Chapter 4: Reactions in Aqueous Solutions
Chapter 4: Reactions in Aqueous Solutions 4.1 – General Properties of Aqueous Solutions • • •
A solution is a homogenous mixture of two or more substances (solute and solvent). A solute is the substance present in a smaller amount and the solvent is the substance present in a larger amount. An aqueous solution is a solution in which the solute is a liquid or solid, and the solvent is water.
Electrolytic Properties • All solutes that can dissolve in water can be classified as electrolytes or nonelectrolytes. o An electrolyte is a solute that, when dissolved in water, results in a solution that can conduct electricity. o A nonelectrolyte is a solute that, when dissolved in water, results in a solution that cannot conduct electricity. • Pure water is a poor conductor of electricity since it does NOT contain any dissolved solutes. However, tap water contains many dissolved solutes and is able to conduct electricity well. • When a light bulb, power source, and two electrodes are immersed in an aqueous solution, the light bulb can turn on brightly, dimly, or not even turn on, depending on the type of solute dissolved in the solution. o Strong electrolytes such as ionic compounds, strong acids, or strong bases dissociate or ionize completely (100%) when dissolved in water. When dissociated, their ions undergo hydration. Strong electrolytes cause the light bulb to shine brightly with light. When NaCl is dissolved in H2O, the H2O molecules cause NaCl to dissociate completely into its ions and then the H2O molecules surround each ion in a specific manner (O ends surround Na+ and H ends surround Cl-). This process is called hydration and prevents the Na+ and Cl- from reacting with one another and reforming the NaCl. Then, Na+ get attracted to the negative electrode and the Cl- get attracted to the positive electrode. Eventually, a pathway of ions will form from one electrode to the other electrode, allowing the current flow through the entire circuit. That lights up the bulb. o Weak electrolytes such as weak acids and weak bases do not completely ionize when dissolved in water. They cause the light bulb to be dimly lit. When HF is dissolved in H2O, it does not completely ionize into H+ and F-.This provides a weak pathway for the flow of electrons between the electrodes, causing the light bulb to be dimly lit. HF is not completely ionized since the ionization of HF stops when equilibrium is reached. That is when the rate of ionization is equal to the rate of HF reforming. HF ↔ H+ + Fo Nonelectrolytes such as molecular compounds do not dissociate into ions When molecular compounds such as methanol (CH3OH) are placed in water, hydrogen bonds form between the O and H atoms of the hydroxyl group and the O and H atoms of the H2O molecules, causing the molecule to dissolve in water. However, since molecular compounds do not dissociate, their solutions do not conduct electricity.
Weak acids/bases have reversible equations (↔), whereas strong acids/bases do not (→). http://highered.mcgraw-
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4.2 – Precipitation Reactions •
Precipitation reactions occur in an aqueous solution and result in the formation of a precipitate.
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A precipitate is an insoluble solid that separates from the solution.
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Many precipitation reactions involve ionic compounds and are double displacement reactions. o
Double displacement reactions can be called metathesis reactions.
Solubility •
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Solubility of a solute is the maximum amount of solute that will dissolve in a given amount of solvent at a given temperature. o
Chemists refer to solutes as soluble, slightly soluble, or insoluble.
o
Insoluble solutes are still soluble, but not that much as slightly soluble or soluble solutes.
The solubility table below can be used to determine if a compound is soluble or insoluble. o
Look at anions first, and then the cation soluble/insoluble exceptions.
Molecular Equations, Ionic Equations, and Net Ionic Equations •
A molecular equation is an equation in which the formulas of all the compounds are written as though all the species existed as molecules or whole units. o
Pb(NO3)2 (aq) + 2KI(aq) → PbI2 (s) + 2KNO3 (aq) (Molecular Equation)
•
Molecular equations are very useful when performing an experiment as they tell us the compounds that are reacting.
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An ionic equation is an equation that shows the dissolved ionic compounds as ions. The precipitate occurs in a solid state, not dissolved, in the reaction. o
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The spectator ions are the ions that do not participate in the overall reaction. o
•
•
Pb2+(aq) + 2NO3- (aq) + 2K+(aq) + 2I-(aq) → PbI2 (s) + 2NO3- (aq) + 2K+(aq) (Ionic Equation)
They exist as ions in the solution, before the reaction starts and after it finishes.
The net ionic equation is an equation that only shows the species that take part in the reaction. It is obtained by cancelling the spectator ions in the ionic equation. o
Pb2+(aq) + 2NO3- (aq) + 2K+(aq) + 2I-(aq) → PbI2 (s) + 2NO3- (aq) + 2K+(aq)
o
Pb2+(aq) + 2I-(aq) → PbI2 (s) (Net Ionic Equation)
Below is a diagramic representation of this reaction:
4.3
– Acids-Base Reactions
Acids and Bases •
Svante Arrhenius defined acids as substances that ionizes in water to produce H+ and bases as substances that ionize in water to produce OH-.
