Solution Concentrations / Acids and Bases /

Oxidation-Reduction

SolutionsA solution is a homogeneous mixture of one substance

dissolved in another.

The substance being dissolved is usually the solute, and the substance in which it is dissolved is called the solvent . Which is which can be confusing, especially if both are in the same phase. The substance in greater quantity is normally called the solvent, and in solid-in-liquid solutions, the liquid is always the solvent.

Types of Solutions

Liquid Liquid

Solid

Gas

Antifreeze

Kool-aid

Soft drinks

Gas Liquid

Solid

Gas

Humidity

Mothballs

Air(O2 in N2)

Solid Liquid

Solid

Gas

Amalgams(Hg in Na)

Steel (C in Fe)Palladium-Hydrogen electrode

Solvent Solute Example

Nature of Solute and SolventWhen making solutions, the general rule is that “Like

dissolves like”. Since there are two main types of compounds, polar(with uneven distribution of charges) and nonpolar(with evenly distributed charges), the decision of which solvent to use comes down to making it agree with the solute.

The issue of solubility(amount of substance that will dissolve in a solvent) is made more complex in that even some which do agree as far as polarity are still insoluble(will not dissolve) due to other factors

When speaking of two liquids, if they mix completely, they are called miscible, if not, like oil and water, they are immiscible.

Solution Process

How does the solution process work?

In water, water molecules pull away the outer layer of molecules from the substance and surround them in solution. This exposes another layer to attack and by layers, the substance dissolves in water without outside help.

04_41

+

+

++

+

+

+

+

++

+

OH

H

O HH

+

+

Cation

Anion

Rate of DissolvingHow fast any solute dissolves depends on the nature of both

solute and solvent

Solids: The rate of dissolving in a liquid can be increased by:

1. Stirring or shaking the solution

2. Crushing the solid

3. Heating the solution-temperature often makes solids

more soluble

Gases: The rate of dissolving in liquids can be increased by:

1. Cooling the solution

2. Increasing the gas pressure above the solution

Energy in Solutions

Many solutes involve energy in their dissolving. If a substance causes the solution to get cold as it dissolves, it is taking energy from the solvent, and it is called endothermic. If the solution gets warmer as the solute dissolves, it is releasing energy to the solvent and is called exothermic.

Solution Concentrations Descriptive

1. Dilute—few solute particles

2. Concentrate—many solute in given amount of solvent

3. Saturated—maximum solute in given amount of solvent

4. Unsaturated—less than maximum solute/solvent ratio

5. Supersaturated-unstable special condition with solute/solvent over normal maximum

Henry’s Law

P = kCP = partial pressure of gaseous solute above

the solutionC = concentration of dissolved gask = a constant

The amount of a gas dissolved in a solution is The amount of a gas dissolved in a solution is directly proportional to the pressure of the gas directly proportional to the pressure of the gas above the solution.above the solution.

Solution Concentrations

Quantitative

Molarity—moles solute per liter of solution

a. Find moles solute

b. Find liters solution

c. Solve for molarity

M = mol/L

Molarity CalculationsM= mol g | 1 mol = mol mols | MM = grams

L | MM | 1 mol

L = mol mol= M x L mass=M x L x MM

M

Dilution Equation

M1V1 = M2V2

ExamplesHow many grams of NaCl are needed to make

1.00 L of a 0.100 M solution?

How many mL of 2 M HCl are needed to get 1 mol of HCl?

% Concentrations

Weight / weight—(grams solute / grams solution) x 100%

Volume/volume--- (volume of solute/ volume of solution) x 100%

Examples

How many grams of NaHCO3 need to be added to water to make a 0.4% solution?

How many mL of ethanol need to be added to 20 mL of solution to make it 9.0 %?

Performing Dilutions

Reducing concentration by adding solventDone to conserve space—concentrated stock is

stored, dilutions made as needed If a researcher is given a stock solution, it is easy to make a new, lesser concentrated solution by using the dilution equation given before:

M1V1=M2V2 Since molarity times liters equals moles, the two

products above will have equal numbers of moles

ExamplesHow many mL of 6.0 M HCl is needed to make

100 mL of 1.5 M HCl?

What is the final concentration of a 1:5 dilution of 100 mL of 15% NaCl?

