At 25°C, number of OH ions in water is 10-X
NA. X can be
Answers
Answer:
At 25°C, number of OH ions in water is 10-X
NA. X can be Hydroxide
Explanation:
Because of its amphoteric nature (i.e., acts as both an acid or a base), water does not always remain as H2O molecules. In fact, two water molecules react to form hydronium and hydroxide ions:
2H2O(l)⇌H3O+(aq)+OH−(aq)(1)
This is also called the self-ionization of water. The concentration of H3O+ and OH− are equal in pure water because of the 1:1 stoichiometric ratio of Equation 1 . The molarity of H3O+ and OH- in water are also both 1.0×10−7M at 25° C. Therefore, a constant of water ( Kw ) is created to show the equilibrium condition for the self-ionization of water. The product of the molarity of hydronium and hydroxide ion is always 1.0×10−14 (at room temperature).
Kw=[H3O+][OH−]=1.0×10−14(2)
Equation 2 also applies to all aqueous solutions. However, Kw does change at different temperatures, which affects the pH range discussed below.
H+ and H3O+ is often used interchangeably to represent the hydrated proton, commonly call the hydronium ion.
Equation 1 can also be written as
H2O⇌H++OH−(3)
As expected for any equilibrium, the reaction can be shifted to the reactants or products:
If an acid ( H+ ) is added to the water, the equilibrium shifts to the left and the OH− ion concentration decreases
If base ( OH− ) is added to water, the equilibrium shifts to left and the H+ concentration decreases.
pH and pOH
Because the constant of water, Kw is 1.0×10−14 (at 25° C), the pKw is 14, the constant of water determines the range of the pH scale. To understand what the pKw is, it is important to understand first what the "p" means in pOH and pH. The addition of the "p" reflects the negative of the logarithm, −log . Therefore, the pH is the negative logarithm of the molarity of H, the pOH is the negative logarithm of the molarity of OH− , and the pKw is the negative logarithm of the constant of water:
pHpOHpKw=−log[H+]=−log[OH−]=−log[Kw](4)(5)(6)
At room temperature,
Kw=1.0×10−14(7)
So
pKw=−log[1.0×10−14]=14(8)(9)
Using the properties of logarithms, Equation 8 can be rewritten as
10−pKw=10−14.(10)
The equation also shows that each increasing unit on the scale decreases by the factor of ten on the concentration of H+ . Combining Equations 4 - 6 and 8 results in this important relationship:
pKw=pH+pOH=14(11)
Equation 11 is correct only at room temperature since changing the temperature will change Kw .
The pH scale is logarithmic, meaning that an increase or decrease of an integer value changes the concentration by a tenfold. For example, a pH of 3 is ten times more acidic than a pH of 4. Likewise, a pH of 3 is one hundred times more acidic than a pH of 5. Similarly a pH of 11 is ten times more basic than a pH of 10.
Properties of the pH Scale
From the simple definition of pH in Equation 4 , the following properties can be identified:
This scale is convenient to use, because it converts some odd expressions such as 1.23×10−4 into a single number of 3.91.
This scale covers a very large range of [H+] , from 0.1 to 10-14. When [H+] is high, we usually do not use the pH value, but simply the [H+] . For example, when [H+]=1.0 , pH = 0. We seldom say the pH is 0, and that is why you consider pH = 0 such an odd expression. A pH = -0.30 is equivalent to a [H+] of 2.0 M. Negative pH values are only for academic exercises. Using the concentrations directly conveys a better sense than the pH scales.
The pH scale expands the division between zero and 1 in a linear scale or a compact scale into a large scale for comparison purposes. In mathematics, you learned that there are infinite values between 0 and 1, or between 0 and 0.1, or between 0 and 0.01 or between 0 and any small value. Using a log scale certainly converts infinite small quantities into infinite large quantities.
The non-linearity of the pH scale in terms of [H+] is easily illustrated by looking at the corresponding values for pH.
Because the negative log of [H+] is used in the pH scale, the pH scale usually has positive values. Furthermore, the larger the pH, the smaller the [H+] .
The Effective Range of the pH Scale
It is common that the pH scale is argued to range from 0-14 or perhaps 1-14, but neither is correct. The pH range does not have an upper nor lower bound, since as defined above, the pH is an indication of concentration of H+. For example, at a pH of zero the hydronium ion concentration is one molar, while at pH 14 the hydroxide ion concentration is one molar. Typically the concentrations of H+ in water in most solutions fall between a range of 1 M (pH=0) and 10-14 M (pH=14). Hence a range of 0 to 14 provides sensible (but not absolute) "bookends" for the scale. One can go somewhat below zero and somewhat above 14 in water, because the concentrations of hydronium ions or hydroxide ions can exceed one molar. Figure 1 depicts the pH scale with common solutions and where they are on the scale.
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