Acids • • • • •
Acids taste sour (i.e. lemons, vinegar) Acids cause colour changes in plant dyes. They change the colour of litmus from blue to red. Aqueous acid solutions conduct electricity. Acids react with metals above H on the activity series (Mg, Zn, Fe) to produce H2 (g). o 2HCl(aq) + Mg(s) → MgCl2(s) + H2 (g) Acids react with metal carbonates and metal bicarbonates to produce H2O (l) and CO2 (g). o Na2CO3(s) + 2HCl(aq) →2 NaCl(s) + H2O(l) + CO2 (g)
Bases • • • •
Bases taste bitter. Bases feel slippery (i.e. soap) Bases cause colour changes in plant dyes. They change the colour of litmus from red to blue. Aqueous base solutions conduct electricity.
Bronstead Acids and Bases • Bronstead defined acids as proton donors, and bases as proton acceptors. o
•
Unlike Arrhenius definition, Bronstead definition showed that acids and bases are not limited to aqueous states.
When an acid is placed in water, it ionizes and forms an anion and H+. In water, the H+ become hydrated by H2O molecules. o
Therefore, in water, H+ attract to many H2O molecules, forming many ions of the form H(H2O)n-, where n represents the amount of water molecules the H+ is attracted to or hydrated by.
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H3O+ are hydrated H+ and are called hydronium ions.
Monoprotic Acids are acids that ionize into only one H+ in water. o
Since the ionization of acetic acid in water is a reverse reaction, it does not completely ionize into H+ (double arrow).
o
This is why acetic acid is a weak acid and a weak electrolyte.
Since HCl (aq) and HNO3 (aq) ionize completely in water, they are strong acids and strong electrolytes (single arrow).
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Diprotic Acids are acids that ionize into two H+ in water. o
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H2SO4 (aq) is a strong acid and electrolyte. However HSO4- (aq) is a weak acid electrolyte.
Triprotic Acids are acids that ionize into three H+ in water. o
H3PO4 (aq), H2PO4- (aq) and HPO42- (aq) are all weak acids and electrolytes.
Base
Acid
Conjugate Acid
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Since OH- can accept a H+ and form H2O, it is a base (conjugate base).
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Since NH4+ can donate a H+ and form NH3, it is an acid (conjugate acid).
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All alkali metal hydroxides are strong bases and electrolytes. o
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Conjugate Base
Ba(OH)2 is the only alkaline earth metal hydroxide that is soluble in water.
NH3 is a weak base because it reacts with H2O to form very few NH4+ and OH-.
Acid-Base Neutralizations •
A neutralization reaction is a reaction between an acid and a base. It usually produces a salt and water.
Strong Acid and Strong Base Neutralization Reactions Chemical Equation: NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l) The strong acids and bases ionize completely in water.
Total Ionic Equation: Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq) → Na+(aq) + Cl-(aq) + H2O(l) Na+(aq) and Cl-(aq) are spectator ions.
Net Ionic Equation: OH-(aq) + H+(aq) → H2O(l) •
If we started with equal molar amounts of acid and base, we would end up with no excess acid or base.
Weak Acid and Strong Base Neutralization Reactions Chemical Equation: NaOH(aq) + HCN(aq) → NaCN(aq) + H2O(l) HCN(aq) is a weak acid and does not ionize completely in H2O.
Total Ionic Equation: Na+(aq) + OH-(aq) + HCN(aq) → Na+(aq) + Cl-(aq) + H2O(l) Na+(aq) is the spectator ions.
Net Ionic Equation: OH-(aq) + HCN(aq) → Cl-(aq) + H2O(l) HNO3 (aq) + NH3 (aq) → NH4NO3 (aq) •
Since NH4NO3 is in an aqueous state, it consists of NH4+ and NO3-. The NH4+ is composed of NH3 and H2O. o
Therefore, water is also produced in this neutralization reaction.
o
HNO3 (aq) + NH3 (aq) → NH3 (aq) + H2O(l) + NO3 (aq)
Acid-Base Reactions Leading To Gas Formation •
Certain salts such as carbonates, bicarbonates, sulphites and sulphides react with acids to form gaseous products.
4 .4 – Reduction Reactions •
Oxidation-
Oxidation-reduction reactions (also called redox reactions) are reactions that involve the transfer of electrons. o
Many redox reactions take place in water (aqueous states), but not all do.
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Half reactions are equations that show the electrons involved in a redox reaction. Every redox reaction consists of an oxidation half reaction and a reduction half reaction. o
Oxidation half reactions show the loss of electrons in the redox reaction.
o
Reduction half reactions show the gain of electrons in the redox reaction.
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Adding both half reactions gives us the chemical equation.
