Aqueous Chemical Reactions Lesson CH9 by Owen Borville 10.19.2025
A solution is a homogenous mixture of two or more substances. The substance present in the largest amount (moles) is referred to as the solvent. The other substances present are called the solutes. A substance that dissolves in a particular solvent is said to be soluble in that solvent.
An electrolyte is a substance that dissolves in water to yield a solution that conducts electricity. An electrolyte undergoes dissociation and breaks apart into its constituent ions, such as when NaCl breaks down: NaCl(s) in water=> Na+ (aq) + Cl-(aq). Ionization is when a molecular compound forms ions when it dissolves.
A nonelectrolyte is a substance that dissolves in water to yield a solution that does not conduct electricity, such as sucrose [C12H22O11(s)] when it dissolves in water. The sucrose molecules remain intact upon dissolving.
Strong electrolytes are electrolytes that dissociate completely, such as water soluble ionic compounds. Examples of strong electrolytes are NaCl (sodium chloride), HCl (hydrochloric acid), and NaOH (sodium hydroxide), KOH (potassium hydroxide), H2SO4 (sulfuric acid), HNO3- (nitric acid), and KI (potassium iodide). Strong electrolytes dissolve and dissociate into its constituent ions. Strong acids like HCl and strong bases like NaOH are strong electrolytes.
Other strong acids in aqueous solutions in addition to HCl (hydrochloric acid) are hydrobromic acid (HBr), hydroiodic acid (HI), nitric acid (HNO3), chloric acid (HClO3), perchloric acid (HClO4), and sulfuric acid (H2SO4).
Weak electrolytes are compounds that produce ions upon dissolving but exist in solution predominantly as molecules that are not ionized. Weak electrolytes include weak acids like HC2H3O2 (acetic acid), where H+ ions partially dissociate from the acid. Weak bases include NH3 (ammonia), whose solution dissolves into ammonia and hydroxide ions.
The double arrow (⇌ ) in the chemical equation denotes a reaction that occurs in both directions in the form A + B ⇌ C + D. When both the forward and reverse reactions occur at the same rate, the reaction is in a state of dynamic chemical equilibrium. The double arrow indicates reversibility of the chemical reaction forward and reverse. The double arrow represents dynamic equilibrium where the rate of the forward reaction is equal to the rate of the reverse reaction. Because the rates are equal, the concentrations of all reactants and products appear to be constant, even though both reactions are still occurring at the molecular level.
Ex. H2 + I2 ⇌ 2HI
Classifying electrolytes involves (1) identifying if the compound is ionic or molecular (2) Is the compound an acid or base or neither? (3) Is the compound one of the seven strong acids mentioned above? Any soluble ionic compound is a strong electrolyte, however, most molecular compounds are nonelectrolytes or weak electrolytes, with the exception of the strong acids mentioned above.
Precipitates are insoluble products that separate from a solution, with the form AB (aq) + CD (aq) => AD (s) + CB (aq). A chemical reaction in which a precipitate forms is called a precipitation reaction.
Water is a good solvent for ionic compounds because it is a polar molecule. The polarity of water molecules results from electron distributions within the molecule. The oxygen atom has an attraction for the hydrogen atom's electrons and is therefore partially negative compared to hydrogen. In the water molecule, the oxygen atom is partially negative and the hydrogen atoms are partially positive.
Hydration occurs when water molecules remove the individual ions from an ionic solid surrounding them so the substances dissolve.
Solubility is defined as the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature.
In a molecular equation, compounds are represented by chemical formulas as though they exist in solution as molecules or formula units. Reactions in which cations in two ionic compounds exchange anions are called metathesis or double replacement reactions.
Ionic Equations: In an ionic equation, compounds that exist completely or predominately as ions in a solution are represented as those ions. In the reaction between aqueous Na2SO4 and Ba(OH)2
Na2SO4 (aq) + Ba(OH)2(aq) => 2Na(OH)(aq) + BaSO4(s)
the aqueous species are represented as follows:
Na2SO4(aq) => 2Na+(aq) + SO4(aq)
Ba(OH)2(aq) => Ba2+(aq) + 2OH-(aq)
NaOH(aq) =>Na+(aq) + OH-(aq)
2Na+(aq) + SO42-(aq) + Ba2+(aq) + 2OH-(aq) => 2Na+(aq) + 2OH-(aq) + BaSO4(s)
Net Ionic Equations: An equation that includes only the species that are actually involved in the reaction is called a net ionic equation. Ions that appear on both sides of the equation are called spectator ions. Spectator ions do not participate in the reaction.