Homework 15a

p. 463 8, 10

p. 464 11, 13

p. 465 14, 16

p. 466 17, 19

p. 468 22, 23

p. 484ff 70, 71, 77, 79, 81, 82

Other Solution Units

1. Mole fraction (A) =

2. Molality (m) =

molestotal moles in solution

A

moles of solutekilograms of solvent

ExamplesWhat is the molality of a solution of 4.5 g of

NaCl in 100.0 g of water?

What is the mole fraction of the same solution?

Colligative Properties

Depend only upon how many particles are present, not on what they are

Osmotic Pressure

Vapor Pressure Lowering

Freezing Point Depression

Boiling Point Elevation

Vapor Pressure Related

1. Vapor pressure is lowered because solute particles compete for space at the surface thus making vapor less above the solution. 2. Freezing point depression

a. Lower vapor pressure causes a lower temperature to freeze the solvent

b. Solute particles also interfere with crystal formation

c. Electrolyte (ionic) solutes have a greater effect due to disassociation into ions making more particles—

Sugar- 1 particle NaCl—2 CaCl2 --3 3. Boiling point elevation—same cause as freezing point

depression—also solute particles interfere with vaporization

Raoult’s Law

Psoln = solventPsolvent

Psoln = vapor pressure of the solution

solvent = mole fraction of the solvent

Psolvent = vapor pressure of the pure solvent

The presence of a nonvolatile solute lowers The presence of a nonvolatile solute lowers the vapor pressure of a solvent.the vapor pressure of a solvent.

11_276

WaterVapor

Water Aqueoussolution

Aqueoussolution

(a) (b)11_279

Vapor pressure of pure B

Vapor pressureof pure A

Vapo

r pre

ssure

B

Vapor pressureof solution Vapor pressure

of solution

(a) (b) (c)

AB

AB

A

Boiling Point Elevation

A nonvolatile solute elevates the boiling point of the solvent.

Tb = Kbmsolutei

Kb = molal boiling point elevation constantm = molality of the solutei = van’t Hoff factor, the number of particles a

substance becomes in water solution

Freezing Point Depression

A nonvolatile solute depresses the freezing point of the solvent.

Tf = Kfmsolutei

Kf = molal freezing point depression constantm = molality of the solutei = van’t Hoff factor, the number of particles a

substance becomes in water solution

11_280

atmP

ress

ure

(atm

)

Tf Tb

Freezingpoint ofsolution

Freezing pointof water

Boiling pointof water

Boiling pointof solution

Temperature (C)

Vapor pressureof pure water

Vapor pressureof solution

Examples

What are the freezing and boiling points of a 0.029 m solution of NaCl?

Osmotic PressurePressure to keep solvent from flowing across a semi-

permeable membrane into a solution

1. Water naturally flows into solutions

2. Can be reversed to make pure water from solutions such as seawater or (Saudi)

3. Natural example—trees lose water from leaves, become more concentrated, water flows up from roots by osmotic pressure

Osmotic PressureBlood solutions have to account for this

a. Isotonic—same osmotic pressure (as blood or tears)

b. Hypotonic—less pressure—causes water to go into cells, swelling them to rupture-- hemolysis

c. Hypertonic—more pressure, water goes out of cells, leaving them dehydrated and shriveled-- crenation

d. Intravenous solutions must always be isotonic with blood

e. Edema (swelling of body parts) happens if blood becomes less concentrated and water remains in tissues. Caused by malnutrition or kidney failure

f. Dialysis for kidney failure takes advantage of osmotic pressure to clean the blood of patients

Homework 15b

p 469-70 24, 25, 26, 27

p. 484-5 50, 55, 57, 58, 86, 88

Common Definitions

Acids conduct electricity in solution, react with active metals, sour taste, change blue litmus to red

Bases conduct electricity in solution, feel slippery, have bitter taste, turn red litmus blue

Some substances act as multiple acids since they have more than one acid H+

--polyprotic— H2SO4 H3PO4

Acids and BasesDefinition Theories Presently Held

1. Bronsted-Lowry

a. Acid is a proton(H+) donor

b. Base is a proton acceptor

c. Explains all bases, but not all substances which act as acids

2. Lewis Theory

a. Acid is electron-pair acceptor—attracts elements with exposed electron pairs

b. Base is electron pair donor (has exposed electron-pair)

c. Explains how substances like Fe+3 can become acids

Acid Base Strength

1. Strong acids completely ionize in solution

HCl HBr HI HNO3 H2SO4 HClO4

2. Weak acids only partially ionize

HC2H3O2 H2CO3 H3PO4

3. Strong bases have large attraction for protons

OH- PO4-3 NH2

-

4. Weak bases have small attraction for protons

NH3 C2H3O2- NO2

-

Naming AcidsIf acid has no oxygen in its formula, it is binary

a. Start the name with hydro-b. Add –ic ending to the name of other element or ionc. Add the word acid

