4:Types Chapter of Chemical Reactions A. Dissolution Reactions B. Precipitation Reactions Solution (A homogeneous mixture of different substances)
(Solvent other than water) Solvent: Component which is in larger amount
Solute: component which is in smaller amount sugar
Aqueous Solution (Water is used as solvent)
Non-aqueous solution
Aqueous solution of sugar with water as solvent water sugar as solute Dissolution is the process where one substance disperses into another at molecular or atomic level. The solid structure is broken and individual molecules or ions are dispersed.
Dissolution When ionic compounds are dissolved in water, each ion is surrounded by water molecules due to ion-dipole interactions. NaCl(s)
H2O
Na+(aq) + Cl-(aq)
Ions are said to be aquated or solvated. When polar molecular (covalent) compounds are dissolved in water, the molecules of compound stay intact (do not break up) but are surrounded by water molecules due to dipole-dipole or hydrogen bonding interactions.
C6H12O6(s) C6H12O6(aq) The molecular compound is said to be aquated or solvated. (aq) indicates dissolution in water (or aqueous solution )
H2O
Solutions
Electrolyte solutions (A solution that conducts current more than the solvent)
Non- or weak electrolyte solutions: Addition of a substance to solvent does not increase, or slightly increases, the conductance
Strong electrolyte:If a compound, upon addition to solvent, dissociates
Example: NaCl(s)
water
completely into constituent ions then that substance is a strong electrolyte. Na+(aq) + Cl-(aq)
chloride ions solvated by water molecules Ion-ion interactions in solid NaCl are broken up in water and replaced by H ion-water interactions H O H water O Na+ O Na+ ClH Cl- H H Ionic compounds (made up of Group I and II elements) are strong water electrolytes K2SO4(s) 2K+(aq) + SO42-(aq)
sodium ions solvated by water molecules
Weak electrolytes: If a compound does not dissociate completely into ions then most of the ions of that compound are not dispersed into solution. Example: lead acetate, Pb(CH3COO)2; Cadmium iodide, CdI2 Non-electrolytes: Substances that do not form any ions in solution Example: sugar solution Covalent compounds do not form ions in solution Solubility: of a substance is the largest amount of substance that will dissolve in a given amount of solvent (expressed as grams/Liter) Soluble: >10g/L; slighlty soluble:0.1g/L-10g/L; Insoluble: <0.1g/L (these ranges are approximate estimates meant to give guidance) Warning: When a substance is soluble that does not mean that the substance dissociates into ions!!
Water soluble ionic compounds (see Table 4-1) (a). All nitrates (NO3-) are soluble in water (b). All acetates (CH3COO-), except Be(CH3COO)2 are soluble (c). All chlorates (ClO3-) are soluble (d). Most perchlorates (ClO4- ),except KClO4, are soluble (e). Most chlorides(Cl-), bromides (Br-) and iodides (I-) are soluble Slightly soluble Insoluble PbCl2 AgCl, Hg2Cl2 PbBr2 AgBr,Hg2Br2 HgBr2 AgI,Hg2I2,PbI2,HgI2 (f). Most sulfates, SO42-, are soluble in water Slightly soluble Insoluble CaSO4 PbSO4 Ag2SO4 SrSO4 Hg SO BaSO
Water insoluble ionic compounds (see Table 4-1) (a). All sulfides (S2-) are insoluble, except Group I, Ba2+ and (NH4)2S (b). All carbonates (CO32-) are insoluble, except Group I carbonates and (NH4)2CO3 (c). All sulfites (SO32-) are insoluble, except Group I sulfites and (NH4)2SO3 (d). All phosphates (PO43-) are insoluble, except Group I phosphates, (NH4)2PO4 , Li3PO4 (e). All hydroxides (OH-) are insoluble, except Group I hydroxides, Ba(OH)2, Sr(OH)2, Ca(OH)2
Note: the boxed compounds are slightly soluble
Solubility Rules in Simplified Form
Mostly soluble: All common compounds of 1. Group I and NH4+ ions 2. NO3-,CH3COO- ClO3- and ClO43. Cl-, Br-, I- except those of Ag+, Pb2+, Hg22+ 4. SO42- except those of Ca2+, Sr2+, Ba2+, Ag+, Pb2+, Hg22+ Mostly Insoluble (or slightly soluble): All common compounds of 1. OH- except those of Group I and Ba2+ 2. CO32-, PO43-, SO32- except those of Group I and NH4+ 3. S2- except those of Group 1(I), NH4+, and Ba2+
Predicting Dissolution Reactions: "like dissolves like" (1). Polar solute molecules dissolve in polar solvents [Polar and ionic substances dissolve in water but non-polar substances, such as benzene, do not dissolve in water] (2). Non-polar solute molecules dissolve in non-polar or slightly polar solvents. [non-polar substances such as benzene, Cl hydrocarbons (oils)dissolve in non-polar CCl4, but not in water] O Cl C H H (3). Solubility of gases in water Cl Cl Water: polar and other solvents decreases (CCl4: non-polar) as temperature is increased. (4). Nearly all ionic compounds of first group metals or of ammonium ion dissolve in water. ((5). Organic acids of low molar mass are soluble in water. organic compounds of low molar mass containing OH or NH2 group are soluble in water.
