chapter 4: solution chemistry and the hydrosphere problems: 4.1-4.80, 4.85-4.96, 4.99-4.100,...

Chapter 4:Solution Chemistry and

the Hydrosphere

Problems: 4.1-4.80, 4.85-4.96, 4.99-4.100, 4.111-4.113, 4.119-4.120, 4.129,

4.131-4.132

Solutions on Earth and Other Places

aqueous solution: a solution where water is the dissolving medium (the solvent)• For example, when table salt (NaCl) is

dissolved in water, it results in an aqueous solution of sodium chloride, NaCl(aq), with Na+ and Cl- ions dissolved in water.

• Note: The physical state aqueous,(aq), indicates an element or compound dissolved in water while the physical state liquid,(l), means a pure substance in the liquid state.– Thus, NaCl(aq) NaCl(l), which is molten

NaCl requiring very high temperatures.A solution consists of a solute dissolved in a solvent.

Solutionssolute: component present in smaller amountsolvent: component present in greater amountThe formation of a solution:As a solute crystal is dropped into a solution, the water molecules begin to pull apart the ionic compound ion by ionSolvent molecules surround the solute particles, forming a solvent cage

– the ions are now hydrated (surrounded by polar water molecules)

– solute is now dissolved in the solvent and cannot be seen because the ions are far apart, like the particles in a gas

Unsaturated, Saturated and Supersaturated

Solutions• In general, if a solid is soluble in a solvent,

more solid dissolves in the solvent at higher temperatures.

unsaturated: contains less than the maximum amount of solute that a solvent can hold at a specific temperaturesaturated: contains the maximum amount of solute that a solvent can hold at a specific temperaturesupersaturated: contains more than the maximum amount of solute that a solvent should be able to hold at specific temperature

• A supersaturated NaC2H3O2 solution recrystallizing after addition of more solute:


Solutions

• How can a solution hold more solute than it should be able to hold? – If a given amount of solute is dissolved in a

solvent at a higher temperature, and the solution is allowed to cool without being disturbed, the solute will remain in solution.

• But the solution is unstable, and the solute will come out of solution (i.e. recrystallize) if the solution is disturbed (e.g. by adding more solute, scratching the glass, etc.)


Solutions

For some substances, recrystallization is exothermic, releasing heat to the surroundings.

– Hot packs used to warm hands and feet in winter (though some of these are oxidation reactions, which we will discuss later)

For other substances, recrystallization is endothermic, absorbing heat and making the surroundings colder.

– Cold packs used for sports injuries


Solutions

How do we measure concentration?

solution: homogeneous mixture of substances present as atoms, ions, and/or molecules solute: component present in smaller amount solvent: component present in greater amount Note: Unless otherwise stated, the solvent for most solutions considered in this class will almost always be water!

Aqueous solutions are solutions in which water is the solvent.

• A concentrated solution has a large quantity of solute present for a given amount of solution.

• A dilute solution has a small quantity of solute present for a given amount of solution.

SOLUTION CONCENTRATION = The more solute in a given amount of solution the more concentrated the solution

Example: Explain the difference between the density of pure ethanol and the concentration of an ethanol solution.


Concentration can be measured a number of ways:• ppm (parts per million) – one part in a million

parts• ppb (parts per billion) – one part in a billion

parts• g/kg (grams per kilogram) – one gram solute

per one kilogram of solventThe chemical standard most used is Molarity

Molarity =

units: M (molar = mol/L)

We’ll come back to concentration later in the chapter…


Evidence of a Chemical Reaction

a) A gas is produced. b) A precipitate forms. c) Heat is released or absorbed

Types of Chemical Reactions

• Precipitation Reactions • Acid-Base Neutralization Reaction• Oxidation-Reduction (Redox) Reactions

– Further classified as:• Combination• Decomposition• Combustion• Single-replacement reactions

Precipitation Reactions• Solubility Rules: Indicate if an ionic compound is

soluble or insoluble in water.• Keep in mind that these are just general guidelines,

and in reality, some ionic compounds are slightly soluble, and solubility may depend on the temperature.

