Aqueous Reactions and Solution Stoichiometry
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Unit 12: Solutions • Topics Covered – General properties of aqueous solutions – Aqueous reactions – Solubility – Concentration – Colligative properties
• Reading Quiz – Read chapters 4 & 13 – Know definitions for all words in bold – Table 4.1 – Sample exercises 4.2, 4.3, 4.11, 4.12, 4.13, 4.14, 13.1, 13.2, 13.3, 13.4
Electrolytic Properties
General Properties of Aqueous Solutions • A solution is a homogenous mixture of two or more substances • The substance present in the greater quantity is usually called the solvent • The other substances in the solution are called the solutes
Electrolytic Properties
• A substance whose aqueous solutions contain ions (such as NaCl) is called an electrolyte • Electrolytic solutions conduct electricity • A substance that does not form ions in solution (such as C6H12O6) is a nonelectrolyte • Nonelectrolytic solutions do not conduct electricity
Strong vs. Weak Electrolytes
Ionic Compounds in Water
• Strong electrolytes are characterized by their nearly complete dissociation in water, while weak electrolytes dissociate to a much smaller extent
• In water, some ionic crystals will dissociate into their component ions • Water is a very effective solvent for many ionic compounds because of its polarity, which creates strong ion-dipole interactions that dissolve the ions into solution • Electrolytes
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Sample Problems
• From the diagrams shown, identify the strongest electrolyte and the weakest electrolyte
Precipitation Reactions • Reactions that result in the formation of an insoluble product are called precipitation reactions • Precipitation reactions occur when certain pairs of oppositely charged ions attract each other so strongly that they form an insoluble ionic compound • Example: – Pb(NO3)2(aq) + 2KI(aq)
PbI2(s) + 2KNO3(aq)
Molecular Compounds in Water • When a molecular compound dissolves in water, the solution usually consists of intact molecules dispersed throughout the solution • Some molecular substances ionize in water to form electrolytic solutions • The most important of these are acids • For example, HCl(g) ionizes in water to form HCl(aq), which is H+(aq) and Cl-(aq) ions
Solubility Guidelines for Ionic Compounds • The solubility of a substance is the amount of that substance that can be dissolved in a given quantity of solvent • For our purposes, any substance with solubility less than 0.01 mol/L will be referred to as insoluble • There are no simple rules that allow us to predict whether or not a compound will be soluble in water • We do have guidelines for predicting the solubility of many compounds • Just remember that all common ionic compounds of the alkali metals and the ammonium ion are soluble in water
• Precipitation reaction
Solubility Guidelines for Ionic Compounds
Sample Problem • Classify the following as soluble or insoluble in water: – – – –
Sodium carbonate Lead sulfate Cobalt(II) hydroxide Barium nitrate
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Exchange Reactions
Sample Problem
• Exchange reactions, double replacement, or metathesis reactions, are reactions in which ionic compounds exchange cations
• Which combinations of anions and cations would produce a precipitate as represented in the diagram?
Net Ionic Equations
PbI2(s) + 2KNO3(aq)
• An equation where all strong electrolytes are shown as ions is called a complete ionic equation – Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2I-(aq) + 2K+(aq) + 2NO3-(aq)
AgCl(s) +
• Precipitation reactions conform to this pattern, as do many acid-base reactions
Net Ionic Equations
• An equation that shows the complete chemical formulas of reactants and products is called a molecular equation – Pb(NO3)2(aq) + 2KI(aq)
– AX + BY AY + BX – AgNO3(aq) + KCl(aq) KNO3(aq)
PbI2(s)
• Ions that appear on both sides of a complete ionic equation are called spectator ions (shown in red)
Sample Problem • Write the net ionic equation for the precipitation reaction that occurs when solutions of calcium chloride and sodium carbonate are mixed
• When spectator ions are omitted, we are left with the net ionic equation – Pb2+(aq) + 2I-(aq)
PbI2(s)
• To write net ionic equations: – Write a balanced molecular equation – Write a complete ionic equation by separating strong electrolytes – Identify and cancel spectator ions
Sample Problem • Write the net ionic equation for the precipitation reaction that occurs when solutions of silver nitrate and potassium phosphate are mixed
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Acid-Base Reactions • Acids and bases that are strong electrolytes (completely ionized in solution) are called strong acids and strong bases • Those that are weak electrolytes (partly ionized) are called weak acids and weak bases
Acid-Base Reactions • When a solution of an acid and a base are mixed, a neutralization reaction occurs – HCl(aq) + NaOH(aq)
H2O(l) + NaCl(aq)
• A neutralization reaction between an acid and a metal hydroxide produces a water and a salt • The term salt is any ionic compound whose cation comes from a base and anion comes from an acid