2Na+(aq) + (SO4)2-(aq) + Ba2+(aq) + 2OH+(aq) => 2Na+(aq) + 2OH+(aq) + BaSO4(s) (Spectator ions are Na+ and OH-)
The reaction without the spectator ions is : Ba2+(aq) + SO42-(aq) =>BaSO4(s)
Precipitation Reactions: To determine the molecular, ionic, and net ionic equations: (1) Write and balance the molecular equation, predicting the products by assuming that the cations trade anions. (2) Write the ionic equation by separating strong electrolytes into their constituent ions. (3) Write the net ionic equation by identifying and canceling spectator ions on both sides of the equation. (4) If both the reactants and products are all strong electrolytes, all the ions in solution are spectator ions. In this case, there is no net ionic equation and no reaction takes place.
Acid-Base Reactions: Acids can be either strong or weak. A strong acid is a strong electrolyte.
Strong acids are: HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4
Strong bases are: LiOH, NaHO, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2
Strong Acids and Bases: Strong bases are strong electrolytes (dissociate completely). Strong bases are the hydroxides of Group 1A and heavy Group 2A. Ex. NaOH(s)=>Na+(aq)+OH-(aq)
A weak acid is a weak electrolyte and it does not dissociate completely. Most acids are weak acids. Ex. HF (aq) ⇌ H+(aq) + F-(aq)
Bronsted Acids and Bases:
Arrhenius acid is an acid that ionizes in water to produce H+ ions. Ex. HCl(g) => H+(aq) + Cl-(aq)
Arrhenius base is a base that dissociates in water to produce OH- ions. Ex. NaOH(s) => Na+(aq) + OH-(aq)
Bronsted acid is a proton donor.
Bronsted base is a proton acceptor.
In these definitions, a proton refers to a hydrogen atom that has lost its electron-also known as a hydrogen ion (H+).
NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq) (H2O is a Bronsted acid because it donates a proton to become OH-. NH3 is a Bronsted base because it accepts a proton to become NH4+)
Bronsted acids donate protons to water to form the hydronium ion (H3O+). The hydrogen ion (H+), the proton, and the hydronium ion (H3O+) all refer to the same aqueous species. HF(aq) + H2O(l) ⇌ H3O (aq) + F-(aq)
Monoprotic acids have one proton to donate. Hydrochloric acid is an example: HCl(g) => H+(aq) + Cl-(aq)
Polyprotic acids have more than one acidic hydrogen atom. Sulfuric acid, H2SO4, is an example of a diprotic acid because there are two acidic hydrogen atoms. Polyprotic acids lose protons in a stepwise fashion: (1) H2SO4(aq) => H+(aq) + HSO4-(aq) In H2SO4, the first ionization is strong. (2) HSO4-(aq) ⇌ H+(aq) + (SO4)2-(aq) In H2SO4, the second ionization occurs only to a very small extent.
Monobasic bases are bases that produce only one mole of hydroxide per mole of compound. Sodium hydroxide is an example. NaOH(s) =>Na+(aq) + OH-(aq) (one equivalent of hydroxide.
Dibasic Base: Some strong bases produce more than one hydroxide per mole of compound. Barium hydroxide is an example of a dibasic base. Ba(OH)2(s) => Ba2+(aq) + 2OH-(aq) (two equivalents of hydroxide).
Acid-Base Neutralization: A neutralization reaction is a reaction between an acid and a base. Generally, a neutralization reaction produces water and a salt. (Ex): HCl(aq)(acid) + NaOH(aq)(base) => H2O(l)(water) + NaCl(aq)(salt). The net ionic equation of many acid-base reactions is: H+(aq) + OH-(aq) +> H2O(l)
Oxidation-Reduction (redox) reactions are chemical reactions in which electrons are transferred from one reactant to another. Oxidation is the loss of electrons. Reduction is the gain of electrons.
Ex: Zn metal loses two electrons and is oxidized to Zn2+. Zn2+ is called the reducing agent: Zn(s) + Cu2+(aq) => Zn(2+)(aq) + Cu(s). (Cu2+ gains two electrons and is reduced to Cu metal. Cu is called the oxidizing agent.
