General Chemistry

2nd Semester review

• 1 mol =

• 1 mol =

• 1 mol =

• How many molecules are in 7.25 mol of carbon dioxide?

• How many moles are in 1.62 x 1023 atoms of argon?

• How many moles are in 83.2 g of H2SO4?

• What is the mass in grams of 1.5 mol of NaCl?

• What is the volume of 1.50 mol of nitrogen at STP?

• How many moles are in 75.3 L of water vapor at STP?

• How many atoms are in a 44.3 g piece of iron?

• What is the volume of 5.24 x 1022 molecules of iodine?

• What is the volume of 68.4 g of F2 gas at STP?

• How many molecules are in 0.75 L of nitrogen gas at STP?

Calculating percent composition

1. Find the total molar mass of each element in the compound.

2. Find the molar mass of the entire compound.3. Divide the total molar mass of each element

by the molar mass of the compound then multiply by 100

4. Check that all your percentages add up to 100

• What is the percent composition of H2SO4?

Calculating empirical formula

1. Change % to g (assume 100 g of compound so 30% = 30 g)

2. Convert each element from g to moles3. Divide each mole amount by the smallest

number from step 24. Change to a whole number = subscript in

empirical formula

Example: 43.7 % P and 56.3 % O

Assumptions of the Kinetic Molecular Theory

• Gasses consist of small particles that take up little volume relative to the volume of empty space around them– Gas molecules are very far apart and therefore

don’t experience attractive or repulsive forces.

• Gas particles move in constant, random straight lines until they collide with other particles or with the walls of the container– Collisions are elastic -

• The energy of gas particles is determined by the particle’s mass and velocity– KE =

Intermolecular forces

• Inter- means between or among

• Intermolecular forces can hold together identical particles or two different types of particles

• Weaker than intramolecular forces (bonds)

Dispersion Forces

• Weak forces that result from temporary shifts in the density of electrons in electron clouds

• Exist between all particles– Weak for small particles– Get stronger as the number of electrons involved

increases– F2

– Cl2

– Br2

– I2

Dipole-dipole forces

• Attraction between oppositely charged regions of polar molecules– Polar molecule =

• Neighboring polar molecules orient themselves so that oppositely charged regions align

Hydrogen Bonds

• Dipole-dipole attraction that occurs between molecules containing a hydrogen atom bonded to a flourine, oxygen, or nitrogen atom

• Viscosity - measure of the resistance of a liquid to flow– Attractive forces – stronger intermolecular forces

= higher viscosity– Particle size – larger molecules = higher viscosity– Temperature – lower temperature = higher

viscosity

9.4 Phase changes

• Matter changes phases when energy is added or removed

Phase changes that require energy

• Melting– Heat flows from an object at a higher temperature

to an object at a lower temperature– Ice absorbs heat which does not raise temperature

but is used to break hydrogen bonds– When hydrogen bonds are broken molecules can

move further apart into the liquid phase

• Vaporization – process by which liquid changes to vapor– Vapor – gaseous state of a substance that is

normally liquid at room temperature– Evaporation – when vaporization occurs only at

the surface of a liquid– Vapor pressure – the pressure exerted by a vapor

over a liquid

• Sublimation – changing from solid to gas without becoming a liquid– Dry ice– Moth balls– Solid air fresheners

Phase changes that release energy

• Freezing – Heat flows out of warmer object into cooler object– Molecules slow down & become less likely to flow

past one another– Intermolecular forces cause the molecules to

become fixed into set positions– Freezing point – temperature in which a liquid

becomes a solid

• Condensation – process by which a gas or vapor becomes a liquid

• Deposition – substance changes from gas or vapor to solid without first becoming a liquid– frost

Phase Diagrams

• Temperature and pressure both effect the phase of a substance– Have opposite effects

• Phase diagram – graph of pressure vs temperature that shows which phase a substance will be in under different conditions.

• Triple point = point at which all three phases exist at the same time

• Gas Laws:

• Remember for all gas laws T must be in K– K = C + 273

• Back of periodic table

• A balloon contains 30.0 L of air at 100.0 kPa. What did the volume change to that caused the pressure to decrease to 25 kPa?