Solubility Rules for Ionic Compounds in Water

Soluble if the ionic compound contains: • Li+, Na+, K+, NH4

+ (ALWAYS!)• C2H3O2

–, NO3–, ClO3

–, ClO4–

• Halide ions (X–): Cl–, Br–, or I–, but AgX, PbX2, HgX, and Hg2X2 are insoluble

• sulfate ion (SO42-), but CaSO4,

SrSO4, BaSO4, Ag2SO4, `and PbSO4 are insoluble.

Insoluble if the ionic compound contains: • carbonate ion, CO3

2-

• chromate ion, CrO42-

• phosphate ion, PO43-

• sulfide ion (S2–), but CaS, SrS, and BaS are all soluble.

• hydroxide ion (OH–), but Ca(OH)2, Sr(OH)2, and Ba(OH)2 are soluble.

Precipitation Reactionssoluble = compound dissolves in water exists as individual ions in solution

physical state is aqueous, (aq)Insoluble = compound does not dissolve in water but remains a solid

physical state is shown as solid, (s)

Precipitation Reactions

• Example 1: Use the Solubility Rules and identify the ionic compounds are soluble or insoluble by indicating the physical state of each compound.

a. NaCl d. LiOH g. Mg(OH)2 j. Ag3PO4

b. MgS e. CaS h. SrSO4 k. BaCO3

c.K3PO4 f. Li2CrO4 i. Na2CO3l.

(NH4)2CrO4


• Example 1: Use the Solubility Rules and identify the ionic compounds are soluble or insoluble by indicating the physical state of each compound.




(NH4)2CrO4

Soluble (aq)

Insoluble (s)




(NH4)2CrO4


• In a precipitation reaction, two solutions react to form a precipitate (an insoluble solid):AX(aq) + BZ(aq) AZ(s) + BX(aq)

precipitate

For example: KI (aq)+ Pb(NO3)2(aq) PbI2 (s) + KNO3 (aq)

Precipitation ReactionsTo balance and complete the precipitation reactions:1. Exchange the anions, writing the formulas for the products based on the charges of the ions!2. Use the Solubility Rules to determine if each product is soluble or insoluble.

– If at least one product is insoluble, a precipitation reaction has occurred. Write the formulas for both products, indicating the precipitate as (s), then balance the equation.

– If both products are soluble, write NR (=no reaction).

3. Keep in mind that the charges on ions do NOT change in precipitation reactions. For metals that can form more than one charge, use the charge on the metal ion from the reactant side of the equation.


Ex 1. MgSO4(aq) + NaOH(aq)

Ex 2. K2CO3(aq) + AlCl3(aq)

Ex 3. SrBr2(aq) + Zn(NO3)2(aq)

Ex 4. CuSO4(aq) + NaOH(aq)

Ex 5. KI(aq) + Pb(NO3)2(aq)

Acid-Base Neutralization Reactions: Proton Transfer

Properties of Acids and Bases

Acids Bases–produce hydrogen ions, H+ –produce hydroxide ions, OH–

–taste sour –taste bitter; feel soapy, slippery

–turn blue litmus paper red –turn red litmus paper blue

Arrhenius Definitions

acid: A substance that releases H+ when dissolved in water

– Some acids are monoprotic (release only H+ per molecule)• e.g. HCl, HBr, HI, HNO3, HClO4

– Some acids are polyprotic (release more than on H+ per molecule)• e.g. H2SO4 and H2CO3 are both diprotic;

H3PO4 is triprotic.

base: A substance that releases OH– when dissolved in water

Acid-Base Reactions

In an acid-base reaction, • H+ from acid reacts with the OH– from base to form

water, H2O• The cation (M+) from base combines with anion

from acid (X–) to form a salt. A general equation for an acid-base neutralization reaction is shown below:

HX(aq) + MOH(aq) H2O(l) + MX

acid base water salt

Because water is always produced, an acid always reacts with a base!

