REGENTS REVIEWUnit 9 – Gas Laws

Unit 10 – Solutions

Unit 11 – Acids & Bases

Unit 12 – Equilibrium

GAS LAWSChapter 14

Ideal Gas• Don’t exist

• Model to explain behavior of all gases

• Review• 1 mole of ANY gas occupies 22.4L of volume at STP

Kinetic Molecular Theory• The particles in a gas are constantly moving in rapid,

random, straight-line motion.• Gas particles have no volume compared to the volume of

the gas.• No attraction between particles• All collisions are completely elastic

Boyle’s Law• Relationship between pressure and volume

• Constant Temperature and amount of gas

• Mathematical relationships• As pressure is increasing, volume is decreasing• As pressure is decreasing, volume is increasing

PV PV1 1 2 2

Boyle’s Law

V

P

Example

• The pressure of a 25 L sample is changed from 2 atm to 0.4 atm. What is the new volume of the gas?

• P1V1 = P2V2

• (2atm)(25L) = (0.4atm)V2

• V2 = 125 L

Charles’s Law• Relationship between volume and temperature

• Constant Pressure and amount of gas

• Mathematical relationships• As temperature is increasing, volume is increasing• As temperature is decreasing, volume is decreasing

• Temperature must be in Kelvin

V

T

V

T1

1

2

2

Charles’s Law

T

V

Example

• A gas occupying 45L at 27°C is cooled until its volume is 15L. What is the new temperature of the gas?

2

15

300

45

T

L

K

LV

T

V

T1

1

2

2

KT 1002

Gay-Lussacs's Law• Mathematical relationships

• As temperature increases, pressure increases• As temperature decreases, pressure decreases

• Also known as Gay-Lussac’s Law

• Temperature must be in Kelvin P

T

P

T1

1

2

2

Gay-Lussacs's Law

T

P

Combined Gas Law

• Combines Boyle’s, Charles’s, and Gay-Lussacs's Laws together

PV

T

PV

T1 1

1

2 2

2

Example

• A gas at 5atm is heated and compressed from 10L at 100K to 5L at 200K. What is the new pressure?

PV

T

PV

T1 1

1

2 2

2

K

LP

K

Latm

200

)5)((

100

)10)(5( 2

atmP 202

Example

• A gas at 2atm and 27°C occupies 10L of space. What is the new volume when it is cooled to STP?

PV

T

PV

T1 1

1

2 2

2

K

Vatm

K

Latm

273

))(1(

300

)10)(2( 2

LV 2.182

Avogadro’s Law• Equal volumes of gas at the same Temperature and

Pressure have the same number of particles

• At the same temperature and pressure, which sample contains the same number of moles of particles as 1 liter of O2(g)?

(1) 1 L Ne(g) (3) 0.5 L SO2(g)

(2) 2 L N2(g) (4) 4 L H2O(g)

Ideal Gases

• Ideal Gases follow assumptions of the Kinetic Molecular Theory

• When do real gases act most like an ideal gas?• High Temperature• Low Pressure

• When do real gases act least like an ideal gas?• Low Temperature• High Pressure

SOLUTIONS AND SOLUBILITY

Chapters 15, 16

Reference Tables F, G

Solution• Homogeneous Mixture

• Uniform Throughout

• Solvent• Dissolving medium in mixture

• Solute• Dissolved particles in solution

• Aqueous Solution• Solution with water as the solvent• NaCl(aq)

Solubility• measure of the amount of solute (how much) that can

dissolve in a given quantity of solvent at certain conditions • Affected by Temperature, Pressure, and Chemical Nature

• Soluble• Solute will dissolve in solvent

• Insoluble• Solute will not dissolve in solvent

Temperature• increasing temperature increases solubility of solids and

liquids in other solids and liquids

• decreasing temperature increases solubility of gases in liquids • Ex: Cold soda is fizzy

Pressure• Increasing partial pressure of gas above liquid increases

solubility of the gas in the liquid

• Example• Soda Bottles

Chemical Nature• “Like dissolves Like”

• Polar and ionic substances will dissolve in polar solvents

• Nonpolar substances will dissolve in nonpolar solvents

Dissolving Speed• Stirring

• stirring increases speed of dissolving

• Temperature• Increasing temperature increases speed of dissolving solids and

liquids

• Particle Size• decrease particle size or increase surface area, increase speed of

dissolving

Table G

• Shows the relationship between temperature and amount of solute for a number of different compounds

Types of Solution

• Saturated• maximum amount of solute for a given quantity of

solvent• At Equilibrium

• Unsaturated• contains less than the maximum amount of solute

• Supersaturated• Contains more solute than it can theoretically hold

Solution Types

• Saturated• On the line

• Unsaturated• Under the line

• Supersaturated• Above the line

Electrolytes• Compounds that conduct an electrical current when

dissolved or in a molten state (melted)

