chapter 4: solution chemistry and the...

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.

Solutions solute: component present in smaller amount

solvent: component present in greater amount

The 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 ion

Solvent 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

temperature

saturated: contains the maximum amount of solute

that a solvent can hold at a specific temperature

supersaturated: 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:

Unsaturated, Saturated and

Supersaturated 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.)

Unsaturated, Saturated and

Supersaturated 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

Unsaturated, Saturated and

Supersaturated 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 = amount of soluteamount of solvent

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.

How do we measure

concentration?

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 solvent

The chemical standard most used is Molarity

Molarity = moles of soluteliters of solution

units: M (molar = mol/L)

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

How do we measure

concentration?

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, CO32-

• 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 Reactions soluble = 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. Na2CO3 l. (NH4)2CrO4

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. Na2CO3 l. (NH4)2CrO4

Soluble

(aq)

Insoluble

(s)

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

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

c. K3PO4 f. Li2CrO4 i. Na2CO3 l. (NH4)2CrO4

Precipitation Reactions

• 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 Reactions To 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.

Precipitation Reactions

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 Series

Activity 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 + B metal 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 > Na

Active 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 Reactions For 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 Bases Know 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 Equations

Guidelines for Writing Net Ionic Equations

1. 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.