Sample Problem • Write a balanced equation for the reaction between aqueous solutions of acetic acid (HC2H3O2) and barium hydroxide and write the net ionic equation
Oxidation-Reduction Reactions • Examples – Ca(s) + 2H+(aq) Ca2+(aq) + H2(g) – 2Ca(s) + O2(g) 2CaO(s)
• In the first case, calcium is being oxidized by the hydrogen ion, which is being reduced • The species being oxidized is often referred to as the reducing agent • The species being reduced can be called the oxidizing agent
Oxidation-Reduction Reactions • Reactions where electrons are transferred between reactants are called oxidation-reduction, or redox, reactions • Loss of electrons by a substance in called oxidation (LEO) • Gaining of electrons by a substance in called reduction (GER)
Oxidation Numbers • In order to identify the species being oxidized and reduced, you must first assign oxidation numbers to each element • Rules for assigning oxidation numbers – For an atom in elemental form the ON is always zero – For monatomic ions, the ON equals the charge of the ion
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Oxidation Numbers
Oxidation Numbers
• Rules for assigning ONs (continued)
• Rules for assigning ONs (continued)
– Nonmetals usually have negative ON, though they can sometimes be positive
– The sum of the ONs of all atoms in a neutral compound is zero – The sum of the ONs of all atoms in a polyatomic equals the charge of the ion
• The ON of oxygen is usually -2, the major exception is peroxides, O22-, where the ON is -1 • The ON of hydrogen is +1 when bonded to nonmetals and -1 when bonded to metals • The ON of fluorine is -1 in all compounds, the other halogens have ONs of -1 in most binary compounds except when combined with O
Sample Problem • Determine the oxidation state of sulfur in each of the following: H2S, S8, SCl2, Na2SO3, SO42-
Oxidation of Metals by Acids and Salts • The reaction of a metal with either an acid or a metal salt conforms to the following pattern: – A + BX AX + B – Zn(s) + 2HBr(aq) ZnBr2(aq) + H2(g) – Mn(s) + Pb(NO3)2(aq) Mn(NO3)2(aq) + Pb(s)
Sample Problem • What is the ON of the boldfaced element in each of the following: P2O5, NaH, Cr2O72-, SnBr4, BaO2
Sample Problem • Write a balanced molecular and net ionic equations for the reactions of aluminum with hydrobromic acid.
• These reactions are called displacement reactions or single replacement because the ion in solution is displaced or replaced through oxidation
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Sample Problem • Write a balanced molecular and net ionic equations for the reaction between magnesium and a solution of cobalt(II) sulfate. What is oxidized and what is reduced?
The Activity Series • A list of metals in order of decreasing ease of oxidation is called an activity series • Hydrogen is also included in this list • Any metal on the list can be oxidized by ions of the elements below it
Sample Problems • Will an aqueous solution of iron(II) chloride oxidize magnesium metal? • Which of the following metals will be oxidized by Ni(NO3)2: Zn, Cu, Fe?
Concentrations of Solutions • Concentration is the amount of solute dissolved in a given amount of solvent or solution • Molarity (M) expresses concentration as the number of moles of solute per liter of solution
Molarity
Sample Problem • Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate in enough water to form 125 mL of solution.
moles solute volumeof solution in liters
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Sample Problem • How many grams of glucose (C6H12O6) should be dissolved to make a 100. mL, 0.278 M solution?
Sample Problem • What is the molar concentration of K+ ions in a 0.015 M solution of potassium carbonate?
Sample Problem • What are the concentrations of each of the ions present in a 0.025 M aqueous solution of calcium nitrate?
Dilution • Solutions are often prepared for laboratory use in a concentrated form called a stock solution • The stock solution can then be diluted to the appropriate concentration by adding water • Moles solute before dilution = moles solute after dilution, therefore:
Mconc Vconc
Sample Problem • How many mL of 3.0 M H2SO4 are needed to make 450 mL of 0.10 M H2SO4?
Mdil Vdil
Sample Problem • If 10.0 mL of a 10.0 M stock solution of NaOH is diluted to 250 mL, what is the concentration of the resulting solution?
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Titration • To determine the concentration of a sample, chemists will often react it with a reagent of known concentration, called a standard solution, in a process called titration • Titrations can be performed for acid-base, precipitation, or redox reactions • The point at which stoichiometrically equivalent quantities are brought together is called the equivalence point
Sample Problem • The quantity of Cl- in a water supply is determined by titrating the sample with Ag+ according to the reaction Ag+(aq) + Cl-(aq) AgCl(s). How many grams of chloride ions are in a sample of the water if 20.2 mL of 0.100 M Ag+ solution is needed to react with all the chloride in the sample?