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The reducing agent is the reactant that causes the reduction of another reactant in the reaction. In the process, it gets oxidized.
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The oxidizing agent is the reactant that causes the oxidation of another reactant in the reaction. It gets reduced in the process. Chemical Equation: 2Mg(s) + O2 (g) → 2MgO(s) o In the reaction below, Mg(s) is the reducing agent Oxidation Half Reaction: 2Mg(s) → 2Mg2+(s) + 4eand O2(g) is the oxidizing agent. Reduction Half Reaction: O2 (g) + 4e- → 2O2-(s) o
Mg(s) is oxidized and O2(g) reduced. e- lost by reducing agent = e- gained by oxidizing agent
Chemical Equation: Zn (s) + CuSO4 (aq) → ZnSO4 (aq) + Cu (s) Net Ionic Equation: Zn (s) + Cu2+ (aq) → Zn2+ (aq) + Cu (s) Oxidation Half Reaction: Zn (s) → Zn2+ (aq) + 2eReduction Half Reaction: Cu2+ (aq) + 2e- → Cu (s) • •
• •
Zn (s) is the reducing agent. It gets oxidized. CuSO4 (aq) is the oxidizing agent. It gets reduced (Cu2+ (aq))
In the formation of ionic compounds (such as MgO), electrons are fully transferred from the metal to the non-metal. However, molecular compounds do not have a full transfer of electrons. Instead, they can only have partial transfer of electrons since the more electronegative atom pulls electrons towards itself and away from the other atom it is bonding with. o Oxidation numbers (also called oxidation states) show the charge each atom in a molecular compound would have if there was a full transfer of electrons. The change in oxidation numbers of an element represents the number of electrons transferred during the reaction. In a reaction oxidation numbers can be used to determine which elements are being oxidized and which ones are being reduced. o The elements that show an increase in oxidation number are oxidized and the elements that show a decrease in oxidation number are reduced.
Oxidation Number Rules • In free elements (uncombined state), each atom has an oxidization number of 0. o Na, H2, Br2, K, O2 and P4 all have an oxidation number of zero. • Monatomic ions have an oxidation number equal to the charge on their ion. o Na+ = +1 , Be2+ = +2 , Cl- = -1 , S2- = -2 • The oxidation number of O in most compounds is -2. However, in hydrogen peroxide (H2O2) and in the peroxide ion (O22-), it is -1. In OF2, it is +2.
• • •
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The oxidation number of hydrogen always is +1, except for the H in metal hydrides (NaH, LiH, CaH2). In metal hydrides, it is -1. Fluorine has an oxidation number of -1 in all its compounds. All halogens have negative oxidation numbers in their compounds. In a neutral molecule, the sum of the oxidation numbers of all the atoms is 0. In an ion, the sum of the oxidation numbers of all the atoms is the charge of the ion. o In NH4+, N = -3 and H = +1 (-3 + 4(+1) = +1). Oxidation numbers do not have to be integers. o The superoxide ion (O2- has an oxidation number of -1/2.
***Always assign oxidation numbers in a compound starting with the most electronegative atom. *** • •
The highest oxidation number of any element in Groups 1A-7A is its group number. o The highest possible oxidation number of halogens is +7. Metals can only have + oxidation numbers. Non-metals can have + or – oxidation numbers.
Types of Redox Reactions • The most common types of redox reactions are combination, decomposition, combustion, displacement, and disproportionation reactions. o However, not all combination and decomposition reactions are also redox reactions. Combination Reactions • A combination reaction is a reaction in which two or more substances combine to form a single product.
Decomposition Reactions • A decomposition reaction is a reaction in which a single compound breaks down into two or more components.
Combustion Reactions • A combustion reaction is a reaction in which a substance reacts with oxygen, usually with the release of heat and light to produce a flame. C3H8 (g) + 5O2 (g) → H2O (g) + CO2 (g) • The combustion of propane is a redox reaction since the oxidation number of O changes from 0 to -2. Displacement Reactions • In a displacement reaction, an ion/atom in a compound is replaced by an atom/ion of another element. • Only metals that are above hydrogen in the activity series will displace it from both water and acids. o
The metals that displace H from water are alkali metals and some alkali earth metals (Ca, Sr, Ba)
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These metals are the most reactive metals (hydrogen displacement)
Any metal in the activity series can only displace metals (in compounds) that are below it (metal displacement) o
Zn (s) + CuSO4 (aq) → ZnSO4 (aq) + Cu (s)
The reducing agent (Zn(s)) is lower than the oxidizing agent (Cu2+(aq)) in the activity series.
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Only halogens higher in the activity series than the halogen in a compound can displace the halogen from a compound (halogen displacement)
Halogen Activity Series
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Halides (KBr) can be converted back into halogens (Br2) by reacting it with a halogen that is a stronger oxidizing agent (more reactive).
