Acids and Bases Lesson 15 by Owen Borville October 28, 2025
Bronsted Acids and Bases: When a Bronsted acid donates a proton, what remains of the acid is known as a conjugate base.
HCl(aq) (acid) + H2O(l)(base) ⇌ H3O+(aq)(conjugate acid) + Cl-(aq)(conjugate base)
The acid (HCl) loses a proton to the conjugate base Cl-.
The conjugate acid H3O+ gains a proton from the base H2O
The two species HCl and Cl- are known as a conjugate acid-base pair or simply a conjugate pair.
When a Bronsted base accepts a proton, the newly formed protonated species is known as a conjugate acid.
Molecular Structure and Acid Strength: The strength of an acid is measured by its tendency to ionize: HX =>H+ + X-
Two factors affect ionization (1) The strength of the H-H bond. (2) The polarity of the H-X bond. 𝛿+ , 𝛿-, H-X
Hydrohalic acid strength:
HF << HCl < HBr < HI
Biggest factor is bond strength.
Only HF is a weak acid. HCl, HBr, and HI are strong acids.
Oxoacids: contains hydrogen, oxygen, and a central nonmetal atom. Ex. Carbonic acid, Nitrous acid, Nitric acid, Phosphorus acid, Phosphoric acid, Sulfuric acid.
To compare oxoacid strength, it is convenient to divide the oxoacids into two groups:
(1) Oxoacids having different central atoms that are from the same group of the periodic table and that have the same oxidation number. HClO3 > HBrO3. Cl is more electronegative. The O-H bond is more polar. (2) Oxoacids having the same central atom, but different numbers of oxygen atoms. Perchloric acid HClO4 > Chloric acid HClO3 > Chlorous acid HClO2 > Hypochlorous acid HClO More oxygens bonded to the central atom of an acid make it a stronger acid, because it increases the polarity of the O-H bond, making the hydrogen easier to donate, oxidation number is higher, and stabilizes the resulting negative charge on the conjugate base.
Carboxylic Acids are an important group of organic acids. R is part of the acid and the other part is the carboxyl group-COOH. The conjugate base can be represented by more than one structure. Stability of the anion is determined by the ability to delocalize electron density. A more stable anion is from a stronger acid.
Acid strength depends on the nature of the R group. The electronegative Cl atom shifts the electron density towards the R group. Chloroacetic acid (C2H3ClO2) is the stronger of the two acids. (than acetic acid, CH3COOH).
Acid-Base Properties of Water: A species that can behave either as a Bronsted acid or Bronsted base is called amphoteric. The acid-base properties of water produces H3O+ and OH- ions in the equilibrium with water in a reaction known as the autoionization of water.
2H2O(l)⇌ H3O+(aq) + OH-(aq)
acid + base ⇌ conjugate acid + conjugate base
The equilibrium expression of the autoionization of water is given by: Kw = [H3O+][OH-] = 1.0 x 10^-14 M^2 (at 25 degrees C)
Since the product of the concentrations of H3O+ and OH- is equal to a constant, the relative amount of each obeys a fixed relationship. Depending on which ion concentration is in excess, the solution will be considered acidic or basic.
When [H3O+] = [OH-], the solution is neutral
When [H3O+] > [OH-], the solution is acidic
When [H3O+] < [OH-], the solution is basic
The acidity of an aqueous solution depends on the concentration of hydronium ions, [H3O+].
The pH of a solution is defined as the negative base 10 logarithm of the hydronium ion concentration (in mol/L)
pH = -log[H3O+]
[H3O+] = 10^-pH M
In pure water at 25 degrees C, [H3O+] = log 1.0 x 10^-7 M
pH = -log(1.0 x 10^-7] = 7.00
pH is a dimensionless quantity
As pH decreases, acid concentration increases.