A redox reaction is the sum of an oxidation half-reaction and a reduction half-reaction. Zn(s) + Cu2+(aq) => Zn(2+)(aq) + Cu(s)
Oxidation half-reaction: Zn(s) => Zn2+(aq) + 2e-
Reduction half-reaction: Cu2+(aq) + 2e- => Cu(s)
Overall redox reaction: Cu2+(aq) + Zn(s) => Zn2+(aq) +Cu(s)
Oxidation numbers: The oxidation number is the charge an atom would have if electrons were transferred completely.
Ex:------------------- H2(g) + F2(g) => 2HF(g)
Oxidation number: 0 0 +1,-1
Total contribution
to charge: 0 0 +1,-1
--------------------------------------------------------
Ex:-------------------N2(g) + 3H2(g) => 2NH3(g)
Oxidation number: 0 0 -3, +1
Total contribution
to charge: 0 0 -3, +3
The oxidation number is sometimes called the oxidation state.
To assign oxidation numbers: (1) The oxidation number of an element, in its elemental form, is zero. (2) The oxidation numbers in any chemical species must sum to the overall charge on the species; must sum to zero for any molecule; must sum to the charge on any polyatomic ion; the oxidation number of a monatomic ion is equal to the charge on the ion. (3) know the elements that nearly always have the same oxidation number. Oxidation numbers can be found on the Periodic Table for each element, and are most commonly identified by their group number (vertical column), however, many elements have multiple oxidation numbers.
Ex: Assign oxidation numbers to KMnO4:
K = +1 (total contribution to charge)
Mn = +7 (total contribution to charge)
O4 = (-2)4 = -8 (total contribution to charge)
The numbers must sum to zero because KMnO4 is a neutral compound.
Ex: Assign oxidation numbers to H2SO4
H2 =(+1)2 = +2 total contribution to charge
S = +6 total contribution to charge
O4 = +2(4) = -8 total contribution to charge
The numbers must sum to zero because H2SO4 is neutral.
Ex: Assign oxidation numbers to ClO3-
Cl = +5
O3 = -2(3) = -6
The numbers sum to negative one, so the chemical is a -1 ion.
Oxidation of Metals in Aqueous Solutions: In a displacement reaction, an atom or an ion in a compound is replaced by an atom of another element. Zn(s) + CuCl2(aq) => ZnCl2(aq) + Cu(s).
Zn = 0
Cu = +2
Cl2 = -1(2) = -2
=>
Zn = +2
Cl2 = -1(2) = -2
Cu = 0
Zinc displaces, or replaces copper in the dissolved salt. Zn is oxidized to Zn2+. Cu2+ is reduced to Cu. When a metal is oxidized by an aqueous solution, it becomes an aqueous ion.
------------------------------------------------------------------------------------
The activity series is a list of metals (and hydrogen) arranged from top to bottom in order of decreasing ease of oxidation. Metals listed at the top are called active metals. Metals listed at the bottom are called noble metals. An element in the series will be oxidized by the ions of any element that appears below it in the table.
Ex: Zinc
Iron
Nickel
Hydrogen
Copper
Silver
Gold
Zn(s) + CuCl2(aq) => ZnCl2(aq) + Cu(s)
Cu(s) + ZnCl2(aq) => no reaction
------------------------------------------------------------------------------
In the following reaction series, which reactions will occur?
Barium
Sodium
Cobalt
Tin
Copper
Silver
Gold
Co(s) + BaI2(aq) = No reaction because cobalt is below barium
Sn(s) + CuBr2(aq) => Cu(s) + SnBr2(aq)
Ag(s) + NaCl(aq) => No reaction because silver is below sodium
Balancing Simple Redox Equations: Redox reactions must have both mass balance and charge balance:
Cr(s) + Ni2+(aq) => Cr3+(aq) + Ni(s)
Oxidation half-reaction: Cr(s) => Cr3+(aq) + 3e-
Reduction half-reaction: Ni2+(aq) + 2e- => Ni (s)
Before adding half-reactions, the electrons must balance.
Prior to adding the two half-reactions, balance the electrons.
2[Cr(s) => Cr3+(aq) + 3e-]
3[Ni2+(aq) +2e- => Ni(s)]
Oxidation half-reaction 2Cr(s) => 2Cr3+(aq) + 6e-
Reduction half reaction 3Ni2+(aq) + 6e- => 3 Ni(s)
------------------------------------------------------------
---------------------------->3Ni2+(aq) + 2Cr(s) => 3Ni(s) + 2Cr3+(aq)
This is the half reaction method.
