AP CHEMISTRY- Final Project

AP CHEMISTRY- Final Project

1. UNIT 1: Atomic Structure and Properties

   • Moles and Molar Mass
     • Molar Mass (grams/mole): Amount of frams there are in a mole
     • Periodic Table: Tabular display of the chemical elements.
     • Avogadro’s Number: Equal to 6.022 x 1023
   • Mass Spectroscopy of Elements
     • Isotope: Elements with same number of protons and elements but different number of neutrons
   • Elemental Composition of Pure Substances
     • Solids: Composed of densely packed atoms, fixed volume and fixed shape
     • Liquids: Atoms assume the shape of the container they are occupying, fixed volume, but no fixed shape.
     • Gasses: Atoms flow freely and expand to fill their container, no fixed volume or shape.
     • Law of Definite Proportions: Ratio of mass of the constituent elements in any pure sample of that compound is always the same.
   • Composition of Mixtures
     • Mixture: Composition of two or more elements or compounds
       • Homogeneous: Composed of one type of element or compound
       • Heterogeneous: Composed of more than one element or compound
   • Atomic Structure and Electron Configuration
     • Dalton’s Theory: Atoms are indivisible and indestructible.
     • Coulomb's Law
     • Bohr Atomic Model
   • Photoelectron Spectroscopy
     • Photoelectric effect: Demonstrates that electrons are emitted from a metal surface when light strikes it if the frequency is high enough
     • Photoelectron spectroscopy (PES): Method to compare the relative energies of atoms ions, and molecules
   • Periodic Trends
   • Valence Electrons

2. UNIT 2: Molecular and Ionic Compound Structure and Properties

   • Types of Chemical Bonds
     • Ionic Bonds: Transferring Electrons (Usually involves a metal and a nonmetal)
     • Covalent Bonds: Sharing Electrons (Usually involves two nonmetals)
       • Nonpolar - electrons shared equally
       • Polar - Electrons shared equally
     • Metallic Bonds: Electrons not associated with a single atom or molecule (delocalized electrons)
     • Bond Polarity: Difference in electronegativity values of two elements.
       • High electronegativity (negative partial charge)
       • Low electronegativity (positive partial charge)
   • Intramolecular Force and Potential Energy
     • Bond Energy: Energy required to break the bond
     • Bond length: Distance between atoms
     • Potential Energy curves for covalent bonds
       • If potential energy is greater than zero, it is an unstable atom arrangement.
       • Less than zero indicates a stable arrangement.
     • Coulomb's Law: Interactions between cations and anions can be explained.
       • Longer Ion Charge: Larger Force of Attraction
       • Greater Distance between charges: Smaller Force of Attraction.
   • Structure of Ionic Solids
     • Conditions of conductivity (Ions) : Charged Particles; Particles must be free to move
     • Solid: Doesn’t conduct electricity Molten (melted) - Conducts electricity
   • Structure of Metals and Alloys
     • Metallic Bonds: Positive metal ions surrounded by a ‘sea of mobile valence electrons’ (delocalized electrons)
       • Properties: Good conductors of electricity; Malleable - bendable; Ductile - pulled in a wire
     • Alloys: Combining two or more metallic elements
       • Substitutional Alloy
       • Interstitial Alloy
   • Lewis Diagrams
     • Octet Rule: Having 8 electrons around an element.
   • Resonance and Formal Charge
     • Resonance: A way of describing valence electrons and its delocalization of electrons within molecules. It involves constructing multiple Lewis Structures that, when combined, represent the full electronic structure of the molecule.
     • Formal Charge: The hypothetical charge the atom would have if all of the atoms had the same electronegativity. The best lewis diagrams will have the minimum number of nonzero formal charges.
     • VSEPR and Bond Hybridization
       • VSEPR: Valence Shell Electron Pair Repulsion. Bonds and lone pairs will arrange themselves in order to minimize repulsion.
       • Molecular Geometry: The geometry of molecules will depend on the number of electron domains (bonds and lone pairs) around the central atom.