A pOH scale is analogous to the pH scale and can be defined as the negative base-10 logarithm of the hydroxide ion concentration.
pOH = -log[OH-] concentration
[OH-] concentration = 10^-pOH M
From the definition of pH and pOH, pH + pOH = 14.00
The pOH scale is the reverse of the pH scale.
pOH Scale Range
Basic 1.00-5.00
Neutral 7.00
Acidic 9.00-13.00
Strong acid dissociations are not treated as equilibria, but rather as processes that go to completion:
Hydrochloric acid HCl(aq) + H2O(l) => H3O+(aq) + Cl-(aq)
Hydrobromic acid HBr(aq) + H2O(l) => H3O+(aq) + Br-(aq)
Hydroiodic acid HI(aq) + H2O(l) => H3O+(aq) + I-(aq)
Nitric acid HNO3(aq) + H2O(l) => H3O+(aq) + NO3-(aq)
Chloric acid HClO3(aq) + H2O(l) => H3O+(aq) + ClO3-(aq)
Perchloric acid HClO4(aq) + H2O(l) => H3O+(aq) + ClO4-(aq)
Sulfuric acid H2SO4(aq) + H2O(l) => H3O+ HSO4-(aq)
Strong Bases consist of the hydroxides of alkali metals and the heaviest alkaline earth metals:
Group 1A Hydroxides
LiOH(aq) => Li+(aq) + OH-(aq)
NaOH(aq) => Na+(aq) + OH-(aq)
KOH(aq) => K+(aq) + OH-(aq)
RbOH(aq) => Rb+(aq)+OH-(aq)
CsOH(aq) => Cs+(aq) + OH-(aq)
Group 2A Hydroxides
Ca(OH)2(aq) => Ca2+(aq) + 2OH-(aq)
Sr(OH)2(aq) => Sr2+(aq) + 2OH- (aq)
Ba(OH)2(aq) => Ba2+(aq) + 2OH-(aq)
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The ionization of a weak monoprotic acid HA in water is represented by: HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq)
Ka = [H3O+][A-]/[HA] Ka is called the acid ionization constant. The larger the value of Ka, the stronger the acid. Ka values can be found in reference tables. To find the pH or concentration, set up equilibrium tables based on the chemical equation and solve for x. Initial concentration is M. Change in concentration is (-x) or reactants and (+x for each product, or x^2). Equilibrium concentration is (M-x). x=change in concentration. Therefore, Ka = (x)(x)/(M-x)= Ka = (x^2)/(M-x)
Percent Ionization is a quantitative measure of the degree of ionization.
Percent ionization = [H3O+]eq/[HA]0 x 100%
Ex. If concentration is = 0.0025 M, [HA]=0.1, then percent ionization = (0.0025)/(0.100) x 100% = 2.5%
Ex. If concentration is = 0.0018 M, [HA]=0.050, then percent ionization = (0.0018)/(0.050) x 100% = 3.6%
On a graphical plot of original concentration of acid (x) versus percent ionization (y), when percent ionization is 100%, this represents a strong acid. Percent ionization depends on the original acid concentration and as this original acid concentration approaches zero, ionization approaches 100%.
Weak base ionization is incomplete and is treated in the same way as the ionization of a weak acid.
B(aq) + H2O(l) ⇌ HB+(aq) + OH-(aq)
Kb = [HB][OH-]/[B]
Kb is the base ionization constant. The larger the value of Kb, the stronger the base. Kb values can be found in reference tables.To find the pH or concentration, set up equilibrium tables based on the chemical equation and solve for x. Initial concentration is M. Change in concentration is (-x) or reactants and (+x for each product, or x^2). Equilibrium concentration is (M-x). x= change in concentration. Therefore, Kb = (x)(x)/(M-x)= Kb = (x^2)/(M-x)
Conjugate Acid-Base Pairs Strength: A strong acid ionizes completely in water: HCl(aq) => H+(aq) + Cl-(aq) (No affinity for the H+ ion)
The chloride ion is a weak conjugate base. Cl-(aq) + H2O(l) =>X> HCl(aq) + OH-(aq)
A weak acid ionizes to a limited degree in water: HF(aq) ⇌ H+(aq) + F-(aq) (strong affinity for the H+ ion).
The fluoride ion is a strong conjugate base. F-(aq) + H2O(l) ⇌ HF(aq) + OH-(aq)
A strong acid has a weak conjugate base.
A weak acid has a strong conjugate base.
A strong base has a weak conjugate acid.
A weak base has a strong conjugate acid.
A simple relationship between the ionization constant of a weak acid (Ka) and the ionization constant of a weak base (Kb) can be derived.