Disproportion Reactions •
Disproportionation reactions are redox reactions in which one element (with one oxidation number) is being simultaneously reduced and oxidized. o
•
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There are at least 3 oxidation numbers for the element undergoing oxidation and reduction. The highest and the lowest oxidation numbers exist for that element in the products. The intermediate oxidation number exists in the reactants.
In this reaction, the O atom in the reactant has an oxidation number of -1, whereas the O in the products has oxidation numbers of 0 and -2. o
Reduction occurs when H2O is formed (-1 → -2)
o
Oxidation occurs when O2 is formed (-1 → 0)
The H atom does not undergo a transfer of electrons.
4.4 •
– Concentration of Solutions The concentration of a solution is the amount of solute present in a given amount of solvent or solution. o
The unit for concentration of a solution is molarity (M) or molar concentration. Molarity (or molar concentration) is the number of moles of solute per litre of solution (mol/L or M). M=
n V
n represents the number of moles of solute in the solution. V represents the volume of the solution in litres.
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Concentration is an intensive property because its value does not depend on the amount of solution present.
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In a 1M KCl solution, [K+(aq)] = 1M and [Cl-(aq)] = 1M.
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In a 1M Ba(NO3)2 solution, [Ba2+(aq)] = 1M and [NO3-] = 2M
Steps to Prepare a Solution of Known Molarity 1. The solute is accurately weighed and transferred to a volumetric flask using a funnel. 2. Then, water is added to the flask and swirled until all the solute is dissolved in it. 3. After all the solute dissolves in the water (solvent), more water is added to the solution until the desired
volume of solution is achieved. •
Stock solutions are concentrated solutions.
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Dilution is the procedure for producing a less concentrated solution from a more concentrated one.
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Adding more solvent to a stock solution reduces the concentration of the solute in the solution without changing the number of moles of solute present in the solution.
Number of Moles of Solute Before Dilution=Number of Moles of Solute After Dilution M 1 V 1=M 2 V 2 This equation is used to calculate the final volume, V2 of a stock solution after it has been diluted to a molarity M2 from a molarity, M1 and an initial volume, V1.
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Quantitative analysis is the determination of the amount (moles) or concentration (molar) of a substance in a sample. o
4.6 •
Examples of quantitative analysis are gravimetric analysis and titrations.
– Gravimetric Analysis Gravimetric Analysis is a technique used to perform the qualitative analysis of a substance. It involves the formation, isolation and the determination of the mass of a precipitate.
Steps Involved In One Type of Gravitational Analysis 1. A sample substance whose % composition is unknown is dissolved in H2O and then reacted with another
substance to form a precipitate. 2. Then, the precipitate is filtered, dried and weighed (mass determination). 3. Using the measured mass of the precipitate, as well its chemical formula, the mass the chemical
component that is present in both the original sample and the precipitate can be measured. 4. Then, the mass of this chemical component and the mass of the original sample can be used to determine
the % composition of this chemical component in the original sample. •
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This procedure can only be used if all the ions in the original sample were removed from the solution and accumulated in the precipitate. o
For example if NaCl (aq) reacts with AgNO3 (aq) and precipitates into AgCl, all the Cl- ions from the NaCl(aq) MUST be removed from the solution and accumulate as precipitate for gravimetric analysis to be used to determine the % composition of Cl in NaCl(aq).
o
This also means that the original sample (NaCl (aq)) has to be the limiting reactant. If it is not, all the Cl- may not accumulate as a precipitate; some may remain in the solution in an aqueous state.
Using gravimetric analysis, the unknown ion in a sample with an unknown ion can be determined. o
4.7 •
Usually, the samples used are ionic compounds.
– Acid-Base Titrations
Quantitative studies of acid-base neutralization reactions are carried out using a technique called titration. o In a titration, a standard solution is gradually added to another solution of an unknown concentration, until the chemical reaction between both solutions is complete. A standard solution is a solution whose concentration is known. Steps for an Acid-Base Titration
1. A known amount of acid is transferred to an Erlenmeyer flask and mixed with distilled water to form a
solution of known concentration. 2. Then, an acid-base indicator, such as phenolphthalein is added to the acid. Acid-base indicators are substances that have different colours in acidic and in basic solutions. The phenolphthalein indicator remains colourless in acidic and in neutral solutions, but changes colour to reddish pink in basic solutions. 3. After a buret containing a base is placed over the Erlenmeyer flask, some of the base is released slowly into the acid. At the beginning, the solution in the Erlenmeyer flask will not change colour since there is more acid present in it than base. When the amount of acid and base in the solution are equal, the acid completely reacts with base, reaching the equivalence point. However, when one more drop of base is added to the solution, the amount of base will exceed the amount of acid, causing the solution to turn pink. Strong acid plus strong base = pH 7 at equivalence Weak acid plus strong base = pH >7 at equivalence Strong acid plus weak base = pH
A solution is a homogenous mixture of two or more substances (solute and solvent). A solute is the substance present in a smaller amount and the solvent is the substance present in a larger amount. An aqueous solution is a solution in which the solute is a liquid or solid, and the solvent is water.