Examples

Complete and balance each of the equations below:

a. HCl(aq) + NaOH(aq)

b. H2SO4(aq) + KOH(aq)

c. H3PO4(aq) + Ca(OH)2(aq)

Acid-Base Reactions with Gas Formation

Some acid-base reactions produce carbon dioxide gas, CO2(g), along with water and salt.

When the base contains carbonate ion (CO32–) or

hydrogen carbonate ion (HCO3–), then the products of

the acid-base reaction are water, carbon dioxide gas, and a salt.

The general equations for the unbalanced acid-base reactions are:

HX(aq) + MCO3(s) H2O(l) + CO2(g) + MX

acid base water carbon dioxide salt HX(aq) + MHCO3(s) H2O(l) + CO2(g) + MX

acid base water carbon dioxide salt

Because water is always produced, an acid always reacts with a base!

Complete and balance each of the equations below:

a. HCl(aq) + Na2CO3(s)

b. HNO3(aq) + CaCO3(s)

c. H2SO4(aq) + KHCO3(s)

d. HClO4(aq) + Sr(HCO3)2(s)

Acid-Base Reactions with Gas Formation

A double-replacement reaction that produces NH4OH(aq) actually produces ammonia, NH3(g).

NH4OH(aq) NH3(g) + H2O(l)

Example: Complete and balance the equation below: (NH4)2SO4(aq) + KOH(aq)

Brønsted-Lowry Definition of Acids and Bases

• Brønsted-Lowry acid: A substance that donates a proton (H+)—i.e., a proton donor

• Brønsted-Lowry base: A substance that accepts a proton (H+)—i.e., a proton acceptor

• Unlike an Arrhenius base, a Brønsted-Lowry base does not need to contain OH–.

Why is H+ called a proton?

A Brønsted-Lowry acid-base reaction simply involves a proton (H+) transfer.

NH3(aq) + H2O(l) ⇄ NH4+(aq) + OH–

(aq)

Brønsted-Lowry Acids and Bases

Note: This reaction simply involves H2O donating a H+ ion to NH3 to produce NH4

+ and OH–.• In this reaction, H2O is the Brønsted-Lowry acid, and NH3 is the

Brønsted-Lowry base.• The conjugate acid-base pairs differ only by a H+.• In this reaction, the conjugate acid-base pairs are NH3 and NH4

+ and H2O and OH–.

Conjugate Acid-Base Pairs

Conjugate acid-base pairs: a Brønsted-Lowry acid/base and its conjugate differ by a H+

HA(aq) + H2O(l) ⇄ H3O+(aq) + A–(aq)

• For the reaction above, when HA donates H+ to H2O, it leaves behind A–, which can act as a base for the reverse reaction.

• An acid and base that differ only by the presence of H+ are conjugate acid-base pairs.

• The general reaction for the dissociation (or ionization) of an acid can be represented as above, where the double-arrow indicates both the forward and reverse reactions can occur.

• Note: The double arrow (⇄) indicates the reaction is reversible (goes in both directions).

Determine the Brønsted-Lowry acid and base in each of the following reactions: a. CH3COOH(aq) + NH3(aq) ⇄ NH4

+(aq) + CH3COO–(aq)

b. NH3(aq) + H2O(l) ⇄ NH4+(aq) + OH–(aq)

c. H2O(l) + H2SO4(aq) ⇄ H3O+(aq) + HSO4

–

(aq)

Conjugate Acid-Base Pairs

Oxidation-Reduction (Redox) Reactions

Types of Redox Reactions

• Combination Reaction• Decomposition Reaction• Single-Replacement (or Displacement)

Reaction• Combustion Reaction

Combination Reactions

A + Z AZ

Usually a meal and a non-metal react to form a solid ionic compound: metal + nonmetal ionic compound(s)

