Chemistry Final Exam Study Guide

5 natures of gases

5 natures of gases

1. Gasses have mass (low density compared to solids and liquids)

2. Gasses are compressible

3. Gasses fill their containers

4. Gasses diffuse (move from areas of high to low concentration)

5. Gasses exert pressure (sum of collisions of gas molecules with themselves and their container)

Things that affect rate of diffusion

Size & energy influence speed of diffusion (heavier atomic mass -> slower diffusion)

Atmospheric pressure

Pressure of the gas in the air on the earth (measured by a barometer)

Gas law units

• Pressure: 1 atm = 760 torr = 760 mmHg = 101.3 kPa

• Temperature: 0ºC = 273.15 K (MUST USE KELVIN!!!)

• Volume: 1 L = 1000 mL = 1000 cm^3

Dalton’s law of partial pressure

P total = P1 +P2 + P3… 

Ideal Gas Law

PV = nRT (number of moles (n) of contained gas)

International Gas Law Constant (R)

• Changes depending on unit of pressure

  • R for atm = 0.0821

  • R for mmHg and torr = 62.4 

  • R for kPa = 8.314

Manometer

Tool used to measure the pressure of contained gas by comparing it to atmospheric pressure

Manometer Problems (P = pressure)

• If Pgas > Patm: Pgas = Patm + ∆z

• If Pgas < Patm: Pgas = Patm - ∆z

Use mole ratio when solving for

• Molar mass

• Na -> Nb

• # molecules a -> # molecules b

Avigadro’s #

1 mole (6.02*10^23) = the amount of molecules in an element as appears on periodic table

4 unusual properties of water

1. High surface tension: Water is polar & very attracted to itself via h-bonding which creates a sort of skin on the surface of the water

2. High boiling point: Water is polar & very attracted to itself via h-bonding, so it takes lots of energy for H2O to break h-bonds and escape as vapor

3. Ice is less dense than water: Frozen water forms a lattice pattern which fills with air and causes ice to be less dense than water

4. Water is the universal solvent: Because water is polar, it has the ability to attract other polar molecules and pull them apart

Like dissolves like

Polar dissolves polar, nonpolar dissolves nonpolar

Factors affecting rate of solvation and why

• Stirring: makes H2O molecules tear the solute apart

• Temperature: increased temp -> increased KE -> increased collisions and contact between solute and solvent

• Particle Size: Decreased particle size -> increased surface area and contact

Factors affecting solubility:

• Gasses

  • Temperature is directly proportional to solubility

    • Increase in temp -> increase in solubility

  • Pressure: N/A

• Solids

  • Temperature is inversely proportional to solubility

    • Increase in temp -> decrease in solubility

  • Pressure is directly proportional to solubility

    • Increase in pressure -> increase in solubility

Ion dipole forces

Exist between an ionic compound in water (ex: Cl and positive poles in water)

Strongest IMF!!!

Hydrogen bonding

Stronger version of DP-DP, hydrogen with strong positive charge due to electronegativity of O, N, or F becomes attracted to the O, N, or Fs of other molecules

Dipole-Dipole

Polar molecules with oppositely charged ends attract each other

LDF (London dispersion forces)

Exists between all molecules; is a temporary force that exists when adjacent electrons are positioned to make the atoms form temporary dipoles

Is the weakest IMF!!!

Colligative properties

properties that depend only on the number of solute particles and not their identity

BP elevation and FP depression

Difference in temp of boiling/freezing point of solution vs boiling/freezing point of pure solvent

Steps to solve colligative property problem

1. Solve for molality

2. Solve for ∆Tf & ∆Tb

   1. ∆Tf = Kf * m

   2. ∆Tb = Kb * m

3. Solve for new boiling or freezing point

   1. BP solution = BP + ∆Tb (elevation)

   2. FP solution = FP + ∆Tf (depression)

Collision theory

Particles that collide with sufficient energy will react to form products

Activation energy

The minimum energy needed by colliding particles in order to react

Activated complex

An unstable arrangement of atoms that exists momentarily at the peak of the activation energy barrier 

Reaction rate 4 factors affecting it

• 4 factors:

  • Concentration: More particles -> more collisions -> faster rate

  • Temperature: Increase in temp -> more collisions -> faster rate

  • Particle Size: Increase in surface area -> more collisions -> faster rate

  • Catalyst: Makes reaction easier by lowering activation energy

The rate law

• shows the relationship of the reaction rate to the rate constant and the concentrations of the reactants raised to some power