3. Unit 3: Intermolecular Forces and Properties

   • Intermolecular Forces
     • Dipole-Dipole Interactions
       • Dipole-Induced Dipole
     • London Dispersion Forces (LDF)
     • Hydrogen Bonding
     • Ion-Dipole Interaction
   • Properties of Solid
     • Ionic: High Melting Point and Boiling Point; Brittle; Poor Conductors
     • Molecular: Relatively low melting point and boiling point; Poor Conductors
     • Metallic: Great conductors of heat and electricity; Malleable and ductile; Melting points vary depending on the metal.
     • Covalent Network: Very high melting point and hardness; Poor conductors of electricity.
   • Solids, Liquids, and Gasses
   • Ideal Gas Law
   • Kinetic Molecular Theory
     • Summarizes the ideal behavior of gasses. 1) Random continuous motion 2) Perfectly elastic collisions (no stickiness) 3) Negligible volume particle volume 4) Constant temperature = Constant average kinetic energy.
   • Deviation from Ideal Gas Law
     • Real gas behaviors: PV is not equal to nRT
     • Happens when: Low temperature; High Pressures; Particles with significant intermolecular forces; Particles with a significant molecular size.
   • Solutions and Mixtures
     • Solution: A physical combination of any state of matter in which macroscopic properties do not vary.
     • Heterogeneous mixtures do have varying properties depending on location in the mixture.
   • Representations of Solutions
     • Mixtures can be represented with ion sizes and orientation of solute ions and solvent particles.
     • Solutions can be represented through relative concentrations of the components in the solution and a drawing that shows interactions among the components.
   • Separation of Solutions and Mixtures Chromatography
     • Chromatography
     • Distillation
   • Solubility
     • Substances with similar intermolecular attractions will tend to be soluble (or miscible) in one another.
       • Ionic compounds will tend to dissolve in polar solvents.
       • Molecular compounds that do not have dipoles and have predominantly London Dispersion forces will tend to dissolve in nonpolar solvents.
     • The degree of polarity and presence of intermolecular forces determines solubility in a particular solvent.
   • Spectroscopy and Electromagnetic Spectrum
     • Spectroscopy: The study of matter’s interactions with electromagnetic radiation
     • Matter can absorb or emit radiation in different regions of the spectrum, and those regions are associated with molecular motion or electronic transitions. Types: Microwave Radiation; Infrared Radiation; UV/Visible Radiation
   • Photoelectric Effect
     • The wavelength of the photon is related to the frequency of the radiation.
     • The energy of a photon is related to the frequency of the wave through Planck’s equat
     • Beer-Lambert Law: Relates the absorption of light by a solution to three variables.

4. Unit 4: Chemical Reactions

   • Net Ionic Equations
     • Chemical equation for a reaction that lists only those species participating in the reaction.
   • Representations of Reactions
     • Law of conservation of mass: Amount of matter stays constant in a closed system.
   • Physical and Chemical Changes
     • Physical Change: Changes an object, but doesn’t change the chemical structure (Ex: Water-> Steam)
     • Chemical Change: Changes chemical structure (Ex: Burning of paper)
   • Stoichiometry
     • Using mole ratios and measurements to manipulate them and get our desired unit.
   • Introduction to Titration
     • Experimental method used to analyze acid-base reactions.
   • Types of Chemical Reactions
   • Introduction to Acid-Base Reactions
     • Bronsted-Lowry Acid/Base: Acids: Proton donors, bases: proton acceptors
   • Oxidation-Reduction (Redox) Reactions
     • Reactions deal with the transferring of electrons

