✍️ Free Response Questions
AP Chemistry Free Response Questions
⚛️ Unit 1 - Atomic Structure and Properties
1.1Moles and Molar Mass
1.2Mass Spectroscopy of Elements
1.3Elemental Composition of Pure Substances
1.4Composition of Mixtures
1.5Atomic Structure and Electron Configurations
1.6Photoelectron Spectroscopy & Graph Interp.
🤓 Unit 2 - Molecular and Ionic Compound Structures and Properties
2.0Unit 2 Overview: Molecular and Ionic Bonding
2.1Types of Chemical Bonds
2.2Intramolecular Force and Potential Energy
2.3Ionic Bonding and Ionic Solids
2.4Metallic Bonding and Alloys
2.5Lewis Dot Diagrams
2.6Resonance and Formal Charge
🌀 Unit 3 - Intermolecular Forces and Properties
3.0Unit 3 Overview: Intermolecular Forces and Properties
3.2Properties of Solids
3.3Solids, Liquids, and Gases
3.4The Ideal Gas Law
3.5The Kinetic Molecular Theory of Gases
3.6Deviations from the Ideal Gas Law
3.7Mixtures and Solutions
3.8Representations of Solutions
3.9Separation of Solids/Mixtures
3.10Solubility and Solubility Rules
3.11Spectroscopy and the Electromagnetic Spectrum
3.12Quantum Mechanics and the Photoelectric Effect
🧪 Unit 4 - Chemical Reactions
4.0Unit 4 Overview: Chemical Reactions
4.1Recognizing Chemical Reactions
4.2Net Ionic Equations
4.4Physical vs. Chemical Changes
4.5Stoichiometry & Calculations
4.6Titrations - Intro and Calculations
4.8Intro to Acid-Base Neutralization Reactions
👟 Unit 5 - Kinetics
5.0Unit 5 Overview: Kinetics
5.1Defining Rate of Reaction
5.2Introduction to Rate Laws
5.3Rate and Concentration Change
5.4Writing Rate Laws
5.5Collision Model of Kinetics
5.6Reaction Energy and Graphs w/ Energy
5.7Reaction Mechanisms and Elementary Steps
5.8Writing Rate Laws Using Mechanisms
🔥 Unit 6 - Thermodynamics
6.0 Unit 6 Overview: Thermochemistry and Reaction Thermodynamics
6.1Endothermic Processes vs. Exothermic Processes
6.2Energy Diagrams of Reactions
6.3Kinetic Energy, Heat Transfer, and Thermal Equilibrium
6.4Heat Capacity and Coffee-Cup Calorimetry
6.5Phase Changes and Energy
6.6Introduction to Enthalpy of Reaction
6.7Bond Enthalpy and Bond Dissociation Energy
6.8Enthalpies of Formation
⚖️ Unit 7 - Equilibrium
🍊 Unit 8 - Acids and Bases
8.0Unit 8 Overview: Acids and Bases
8.1Introduction to Acids and Bases
Unit 9 - Applications of Thermodynamics
🤺 AP Chemistry Essentials
🧐 Multiple Choice Questions
AP Chemistry Self-Study and Homeschool
⏱️ 3 min read
August 6, 2020
The amount of solute needed to form a saturated solution at any particular temperature is the solubility of that solute at that temperature🌡️.
The solubility of one substance in another depends on:
The tendency of systems to become more random (by becoming more dispersed in space)
The relative intermolecular solute-solute energies compared with solute-solvent interactions.
Polar and ionic solutes tend to dissolve in polar solvents, and non-polar solutes tend to dissolve in non-polar solvents. (Remember “like dissolves like”!)
There are a few solubility rules that will be helpful on the AP Exam (though all but few of these are necessary to commit to memory):
Image Courtesy of Quizlet
💡You can easily notice that many of the same ions (Notably Ag, Hg2 2+, and Pb2+) are often exceptions to these solubility rules. However, do not try to memorize these completely! These, like things like polyatomic ions, are memorized implicitly through usage! That's why by the end of AP Chemistry, most students don't think about charts like these, they simply know them through doing a bajillion problems.
Every solution, no matter the solute and solvent, has something called a saturation point. Essentially, this is the point at which no more solute can be dissolved in the solvent.
It depends on three things: the temperature, the solvent, and the solute. At a higher temperature, more solute can be dissolved (for the majority of solutes). This can be seen in the following solubility curve:
Image Courtesy of Dynamic Science
For example, looking at a saturated solution (note, saturated is important here, we'll get to that in a moment) of 100g H2O and KCl at approximately 70C, there will be around 50g KCl dissolved in the solution.
A saturated solution is a solution in which all possible solute has been dissolved in a solution. Once you have reached the saturation point, any additional solute (at that temperature) will fall out of solution and not dissolve. Essentially, with saturated solutions, you are on the solubility curve.
Undersaturated solutions are what their name implies - a solution that has not yet reached saturation. With an undersaturated solution, you can add more solute at the current temperature that still dissolves. Like saturated solutions, an undersaturated solution can be modeled using a solubility curve, representing a point below the solubility curve.
Supersaturated solutions are a type of solution where you have more solute than you can dissolve in the solvent. In order to accomplish this, a chemist heats up the solution and then adds enough solute to saturate it at that temperature, and then slowly cools it down to the temperature they need. This accomplishes a supersaturated solution that when agitated will produce crystals as the extra salt falls out of solution.
A supersaturated solution of CH3COONa, GIF Courtesy of Gyfcat
As mentioned before, polarity has a huge effect on solubility. Like dissolves like! Nonpolar substances are more likely to dissolve in nonpolar solvents and polar substances are more likely to dissolve in polar solvents.
Pressure only affects the solubility of gases. As pressure increases, solubility increases as well. This phenomenon explains why soda🥤 becomes flat overtime.
This can also be seen using Henry's Law: C = kP, or (solubility of gas in M) = (k Constant) (Partial pressure of [g] above solution). This equation is just an extra piece of information - not necessary for the AP.
For solids and liquids, pressure does not affect solubility.
For gases only: as temperature increases, the solubility of the gas decreases. As mentioned before, as temperature increases, the solubility of liquids and solids increase as well.
👉More about this: Unit 7 later in the course!
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