Sample Problem • A sample of 70.5 mg of potassium phosphate is added to 15.0 mL of 0.050 M of silver nitrate, resulting in the formation of a precipitate. (a) Write the molecular equation for the reaction. (b) What is the limiting reagent in the reaction? (c) Calculate the theoretical yield, in grams, of the precipitate that forms.
Titration • For acid-base titrations, a color changing indicator is used to determine when the reaction has reached its end point
Sample Problem • What is the molarity of an NaOH solution if 48.0 mL is needed to neutralize 35.0 mL of 0.144 M H2SO4?
Saturated Solutions and Solubility • A solution that is in equilibrium with undissolved solute is called saturated • Additional solute will not dissolve if added to a saturated solution • The amount of solute needed to form a saturated solution in a given quantity of solvent is known as solubility
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Saturated Solutions and Solubility • Under certain conditions it is sometimes possible to form solutions with a greater amount of solute than needed to form a saturated solution • Such solutions are called supersaturated • Supersaturated solutions are usually created by saturating a solution at high temperatures and then cooling it very slowly and carefully
Factors Affecting Solubility • The stronger the attractions between solute and solvent molecules, the greater the solubility • Polar liquids tend to dissolve readily in polar liquids, and nonpolar liquids tend to dissolve readily in nonpolar liquids • Nonpolar liquids tend to be insoluble in polar liquids • Pairs of liquids that mix in all proportions are called miscible, whereas those that do not are called immiscible • Like dissolves like
Sample Problem • Rank the following in terms of increasing solubility in water: C5H10(OH)2, C5H11Cl, C5H12, C5H10OH
Saturated Solutions and Solubility • Supersaturated solutions are unstable • The addition of a small seed crystal or simple agitation will cause the crystallization of the excess solute
Sample Problem • Predict whether each of the following substances is more likely to dissolve in carbon tetrachloride, CCl4, or water: C7H16, Na2SO4, HCl, and I2.
Pressure Effects • The solubility of solids and liquids are not affected by pressure • The solubility of a gas in any solvent is increased as the pressure over the solvent is increased
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Henry’s Law • The solubility of the gas increases in direct proportion to its partial pressure above the solution • This relationship is known as Henry’s law:
Sg
kPg
Sample Problem • Calculate the concentration of CO2 in a soft drink that is bottled with a partial pressure of CO2 of 4.0 atm over the liquid at 25 C. The Henry’s law constant for CO2 in water at this temperature is 3.1 x 10-2 mol/L-atm.
• Sg is the solubility of the gas, usually in molarity, k is called the Henry’s law constant, and Pg is the partial pressure of the gas over the solution
Sample Problem
Temperature Effects
• Calculate the concentration of CO2 in a soft drink after the bottle is opened and equilibrates at 25 C under a CO2 partial pressure of 3.0 x 10-4 atm.
• The solubility of most solid solutes in water increases as the temperature of the solution increases
Temperature Effects
Mass Percentage • One of the simplest ways to express concentration is mass percentage: Mass% of component
• Most gases become less soluble in water as the temperature rises
mass of component in solution 100% total mass of solution
• For example, a solution that is 36% HCl by mass contains 36 g of HCl for every 100 g of solution
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Parts Per Million • For very dilute solutions, concentration is often expressed in parts per million (ppm): ppm of component
Sample Problem • A solution is made by dissolving 13.5 g of glucose in 0.100 kg of water. What is the mass percentage of solute in the solution?
mass of component in solution 106 total mass of solution
• A solution whose solute concentration is 1 ppm contains 1 g of solute for every million (106) g of solution
Sample Problem • A 2.5 g sample of groundwater was found to contain 5.4 g of Zn2+. What is the concentration in ppm?
Mole Fraction • Recall that the mole fraction (X) of a component of a solution is: X of component
moles of component total moles of all components
• Thus a solution containing 1.00 mol of HCl in 8.00 mol of water has a mole fraction of XHCl = (1.00 mol)/(1.00 mol + 8.00 mol) = 0.111 (no units)
Molality • The molality (m) of a solution equals the number of moles of solute per kg of solvent: Molality
Sample Problem • A solution is made by dissolving 4.35 g glucose in 25.0 mL of water. Calculate the molality of glucose in the solution.
moles solute kg solvent
• This is different from molarity, which depends on liters of solution instead of kg of solvent
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Sample Problem • What is the molality of a solution made by dissolving 36.5 g of napthalene (C10H8) in 425 g of toluene (C7H8)?