Combination reactions are another type of Redox reactions. Combination reactions can involve oxidation and reduction:
N2(g) + 3H2(g) => 2NH3(g)
N2=0,0
H2=0,0
=>
N=-3, -3
H3=+1, +3
Hydrogen is oxidized from 0 to +1.
Nitrogen is reduced from 0 to -3.
Decomposition can also be a redox reaction.
NaH(s) => 2Na(s) + 3H2(g)
Na=+1, +1
H=-1, -1
=>
Na=0,0
H2=0,0
Na+ is reduced to Na
H- is oxidized to H2
Disproportionation reactions occur when one element undergoes both oxidation and reduction.
2H2O2(aq) => 2H2O(l) + O2(g)
H2=+1, +2
O2=-1, -2
=>
H2=+1, +2
O=-2, -2
O2=0,0
Oxygen in H2O2 (and other peroxides) has an oxidation number of -1.
Combustion is also a redox process.
CH4(g) + 2O2(g) => CO2 + 2H2O(l)
-4,+4 + 0,0 => +4, -4 + +2, -2
Molarity (M) or molar concentration is defined as the number of moles of solute per liter of solution. M is a measure of concentration of moles of solute per liter of solution.
molarity = (moles of solute)/(liters of solution L)
L=(mol)/(M)
Dilution is the process of preparing a less concentrated solution from a more concentrated one. The moles of solute before dilution = moles of solute after dilution.
Ex. Convert 1L of a 1.00 M solution of KMnO4 to a 0.4 M solution of KMnO4.
Mc x Lc = Md x Ld
(1.00 M KMnO4)(Lc) = (O.400 M KMnO4)(1.00 L)
Lc = O.400 L or 400 mL
To make the solution, pipet 400 mL of stock solution into a 1.00 L volumetric flask and carefully dilute to the calibration mark.
Dilution and Serial Dilution: Because most volumes in the laboratory are in millimeters rather than liters, the equation can be written as: Mc x mLc = Md x mLd
A series of dilutions that may be used to prepare a number of increasingly dilute solutions is called a Serial Dilution: Step 1) Prepare a dilute solution from the stock. Step 2) Dilute a portion of the prepared solution to make a more dilute solution. Step 3) Repeat as needed.
The pH Scale: 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
In pure water at 25 degrees C, [H3O+] = log 1.o x 10^-7
pH = -log(1.0 x 10^-7) = 7.00. pH is a dimensionless quantity. pH ranges from 1.00 to 14.00, where lower than 7.00 is considered increasingly acidic, 7.00 is considered neutral, and greater than 7.00 is considered increasingly basic pH.
pH values of some common fluids are stomach acid (1.0), lemon juice (2.0), vinegar (3.0), orange juice (3.5), urine (4.8-7.5), rainwater (5.5), saliva (6.4-6.9), milk (6.5), pure water (7.0), blood (7.4), tears (7.4), milk of magnesia (10.6), ammonia (11.5).
Aqueous Reactions and Chemical Analysis: Gravimetric analysis is an analytical technique based on the measurement of mass. Gravimetric analysis is highly accurate and it is applicable only to reactions that go to completion or have nearly 100 percent yield.
Acid-Base Titrations: Quantitative studies of acid-base neutralization reactions are most conveniently carried out using a technique known as a titration. A titration is a volumetric technique that uses burets. The point in the titration technique where the acid has been neutralized is called the equivalence point.
The equivalence point is usually signaled by a color change. The color change is brought about by the use of an indicator. Indicators have distinctly different colors in acidic and basic media. The indicator is chosen so that the color change, or endpoint, is very close to the equivalence point. Phenolphthalein is a common indicator.
Sodium hydroxide solutions are commonly used in titrations. NaOH solutions must be standardized as the concentrations change over time. (NaOH reacts with CO2 that slowly dissolves into the solution forming a carbonic acid).
The acid potassium hydrogen phthalate (KHP) is frequently used to standardize NaOH solutions. KHP is a monoprotic acid because it can donate only one proton (H+) per molecule, making it useful for standardizing solutions in titrations due to its chemical structure, which contains a single acidic hydrogen atom attached to the phthalate ion (HC8H4O4-). Monoprotic means that KHP can lose a single proton (H+) in a chemical reaction.
In a neutralization reaction with a base like sodium hydroxide (NaOH), KHP and NaOH react in a 1:1 mole ratio because both are monoprotic. Because it has a large molar mass and is a solid, KHP is frequently used as a primary standard to determine the concentration of a base solution in an acid-base titration technique.
