Aqueous solutions Types of reactions

Aqueous solutions Types of reactions
Chapter 4 Aqueous solutions Types of reactions

Parts of Solutions Solution – Homogeneous mixture.
Solute – What gets dissolved. Solvent – What does the dissolving. Soluble – Can be dissolved. Insoluble – Can’t be dissolved Miscible – Liquids dissolve in each other. Immiscible – Liquids that don’t dissolve in each other

Aqueous solutions Dissolved in water.
Water is a good solvent because the molecules are polar. The oxygen atoms have a partial negative charge. The hydrogen atoms have a partial positive charge.

Hydration Hydration is the process of breaking the ions of salts apart. Ions have charges and attract the opposite charges on the water molecules. An ion is surrounded by other ions of the opposite charge.

Hydration Also check out the animation on the website.

Solubility How much of a substance will dissolve in a given amount of water. Usually g of solute / 100 grams of H2O Varies greatly, but if they do dissolve the ions are separated, and they move around. Water can also dissolve non-ionic compounds if they have polar bonds, such as soap.

Electrolytes Electricity is moving charges.
The ions that are dissolved can move. Solutions of ionic compounds can conduct electricity. Solutions are classified three ways in their ability to conduct electricity.

Types of solutions Strong electrolytes - completely dissociate
Examples: NaCl Many ions - Conduct well. Weak electrolytes - Partially split into ions. Examples: Vinegar, ammonia Few ions - Conduct slightly. Non-electrolytes - Don’t split apart. Examples: Sand, plastic No ions - Don’t conduct.

Types of solutions - Acidic
Acids – form H+1 ions when dissolved. Strong acids ionize completely. The following are strong acids H2SO4 HNO3 HCl HClO4 HBr HI HClO3 Weak acids – partial ionization Example: HC2H3O2

Types of Solutions - Basic
Bases - form OH-1 ions when dissolved. Strong bases – ionize completely. Any Group 1 hydroxide is considered to be a strong base. Examples: NaOH and KOH Transition metal hydroxides will be weaker bases. Fe(OH)2 and CuOH

Point of Emphasis Strong v. Weak & Concentrated v. Dilute
These terms often get misapplied or confused. Strong or Weak is a description of how much a species will dissociate into ions and is an intensive property. Concentrated or Dilute is a description of how much solute is used by the experimenter.

Some examples Strong and Concentrated means that the substance will dissociate fully and the experimenter used a lot of it – it is in a high concentration. Ex. 8 M HCl Strong and Dilute means that the substance will dissociate fully and the experimenter used a small concentration. Ex M HCl

Some more examples Weak and concentrated means that the substance doesn’t dissociate fully and the experimenter used a lot of it Ex. 8 M HC2H3O2 (acetic acid) Weak and dilute means that the substance doesn’t dissociate fully and the experimenter used a small concentration. Ex M HC2H3O2

Measuring Solutions Concentration - how much is dissolved.
Molarity = Moles of solute n Liters of solution L abbreviated M (Capital M) 1 M = 1 mole solute / 1 liter solution Calculate the molarity of a solution with 34.6 g of NaCl dissolved in 125 mL of solution.

Solved Answer Calculate the molarity of a solution with 34.6 g of NaCl dissolved in 125 mL of solution. First – find moles of NaCl  moles Second – set up formula M = n / L M = .592 / .125 L M = M

Molarity How many grams of HCl would be required to make 50.0 mL of a 2.7 M solution ? What would the concentration be if you used 27g of CaCl2 to make 500. mL of solution ? What is the concentration of each ion?

Solved Answer 2.7 = n / .050 L n = .135 moles
How many grams of HCl would be required to make 50.0 mL of a 2.7 M solution ? Use M = n / L 2.7 = n / .050 L n = .135 moles Use n = g / M  g of HCl This is added with enough water to make 50 mL of soln.

More Solved Answers What would the concentration be if you used 27 g. of CaCl2 to make 500. mL of solution ? (FW of CaCl2 = g) Find moles  27 / = .243 moles Use M = n / L M = .243 / .5 L M = .487 M (can be read as .487 Molar)

Molarity Calculate the concentration of a solution made by dissolving 45.6 g of Na2SO4 to 475 mL. (FW of Na2SO4 = ) What is the concentration of each ion ? A balanced equation is a good place to start Na2SO4  2 Na+1 + SO4-2 This shows 2 Na+1 ions for every sulfate.