Ka * Kb = Kw (autoionization constant of water, 1x10^-14 mol2*dm-6 at room temperature), so Ka = Kw/Kb and Kb = Kw/Ka
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Diprotic and Polyprotic Acids undergo successive ionizations, losing one proton at a time, and each has a Ka associated with it.
H2CO3(aq) ⇌ H+(aq) + HCO3-(aq)
Ka1 = [H+][HCO3-]/[H2CO3]
HCO3-(aq) ⇌ H+(aq) + (CO3)2-(aq)
Ka2 = [H+][(CO3)2-]/[HCO3-]
Ka1 > Ka2
For a given acid, the first ionization constant is much larger than the second, and so on.
To calculate concentrations, set up an equilibrium table using the chemical equation and x as the unknown for the first ionization and y as the unknown in the second ionization.
M= initial concentration, y=change in concentration
So Ka1 = (x^2)/(M-x)
and Ka2 = (M+y)(y)/(M-y)
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Salt hydrolysis occurs when ions produced by the dissociation of a salt react with water to produce either hydroxide ions or hydronium ions.
Basic salts (conjugates of weak acids): F-(aq) + H2O(l) ⇌ HF(aq) + OH-(aq)
Acidic salts (conjugates of weak bases): NH4+(aq) + H2O(l) ⇌ NH3(aq) + H3O+(aq)
Set up an equilibrium table. M=initial concentration, x=change in concentration. Using the chemical formula and Kb=Kw/Ka to find Kb.
Then, Kb = (x^2)/(M-x). Then find pOH and pH using the respective log equations pH = -log[H3O+] and pOH = -log[OH-]
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Acidic Salt Solutions: Small, highly charged metal ions can react with water to produce an acidic solution.
Acid-Base Properties of Salt Solutions: The pH of salt solutions can be qualitatively predicted by determining which ions facilitate hydrolysis.
A cation that will make a solution acidic is: The conjugate acid of a weak base or a small, highly charged metal ion (other than Group 1A or 2A) Ex. NH4+, Al3+, Cr3+, Fe3+, Bi3+, CH3NH3+, C2H5NH3+
An anion that will make a solution basic is: The conjugate base of a weak acid. Ex. CN-, NO2-, CH3COO-
A cation that will not affect the pH of a solution is: Group 1A or heavy Group 2A (except Be2+) Ex. Li+, Na+, Ba+
An anion that will not affect the pH of a solution is: The conjugate base of a strong acid. Ex. Cl-, NO3-, ClO4-
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Salts When Both the Cation and Anion Hydrolyze: The pH of a solution that contains a salt in which both the cation and the anion hydrolyze depends on the relative strengths of the weak acid and base. Qualitative predictions can be made using the Kb (of the salts anion) and the Ka (of the salts cation).
When Kb>Ka, the solution is basic.
When Kb<Ka, the solution is acidic.
When Kb is approximately = Ka, the solution is neutral or nearly neutral.
Basic metallic oxides react with water to form metal hydroxides:
Na2O(s) + H2O(l) => 2NaOH(aq)
BaO(s) + H2O(l) =>Ba(OH)2(aq)
Acidic oxides reaction with water as follows:
CO2(g) + H2O(l)⇌ H2CO3(aq)
SO3(g) + H2O(l) ⇌H2SO4(aq)
Reactions between acidic oxides and bases and those between basic oxides and acids resemble normal acid-base reactions that produce a salt and water.
CO2(g) + 2NaOH(aq) =>Na2CO3(aq) + H2O(l)
BaO(s) + 2HNO3(aq) => Ba(NO3)2(aq) + H2O(l)
Aluminum Oxide (Al2O3) is amphoteric: because it can act as an acid or base:
As acid: Al2O3(s) + 6HCl(aq) => 2AlCl3(aq) + 3H2O(l)
As base: Al2O3(s) + 2NaOH(aq) + 3H2O(l) =>2NaAl(OH)4(aq)
All the alkali and alkaline earth metal hydroxides, except Be(OH)2 are basic.
Amphoteric: Be(OH)2, Al(OH)3, Sn(OH)2, Pb(OH)2, Cr(OH)3, Cu(OH)2, Zn(OH)2, Cd(OH)2
Acid: Be(OH)2(s) + 6H+(aq) => 2Be2+(aq) + 2H2O(l)
Base: Be(OH)2(s) + 2OH-(aq) =>Be(OH)4(2-)(aq)
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.