A solution is a homogenous mixture of two or more substances. The substance present in the largest amount (moles) is referred to as the solvent. The other substances present are called the solutes. A substance that dissolves in a particular solvent is said to be soluble in that solvent.
An electrolyte is a substance that dissolves in water to yield a solution that conducts electricity. An electrolyte undergoes dissociation and breaks apart into its constituent ions, such as when NaCl breaks down: NaCl(s) in water=> Na+ (aq) + Cl-(aq). Ionization is when a molecular compound forms ions when it dissolves.
A nonelectrolyte is a substance that dissolves in water to yield a solution that does not conduct electricity, such as sucrose [C12H22O11(s)] when it dissolves in water. The sucrose molecules remain intact upon dissolving.
Strong electrolytes are electrolytes that dissociate completely, such as water soluble ionic compounds. Examples of strong electrolytes are NaCl (sodium chloride), HCl (hydrochloric acid), and NaOH (sodium hydroxide), KOH (potassium hydroxide), H2SO4 (sulfuric acid), HNO3- (nitric acid), and KI (potassium iodide). Strong electrolytes dissolve and dissociate into its constituent ions. Strong acids like HCl and strong bases like NaOH are strong electrolytes.
Other strong acids in aqueous solutions in addition to HCl (hydrochloric acid) are hydrobromic acid (HBr), hydroiodic acid (HI), nitric acid (HNO3), chloric acid (HClO3), perchloric acid (HClO4), and sulfuric acid (H2SO4).
Weak electrolytes are compounds that produce ions upon dissolving but exist in solution predominantly as molecules that are not ionized. Weak electrolytes include weak acids like HC2H3O2 (acetic acid), where H+ ions partially dissociate from the acid. Weak bases include NH3 (ammonia), whose solution dissolves into ammonia and hydroxide ions.
The double arrow (⇌ ) in the chemical equation denotes a reaction that occurs in both directions in the form A + B ⇌ C + D. When both the forward and reverse reactions occur at the same rate, the reaction is in a state of dynamic chemical equilibrium. The double arrow indicates reversibility of the chemical reaction forward and reverse. The double arrow represents dynamic equilibrium where the rate of the forward reaction is equal to the rate of the reverse reaction. Because the rates are equal, the concentrations of all reactants and products appear to be constant, even though both reactions are still occurring at the molecular level.
Ex. H2 + I2 ⇌ 2HI
Classifying electrolytes involves (1) identifying if the compound is ionic or molecular (2) Is the compound an acid or base or neither? (3) Is the compound one of the seven strong acids mentioned above? Any soluble ionic compound is a strong electrolyte, however, most molecular compounds are nonelectrolytes or weak electrolytes, with the exception of the strong acids mentioned above.
Precipitates are insoluble products that separate from a solution, with the form AB (aq) + CD (aq) => AD (s) + CB (aq). A chemical reaction in which a precipitate forms is called a precipitation reaction.
Water is a good solvent for ionic compounds because it is a polar molecule. The polarity of water molecules results from electron distributions within the molecule. The oxygen atom has an attraction for the hydrogen atom's electrons and is therefore partially negative compared to hydrogen. In the water molecule, the oxygen atom is partially negative and the hydrogen atoms are partially positive.
Hydration occurs when water molecules remove the individual ions from an ionic solid surrounding them so the substances dissolve.
Solubility is defined as the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature.
In a molecular equation, compounds are represented by chemical formulas as though they exist in solution as molecules or formula units. Reactions in which cations in two ionic compounds exchange anions are called metathesis or double replacement reactions.
Ionic Equations: In an ionic equation, compounds that exist completely or predominately as ions in a solution are represented as those ions. In the reaction between aqueous Na2SO4 and Ba(OH)2
Na2SO4 (aq) + Ba(OH)2(aq) => 2Na(OH)(aq) + BaSO4(s)
the aqueous species are represented as follows:
Na2SO4(aq) => 2Na+(aq) + SO4(aq)
Ba(OH)2(aq) => Ba2+(aq) + 2OH-(aq)
NaOH(aq) =>Na+(aq) + OH-(aq)
2Na+(aq) + SO42-(aq) + Ba2+(aq) + 2OH-(aq) => 2Na+(aq) + 2OH-(aq) + BaSO4(s)
Net Ionic Equations: An equation that includes only the species that are actually involved in the reaction is called a net ionic equation. Ions that appear on both sides of the equation are called spectator ions. Spectator ions do not participate in the reaction.
