Matter changes states through phase transitions. These shifts between solid, liquid, and gas are driven by temperature and pressure changes. Understanding these transitions helps explain everyday phenomena like boiling water or why ice cubes melt.
play a crucial role in phase transitions. Stronger forces require more energy to break, leading to higher and boiling points. This explains why water, with its strong hydrogen bonds, boils at a higher temperature than many other substances.
Phase Transitions
Phase transitions and key temperatures
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Phase transitions change matter between solid, liquid, and gas phases
converts solid to liquid by overcoming
Key melting points at 1 atm pressure: water (0°C), ethanol (-114°C), gold (1064°C)
converts liquid to gas when equals atmospheric pressure
Boiling vaporizes throughout the liquid at the
Key boiling points at 1 atm pressure: water (100°C), ethanol (78°C), gold (2856°C)
converts solid directly to gas, skipping the liquid phase
Occurs with substances such as dry ice (solid CO2), iodine, naphthalene
converts gas to liquid by removing heat energy
converts liquid to solid at the same temperature as the melting point for a given pressure
converts gas directly to solid, the reverse of
Intermolecular forces in phase transitions
Intermolecular forces (IMFs) attract molecules to each other
Stronger IMFs lead to higher temperatures needed for melting and boiling
IMFs increase in strength from to to
London dispersion forces (LDFs) exist between all molecules due to temporary dipoles, increasing with molecular size and surface area
Dipole-dipole forces occur between polar molecules with permanent dipoles from uneven charge distribution
Hydrogen bonding, the strongest IMF, occurs when H bonds to highly electronegative N, O, or F
Substances with stronger IMFs require more energy to overcome attractions and change phase
Water has strong hydrogen bonding, resulting in high melting (0°C) and boiling (100°C) points compared to similar-sized nonpolar methane (-182°C and -161°C)
Interpretation of heating and cooling curves
Heating curves show temperature change as heat is added at a constant rate
Plateaus represent phase transitions where heat is added but temperature remains constant
Heat added during a equals the enthalpy (H) of that transition
Hfus represents the for melting/
Hvap represents the for boiling/condensing
Hsub represents the
Cooling curves show temperature change as heat is removed at a constant rate, the reverse of heating curves
Calculate heat flow (q) and enthalpy changes (ΔH) using:
q=mcΔT for temperature changes (m = mass, c = specific heat capacity, ΔT = temperature change)
q=mΔHfus or q=mΔHvap for phase transitions (m = mass, ΔH = enthalpy of fusion or )
Enthalpy change (ΔH) equals the heat flow (q) at constant pressure
The energy absorbed or released during a phase change without temperature change is called
Phase diagrams and critical points
Phase diagrams graphically represent the relationship between temperature, pressure, and physical state of a substance
Key features of a include:
: where solid, liquid, and gas phases coexist in equilibrium
: the highest temperature and pressure at which liquid and gas phases can coexist
Vapor pressure curve: shows the pressure at which a liquid and its vapor are in equilibrium at a given temperature