Ch 13 Solids & Liquids Test

Intermolecular Forces

attractions between separate molecules (not the bonds inside a molecule); much weaker than real bonds; they decide whether a substance is solid, liquid, or gas at a certain temperature

Dipole-Dipole Interactions

attractions between the slightly positive end of one polar molecule and the slightly negative end of another polar molecule; weaker than ionic bonds so these substances have lower melting points than ionic solids

Hydrogen Bonds

very strong attraction that happens when hydrogen is bonded to N, O, F, or Cl; the hydrogen becomes very positive and gets pulled toward the negative part of another molecule

Dispersion (London) Forces

temporary attractions caused by momentary uneven electron distribution; happens in ALL molecules (polar and nonpolar); explains why wax is solid and gasoline is liquid at room temperature

All intermolecular forces possible

one, two, or all three types of forces (dipole-dipole, hydrogen bonds, dispersion) can be working together to give a molecule its physical properties

Ionic Solids (table)

made of ions; held together by ionic bonds; very high melting/boiling points; dissolves in polar solvents; conducts electricity when melted or dissolved

Metallic Solids (table)

made of metal cations and free-moving electrons; held by metallic bonds; high melting/boiling points; does not dissolve easily; conducts electricity in solid or liquid state

Polar Molecular Solids (table)

made of polar molecules; held by dispersion + dipole-dipole forces (and sometimes hydrogen bonds); slightly higher melting/boiling points than nonpolar; dissolves in polar solvents; does not conduct electricity

Nonpolar Molecular Solids (table)

made of nonpolar molecules; held only by dispersion forces; low melting/boiling points; dissolves only in nonpolar solvents; does not conduct heat or electricity

Boiling Point Trend in Hydrides (Groups 15, 16, 17)

boiling points normally increase as molecules get heavier (stronger dispersion forces), BUT NH₃, H₂O, and HF have unusually high boiling points because of hydrogen bonding

Crystalline Solids

particles arranged in a very ordered, repeating 3D pattern; have a sharp melting point; example = rock salt

Amorphous Solids

particles arranged in a random, messy way; no sharp melting point, they just soften gradually; example = glass

Heat of Fusion

amount of energy needed to break the crystal lattice and melt a solid into liquid (temperature stays the same while melting)

Crystalline Melting Curve

energy first warms the solid → then heat of fusion melts it (flat line) → then energy warms the liquid

Amorphous Warming Curve

temperature rises gradually the whole time; no flat melting section because there is no regular crystal structure

Phase Changes

Melting ↔ Freezing; Evaporation/Vaporization ↔ Condensation; Sublimation ↔ Deposition

Unit Cell

the smallest repeating block of a crystal that shows the whole pattern; the full crystal is made by repeating this block over and over in 3D

Crystal Shapes

7 basic shapes plus 14 modified versions (cubic, tetragonal, rhombohedral, triclinic, monoclinic, hexagonal, orthorhombic)

Body-centered or Face-centered

ways to change a basic crystal shape by adding extra atoms in the center of the cube or on the faces of the cube

Polymorphs

same compound can form different crystal shapes depending on temperature; example: calcium carbonate = calcite (low temp) and aragonite (high temp)

Allotropes

different structural forms of the SAME element in the same state; carbon has four: graphite, diamond, fullerene C₆₀, and graphene; sulfur also has allotropes

Lattice Energy

energy released when gas particles come together to form a solid crystal; the stronger the lattice energy, the harder it is to melt the crystal

Factors Affecting Crystalline Strength (Lattice Energy)

1. bigger charges = stronger attraction; 2. smaller particles = stronger attraction; 3. ions packed closer together = stronger attraction

Types of Crystalline Solids

Atomic (frozen noble gases), Covalent Molecular (ice), Covalent Network (diamond), Ionic (NaCl), Metallic (metal ions in a sea of moving electrons)

Liquids – General Properties

molecules held together by intermolecular forces that balance their movement; they can roll and slide past each other; take the shape of their container; higher density than gases; hard to compress or expand

Cohesion

attraction between molecules of the SAME liquid (they stick to each other)

Adhesion

attraction between the liquid molecules and a different surface (like glass)

Surface Tension

unbalanced forces at the surface pull molecules inward; makes droplets round and lets some bugs walk on water

Meniscus

the curved top surface of a liquid in a tube (concave for water, convex for mercury)

Capillary Action

liquid climbs up a narrow tube when adhesion to the tube is stronger than the surface tension inside the liquid

Viscosity

how thick or “sticky” a liquid is and how much it resists flowing; stronger forces = higher viscosity; gets lower when temperature rises (cold syrup is thicker)

Surfactants

chemicals that lower surface tension (example: soap or detergent in water)

Liquids are wet only to…

substances that have similar polarity (water wets polar things but beads up on nonpolar things like oil)

Miscible

two liquids that mix completely with each other (example: alcohol + water)

Immiscible

two liquids that do not mix and stay separate (example: oil + water)

Evaporation

faster-moving molecules escape from the liquid surface into gas at any temperature; this cools the liquid because the highest-energy molecules leave

Boiling

liquid turns to gas quickly throughout the whole liquid (bubbles form inside) when its vapor pressure equals the outside pressure

Vapor Pressure

pressure created by the gas molecules pushing on the surface of their own liquid; increases as temperature increases; stronger attractions = lower vapor pressure

Normal Boiling Point

temperature where vapor pressure = 760 torr (normal air pressure at sea level)

Dynamic Equilibrium

rate of molecules leaving the liquid equals the rate of molecules returning to the liquid (closed container)

Heat of Vaporization

energy needed to turn liquid at its boiling point into gas at the same temperature (no temperature change during this)

Heat of Fusion

energy needed to turn solid at its melting point into liquid at the same temperature (no temperature change during this)

Latent Heat

hidden energy used during a phase change (melting, boiling, etc.) where temperature does NOT change

Sensible Heat

regular heat that makes the temperature go up or down (the sloped parts of a heating curve)

Heating Curve of Water

A = warm solid; B = melting (fusion); C = warm liquid; D = boiling (vaporization); E = warm gas; flat lines = latent heat

Heat of Condensation

energy released when gas turns back into liquid; same amount as heat of vaporization but opposite direction

Distillation

separating liquids by heating them to their different boiling points; the one with the lower boiling point evaporates first, then is cooled and collected

Fractional Distillation of Crude Oil

heating crude oil and collecting different products at different temperature ranges: gas (

Phase Diagram (Water)

graph that shows which phase (solid, liquid, gas) exists at any combination of temperature and pressure

Triple Point (Water)

the single point where solid, liquid, and gas can all exist together at the same time (0.01°C and very low pressure)

Normal Melting Point

temperature where solid turns to liquid at normal air pressure (0°C for water)

Normal Boiling Point

temperature where liquid turns to gas at normal air pressure (100°C for water)

Critical Temperature

highest temperature at which a gas can still be turned into liquid by applying pressure

Critical Pressure

the exact pressure needed to turn a gas into liquid at its critical temperature

Lower temperature → less pressure needed to liquefy gas

because temperature and pressure work together when turning gas into liquid