2Na+(aq) + (SO4)2-(aq) + Ba2+(aq) + 2OH+(aq) => 2Na+(aq) + 2OH+(aq) + BaSO4(s) (Spectator ions are Na+ and OH-)
The reaction without the spectator ions is : Ba2+(aq) + SO42-(aq) =>BaSO4(s)
Precipitation Reactions: To determine the molecular, ionic, and net ionic equations: (1) Write and balance the molecular equation, predicting the products by assuming that the cations trade anions. (2) Write the ionic equation by separating strong electrolytes into their constituent ions. (3) Write the net ionic equation by identifying and canceling spectator ions on both sides of the equation. (4) If both the reactants and products are all strong electrolytes, all the ions in solution are spectator ions. In this case, there is no net ionic equation and no reaction takes place.
Acid-Base Reactions: Acids can be either strong or weak. A strong acid is a strong electrolyte.
Strong acids are: HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4
Strong bases are: LiOH, NaHO, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2
Strong Acids and Bases: Strong bases are strong electrolytes (dissociate completely). Strong bases are the hydroxides of Group 1A and heavy Group 2A. Ex. NaOH(s)=>Na+(aq)+OH-(aq)
A weak acid is a weak electrolyte and it does not dissociate completely. Most acids are weak acids. Ex. HF (aq) ⇌ H+(aq) + F-(aq)
Bronsted Acids and Bases:
Arrhenius acid is an acid that ionizes in water to produce H+ ions. Ex. HCl(g) => H+(aq) + Cl-(aq)
Arrhenius base is a base that dissociates in water to produce OH- ions. Ex. NaOH(s) => Na+(aq) + OH-(aq)
Bronsted acid is a proton donor.
Bronsted base is a proton acceptor.
In these definitions, a proton refers to a hydrogen atom that has lost its electron-also known as a hydrogen ion (H+).
NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq) (H2O is a Bronsted acid because it donates a proton to become OH-. NH3 is a Bronsted base because it accepts a proton to become NH4+)
Bronsted acids donate protons to water to form the hydronium ion (H3O+). The hydrogen ion (H+), the proton, and the hydronium ion (H3O+) all refer to the same aqueous species. HF(aq) + H2O(l) ⇌ H3O (aq) + F-(aq)
Monoprotic acids have one proton to donate. Hydrochloric acid is an example: HCl(g) => H+(aq) + Cl-(aq)
Polyprotic acids have more than one acidic hydrogen atom. Sulfuric acid, H2SO4, is an example of a diprotic acid because there are two acidic hydrogen atoms. Polyprotic acids lose protons in a stepwise fashion: (1) H2SO4(aq) => H+(aq) + HSO4-(aq) In H2SO4, the first ionization is strong. (2) HSO4-(aq) ⇌ H+(aq) + (SO4)2-(aq) In H2SO4, the second ionization occurs only to a very small extent.
Monobasic bases are bases that produce only one mole of hydroxide per mole of compound. Sodium hydroxide is an example. NaOH(s) =>Na+(aq) + OH-(aq) (one equivalent of hydroxide.
Dibasic Base: Some strong bases produce more than one hydroxide per mole of compound. Barium hydroxide is an example of a dibasic base. Ba(OH)2(s) => Ba2+(aq) + 2OH-(aq) (two equivalents of hydroxide).
Acid-Base Neutralization: A neutralization reaction is a reaction between an acid and a base. Generally, a neutralization reaction produces water and a salt. (Ex): HCl(aq)(acid) + NaOH(aq)(base) => H2O(l)(water) + NaCl(aq)(salt). The net ionic equation of many acid-base reactions is: H+(aq) + OH-(aq) +> H2O(l)
Oxidation-Reduction (redox) reactions are chemical reactions in which electrons are transferred from one reactant to another. Oxidation is the loss of electrons. Reduction is the gain of electrons.
Ex: Zn metal loses two electrons and is oxidized to Zn2+. Zn2+ is called the reducing agent: Zn(s) + Cu2+(aq) => Zn(2+)(aq) + Cu(s). (Cu2+ gains two electrons and is reduced to Cu metal. Cu is called the oxidizing agent.
