Ionic Equilibrium
Chemical equilibrium and ionic equilibrium are two major concepts in chemistry. Ionic equilibrium deals with the equilibrium involved in an ionization process while chemical equilibrium deals with the equilibrium during a chemical change. Ionic equilibrium is established between the ions and unionized species in a system. Understanding the concept of ionic equilibrium is very important to answer the questions related to certain chemical reactions in chemistry.
Arrhenius Acid
Arrhenius acid act as a good electrolyte as it dissociates to its respective ions in the aqueous solutions. Keeping it similar to the general acid properties, Arrhenius acid also neutralizes bases and turns litmus paper into red.
Bronsted Lowry Base In Inorganic Chemistry
Bronsted-Lowry base in inorganic chemistry is any chemical substance that can accept a proton from the other chemical substance it is reacting with.
![**Problem Statement:**
A weak base has a base hydrolysis constant, \( K_b \), of \( 3.1 \times 10^{-6} \). What is the pH of a 0.11 M solution of the weak base?
**Solution:**
To find the pH of the solution, we first need to determine the concentration of hydroxide ions, \([OH^-]\), in the solution. We can use the expression for the base hydrolysis constant:
\[ K_b = \frac{[BH^+][OH^-]}{[B]} \]
Assuming the initial concentration of the base, [B], is 0.11 M, and \([BH^+]\) and \([OH^-]\) are both \( x \) at equilibrium, we have:
\[ K_b = \frac{x \cdot x}{0.11 - x} \]
Given the small value of \( K_b \), we can approximate \( 0.11 - x \approx 0.11 \), thus:
\[ 3.1 \times 10^{-6} = \frac{x^2}{0.11} \]
Solving for \( x \):
\[ x^2 = 3.1 \times 10^{-6} \cdot 0.11 \]
\[ x^2 = 3.41 \times 10^{-7} \]
\[ x = \sqrt{3.41 \times 10^{-7}} \]
\[ x \approx 5.84 \times 10^{-4} \]
This \( x \) is the concentration of \( OH^- \) ions. To find the pOH:
\[ \text{pOH} = -\log(5.84 \times 10^{-4}) \]
\[ \text{pOH} \approx 3.23 \]
Finally, to find the pH:
\[ \text{pH} = 14 - \text{pOH} \]
\[ \text{pH} \approx 14 - 3.23 \]
\[ \text{pH} \approx 10.77 \]
Thus, the pH of the 0.11 M solution of the weak base is approximately 10.77.](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2Faf8dd5bb-1368-4a10-9929-9294802d0b74%2Fc9584a44-4155-4487-add9-c58ae4073788%2F67p4mkt_processed.png&w=3840&q=75)
Given,
Molarity of a weak base = 0.11 M
Kb of base = 3.1 x 10-6
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