Key Concepts of Chapter Components of a solution. Electrolytes, Weak Electrolytes, Nonelectrolytes...

Key Concepts of Chapter

• Components of a solution.

• Electrolytes, Weak Electrolytes, Nonelectrolytes

• Precipitation Reactions and Solubility Guidelines

• Molecular, Complete Ionic, Net Ionic Equations

• A Few Concepts Related to Acids and Bases

• Redox Reactions and Determining Oxidation Numbers

• Activity Series of Metals

• Molarity / Solution Preparation / Dilution

Properties of water• High melting and boiling point.• Expands upon freezing.• Dissolves a variety of substances easily.• Pure water is rare in nature.• These properties are mainly due to polarity, which will be

discussed a bit later.

Most solids—molecules are tightly packed, and therefore they contract upon freezing.

Ice—molecules take on an open hexagonal arrangement, and therefore it expands upon freezing. (More about this in Chapter 11)

Reactions in aqueous solutions• Aqueous solution:

– Solution in which water is the solvent (dissolving agent).

• 3 major types of chemical processes of aqueous solutions:– Precipitation reactions

– Acid-base reactions

– Redox reactions

• Solution:–Homogenous mixture of 2 or

more substances.

• Solvent:–Dissolving medium, usually

present in greater quantity.

• Solute:–The other substance(s) in the

solution.

• Electrolyte:– A substance whose aqueous solution

forms ions; conducts electricity.– Ionic compounds.

• Nonelectrolyte:– Substance that does not form ions in

an aqueous solution; poor conductor.– Molecular compounds.

• Ionic compounds in water:– Dissociate into its component

ions.

dissolving_NaCl_probe.swf

A closer Look

Unequal sharing of electrons leads to partial positive and negative charges in a water molecule. These charges attract the ions which causes dissociation of the ionic compound in water.

Do not get to wrapped up in the difference between the terms ionization and dissociation. Consider them to mean the same thing, the separation

of a substances ions.

Equations showing ionization or dissociation

)()()( aqClaqNaaqNaCl

)(2)()( 22 aqClaqCuaqCuCl

)()(3)( 343 4 aqPOaqNaaqPONa

Molecular compounds in H2O

Molecular compounds – nonmetal + nonmetal

Structural integrity of molecule is usually maintained meaning no ions form (C12H22O11)

Exception:Some molecular solutes interact with water to form ions. These would be electrolytes.

Examples: Acids HCl, H2C3O2 Ammonia NH3

dissolving_sugar_probe.swf

Strong Electrolytes

• Exists in solution completely or almost completely as ions.

• All ionic compounds and a few molecular compounds.(Ex: Strong Acids)

)()()( aqaqaq ClHHCl

)()()( aqaqs ClNaNaCl

Weak Electrolytes• Molecular compounds that produce a

small concentration of ions when dissolved in H2O.

Ex: Acetic acid (HC2H3O2) only slightly ionizes when dissolved in water.HC2H3O2(aq) H+

(aq)+ C2H3O2-(aq)

Weak acids are better conductors if they are dilute, as you will see in lab. Explain.

• Reactions that result in an insoluble product.

• Insoluble:–Substance with solubility less than

0.01 mol/L–Water molecules cannot overcome

the attraction between the ions.

KI (aq) + Pb(NO3)2 (aq) PbI2(s) + KNO3 (aq)

You must be able to determine whether a substance is soluble in water by simple examination of the chemical formula.

To do so, you must memorize how specific polyatomic ions act in water.

Not as hard as it sounds. We will focus mainly on 10 anions. This will give you the tools to predict the

solubility of many compounds.

Solubility of Ionic Compounds

• All acetates and nitrates are soluble in water.

• All ionic compounds of alkali metals and ammonium are soluble.

• Table on next slide is on page 118 of your textbook.