A Little Stoichiometry
Calculate the concentration of a solution made by dissolving 45.6 g of Na2SO4 to 475 mL. Solve as before: 45.6 /  moles Use M = n / L / .475 = .676 M Since the Sodium Ions are in a 2 to 1 ratio there would be M Na+1 and there would be .676 M SO4-2

Look at the Stoichiometry
Let’s look at the dissociation Na2SO4  2 Na+1 + SO4-2 We see the 2 : 1 ratio of the two ions, so that is how we get the values from the slide above.

Making solutions # 1 Describe how to make mL aliquot of a 1.0 M K2Cr2O4 solution. The formula wt. of the solute is g Using M = n / L  = n / .1 L So we need .1 moles of solute. Using n = g / m  .1 = g / 246.2 We need grams of solute

Making Solutions # 2 Describe how to make 250. mL of an 2.0 M copper (II) sulfate solution. The formula wt. of the solute is g. Using M = n / L  2 = n / .25 So we need .5 moles of solute Using n = g / M  .5 = g / 159.6 We need 79.8 grams of solute

Dilution Adding more solvent to a known solution.
The moles of solute stay the same. moles = M x L M1 V1 = M2 V2 (this is a very handy formula) moles = moles Stock solution is a solution of known concentration used to make more dilute solutions

Dilution # 1 What volume of a 1.7 M solutions is needed to make 250 mL of a 0.50 M solution ? 1. Given info. M1 = 1.7 M2 = V2 = .25 2. Formula M1 V1 = M2 V2 3. Substitute 1.7 V1 = .5 (.25) 4. Solve L (or 73.5 mL)

Dilution # 2 18.5 mL of 2.3 M HCl is added to 250 mL of water. What is the concentration of the solution ? 1. Given info. V1 = M1 = 2.3 V2 = .25 2. Formula M1 V1 = M2 V2 3. Substitute (2.3) = x (.25) 4. Solve M

Dilution You have a 4.0 M stock solution. Describe how to make 1.0 L of a .75 M solution. Using M1V1 = M2V2 1.0 (.75) = 4.0 V2 V2 = .188 L (or mL) Take mL of the 4.0 M Stock soln. and add enough water to make 1.0 L

Types of Reactions Precipitation reactions
When aqueous solutions of ionic compounds are poured together a solid forms. A solid that forms from mixed solutions is a precipitate If you’re not a part of the solution, your part of the precipitate (old chem joke)

Precipitations Reactions
Only happens if one of the products is insoluble. Otherwise all the ions stay in solution- nothing has happened. Need to memorize the rules for solubility.

Some Solubility Rules We need to know these.
All nitrates are soluble Alkali metals ions and NH4+1 ions are soluble Halides are soluble except Ag+1, Pb+2, and Hg2+2 Most sulfates are soluble, except Pb+2, Ba+2, Hg+2,and Ca+2

Solubility Rules Most hydroxides are slightly soluble (insoluble) except NaOH and KOH Sulfides, carbonates, chromates, and phosphates are generally insoluble

Precipitation reactions
NaOH(aq) + FeCl3(aq) ® NaCl(aq) + Fe(OH)3(s) is really Na+1(aq)+OH-1(aq) + Fe+3(aq) + Cl-1(aq) ® Na+1(aq) + Cl-1 (aq) + Fe(OH)3 (s) So all that really happens is 3 OH-1(aq) + Fe+3(aq) ® Fe(OH)3(s) Essentially a more complex double replacement reaction.

Precipitation reaction
We can predict the products by knowing types of reactions and the rules for solubility AgNO3 (aq) + KCl (aq) ® Zn(NO3)2 (aq) + BaCr2O7 (aq) ® CdCl2 (aq) + Na2S (aq) ®

Answers AgNO3 (aq) + KCl (aq)  KNO3 (aq) + AgCl (s)
We know the KNO3 is soluble because all Group 1 metals are soluble. Zn(NO3)2 (aq) + BaCr2O7 (aq)  Ba(NO3)2 (aq) + ZnCr2O7 (s) The nitrates are soluble. CdCl2 (aq) + Na2S (aq)  NaCl (aq) + CdS (s) The Group 1 Sodium ion is soluble.