Precipitation When a compound "falls out" (or settles down) from a clear solution, this process is called precipitation. This process can occur in several different ways. (a). Evaporation of solvent: If you leave NaCl solution in a open container, water solvent evaporates leaving solid NaCl behind. (b). Addition of a second solvent: NaCl has high solubility in water, but poor solubility in ethanol. When ethanol is added to a saturated NaCl solution, solid NaCl "falls out" of solution. (c). Changing temperature: The solubility of a substance in a given solvent depends on temperature. In most cases, decrease of temperature decreases solubility, so a substance can be precipitated by lowering the temperature. (d). Precipitation reactions: Combination with a different kind of ion can cause precipitation .....Contd
Precipitation Reactions (b). K2SO4 dissolves in water (a). BaCl2 dissolves in water
BaCl2(s) K2SO4(s)
water
Ba2+(aq) +2 Cl-(aq)
water
2K+(aq) + SO42-(aq) 2K+(aq) +2Cl-(aq)
(c). Mix (a) and (b) +
Ba2+(aq) +2 Cl-(aq) 2K+(aq) + SO4
2-(aq)
BaSO4(s)
BaSO4(s) + 2K+(aq) + 2Cl-(aq) Spectator ions (ions present on both sides of equation): K+ and ClNet ionic equation: Ba2+(aq) + SO42-(aq) BaSO4(s), precipitation reaction
Ba2+(aq) + 2Cl-(aq)+ 2K+(aq)+ SO42-(aq)
Precipitation Reactions contd...
water
Another example: (a). Na2CO3(s)
(b). (a) + (b): CaCl2(s)
2Na+(aq)+ CO32-(aq)
water Ca2+(aq) + 2Cl-(aq)
2Na+(aq)+ CO32-(aq) + Ca2+(aq) + 2Cl-(aq) CaCO3(s)+ 2Na+(aq)+ 2Cl-(aq) Spectator ions: Na+ and ClNet ioinc equation: Ca2+(aq) + CO32-(aq) CaCO3(s)
Metathesis reaction: Reaction where atoms or groups of atoms are interchanged among reactants.
Procedure to Predict Precipitation Reaction and net ioinc equation Predict what happens when aluminum sulfate and sodium hydroxide are mixed together in water. (a). Write a balanced chemical equation. Al2(SO4)3 + 6 NaOH
3 Na2SO4 + 2 Al(OH)3
(b). Identify the substances as soluble/insoluble in water Al2(SO4)3,NaOH and Na2SO4 are soluble in water. Add (aq) to these three. Al(OH)3 is insoluble in water. Add (s) to this.
Al2(SO4)3(aq) + 6NaOH(aq) 3Na2SO4(aq)+2Al(OH)3(s) Contd....
Procedure to Predict Precipitation Reaction and net ioinc equation contd..
