Brønsted acid is a proton donor• A Brønsted base is a proton acceptor• Example:
acid base acid base
• A Brønsted acid must contain at least one ionizable proton!
DEFINITION OF ACID AND BASE BY BRØNSTED-LOWRY
HCl (aq) +H2O (l) → H3O+ (aq) + Cl- (aq)
• HCl is a acid because it donates proton to H2O
• H2O is a base because it accepts proton from HCl
Brønsted acids and bases
• Conjugate acid-base pair:
i) Conjugate base of a Brønsted acid
- the species that remains when one proton has been removed from the acid
ii) Conjugate acid
- addition of a proton to a Brønsted base
Examples:
HCl (aq) +H2O (l) H3O+ (aq) + Cl- (aq)
acid1 base2 acid2 base1
• Cl- is a conjugate base of HCl and HCl is a conjugate acid of Cl-
• H2O is a base conjugate of H3O+ and H3O+ is a acid conjugate of
H2O
NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq)
base1 acid2 acid1 base2
• subscripts 1 and 2 = two conjugate acid-base pair
• When a strong acid react with a strong base in Brønsted acid-base
reaction, it will give a weak conjugate acid and conjugate base.
•Examples:
HCl (aq) + H2O (l) H3O+ (aq) + Cl- (aq) strong acid strong base weak conjugate weak conjugate acid base NH3 (aq) + H2O (l) NH4
+ (aq) + OH- (aq) Weak base weak acid strong conjugate strong conjugate acid base
• H2O can function as acid or base which called amphoteric
• Amphoteric or amphiprotic substance is one that can react as either an acid or base
Identify each of the following species as a Brønsted acid, base, or both. (a) HI, (b) CH3COO-, (c) H2PO4
-
HI (aq) H+ (aq) + I- (aq) Brønsted acid
CH3COO- (aq) + H+ (aq) CH3COOH (aq) Brønsted base
H2PO4- (aq) H+ (aq) + HPO4
2- (aq)
H2PO4- (aq) + H+ (aq) H3PO4 (aq)
Brønsted acid
Brønsted base
• A Lewis acid is a substance that can accept a pair of electrons
• A Lewis base is a substance that can donate a pair of electrons
H+ H O H••••
+ OH-••••••
acid base
N H••
H
H
H+ +
acid base
N H
H
H
H+
DEFINITION OF ACID AND BASE BY LEWIS
Examples of Lewis Acids and Bases reactions:
N H••
H
H
acid base
F B
F
F
+ F B
F
F
N H
H
H
b) Ag+ (aq) + 2NH3 (aq) Ag(NH3)2+ (aq)
acid base
c) Cd+ (aq) + 4I- (aq) CdI2-4 (aq)
acid base
d) Ni (s) + 4CO (g) Ni(CO)4 (g)
acid base
a)
• Acids
i) Strong acids:- Acids that completely ionized in solution. - Example:
HCl (aq) → H+ (aq) + Cl- (aq)
ii) Weak acids- Acids that incompletely ionized in solution- Example:
CH3COOH (aq) CH3COO- (aq) + H+ (aq)
TYPES OF ACIDS-BASES
Monoprotic acid:
- each unit of the acid yields one hydrogen ion upon ionization
HCl H+ + Cl-
HNO3 H+ + NO3-
CH3COOH H+ + CH3COO-
Strong electrolyte, strong acid
Strong electrolyte, strong acid
Weak electrolyte, weak acid
Diprotic acid:
- each unit of the acid gives up two H+ ions, in two separate steps
H2SO4 H+ + HSO4-
HSO4- H+ + SO4
2-
Strong electrolyte, strong acid
Weak electrolyte, weak acid
Triprotic acids: - yield three H+ ions
H3PO4 H+ + H2PO4-
H2PO4- H+ + HPO4
2-
HPO42- H+ + PO4
3-
Weak electrolyte, weak acidWeak electrolyte, weak acidWeak electrolyte, weak acid
• Bases
i) Strong bases:- Bases that completely ionized in solution. - Example:
NaOH (s) → Na+ (aq) + OH- (aq)
ii) Weak bases- bases that incompletely ionized in solution- Example:
NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq)
Acids and bases as electrolytesAcids and bases as electrolytes
• Strong acids such as HCl and HNO3 are strong electrolytes, while weak acid such as acetic acid (CH3COOH) is a weak electrolyte.