Sample Problem • A commercial bleach solution contains 3.62 mass % NaOCl in water. (a) Calculate the molality of NaOCl in the solution. (b) Calculate the mole fraction.
Sample Problem • A solution containing equal masses of glycerol (C3H8O3) and water has a density of 1.10 g/mL. Calculate (a) the molality of glycerol, (b) the mole fraction of glycerol, and (c) the molarity of glycerol in the solution.
Sample Problem • A solution of HCl contains 36% HCl by mass. (a) Calculate the mole fraction of HCl in the solution. (b) Calculate the molality of HCl in the solution.
Sample Problems • A solution contains 5.0 g of toluene (C7H8) and 225 g of benzene and has a density of 0.876 g/mL. Calculate the molarity.
Colligative Properties • Physical properties of solutions can vary greatly from those of the pure solvent • Lowering vapor pressure and freezing point and raising the boiling point are properties of solutions that do not depend on the type or identity of solute particles but only on the quantity (concentration) of solute particles • These properties are known as colligative properties
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Lowering Vapor Pressure
Lowering Vapor Pressure
• Adding a nonvolatile solute always lowers the vapor pressure of a solvent
• The extent to which this happens depends on the concentration of the solute and follows Raoult’s law:
PA
X A PA
• The partial pressure exerted by the solvent above a solution, PA, equals the product of the mole fraction of the solvent, XA, times the vapor pressure of the pure solvent, PA
Sample Problem • Glycerin (C3H8O3) is a nonvolatile nonelectrolyte with a density of 1.26 g/mL at 25 C. Calculate the vapor pressure at 25 C of a solution made by adding 50.0 mL of glycerin to 500.0 mL of water. The vapor pressure of pure water at 25 C is 23.8 torr.
Boiling-Point Elevation • The increase in bp relative to that of the pure solvent (∆Tb) can be found using the following:
Tb
iK b m
• The value of Kb, the molal boiling-pointelevation constant, depends only on the solvent • m is the molal concentration of the solute • i is the van’t Hoff factor, the number of particles generated when the solute dissolves
Boiling-Point Elevation • Since the vapor pressure of the solvent is lowered, the temperature needed to increase the vapor pressure until it equals external pressure (i.e. the boiling point) is raised • Thus the boiling-point of a solution is higher than that of the pure solvent • This is known as boiling-point elevation
Freezing-Point Depression • When a solution freezes, solute particles “get in the way” and make it more difficult for solvent particles to form crystals • Because of this, the freezing point of a solution is lower than that of the pure solvent • This is known as freezing-point depression
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Freezing-Point Depression • The decrease in freezing-point, ∆Tf, is directly proportional to the molality of the solute:
Tf
iK f m
• Kf is the molal freezing-pointdepression constant
Sample Problem • Calculate the normal freezing point and boiling point of an aqueous solution of 0.50 m CaCl2. Kb and Kf for water are 0.51 C/m and 1.86 C/m, respectively.
Osmosis • Certain materials are semi-permeable, they allow certain molecules to pass through but not others • They often allow small molecules such as water to pass but block larger solute particles or ions
Sample Problem • Antifreeze consist of ethylene glycol (C2H6O2), a nonvolitile nonelectrolyte. Calculate the normal boiling point and freezing point of a 25.0% by mass solution of ethylene glycol in water. Kb and Kf for water are 0.51 C/m and 1.86 C/m, respectively.
Sample Problem • List the following aqueous solutions in order of their expected freezing points: 0.050 m CaCl2; 0.15 m NaCl; 0.10 m HCl; 0.050 m HC2H3O2; 0.10 m C12H22O11.
Osmosis • Consider a situation in which only solvent particles are allowed to pass through a membrane • If this membrane in placed between two solutions of different concentration, solvent will move through the membrane to equalize concentration • In this process, called osmosis, the net movement of solvent is always toward the more concentrated solution
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Osmosis
Pressure difference
Osmotic Pressure • The amount of pressure that must be applied to prevent osmosis is called osmotic pressure ( ) • The osmotic pressure can be calculated using a law that is similar to the ideal gas equation:
n RT V
MRT
• V is the volume of the solution, n is the number of moles of solute, R is the ideal gas constant, and T is Kelvin temperature
Sample Problem Sample Problem • What is the osmotic pressure at 20 C of a 0.0020 M sucrose (C12H22O11) solution?
• The osmotic pressure of a certain protein was measured in order to determine its molar mass. The solution contained 3.50 mg of protein dissolved in sufficient water to form 5.00 mL of solution. The osmotic pressure of the solution at 25 C was found to be 1.54 torr. Calculate the molar mass of the protein.
