Embedded Files
Mon P5 AG20
Mon P5 AG20
Wed P4 AG20
Wed P4 AG20
A strong base is an atom that is less electronegative and less able to hold onto electrons, making it more reactive and able to accept protons.
A strong base, within the Brønsted-Lowry framework, is a substance with a high affinity for protons that dissociates or ionises completely in aqueous solution. Strong bases generally feature atoms with low electronegativity bearing a negative charge, as less electronegative atoms are better at donating their lone pair to a proton.
A strong base, in terms of both Lewis theory and electronegativity, is a chemical species that possesses a high electron density, usually as a lone pair, that it can readily donate to a Lewis acid. Strong bases are characterized by lower electronegativity of the donor atom, which makes them less likely to hold onto their electron pair.
Examples of Strong Bases
Group 1 and 2 Hydroxides: NaOH, KOH, Ca(OH).
Oxide Ion: O2-
Hydride Ion: H-
Amide Ion: NH2-
A weak base is a highly electronegative atom (like oxygen or fluorine). It holds onto its lone pair of electrons tightly and is unwilling to share them to form a bond with a proton. An atom that strongly attracts electrons. keeps its electrons close, making it a poor electron donor and, therefore, a weak base.
A weak base reacts only partially with water to form conjugate acid (BH+) and hydroxide ions (OH-).
For polyprotic acids, we only need to do the first dissociation. In case you are curious, there is a proof here: https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Acids_and_Bases/Ionization_Constants/Weak_Acids_and_Bases
12C6 2.3a Strong Weak 2 PPT.pptxThurs P5 AG20
Thurs P5 AG20
What is a buffer?
A buffer solution is an aqueous solution that resists significant changes in pH when small amounts of acid or base are added.
It maintains a stable pH, crucial for chemical reactions and biological systems (e.g., blood). It is composed of:
a weak acid and its conjugate base, or
a weak base and its conjugate acid.
The weak acid doesn’t fully dissociate. The conjugate base is only slightly stronger than its acid, so it doesn’t fully grab protons. The weak acid and (slightly less weakly) basic conjugate exist in a dynamic equilibrium.
HA ⇌ H⁺ + A⁻
If You Add Acid (H⁺) the conjugate base removes it:
H⁺ + A⁻ → HA
So the added acid gets neutralised.
If You Add Base (OH⁻) the weak acid removes it:
OH⁻ + HA → A⁻ + H₂O
So the added base gets neutralised.
A buffer works because of partial strength on both sides:
The weak acid partially dissociates.
Its conjugate base is strong enough to react with added H⁺, but weak enough not to destroy the equilibrium.
The buffer works best when [A-] = [HA], which is pH = pKa (from Henderson-Hasselback equation, more on that later).
The purpose of a buffer is to maintain the pH of a solution.
The purpose of a buffer is to maintain the pH of a solution.
Buffers work to maintain the pH of a solution by reacting with and neutralising any strong acid or base that is added to the solution. To do this, buffer solutions must contain
a weak acid, to react with strong bases, and
a weak base, to react with strong acids.
The weak acid and weak base in a buffer solution are a conjugate acid-base pair. A conjugate acid-base pair is a pair of chemicals that form each other when a proton is gained or lost.
When carbonic acid H2CO3 loses a proton, it becomes its conjugate base, the bicarbonate ion HCO3–.
If HCO3– gains a proton, it reforms H2CO3.
Because both the acid and base in a buffer solution are conjugates of each other, the reaction can move in either direction to counteract the addition of strong acids and bases, and so maintain the pH of the solution.
A strong acid, like hydrochloric acid HCl, will dissociate completely in a carbonic acid H2CO3 buffer solution, causing the bicarbonate ions HCO3– to accept the excess protons and form carbonic acid molecules (the reverse of the diagram above).
Carbonic acid is a much weaker acid than the hydrochloric acid, tTherefore, the pH rises back to its original range.
A strong base, like sodium hydroxide NaOH, will dissociate completely in a carbonic acid buffer solution.
The added hydroxide ions react with carbonic acid, causing the carbonic acid to release protons. The protons released from carbonic acid bond to the hydroxide ions, forming water. This prevents a significant increase in the pH of the solution.
The bicarbonate ions produced when the carbonic acid releases protons act as a reservoir of base, available to react with any additional acid that may be added to the solution.
Remember:
the general reaction for an acid in aqueous solution (below), and
the strength of an acid is how likely it is to donate a proton. This means that for a weak acid, the reversible reaction (below) will establish an equilibrium based on the strength of the acid.