Δ

Combination Reactions:A + Z AZ

Complete and balance each of the equations below: a. Na(s) + Cl2(g)

b. Al(s) + O2(g)

c. Zn(s) + S8(s)

d. Mg(s) + N2(g)

Δ

Δ

Δ

Δ

Decomposition Reactions:AZ A + Z

Be able to classify and balance decomposition reactions. (You won’t need to predict products.) a. ___ KHCO3(s) ___ K2CO3(s) + ___ H2O(l) + ___ CO2(g)

b. ___ Al2(CO3)3(s) _____ Al2O3(s) + _____ CO2(g)

c. ___ KClO3(s) _____ KCl(s) + _____ O2(g)

Δ

Δ

Δ, MnO2

Single-Replacement Reactions and the Activity

SeriesActivity Series: Relative order of elements arranged by their ability to replace cations in aqueous solution Li > K > Ba > Sr > Ca > Na > Mg > Al > Mn >

Zn > Fe > Cd > Co > Ni > Sn > Pb > (H) > Cu > Ag >

Au Note: The Activity Series will be given to you on quizzes and exams.

Single Replacement Reactions:

A + BZ AZ + Bmetal A + aqueous solution B aqueous solution A + metal B

Zn(s) + Sn2+(aq) Sn(s) + Zn2+(aq)Cu(s) + 2 Ag+(aq) 2 Ag(s) + Cu2+(aq)

To balance and complete the following rxns:

• Check the Activity Series to see which metal is more active.– The more active metal replaces the less active

by going into solution as an ion and the less active metal ion comes out as a solid.

1. Mg(s) + AlCl3(aq)

2. Al(s) + CdSO4(aq)

3. Cd(s) + AgNO3(aq)

4. Ag(s) + Mg(NO3)2(aq)

Activity Series: Li > K > Ba > Sr > Ca > Na > Mg > Al > Mn > Zn > Fe > Cd > Co > Ni > Sn > Pb > (H) > Cu

> Ag > Au

To balance and complete the following reactions:

• Check the Activity Series to see which metal is more active, the metal or H.– The more active metal replaces the less active

by going into solution as an ion and the H comes out as hydrogen gas, H2(g).

metal A + acid solution aqueous solution A + H2(g)

1. Zn(s) + HCl(aq) 2. Al(s) + HNO3(aq)

3. Cu(s) + HI(aq)

Activity Series: Li > K > Ba > Sr > Ca > Na > Mg > Al > Mn > Zn > Fe > Cd > Co > Ni > Sn > Pb > (H) > Cu

> Ag > Au

Active Metals: Li > K> Ba > Sr > Ca >

NaActive metals react directly with water:The active metal replaces the less active by going into solution as an ion with hydroxide, OH–, and the H comes out as hydrogen gas, H2(g).

active metal + H2O(l) metal hydroxide + H2(g)

1. Ca(s) + H2O(l)

2. Na(s) + H2O(l)

3. Fe(s) + H2O(l)

Combustion Reactions

CxHy + O2(g) H2O(g) + CO2(g)

CxHyOz + O2(g) H2O(g) + CO2(g)

1. C3H8(g) + O2(g)

2. C6H6O(l) + O2(g)

3. C2H2(g) + O2(g)

Δ

Δ

Δ

Identify the ReactionsFor each of the following,

1. Identify the type of reaction using the letters designated below: Combination (C) Precipitation (P) Decomposition (D) Acid-Base Neutralization (N) Combustion (B) Single Replacement/Displacement (SR)

2. Balance the equations:

____a. ___ Mg(NO3)2(aq) + ___ K3PO4(aq) ___ Mg3(PO4)2(s) + ___ KNO3(aq)

____b. ___ Ni(OH)3(s) + ___ HCl(aq) ___ H2O(l) + ___ NiCl3(aq)

____c. ___ Al(HCO3)3(aq) ___ CO2(g) + ___ H2O(g) + ___ Al2(CO3)3(s)

____d. ___ Fe(s) + ___ Pb(NO3)2(aq) ___ Pb(s) + ___ Fe(NO3)3(aq)

Δ

Electrolytes and Non-Electrolytes

Electrolytes and electrical conductivity• If a solution conducts electricity, it contains ions• A solution that contains many ions is a strong

electrolyte.• A solution that contains only a few ions is a weak

electrolyte.• A solution that contains only a no ions is a

nonelectrolyte.