Colligative Properties
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Unit 12: Solutions • Topics Covered – General properties of aqueous solutions – Aqueous reactions – Solubility – Concentration – Colligative properties
• Reading Quiz – Read chapters 4 & 13 – Know definitions for all words in bold – Table 4.1 – Sample exercises 4.2, 4.3, 4.11, 4.12, 4.13, 4.14, 13.1, 13.2, 13.3, 13.4
Electrolytic Properties
General Properties of Aqueous Solutions • A solution is a homogenous mixture of two or more substances • The substance present in the greater quantity is usually called the solvent • The other substances in the solution are called the solutes
Electrolytic Properties
• A substance whose aqueous solutions contain ions (such as NaCl) is called an electrolyte • Electrolytic solutions conduct electricity • A substance that does not form ions in solution (such as C6H12O6) is a nonelectrolyte • Nonelectrolytic solutions do not conduct electricity
Strong vs. Weak Electrolytes
Ionic Compounds in Water
• Strong electrolytes are characterized by their nearly complete dissociation in water, while weak electrolytes dissociate to a much smaller extent
• In water, some ionic crystals will dissociate into their component ions • Water is a very effective solvent for many ionic compounds because of its polarity, which creates strong ion-dipole interactions that dissolve the ions into solution • Electrolytes
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Sample Problems
• From the diagrams shown, identify the strongest electrolyte and the weakest electrolyte
Precipitation Reactions • Reactions that result in the formation of an insoluble product are called precipitation reactions • Precipitation reactions occur when certain pairs of oppositely charged ions attract each other so strongly that they form an insoluble ionic compound • Example: – Pb(NO3)2(aq) + 2KI(aq)
PbI2(s) + 2KNO3(aq)
Molecular Compounds in Water • When a molecular compound dissolves in water, the solution usually consists of intact molecules dispersed throughout the solution • Some molecular substances ionize in water to form electrolytic solutions • The most important of these are acids • For example, HCl(g) ionizes in water to form HCl(aq), which is H+(aq) and Cl-(aq) ions
Solubility Guidelines for Ionic Compounds • The solubility of a substance is the amount of that substance that can be dissolved in a given quantity of solvent • For our purposes, any substance with solubility less than 0.01 mol/L will be referred to as insoluble • There are no simple rules that allow us to predict whether or not a compound will be soluble in water • We do have guidelines for predicting the solubility of many compounds • Just remember that all common ionic compounds of the alkali metals and the ammonium ion are soluble in water
• Precipitation reaction
Solubility Guidelines for Ionic Compounds
Sample Problem • Classify the following as soluble or insoluble in water: – – – –
Sodium carbonate Lead sulfate Cobalt(II) hydroxide Barium nitrate
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Exchange Reactions
Sample Problem
• Exchange reactions, double replacement, or metathesis reactions, are reactions in which ionic compounds exchange cations
• Which combinations of anions and cations would produce a precipitate as represented in the diagram?
Net Ionic Equations
PbI2(s) + 2KNO3(aq)
• An equation where all strong electrolytes are shown as ions is called a complete ionic equation – Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2I-(aq) + 2K+(aq) + 2NO3-(aq)
AgCl(s) +
• Precipitation reactions conform to this pattern, as do many acid-base reactions
Net Ionic Equations
• An equation that shows the complete chemical formulas of reactants and products is called a molecular equation – Pb(NO3)2(aq) + 2KI(aq)
– AX + BY AY + BX – AgNO3(aq) + KCl(aq) KNO3(aq)
PbI2(s)
• Ions that appear on both sides of a complete ionic equation are called spectator ions (shown in red)
Sample Problem • Write the net ionic equation for the precipitation reaction that occurs when solutions of calcium chloride and sodium carbonate are mixed
• When spectator ions are omitted, we are left with the net ionic equation – Pb2+(aq) + 2I-(aq)
PbI2(s)
• To write net ionic equations: – Write a balanced molecular equation – Write a complete ionic equation by separating strong electrolytes – Identify and cancel spectator ions
Sample Problem • Write the net ionic equation for the precipitation reaction that occurs when solutions of silver nitrate and potassium phosphate are mixed
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Acid-Base Reactions • Acids and bases that are strong electrolytes (completely ionized in solution) are called strong acids and strong bases • Those that are weak electrolytes (partly ionized) are called weak acids and weak bases
Acid-Base Reactions • When a solution of an acid and a base are mixed, a neutralization reaction occurs – HCl(aq) + NaOH(aq)
H2O(l) + NaCl(aq)
• A neutralization reaction between an acid and a metal hydroxide produces a water and a salt • The term salt is any ionic compound whose cation comes from a base and anion comes from an acid