(c). Identify the reactants/products that dissociate into ions Al2(SO4)3,NaOH and Na2SO4 dissociate into ions in water (d). Write ionic equation 2Al3+(aq) + 3SO42-(aq)+ 6 Na+(aq)+ 6 OH-(aq) 3SO42-(aq) + 6 Na+(aq)+ 2Al(OH)3(s)
(e). Identify the spectator ions SO42-(aq) and Na+(aq) are spectator ions (common to both sides)
(f). Write net ionic equation: Remove the spectator ions from ionic equation 2Al3+(aq) + 6 OH-(aq) 2Al(OH)3(s)
Predict Precipitation Reaction and net ioinc equation Write down the reaction between sodium chloride and calcium nitrate in water. (a). 2 NaCl + Ca(NO3)2 CaCl2 + 2 NaNO3 (b). NaCl, Ca(NO3)2, CaCl2, NaNO3 are soluble in water. So,
2 NaCl(aq) + Ca(NO3)2(aq) CaCl2 (aq)+ 2 NaNO3 (aq) (c). NaCl, Ca(NO3)2, CaCl2, NaNO3 dissociate into ions in water
(d). Ionic equation: Ca2+(aq)+ 2NO3-(aq)+ 2Na+(aq)+2Cl-(aq) 2Na+(aq)+2NO3-(aq) + Ca2+(aq)+ 2Cl-(aq)
(e). All ions are spectator ions because, there is no difference bewteen left hand side and right hand side of this equation.
(f). Net ionic equation: there is no net reaction.
Predict Precipitation Reaction and net ioinc equation When two clear aqueous solutions were mixed, it was found that the final solution contained Na+(aq) and CH3COO-(aq) ions and SrSO4(s). Predict the reaction and net ionic equation. Na2SO4(aq)+ Sr(CH3COO)2(aq) Na+(aq) + CH3COO-(aq)+ SrSO4(s) (a). Balance the equation Na2SO4(aq)+ Sr(CH3COO)2(aq) 2Na+(aq) + 2CH3COO-(aq)+ SrSO4(s) (b). Na2SO4 and Sr(CH3COO)2 dissociate into ions in water (c). Ionic equation: 2Na+(aq) +SO42-(aq)+ Sr2+(aq)+2CH3COO-(aq) 2Na+(aq) + 2CH3COO-(aq)+ SrSO4(s) (d). Spectator ions: Na+(aq) and CH3COO-(aq) (e). Net ionic equation: SO42-(aq)+ Sr2+(aq) SrSO4(s)
Predict Precipitation Reaction and net ioinc equation Predict what happens when sodium chloride and iron(II)nitrate are mixed together in water (a). Balanced chemical equation: 2 NaCl + Fe(NO3)22 NaNO3 + FeCl2 (b). All nitrates are soluble; NaCl and FeCl2 are soluble. So, 2 NaCl(aq) + Fe(NO3)2(aq) 2 NaNO3(aq) + FeCl2(aq) (c). All substances dissociate into ions. (d). Ionic equation: 2Na+(aq) + 2Cl-(aq) + Fe2+(aq)+2NO3-(aq) 2 Na+(aq) + 2NO3-(aq) + Fe2+(aq)+2Cl-(aq) (e). All ions are spectator ions because, there is no difference bewteen left hand side and right hand side of this equation.
(f). Net ionic equation: there is no net reaction.
Weak Electrolytes in Precipitation Reactions Write down the reaction when mercury(II)chloride and potassium iodide are mixed together in water solvent
(a). 2KI(aq) (b). HgCl2(aq) Add (a) and (b):
2K+(aq) +2 I-(aq) HgCl2(aq) [weak electrolyte]
2K+(aq) + 2I-(aq) + HgCl2(aq) 2K+(aq) + 2Cl-(aq) + HgI2(s) Net ionic Equation: 2I-(aq) + HgCl2(aq) 2Cl-(aq) + HgI2(s)
Acids, Bases and their Reactions Autoionization of water: H2O(l) H+(aq) + OH-(aq) In pure liquid water, a small portion of water molecules dissociate into ions. This is called autoionization of water. The H+ and OH- ions are essential for acid-base behaviour Arhenius acid: is a substance that dissociates to produce H+ ions when dissolved in water. Acid is a H+ ion donor; HCl(g) water H+(aq) + Cl-(aq) Arhenius base: is a substance that dissociates to produce OHions when dissolved in water. Base is OH- ion donor; NaOH(s) water Na+(aq) + OH-(aq)
Amphoteric: a substance that behaves both as an acid as well as a base. Water is amphoteric.
Strong acid: A substance that donates 1 mol H+ ions for every mol of acid dissolved in water. That means, all of the acid molecules dissociate into ions.