• For aA + bB -> cC + dD: Rate = k[A]^x[B]^y

Le chatlier’s principle

f stress is applied to a system (reaction) at equilibrium, the system changes to relieve the stress

Stresses that exist

• Change in concentration: If you remove a reactant or product, the reaction will favor the reaction making it (opposite also true)

• Change in temperature: If you heat a system it will favor the reaction using heat (opposite also true)

  • Exothermic: absorb heat (heat is a reactant)

  • Endothermic: give of heat (heat is a product)

• Change in pressure/volume:  If pressure is increased the system will favor the reaction making less moles of gas (opposite also true)

IF NO GASSES, NO EFFECT!!!

Acid properties

Acids contain an ionizable hydrogen, are sour, and have a pH of less than 7

Bases properties

Bases contain an ionizable hydroxide, are bitter, slippery when wet, and have a pH of more than seven

6 strong acids & bases: (Ionize completely in water)

6 strong acids	6 strong bases
HClO4 (perchloric acid)	LiOH (lithium hydroxide)
HCl (hydrochloric acid)	AQZNaOH (sodium hydroxide)
HBr (hydrobromic acid	KOH (potassium hydroxide)
HI (hydroiodic acid)	Ca(OH)2 (calcium hydroxide)
HNO3 (nitric acid)	Sr(OH)2 (strontium hydroxide)
H2SO3 (sulfuric acid)	Ba(OH)2 (barium hydroxide)

Arrhenius acids and bases

• Arrhenius acids give off hydrogen in water

• Arrhenius bases give off hydroxide in water

• Mono, di, & triprotic acids: Acids with one, two and three ionizable hydrogens

Bronstead lowry acids and bases: (relationship between them that does not involve water)

• Acids give off hydrogen ions

• Bases can gain hydrogen ions

• Amphoteric substances ca lose or gain a hydrogen ion

• Conjugate acids: The product formed when a base gains a hydrogen

• Conjugate bases: The product formed when an acid loses a hydrogen

Ion product constant (Kw)

At 25ºC, Kw = 1.0*10^-14

(In pure water [H] = 1.0*10^-7)

pH equations

• If given [H]

  • pH = -log([H])

  • [OH] = kw/[H]

• If given pH

  • pOH = 14-pH

  • [H] = 10 ^-pH

• If given [OH]

  • pOH = -log([OH])

  • [H] = kw/[OH]

• If given pOH

  • pH = 14-pOH

  • [0H] = 10 ^-pOH

Weak acids & bases

• Do not fully dissociate in water

• Weak acids → strong conj bases (opposite also true)

Ka & Kb

• Extent of proton transfer between the acid/base and H2O

• Determines strength of acid/base (smaller ka = weaker acid)

Ka equation

Keq * [H20] = [H3O][A]/[HA] (products/reactants)

Solving Ka/Kb steps (given: M and pH)

1. Is it a strong acid? if not, make ICE table

2. create ICE table & Ka equation

3. use pH to find [H]

4. plug in [H] for x in Ka equation

Solving [OH] & pH steps (given Ka/Kb, M)

1. check to see if [HA]/Kb > 500

2. create ICE table & Kb equation

3. use Kb equation to find Kb using algebra

4. use [OH] to find pH

Ha and Ka relationship

• if [HA]/Kb > 500: change in initial concentration of x is negligible (can remove x from E row)

• if [HA]/Kb < 500: change in initial concentration of x is not negligible (must keep x in E row)

% dissociation

• Acids: final [H30]/initial acid * 100

• Bases: final [OH]/initial base * 100

Neutralization reaction

when an acid and a base react and neutralize each other (moles H = moles OH, produces water)

Neutralized products formula

combine cation from base and anion from acid + HOH to make product formulas

Neutralization problems

Given V & M of one substance, V of another (use train tracks to solve for M)

Buret

a graduated glass tube with a tap at one end, for delivering known volumes of a liquid, especially in titrations

Titrant/standard solution

The substance of known concentration added to the analyte in a titration

Analyte

The substance of unknown concentration

Equivalence point

Point in a titration where neutralization occurs

End point

Point at which the indicator changes color in titration