Precipitation Stoichiometry
25 mL of 0.67 M H2SO4 is added to 35 mL of 0.40 M CaCl2 . What mass CaSO4 Is formed ? First we need a balanced equation and moles of each reactant. H2SO4 + CaCl2  CaSO4 + 2 HCl

Precipitation Stoichiometry
H2SO4 + CaCl2  CaSO4 + 2 HCl The stoichiometry is 1:1 so CaCl2 is the limiting reagent. The ratio would be: 1 n CaCl2 / 1 CaSO4 = .014 / x x = .014 n of CaSO4 Convert to grams n g / M .014 = g / g = 1.9 grams

Three Types of Equations
Molecular Equation - written as whole formulas, not the ions. K2CrO4 (aq) + Ba(NO3)2 (aq) ® Complete Ionic Equation show dissolved electrolytes as the ions. 2 K+1 + CrO4-2 + Ba NO3-1 ® BaCrO4 (s) + 2 K NO3-1 Spectator ions are those that don’t react.

Three Type of Equations
Net Ionic Equations show only those ions that react, not the spectator ions Ba+2 + CrO4-2 ® BaCrO4 (s) Write the three types of equations for the reactions when these solutions are mixed. a. Iron (III) sulfate and potassium sulfide b. Lead (II) nitrate and sulfuric acid.

Three Types of Equations
a. Iron (III) sulfate and potassium sulfide Fe2(SO4)3 (aq) + K2S (aq)  K2SO4 (aq) + Fe2S3 (s) Fe+3 + SO4-2 + K+1 + S-2  K+1 + SO4-2 + Fe2S3 (s) Fe+3 (aq) + S-2 (aq)  Fe2S3 (s) b. Lead (II) nitrate and sulfuric acid. Pb(NO3)2 (aq) + H2SO4 (aq)  HNO3 (aq) + PbSO4 (s) Pb+2 + NO3-1 + H+1 + SO4-2  NO3-1 + H+1 + PbSO4 (s) Pb+2 + SO4-2  PbSO4 (s)

Stoichiometry of Precipitation
Exactly the same as regular stoichiometry, except you are finding moles in a different way. Concepts of Limiting Reagent and materials in excess are also the same as before.

Stoichiometry of Precipitation – Problem # 1
What mass of solid is formed when mL of M Barium chloride is mixed with mL of M sodium hydroxide ? Balanced Equation: BaCl2 + 2 NaOH  Ba(OH)2 + 2 NaCl Calculate Moles of each reactant BaCl2  M = n / L .1 = n / .1 L n = .01 n NaOH  same n

Look at ratios to determine limiting reagent. BaCl2 + 2 NaOH  Ba(OH)2 (s) + 2 NaCl (aq) 2 / 1 = .01 NaOH / x x = .005 moles So NaOH is the limiting reagent Proportion 2 n NaOH / 1 n Ba(OH)2 = .01 / x x = .005 moles

What volume of M HCl is needed to precipitate all the silver from 50.ml of M silver nitrate solution ? HCl (aq) + AgNO3 (aq)  AgCl (s) + HNO3 (aq) .003 n It is a 1:1 ratio so the moles are the same. Use M = n / L = .003 / L L = L

Types of Reactions Acid-Base For our purposes an acid is a H+1 donor.
A base is a OH-1 donor. A neutralization is when there is equal numbers of moles of acid and moles of base. A titration is the lab procedure used to achieve this.

Always remember the following when dealing with Acid-Base Reactions
General Formula: Acid + Base  salt + water Net Ionic Form: H+1 + OH-1  H2O Dissociation of Water: H2O  H+1 + OH-1

Acid - Base Reactions Often called a neutralization reaction Because the acid neutralizes the base. Often titrate to determine concentrations. Solution of known concentration (titrant), Is added to the unknown (analyte), Until the equivalence point is reached where enough titrant has been added to neutralize it.

