Definations of Brønsted acids and bases are givin in Chapter 4. A Brønsted acid is a substance capable of donating a proton, and a Brønsted base is a substance capable of accepting a proton.
Conjugate acid-base pair The conjugate base of a Brønsted acid is the species that remains when one proton has been removed from the acid. Conversely, a conjugate acid results from the addition of a proton to a Brønsted base.
Diprotic and Polyprotic Acids
or
$$\ce{H2O <=> H+ + OH-}$$at $25\,^\circ\text{C}$
$$K_w = [\ce{H+}][\ce{OH-}]=1.0\times10^{-14}$$pH - A Measure of Acidity $$\ce{pH}=-\log [\ce{H+}]$$
In water, at $25\,^\circ\text{C}$, $\ce{pH}=7$ means the solution is neutral; $\ce{pH}<7$, acidic; $\ce{pH}>7$, basic.
$\ce{pOH}$ can be defined as $$\ce{pOH}=-\log [\ce{OH-}]$$
Since $K_w = [\ce{H+}][\ce{OH-}]=1.0\times10^{-14}$, $\ce{pH}+\ce{pOH} = 14$. Knowing one of them means the other is also known. Only $\ce{pH}$ is commonly used.
Strong acids and strong bases are strong electrolytes which are assumed to ionize completely in water.
Weak acids and weak bases are weak electrolytes which would partially ionize in water.
Important! Memorize the following common strong acids and weak acids.
Strong acids: $\ce{HClO4}$, $\ce{HI}$, $\ce{HBr}$, $\ce{HCl}$, $\ce{H2SO4}$, $\ce{HNO3}$.
Weak acids: $\ce{HSO4-}$, $\ce{HF}$, $\ce{HNO2}$, $\ce{HCOOH}$, $\ce{CH3COOH}$, $\ce{NH4+}$, $\ce{HCN}$.
Remember the typical weak bases that would be considered in this general chemistry are mainly amines, such as $\ce{NH3}$, $\ce{N(CH3)3}$, $\ce{C6H5NH2}$, etc.
Requirements
- Understand the concepts, understand the meaning of $\ce{pH}$ value;
- Remember the relavent strength of common acids and bases;
- Given an acid/base reaction, recognize which is the acid and which is the base, and tell what are their conjugate base and conjugate acid.
Hydrohalic acids Stronger bond, weaker acid.
$$\ce{HF}\ll\ce{HCl}<\ce{HBr}<\ce{HI}$$
Oxoacids
Salts that produce neutral solutions A strong acid and a strong base generate a salt that produces neutral solution in water. For example, $\ce{NaOH}$ and $\ce{HCl}$ generate $\ce{NaCl}$ which produces neutral solution in water. The $\ce{pH}$ of $\ce{NaCl}$ solution is $7$.
Salts that produce basic solutions A strong base and a weak acid generate a salt that produces basic solution in water. For example, $\ce{NaOH}$ and $\ce{CH3COOH}$ generate $\ce{CH3COONa}$ which produces basic solution in water because $\ce{CH3COO-}$ can hydrolyze. The $\ce{pH}$ is larger than $7$.
Salts that produce acidic solutions A strong acid and a weak base generate a salt that produces acidic solution in water. For example, $\ce{HCl}$ and $\ce{NH3}$ generate $\ce{NH4Cl}$ which produces acidic solution in water. The $\ce{pH}$ is smaller than $7$. Small and highly charged metal cations can hydrolyze to produce $\ce{H+}$. For example, $\ce{AlCl3, CrCl3, FeCl3, BiCl3, BeCl2}$ solutions are all acidic.
Note that $\ce{Al(H2O)6^3+}$ is roughly as strong an acid as $\ce{CH3COOH}$.
Salts in which both the cation and the anion hydrolyze If the salt is generated from a weak acid and a weak base, the $\ce{pH}$ of the solution depends on the $K_a$ of the cation and the $K_b$ of the anion.
Oxides can be classified as acidic, basic, or amphoteric, see following Figure.
The basic metallic oxides react with water to form metal hydroxides. Acidic oxides and water usually generate acids.
$$\ce{Na2O(s) + H2O(l) -> 2NaOH(aq)}$$
$$\ce{SO3(g) + H2O(l) -> H2SO4(aq)}$$
Reactions between acidic oxides and bases and those between basic oxides and acids resemble normal acid-base reactions
$$\ce{CO2(g) + 2NaOH(aq) -> Na2CO3(aq) + H2O(l)}$$
$$\ce{BaO(s) + 2HNO3(aq)-> Ba(NO3)2(aq) + H2O(l)}$$
aluminum oxide ($\ce{Al2O3}$) is amphoteric. Depending on the reaction conditions, it can behave either as an acidic oxide or as a basic oxide.
$$\ce{Al2O3(s) + 6HCl(aq) -> 2AlCl3(aq) + 3H2O(l)}$$
$$\ce{Al2O3(s) + 2NaOH(aq) + 3H2O(l) -> 2NaAl(OH)4(aq)}$$
Requirements
- Remember the trends of the acidity strength of hydrohalic acids and oxoacids.
- Remember the examples of salts those produce neutral, acidic or basic solutions. Know how to determine the acidity of a salt solution when both the cation and the anion hydrolyze.
