Chem Module 11 Study Guide

Chemistry Notes - Module 11.1

Chapter 11 - States of Matter

11.1 - Gases

Around the year 1860, a model was developed to explain the different properties of gases. This model was called the Kinetic Molecular Theory (KMT).

The kinetic molecular theory describes the behavior of matter in terms of particles in motion. The model makes several assumptions about the size, motion, and energy of gas particles:

Assumption 1: Particle Size The size of gas particles is small relative to the volume of the empty space. Therefore, we ignore the volume of the particles themselves. Because of the amount of empty space between the particles, we can assume there are no significant attractive or repulsive forces between the particles.

Assumption 2: Particle Motion Gas particles are in constant random motion and their collisions are elastic, meaning no kinetic energy is lost.

Assumption 3: Particle Energy Since the gas particles have mass and are in motion, they therefore have kinetic energy, given by the equation: KE = 1/2 mv²Where KE is kinetic energy, m is mass, and v is velocity (speed). Temperature is a measure of the average kinetic energy of the particles in a sample of matter.

Explaining the Behavior of Gases

• Low Density: Recall that Density = mass/volume. Since gases are mostly empty space, their volumes are large, making their densities small compared to solids and liquids.

• Compression and Expansion: Since gases are mostly empty space, this space can be easily compressed or expanded, unlike solids or liquids where the particles are next to each other.

• Diffusion and Effusion:

• Diffusion is the movement of one material through another, or the mixing of materials due only to the particle motions of the substance.

• Effusion is the diffusion of a substance through a small opening, like a leak in a bike tire. Note: You do not need to know effusion for the test.

Gas Pressure

• Pressure is defined as force per unit area: Pressure = Force/Area

• Vacuum: A space devoid of matter.

• Air Pressure: The Earth is surrounded by miles of atmosphere, and those atmospheric particles have mass, therefore, they exert pressure on us here on Earth.

• Measuring Air Pressure: A barometer is an instrument used to measure atmospheric pressure. A simple barometer can be made by inverting a closed-end tube filled with a liquid upside-down into a pool of that same liquid. The weight of the column of liquid will push down,n but will be held up by the pressure of the atmosphere pushing on the pool of liquid. The difference between the two liquid levels, measured in distance, gives the atmospheric pressure (e.g., 760 mm of mercury).

• A manometer is an instrument used to measure gas pressure in a closed container. A simple manometer can be made with a U-shaped tube, partially filled with liquid, attached to a closed container of gas. Upon opening the closed container, the gas pressure pushes the liquid down on one side of the tube and up on the other side. As with the barometer, the pressure is measured in the distance between the two liquid levels.

Units of Pressure

The SI unit of pressure is the pascal (Pa), defined as one newton (N) per square meter: 1 Pa = 1 N/m². There are various other pressure units, and conversions are important:

• One Atmosphere (1 atm) = 760 mmHg = 760 torr = 101.3 kPa = 14.7 pounds per square inch (psi) = 1.01 bar.

Dalton's Law of Partial Pressures states that the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture: Total Pressure = P1 + P2 + P3 + ... Using Dalton's Law, partial pressures can be used to determine the amount of gas produced by a reaction. A gas can be collected in an inverted container of water, measuring its amount after subtracting the partial pressure of the water vapor.

Chemistry Notes - Module 11.2

Chapter 11 - States of Matter

11.2 - Forces of Attraction

Intermolecular Forces are forces between molecules or atoms. The prefix "Inter-" literally means between. We will learn three types: (a) Dispersion forces, (b) Dipole-dipole forces, (c) Hydrogen bonds (which are a type of dipole-dipole).

Intramolecular forces are forces that occur within a molecule, namely, chemical bonds such as ionic, covalent, and metallic bonds. These bonds are much stronger than intermolecular forces.

Table 1 - Intramolecular forces (also known as chemical bonds)

Table 2 - Three Types of Intermolecular Forces:

Dispersion Forces are weak forces resulting from temporary shifts in the density of electrons in electron clouds. They are very weak for small particles but become stronger as the number of electrons increases.

Dipole-Dipole Forces occur between oppositely charged regions of polar molecules.

Hydrogen bonds are a special type of dipole-dipole attraction that occurs in molecules with a hydrogen atom bonded to nitrogen, oxygen, or fluorine. Examples include H₂O, NH₃, HF, and alcohols (due to the -OH functional group). Hydrogen bonds explain why water is a liquid at room temperature, unlike other molecules of comparable mass.

Chemistry Notes - Module 11.3

Chapter 11 - States of Matter

11.3 - Liquids and Solids

When applying kinetic molecular theory to solids and liquids, you must consider the forces of attraction between particles as well as their energy of motion.