Electrolytic Properties • All solutes that can dissolve in water can be classified as electrolytes or nonelectrolytes. o An electrolyte is a solute that, when dissolved in water, results in a solution that can conduct electricity. o A nonelectrolyte is a solute that, when dissolved in water, results in a solution that cannot conduct electricity. • Pure water is a poor conductor of electricity since it does NOT contain any dissolved solutes. However, tap water contains many dissolved solutes and is able to conduct electricity well. • When a light bulb, power source, and two electrodes are immersed in an aqueous solution, the light bulb can turn on brightly, dimly, or not even turn on, depending on the type of solute dissolved in the solution. o Strong electrolytes such as ionic compounds, strong acids, or strong bases dissociate or ionize completely (100%) when dissolved in water. When dissociated, their ions undergo hydration. Strong electrolytes cause the light bulb to shine brightly with light. When NaCl is dissolved in H2O, the H2O molecules cause NaCl to dissociate completely into its ions and then the H2O molecules surround each ion in a specific manner (O ends surround Na+ and H ends surround Cl-). This process is called hydration and prevents the Na+ and Cl- from reacting with one another and reforming the NaCl. Then, Na+ get attracted to the negative electrode and the Cl- get attracted to the positive electrode. Eventually, a pathway of ions will form from one electrode to the other electrode, allowing the current flow through the entire circuit. That lights up the bulb. o Weak electrolytes such as weak acids and weak bases do not completely ionize when dissolved in water. They cause the light bulb to be dimly lit. When HF is dissolved in H2O, it does not completely ionize into H+ and F-.This provides a weak pathway for the flow of electrons between the electrodes, causing the light bulb to be dimly lit. HF is not completely ionized since the ionization of HF stops when equilibrium is reached. That is when the rate of ionization is equal to the rate of HF reforming. HF ↔ H+ + Fo Nonelectrolytes such as molecular compounds do not dissociate into ions When molecular compounds such as methanol (CH3OH) are placed in water, hydrogen bonds form between the O and H atoms of the hydroxyl group and the O and H atoms of the H2O molecules, causing the molecule to dissolve in water. However, since molecular compounds do not dissociate, their solutions do not conduct electricity.
Weak acids/bases have reversible equations (↔), whereas strong acids/bases do not (→). http://highered.mcgraw-
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4.2 – Precipitation Reactions •
Precipitation reactions occur in an aqueous solution and result in the formation of a precipitate.
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A precipitate is an insoluble solid that separates from the solution.
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Many precipitation reactions involve ionic compounds and are double displacement reactions. o
Double displacement reactions can be called metathesis reactions.
Solubility •
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Solubility of a solute is the maximum amount of solute that will dissolve in a given amount of solvent at a given temperature. o
Chemists refer to solutes as soluble, slightly soluble, or insoluble.
o
Insoluble solutes are still soluble, but not that much as slightly soluble or soluble solutes.
The solubility table below can be used to determine if a compound is soluble or insoluble. o
Look at anions first, and then the cation soluble/insoluble exceptions.
Molecular Equations, Ionic Equations, and Net Ionic Equations •
A molecular equation is an equation in which the formulas of all the compounds are written as though all the species existed as molecules or whole units. o
Pb(NO3)2 (aq) + 2KI(aq) → PbI2 (s) + 2KNO3 (aq) (Molecular Equation)
•
Molecular equations are very useful when performing an experiment as they tell us the compounds that are reacting.
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An ionic equation is an equation that shows the dissolved ionic compounds as ions. The precipitate occurs in a solid state, not dissolved, in the reaction. o
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The spectator ions are the ions that do not participate in the overall reaction. o
•
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Pb2+(aq) + 2NO3- (aq) + 2K+(aq) + 2I-(aq) → PbI2 (s) + 2NO3- (aq) + 2K+(aq) (Ionic Equation)
They exist as ions in the solution, before the reaction starts and after it finishes.
The net ionic equation is an equation that only shows the species that take part in the reaction. It is obtained by cancelling the spectator ions in the ionic equation. o
Pb2+(aq) + 2NO3- (aq) + 2K+(aq) + 2I-(aq) → PbI2 (s) + 2NO3- (aq) + 2K+(aq)
o
Pb2+(aq) + 2I-(aq) → PbI2 (s) (Net Ionic Equation)
Below is a diagramic representation of this reaction:
4.3
– Acids-Base Reactions
Acids and Bases •
Svante Arrhenius defined acids as substances that ionizes in water to produce H+ and bases as substances that ionize in water to produce OH-.