- Remember the examples and properties of acidic, basic and amphoteric oxides.
A Lewis Base is a substance that can donate a pair of electrons;
A Lewis Acid is a substance that can accept a pair of electrons.
Learn how to recognize Lewis acids and bases from the following examples:
Requirements
- Understand the defination of Lewis acid and base. Recognize the Lewis acids and bases in reactions.
Calculate $\ce{pH}$ and $\ce{pOH}$ from $[\ce{H+}]$ or $[\ce{OH-}]$
Remember $[\ce{H+}][\ce{OH-}]=1.0\times10^{-14}$ and $\ce{pH}+\ce{pOH} = 14$. Knowing the value of one of the following: $\ce{pH}$, $\ce{pOH}$, $[\ce{H+}]$ and$[\ce{OH-}]$, all other three values can be calculated.
Strong acids and strong bases
Strong acids and strong bases ionize completely in water.
Weak acids and weak bases
Weak acids and weak bases ionize partially in water.
The concentration of $\ce{H+}$ depends on the initial concentraion of $\ce{HA}$ and the value of $K_a$.
\begin{align*}
& \ce{HA(aq) + H2O(l) <=> H3O+(aq) + A- (aq)} \quad K_a = \frac{[\ce{H3O+}][\ce{A-}]}{[\ce{HA}]}\\
\end{align*}
or
\begin{align*}
& \ce{HA(aq) <=> H+(aq) + A- (aq)}\quad K_a=\frac{[\ce{H+}][\ce{A-}]}{[\ce{HA}]}
\end{align*}
Percent Ionization
\begin{align*}
& \text{percent ionization}\\
& =\frac{\text{ionized acid concentration at equilibrium}}{\text{initial concentration of acid}}\times 100\%
\end{align*}
| $\ce{HF}$ | $\ce{H+}$ | $\ce{F-}$ | |
|---|---|---|---|
| Initial (M): | $0.50$ | $0$ | $0$ |
| Change (M): | $-x$ | $+x$ | $+x$ |
| Equilibrium (M): | $(0.50-x)$ | $x$ | $x$ |
The concentration of $\ce{OH-}$ depends on the initial concentraion of the base and the value of $K_b$. Such as:
$$\ce{NH3(aq) + H2O(l) <=> NH4+ (aq) + OH- (aq)}$$
$$K_b = \frac{[\ce{NH4+}][\ce{OH-}]}{[\ce{NH3}]}$$
| $\ce{NH3}$ | $\ce{H2O}$ | $\ce{NH4+}$ | $\ce{OH-}$ | |
|---|---|---|---|---|
| Initial (M): | $0.40$ | N/A | $0$ | $0$ |
| Change (M): | $-x$ | N/A | $+x$ | $+x$ |
| Equilibrium (M): | $(0.40-x)$ | N/A | $x$ | $x$ |
from sympy import *
x,y= symbols('x y')
# Solve an equation to get the value of x
# 1. In calculation of the pH of a HF solution
# You can go to the website [http://live.sympy.org/]
# and try the following codes
Ka = 0.0007
C0=0.5
solve(Ka*(C0-x)-x**2)
# Ignore the nagtive value.
[-0.0190615605976626, 0.0183615605976626]
# 2. In calculatation of the pH of an ammonia solution
Kb=0.000018
C0=0.4
solve(Kb*(C0-x)-x**2)
# Ignore the nagtive value.
[-0.00269229666641615, 0.00267429666641615]
As discussed above, diprotic and polyprotic acids Ionize in a stepwise manner. Each step corresponds to an equilibrium constant. To understand the method to calculate the pH and the concentrations of all species, see the following example.
| $\ce{HC2O4-}$ | $\ce{H+}$ | $\ce{C2O4^2-}$ | |
|---|---|---|---|
| Initial (M): | $0.054$ | $0.054$ | $0$ |
| Change (M): | $-y$ | $+y$ | $+y$ |
| Equilibrium (M): | $(0.054-y)$ | $(0.054-y)$ | $y$ |
Ka1=0.065
C0=0.1
solve(Ka1*(C0-x)-x**2)
[-0.119426693253569, 0.0544266932535686]
Ka2=0.000061
C1=0.054
solve(Ka2*(C1-y)-(C1+y)*y)
[-0.0541218626502987, 6.08626502987110e-5]
As discussed above, an anion from a weak acid or a cation from a weak base can hydrolyze in water to generate either $\ce{OH-}$ or $\ce{H+}$.
Kb= 0.00000000056
C0= 0.24
solve(Kb*(C0-x)-x**2)
[-1.15933814003329e-5, 1.15928214003329e-5]
Requirements
- Calculate $\ce{pH}$ and $\ce{pOH}$ from $[\ce{H+}]$ or $[\ce{OH-}]$;
- Calculate $\ce{pH}$ for a given concentration of a strong acid or a strong base;
- Calculate $\ce{pH}$ for a given concentration of a weak acid ($K_a$) or a weak base ($K_b$);
- Calculate the concentrations of ions for a given concentration of a weak diprotic acid ($K_a$);
- Calculate $\ce{pH}$ of a salt solution.
- Remember the correlation between $K_a$ and $K_b$ of conjugate acid-base.
from IPython.core.display import HTML
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styles = open("custom.css", "r").read()
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