Question

What is the pH of a solution with a 4.60 × 10−4 M hydroxide ion concentration?

Answer 1

Answer:

The correct answer is option :10.7

Explanation:

The pOH of the solution is the negative logarithm of hydroxide ions in a solution.

$pOH=-\loh[OH^-]$

$pOH=-\log[4.60\times 10^{-4}M]=3.33$

pH+pOh=14

$pH=14-pOH=14-3.33=10.67\approx 10.7$

The correct answer is option :10.7

Answer 2

Answer:

pH

=

10.66

Explanation:

For pure water at  

25

∘

C

, the concentration of hydronium ions,  

H

3

O

+

, is equal to the concentration of hydroxide ions,  

OH

−

.

More specifically, water undergoes a self-ionization reaction that results in the formation of equal concentrations of hydronium and hydroxide anions.

2

H

2

O

(l]

⇌

H

3

O

+

(aq]

+

OH

−

(aq]

At room temperature, the self-ionization constant of water is equal to

K

W

=

[

H

3

O

+

]

⋅

[

OH

−

]

=

10

−

14

This means that neutral water at this temperature will have

[

H

3

O

+

]

=

[

OH

−

]

=

10

−

7

M

As you know, pH and pOH are defined as

pH

=

−

log

(

[

H

3

O

+

]

)

pOH

=

−

log

(

[

OH

−

]

)

and have the following relationship

pH

+

pOH

=

14

In your case, the concentration of hydroxide ions is bigger than  

10

−

7

M

, which tells you that you're dealing with a basic solution and that you can expect the pH of the water to be higher than  

7

.

A pH equal to  

7

is characteristic of a neutral aqueous solution at room temperature.

So, you can use the given concentration of hydroxide ions to determine the pOH of the solution first

pOH

=

−

log

(

4.62

⋅

10

−

4

)

=

3.34

This means that the solution's pH will be

pH

=

14

−

pOH

pH

=

14

−

3.34

=

10.66

Indeed, the pH is higher than  

7

, which confirms that you're dealing with a basic solution.