Titration Where the indicator changes color is the endpoint.
Not always at the equivalence point. A mL sample of aqueous Ca(OH)2 requires mL of M Nitric acid for neutralization. What is [Ca(OH)2 ] ? # of H+1 = # of OH-1

Titration Solved Problem
50.00 mL sample of aqueous Ca(OH)2 requires mL of M Nitric acid for neutralization. What is [Ca(OH)2 ] ? Ca(OH)2 + 2 HNO3  Ca(NO3)2 + H2O .098 = n / n = moles H+1 Ca(OH)2  Ca OH-1

Acid – Base Reaction This is an important problem
75 mL of 0.25 M H2SO4 is mixed with 225 mL of M NaOH . What is the concentration of the excess H+1 or OH-1 ? Think about what you are being asked to solve and about the process that is going on. In other words, don’t just dive in before you plan the problem.

Acid – Base Reaction First lets get a balanced equation.
H2SO4 + 2 NaOH  Na2SO4 + 2 H2O Notice this follows the general form that we talked about a few slides back. Acid + Base  Salt + Water

Acid – Base Reaction We need to determine moles of each reactant.
H2SO4: M = n / L .25 = n / .075 n = n NaOH: same format moles We need to set up a proportion to determine which species is the Limiting Reagent and which is in excess.

Acid – Base Reaction 1 n H2SO4 / 2 n NaOH : x / .0124 n NaOH
Solving for x gives n of H2SO4 When we compare that to the given we see that we have a lot more of the acid than what is needed so we can conclude that H2SO4 is in excess and the NaOH is the limiting reagent. So what is the next step ?

Acid – Base Reaction Since we are given n of H2SO4 and only need moles we are left with = n of H2SO4 in excess However, each molecule of H2SO4 liberates 2 ions of H+1 so we need to factor that in. H2SO4  2 H+1 + SO4-2

Acid – Base Reaction This is a limiting reagent problem as well as an acid base problem. After the reaction takes place and accounting for the dissociation of the acid we have the following: n of H+1 But we were asked for concentration. Clearly we need to use M = n / L , but there is a catch. Solve for M in your binder.

Acid – Base Reaction We need moles and we just solved for that and got moles of H+1 . What about the volume ? Both reagents are in the solution together, so the volumes need to be combined. The total volume is mL. M = / .3  M = .084 M H+1 Forgetting to add the volumes together is a common mistake in these types of problems.

Types of Reaction - Redox
Oxidation-Reduction called Redox Ionic compounds are formed through the transfer of electrons. An Oxidation – Reduction reaction involves the transfer of electrons. We need a way of keeping track.

Oxidation States A way of keeping track of the electrons.
You need to memorize the rules for assigning oxidation states. The oxidation state of elements in their standard states is zero. Oxidation state for monoatomic ions are the same as their charge.

Oxidation states Oxygen is assigned an oxidation state of -2 in its covalent compounds except as a peroxide. In compounds with nonmetals hydrogen is assigned the oxidation state +1. In its compounds fluorine is always –1. 6 The sum of the oxidation states must be zero in compounds or equal the charge of the ion.

Oxidation States Assign the oxidation states to each element in the following. CO2 NO3-1 H2SO4 Fe2O3

Oxidation States Assign the oxidation states to each element in the following. CO2 C = O = -2 NO3-1 N = O = -2 H2SO4 H = S = O = -2 Fe2O3 Fe = O = -2

Oxidation-Reduction Transfer electrons, so the oxidation states change. 2 Na + Cl2 ® 2 NaCl CH4 + 2 O2 ® CO2 + 2 H2O Oxidation is the loss of electrons. Reduction is the gain of electrons. OIL RIG or LEO GER or OLE

Oxidation-Reduction Oxidation means an increase in oxidation state - lose electrons. Reduction means a decrease in oxidation state - gain electrons. The substance that is oxidized is called the reducing agent. The substance that is reduced is called the oxidizing agent.

Redox Reactions

Redox Vocabulary Oxidizing agent gets reduced. Gains electrons.
More negative oxidation state. Reducing agent gets oxidized. Loses electrons. More positive oxidation state. Both have to happen in a reaction for it to occur.