Cl-

Acetate ion C2H3O2-

Br-

I-

NO3-

SO42-

Soluble

Exceptions

Ag+, Hg22+, Pb2+

None

Ag+, Hg22+, Pb2+

Ag+, Hg22+, Pb2+

None

Sr2+, Ba2+, Hg22+, Pb2+

CO32-

OH-

PO43-

S2-

Insoluble

Exceptions

NH4+ and Group 1 metals

Group 1 metals and Ba2+, Sr2+, Ca2+

NH4+ and Group 1 metals

NH4+ and Group 1 metals

and Ba2+, Sr2+, Ca2+

Equation Types• Molecular

• Complete Ionic

• Net ionic equation

)3)(2)(23 22)( aqsaq KNOPbIKINOPb

)(3)()(2)()(2

)( 22222)( aqaqsaqaqaq NOKPbIIKNOPb

aq

)(2)(2

)( 2 saqaq PbIIPb

3

Ionic Equations

)()()(3

2)( 222 aqaqaqaq IKNOPb )(3)()(2 22 aqaqs NOKPbI

•Those ions that appear on both sides of a complete ionic equation are known as Spectator Ions.

•Net ionic equations do not include spectator ions.

• Exchange Reactions

• Metathesis reactions

• Double displacement

• Double replacement

BXAYBYAX

)(3)()()(3 aqsaqaq KNOAgClKClAgNO

Writing Net Ionic Equations

1) Write a balanced molecular equation.

2) Rewrite the equation showing ions of strong electrolytes only.

3) Identify and cancel all spectator ions.

Acid-Base Reactions

Acids:• Ionize in H2O, causes

increase in H+ ions.

• H+ ions are bare protons.

• Acids are proton donors.

All this acid rain is

killing my complexion!

• Monoprotic Acids: (HCl, HNO3)

–Acids that can only yield one H+ per molecule upon ionization.

HCl H+ + Cl-

Diprotic Acids: (H2SO4)Ionization occurs in 2 steps.

Only the first ionization is complete.

)(4)()(42 aqaqaq HSOHSOH

2)(4)()(4

aqaqaq SOHHSO

Is HF a weak or strong acid?

weak acid

Although it is a weak acid, this acid is extremely reactive because of the F- ion.

Must be kept in special polypropylene container because it eats through glass.

Used to etch glass.

Has caused major accidents in lab.

• Substances that increase the OH- when added to water. (NaOH)

• NH3 is a base. In water it accepts an H+ ion from HOH, leaving an OH- in solution.– NH3 is a weak electrolyte– About 1% ionizes to form NH4

+/OH-

Strong acids and bases

• Acids and bases that ionize completely in solution are strong acids and bases.

• Those that only ionize partially are weak acids and bases.

• You must memorize these.

Strong Acids

Hydrochloric Acid –

Hydrobromic Acid –

Hydroiodic Acid –

Nitric Acid –

Sulfuric Acid –

Chloric Acid –

Perchloric Acid –

HCl

HBr

HI

HNO3

H2SO4

HClO3

HClO4

Strong Bases

All group 1 Metal Hydroxides

(LiOH, NaOH, KOH, RbOH, CsOH)

Heavy Group 2 Metal Hydroxides

Ca(OH)2, Sr(OH)2, Ba(OH)2

Once you memorize the strong acids and bases, you will have enough information to determine if a

substance is a strong or weak electrolyte.

Soluble

Ionic?

YES NOStrong

Electrolyte Acid?

YES NO

Strong

Electrolyte

StrongAcid?

YES

NH3 ?

NO

Weak

Electrolyte

YES

Nonelectrolyte

NO

Another Flowchart (Choose the one you like for practice)

Acid/base?

Yes No

Strong or weak acid/base?

Ionic or Molecular Compound?

Strong Weak

Strong Electrolyte Weak Electrolyte

Ionic

(metal + nonmetal)

Molecular

(nonmetals)

Soluble?

Yes

Strong Electrolyte

nonelectrolyte

Example problems:KFNa3PO4

NH3

CH3CH2OHHClNO2

HC2H3O2

CH4

NH4ClCH3Cl

strong electrolyte

strong electrolyte

strong electrolyte

strong electrolyte

weak electrolyte

nonelectrolyte

nonelectrolyte

nonelectrolyte

nonelectrolyte

weak electrolyte

Handout: Flowchart

Homework:

Memorize Strong Acids and Bases

4.16 - 4.24

4.27 – 4.28

worksheet

Acid Propertieso Sour tasteo Turn blue litmus

redo pH < 7

Base propertieso Bitter tasteo Turns red litmus

blueo pH >7o slippery

Some Properties of Acids and Bases

Acid + Base Neutralization

• Products of a neutralization reaction have none of the properties of an acid or a base.