Bronsted Acids and Bases: When a Bronsted acid donates a proton, what remains of the acid is known as a conjugate base.
HCl(aq) (acid) + H2O(l)(base) ⇌ H3O+(aq)(conjugate acid) + Cl-(aq)(conjugate base)
The acid (HCl) loses a proton to the conjugate base Cl-.
The conjugate acid H3O+ gains a proton from the base H2O
The two species HCl and Cl- are known as a conjugate acid-base pair or simply a conjugate pair.
When a Bronsted base accepts a proton, the newly formed protonated species is known as a conjugate acid.
Molecular Structure and Acid Strength: The strength of an acid is measured by its tendency to ionize: HX =>H+ + X-
Two factors affect ionization (1) The strength of the H-H bond. (2) The polarity of the H-X bond. 𝛿+ , 𝛿-, H-X
Hydrohalic acid strength:
HF << HCl < HBr < HI
Biggest factor is bond strength.
Only HF is a weak acid. HCl, HBr, and HI are strong acids.
Oxoacids: contains hydrogen, oxygen, and a central nonmetal atom. Ex. Carbonic acid, Nitrous acid, Nitric acid, Phosphorus acid, Phosphoric acid, Sulfuric acid.
To compare oxoacid strength, it is convenient to divide the oxoacids into two groups:
(1) Oxoacids having different central atoms that are from the same group of the periodic table and that have the same oxidation number. HClO3 > HBrO3. Cl is more electronegative. The O-H bond is more polar. (2) Oxoacids having the same central atom, but different numbers of oxygen atoms. Perchloric acid HClO4 > Chloric acid HClO3 > Chlorous acid HClO2 > Hypochlorous acid HClO More oxygens bonded to the central atom of an acid make it a stronger acid, because it increases the polarity of the O-H bond, making the hydrogen easier to donate, oxidation number is higher, and stabilizes the resulting negative charge on the conjugate base.
Carboxylic Acids are an important group of organic acids. R is part of the acid and the other part is the carboxyl group-COOH. The conjugate base can be represented by more than one structure. Stability of the anion is determined by the ability to delocalize electron density. A more stable anion is from a stronger acid.
Acid strength depends on the nature of the R group. The electronegative Cl atom shifts the electron density towards the R group. Chloroacetic acid (C2H3ClO2) is the stronger of the two acids. (than acetic acid, CH3COOH).
Acid-Base Properties of Water: A species that can behave either as a Bronsted acid or Bronsted base is called amphoteric. The acid-base properties of water produces H3O+ and OH- ions in the equilibrium with water in a reaction known as the autoionization of water.
2H2O(l)⇌ H3O+(aq) + OH-(aq)
acid + base ⇌ conjugate acid + conjugate base
The equilibrium expression of the autoionization of water is given by: Kw = [H3O+][OH-] = 1.0 x 10^-14 M^2 (at 25 degrees C)
Since the product of the concentrations of H3O+ and OH- is equal to a constant, the relative amount of each obeys a fixed relationship. Depending on which ion concentration is in excess, the solution will be considered acidic or basic.
When [H3O+] = [OH-], the solution is neutral
When [H3O+] > [OH-], the solution is acidic
When [H3O+] < [OH-], the solution is basic
The acidity of an aqueous solution depends on the concentration of hydronium ions, [H3O+].
The pH of a solution is defined as the negative base 10 logarithm of the hydronium ion concentration (in mol/L)
pH = -log[H3O+]
[H3O+] = 10^-pH M
In pure water at 25 degrees C, [H3O+] = log 1.0 x 10^-7 M
pH = -log(1.0 x 10^-7] = 7.00
pH is a dimensionless quantity
As pH decreases, acid concentration increases.