Identify the following
Oxidizing agent & reducing agent Substance oxidized & Substance reduced for the following reactions Fe (s) + O2(g) ® Fe2O3(s) Fe2O3(s)+ 3 CO(g) ® 2 Fe(s ) + 3 CO2(g) SO3-2 + H+1 + MnO4-1 ® SO4-2 + H2O + Mn+2

Answers Fe (s) + O2(g) ® Fe2O3(s) Iron is oxidized, Oxygen is reduced
Iron is the reducing agent, Oxygen is the oxidizing agent Fe2O3(s)+ 3 CO(g) ® 2 Fe(s ) + 3 CO2(g) Fe2O3 is reduced, CO is oxidized Fe2O3 is the oxidizing agent, CO is the reducing agent SO3-2 + H+1 + MnO4-1 ® SO4-2 + H2O + Mn+2 S is oxidized, Mn is reduced S is the reducing agent, Mn is the oxidizing agent

Balancing Half-Reactions is being de-emphasized for the AP Test
Balancing Half-Reactions is being de-emphasized for the AP Test. (But it is good to know the theory behind it) All redox reactions can be thought of as happening in two halves. One produces electrons - Oxidation half. The other requires electrons - Reduction half. Write the half reactions for the following. Na + Cl2 ® Na+1 + Cl-1 SO3-2 + H+1 + MnO4-1 ® SO4-2 + H2O + Mn+2

Balancing Half-reactions
Na + Cl2  Na+1 + Cl-1 Na  Na+1 + e (oxidation ½ reaction) Cl2 + 2 e-1  2 Cl-1 (reduction ½ reaction) SO3-2 + H+1 + MnO4-1  SO4-2 + H2O + Mn+2 SO3-2 + H2O  SO H e-1 8 H+1 + MnO e-1  Mn H2O

Redox Theory You need to remember that while these are studied separately, the oxidation half-reaction and the reduction half-reaction must always happen together. You can’t have a species gaining electrons unless you have a species there to lose those same electrons.

Balancing Redox Equations
In aqueous solutions the key is the number of electrons produced must be the same as those required. For reactions in acidic solution an 8 step procedure. Write separate half reactions For each half reaction balance all reactants except H and O Balance O using H2O

Acidic Solution Balance H using H+1 Balance charge using e-1
Multiply equations to make electrons balance. Add equations and cancel identical species. Check that charges and elements are balanced.

Summary of Rules in Acidic Medium
Balance the atoms in each half-reaction. This refers to atoms other than H and O Balance the Oxygen by adding H2O where needed. Balance the Hydrogen by adding H+1 where needed. Balance by charge by adding e-1

Balancing in Acidic Medium Problem # 1
Fe (s) + O2(g) ® Fe2O3(s) Determine the charge on each species So: Fe  Fe+3 (Fe is oxidized) & O2  O-2 (O2 is reduced)

Fe  Fe+3 (Fe is oxidized) Fe  Fe e-1 O2  O-2 (O2 is reduced) O2  2 O-2 4 e-1 + O2  2 O-2

The final step is to combine the two ½ reactions. But first we must get rid of the electrons 4 ( Fe  Fe e-1 ) 3 ( 4 e-1 + O2  2 O-2 ) We get: 4 Fe + 3 O2  4 Fe O-2

Fe2O3(s)+ 3 CO(g) ® 2 Fe(s ) + 3 CO2(g) Fe e-1  Fe (Fe is reduced) a. CO  CO2 b. H2O + CO  CO2 c. H2O + CO  CO H+1 d. H2O + CO  CO H e-1

Fe e-1  Fe (Fe is reduced) The CO is oxidized. Here are the steps. Atoms: CO  CO2 Oxygen (with water) H2O + CO  CO2 Hydrogen (with H+1) H2O + CO  CO2 + 2 H+1 Charge (with e-1) H2O + CO  CO H e-1

Once again, the final step is to factor out the electrons and combine the two ½ reactions. 2 ( Fe e-1  Fe ) 3 (H2O + CO  CO H e-1 ) 2 Fe H2O + 3 CO  2 Fe +3 CO2 + 6 H+1

SO3-2 + H+1 + MnO4-1  SO4-2 + H2O + Mn+2 We can see that the S is being oxidized and the Mn is being reduced.

Oxidation ½ reaction SO3-2 + H+1 + MnO4-1  SO4-2 + H2O + Mn+2 a. SO3-2  SO4-2 b. H2O + SO3-2  SO4-2 c. H2O + SO3-2  SO H+1 d. H2O + SO3-2  SO H e-1 Sulfur is oxidized. How can you tell ?