Acids • • • • •
Acids taste sour (i.e. lemons, vinegar) Acids cause colour changes in plant dyes. They change the colour of litmus from blue to red. Aqueous acid solutions conduct electricity. Acids react with metals above H on the activity series (Mg, Zn, Fe) to produce H2 (g). o 2HCl(aq) + Mg(s) → MgCl2(s) + H2 (g) Acids react with metal carbonates and metal bicarbonates to produce H2O (l) and CO2 (g). o Na2CO3(s) + 2HCl(aq) →2 NaCl(s) + H2O(l) + CO2 (g)
Bases • • • •
Bases taste bitter. Bases feel slippery (i.e. soap) Bases cause colour changes in plant dyes. They change the colour of litmus from red to blue. Aqueous base solutions conduct electricity.
Bronstead Acids and Bases • Bronstead defined acids as proton donors, and bases as proton acceptors. o
•
Unlike Arrhenius definition, Bronstead definition showed that acids and bases are not limited to aqueous states.
When an acid is placed in water, it ionizes and forms an anion and H+. In water, the H+ become hydrated by H2O molecules. o
Therefore, in water, H+ attract to many H2O molecules, forming many ions of the form H(H2O)n-, where n represents the amount of water molecules the H+ is attracted to or hydrated by.
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H3O+ are hydrated H+ and are called hydronium ions.
Monoprotic Acids are acids that ionize into only one H+ in water. o
Since the ionization of acetic acid in water is a reverse reaction, it does not completely ionize into H+ (double arrow).
o
This is why acetic acid is a weak acid and a weak electrolyte.
Since HCl (aq) and HNO3 (aq) ionize completely in water, they are strong acids and strong electrolytes (single arrow).
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Diprotic Acids are acids that ionize into two H+ in water. o
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H2SO4 (aq) is a strong acid and electrolyte. However HSO4- (aq) is a weak acid electrolyte.
Triprotic Acids are acids that ionize into three H+ in water. o
H3PO4 (aq), H2PO4- (aq) and HPO42- (aq) are all weak acids and electrolytes.
Base
Acid
Conjugate Acid
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Since OH- can accept a H+ and form H2O, it is a base (conjugate base).
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Since NH4+ can donate a H+ and form NH3, it is an acid (conjugate acid).
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All alkali metal hydroxides are strong bases and electrolytes. o
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Conjugate Base
Ba(OH)2 is the only alkaline earth metal hydroxide that is soluble in water.
NH3 is a weak base because it reacts with H2O to form very few NH4+ and OH-.
Acid-Base Neutralizations •
A neutralization reaction is a reaction between an acid and a base. It usually produces a salt and water.
Strong Acid and Strong Base Neutralization Reactions Chemical Equation: NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l) The strong acids and bases ionize completely in water.
Total Ionic Equation: Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq) → Na+(aq) + Cl-(aq) + H2O(l) Na+(aq) and Cl-(aq) are spectator ions.
Net Ionic Equation: OH-(aq) + H+(aq) → H2O(l) •
If we started with equal molar amounts of acid and base, we would end up with no excess acid or base.
Weak Acid and Strong Base Neutralization Reactions Chemical Equation: NaOH(aq) + HCN(aq) → NaCN(aq) + H2O(l) HCN(aq) is a weak acid and does not ionize completely in H2O.
Total Ionic Equation: Na+(aq) + OH-(aq) + HCN(aq) → Na+(aq) + Cl-(aq) + H2O(l) Na+(aq) is the spectator ions.
Net Ionic Equation: OH-(aq) + HCN(aq) → Cl-(aq) + H2O(l) HNO3 (aq) + NH3 (aq) → NH4NO3 (aq) •
Since NH4NO3 is in an aqueous state, it consists of NH4+ and NO3-. The NH4+ is composed of NH3 and H2O. o
Therefore, water is also produced in this neutralization reaction.
o
HNO3 (aq) + NH3 (aq) → NH3 (aq) + H2O(l) + NO3 (aq)
Acid-Base Reactions Leading To Gas Formation •
Certain salts such as carbonates, bicarbonates, sulphites and sulphides react with acids to form gaseous products.
4 .4 – Reduction Reactions •
Oxidation-
Oxidation-reduction reactions (also called redox reactions) are reactions that involve the transfer of electrons. o
Many redox reactions take place in water (aqueous states), but not all do.
•
Half reactions are equations that show the electrons involved in a redox reaction. Every redox reaction consists of an oxidation half reaction and a reduction half reaction. o
Oxidation half reactions show the loss of electrons in the redox reaction.
o
Reduction half reactions show the gain of electrons in the redox reaction.
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Adding both half reactions gives us the chemical equation.
