Gas Laws

• The 5 natures of gasses:

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)

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

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

• Atmospheric pressure: Pressure of gasses in the air, varies with altitude

  • Measured with a barometer

• Definition and units of the following (conversion factors will be provided):

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

  • Temperature: 0ºC = 273.15 K

  • Volume: 1 L = 1000 mL = 1000 cm^3

• Standard temperature and pressure (STP): 0ºC & 1 atm (freezing point of water at sea level)

• The gas laws (MUST USE KELVIN!!!)

  • Boyles: P1*V1 = P2*V2 (pressure and volume are inversely proportional)

  • Charles: V1/T1 = V2/T2 (Volume and temp are directly proportional)

  • Gay Lussac's: P1/T1 = P2/T2 (pressure and temp are directly proportional)

  • Combined: P1*V1/T1 = P2*V2/T2

  • Dalton’s law of partial pressure: Ptotal = P1 +P2 + P3…

  • Ideal: 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

  • Gas stoichiometry

• 

  • Use mole ratio when solving for

    • Molar mass

    • Na -> Nb

    • # molecules a -> # molecules b

  • Avigadro’s number in stoichiometry

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

• Anything from the lab with all of the stations where we explored real world concepts

Solutions Concepts

• All things water

  • Structure: Water is a polar molecule because it has an uneven distribution of electrons

• 4 unusual properties and why

  • 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

  • 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

  • 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

  • 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

  • Cohesion: ability of a molecule to stick to itself

  • Adhesion: ability of a molecule to stick to something else

• Solution formation

  • Hydration: Solvation where water is the solvent

  • Solvation: Word for dissolving

  • Solute vs solvent: Solute is being dissolved, solvent is doing the dissolving

  • 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

  • Solubility: amount of solute that can dissolve at given temp (g solute/100 g solvent)

  • 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

  • Saturation: number of solute particles that are dissolved in a solvent at a given temperature

    • Saturated solution: max amount of solute dissolved at given temp

    • Unsaturated: less solute that can theoretically be dissolved at given temp

    • Supersaturated: through a process, more solute dissolved than theoretically possible

• Measurements of solution

  • Concentration: amount of solute in a solution to either solvent or total solution

    • Dilute solution: small amount of solute

    • Concentrated solution: max amount of solute

  • Molarity (M): moles of solute/liters of solution

  • Molality (m): moles of solute/kilograms of solvent

• Intermolecular forces: forces of attraction between molecules due to polarity and movement of electrons within the bond

  • The fundamental difference between states of matter is the strength if the IMFs holding them together

• Types of IMFs:

  • Ion dipole: 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

Medium IMF

• London dispersion: 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!!!

• Polarity: Uneven distribution of electrons creates positive and negative poles on a bond

• How IMFs influence boiling point & melting point: Stronger IMFs -> more energy required for state change -> higher boiling point & lower melting point

  • If same IMF, increase in mass of molecule -> increase in strength of dispersion force

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

  • BP elevation: Difference in temp of boiling point of solution vs boiling point of pure solvent

  • FP depression vs FP of solution: Difference in temp between freezing point of solution vs freezing point of pure solvent

  • Molal freezing point and boiling point constants: depend on solvent

    • Kf for water is 1.86º C/m

    • Kb for water is 0.521º C/m

• Steps to solve colligative property problems:

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)

Rate and Equilibrium

• Collision theory: Atoms, ions, and molecules with sufficient energy can react to form products when they collide

• 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:

  • What is it?

  • 4 factors affecting it

    • 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

• How to solve for rate law equations

  • How to determine the order of a reactant

  • What is “overall order”: reaction is (x+y)th order overall

  • How to set up rate law expression

  • How to solve for rate law constant

• Reversible reaction

  • Forward reaction

  • Reverse reaction

• Equilibrium: when the rate of the forward and reverse reactions are equal and the concentrations of reactants and products don’t change

  • Le chatlier’s principle: If stress is applied to a system (reaction) at equilibrium, the system changes to relieve the stress

    • What does it mean

• The 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!!!

• Keq expression

  • How to set it up

  • What the value tells you

  • ICE tables: stand for:

Initial concentration

Change in concentration

Equilibrium concentration

They are set up in the following manner:

• Given Keq

  • Asked for Keq

Acids and Bases

• Properties:

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

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

  • Both conduct electricity

• 6 strong acids & bases: Complete ionization in water

• 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

• pH

  • Ion product constant (Kw): At 25ºC, Kw = 1.0*10^-14

    • In pure water [H] = 1.0*10^-7

  • 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

  • pH and pOH scale: Tells you concentration of hydrogen

    • Low pH value = stronger acid

  • Indicators:

    • Acid/base indicators: Very specific, ~ 1.5 pH range

    • Universal indicators: Unspecific, 1-12 pH ranger

• Weak acids and bases: Do not fully dissociate (dissolve) in water

  • Pairing with conj acids & bases: 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)

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

6. create ICE table & Kb equation

7. use Kb equation to find Kb using algebra

8. 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

• Titration

  • Neutralization reactions: 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