A redox reaction is the sum of an oxidation half-reaction and a reduction half-reaction. Zn(s) + Cu2+(aq) => Zn(2+)(aq) + Cu(s)
Oxidation half-reaction: Zn(s) => Zn2+(aq) + 2e-
Reduction half-reaction: Cu2+(aq) + 2e- => Cu(s)
Overall redox reaction: Cu2+(aq) + Zn(s) => Zn2+(aq) +Cu(s)
Oxidation numbers: The oxidation number is the charge an atom would have if electrons were transferred completely.
Ex:------------------- H2(g) + F2(g) => 2HF(g)
Oxidation number: 0 0 +1,-1
Total contribution
to charge: 0 0 +1,-1
--------------------------------------------------------
Ex:-------------------N2(g) + 3H2(g) => 2NH3(g)
Oxidation number: 0 0 -3, +1
Total contribution
to charge: 0 0 -3, +3
The oxidation number is sometimes called the oxidation state.
To assign oxidation numbers: (1) The oxidation number of an element, in its elemental form, is zero. (2) The oxidation numbers in any chemical species must sum to the overall charge on the species; must sum to zero for any molecule; must sum to the charge on any polyatomic ion; the oxidation number of a monatomic ion is equal to the charge on the ion. (3) know the elements that nearly always have the same oxidation number. Oxidation numbers can be found on the Periodic Table for each element, and are most commonly identified by their group number (vertical column), however, many elements have multiple oxidation numbers.
Ex: Assign oxidation numbers to KMnO4:
K = +1 (total contribution to charge)
Mn = +7 (total contribution to charge)
O4 = (-2)4 = -8 (total contribution to charge)
The numbers must sum to zero because KMnO4 is a neutral compound.
Ex: Assign oxidation numbers to H2SO4
H2 =(+1)2 = +2 total contribution to charge
S = +6 total contribution to charge
O4 = +2(4) = -8 total contribution to charge
The numbers must sum to zero because H2SO4 is neutral.
Ex: Assign oxidation numbers to ClO3-
Cl = +5
O3 = -2(3) = -6
The numbers sum to negative one, so the chemical is a -1 ion.
Oxidation of Metals in Aqueous Solutions: In a displacement reaction, an atom or an ion in a compound is replaced by an atom of another element. Zn(s) + CuCl2(aq) => ZnCl2(aq) + Cu(s).
Zn = 0
Cu = +2
Cl2 = -1(2) = -2
=>
Zn = +2
Cl2 = -1(2) = -2
Cu = 0
Zinc displaces, or replaces copper in the dissolved salt. Zn is oxidized to Zn2+. Cu2+ is reduced to Cu. When a metal is oxidized by an aqueous solution, it becomes an aqueous ion.
------------------------------------------------------------------------------------
The activity series is a list of metals (and hydrogen) arranged from top to bottom in order of decreasing ease of oxidation. Metals listed at the top are called active metals. Metals listed at the bottom are called noble metals. An element in the series will be oxidized by the ions of any element that appears below it in the table.
Ex: Zinc
Iron
Nickel
Hydrogen
Copper
Silver
Gold
Zn(s) + CuCl2(aq) => ZnCl2(aq) + Cu(s)
Cu(s) + ZnCl2(aq) => no reaction
------------------------------------------------------------------------------
In the following reaction series, which reactions will occur?
Barium
Sodium
Cobalt
Tin
Copper
Silver
Gold
Co(s) + BaI2(aq) = No reaction because cobalt is below barium
Sn(s) + CuBr2(aq) => Cu(s) + SnBr2(aq)
Ag(s) + NaCl(aq) => No reaction because silver is below sodium
Balancing Simple Redox Equations: Redox reactions must have both mass balance and charge balance:
Cr(s) + Ni2+(aq) => Cr3+(aq) + Ni(s)
Oxidation half-reaction: Cr(s) => Cr3+(aq) + 3e-
Reduction half-reaction: Ni2+(aq) + 2e- => Ni (s)
Before adding half-reactions, the electrons must balance.
Prior to adding the two half-reactions, balance the electrons.
2[Cr(s) => Cr3+(aq) + 3e-]
3[Ni2+(aq) +2e- => Ni(s)]
Oxidation half-reaction 2Cr(s) => 2Cr3+(aq) + 6e-
Reduction half reaction 3Ni2+(aq) + 6e- => 3 Ni(s)
------------------------------------------------------------
---------------------------->3Ni2+(aq) + 2Cr(s) => 3Ni(s) + 2Cr3+(aq)
This is the half reaction method.