• An acid reacts with a metal hydroxide to form a salt plus water.

Neutralization Reactions

• Acid + Base (Metal Hydroxide) Salt + Water

• HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

• H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O(l)

H+ + OH- H2O(l)

Potassium Hydroxide + Sulfuric Acid

• Ionic equation:

• Net Ionic equation:

2)(4)()()( 222 aqaqaqaq SOHOHK 2

)(4)()( 22 aqaql SOKHOH

)(2)()( laqaq OHOHH

Write the net ionic equation for the following reaction. It might help to first write the molecular

equation, and then the complete ionic equation, followed by the net ionic equation.

Neutralization Reaction of Weak Acid

HC2H3O2(aq) + NaOH(aq) NaC2H3O2(aq) + H2O(l)

Weak acid strong base soluble salt water

HC2H3O2 + Na+ + OH- Na+ + C2H3O2- + H2O(l)

HC2H3O2(aq) + OH-(aq) C2H3O2- (aq)+ H2O(l)

*Remember, only strong electrolytes are written as ions.*

Acetic Acid + Sodium Hydroxide

Acid/Base Rx’s with gas formation

• Other bases besides OH- react with H+ to form molecular compounds.Two common bases are CO3

-2 and S-2.

• Carbonates and bicarbonates react with acid to form CO2.

2HCl (aq) + Na2S(aq) H2S(g) + 2NaCl(aq)Hydrochloric acid + Sodium Sulfide

2H+ (aq) + S2-(aq) H2S(g)

HCl (aq) + NaHCO3(aq) NaCl(aq) + H2CO3(aq)

H2CO3(aq) H2O(l) + CO2(g)

HCl (aq) + NaHCO3(aq) NaCl(aq) + H2O(l) + CO2(g)

Hydrochloric acid + Sodium Hydrogen Carbonate

H+ (aq) + HCO3-(aq) H2O(l) + CO2(g)

Acid Spills and Antacids

• NaHCO3 and Na2CO3 are used to clean up acid spills in the lab.

• See page 127 for a list of antacids used in over the counter medications.

Reactions in which electrons are transferred between substances

Use of Oxidation numbers in determining redox reactions is

basically a bookkeeping method for keeping track of electrons

You must be able to identify an oxidation-reduction reaction. But first, we must learn the rules for

assigning oxidation #’s to different species.

Rules for oxidation numbers1) Atoms in elemental form are 0.2) Monatomic ion; charge of the ion is its oxidation

number.3) Nonmetals; usually negative numbers.

a.) oxygen = -2 unless a peroxide = -1b.) Hydrogen +1 with nonmetals, -1 with metalsc.) Halogens (-1) unless bonded to oxygen (+)

in a polyatomic ion (Ex: ClO3-; Cl = +5)

4) Sum of oxidation numbers must = 05) Most electronegative (furthest to right and up)

element gets a negative charge.

1) Atoms in elemental form are 0.

ExamplesAgPbCl2

O2

Oxidation # = 0 for 7 diatomic elements and for allother elements when by themselves.

2) Monatomic ion-- charge of the ion is its oxidation number.

ExamplesAgCl Ag = +1 Cl = -1PbI2 Pb = +2 I = -1Fe2O3 Fe = +3 O = -2

3) Nonmetals; usually negative numbers.a.) oxygen = -2 unless a peroxide = -1b.) Hydrogen +1 with nonmetals, -1 with metalsc.) Halogens (-1) unless bonded to oxygen (+)

in a polyatomic ion (Ex: ClO3-; Cl = +5)

ExamplesPbO oxygen = -2 Na2O2 oxygen = -1H2S hydrogen = +1 NaH hydrogen = -1KI iodine = -1 KIO2 iodine = + 3

Determine Oxidation # of element red element in each of the following:

MnO2

+4

KMnO4

+7

BrO2-

+3

BrO3-

+5

Br2

0

HClO4

+7

H2SO4

+6

PO33-

+3

CaH2

-1

SO42-

+6

Na2S

-2

Mg(NO3)2

+5

• If one reactant gains electrons another must lose electrons.

