AP Chemistry Study Guide Unit 3 PDF

Unit 3 Properties of Substances and Mixtures According to the College Board, "Transformations of matter can be observed in ways that are generally categorized as either a chemical or physical change. The shapes of the particles involved and the space between them are key factors in determi...

Unit 3 Properties of Substances and Mixtures According to the College Board, "Transformations of matter can be observed in ways that are generally categorized as either a chemical or physical change. The shapes of the particles involved and the space between them are key factors in determining the nature of physical changes. The properties of solids, liquids, and gases reflect the relative orderliness of the arrangement of particles in those states, their relative freedom of motion, and the nature and strength of the interactions between them. There is a relationship between the macroscopic properties of solids, liquids, and gases, as well as the structure of the constituent particles of those materials on the molecular and atomic scale. In subsequent units, students will explore chemical transformations of matter." In this unit, we dive deep into what happens when you put a bunch of molecules (or atoms, but mostly molecules) together. We begin with a look at intermolecular forces and the states of matter, then move into gases and solutions. This unit is super exciting because there are tons of real-world, out-of-lab experiences you have with these concepts every day and we'll come to learn all about them. About 18-22% of the AP exam is about this unit! It's the largest unit on the exam but really interesting, so let's get to it. 3.1 Intermolecular Forces The best way to remember the difference is that "inter" means between, while "intra" means within. From here, you can remember that intermolecular forces hold molecules together, while intramolecular forces hold atoms within a molecule together. 1. London Dispersion Forces (LDFs): These are the weakest IMFs and are present in all molecules. They are the only forces that hold together non-polar molecules and noble gases when they are in liquid or solid forms. LDFs happen when a molecule with a temporary uneven charge (called a temporary dipole) causes a similar dipole in a nearby molecule. The slightly negative side of one molecule attracts the slightly positive side of the other. 2. Dipole-Dipole Interactions: These are stronger than LDFs and happen in polar molecules. Since polar molecules have permanent dipoles, their positive and negative ends are always attracting each other.  3. Hydrogen Bonding: This is the strongest IMF in pure substances. It happens only when hydrogen is directly bonded to fluorine (F), oxygen (O), or nitrogen (N). 4. Ion-Dipole Interactions: These are even stronger than hydrogen bonds, but they only occur in mixtures where an ionic compound (like salt) interacts with polar molecules (like water). 3.2 Properties of Solids There are two main types of solids: amorphous solids and crystalline solids. 1. Amorphous solids don’t have an organized structure. Their particles are arranged in a random, disordered way. 2. Crystalline solids have their particles arranged in a neat, repeating pattern, like a grid. 3.3 Solids, Liquids, and Gases All solids, no matter the type, keep their own shape and volume. Unlike liquids or gases, solids don’t expand to fill their container because their particles are packed very tightly together and can’t move. The forces holding the particles together are strong enough to keep them in place. Liquids are different because their particles aren’t packed as tightly. This allows them to flow past each other, which is called fluidity. The forces between the particles are strong enough to keep them close but not strong enough to lock them in place. Liquids also: Minimize their surface area (like when water forms droplets). Can move upward in small spaces through capillary action (like water traveling up a plant stem). Have different levels of thickness, which is called viscosity (like syrup being thicker than water). Gases assume the volume and shape of their container. Gas particles move rapidly in straight lines; more about the behavior of gases is covered in the rest of this unit! The molecules have enough energy to overcome any intermolecular forces that exist, allowing them to move freely. They are compressible, flow readily, and expand to fill the container. 3.4 Ideal Gas Law Gases can move freely, so chemists have studied their behavior to understand them better. In this course, you’ll learn about several gas laws that describe how ideal gases behave. Ideal gases follow specific rules explained by the ideal gas law and the Kinetic Molecular Theory. One of the most important equations you’ll need to know is the ideal gas law: PV = nRT, where: P is pressure (in atm), V is volume (in liters), n is the amount of gas (in moles), R is the universal gas constant, T is temperature (in Kelvin). Another important concept is Dalton’s law of partial pressures, which explains that the total pressure of a gas mixture is the sum of the pressures of each individual gas in the mixture. This is especially useful because most gases naturally exist as mixtures—for example, the air around us! 3.5 Kinetic Molecular Theory The Kinetic Molecular Theory (KMT) explains the behavior of ideal gases and it has five key assumptions: 1. There are no interactions between gas particles. 