For the reaction between an acidic compound and water, the equilibrium constant (or acid dissociation constant) is:
The pH of a buffer solution can be determined using the Henderson-Hasselbach Equation.
A buffer works best when pH = pKa.
What is the real effect?
1.0 L of buffer containing:
0.50 mol acetic acid (HA) [HA] = 0.5M
0.50 mol acetate (A⁻) [A-] = 0.5M
pKa = 4.76
Add 0.1M HCl
A⁻ decreases by 0.10 mol: 0.5 - 0.1 = 0.4M
HA increases by 0.10 mol: 0.5 + 0.1 = 0.6M
Adding that same amount of acid to water is a drop from 7 to 1. The buffer shows a pH drop of 0.18.
Q1. Weak Acid
Calculate the pH of a solution containing 0.15 M methanoic acid (HCOOH) and 0.10 M sodium methanoate (HCOONa). The pKa is 3.75.
This is a buffer solution containing a weak acid, methanoic acid (HCOOH), and its conjugate base, sodium methanoate (HCOONa).
Solve it by applying the Henderson-Hasselbalch equation for weak acids:
pH = pKa + log ([A–]/[HA])
[A–] is the concentration of the conjugate base (methanoate ion, HCOO-) and [HA] is the concentration of the weak acid (methanoic acid, HCOOH).
Since sodium methanoate is a soluble salt, it completely dissociates in water, providing the same concentration of methanoate ions as the initial concentration of the salt:
[A-] = [HCOO-] = 0.10 M
The concentration of methanoic acid, the weak acid, is:
[HA] = [HCOOH] = 0.15 M
Now, insert these values into the Henderson-Hasselbalch equation, along with the pKa value of methanoic acid:
pH = 3.75 + log (0.10/0.15)
Calculating the logarithm and adding it to the pKa:
pH = 3.75 – 0.18
pH ≈ 3.57
Thus, the pH of the solution containing 0.15 M methanoic acid and 0.10 M sodium methanoate is approximately 3.57.
Weak Base
Calculate the pH of a solution containing 0.25 M ammonia (NH3) and 0.10 M ammonium chloride (NH4Cl). The pKb of ammonia is 4.75.
This is a buffer solution containing a weak base, ammonia NH3, and its conjugate acid ammonium ion NH4+. To find the pH of this solution, apply the Henderson-Hasselbalch equation for weak bases:
pOH = pKb + log ([B]/[HB+])
[B] is the concentration of the weak base (ammonia, NH3) and [HB+] is the concentration of the conjugate acid (ammonium ion, NH4+).
Ammonium chloride is a salt and so it completely dissociates in water, providing the same concentration of ammonium ions as the initial concentration of the salt:
[HB+] = [NH4+] = 0.10 M
The concentration of ammonia, the weak base, is:
[B] = [NH3] = 0.25 M
Now, plug these values into the Henderson-Hasselbalch equation for weak bases, along with the pKb value of ammonia:
pOH = 4.75 + log (0.25/0.10)
Calculate the logarithm and add it to the pKb:
pOH = 4.75 + 0.70 pOH ≈ 5.45
Now, convert pOH to pH. The sum of pH and pOH equals 14:
pH + pOH = 14
Therefore, the pH of the solution is:
pH = 14 – pOH pH = 14 – 5.45 pH ≈ 8.55
Thus, the pH of the solution containing 0.25 M ammonia and 0.10 M ammonium chloride is approximately 8.55.
A buffer with a pH of 5.00 is needed using acetic acid/sodium acetate. The pKa of acetic acid is 4.75. What is the ratio of [A-] / [HA]?
Fri P1 AG20
Fri P1 AG20
Research two natural buffering systems (one for the human body e.g. carbonic acid/carbonate in blood, and one for the natural environment) and outline the following:
Importance of maintaining a specific pH
Changes to the environment which could affect the pH
Indicate the shift/direction to minimise the disturbance.
Some ideas:
Carbonic Acid-Bicarbonate Buffer: The most crucial extracellular buffer. It manages metabolic acids by converting them to carbon dioxide, which is exhaled, or by using bicarbonate to neutralize acids.
Protein Buffer System: Proteins, particularly hemoglobin in red blood cells and albumin in plasma, act as buffers by binding hydrogen ions when pH drops.
Phosphate Buffer System: Operates mainly inside cells and in the kidneys to maintain, respectively, intracellular fluid and urine pH.
Buffers in Nature https://www.youtube.com/watch?v=xVob4KYHvRg
Marine/Water Systems: The carbonic acid/hydrogen carbonate system helps regulate the pH of oceans and water bodies to protect marine life.
Page updated
Google Sites
Report abuse