• Ionic compounds• NaCl, KNO3, HCl

Dissolving vs. Dissociation• Dissolving

• Molecules separate as solvent molecules mix• C6H12O6(s) + H2O(l) C6H12O6(aq)

• Dissociation• Ions separate as solvent molecules mix• NaCl(s) + H2O(l) Na+(aq) + Cl-(aq)

Separating Mixtures

• Mixtures can be separated based on their physical properties• Boiling Pt, Freezing Pt, Density, Molecular Polarity,

Particle Size

• Process Examples• Filtering, Distillation, Evaporation, Crystallization,

Chromatography, Desalination, Extraction

Molarity• Molarity = Moles of Solute

Liters of Solution• 1 mol/L = 1 M

• Often used for solids dissolved into liquids• Most common concentration system

Molarity Example

• What is the molarity of 2 moles of glucose dissolved in 5 Liters of solution?

• Molarity = Moles of Solute Liters of Solution

2

5

mol

Liter 0 4. M

Example• How many moles of HCl are dissolved in 4L of a 3M

solution of HCl?• Molarity = Moles of Solute

Liters of Solution

L

XM

43 molX 12

Percent Composition• Percent Comp = Part x 100%

Whole

Example

• A solution contains 80g of NH4Cl in a 1000g solution, what is the percent by mass composition of this solution?

ClNHg

g4%8%100*

1000

80

Example

• A solution contains 60 mL of NH3 in a 1 Liter solution, what is the percent by volume composition of this solution?

3%6%100*1000

60NH

mL

mL

Parts Per Million (ppm)• Used for very small concentrations

• ppm = grams of solute x 1,000,000

grams of solution

• Units = ppm

ppm example

• A 2 kg bar of silver contains 0.05 g of gold, what is the parts per million concentration of gold in the silver bar?

0 05

20001 000 000

.* , ,

g

g 0 000025 1 000 000. * , ,

25 ppm Au

Colligative Properties of Solutions• Properties of a solution that depend only on the number of

particles dissolved

• Vapor Pressure• Boiling Point• Melting Point

• Added solute particles get in the way of the solvent molecules changing the above properties

Colligative Properties of Solutions

• Adding more solute particles to pure solvent decreases vapor pressure

• Adding more solute particles to pure solvent increases boiling point

• Adding more solute particles to pure solvent decreases freezing point

Number of Particles• When a covalent compound dissolves the compound

stays intact.

• When an ionic compound dissociates the compound splits into its ions.

• More particles, larger change in property

ACIDS AND BASESChapter 19

Reference Tables K, L, M

Electrolyte• A substance that conducts an electrical current when

melted or in solution • Ionic compounds• Acids and bases

Acid-Base Theories• Arrhenius

• Acid• Compounds that ionize to produce hydrogen ions (H+) in aqueous

solutions• Examples: HCl, HBr, H2SO4, CH3COOH

• Base• Compounds that ionize to produce hydroxide ions (OH-) in aqueous

solutions• Examples: KOH, NaOH, LiOH

Acid - Base Theories• Bronsted-Lowry

• Acid• Hydrogen ion donor

• Base• Hydrogen ion acceptor

• Lewis• Acid

• Accepts a pair of electrons

• Base• Donates a pair of electrons

Acid-Base Theory

Arrhenius

Bronsted-Lowry

Lewis

Properties • Acids

• Taste Sour• Will change color of acid – base indicator• Can be strong or weak electrolytes in an aqueous solution

• Bases• Taste Bitter• Feel Slippery• Will change color of acid – base indicator• Can be strong or weak electrolytes in an aqueous solution

Ionization• Electrolytes will dissociate into ions when dissolved in

water

• Strong Electrolytes will completely dissociate• Weak Electrolytes will only partially dissociate

Ionization of Water• Water can be split into 2 ions

• H+ and OH-

• Ionization of Water• H2O H+ + OH-

• H2O + H2O H3O+ + OH-

OHH 3

Strong Acids• Completely dissociate when in solution

• HCl(s) H+(aq) + Cl-(aq)

• HNO3(s) H+(aq) + NO3

-(aq)

Polyprotic Acids• Acids that have more than one H

• Examples:H2SO4, H3PO4

• Can release more than one H+ into solution• H2SO4(s) 2H+

(aq) + SO42-

(aq)

Strong Bases• Completely dissociate when in solution

• NaOH(s) Na+(aq) + OH-

(aq)

• KOH(s) K+(aq) + OH-

(aq)

Neutral Solutions• For neutral solutions

• [H+] = [OH-]

• For all aqueous solutions• [H+] * [OH-] = 1.0 x 10-14

Measuring Acidity (Alkalinity)• Traditionally we measure [H+]