HCl – hydro chlor –ic acid

H2S -- hydrosufuric acidIf acid has oxygen in its formula, it is an oxyacid

a. Name the acid from the name of the polyatomic ion in it—DO NOT NAME THE HYDROGEN!!!

b. Add –ic if the ion ends in –ate c. Add –ous if the ion ends in –ite

H2SO4 – sufur –ic acid (SO4-2 is sulfate)

HNO2 – nitrous acid (NO2- is nitrite)

Neutralization of Acids With Bases

1. Acid plus base makes salt plus water

HCl + NaOH NaCl + H2O typical example

 

2. The reaction cancels both acid and base if concentrations are equal

Ionization of Water

H2O OH- + H+

1. Water is both weak acid and weak base 2. Equilibrium constant K for this reaction becomes Kw = [H+][OH-] since H2O is pure liquid

3. Kw is constant at room temperature in pure water and equals 1 x 10-14 since

[H+]=[OH]=1x10-7

Measuring Acid Base ConcentrationspH Scale

1. Hydrogen ion concentrations can be measured by electric current flow or by the affect on certain indicators.

2. pH scale was developed to easily measure [H+] by taking its negative log.

pH= -log [H+] [H+]=10-pH

3. Acids have pH <7 Bases have pH>7 Neutral pH=7 where[H+] = 1x10-7

ExamplesWhat are the pH values for the following

solutions?

a. [H+] = 1 x 10-2 M

b. [H+] = 3.0 x 10-6 M

c. [OH-] = 8.2 x 10-6 M

pH and pOH

Since the p stands for a function, as in math, it can be applied to any value and mean the same thing. If applied to [OH-], the new value is called pOH. A special relationship occurs in any water solution:

pH + pOH = 14.00

Examples

Household ammonia has a hydroxide concentration of 4.0 x 10-3 M. What are its pH and pOH values?

pH and pOH to Concentration

The inverse relation can also be calculated:

If blood is pH 7.40, what are the [H+] and [OH-] values?

Homework 19a

p. 596 #1 p. 601 #5 p. 607 #12

p. 609 #18 p. 612 #20 p. 614 #21

p 631ff 88, 89, 90

Acid-Base ReactionsAcids and bases are opposite- one donates protons, the

other takes them. So when they get together, a canceling or neutralization happens, where an acid plus a base make water and a salt (formed in acid-base reactions from positive base ion and negative acid ion)

This reaction can take place even when the acid or base is not complete. Such incomplete acids or bases are called anhydrides, an acid or base without water. Metal oxides are base anhydrides and nonmetal oxides are acid anhydrides

Titrations

Measure acid or base by combining a known quantity of one with a continually measured volume of the other. The pH is monitored until an indicator changes or pH meter measures neutral condition.

MaVa = MbVb

Titration (pH) Curve

A plot of pH of the solution being analyzed as a function of the amount of titrant added.

Equivalence (stoichiometric) point: Enough titrant has been added to react exactly with the solution being analyzed.

15_327

01.0

Vol NaOH added (mL)

50.0

7.0

13.0pH

100.0

Equivalencepoint

Acid-Base Indicator

. . . marks the end point of a titration by changing color.

The equivalence point is not necessarily the same as the end point.

15_333

– O

C

O

C O–

O

(Pink base form, In– )

HO

COH

C O–

O

(Colorless acid form, HIn)

OH

15_3340 1 2 3 4 5 6 7 8 9 10 11 12 13

The pH ranges shown are approximate. Specific transition ranges depend on the indicator solvent chosen.

pH

Crystal Violet

Cresol Red

Thymol Blue

Erythrosin B

2,4-Dinitrophenol

Bromphenol Blue

Methyl Orange

Bromcresol Green

Methyl Red

Eriochrome* Black T

Bromcresol Purple

Alizarin

Bromthymol Blue

Phenol Red

m - Nitrophenol

o-Cresolphthalein

Phenolphthalein

Thymolphthalein

Alizarin Yellow R

* Trademark CIBA GEIGY CORP.