Sample Problem • Write a balanced equation for the reaction between aqueous solutions of acetic acid (HC2H3O2) and barium hydroxide and write the net ionic equation
Oxidation-Reduction Reactions • Examples – Ca(s) + 2H+(aq) Ca2+(aq) + H2(g) – 2Ca(s) + O2(g) 2CaO(s)
• In the first case, calcium is being oxidized by the hydrogen ion, which is being reduced • The species being oxidized is often referred to as the reducing agent • The species being reduced can be called the oxidizing agent
Oxidation-Reduction Reactions • Reactions where electrons are transferred between reactants are called oxidation-reduction, or redox, reactions • Loss of electrons by a substance in called oxidation (LEO) • Gaining of electrons by a substance in called reduction (GER)
Oxidation Numbers • In order to identify the species being oxidized and reduced, you must first assign oxidation numbers to each element • Rules for assigning oxidation numbers – For an atom in elemental form the ON is always zero – For monatomic ions, the ON equals the charge of the ion
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Oxidation Numbers
Oxidation Numbers
• Rules for assigning ONs (continued)
• Rules for assigning ONs (continued)
– Nonmetals usually have negative ON, though they can sometimes be positive
– The sum of the ONs of all atoms in a neutral compound is zero – The sum of the ONs of all atoms in a polyatomic equals the charge of the ion
• The ON of oxygen is usually -2, the major exception is peroxides, O22-, where the ON is -1 • The ON of hydrogen is +1 when bonded to nonmetals and -1 when bonded to metals • The ON of fluorine is -1 in all compounds, the other halogens have ONs of -1 in most binary compounds except when combined with O
Sample Problem • Determine the oxidation state of sulfur in each of the following: H2S, S8, SCl2, Na2SO3, SO42-
Oxidation of Metals by Acids and Salts • The reaction of a metal with either an acid or a metal salt conforms to the following pattern: – A + BX AX + B – Zn(s) + 2HBr(aq) ZnBr2(aq) + H2(g) – Mn(s) + Pb(NO3)2(aq) Mn(NO3)2(aq) + Pb(s)
Sample Problem • What is the ON of the boldfaced element in each of the following: P2O5, NaH, Cr2O72-, SnBr4, BaO2
Sample Problem • Write a balanced molecular and net ionic equations for the reactions of aluminum with hydrobromic acid.
• These reactions are called displacement reactions or single replacement because the ion in solution is displaced or replaced through oxidation
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Sample Problem • Write a balanced molecular and net ionic equations for the reaction between magnesium and a solution of cobalt(II) sulfate. What is oxidized and what is reduced?
The Activity Series • A list of metals in order of decreasing ease of oxidation is called an activity series • Hydrogen is also included in this list • Any metal on the list can be oxidized by ions of the elements below it
Sample Problems • Will an aqueous solution of iron(II) chloride oxidize magnesium metal? • Which of the following metals will be oxidized by Ni(NO3)2: Zn, Cu, Fe?
Concentrations of Solutions • Concentration is the amount of solute dissolved in a given amount of solvent or solution • Molarity (M) expresses concentration as the number of moles of solute per liter of solution
Molarity
Sample Problem • Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate in enough water to form 125 mL of solution.
moles solute volumeof solution in liters
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Sample Problem • How many grams of glucose (C6H12O6) should be dissolved to make a 100. mL, 0.278 M solution?
Sample Problem • What is the molar concentration of K+ ions in a 0.015 M solution of potassium carbonate?
Sample Problem • What are the concentrations of each of the ions present in a 0.025 M aqueous solution of calcium nitrate?
Dilution • Solutions are often prepared for laboratory use in a concentrated form called a stock solution • The stock solution can then be diluted to the appropriate concentration by adding water • Moles solute before dilution = moles solute after dilution, therefore:
Mconc Vconc
Sample Problem • How many mL of 3.0 M H2SO4 are needed to make 450 mL of 0.10 M H2SO4?
Mdil Vdil
Sample Problem • If 10.0 mL of a 10.0 M stock solution of NaOH is diluted to 250 mL, what is the concentration of the resulting solution?
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Titration • To determine the concentration of a sample, chemists will often react it with a reagent of known concentration, called a standard solution, in a process called titration • Titrations can be performed for acid-base, precipitation, or redox reactions • The point at which stoichiometrically equivalent quantities are brought together is called the equivalence point
Sample Problem • The quantity of Cl- in a water supply is determined by titrating the sample with Ag+ according to the reaction Ag+(aq) + Cl-(aq) AgCl(s). How many grams of chloride ions are in a sample of the water if 20.2 mL of 0.100 M Ag+ solution is needed to react with all the chloride in the sample?