Forces of attraction between particles in the liquid limit their range of motion, so the particles remain closely packed in a fixed volume.

Density and Compression: Liquids are much denser than gases, with higher density due to intermolecular forces holding particles together. Like solids, liquids are considered incompressible.

Fluidity:Gases and liquids are classified as fluids because they can flow and diffuse. Liquids are less fluid than gases because intermolecular attractions are greater in liquids.

Viscosity:Viscosity is a measure of the resistance of a liquid to flow. It is determined by the type of intermolecular forces in the liquid, the size and shape of particles, and temperature.

• The stronger the intermolecular attractive forces, the higher the viscosity of the liquid.

• The greater the mass of molecules, the greater the viscosity.

• Molecules with long chains of atoms (like cooking oils and motor oil) have higher viscosity than shorter, compact molecules.

• Viscosity decreases with temperature.

Surface Tension:Surface tension measures the inward pull by particles in the interior of a liquid. Water has a high surface tension because its molecules can form multiple hydrogen bonds. Compounds that lower surface tension (called surfactants) are used in detergents and soaps to dissolve dirt in water.

Cohesion and Adhesion:

• Cohesion is the force of attraction between identical molecules.

• Adhesion describes the force of attraction between different molecules. Water in a graduated cylinder forms a meniscus due to these forces.

Capillary Action:Narrow tubes, called capillary tubes, allow water to climb up due to adhesion.

Solids

Solids occur when attractive forces (intermolecular forces) are stronger than the kinetic energy of the particles, holding them in a fixed position.

Density of Solids: Generally, solids are more densely packed than their liquid forms, causing them to sink in their own liquid. An exception is water, which freezes in a slightly larger volume than its liquid state, causing solid ice to float in liquid water.

Crystalline Solids:A crystalline solid has atoms, molecules, or ions arranged in an orderly, geometric structure. The Unit Cell is the smallest arrangement of atoms in a crystal lattice.

Covalent Network Solid is a crystalline solid whose structure is held together by covalent bonds instead of intermolecular forces. Examples include carbon and silicon.

An allotrope is an element that has more than one form in the same state; for example, solid carbon can exist as diamond or graphite, two different types of pure carbon.

An amorphous solid has particles not arranged in a regular repeating pattern. Examples include glass, rubber, and many plastics.

Chemistry Notes - Module 11.4

Chapter 11 - States of Matter

11.4 - Phase Changes

• Deposition: Exothermic; releases energy

  Sublimation: Endothermic; absorbs energy (gas to solid)

  Boiling: Endothermic; absorbs energy

  Vaporization: Endothermic; absorbs energy

  Evaporation: Endothermic; absorbs energy

  Condensation: Exothermic; releases energy

  Freezing: Exothermic; releases energy

  Melting: Endothermic; absorbs energy

  Solid: Not applicable to phase change direction

In the above diagram, you should notice that boiling, melting, and sublimation are endothermic processes, while condensation, freezing, and deposition are exothermic processes.

The Melting Point of a crystalline solid is the temperature at which the forces holding the crystal lattice together are broken and it becomes a liquid.

Even if the average kinetic energy (or temperature) is less than the boiling point, some portion of the molecule population may have sufficient energy for vaporization. This is how the evaporation of a puddle occurs even if the temperature never reaches boiling point.

Vaporization is the process by which a liquid changes to gas or vapor. When vaporization occurs only at the surface of the liquid, it is termed evaporation.In a partially filled container, water vapor collects above liquid water and exerts pressure on the surface, known as vapor pressure.

The boiling point is the temperature at which vapor pressure of a liquid equals external or atmospheric pressure, causing it to boil.

Sublimation is the process by which a solid changes directly to a gas without first becoming liquid.

The freezing point is the temperature at which a liquid is converted into a crystalline solid.

Condensation is the process through which a gas or vapor becomes a liquid, while deposition is the process by which a substance changes from a gas or vapor to a solid without first becoming liquid.

Phase Diagrams

A phase diagram is a graph of pressure versus temperature that shows in which phase a substance exists under different conditions.

In the phase diagram for water:

• The vertical curved line represents conditions for melting.

• The normal freezing or melting point is labeled where this line crosses 1 atm of pressure.

• The upward curving line from the triple point indicates boiling.

• The normal boiling point where this line crosses 1 atm is also indicated.

• The short segment between the "solid" and "vapor" areas shows conditions for sublimation.

The triple point is where three phases of a substance coexist. The critical point indicates temperature and pressure above which a substance cannot exist as a liquid. Above this point, gas and liquid become indistinguishable, forming a supercritical fluid.

potential question: higher boiling point is because of more electrons which cause stronger dispersion forces

polar means dipole-dipole, non-polar means dispersion