Brønsted-Lowry Base: A proton acceptor that takes hydrogen ions (H+) from water.
Reaction Example: B(aq) + H2O → BH+(aq) + OH-(aq)
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Conjugate Acid-Base Pairs
Brønsted-Lowry Theory defines acids and bases as conjugate pairs, capable of transforming into each other through proton transfer.
Acid-base reactions are equilibria.
General Reaction: HA ⇌ H+ + A-
In the reaction, HA donates a proton and transforms into its conjugate base A-, whereas the conjugate base A- accepts a proton to form back the conjugate acid HA.
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Conjugate Acid-Base Examples
NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq)
HCO3-(aq) + S2-(aq) ⇌ HS-(aq) + CO32-(aq)
Indication of acid (HA) and base (B) along with their conjugate pairs.
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Strong Acids and Bases
Strong Acids: Dissociate almost completely in water and release most H+.
Example: HCl → H+ + Cl-
Other strong acids include HNO3 (nitric acid), H2SO4 (sulfuric acid).
Strong Bases: Also dissociate almost completely in water.
Example: NaOH → Na+ + OH-
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Weak Acids and Bases
Most acids and bases are weak, partially ionizing in water and maintaining an equilibrium.
Weak Acids: Slightly dissociate in water with a shift in equilibrium to the left.
Example: CH3COOH ⇌ CH3COO- + H+
Weak Bases: Also slightly dissociate in water.
Example: NH3 + H2O ⇌ NH4+ + OH-
Only about 2% of ethanoic acid (CH3COOH) dissociates in solution.
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Acid-Base Equilibria
An acid requires a base to accept a proton.
Proton transfer occurs between acids (HA) and bases (B).
Reaction: HA(aq) + B(aq) ⇌ BH+(aq) + A-(aq)
Changes in concentrations of acid or base will shift the equilibrium to maintain balance:
Adding HA or B shifts equilibrium right.
Adding BH+ or A- shifts equilibrium left.
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Dissociation of Water
Water can also dissociate slightly, creating an equilibrium:
Reaction: H2O + H2O ⇌ H3O+ + OH-
Alternative representation: H2O ⇌ H+ + OH-
The equilibrium constant expression can be determined, showing very low concentrations of ions lead to constant water concentrations.
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Kw
The ionic product of water (Kw) is a constant, given by:
Kw = 1.0 x 10^-14 mol² dm⁻⁶ at 298 K.
The value of Kw varies with temperature, and in pure water, [H+] = [OH-]. Thus, Kw = [H+]².
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Determination of pH
pH is defined as the negative logarithm of the molar hydrogen-ion concentration:
Formula: pH = -log10[H+]
The pH scale ranges from 0 (very acidic) to 14 (very basic), with 7 being neutral.
Example Calculation: For a 0.005 mol dm⁻³ hydrogen ion concentration:
pH = -log10(0.005) = 2.3
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Calculating [H+]
Hydrogen ion concentration can be derived from pH:
Formula: [H+] = 10^-pH.
Example Calculation: Given a solution of hydrochloric acid with a pH of 2.0:
[H+] = 10^-2.0 = 0.01 mol dm⁻³.
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Monoprotic Acids
Strong acids ionize fully in solution; examples include HCl and HNO3.
Monoprotic acids release one proton when dissociated:
One mole of acid yields one mole of hydrogen ions.
Hydrogen ion concentration equals acid concentration.
Example Calculation: For a 0.05 mol dm⁻³ HCl:
pH = -log10(0.05) = 1.3.
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Diprotic Acids
Strong diprotic acids release two protons per dissociation.
Each mole of diprotic acid yields two moles of hydrogen ions.
Example Calculation: For a 0.01 mol dm⁻³ sulfuric acid (H2SO4):
[H2SO4] = 0.01 mol dm⁻³; therefore,
[H+] = 2 × 0.01 mol dm⁻³ = 0.02 mol dm⁻³;
pH = -log10(0.02) = 1.7.