• A solution that contains many ions is a strong electrolyte.Light bulb burns brightly in a light bulb

conductivity apparatus.• A solution that contains only a few ions is a

weak electrolyte.Light bulb burns dimly in a light bulb

conductivity apparatus.• A solution that contains only a no ions is a

nonelectrolyte.Light bulb does not light in a light bulb

conductivity apparatus.

Electrolytes and Non-Electrolytes

Strong Electrolytes

strong electrolytes: substances that are good conductors of electricity• These substances break up to produce many ions

in water– many ions present to move electrons/conduct

electricity strong electrolyte For example,

NaCl(s) Na+(aq) + Cl–(aq) KOH(s) K+(aq) + OH–(aq) HBr(aq) H+(aq) + Br–(aq) Examples: strong acids, strong bases, all soluble ionic compounds

H2O

H2O

H2O

Weak Electrolytes

weak electrolytes: substances that are weak/poor conductors of electricity• These substances mostly remain intact as

compounds, producing very few ions in water– only a few ions present to move

electrons/conduct electricity weak electrolyte

For example,Mg(OH)2(s) Mg(OH)2(s)

HNO2(aq) HNO2(aq)

Examples: weak acids, weak bases, insoluble ionic compounds

H2O

Non-electrolytes

nonelectrolytes: substances that cannot conduct electricity• These molecules never break down into ions.

– They always remain intact as neutral molecules that have no charge no ions to move electrons/conduct electricity

For example,C12H22O11(s) C12H22O11(aq)

Examples: sugar (e.g. sucrose), ethanol (C2H5OH), and all other molecules that are not acids

H2O

Acids and BasesKnow the following acids and bases. All other acids and bases are weak!

Strong acids and bases dissolve in water to form ions (or species) in solution.

HNO3(aq) H+(aq) + NO3–(aq)

Ca(OH)2(aq) Ca2+(aq) + 2 OH–(aq)

Note: H2SO4(aq) is a strong acid and diprotic (able to release 2 H+ ions), but it generally ionizes to release only one H+ ion in water: H2SO4(aq) H+(aq) + HSO4

–(aq)

Recognize that both protons are not released in water!

Strong Acids Strong Bases

HCl, HBr , HI, HNO3, HClO4, H2SO4

LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2,

Ba(OH)2

Molecular, Ionic and Net Ionic Equations

• molecular equation: equation showing reactants and products as compounds

• total/complete ionic equation: shows strong electrolytes as individual ions while all solids, liquids, gases, and weak electrolytes remain intact as compounds

• spectator ions: ions that do not form solids, liquids, gases, weak electrolytes – appear on both sides of total ionic equation as ions

• net ionic equations: show only solids, liquids, gases, weak electrolytes (weak acids and weak bases), and ions undergoing a chemical change/reaction – excludes spectator ions

Net Ionic EquationsGuidelines for Writing Net Ionic Equations1. Balance the chemical/molecular equation.2. Convert the molecular equation to total ionic

equation– Leave solids, liquids, gases, and weak acids and

bases as compounds– Show strong acids and all aqueous ionic

compounds as ions in solution.3. Cancel spectator ions to get net ionic equation

– If canceling spectator ions eliminates all ions NO REACTION (NR)

– If coefficients can be simplified, do so to get the lowest ratio.

4. Make sure total charges on both sides of the equation are equal.

chapter 4: solution chemistry and the hydrosphere problems: 4.1-4.80, 4.85-4.96, 4.99-4.100,...

Documents

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