Phases of Matter: Solids, Liquids, and Gases

Unit 11: Phases of Matter

• Preparation involves completing assignments, staying on pace, and having strategies for assessments.
• Participation includes attendance, note-taking, practice, activities, and collaboration.

Overview of States of Matter

• A substance's physical state is determined by the balance between the kinetic energy of its particles and the strength of attraction between them.
• Increasing temperature increases the average kinetic energy of molecules; if sufficient, it can overcome attractive forces and cause a change in state.
• Increasing pressure causes molecules to move closer, increasing the attraction between them, which lowers kinetic energy and can induce a state change.
• The state a substance is in at room temperature is determined by the strength of intermolecular forces.

Comparison of Solids, Liquids, and Gases

• Solids: Have a condensed state of matter, are virtually incompressible, and maintain their shape and volume. Density is very high and diffusion is very low
• Liquids: Have a condensed state of matter, maintain volume, but assume the container's shape. Density is very high, diffusion occurs slowly, and they flow readily as molecules slide past each other.
• Gases: Have large spaces between particles and assume both the volume and shape of the container. Density is very low (influenced by pressure), diffusion is very rapid, and they mix homogeneously with molecules having random dynamic motion.

Interatomic vs. Intermolecular Forces

• Interatomic forces are bonds within molecules.
• Intermolecular forces are attractions between molecules.

Interatomic Forces (Bonds)

• Covalent (Molecular) Bonds: Involve shared electron pairs with overlapping electron clouds, creating a strong bond and resulting in low boiling points, low melting points, non-malleability, poor electrical conductivity, and significant vapor pressure if the molecule is under 500 g/mol.
• Ionic Bonds: Result from the transfer of electrons, leading to electrostatic attraction between cations and anions. They form "salt crystals" that are solid at room temperature with high melting and boiling points.
• Metallic Bonds: Feature atoms of the same metallic element sharing valence electrons, forming metallic ions fixed in a moving electron cloud, giving metals conductivity, strength, and luster.

Intermolecular Forces

• Intermolecular forces arise from attractions between positive and negative species and include:
  • Ion-dipole forces
  • Van der Waals forces (Dipole-dipole forces, London dispersion forces, and Hydrogen bonds)

Types of Intermolecular Forces (IMFs)

• Ion-Dipole: Occurs between an ion and the partial charge of a polar molecule; attraction increases with the charge of the ion or the magnitude of the dipole moment.
• Dipole-Dipole: Exists between polar molecules when the positive end of one attracts the negative end of another; generally weaker than ion-dipole forces.
• London Dispersion Forces (LDF): A temporary force of attraction caused by the random motion of electrons, creating an instantaneous dipole moment.
• Present in all atoms, compounds, and molecules.
• They increase with molecular weight (molar mass), providing more area for distortion and a larger instantaneous dipole moment.
• Hydrogen Bonding (H-Bond): A special attraction between a hydrogen atom in a polar bond (H-F, H-O, or H-N) and an unshared electron pair on a nearby electronegative atom; hydrogen bonds are strongest of the Van der Waals forces.

Polarizability

• It is a measure of the "squashiness" of the electron cloud of an atom, which is the ease with which the charge distribution can be distorted by an external electric field.
• Shapes that allow for more contact have stronger IMFs.

Biochemical Role of Hydrogen Bonding

• H-bonds play a vital role in biochemical systems.
• They facilitate protein folding, and are essential for the 3-D shapes to fit together.
• H-bonds create the "ladder" of DNA.
• The presence of hydrogen, fluorine, oxygen, or nitrogen (HFON) on biomolecules allows them to dissolve in the water present in organisms.

Properties of Water and Hydrogen Bonds

• H bonds are what cause the density of ice to be less than that of liquid water
• Account for the basic hexagonal shape of snowflakes.

Lake Turnover and Ecosystems

• Lake turnover refers to the seasonal movement of water in a lake
• Lake turnover has a profound effect on aquatic life.
• Mixing causes aeration of the entire water column, distributing settled nutrients.