2. Ideal gases are negligible or have no volume. 3. Ideal gas particles move in random, constant, straight-line motion. 4. Collisions between ideal gas particles are elastic. 5. When observing particles, their kinetic energy is directly related to their velocity. All gases have the same average kinetic energy at a given temperature. Maxwell-Boltzmann distributions display the energy at given temperatures for a gas and are based on the fifth concept of the KMT. They generally show that as temperature increases, the range of velocities becomes larger as particles move at a higher speed. 3.6 Deviation from Ideal Gas Law In the real world, gases don’t always behave as defined by the kinetic molecular theory. This is because gas particles can become attracted to each other and they make up a significant portion of a gas sample's volume. Conditions of low temperatures and high pressures can cause gases to deviate from ideal gas behavior in these ways. The Van der Waals equation makes corrections to the pressure and volume terms of the ideal gas law to represent real gases. It adds to pressure since the pressure is lower in real gases, and it subtracts from volume since volume is higher in real gases. 3.7 Solutions and Mixtures When you see (aq) in a chemical reaction, it means the substance is aqueous, or dissolved in water. A solution is a mixture where one substance (the solute) is evenly distributed in another (the solvent). The solute is what gets dissolved (like sugar in water). The solvent is what dissolves (like the water). Concentration tells us how much solute is dissolved in a certain amount of solvent. One way to measure concentration is molarity (M), which is the number of moles of solute in one liter of solution. In labs, chemists often need to make solutions less concentrated, or dilute them. To do this, they can: 1. Lower the amount of solute in the solution. 2. Add more solvent to increase the total volume. To figure out how much solvent or solute to use, we use the formula M₁V₁ = M₂V₂, where: M₁ and V₁ are the molarity and volume of the starting solution. M₂ and V₂ are the molarity and volume of the diluted solution. 3.8 Representations of Solutions We can use particle or molecule diagrams to show what happens in a solution. These diagrams display how the solvent and solute interact. You might see questions about this on the AP exam to test your understanding of what solutions are made of. Another important topic is electrolytes. Electrolytes are substances that can generate electricity when dissolved in water. This happens because they produce ions (charged particles) that can move and carry an electric current. 3.9 Separation of Solutions and Mixtures Chromatography When working with solutions, you have a solute (or more than one solute) dissolved in a solvent. After a chemical reaction, chemists often need to separate the solutes from the solvent. This can be done using different methods based on physical properties and the strength of intermolecular forces. Here are some common methods of separation: 1. Evaporation: The solvent is boiled and evaporates, leaving the solute behind. 2. Filtration: This removes insoluble substances from a solution by trapping them in a filter, while the liquid (the filtrate) passes through. 3. Chromatography: This method separates substances based on how they interact with a stationary phase (like paper or a solid surface). The AP Chemistry exam usually focuses on paper chromatography and thin-layer chromatography, but column chromatography is another method. 4. Distillation: This separates liquids in a solution by their boiling points. The liquid with the lower boiling point evaporates first and is collected. Each technique is chosen based on the properties of the substances being separated! 3.10 Solubility Solubility is how well a substance can dissolve in a solvent to form a uniform mixture, called a solution. If a substance is soluble, it will dissolve completely in the solvent. If a substance is insoluble, it won’t dissolve and will stay separate. Every solution has a saturation point, which is the limit where no more solute can dissolve in the solvent. 3.11 Spectroscopy and the Electromagnetic Spectrum Light is an interesting quantum idea because of the fact that it acts both as a particle (the photon) and as a wave. This is called particle-wave duality. When thinking of a wave, it is useful to visualize it as a sine wave, oscillating back and forth periodically. The amplitude of a wave refers to its vertical height and determines the light's brightness. Wavelength refers to the length of one period of the wave and it determines the color of the light we see. Frequency describes the number of waves that pass a fixed place in a given amount of time and it is inversely proportional to wavelength. Visible light, which is the light we can see, is only a small category of light. The electromagnetic spectrum includes all wavelengths of electromagnetic radiation, ranging from very short gamma rays to very long radio waves. A key trend to note is that electromagnetic radiation can be characterized by its wavelength; the shorter the wavelength, the higher the frequency. 