Sample Problem • A sample of 70.5 mg of potassium phosphate is added to 15.0 mL of 0.050 M of silver nitrate, resulting in the formation of a precipitate. (a) Write the molecular equation for the reaction. (b) What is the limiting reagent in the reaction? (c) Calculate the theoretical yield, in grams, of the precipitate that forms.
Titration • For acid-base titrations, a color changing indicator is used to determine when the reaction has reached its end point
Sample Problem • What is the molarity of an NaOH solution if 48.0 mL is needed to neutralize 35.0 mL of 0.144 M H2SO4?
Saturated Solutions and Solubility • A solution that is in equilibrium with undissolved solute is called saturated • Additional solute will not dissolve if added to a saturated solution • The amount of solute needed to form a saturated solution in a given quantity of solvent is known as solubility
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Saturated Solutions and Solubility • Under certain conditions it is sometimes possible to form solutions with a greater amount of solute than needed to form a saturated solution • Such solutions are called supersaturated • Supersaturated solutions are usually created by saturating a solution at high temperatures and then cooling it very slowly and carefully
Factors Affecting Solubility • The stronger the attractions between solute and solvent molecules, the greater the solubility • Polar liquids tend to dissolve readily in polar liquids, and nonpolar liquids tend to dissolve readily in nonpolar liquids • Nonpolar liquids tend to be insoluble in polar liquids • Pairs of liquids that mix in all proportions are called miscible, whereas those that do not are called immiscible • Like dissolves like
Sample Problem • Rank the following in terms of increasing solubility in water: C5H10(OH)2, C5H11Cl, C5H12, C5H10OH
Saturated Solutions and Solubility • Supersaturated solutions are unstable • The addition of a small seed crystal or simple agitation will cause the crystallization of the excess solute
Sample Problem • Predict whether each of the following substances is more likely to dissolve in carbon tetrachloride, CCl4, or water: C7H16, Na2SO4, HCl, and I2.
Pressure Effects • The solubility of solids and liquids are not affected by pressure • The solubility of a gas in any solvent is increased as the pressure over the solvent is increased
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Henry’s Law • The solubility of the gas increases in direct proportion to its partial pressure above the solution • This relationship is known as Henry’s law:
Sg
kPg
Sample Problem • Calculate the concentration of CO2 in a soft drink that is bottled with a partial pressure of CO2 of 4.0 atm over the liquid at 25 C. The Henry’s law constant for CO2 in water at this temperature is 3.1 x 10-2 mol/L-atm.
• Sg is the solubility of the gas, usually in molarity, k is called the Henry’s law constant, and Pg is the partial pressure of the gas over the solution
Sample Problem
Temperature Effects
• Calculate the concentration of CO2 in a soft drink after the bottle is opened and equilibrates at 25 C under a CO2 partial pressure of 3.0 x 10-4 atm.
• The solubility of most solid solutes in water increases as the temperature of the solution increases
Temperature Effects
Mass Percentage • One of the simplest ways to express concentration is mass percentage: Mass% of component
• Most gases become less soluble in water as the temperature rises
mass of component in solution 100% total mass of solution
• For example, a solution that is 36% HCl by mass contains 36 g of HCl for every 100 g of solution
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Parts Per Million • For very dilute solutions, concentration is often expressed in parts per million (ppm): ppm of component
Sample Problem • A solution is made by dissolving 13.5 g of glucose in 0.100 kg of water. What is the mass percentage of solute in the solution?
mass of component in solution 106 total mass of solution
• A solution whose solute concentration is 1 ppm contains 1 g of solute for every million (106) g of solution
Sample Problem • A 2.5 g sample of groundwater was found to contain 5.4 g of Zn2+. What is the concentration in ppm?
Mole Fraction • Recall that the mole fraction (X) of a component of a solution is: X of component
moles of component total moles of all components
• Thus a solution containing 1.00 mol of HCl in 8.00 mol of water has a mole fraction of XHCl = (1.00 mol)/(1.00 mol + 8.00 mol) = 0.111 (no units)
Molality • The molality (m) of a solution equals the number of moles of solute per kg of solvent: Molality
Sample Problem • A solution is made by dissolving 4.35 g glucose in 25.0 mL of water. Calculate the molality of glucose in the solution.
moles solute kg solvent
• This is different from molarity, which depends on liters of solution instead of kg of solvent
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Sample Problem • What is the molality of a solution made by dissolving 36.5 g of napthalene (C10H8) in 425 g of toluene (C7H8)?