HCl (aq) + H2O (l) H3O+ (aq) + Cl- (aq)
HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq)
HClO4 (aq) + H2O (l) H3O+ (aq) + ClO4- (aq)
H2SO4 (aq) + H2O (l) H3O+ (aq) + HSO4- (aq)
Strong Acids are strong electrolytes
HCl (aq) + H2O (l) H3O+ (aq) + Cl- (aq)
HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq)
HClO4 (aq) + H2O (l) H3O+ (aq) + ClO4- (aq)
H2SO4 (aq) + H2O (l) H3O+ (aq) + HSO4- (aq)
Acids and bases as electrolytesAcids and bases as electrolytes
HF (aq) + H2O (l) H3O+ (aq) + F- (aq)
Weak Acids are weak electrolytes
HNO2 (aq) + H2O (l) H3O+ (aq) + NO2- (aq)
HSO4- (aq) + H2O (l) H3O+ (aq) + SO4
2- (aq)
H2O (l) + H2O (l) H3O+ (aq) + OH- (aq)
Strong Bases are strong electrolytes
NaOH (s) Na+ (aq) + OH- (aq)H2O
KOH (s) K+ (aq) + OH- (aq)H2O
Ba(OH)2 (s) Ba2+ (aq) + 2OH- (aq)H2O
F- (aq) + H2O (l) OH- (aq) + HF (aq)
Weak Bases are weak electrolytes
NO2- (aq) + H2O (l) OH- (aq) + HNO2 (aq)
Conjugate acid-base pairs:
• The conjugate base of a strong acid has no measurable strength.
• H3O+ is the strongest acid that can exist in aqueous solution.
• The OH- ion is the strongest base that can exist in aqueous solution.
H2O (l) H+ (aq) + OH- (aq)
autoionization of water
• Can act either as a acid or as a base.
• Water functions as a base with acids such as HCl and CH3COOH and function as acid in reaction with bases.
• Water is a very weak electrolyte and undergo ionization to a small extent:
ACID-BASE PROPERTIES OF WATER
pH = -log [H+]
[H+] = [OH-]
[H+] > [OH-]
[H+] < [OH-]
Solution Is
neutral
acidic
basic
[H+] = 1 x 10-7
[H+] > 1 x 10-7
[H+] < 1 x 10-7
pH = 7
pH < 7
pH > 7
At 250C
pH [H+]
pH-A MEASURE OF ACIDITY
• pH – the negative logarithm of the hydrogen in
concentration (in mol/L)
pOH = -log [OH-]
[H+][OH-] = Kw = 1.0 x 10-14
-log [H+] – log [OH-] = 14.00
pH + pOH = 14.00
Other important relationships
pH Meter
1) The pH of rainwater collected in a certain region of the northeastern United States on a particular day was 4.82. What is the H+ ion concentration of the rainwater?
pH = -log [H+]
[H+] = 10-pH = 10-4.82 = 1.5 x 10-5 M
2) The OH- ion concentration of a blood sample is 2.5 x 10-7 M. What is the pH of the blood?
pH + pOH = 14.00
pOH = -log [OH-] = -log (2.5 x 10-7) = 6.60
pH = 14.00 – pOH = 14.00 – 6.60 = 7.40
CALCULATION OF pH FOR CALCULATION OF pH FOR SOLUTION CONTAINING A SOLUTION CONTAINING A
STRONG ACID AND A STRONG ACID AND A SOLUTION OF A STRONG SOLUTION OF A STRONG
BASEBASE
1) What is the pH of a 2 x 10-3 M HNO3 solution?
HNO3 is a strong acid – 100% dissociation.
HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq)
pH = -log [H+] = -log [H3O+] = -log(0.002) = 2.7
Start
End
0.002 M
0.002 M 0.002 M0.0 M
0.0 M 0.0 M
2) What is the pH of a 1.8 x 10-2 M Ba(OH)2 solution?
Ba(OH)2 is a strong base – 100% dissociation.
Ba(OH)2 (s) Ba2+ (aq) + 2OH- (aq)Start
End
0.018 M
0.018 M 0.036 M0.0 M
0.0 M 0.0 M
pH = 14.00 – pOH = 14.00 + log(0.036) = 12.6
• The concentration of a solution is the amount of solute present in a given quantity of solvent or solution.
M = molarity =moles of solute
liters of solution
1) What mass of KI is required to make 500. mL of a 2.80 M KI solution?
volume of KI solution moles KI grams KIM KI M KI
500. mL = 232 g KI166 g KI
1 mol KIx
2.80 mol KI
1 L solnx
1 L
1000 mLx
CONCENTRATION OF SOLUTION
Preparing a Solution of Known Concentration
Dilution is the procedure for preparing a less concentrated solution from a more concentrated solution.
DilutionAdd Solvent
Moles of solutebefore dilution (i)
Moles of soluteafter dilution (f)=
MiVi MfVf=
DILUTION OF DILUTION OF SOLUTIONSSOLUTIONS
EXAMPLE:
1) How would you prepare 60.0 mL of 0.200 M HNO3 from a stock solution of 4.00 M HNO3?