So, a strong acid is also a strong electrolyte Seven strong acids H2SO4 acid sulfuric HCl hydrochloric acid/hydrogen chloride HNO3 nitric acid HBr hydrobromic acid/ hydrogen bromide HI hydroiodic acid/hydrogen iodoide HClO3 chloric acid HClO4 perchloric acid
Weak acid:
An acid where all of the molecules do not dissociate to give H+ ions when dissolved in water. That means, only a few of the acid molecules dissociate into ions. So, a weak acid is also a weak electrolyte. water Example: CH3COOH(s) CH3COOH(aq) 10% CH3COO-(aq) + H+(aq) Only 10% of acetic acid molecules dissociate into ions. The Commonly encountered weak acids: HF hydrofluoric acid/hydrogen fluoride HCOOH formic acid Bold underlined H is CH3COOH acetic acid acid hydrogen H2C2O4 oxalic acid H3PO4 phosphoric acid
remaining 90% remain in molecular form (so, it is not written in ionic form).
Naming of acids Binary acids formed by hydrogen and one other element/group (Example: HF, HCl, HBr, HI, H2S , HCN)
(a). Start with "hydro" (b). add the name of second element (c). Add ic to the stem of second element Covalent compd name Formula Acid name Hydro+fluor+ic Hydro+chlor+ic Hydro+brom+ic Hydrofluoric acid Hydrochloric acid Hydrobromic acid
hydrogen fluroide hydrogen chloride hydrogen bromide
HF HCl HBr
HI Hydro+iod+ic Hydroiodic acid hydrogen iodide H2S Hydro+sulfur+ic Hydrosulfuric acid hydrogen sulfide ...................................................................... HCN Hydro+cyan+ic Hydrocyanic acid hydrogen cyanide
Oxoacids If the name of anion ends with ate then acid name ends with -ic If the name of anion ends with ite then acid name ends with -ous Anion Acid NO3-, nitrate HNO3 , nitric acid NO2-, nitrite HNO2 , nitrous acid CrO4-, chromate HCrO4, chromic acid SO42-, sulfate H2SO4, sulfuric acid SO32-, sulfite H2SO3, sulfurous acid Organic acids: names must be memorized CH3COO-, acetate ion CH3COOH, acetic acid HCOO-, formate ion HCOOH, formic acid
Naming of acids
..contd
Neutralization reaction: Acid and base react to give salt and water
Reaction with weak acid Reaction with strong acid Neutralization reaction with strong acid: HCl(aq)+ NaOH(aq) NaCl(aq) + H2O(l)
H+ (aq) + Cl-(aq)+ Na+ (aq) +OH-(aq) Na+ (aq) + Cl-(aq) + H2O(l) Net ionic equation: H+ (aq) + OH-(aq) H2O(l)
"One mole of acid is neutralized by one mole of base"
Neutralization reaction with weak acid: CH3COOH(aq)+ NaOH(aq) CH3COONa(aq) + H2O(l) CH3COOH(aq)+ Na+ (aq) +OH-(aq) Na+ (aq) + CH3COO-(aq) + H2O(l) Net ionic equation: CH3COOH(aq)+ OH-(aq) CH3COO-(aq) + H2O(l) Example: Write net ionic equation for neutralization reaction between selenous acid, a weak acid, and sodium hydroxide. (a). Se is in the same group as S. Following sulfurous acid, H2SO3, selenous acid is H2SeO3. Balanced equation is: (b). H2SeO3(aq) + 2NaOH(aq) Na2SeO3 (aq) + 2H2O(l) (c). ionic equation: H2SeO3(aq) + 2Na+ (aq) + 2OH-(aq) 2Na+ (aq)+SeO32-(aq)+2H2O(l) (d). Net ionic equation: H2SeO3(aq) + 2OH-(aq) SeO32-(aq)+2H2O(l)
Bases Strong bases
Weak bases Group I hydroxides and Ba(OH)2 Other ionic hydroxides are are readily soluble in water and slightly or insoluble in water. They can be made to dissolve dissociate into ions in water in acidic solution. Example: Mg(OH)2 dissolves poorly in water, but dissolves readily in acidic solution. Balanced chemical equation is: 2HCl(aq)+Mg(OH)2(s) MgCl2(aq) + 2 H2O(l)
Ionic eq: 2H+(aq)+ 2Cl-(aq)+Mg(OH)2(s) Mg2+(aq) + 2Cl-(aq) + 2 H2O(l) Net ionic eq: 2H+(aq) + Mg(OH)2(s) Mg2+(aq) +2 H2O(l) Mg(OH)2, a white solid, disappears upon adding it to acidic solution.