Sample Problem • A commercial bleach solution contains 3.62 mass % NaOCl in water. (a) Calculate the molality of NaOCl in the solution. (b) Calculate the mole fraction.
Sample Problem • A solution containing equal masses of glycerol (C3H8O3) and water has a density of 1.10 g/mL. Calculate (a) the molality of glycerol, (b) the mole fraction of glycerol, and (c) the molarity of glycerol in the solution.
Sample Problem • A solution of HCl contains 36% HCl by mass. (a) Calculate the mole fraction of HCl in the solution. (b) Calculate the molality of HCl in the solution.
Sample Problems • A solution contains 5.0 g of toluene (C7H8) and 225 g of benzene and has a density of 0.876 g/mL. Calculate the molarity.
Colligative Properties • Physical properties of solutions can vary greatly from those of the pure solvent • Lowering vapor pressure and freezing point and raising the boiling point are properties of solutions that do not depend on the type or identity of solute particles but only on the quantity (concentration) of solute particles • These properties are known as colligative properties
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Lowering Vapor Pressure
Lowering Vapor Pressure
• Adding a nonvolatile solute always lowers the vapor pressure of a solvent
• The extent to which this happens depends on the concentration of the solute and follows Raoult’s law:
PA
X A PA
• The partial pressure exerted by the solvent above a solution, PA, equals the product of the mole fraction of the solvent, XA, times the vapor pressure of the pure solvent, PA
Sample Problem • Glycerin (C3H8O3) is a nonvolatile nonelectrolyte with a density of 1.26 g/mL at 25 C. Calculate the vapor pressure at 25 C of a solution made by adding 50.0 mL of glycerin to 500.0 mL of water. The vapor pressure of pure water at 25 C is 23.8 torr.
Boiling-Point Elevation • The increase in bp relative to that of the pure solvent (∆Tb) can be found using the following:
Tb
iK b m
• The value of Kb, the molal boiling-pointelevation constant, depends only on the solvent • m is the molal concentration of the solute • i is the van’t Hoff factor, the number of particles generated when the solute dissolves
Boiling-Point Elevation • Since the vapor pressure of the solvent is lowered, the temperature needed to increase the vapor pressure until it equals external pressure (i.e. the boiling point) is raised • Thus the boiling-point of a solution is higher than that of the pure solvent • This is known as boiling-point elevation
Freezing-Point Depression • When a solution freezes, solute particles “get in the way” and make it more difficult for solvent particles to form crystals • Because of this, the freezing point of a solution is lower than that of the pure solvent • This is known as freezing-point depression
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Freezing-Point Depression • The decrease in freezing-point, ∆Tf, is directly proportional to the molality of the solute:
Tf
iK f m
• Kf is the molal freezing-pointdepression constant
Sample Problem • Calculate the normal freezing point and boiling point of an aqueous solution of 0.50 m CaCl2. Kb and Kf for water are 0.51 C/m and 1.86 C/m, respectively.
Osmosis • Certain materials are semi-permeable, they allow certain molecules to pass through but not others • They often allow small molecules such as water to pass but block larger solute particles or ions
Sample Problem • Antifreeze consist of ethylene glycol (C2H6O2), a nonvolitile nonelectrolyte. Calculate the normal boiling point and freezing point of a 25.0% by mass solution of ethylene glycol in water. Kb and Kf for water are 0.51 C/m and 1.86 C/m, respectively.
Sample Problem • List the following aqueous solutions in order of their expected freezing points: 0.050 m CaCl2; 0.15 m NaCl; 0.10 m HCl; 0.050 m HC2H3O2; 0.10 m C12H22O11.
Osmosis • Consider a situation in which only solvent particles are allowed to pass through a membrane • If this membrane in placed between two solutions of different concentration, solvent will move through the membrane to equalize concentration • In this process, called osmosis, the net movement of solvent is always toward the more concentrated solution
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Osmosis
Pressure difference
Osmotic Pressure • The amount of pressure that must be applied to prevent osmosis is called osmotic pressure ( ) • The osmotic pressure can be calculated using a law that is similar to the ideal gas equation:
n RT V
MRT
• V is the volume of the solution, n is the number of moles of solute, R is the ideal gas constant, and T is Kelvin temperature
Sample Problem Sample Problem • What is the osmotic pressure at 20 C of a 0.0020 M sucrose (C12H22O11) solution?
• The osmotic pressure of a certain protein was measured in order to determine its molar mass. The solution contained 3.50 mg of protein dissolved in sufficient water to form 5.00 mL of solution. The osmotic pressure of the solution at 25 C was found to be 1.54 torr. Calculate the molar mass of the protein.
Colligative Properties
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