A pOH scale is analogous to the pH scale and can be defined as the negative base-10 logarithm of the hydroxide ion concentration.
pOH = -log[OH-] concentration
[OH-] concentration = 10^-pOH M
From the definition of pH and pOH, pH + pOH = 14.00
The pOH scale is the reverse of the pH scale.
pOH Scale Range
Basic 1.00-5.00
Neutral 7.00
Acidic 9.00-13.00
Strong acid dissociations are not treated as equilibria, but rather as processes that go to completion:
Hydrochloric acid HCl(aq) + H2O(l) => H3O+(aq) + Cl-(aq)
Hydrobromic acid HBr(aq) + H2O(l) => H3O+(aq) + Br-(aq)
Hydroiodic acid HI(aq) + H2O(l) => H3O+(aq) + I-(aq)
Nitric acid HNO3(aq) + H2O(l) => H3O+(aq) + NO3-(aq)
Chloric acid HClO3(aq) + H2O(l) => H3O+(aq) + ClO3-(aq)
Perchloric acid HClO4(aq) + H2O(l) => H3O+(aq) + ClO4-(aq)
Sulfuric acid H2SO4(aq) + H2O(l) => H3O+ HSO4-(aq)
Strong Bases consist of the hydroxides of alkali metals and the heaviest alkaline earth metals:
Group 1A Hydroxides
LiOH(aq) => Li+(aq) + OH-(aq)
NaOH(aq) => Na+(aq) + OH-(aq)
KOH(aq) => K+(aq) + OH-(aq)
RbOH(aq) => Rb+(aq)+OH-(aq)
CsOH(aq) => Cs+(aq) + OH-(aq)
Group 2A Hydroxides
Ca(OH)2(aq) => Ca2+(aq) + 2OH-(aq)
Sr(OH)2(aq) => Sr2+(aq) + 2OH- (aq)
Ba(OH)2(aq) => Ba2+(aq) + 2OH-(aq)
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The ionization of a weak monoprotic acid HA in water is represented by: HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq)
Ka = [H3O+][A-]/[HA] Ka is called the acid ionization constant. The larger the value of Ka, the stronger the acid. Ka values can be found in reference tables. To find the pH or concentration, set up equilibrium tables based on the chemical equation and solve for x. Initial concentration is M. Change in concentration is (-x) or reactants and (+x for each product, or x^2). Equilibrium concentration is (M-x). x=change in concentration. Therefore, Ka = (x)(x)/(M-x)= Ka = (x^2)/(M-x)
Percent Ionization is a quantitative measure of the degree of ionization.
Percent ionization = [H3O+]eq/[HA]0 x 100%
Ex. If concentration is = 0.0025 M, [HA]=0.1, then percent ionization = (0.0025)/(0.100) x 100% = 2.5%
Ex. If concentration is = 0.0018 M, [HA]=0.050, then percent ionization = (0.0018)/(0.050) x 100% = 3.6%
On a graphical plot of original concentration of acid (x) versus percent ionization (y), when percent ionization is 100%, this represents a strong acid. Percent ionization depends on the original acid concentration and as this original acid concentration approaches zero, ionization approaches 100%.
Weak base ionization is incomplete and is treated in the same way as the ionization of a weak acid.
B(aq) + H2O(l) ⇌ HB+(aq) + OH-(aq)
Kb = [HB][OH-]/[B]
Kb is the base ionization constant. The larger the value of Kb, the stronger the base. Kb values can be found in reference tables.To find the pH or concentration, set up equilibrium tables based on the chemical equation and solve for x. Initial concentration is M. Change in concentration is (-x) or reactants and (+x for each product, or x^2). Equilibrium concentration is (M-x). x= change in concentration. Therefore, Kb = (x)(x)/(M-x)= Kb = (x^2)/(M-x)
Conjugate Acid-Base Pairs Strength: A strong acid ionizes completely in water: HCl(aq) => H+(aq) + Cl-(aq) (No affinity for the H+ ion)
The chloride ion is a weak conjugate base. Cl-(aq) + H2O(l) =>X> HCl(aq) + OH-(aq)
A weak acid ionizes to a limited degree in water: HF(aq) ⇌ H+(aq) + F-(aq) (strong affinity for the H+ ion).
The fluoride ion is a strong conjugate base. F-(aq) + H2O(l) ⇌ HF(aq) + OH-(aq)
A strong acid has a weak conjugate base.
A weak acid has a strong conjugate base.
A strong base has a weak conjugate acid.
A weak base has a strong conjugate acid.
A simple relationship between the ionization constant of a weak acid (Ka) and the ionization constant of a weak base (Kb) can be derived.
Ka * Kb = Kw (autoionization constant of water, 1x10^-14 mol2*dm-6 at room temperature), so Ka = Kw/Kb and Kb = Kw/Ka
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Diprotic and Polyprotic Acids undergo successive ionizations, losing one proton at a time, and each has a Ka associated with it.