Modified Arhenius acid/base definition Many compounds such as NH3, CH3NH2 are also bases, but they do not have OH- ions. So the definition of Arhenius acid/base is modified as follows: Arhenius base: is a substance that, when dissolved in water, increases the amount of OH- ions over what is present in pure water. Example: NH3+ H2O NH4+ + OHArhenius acid: is a substance that, when dissolved in water, increases the amount of H+ ions over what is present in pure water. Neutralization reaction: NH3(aq) + HCl(aq) NH4Cl(aq) Ionic equation: NH3(aq) + H+(aq)+ Cl-(aq) NH4+ (aq) + Cl-(aq) Net ionic eq: NH3(aq) + H+(aq) NH4+ (aq)
Note: water is not generated in this neutralization reaction
Acid anhydrides SO3(g) + H2O(l) H2SO4(aq) 2H+(aq)+SO42-(aq) SO3 is an acid because it increases the number of H+ ions SO3 is acid anhydride of H2SO4 (H2SO4-H2O=SO3)
N2O3 is acid anhydride of HNO2 I2O7 is acid anhydride of H5IO6 Base anhydrides (2HNO2-H2O= N2O3) (2H5IO6-5H2O= I2O7)
Na2O(s) + H2O(l)
2NaOH(aq)
2Na+(aq)+2OH-(aq)
Na2O acts as a base because it increases the number of OH- ions
Na2O is base anhydride of NaOH (2NaOH-H2O= Na2O)
Acid/Base anhydrides In general, Ionic oxides are base anhydrides So oxides of Group I and II metals are base anhydrides Transition metal oxides show a range of acidic/basic properties. Acidity increases with oxidation state of metal In general, covalent oxides are acid anhydrides So oxides of non-metals are acid anhydrides
I
II
III
IV
V VI acidity
VII
acidity
basicity
Oxides of main group elements (Figure 4-8 in the text book)
Basicity
Acid-Base titrations Titration is a procedure where the amount of one substance is determined by adding a carefully measured volume of a solution of a second substance with known concentration. HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
buret
NaOH solution
flask
End-point: When HCl is completely neutralized by NaOH, that particular point is called end-point.
Phenolphthalein indicator is colorless in HCl soln, but turns pink when HCl is reacted completely
HCl solution + indicator
Acid-Base titrations
..contd
Equivalence point: chemical amount of NaOH dispensed exactly equals the chemical amount of HCl originally present in the flask.
Example: Determine the amount of HCl in a solution that requires 10 mL of 0.1 M NaOH solution for complete neutralization.
HCl(aq) + NaOH(aq) Remember n = c v
NaCl(aq)+ H2O(l)
Moles of NaOH = (0.1mol/L)(10mL)(L/1000mL) = 0.001 mol moles HCl=moles NaOH
Chemical amount of HCl= 0.001 mol
Problem: A 25.00 mL volume of NaOH solution is titrated with 0.8367 M HCl. It takes 22.48 mL of this acid to reach end point. Find the concentration of NaOH solution.
Remember, n = c v mol x 22.48 mL x __ L __ Moles of HCl used = 0.8367 L 1000 mL = 0.01881 mol Stoichiometry: 1 mol HCl is neutralized by 1 mol NaOH Moles of NaOH in 25.00 mL solution =0.01881 Concentration of NaOH solution = 0.01881 mol _______________ = 0.7524 ____ mol 25.00 mL x ____ 1L L 1000 mL
Further Reactions of Acids and Bases
1. Acids react with carbonates to liberate CO2(g) and H2O(l)
K2CO3(s) + 2HCl(aq) 2KCl(aq) + H2O(l) + CO2(g) Net ionic Eq: K2CO3(s) + 2H+(aq) 2K+(aq) + H2O(l) + CO2(g) 2. Acids react with metal oxides to form salt and water CuO(s) + 2HCl(aq) CuCl2(aq) + H2O(l) Net ionic Eq: CuO(s) + 2H+(aq)Cu2+(aq) + H2O(l) 3. Acids react with many metals to liberate H2(g)
Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)
Net ionic Eq: Zn(s) + 2H+(aq)
Zn2+(aq) + H2(g)
Characteristic Reactions of Bases 1. Bases, with the exception of NH3, react with ammonium salts to generate NH3 NH4Cl(s) + NaOH(aq) NaCl(aq) + H2O(l) + NH3(g) Net ionic Eq: NH4Cl(s) + OH-(aq) Cl-(aq) + H2O(l) + NH3(g)
2. Bases react with oxides of non-metals to produce salt and water SO3(g) + 2NaOH(aq) Na2SO4(aq) + H2O(l)
SO42-(aq) + H2O(l)
Net ionic Eq: SO3(g) + 2OH-(aq)
Oxidation-Reduction Reactions: Electrons are transferred from one species to another, during these reactions.