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The reducing agent is the reactant that causes the reduction of another reactant in the reaction. In the process, it gets oxidized.
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The oxidizing agent is the reactant that causes the oxidation of another reactant in the reaction. It gets reduced in the process. Chemical Equation: 2Mg(s) + O2 (g) → 2MgO(s) o In the reaction below, Mg(s) is the reducing agent Oxidation Half Reaction: 2Mg(s) → 2Mg2+(s) + 4eand O2(g) is the oxidizing agent. Reduction Half Reaction: O2 (g) + 4e- → 2O2-(s) o
Mg(s) is oxidized and O2(g) reduced. e- lost by reducing agent = e- gained by oxidizing agent
Chemical Equation: Zn (s) + CuSO4 (aq) → ZnSO4 (aq) + Cu (s) Net Ionic Equation: Zn (s) + Cu2+ (aq) → Zn2+ (aq) + Cu (s) Oxidation Half Reaction: Zn (s) → Zn2+ (aq) + 2eReduction Half Reaction: Cu2+ (aq) + 2e- → Cu (s) • •
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Zn (s) is the reducing agent. It gets oxidized. CuSO4 (aq) is the oxidizing agent. It gets reduced (Cu2+ (aq))
In the formation of ionic compounds (such as MgO), electrons are fully transferred from the metal to the non-metal. However, molecular compounds do not have a full transfer of electrons. Instead, they can only have partial transfer of electrons since the more electronegative atom pulls electrons towards itself and away from the other atom it is bonding with. o Oxidation numbers (also called oxidation states) show the charge each atom in a molecular compound would have if there was a full transfer of electrons. The change in oxidation numbers of an element represents the number of electrons transferred during the reaction. In a reaction oxidation numbers can be used to determine which elements are being oxidized and which ones are being reduced. o The elements that show an increase in oxidation number are oxidized and the elements that show a decrease in oxidation number are reduced.
Oxidation Number Rules • In free elements (uncombined state), each atom has an oxidization number of 0. o Na, H2, Br2, K, O2 and P4 all have an oxidation number of zero. • Monatomic ions have an oxidation number equal to the charge on their ion. o Na+ = +1 , Be2+ = +2 , Cl- = -1 , S2- = -2 • The oxidation number of O in most compounds is -2. However, in hydrogen peroxide (H2O2) and in the peroxide ion (O22-), it is -1. In OF2, it is +2.
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The oxidation number of hydrogen always is +1, except for the H in metal hydrides (NaH, LiH, CaH2). In metal hydrides, it is -1. Fluorine has an oxidation number of -1 in all its compounds. All halogens have negative oxidation numbers in their compounds. In a neutral molecule, the sum of the oxidation numbers of all the atoms is 0. In an ion, the sum of the oxidation numbers of all the atoms is the charge of the ion. o In NH4+, N = -3 and H = +1 (-3 + 4(+1) = +1). Oxidation numbers do not have to be integers. o The superoxide ion (O2- has an oxidation number of -1/2.
***Always assign oxidation numbers in a compound starting with the most electronegative atom. *** • •
The highest oxidation number of any element in Groups 1A-7A is its group number. o The highest possible oxidation number of halogens is +7. Metals can only have + oxidation numbers. Non-metals can have + or – oxidation numbers.
Types of Redox Reactions • The most common types of redox reactions are combination, decomposition, combustion, displacement, and disproportionation reactions. o However, not all combination and decomposition reactions are also redox reactions. Combination Reactions • A combination reaction is a reaction in which two or more substances combine to form a single product.
Decomposition Reactions • A decomposition reaction is a reaction in which a single compound breaks down into two or more components.
Combustion Reactions • A combustion reaction is a reaction in which a substance reacts with oxygen, usually with the release of heat and light to produce a flame. C3H8 (g) + 5O2 (g) → H2O (g) + CO2 (g) • The combustion of propane is a redox reaction since the oxidation number of O changes from 0 to -2. Displacement Reactions • In a displacement reaction, an ion/atom in a compound is replaced by an atom/ion of another element. • Only metals that are above hydrogen in the activity series will displace it from both water and acids. o
The metals that displace H from water are alkali metals and some alkali earth metals (Ca, Sr, Ba)
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These metals are the most reactive metals (hydrogen displacement)
Any metal in the activity series can only displace metals (in compounds) that are below it (metal displacement) o
Zn (s) + CuSO4 (aq) → ZnSO4 (aq) + Cu (s)
The reducing agent (Zn(s)) is lower than the oxidizing agent (Cu2+(aq)) in the activity series.
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Only halogens higher in the activity series than the halogen in a compound can displace the halogen from a compound (halogen displacement)
Halogen Activity Series
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Halides (KBr) can be converted back into halogens (Br2) by reacting it with a halogen that is a stronger oxidizing agent (more reactive).