–P1 =

–V1 =

–P2 =

–V2 =

• A balloon at 27 oC has a volume of 4.0 L. What happens to the volume of the gas when it is heated to 57 oC?

• The gas in an aerosol can has a pressure of 100 kPa at a temperature of 27 oC. What is the new pressure if the temperature is raised to 927 oC?

• A balloon has a volume of 20.0 L, a pressure of 150 kPa, and a temp of 40oC. What is the new volume when the gas is at 101.3 kPa and 0oC?

• A container with a volume of 20.0 L of N2 gas reaches a pressure of 20,000 kPa at 300 K. How many moles of N2 gas does the container hold?

Heterogeneous mixtures

• Suspensions –

– Particle size:

• Colloids –

– Particle size:

– Brownian motion:

– Tyndall effect:

Homogeneous Mixtures

• Solutions

– Solute

– Solvent

– Can be solid, liquid, or gas

• Soluble -

• Insoluble -

• Miscible -

• Immiscible -

• Molarity (M)

• What is the molarity of a solution containing 3 moles of solute in 1.5 L of solution?

• What is the molarity of 155 mL of solution containing 1.55 g dissolved KBr?

• Diluting molar solutions:

• What volume of 2.00M CaCl2 stock solution would you use to make 0.50L of 0.300M CaCl2 solution?

Colligative Properties of Solutions

• Colligative properties depend on number of solute particles in a solution

• Boiling point elevation

• Freezing point depression

Acids and Bases

• Acids produce H+ ions which react with water to form hydronium (H3O+) ions

• Bases produce OH- ions

Macroscopic properties of acids and bases

• Taste and feel– Acids taste sour (lemon juice, vinegar)– Bases taste bitter– Bases are slippery (soap)

• Litmus test and other color changes– Indicators change colors in the presence of an acid

or a base– Litmus: acid = red, base = blue

• Acids & bases are electrolytes– Substance that dissolve in pure water to form ions

& conduct electricity – Not all acids & bases conduct electricity equally

well• Strong acids & bases conduct electricity better than

weak acids & bases

Strengths of Acids and Bases

• Strong acids and bases ionize completely while weak acids and bases ionize only partially

The pH Scale

• A mathematical scale in which the concentration of H+ ions in a solution is expressed as a number from 0 – 14

Interpreting the pH scale

• pH < 7 = acidic• pH = 7 = neutral• pH > 7 = basic

• Each unit of pH represents a power of 10– Something with pH of 2 is 10 times more acidic

than something with a pH of 3

pH = -log [H+]

• What is the pH of solutions having the following ion concentrations?

1. [H+] = 1.0 x 10-2 M

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

• Relating H+ and OH- ion concentration

pOH = -log [OH-]

• What is the pOH of a solution having the following ion concentration?

1. [OH-] = 1.0 x 10-6 M

2. [OH-] = 6.5 x 10-4 M

pH + pOH = 14

• What is the pOH of a solution whose pH is 5?

Strong Acid + Strong Base

• Strong acids completely ionize

• Strong bases completely ionize

• Ionic equation – everything (aq) written as ions– Spectator ions– Net ionic equation

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

• H2SO4(aq) + 2KOH(aq) K2SO4(aq) + H2O(l)

Strong Acid + Weak Base

• Weak bases do not completely ionize

• 3HBr(aq) + Al(OH)3(s) AlBr3(aq) + H2O(l)

Weak Acid & Strong Base

• HC2H3O2(aq) + NaOH NaC2H3O2 + H2O

Bronsted-Lowry Acids & Bases

• Acid – H+ donor• Base – H+ acceptor

• HC2H3O2 + NH3 NH4 + + C2H3O2-

• Conjugate base – formed when acid donates proton

• Conjugate acid – formed when base accepts a proton

• HC2H3O2 + NH3 NH4 + + C2H3O2-

• Identify the acid, base, conjugate acid, and conjugate base in the following:

• HCO3 -1 + H2O CO3 -2 + H3O +1

• A 15.0 mL sample of a solution of H2SO4 with unknown molarity is titrated with 32.4 mL of 0.145 M NaOH to the endpoint. What is the molarity of the sulfuric acid solution?