• Reduction is always accompanied by oxidation.

Again, oxidation reduction reactions occur when there is a transfer of electrons from

one species to another in a reaction.

Oxidation-Reduction Reactions

• An atom that becomes more positively charged is oxidized.–This is due to loss of e-.

• The gain of electrons by an atom is called reduction.

OIL -- RIGOxidation Involves Loss -- Reduction Involves Gain

Two mnemonics for remembering which substance is undergoing oxidation and which is undergoing reduction?

“Leo the lion says Ger”

Loss of electrons oxidation -- Gain of electrons reduction

Many metals react with O2 in the air to form metal oxides.

Metals lose electrons to oxygen.2 Fe + O2 2 FeO

As Fe is oxidized (loses e-), oxygen is reduced (gains e-).

Reduction is gain

2 Fe + O2 2 FeO

oxidation reduction

Oxidation of metals by acids and salts

• Reaction of a metal with either an acid or metal salt follows general form of:

A + BX AX +B

• Single displacement reaction

+2+6-2 0 +2+6-2 0

CuSO4(aq) + Zn(s)

What are the charges on each species?

What is oxidized and what is reduced?

ZnSO4(aq) + Cu(s)

What are the products?

OxidizedReduced

For the following reactants:

1) Write the reaction that occurs.

2) Identify what is being oxidized and reduced.

Magnesium + Hydrochloric Acid

Aluminum + Cobalt(II) Nitrate

Mg(s) + HCl (aq) MgCl2 (aq) + H2(g)

oxidation reduction

Al(s) + Co(NO3)2 (aq) Al(NO3)3(aq) + Co(s)

oxidation reduction

Types of Redox Reactions

• Combination (synthesis)• Decomposition• Displacement

– hydrogen, metal, halogen

• Disproportionation (When an element is

simultaneously oxidized and reduced).• Ex: H2O2 H2O + O2

Activity Series

• List of metals in order of decreasing ease of oxidation.

• Alkali and alkaline earth metals are at the top. (active metals)

• Gold, Silver, Platinum, and palladium are considered to be (noble metals) because they resist oxidation.

Using activity series to predict reactions

• Activity series can be used to predict reactions between metals and metal salts or acids.

• Any metal listed on the series can be oxidized by the ions of elements below it on the list.

Using activity series of metals, which metals from the list below can be oxidized by H+?

Ni

Al

Cu

Pb

Ag

Mg

Au

Chemical Reaction Types

• Decomposition• Synthesis• Single Replacement

• Precipitation• Neutralization

• Combustion

This should be a review before handing out All reaction types worksheet. Must edit all reaction type worksheet. It has to many weak acids reacting with strong bases, which is confusing for students.

Make sure most have reactions (see note in folder on handout).

Types of Redox Reactions

Types of Double Replacement Reactions

What are the 7 Diatomic Elements

H2 – hydrogen

N2 – nitrogen

O2 – oxygen

F2 - fluorine

Cl2 – chlorine

Br2 – bromine

I2 – Iodine

SynthesisA + B AB

Examples

H2 (g) + O2 (g) H2O (g)

Mg (s) + O2 (g) MgO (s)

Na (s) + Cl2 (g) NaCl (s)

DecompositionAB A + B

Examples

NaCl (s) Na (s) + Cl2 (g)

KClO3 (s) KCl (s) + O2 (g)

elec

Single Replacement

A + BC AC + B

Examples

Na (s) + HOH (l) NaOH (aq) + H2 (g)

sodium replaces hydrogen in water

Cl2 (g) + NaBr (aq) Br2 (l) + NaCl (aq)

chlorine replaces bromine in sodium bromide

Double Replacement (Metathesis)

AB + CD AD + BC

Examples AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq)

HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)NH4Cl (aq) + NaOH (aq) NH3 (g) + H2O (l) + NaCl (aq)

Blue color for the products represents the driving force which allows the chemical reaction to occur.