Combination reactions are another type of Redox reactions. Combination reactions can involve oxidation and reduction:
N2(g) + 3H2(g) => 2NH3(g)
N2=0,0
H2=0,0
=>
N=-3, -3
H3=+1, +3
Hydrogen is oxidized from 0 to +1.
Nitrogen is reduced from 0 to -3.
Decomposition can also be a redox reaction.
NaH(s) => 2Na(s) + 3H2(g)
Na=+1, +1
H=-1, -1
=>
Na=0,0
H2=0,0
Na+ is reduced to Na
H- is oxidized to H2
Disproportionation reactions occur when one element undergoes both oxidation and reduction.
2H2O2(aq) => 2H2O(l) + O2(g)
H2=+1, +2
O2=-1, -2
=>
H2=+1, +2
O=-2, -2
O2=0,0
Oxygen in H2O2 (and other peroxides) has an oxidation number of -1.
Combustion is also a redox process.
CH4(g) + 2O2(g) => CO2 + 2H2O(l)
-4,+4 + 0,0 => +4, -4 + +2, -2
Molarity (M) or molar concentration is defined as the number of moles of solute per liter of solution. M is a measure of concentration of moles of solute per liter of solution.
molarity = (moles of solute)/(liters of solution L)
L=(mol)/(M)
Dilution is the process of preparing a less concentrated solution from a more concentrated one. The moles of solute before dilution = moles of solute after dilution.
Ex. Convert 1L of a 1.00 M solution of KMnO4 to a 0.4 M solution of KMnO4.
Mc x Lc = Md x Ld
(1.00 M KMnO4)(Lc) = (O.400 M KMnO4)(1.00 L)
Lc = O.400 L or 400 mL
To make the solution, pipet 400 mL of stock solution into a 1.00 L volumetric flask and carefully dilute to the calibration mark.
Dilution and Serial Dilution: Because most volumes in the laboratory are in millimeters rather than liters, the equation can be written as: Mc x mLc = Md x mLd
A series of dilutions that may be used to prepare a number of increasingly dilute solutions is called a Serial Dilution: Step 1) Prepare a dilute solution from the stock. Step 2) Dilute a portion of the prepared solution to make a more dilute solution. Step 3) Repeat as needed.
The pH Scale: 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
In pure water at 25 degrees C, [H3O+] = log 1.o x 10^-7
pH = -log(1.0 x 10^-7) = 7.00. pH is a dimensionless quantity. pH ranges from 1.00 to 14.00, where lower than 7.00 is considered increasingly acidic, 7.00 is considered neutral, and greater than 7.00 is considered increasingly basic pH.
pH values of some common fluids are stomach acid (1.0), lemon juice (2.0), vinegar (3.0), orange juice (3.5), urine (4.8-7.5), rainwater (5.5), saliva (6.4-6.9), milk (6.5), pure water (7.0), blood (7.4), tears (7.4), milk of magnesia (10.6), ammonia (11.5).
Aqueous Reactions and Chemical Analysis: Gravimetric analysis is an analytical technique based on the measurement of mass. Gravimetric analysis is highly accurate and it is applicable only to reactions that go to completion or have nearly 100 percent yield.
Acid-Base Titrations: Quantitative studies of acid-base neutralization reactions are most conveniently carried out using a technique known as a titration. A titration is a volumetric technique that uses burets. The point in the titration technique where the acid has been neutralized is called the equivalence point.
The equivalence point is usually signaled by a color change. The color change is brought about by the use of an indicator. Indicators have distinctly different colors in acidic and basic media. The indicator is chosen so that the color change, or endpoint, is very close to the equivalence point. Phenolphthalein is a common indicator.
Sodium hydroxide solutions are commonly used in titrations. NaOH solutions must be standardized as the concentrations change over time. (NaOH reacts with CO2 that slowly dissolves into the solution forming a carbonic acid).
The acid potassium hydrogen phthalate (KHP) is frequently used to standardize NaOH solutions. KHP is a monoprotic acid because it can donate only one proton (H+) per molecule, making it useful for standardizing solutions in titrations due to its chemical structure, which contains a single acidic hydrogen atom attached to the phthalate ion (HC8H4O4-). Monoprotic means that KHP can lose a single proton (H+) in a chemical reaction.
In a neutralization reaction with a base like sodium hydroxide (NaOH), KHP and NaOH react in a 1:1 mole ratio because both are monoprotic. Because it has a large molar mass and is a solid, KHP is frequently used as a primary standard to determine the concentration of a base solution in an acid-base titration technique.