15_335AB

pH

00

Vol 0.10 M NaOH added (mL)

2

4

6

8

10

12

14

20 40 60 80 100 120

Equivalencepoint pH

00

Vol 0.10 M NaOH added (mL)

2

4

6

8

10

12

14

20 40 60 80 100 120

Equivalencepoint

Phenolphthalein

Methyl red

Phenolphthalein

Methyl red

Calculating Acid-Base in Titration

Acid or base molarity can be found with the equation:

MaVai = MbVbi

Where i is the number of H+ or OH- ions in the formula of the acid or base

Example

What is the molarity of 30.0 mL CsOH solution neutralized by 26.4 mL of 0.250 M H2SO4?

pH and Body Systems

pH has serious effects on biological materials if it is out of correct range

a. Enzymes work only in specific pH ranges

b. pH out of normal range can indicate disease of injury

Buffers

pH changes in biological situations can be very dangerous

Buffers protect against changes in pH

Best systems have a weak acid or base and a salt of that acid or base

HC2H3O2 and NaC2H3O2

NH3 and NH4Cl

A Buffered Solution. . . resists change in its pH when either H+ or OH are added.

1.0 L of 0.50 M H3CCOOH

+ 0.50 M H3CCOONa

pH = 4.74

Adding 0.010 mol solid NaOH raises the pH of the solution to 4.76, a very minor change.

Buffered Solution Characteristics

Buffers contain relatively large amounts of weak acid and corresponding base.

Added H+ reacts to completion with the weak base.

Added OH reacts to completion with the weak acid.

The pH is determined by the ratio of the concentrations of the weak acid and weak base.

Buffers

Blood buffer system is based on H2CO3 , HCO3- and CO2

1. Too high CO2 causes acidosis (p H in blood too low)—caused by

a. Hypoventilation b. Emphysema c. Heart failure

2. Too high [H+] caused by

a. Diabetes b. Kidney failure c. Acid drugs:aspirin or acid foods

3. Too low CO2 causes alkalosis (blood pH too high)—caused by

a. Hyperventilation b. High fevers c. Hysteria

4. Too low [H+] caused by

a. Severe vomiting losing acid from stomach

b. Kidney disease

Homework 19b

p. 621 31, 32p. 630 ff 42, 46, 52, 54, 55, 63,

70, 73, 76, 95, 96

Oxidation Numbers

Every atom, ion or polyatomic ion has a formal oxidation number associated with it. This value compares the number of protons in an atom (positive charge) and the number of electrons assigned to that atom (negative charge).

In many cases, the oxidation number reflects the actual charge on the atom, but there are many cases where it does not. Think of oxidation numbers as a bookkeeping exercise simply to keep track of where electrons go.

Reduction

Reduction means what it says: the oxidation number is reduced in reduction.

This is accomplished by adding electrons. The electrons, being negative, reduce the overall oxidation number of the atom receiving the electrons.

Oxidation

Oxidation is the reverse process: the oxidation number of an atom is increased during oxidation.

This is done by removing electrons. The electrons, being negative, make the atom that lost them more positive.

Oxidation versus Reduction: Other Definitions

Oxidation is:

-loss of electrons

-increase of oxidation state

-gain of oxygen/loss of hydrogen

Reduction is:

-gain of electrons

-decrease of oxidation

state

-loss of oxygen/gain of hydrogen

Agents

Oxidizing agent:

--gains electrons because it is stealing them from the other substance

--is the substance which is being reduced in the reaction

Reducing agent:

--loses electrons by giving them away to the other substance

--is the substance which is being oxidized in the reaction

Rules for Assigning Oxidation States

1. Oxidation state of an atom in an element = 0

2. Oxidation state of monatomic element = charge

3. Oxygen = 2 in covalent compounds (except in peroxides where it = 1)

4. H = +1 in covalent compounds

5. 5. Fluorine = 1 in compounds

6. Sum of oxidation states = 0 in compounds

Sum of oxidation states = charge of the ion

Homework 20

p. 640 1, 3

p. 642-3 4, 5, 8, 9

p. 658 ff 44, 45, 47, 49, 51, 52