Rucete ✏ SAT Chemistry In a Nutshell
7. Liquids, Solids, and Phase Changes
This chapter explores how matter behaves in its liquid and solid states, and how substances change between phases. It introduces key physical properties like surface tension, viscosity, vapor pressure, hydrogen bonding, and explains phase diagrams, solubility, and calorimetry—vital ideas for understanding everyday phenomena and solving SAT Chemistry problems.
Liquids and Intermolecular Forces
- Liquids have a definite volume and take the shape of their container.
- Molecules in liquids are much closer together than in gases, so intermolecular forces (IMFs) like hydrogen bonding are very important.
- Liquids are incompressible and change volume very slightly with temperature.
Kinetics and Brownian Motion
- Molecules in liquids still move and collide, as confirmed by Brownian movement—a zigzag path seen under a microscope.
- As temperature increases, the average kinetic energy increases, and molecules move faster.
- If a surface molecule gains enough energy, it can evaporate into gas, leaving behind cooler molecules (why evaporation cools liquids).
Viscosity
- Viscosity = internal friction between liquid molecules.
- Stronger IMFs = higher viscosity (e.g., water > alcohol).
- Liquids with weak forces (e.g., gasoline) flow easily.
Surface Tension
- Caused by unequal attractions on surface molecules (inward and sideways).
- Makes the liquid surface behave like a stretched film (e.g., water beading on car paint).
Capillary Action
- Occurs when liquid is attracted to a solid surface (e.g., water climbing up paper or inside plant stems).
- Related to surface tension and used in paper chromatography.
Phase Equilibrium
- In a closed container, molecules constantly evaporate and condense.
- When the rate of evaporation = rate of condensation, the system reaches dynamic equilibrium.
- The pressure exerted by the vapor at this point is called vapor pressure.
Vapor Pressure and Boiling Point
- Boiling occurs when vapor pressure = atmospheric pressure.
- Lower external pressure → lower boiling point (e.g., water boils below 100°C at high altitudes).
- Stronger intermolecular forces = lower vapor pressure and higher boiling point.
Phase Diagrams
- Show the relationship between temperature, pressure, and phase.
-
Key points:
- Triple point: All three phases coexist in equilibrium.
- Critical point: Above this, gas and liquid phases are indistinguishable (supercritical fluid).
-
The slope of the solid-liquid line tells whether the solid is more or less dense than the liquid:
- For water: slope is negative → ice is less dense than liquid water.
Solids and Crystal Structure
- Solids have definite shape and volume, with tightly packed particles.
- Crystalline solids: regular repeating patterns (e.g., NaCl, quartz, ice).
- Amorphous solids: no definite structure (e.g., glass, plastic).
-
Crystal types:
- Ionic: strong bonds, high melting point
- Molecular: weak forces, low melting point
- Covalent network: very strong (e.g., diamond)
- Metallic: electrons move freely
Properties of Water
- High boiling point, strong hydrogen bonding
- Ice is less dense than liquid water (expands on freezing)
- Excellent solvent for polar substances and ionic compounds
Solubility and Dissolving
- Solubility = how much solute dissolves in a solvent at a given T and P.
-
General rule: “Like dissolves like”
- Polar solutes dissolve in polar solvents (e.g., salt in water)
- Nonpolar solutes dissolve in nonpolar solvents (e.g., grease in oil)
-
Temperature affects solubility:
- Solids: usually increase with T
- Gases: decrease with T
- Pressure affects gas solubility only: more pressure = more gas dissolves (Henry’s Law)
Heating and Cooling Curves
-
A heating curve shows temperature change as heat is added:
- Solid warms (T ↑)
- Melting (constant T; energy breaks IMFs)
- Liquid warms (T ↑)
- Boiling (constant T; energy breaks more IMFs)
- Gas warms (T ↑)
- Plateaus on the curve = phase changes, where added energy goes into breaking bonds, not raising temperature.
Calorimetry and Heat Equations
- Specific heat (c) = amount of energy needed to raise 1g of substance by 1°C.
- Equation:
q = mcΔT -
For phase changes:
- q = mHf for melting/freezing (heat of fusion)
- q = mHv for boiling/condensing (heat of vaporization)
- Total heat = sum of q values for each segment of heating curve.
Concentration Units
-
Molarity (M) = moles of solute / liters of solution
- M = mol / L
- Dilution:
M₁V₁ = M₂V₂ - Percent by mass or volume is also used in some contexts.
Colligative Properties
- Depend only on number of solute particles, not their identity.
-
Adding a solute to a solvent:
- Lowers vapor pressure
- Raises boiling point (boiling point elevation)
- Lowers freezing point (freezing point depression)
- Electrolytes (ionic solutes) have a greater effect because they produce more particles in solution.
In a Nutshell
This chapter explains how liquids and solids behave, how phase changes occur, and how energy affects these transitions. It covers key physical concepts like surface tension, vapor pressure, phase diagrams, and solubility. By mastering heating curves, calorimetry, and colligative properties, you gain tools to understand the behavior of matter in different states—essential for real-world and SAT Chemistry applications.