5. Unit 5: Kinetics

   • Reaction Rates
     • Kinetics: Rate of reaction, how fast a reaction happens.
     • Rate = -Δ[Reactant]/t or as Rate = Δ[Product]/t
   • Introduction to Rate Law
     • Zeroth order: k is in M/s
     • First Order: k is in s^-1
     • Second Order: k is in 1/Ms
       • Rate Law
   • Concentration Changes Over Time
   • Elementary Reactions
     • Chemical reaction in which one or more chemical species react directly to form products in a single reaction step and with a single transition state.
   • Collision Model
   • Reaction Energy Profile
     • Exothermic Reaction: System loses energy
     • Endothermic Reaction: System gains energy
     • Activation energy: Energy required to break the bonds in a reaction to go from the reactants to the activated complex
   • Introduction to Reaction Mechanisms
     • Takes a full reaction and splits up into elementary steps, what actually happens in a reaction.
   • Reaction Mechanism and Rate Law
   • Steady-State Approximation
   • Multistep Reaction Energy Profile
     • Energy profile is used to compare activation energies of the different steps and identify the rate determining step.
   • Catalysis
     • Process of increasing the rate of a chemical reaction by adding a substance known as a catalyst.

6. Unit 6: Thermodynamics

   • Endothermic and Exothermic Processes
     • Kinetic Energy: Energy that results from the motion of an object Potential Energy: Stored energy
   • Energy Diagrams
   • Heat Transfer and Thermal Equilibrium
   • Heat Capacity and Calorimetry
     • Heat Capacity: Amount of heat required to raise the temperature of an object by 1oC
   • Energy of Phase Changes
   • Introduction to Enthalpy of Reaction
     • Enthalpy: Total heat content of a system
   • Bond Enthalpies
     • Bond Enthalpy: Describes the amount of energy stored in a bond between atoms in a molecule
   • Enthalpy of Formation
     • Change of enthalpy during the formation of 1 mole of the substance from its constituent elements in their most stable state at standard conditions.
   • Hess's Law
     • Energy change in an overall chemical reaction is equal to the sum of the energy changes in the individual reactions comprising it

7. Unit 7: Equilibrium

   • Introduction to Equilibrium
   • Direction of Reversible Reactions
     • Reversible reactions: Reactions that can go forward and backward.
   • Reaction Quotient and Equilibrium Constant
     • Reaction Quotient
     • Equilibrium Constant
   • Calculating the Equilibrium Constant
   • Magnitude of the Equilibrium Constant
     • k>1 Reaction is product favored
     • k=1 Reaction is product favored or reactant favored.
     • k<1 Reaction is reactant favored
   • Properties of the Equilibrium Constant
   • Calculating Equilibrium Concentrations
     • To apply this you need the equilibrium constant, initial concentration
   • Representations of Equilibrium
   • Introduction to Le Châtelier’s Principle
     • Le Châtelier’s Principle: If a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifhts to counteract the change to reestablish an equilibrium.
   • Introduction to Solubility Equilibria
     • Solubility Equlibrium: Type of dynamic equilibrium that exists when a chemical compound in the solid state is in chemical equilibrium with a solution of that compound.
   • Common-Ion Effect
     • Common-Ion Effect: Describes how a common ion can suppress the solubility of a substance
     • Used when dissolution takes place in another solution.
   • pH and Solubility
     • pH: Measure of how acidic/basic water is
     • Effect of acidic solution: Increase the solubility of a compound that forms the conjugate base of a weak acid.
     • Effect of basic solution: Compounds that dissolve into significant conjugate acids will be more soluble.
   • Free Energy of Dissolution
     • Entropy: amount of disorder in a system
     • Gibbs Free Energy
       • Negative: Thermodynamically favorable
       • Positive: Thermodynamically unfavorable