Disproportion Reactions •
Disproportionation reactions are redox reactions in which one element (with one oxidation number) is being simultaneously reduced and oxidized. o
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There are at least 3 oxidation numbers for the element undergoing oxidation and reduction. The highest and the lowest oxidation numbers exist for that element in the products. The intermediate oxidation number exists in the reactants.
In this reaction, the O atom in the reactant has an oxidation number of -1, whereas the O in the products has oxidation numbers of 0 and -2. o
Reduction occurs when H2O is formed (-1 → -2)
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Oxidation occurs when O2 is formed (-1 → 0)
The H atom does not undergo a transfer of electrons.
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– Concentration of Solutions The concentration of a solution is the amount of solute present in a given amount of solvent or solution. o
The unit for concentration of a solution is molarity (M) or molar concentration. Molarity (or molar concentration) is the number of moles of solute per litre of solution (mol/L or M). M=
n V
n represents the number of moles of solute in the solution. V represents the volume of the solution in litres.
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Concentration is an intensive property because its value does not depend on the amount of solution present.
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In a 1M KCl solution, [K+(aq)] = 1M and [Cl-(aq)] = 1M.
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In a 1M Ba(NO3)2 solution, [Ba2+(aq)] = 1M and [NO3-] = 2M
Steps to Prepare a Solution of Known Molarity 1. The solute is accurately weighed and transferred to a volumetric flask using a funnel. 2. Then, water is added to the flask and swirled until all the solute is dissolved in it. 3. After all the solute dissolves in the water (solvent), more water is added to the solution until the desired
volume of solution is achieved. •
Stock solutions are concentrated solutions.
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Dilution is the procedure for producing a less concentrated solution from a more concentrated one.
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Adding more solvent to a stock solution reduces the concentration of the solute in the solution without changing the number of moles of solute present in the solution.
Number of Moles of Solute Before Dilution=Number of Moles of Solute After Dilution M 1 V 1=M 2 V 2 This equation is used to calculate the final volume, V2 of a stock solution after it has been diluted to a molarity M2 from a molarity, M1 and an initial volume, V1.
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Quantitative analysis is the determination of the amount (moles) or concentration (molar) of a substance in a sample. o
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Examples of quantitative analysis are gravimetric analysis and titrations.
– Gravimetric Analysis Gravimetric Analysis is a technique used to perform the qualitative analysis of a substance. It involves the formation, isolation and the determination of the mass of a precipitate.
Steps Involved In One Type of Gravitational Analysis 1. A sample substance whose % composition is unknown is dissolved in H2O and then reacted with another
substance to form a precipitate. 2. Then, the precipitate is filtered, dried and weighed (mass determination). 3. Using the measured mass of the precipitate, as well its chemical formula, the mass the chemical
component that is present in both the original sample and the precipitate can be measured. 4. Then, the mass of this chemical component and the mass of the original sample can be used to determine
the % composition of this chemical component in the original sample. •
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This procedure can only be used if all the ions in the original sample were removed from the solution and accumulated in the precipitate. o
For example if NaCl (aq) reacts with AgNO3 (aq) and precipitates into AgCl, all the Cl- ions from the NaCl(aq) MUST be removed from the solution and accumulate as precipitate for gravimetric analysis to be used to determine the % composition of Cl in NaCl(aq).
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This also means that the original sample (NaCl (aq)) has to be the limiting reactant. If it is not, all the Cl- may not accumulate as a precipitate; some may remain in the solution in an aqueous state.
Using gravimetric analysis, the unknown ion in a sample with an unknown ion can be determined. o
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Usually, the samples used are ionic compounds.
– Acid-Base Titrations
Quantitative studies of acid-base neutralization reactions are carried out using a technique called titration. o In a titration, a standard solution is gradually added to another solution of an unknown concentration, until the chemical reaction between both solutions is complete. A standard solution is a solution whose concentration is known. Steps for an Acid-Base Titration
1. A known amount of acid is transferred to an Erlenmeyer flask and mixed with distilled water to form a
solution of known concentration. 2. Then, an acid-base indicator, such as phenolphthalein is added to the acid. Acid-base indicators are substances that have different colours in acidic and in basic solutions. The phenolphthalein indicator remains colourless in acidic and in neutral solutions, but changes colour to reddish pink in basic solutions. 3. After a buret containing a base is placed over the Erlenmeyer flask, some of the base is released slowly into the acid. At the beginning, the solution in the Erlenmeyer flask will not change colour since there is more acid present in it than base. When the amount of acid and base in the solution are equal, the acid completely reacts with base, reaching the equivalence point. However, when one more drop of base is added to the solution, the amount of base will exceed the amount of acid, causing the solution to turn pink. Strong acid plus strong base = pH 7 at equivalence Weak acid plus strong base = pH >7 at equivalence Strong acid plus weak base = pH