Strength of Intermolecular Forces and Physical Properties

• Used to determine and explain a substance's physical properties.
• Stronger attraction forces indicate that substances are more likely to be solids at room temperature.
• Weaker forces indicate gases.
• Metallic/ionic/covalent bonds>>> H-bonds>Dipole>LDF.
• Higher boiling/melting point, surface tension, and viscosity indicated stronger attraction forces.
• Higher vapor pressure and greater volatility indicated weaker attraction forces.

Liquids - Viscosity and Surface Tension

• Viscosity is the resistance of a liquid to flow; stronger IMFs cause higher viscosity.
• Surface tension is the energy required to increase a liquid's surface area; water has high surface tension due to its strong IMFs.
• Cohesive forces attract similar molecules, while adhesive forces attract a substance to a surface.

Capillary Action

• Due to adhesive and cohesive forces and is important in moving water and nutrients in plants.
• The stronger the IMFs in a liquid, the higher it rises in a tube, and the narrower the tube, the higher the liquid rises.

Vapor Pressure

• This is the pressure exerted by a vapor in dynamic equilibrium with its liquid state in a closed container.
• Equilibrium occurs when evaporation and condensation rates are equal.
• High vapor pressure is associated with weak IMFs and low boiling points.
• Low vapor pressure is associated with strong IMFs and high boiling points.
• Vapor pressure increases as temperature increases

Volatility

• This is the tendency of a liquid to evaporate quickly.
• More volatile liquids have higher vapor pressure, weaker IMFs, and lower boiling points.

Boiling Point

• This is point at which a liquid's vapor pressure equals the external pressure.
• The normal boiling point is at 1 atmosphere of pressure (100 °C for water).
• High elevations have lower pressure.
• If pressure is low, the boiling point of water is lower, so the temperature of water is lower, and food takes longer to cook.
• Pressure cookers use sealed containers with water and high temperature. The boiling point and temperature inside are higher, and the food takes less time to cook.

Solids

• Solids are less chaotic (↓KE) and have organized structures that are dense and incompressible.

Classification of Solids

• Metallic solids are held together by shared valence electrons.
• Ionic solids are held together by electrostatic attraction.
• Molecular solids are held together by intermolecular forces.
• Covalent-network solids are held together by covalent bonds.

Types of Solids

• Crystalline solids: Are crystalline solids made of repeating patterns. The repeating pattern is called a crystal lattice.
• Amorphous solids: Do not have an orderly structure and tend to soften before melting.
• Allotropes: Are different forms of the same element in the same physical state.

Melting Point

• As temperature increases, kinetic energy increases, particles disrupt orderly patterns, and flow occurs.
• A substance with strong IMFs will have a high melting point.

Phase Changes

• Heat effects the phase of matter

Liquid/Gas Equilibrium and Transitions

• Evaporation/Vaporization: Requires energy for liquid to transform to gas
• Heat of Vaporization: The heat needed for the vaporization of a liquid

Solid/Liquid Equilibrium and Transitions

• Melting/Fusion: Requires energy for solid to transform to a liquid.
• Heat of Fusion: The heat needed for the melting of a solid.

Condensation and Freezing

• Are the reverse processes of evaporation and melting.

Heating Curves

• Shows the transitions of a substance (solid to liquid to gas) over time with increasing temperature and added energy.
• Changes in added heat results in the pure substance changing phases of matter under constant pressure.
• The change happens at the same temperature until all of the matter has changed phases.

Phase Diagrams

• It is a graphical way to summarize the conditions for equilibria between states of matter.
• Solid, liquid, and vapor phases are exhibited with three curves at different temperatures and pressures.
• A-B represents the vapor pressure curve of the liquid and the equilibrium between the liquid and gas phases. -The vapor-pressure curve ends at the critical point (B), which is the point where the gas or liquid phase becomes indistinguishable and forms a supercritical fluid.
• The line AC represents the variation in the vapor pressure. The sublimation point of a solid is identical to its deposition point.
• The line from A- D represents the change in melting point of the solid with increasing pressure. The melting point of a solid is identical to its freezing point.
• Point A marks the Triple point- where all three phases are at dynamic equilibrium at this temperature and pressure.