3.12 Photoelectric Effect Albert Einstein discovered the photoelectric effect, which supported Max Planck’s theory. He proposed that light is made of packets of energy, called photons, and the energy of these photons depends on the frequency of the light. Einstein’s experiments showed that when light shines on certain metals, electrons are ejected from the surface if the light’s frequency is above a certain threshold. This means: If the light’s frequency is high enough (above the threshold), electrons are released. If the light’s frequency is too low, no electrons are ejected, and the light is just absorbed by the metal. This discovery was key to understanding how energy and light interact with matter! 3.13 Beer-Lambert Law Spectrophotometry is a method used to measure how much of a specific substance is in a sample by analyzing how much light it absorbs. A device called a spectrophotometer is used to measure the amount of light absorbed at a certain wavelength. The Beer-Lambert Law explains the relationship between how much light is absorbed and the concentration of the substance absorbing it. The equation is: A = εbc, where: A is the absorbance (how much light is absorbed). ε is the molar absorptivity (a constant that depends on the substance and wavelength). b is the path length (the distance light travels through the sample, usually in cm). c is the concentration of the substance (in moles per liter, or M). This law helps chemists determine the concentration of a solution by measuring its absorbance. Here’s a cheat sheet for Unit 3: Intermolecular Forces & Properties to help Intermolecular Forces (IMFs) London Dispersion Forces (LDFs): ○ Weakest IMF. ○ Present in all molecules, especially nonpolar ones. ○ Example: Noble gases (He, Ne) and nonpolar molecules like CO₂. ○ Hint: Think of LDFs as "temporary" forces that only happen when molecules come close. Dipole-Dipole Interactions: ○ Occur between polar molecules. ○ Example: HCl, H₂O (water). ○ Hint: Polar molecules have "permanent" charges on opposite ends, like a magnet. Hydrogen Bonding: ○ Strong IMF, only happens when hydrogen is bonded to fluorine (F), oxygen (O), or nitrogen (N). ○ Example: Water (H₂O), HF (hydrofluoric acid). ○ Hint: "Hydrogen + F, O, or N" = "H-bonding." Ion-Dipole Interactions: ○ Stronger than hydrogen bonding. ○ Occurs between an ionic compound and polar molecules. ○ Example: NaCl dissolved in water. ○ Hint: Think "ion" (charged) + "polar molecule" = ion-dipole. Solids & Liquids Amorphous Solids: ○ No regular pattern, disordered structure. ○ Example: Glass, rubber. ○ Hint: "Amorphous" = "a mess" (random structure). Crystalline Solids: ○ Regular, repeating pattern of particles. ○ Example: Salt (NaCl), diamonds. ○ Hint: "Crystalline" = "clear, ordered" structure. Saturation Point: ○ The point where no more solute can dissolve in a solvent. ○ Example: Sugar dissolving in tea – if you add too much, it won't dissolve. ○ Hint: "Saturation" = "saturated" = "full." Gas Laws & Solutions Ideal Gas Law (PV = nRT): ○ A key equation for gas behavior. ○ P = Pressure, V = Volume, n = Moles of gas, R = Gas constant, T = Temperature. ○ Hint: "PVTnR" – remember "PV = nRT" by saying "Pressure and volume, moles and temperature!" Dalton’s Law of Partial Pressures: ○ Total pressure of a gas mixture = sum of the partial pressures of each gas. ○ Example: Air is a mixture of gases (oxygen, nitrogen, etc.). ○ Hint: Think of it like adding up the pressure from each gas in the air. Separation Techniques Evaporation: ○ Solvent is boiled off, leaving the solute behind. ○ Example: Salt water evaporating to leave salt behind. ○ Hint: "Evaporation" = "heat" to "separate." Filtration: ○ Separates insoluble solids from liquids. ○ Example: Coffee filter separating coffee grounds. ○ Hint: Think of a "filter" catching solids while letting liquids through. Chromatography: ○ Separates substances based on how they interact with a surface. ○ Example: Ink separating on paper in paper chromatography. ○ Hint: "Chromatography" = "colors" (think ink) separating on paper. Distillation: ○ Separates liquids based on boiling points. ○ Example: Separating alcohol from water using distillation. ○ Hint: "Distillation" = "boiling" different substances. Solubility & Concentration Solubility: ○ How well a substance dissolves in a solvent. ○ Example: Sugar dissolves in tea but sand does not dissolve. ○ Hint: "Solubility" = "how well something dissolves."  Concentration (Molarity, M): ○ Molarity (M) = moles of solute / liters of solution. ○ Hint: "Concentration" = "how much stuff is in a solution." Dilution (M₁V₁ = M₂V₂): ○ Used to dilute a solution to a lower concentration. ○ Hint: "M₁V₁ = M₂V₂" = "Mixing" for desired concentration! Photoelectric Effect (Einstein’s Discovery) Photoelectric Effect: ○ Electrons are ejected from metals when exposed to light with a frequency above a certain threshold. ○ Example: Solar panels use this principle to convert light to energy. ○ Hint: "Light above a certain frequency" = "electrons jumping off metal."