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Calculating pH of Strong Base
Strong bases fully ionize in water (e.g., NaOH → Na+ + OH-).
Concentrations of bases align with hydroxide ion concentrations:
For a strong base, [Strong base] = [OH-].
To find pH, calculate [H+] using Kw:
Kw = [H+][OH-].
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Base pH Calculation Example
Example: Calculate the pH of 0.20 mol dm⁻³ NaOH at 298 K:
Given Kw = 1.0 x 10^-14, use Kw equation:
1.0 x 10^-14 = [H+] × 0.20 mol dm⁻³
Rearrange:
[H+] = 1.0 x 10^-14 / 0.20 = 5 x 10^-14 mol dm⁻³;
pH = -log10[H+] = -log10(5 x 10^-14) = 13.3.
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Ka
Ka is the acid dissociation constant; it requires two assumptions:
Weak acids dissociate slightly in solutions; [H+] is not equal to [acid].
Use Ka to find pH:
Dissociation: HA(aq) ⇌ H+(aq) + A-(aq).
Assume equilibrium concentration of HA = initial concentration of HA.
Assumption: [H+] = [A-].
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Using Ka to find pH
To calculate the pH of a weak acid using Ka:
Ka is a specific constant for particular acid at fixed temperature.
Example: For 0.030 mol dm⁻³ hydrogen fluoride at 298 K with Ka = 6.6 x 10^-4:
The Ka expression: Ka = [H+]² / [HA].
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Using Ka to find pH
Rearranging the Ka expression to find [H+]:
[H+] = √(Ka × [HA])
Calculation:
[H+] = √(6.6 x 10^-4 × 0.030) = 1.98 x 10^-5.
Find pH:
pH = -log10[H+] = -log10(1.98 x 10^-5) = 4.70.
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Using Ka to find Concentration
Ka allows the calculation of the concentration of a weak acid when given pH and Ka.
Example: Calculate molar concentration of propanoic acid solution with pH 3.30 and Ka = 1.30 x 10^-5:
First calculate [H+]: [H+] = 10^-pH = 10^-3.30 = 5.01 x 10^-4.
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Using Ka to find Concentration
Write the Ka expression to find [CH3CH2COOH]:
Ka = [H+]² / [CH3CH2COOH], rearranging gives [CH3CH2COOH] = [H+]² / Ka.
Substitute into equation:
[CH3CH2COOH] = (5.01 x 10^-4)² / (1.30 x 10^-5) = 0.0193 mol dm⁻³.
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pKa
pKa indicates acid strength:
Lower pKa values indicate stronger acids.
Conversion equation:
pKa = -log10(Ka);
Alternatively, Ka = 10^-pKa.
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pKa Calculations
Example for propanoic acid with Ka = 1.30 x 10^-5:
Calculate pKa:
pKa = -log10(1.30 x 10^-5) = 4.89.
Example for HF with pKa 3.18 to find Ka:
Ka = 10^-3.18 = 6.60 x 10^-4.
Frequently, concentration calculations from pH may require conversion of pKa to Ka first.
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Titration Overview
Description of acid-base titration process:
Unknown acid concentration in a conical flask, while base with known concentration added from a burette.
Use of indicator to signal neutralization of acid by base.
The amount of base used provides the means to calculate the acid's concentration.
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Titrations Best Practices
To achieve precision in titration results:
Measure the unknown volume accurately.
Repeat the titration at least three times for mean calculation.
Ensure all results are within 0.1 cm³ of each other, disregarding any anomalies.
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pH Curves
Visualization of pH changes as bases titrate with acids.
Key curves to note:
Strong acid - strong base
Strong acid - weak base
Weak acid - strong base
Weak acid - weak base
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pH Curves Analysis
Strong Acid - Strong Base:
Starts at pH of 13 before acid addition; drops to 1.
Weak Acid - Strong Base:
Begins at pH of 13, with a decrease to a pH around 5 after adding weak acid.