H2CO3(aq) ⇌ H+(aq) + HCO3-(aq)
Ka1 = [H+][HCO3-]/[H2CO3]
HCO3-(aq) ⇌ H+(aq) + (CO3)2-(aq)
Ka2 = [H+][(CO3)2-]/[HCO3-]
Ka1 > Ka2
For a given acid, the first ionization constant is much larger than the second, and so on.
To calculate concentrations, set up an equilibrium table using the chemical equation and x as the unknown for the first ionization and y as the unknown in the second ionization.
M= initial concentration, y=change in concentration
So Ka1 = (x^2)/(M-x)
and Ka2 = (M+y)(y)/(M-y)
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Salt hydrolysis occurs when ions produced by the dissociation of a salt react with water to produce either hydroxide ions or hydronium ions.
Basic salts (conjugates of weak acids): F-(aq) + H2O(l) ⇌ HF(aq) + OH-(aq)
Acidic salts (conjugates of weak bases): NH4+(aq) + H2O(l) ⇌ NH3(aq) + H3O+(aq)
Set up an equilibrium table. M=initial concentration, x=change in concentration. Using the chemical formula and Kb=Kw/Ka to find Kb.
Then, Kb = (x^2)/(M-x). Then find pOH and pH using the respective log equations pH = -log[H3O+] and pOH = -log[OH-]
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Acidic Salt Solutions: Small, highly charged metal ions can react with water to produce an acidic solution.
Acid-Base Properties of Salt Solutions: The pH of salt solutions can be qualitatively predicted by determining which ions facilitate hydrolysis.
A cation that will make a solution acidic is: The conjugate acid of a weak base or a small, highly charged metal ion (other than Group 1A or 2A) Ex. NH4+, Al3+, Cr3+, Fe3+, Bi3+, CH3NH3+, C2H5NH3+
An anion that will make a solution basic is: The conjugate base of a weak acid. Ex. CN-, NO2-, CH3COO-
A cation that will not affect the pH of a solution is: Group 1A or heavy Group 2A (except Be2+) Ex. Li+, Na+, Ba+
An anion that will not affect the pH of a solution is: The conjugate base of a strong acid. Ex. Cl-, NO3-, ClO4-
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Salts When Both the Cation and Anion Hydrolyze: The pH of a solution that contains a salt in which both the cation and the anion hydrolyze depends on the relative strengths of the weak acid and base. Qualitative predictions can be made using the Kb (of the salts anion) and the Ka (of the salts cation).
When Kb>Ka, the solution is basic.
When Kb<Ka, the solution is acidic.
When Kb is approximately = Ka, the solution is neutral or nearly neutral.
Basic metallic oxides react with water to form metal hydroxides:
Na2O(s) + H2O(l) => 2NaOH(aq)
BaO(s) + H2O(l) =>Ba(OH)2(aq)
Acidic oxides reaction with water as follows:
CO2(g) + H2O(l)⇌ H2CO3(aq)
SO3(g) + H2O(l) ⇌H2SO4(aq)
Reactions between acidic oxides and bases and those between basic oxides and acids resemble normal acid-base reactions that produce a salt and water.
CO2(g) + 2NaOH(aq) =>Na2CO3(aq) + H2O(l)
BaO(s) + 2HNO3(aq) => Ba(NO3)2(aq) + H2O(l)
Aluminum Oxide (Al2O3) is amphoteric: because it can act as an acid or base:
As acid: Al2O3(s) + 6HCl(aq) => 2AlCl3(aq) + 3H2O(l)
As base: Al2O3(s) + 2NaOH(aq) + 3H2O(l) =>2NaAl(OH)4(aq)
All the alkali and alkaline earth metal hydroxides, except Be(OH)2 are basic.
Amphoteric: Be(OH)2, Al(OH)3, Sn(OH)2, Pb(OH)2, Cr(OH)3, Cu(OH)2, Zn(OH)2, Cd(OH)2
Acid: Be(OH)2(s) + 6H+(aq) => 2Be2+(aq) + 2H2O(l)
Base: Be(OH)2(s) + 2OH-(aq) =>Be(OH)4(2-)(aq)
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.