Reduction ½ reaction SO3-2 + H+1 + MnO4-1  SO4-2 + H2O + Mn+2 a. MnO4-1  Mn+2 b. MnO4-1  Mn H2O c. 8 H+1 + MnO4-1  Mn H2O d. 5 e H+1 + MnO4-1  Mn H2O Factor out the electrons and then combine the two half – reactions.

H2O + SO3-2  SO H e-1 5 e H+1 + MnO4-1  Mn H2O Multiply the first ½ rxn. By 5 and the second ½ rxn. by 2 to factor out the e-1. 5 H2O + 5 SO3-2  5 SO H e-1 10 e H MnO4-1  2 Mn H2O We can factor out species other than e-1

5 H2O + 5 SO3-2  5 SO H e-1 10 e H MnO4-1  2 Mn H2O We can start with the water and electrons. 10 e H MnO4-1  2 Mn H2O Next we can factor out the H+1 ions. 5 SO3-2  5 SO H+1 16 6 H MnO4-1  2 Mn H2O

5 SO3-2  5 SO4-2 6 H MnO4-1  2 Mn H2O Combining the two gives us: 6 H MnO SO3-2  Mn H2O + 5 SO4-2

Basic Solution Do everything you would with acid, but add one more step. Add enough OH-1 to both sides to neutralize the H+1 Determine the species undergoing redox in the problems below. CrI3 + Cl2 ® CrO4-1 + IO4-1 + Cl-1

Redox – Basic Solutions
CrI3 + Cl2 ® CrO4-1 + IO4-1 + Cl-1 Cr+3  Cr+7 (CrO4-1) Cr is oxidized I-1  I I is oxidized Cl2  Cl Cl is reduced

Practice MnO4-1 + Fe+2 ® Mn+2 + Fe+3 Cu + NO3-1 ® Cu+2 + NO(g)
Cr(OH)3 + OCl-1 + OH-1  CrO4-2 + Cl-1 + H2O MnO4-1 + Fe+2 ® Mn+2 + Fe+3 Cu + NO3-1 ® Cu+2 + NO(g) Pb + PbO2 + SO4-2 ® PbSO4 Mn+2 + NaBiO3 ® Bi+3 + MnO4-1

Balancing in Basic Medium Problem # 1
Mn+2 + NaBiO3 ® Bi+3 + MnO4-1 You must know the procedure to do this or any problem. Get the oxidation states.

Mn+2 + NaBiO3 ® Bi+3 + MnO4-1 Do the ½ reactions Mn+2  MnO4-1 4 H2O + Mn+2  MnO H+1 8 OH H2O + Mn+2  MnO H2O Notice how the OH-1 has been added above 8 OH H2O + Mn+2  MnO H2O + 5 e-1

BiO3-1  Bi+3 6 H+1 + BiO3-1  Bi H2O 6 H2O + BiO3-1  Bi H2O + 6 OH-1 2 e H2O + BiO3-1  Bi H2O + 6 OH-1

We now combine the two ½ reactions. 8 OH H2O + Mn+2  MnO H2O + 5 e-1 2 e H2O + BiO3-1  Bi H2O + 6 OH-1 We need to take care of the electrons. 16 OH H2O + 2 Mn+2  2 MnO H2O 30 H2O + 5 BiO3-1  5 Bi H2O + 30 OH-1

16 OH H2O + 2 Mn+2  2 MnO H2O 30 H2O + 5 BiO3-1  5 Bi H2O + 30 OH-1 Combine 16 OH H2O + 2 Mn BiO3-1  2 MnO H2O + 5 Bi+3 30 OH-1 Reduce 7 H2O + 2 Mn BiO3-1  2 MnO Bi OH-1

Redox Titrations Same as any other titration.
The permanganate ion is used often because it is its own indicator. MnO4-1 is purple, Mn+2 is colorless. When reaction solution remains clear, MnO4-1 is gone. Chromate ion is also useful, but color change, orange yellow to green, is harder to detect.

Example The iron content of iron ore can be determined by titration with standardized KMnO4. The iron ore is dissolved in HCl, and the iron reduced to Fe+2 ions. This solution is then titrated with KMnO4 soln., producing Fe+3 and Mn+2 ions in acidic solution. If it requires mL of M KMnO4 to titrate a solution made with g of iron ore, what percent of the ore was iron ?

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