• pH = -log [H+]• Neutral solution [H+] = 1.0 x 10-7

• pH = 7

• Acidic Solutions pH < 7.0• [H+] > 1.0 x 10-7

• Basic Solutions pH > 7.0• [H+] < 1.0 x 10-7

Changes in pH• pH increases by 1 for every decrease in [H+] by a

magnitude of 10

Measuring pH• Litmus paper

• Red in acid• Blue in base

• pH paper• pH Meter• Acid – Base Indicators (Table M)

Neutralization

• Acid + Base Water + Salt• Double Replacement Reaction

• HA + BOH HOH + BA

• Examples:• HCl + NaOH H2O + NaCl

• HNO3 + LiOH H2O + LiNO3

• H2SO4 + 2KOH 2 H2O + K2SO4

Titration

• Process in which a volume of solution known concentration is used to determine the concentration of another solution

• Usually shown by a color change of an indicator (end point)

Titration Example• 2 Liters of an unknown conc. of NaOH is titrated with 1 Liter

of 6M HCl, what is the concentration of the base?

MAVA = MBVB

(6M) * (1L) = (X) * (2L)

X =3M NaOH

Polyprotic Acids

• Acids that have more than one H• Examples:H2SO4, H3PO4

• [H2SO4] = 1M

• [H+] = 2M

MAVA (#of H’s)= MBVB (# of OH’s)

Another Titration Example• 500 milliliters of an unknown conc. of Ca(OH)2 is titrated with

1 Liter of 1M H3PO4, what is the concentration of the base?

MAVA = MBVB

(1M) (1L) (3) = (X) (0.5L) (2)

X = 3M Ca(OH)2

KINETICS AND EQUILIBRIUMChapter 18

Reference Table I

Collision Theory• a reaction is more likely to occur if reactant particles

collide with proper energy and orientation

• Reaction Rate• How fast the reaction proceeds

• Activation Energy (EA)• Minimum energy that colliding particles must have in order to react

Factors Affecting Reaction Rates• Temperature

• Increasing temperature increases the reaction rate

• Concentration• increasing the number of particles in a given volume

(concentration) increases the reaction rate

• Surface Area• increasing surface area increases reaction rate

• Catalyst• the presence of a catalyst will often increase reaction rate• Catalysts are not used up during a reaction

Potential Energy• Potential Energy

• Energy stored in chemical bonds

• Heat of Reaction (ΔH)• Energy absorbed or released during a chemical reaction• PEProducts – PEReactants

Potential Energy Diagram• Graphically shows energy released or absorbed during a

reaction

Reaction Process

Ene

rgy

Ene

rgy

Reaction Process

Exothermic Reaction

Reaction Process

Ene

rgy

EA

ΔH

PEProducts

PEReactants

Endothermic Reaction

Ene

rgy

Reaction Process

EA

ΔHPEProducts

PEReactants

Table I• Shows reactions with ΔH

• ΔH =PEProducts – PEReactants

• Endothermic• ΔH = (+)

• Exothermic• ΔH = (-)

Affect of a Catalyst• Provides an alternate pathway for the reaction to proceed

• Decreases activation energy

• Increases reaction rateE

nerg

y

Reaction Process

Without Catalyst

With Catalyst

Entropy• Measure of randomness or disorder

• Systems in nature tend to undergo changes towards lower energy and higher entropy• The universe is lazy and disorganized

• Increasing Entropy• Solid Liquid Gas• Solid Dissolved

Equilibrium• Reversible Reaction

• Reaction in which conversion of reactants to products and conversion of products to reactants occurs simultaneously

• Chemical Equilibrium • Rate of forward reaction is equal to rate of reverse reaction

• At Chemical Equilibrium, there is no net change in the actual amounts of the reactants and products.• Amounts remain constant

Equilibrium• Saturated Solution

• Solid in equilibrium with dissolved particles

)()()(2

aqaqOH

s ClKKCl

LeChatelier’s Principle• If a stress is applied to a system in dynamic equilibrium,

the system changes in a way that relieves the stress

• Stresses• Change in concentration of reactant or product

• Change in temperature

• Change in pressure• Only applies to reactions in the gas phase with unequal number of

moles of gas.

Stresses against Dynamic Equilibrium

• Change in concentration of reactant or product• Adding a substance will shift equilibrium away from that side• Removing a substance will shift equilibrium towards that side

• Change in temperature• Increasing temperature will shift equilibrium away from heat term• Decreasing temp. will shift equilibrium towards heat term

• Change in pressure• Only applies to reactions in the gas phase with unequal number of

moles of gas.• Increasing pressure shifts equilibrium towards the side with least

number of moles of gas

How does equilibrium shift?

• After a stress…..• When equilibrium shifts to the right (Product side)

• Forward reaction speeds up• Concentration of Reactants decrease• Concentration of Products increase

• When equilibrium shifts to the left (Reactant side)• Reverse reaction speeds up• Concentration of Reactants increase• Concentration of Products decrease

….until equilibrium is reestablished