Combustionhydrocarbon + oxygen carbon dioxide + water

Examples

CH4 (g) + O2 (g) CO2 (g) + H2O (l)

C3H8 (g) + O2 (g) CO2 (g) + H2O (l)

CH3OH (g) + O2 (g) CO2 (g) + H2O (l)

CO2 + H2O

Neutralization ReactionsStrong Acid + Strong Base Salt + Water

Example

HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)

H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O(l)

H+ + OH- H2O

Concentration (Molarity)

• Concentration — the amount of solute per unit of solution.

• Molarity (M) — expresses concentration:

LitersinsolutionofVolume

soluteofMolesM )(

Calculate the molarity of a solution that contains 20.0g copper(I) chloride and has a total volume of 300.0 mL.

L

molM Must have the

correct units.

Given: 20.0 grams CuCl; 300.0 ml solution

Need: Moles CuCl; L of solution

20.0 g CuCl

g CuCl

mol CuCl

99.0

1=

300.0 mL

mL

1

1000

L

0.202 mol CuCl

= 0.3000 L

0.673 M CuCl

L

CuClmolM

3000.0

202.0

L

molM

How many grams of NaCl are needed to make 2.5 L of 0.20 molar solution?

L

molM

Given: 0.20 M solution; 2.5 L solution

Need: grams of NaCl

Must calculate # of moles, and then convert it into grams.

(L)(L)

mol = ML

Mol = ML

Mol = 0.20 mol NaCl x 2.5 L = 0.50 mol NaCl

L

0.50 mol NaCl

mol NaCl1= 28 g NaCl

55.84 g NaCl

Dilution of Stock Solutions

• Chemicals are purchased in concentrated form. They need to be diluted for most lab use.

• Formula for dilution:

Mi Vi = Mf Vf

Mi Vi = Mf Vf

How much stock (12 M) HCl (aq) is

required to make 200.0 mL of 3M HCl (aq)?

Vf = 200.0 mL

Mi = 12 M

Mf = 3 M

Vi = ? mL

Must rearrange equation above to solve for initial volume.

50. mL

i

ffi M

VMV

)12(

)0.200)(0.3(

M

mlMVi

Measure 150 mL of water in a beaker. Slowly add 50.0 mL of 12 M HCl for a final volume of 200.0 mL.

Two things to note:1) Always add concentrated acid to water, and not

the reverse to avoid unwanted splashing due to the heat generated.

2) When diluting a solution, the amount of solute doesn’t change, only the final volume.

If diluting a solution other than acids, start with initial volume of concentrated solution, and then dilute with distilled water until you have the desired volume.

You try one!

Mi Vi = Mf Vf

We want to prepare 500. mL of 1.00 M acetic acid from a 17.5 M stock solution of acetic acid. What volume of the stock solution is required?

Vf = 500. mL

Mi = 17.5 M

Mf = 1.00 M

Vi = ? mL

28.6 mL

i

ffi M

VMV

)5.17(

)0.500)(00.1(

M

mlMVi

Pour 471.4 mL of distilled water into a beaker. Slowly pour the 28.6 ml of acid into the water and swirl. Fill

the container with distilled water to 500. mL.

In a solution of 0.25 M MgCl2 you have:

There may be times when you must consider the concentration of ions in a solution.

(You must consider the subscripts for this)

M of Mg2+ = 0.25 M

M of Cl- = 2 x 0.25 M = 0.50 M

MgCl2 Mg2+ + 2Cl-

What is the concentration of each ion in the following?

0.15 Na3PM of Na+ = 0.45 MM of P3- = 0.15 M

Titrations• Determining the concentration of an unknown

solution.• Use a 2nd solution of known concentration (standard

solution) that undergoes a reaction with the unknown solution.

• Use the ratios in the balanced equation along with the M = mol/L equation to determine molarity of unknown.

• The point at which the two solutions are stoichiometrically equal is known as the equivalence point. –The reaction is complete and no

excess reactant is present.

–How do we know when this occurs during the reaction?

• In acid base reactions dyes known as indicators are used.– Phenolphthalein is colorless in acid

solution, and pink in basic solution.

– End point is reached when a drop of the base remains pink. There is no acid for this drop to react with and the solution is now basic.