Advanced Chemistry Final Exam Review Guide

Sig Figs

• Placeholder zeros are removed.
• All zeros after the decimal point are significant.

Percent Error

• PercentError = |(Result / TheoreticalResult)| * 100

Nomenclature

• Ionic:
  • Cation (positive ion) full name first.
  • Anion (negative ion) with -ide ending.
  • Consider oxidation numbers.
• Covalent:
  • Cation first (use prefix unless only one atom, then omit mono).
  • Anion second (use prefix and -ide ending).

Phase Changes

• Solid to gas: sublimation
• Solid to liquid: melting
• Liquid to solid: freezing
• Liquid to gas: evaporation/boiling
• Gas to liquid: condensation
• Gas to solid: deposition

Chemical Formula

• Elements with subscripts to achieve electrical neutrality (e.g., H2O, CO2).

Equilibria

• Reaction/process with a double arrow indicates equal forward and backward reaction.
• Represented in phase diagrams for state changes.

Kinetic Factors

• Factors affecting reactivity: concentration, catalyst, temperature, pressure.

Calculations with Sig Figs

• Keep non-zero numbers until only one number is in the ones place.

Empirical Formula

• Simplified molecular formula; ratio of elements in a compound.

Grams -> Moles

• Divide by molar mass.

Molar Ratio

• Moles of one part / total moles (percent).
• Moles of one substance / moles of another (conversion factor).

Balancing Chemical Equations

• Number of products and reactants must be equal.

Stoichiometry

• PV = nRT (Pressure, Volume, moles,Ideal gas constant, Temperature).
• Ideal gas law: (P1)(V1)/(T1) = (P2)(V2)/(T2).
• PV = (mRT)/M
• PM = DRT

IMFs (Intermolecular Forces)

• LDFs (London Dispersion Forces):
  • Minor forces from unequal electron dispersion.
  • Present in all molecules.
• Dipole-Dipole:
  • Force between oppositely charged polar molecules.
  • Occurs only between polar molecules.
• Hydrogen Bonds:
  • Formed between hydrogen and electronegative atom (O or F) on different molecules.

Strengths of IMFs (Weak to Strong)

• London dispersion forces -> dipole-dipole -> hydrogen bonds

Phase Changes (Enthalpy)

• MassReactants * SpecificHeat * ChangeInTemperature (MSHAT)

Phase Diagrams

• Triple point: all three states of matter coexist.
• Critical point: liquid can no longer exist.
• Axes: Pressure vs. Temperature.
  • Left: Solid
  • Middle: Liquid
  • Right: Gas

STP (Standard Temperature and Pressure)

• Standard temperature: 0°C or 273 K
• Standard pressure: 1 atm, 101.3 kPa, or 760 mmHg

Molar Volume at STP

• One mole of any gas at STP is 22.4 L.

Hydrogen Bonds

• Strongest IMF.
• Form between H atom of one molecule and F or O atom of another.
• Water forms hydrogen bonds.

Temperature / Kinetic Energy

• Higher temperature = higher kinetic energy.

Molar Volume

• 22.4 L of one mole of any gas at STP.

Dalton's Law

• Total pressure equals the sum of partial pressures: P = P1 + P2

Enthalpy

• \Delta h is the total energy of a reaction.
• Sum of products minus sum of reactants.

Endothermic/Exothermic

• Endothermic: reaction gains energy.
• Exothermic: energy is expelled in a reaction.

Ideal Gas Law

• (P1)(V1)/(T1) = (P2)(V2)/(T2)

Calculating Enthalpy (∆h)

• Sum of products minus sum of reactants.

Calculating Entropy (∆s)

• Sum of products minus sum of reactants.

Energy Diagrams

• Endothermic: energy of reaction increases.
• Exothermic: energy of system decreases.
• Energy spikes during reaction for activation energy.

Kinetic Factors of Reaction Rate

• Temperature
• Pressure
• Surface area

Q=Msh∆t

• Q = total heat used
• M = mass (g)
• Sh = specific heat of substance (energy to raise temp by 1°C)
• ∆t = change in temperature (°C)

Behavior of Gases

• Compressible
• High kinetic energy
• Take shape/volume of container
• No interactions
• Ideal gases (Under STP):
  • Elastic collision
  • No volume, no interactions (Kinetic molecular theory)

IMF, MF, and Properties (BP/MP)

• Molecular formula -> structure -> IMFs and properties.

Reaction of Gases Particle Diagram

• Picture -> formula -> stoichiometry.

Empirical Formula

• Simplified molecular formula (ratio of elements).