M1V1 = M2V2
M1 = 4.00 M M2 = 0.200 M V2 = 0.0600 L V1 = ? L
V1 =M2V2
M2
=0.200 M x 0.0600 L
4.00 M = 0.00300 L = 3.00 mL
Dilute 3.00 mL of acid with water to a total volume of 60.0 mL.
Concentration UnitsConcentration UnitsThe concentration of a solution is the amount of solute present in a given quantity of solvent or solution.
Percent by Mass
% by mass = x 100%mass of solute
mass of solute + mass of solvent
= x 100%mass of solutemass of solution
Percent by Volume (%v/v)
% by volume = x 100%Volume of soluteVolume of solution
M =moles of solute
liters of solution
Molarity (M)
Molality (m)
m =moles of solute
mass of solvent (kg)
• Quantitative analytical process based on measuring volumes.
• The most common form of VA is the titration, a process whereby a standard solution of known concentration is chemically reacted with a solution of unknown concentration in order to determine the concentration of the unknown.
VOLUMETRIC ANALYSIS (VA)
• In a titration a solution of accurately known concentration (standard solution) is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete.
• Equivalence point – the point at which the reaction is complete
• Indicator – substance that changes color at (or near) the equivalence point
TITRATIONS
• Titrations can be used in the analysis of acid-base reactions
H2SO4 + 2NaOH 2H2O + Na2SO4
Slowly add baseto unknown acid
UNTIL
the indicatorchanges color
PROCEDURE FOR THE TITRATION PROCEDURE FOR THE TITRATION
EXAMPLE:1) What volume of a 1.420 M NaOH solution is required to titrate 25.00 mL of a 4.50 M H2SO4 solution?
WRITE THE CHEMICAL EQUATION!
H2SO4 + 2NaOH 2H2O + Na2SO4
MaVa
MbVb
=a
b
Ma = concentration of acidMb = concentration of baseVa = volume of acidVb = volume of basea = coefficient of acidb = coefficient of base
(4.50 M) (25 mL)
(1.420 M) (Vb)=
1
2
Vb = 158 mL
ACID-BASE TITRATIONSStrong Acid-Strong Base Titrations
NaOH (aq) + HCl (aq) H2O (l) + NaCl (aq)
OH- (aq) + H+ (aq) H2O (l)
pH PROFILE OF THE TITRATION (TITRATION CURVE)pH PROFILE OF THE TITRATION (TITRATION CURVE)
• Before addition of NaOH
- pH = 1.00
• When the NaOH added
- pH increase slowly at first
• Near the equivalence point (the point which equimolar amounts of acid and base have reacted)
- the curve rises almost vertically
• Beyond the equivalence point
- pH increases slowly
pH PROFILE OF THE TITRATION pH PROFILE OF THE TITRATION (TITRATION CURVE) (TITRATION CURVE)
CALCULATION OF pH AT EVERY STAGE OF CALCULATION OF pH AT EVERY STAGE OF TITRATIONTITRATION
1) After addition of 10.0 mL of 0.100 M NaOH to 25.0 mL of 0.100 M HCl
Total volume = 35.0 mL
Moles of NaOH in 10.0 mL
= 10.0 mL x 0.100 mol NaOH x 1L
1L NaOH 1000 mL
= 1.00 x 10-3 mol
Moles of HCl in 25.0 mL
= 25.0 mL x 0.100 mol HCl x 1L
1 L HCl 1000 mL
= 2.50 x 10-3 mol
Amount of HCl left after partial neutralization
= (2.50 x 10-3)-(1.00 x 10-3)
= 1.50 x 10-3 mol
Concentration of H+ ions in 35.0 mL
1.50 x 10-3 mol HCl x 1000 mL = 0.0429 M HCl
35.0 mL 1L
[H+] = 0.0429 M,
pH = -log 0.0429 = 1.37
2) After addition of 25.0 mL of 0.100 M NaOH to 25.0 mL 0f 0.100 M HCl[H+] = [OH-] = 1.00 x 10-7pH = 7.00
3) After addition of 35.0 mL of 0.100 M NaOH to 25.0 mL of 0.100 mL of HCl
Total volume = 60.0 mLMoles of NaOH added= 35.0 mL x 0.100 mol NaOH x 1L 1 L NaOH 000 mL= 3.50 x 10-3 mol
Moles of HCl in 25.0 mL solution = 2.50 x 10-3 molAfter complete neutralization of HCl, no of moles of NaOH left = (3.50 x 10-3)-(2.50x10-3)= 1.00 x 10-3 mol
Concentration of NaOH in 60.0 mL solution= 1.00 x 10-3 mol NaOH x 1000 mL 60.0 mL 1L= 0.0167 M NaOH[OH-] = 0.0167 MpOH = -log 0.0167 = 1.78pH = 14.00 – 1.78 = 12.22