Strong Acid - Weak Base:
Initiates from the weak base's pH (around 9), dropping to strong acid pH (around 1).
Weak Acid - Weak Base:
Begins at weak base's pH (around 9), moving towards the weak acid's pH (around 5).
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pH Curve Details
Each graph (except for weak acid - weak base) exhibits a vertical section indicating neutralization ends.
A small addition of acid causes drastic pH changes, whereas weak acid - weak base shows gradual shifts.
pH meters are preferred for determining endpoints over color-change indicators.
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Indicators in Titration
An indicator changes color to indicate titration endpoints; it must occur precisely at the end point.
The selected indicator should transition within a narrow pH range, located at the steep part of the pH curve.
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Common Indicators Used
Strong Acid – Strong Base: Methyl orange and Phenolphthalein can be used.
Strong Acid – Weak Base: Methyl Orange is preferred.
Weak Acid – Strong Base: Use Phenolphthalein.
Weak Acid – Weak Base: Neither indicator is suitable due to gradual pH changes.
Indicator
Color at Low pH
pH of Color Change
Color at High pH
Methyl Orange
RED
3.1 – 4.4
YELLOW
Phenolphthalein
COLOURLESS
8.3 – 10
PINK
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Titration Calculations
The results from acid/base titrations yield unknown concentration:
Example: 55 cm³ of HCl neutralized by 30 cm³ of 0.8 mol dm⁻³ NaOH.
Buffers resist changes in pH when small amounts of acid or base are added or when diluted.
Even a tiny amount of acid can drastically adjust pH in water.
Buffers diminish large pH fluctuations but do not completely avoid pH changes.
Both acidic and basic buffers are applicable.
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Acidic Buffers
Acidic buffer solutions comprise a weak acid and its salt; they possess a pH <7.
Example: Ethanoic acid (CH3COOH) and Sodium ethanoate (CH3COO-Na+).
The solution consists of undissociated ethanoic acid and dissociated ethanoate ions.
Adding acid (H+) increases concentration; equilibrium shifts left to counteract.
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Acidic Buffer Responses
For small acid additions, excess H+ react with CH3COO- to form CH3COOH, stabilizing pH.
For small base additions, OH- reacts with H+ to form water, shifting equilibrium right to produce more H+ until pH stabilizes.
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Basic Buffers
Basic buffer solutions consist of a weak base and the salt of that weak base; they have a pH >7.
Example: Ammonia (NH3) and Ammonium chloride (NH4Cl).
The solution contains undissociated ammonia and ammonium ions from dissociated salt.
Adding acid (H+) reduces OH- concentration; equilibrium shifts right to maintain pH.
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Basic Buffers Response
Acid addition increases H+, which forms water with OH-, reducing OH- concentration and shifting equilibrium to create more OH- until returning to original pH.
Base addition raises OH- levels, reacting with NH4+ to yield NH3 and H2O, thus shifting equilibrium left to neutralize excess OH-.
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Calculating Buffer pH
To compute pH of a buffer, knowledge of Ka and concentrations of weak acid and salt is necessary:
Example: pH of buffer containing 0.30 mol dm⁻³ ethanoic acid and 0.50 mol dm⁻³ sodium ethanoate, with Ka = 1.7 x 10^-5.
Ka expression for the buffer:
CH3COOH(aq) ⇌ H+(aq) + CH3COO-(aq).
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Buffer pH Calculation Steps
Rearranging the Ka expression to solve for [H+]:
[H+] = Ka × [CH3COOH] / [CH3COO-].
Insert known values:
[H+] = (1.7 x 10^-5) × (0.30) / (0.50) = 1.02 x 10^-5.
Find pH from [H+]:
pH = -log10(1.02 x 10^-5) = 4.92.
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Buffer Applications
Biological buffers are crucial to maintain blood pH near 7.4 to protect cells and organs.
Hair care products often utilize buffers to prevent damage via alkaline shampoos, maintaining a pH of about 5.5.
Biological washing powders integrate buffers to optimize enzyme functionality for effective cleaning.