Combustion

• Exothermic.
• Uses oxygen and produces CO2.

Ideal Gases

• React completely.
• Are under STP.

Properties of Ionic Compounds

• High melting points.
• Hard and brittle.
• Dissociate into ions in water.
• Solutions/melted compounds conduct electricity, solids do not.

Coulomb's Law

• F = K((Q1)(Q2) / R^2)
• Q = charges, K = constant, F = force, R = distance

Heat Curve ∆H, ∆S

• Q = (Mass)(Specific Heat)(∆T)
• Q = (Moles)(Enthalpy)

Hybridization of Atoms

• Covalent bonding → Lewis structure → Shape → Hybridization
  • Tetrahedral Electronic Geometry -> Bent, sp3
  • Trigonal Planar Electronic Geometry -> Linear, sp2
  • Tetrahedral Electronic Geometry -> Trigonal Pyramidal, sp3

Polarity/ Like Dissolves Like

• Polar dissolves in polar, nonpolar in nonpolar.

Reaction / ∆S (Particle Diagrams)

• Bond forming: entropy decreases (more order).
• Bond breaking: entropy increases (less order).
• Nature seeks higher entropy.

Reaction Rate Factors

• Size of object
• Surface area
• Pressure
• Temperature

Solvation Particle/ ∆h, ∆s Diagram

• Breaking bonds: ∆h positive (energy in), ∆s positive (more randomness).
• Forming bonds: ∆h negative (energy out), ∆s negative (less randomness).
• Endothermic: energy added.
• Exothermic: energy leaves.

Kinetic, Activated Complex

• ∆h: difference between start and end of reaction.
• Activation energy: energy needed for reaction to occur.
• Peak: activated complex.

Ionic Compounds diagram

• When force is used to hit an ionic compound they split because they break when same forces are lined up

Particle Diagram Reaction

• # of molecules of each type are equal on both sides ALWAYS.
• Remember diatomics (H2, N2, F2, O2, I2, Cl2, Br2)

Dimensional Analysis

• Use conversion factors to cancel units and obtain desired unit.

EF and MF

• EF:
  • Assume 100g, % becomes grams.
  • Divide by molar mass to get moles.
  • Set up equation with mole subscripts.
  • Divide each subscript by smallest.
  • Write out results in EF.
• MF:
  • Determine molar mass of empirical formula.
  • Divide molar mass of compound by EF molar mass.
  • Multiply subscripts of EF by this new #.
  • Write out MF.

Stoichiometry

• Grams -> moles, divide by molar mass to get moles, divide by molar ratio then multiply by molar mass to get grams of desired substance

Ideal Gas Law

• (P1)(V1)/(T1) = (P2)(V2)/(T2)

∆G

• \Delta g = \Delta h - T\Delta s
  • T in Kelvin, ∆s in kJ (divide by 1000).
• Spontaneity:
  • \Delta g > 0: non-spontaneous.
  • \Delta g < 0: spontaneous.
  • \Delta g = 0: equilibrium.

∆h & ∆s

• Sum of products minus sum of reactants.

Gas Stoichiometry

• Convert grams of gas to moles.
• Convert to moles of desired substance.
• Use PV = nRT to get desired value.

Calorimetry

• -msh\Delta t (metal) = msh\Delta t (water)
• -(mass metal)(specific heat of metal)(final temp - initial temp)= (mass water)(specific heat water)( final temp water- initial temp water)

Hydrate

• Find grams of water (mass of hydrate - mass of anhydrous).
• Divide by molar mass of water to get moles of water.
• Divide anhydrous mass by its molar mass to get moles.
• Divide moles anhydrous by moles water to get # of water molecules.
• Write out the hydrate (anhydrous should always have a coefficient of one
• Na4cl3 * 5H2O

Identifying IMFs/Solubility

• Nonpolar: only LDFs.
• Polar: LDFs, dipole-dipole.
• Hydrogen bonding: LDFs, dipole-dipole, hydrogen bonds (H bonded to F, N, O).
• Solubility: nonpolar dissolves nonpolar, polar dissolves polar.

IMFs and Properties

• IMFs influence melting point, boiling point, viscosity, and vapor pressure.
• Boiling Point: temperature at which vapor pressure equals atmospheric pressure.
• Melting Point: solid changes to liquid at atmospheric pressure.
• Viscosity: liquid's resistance to flow (stronger IMFs = higher viscosity).
• Vapor Pressure: increases exponentially with temperature. When boiling, vapor pressure equals external pressure.