Oxidation numbers(O.N): In covalent compounds, it is often unclear how many electrons can be counted as transferred from one element to another. So a set of oxidation numbers are assigned and these numbers are monitored in a chemical reaction.
Oxidized: A species is said to be oxidized when it gives up electrons Na Na+ + esodium is oxidized Reduced: A species is said to be reduced when it gains electrons Cl + eCl- chlorine is reduced
Oxidation-Reduction Reactions ..contd Ionic compounds: 2Na + Cl2 2NaCl 2Na+ + 2ClAtoms of elements are assigned an oxidation number of zero. So in the above reaction, Na and Cl atoms have O.N of zero before reaction. After reaction, Na lost one electron and Cl gained one electron. O.N of Na increased from 0 to +1; O.N of Cl is reduced from 0 to 1.
0 0 +1 -1
2Na + Cl2
2NaCl
Covalent compounds: In compounds such as CCl4, we cannot write it as C+4 and Cl- because carbon does not completely give up its electrons. So we assign O.N to each atom according to the following rules.
O.N Rules for compounds
(1). Atoms of elements are assigned an oxidation number of zero
(2). The O.Ns of all atoms in a neutral molecule should add up to zero. In ions they should add up to the net charge on the ion.
No. Atom (3). Group I (4). Group II (5). Group III (6). F (7). Halogens
O.N +1 +2 +3 -1 -1
Situation always
(8). (9).
Hydrogen Oxygen
always always always except in compounds with O and other halogens where they can have positive O.N +1 except in metal hydrides where H can have 1 -2 except in: (a). Compounds with F where it has positive O.N (b). In peroxides it has O.N of 1 (c). Super oxides where it has 0.5
O.N Rules for compounds Examples:
0 0 0 0 0 0 0 0 0
..contd
H-H ; O-O ; F-F ; Cl-Cl; Na
+1 -1
+2
[Rule 1]
[Rules 3,4 and 5] [Rules 6,7] [Rule 8]
-1
+3 -1
NaCl
+5 -1
MgCl2 AlCl3
+1 -1 +3 -2
IF5
+1 -1
ClF
+1 -1
Cl2O3
+3 -1
HCl
+1 -2
NaH
+1 -1
BH3
+1 -1
+1 -0.5
H2O
H2O2 Na2O2
KO2
[Rule 9]
Redox reactions 5Cu(s) + 2MnO4-(aq) + 16 H+(aq) +2 +2 5Cu2+(aq) + 2Mn2+(aq) + 8H2O(l)
0 +7
(a). Copper goes from 0 to +2 oxidation state ( so copper atom lost two electrons; total electrons lost =5 atoms x 2 = 10) (b). Manganese goes from +7 to +2 oxidation state ( so Mn atom gains 5 electrons; total electrons gained=2 atoms x 5=10) (c). 8 oxygen atoms go from 2 to 2 oxidation state. (d). 16 hydrogen atoms go from +1 to +1 oxidation state. Atoms of Cu(s) are oxidized (by MnO4-) Net reaction: Manganese atoms in MnO4- are reduced (by Cu(s)) MnO4- is oxidizing agent; Cu(s) is reducing agent Atoms in reducing agent are oxidized Atoms in oxidizing agent are reduced.
Titrations using Redox Reactions A 100 mL solution of Fe2+ is given. Titration using 0.09625 M potassium permanganate solution required 26.3 mL to reach the end point. If the reaction is given as, Pale yellow MnO4-(aq) + 5 Fe2+(aq) + 8H+(aq) Mn2+(aq) + 5 Fe3+(aq) + 4 H2O(l) deep purple What is the concentration of Fe2+ solution?
(a). Moles of MnO4 - used, n= (0.09625 mol/L) x 26.3 mL x (L/1000 mL)= 0.00253
(b). Stoichiometry: 1 mol MnO4- reacts with 5 mol Fe2+ (c). Moles of Fe2+ reacted= 0.00253 mol MnO4- x 5 mol Fe2+/ 1 mol MnO4- = 0.01266 (d). Concentration, c = (0.01266 mol/ 100 mL) x (1000 mL/L) =0.1266 M