8. Unit 9: Applications of Thermodynamics

   • Introduction to Entropy
   • Absolute Entropy and Entropy Change
     • The entropy change for a process can be calculated from the absolute entropies of the species involved before and after the process occurs.
     • Entropy is often expressed in joules per kelvin rather than kilojoules.
   • Gibbs Free Energy and Thermodynamic Favorability
     • Gibbs free energy change describes whether a reaction is thermodynamically favorable or unfavorable.
       • Negative: Thermodynamically Favorable
       • Positive: Thermodynamically Unfavorable
   • Thermodynamic and Kinetic Control
     • Some reactions are thermodynamically favorable but do not appear to make products at a measurable rate. These processes are under “kinetic control” and generally have a large activation energy that accounts for the slow rate of reaction.
     • Catalysts may be used to decrease activation energy and increase reaction rate, but this will have no effect on the thermodynamic favorability of the reaction.
   • Free Energy and Equilibrium
     • At equilibrium, no net change in concentration of reactants and products occurs
   • Coupled Reactions
     • Two reactions that share a common intermediate (a product of one reaction is a reactant in the other) can be coupled.
     • Hess’s law is applied to determine the free energy change, for the coupled reaction.
     • The sum of the reactions’ G values is negative, making the overall process thermodynamically favorable.
   • Galvanic (Voltaic) and Electrolytic Cells
     • The electrochemical cell is a device that can convert energy released by a thermodynamically favorable redox reaction to electrical energy, or one that, conversely, can use electrical energy to drive a thermodynamically unfavorable redox reaction.
     • Galvanic (or voltaic) cells involve a thermodynamically favorable reaction, while electrolytic cells involve a thermodynamically unfavorable reaction.
     • Both types of cells contain an anode, where oxidation occurs, and a cathode, where reduction occurs.
   • Cell Potential and Free Energy
     • Cell Potential (V) : Forces electrons traveling through a wire from the anode to the cathode feel.
     • Standard Reduction Potentials
   • Cell Potential Under Nonstandard Conditions
     • Standard conditions for electrochemical cell: 298.15K, 1 atm, 1M concentration of both solutions in either half cell, and Q=1
     • Nernst Equation: Shows relationship between Eo cell and Ecell and Q
   • Electrolysis and Faraday’s Law
     • Faraday’s Law of Induction
     • Electrolysis: technique that uses direct electric current to drive an otherwise non-spontaneous chemical reaction

9. Unit 8: Acids and Bases

   • Introduction to Acids and Bases
     • Arrhenius definition
       • Acid: Any compound that increases the concentration of H+
       • Base: Any compound that increases the concentration of OH-
     • Brønsted-Lowry definition
       • Acid: H+ donator
       • Base: H+ Acceptor
   • pH and pOH of Strong Acids and Bases
     • pH: Concentration of protons in a solution, measures concentration of H+
       • pOH: Measures concentration of OH-
         • pOH = -log([OH-])
       • pH = -log([H+])
     • Strong acid or base is one that completely dissociates in water
   • Weak Acid and Base Equilibria
     • Weak acids and bases incompletely dissociate in water
   • Acid-Base Reactions and Buffers
     • Buffers: Mixture of an acid (HA) and its conjugate base (A-) in a solution. Resistant to changes in pH
     • Strong Acid and Base reaction: acid fully dissociates
     • Weak Acid Strong Base reaction: Acid doesn’t fully dissociate
   • Acid-Base Titrations
     • Titration: Procedure taken to find the concentration of an unknown “analyte” by dripping a solution of known concentration into the analyte until an equivalence point is met.
     • Titration Curve
       • Strong Base Weak Acid
       • Strong Acid Weak Base
   • Molecular Structure of Acids and Bases
     • Acid-base strength of a molecule depends strongly on its structure
     • Stronger the bond = Weaker the acid
   • pH and pKa
     • Henderson-Hasselbalch Equation: Used to find pH of a buffer
     • Acid-Base indicators
   • Properties of Buffers
     • Buffers have pH resistance due to having both acid and a base that do not actively react together at equilibrium
   • Henderson-Hasselbalch Equation
     • Finds pH of buffer
   • Buffer Capacity
     • Helps us see how much acid/base one can add until there is a significant change in pH.
     • Determined by the magnitudes of the concentrations.

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