Explain kinetic molecular theory, phase changes, heating curves, vapor pressure, and intermolecular forces in solids, liquids, and gases.
A lesson of Chemistry, in Energy, States of Matter, and Solutions, about 35 minutes.
By the end of this lesson, you will be able to explain each state in terms of energy against attraction, name and rank the intermolecular forces, read a heating curve, and say why temperature stops rising during a phase change.
Energy Against Attraction
Every state of matter is one contest: how much kinetic energy the particles have,
against how strongly they attract each other.
Solid. Attraction wins. Particles are locked in fixed positions and vibrate in
place. Definite shape, definite volume.
Liquid. Roughly balanced. Particles touch but slide past each other. Definite
volume, no definite shape.
Gas. Energy wins completely. Particles are far apart and barely interact.
Neither shape nor volume is definite, and a gas can be compressed because it is
mostly empty space.
This explains why some substances are solid at room temperature and others are
gas. It is not about the size of the molecules. It is about how strongly they
attract each other compared with the energy available at $25°C$.
Plasma is the fourth state, an ionized gas, and it is the most common state in
the universe by a wide margin: stars are plasma. It just does not occur naturally
at Earth's surface temperatures.
Intermolecular Forces, Ranked
These are attractions BETWEEN molecules, not the bonds within them, and they are
what a phase change actually overcomes.
London dispersion forces. Present in everything. Electrons move, so at any
instant a molecule has a slightly uneven charge distribution, and that induces a
matching unevenness in its neighbor. Weak individually, but they grow with
molecular size, because bigger electron clouds are easier to distort.
Dipole-dipole. Between permanently polar molecules, positive end to negative
end. Stronger than dispersion at comparable size.
Hydrogen bonding. A special, much stronger dipole-dipole, occurring only when
hydrogen is bonded to N, O or F. Those three are small and highly
electronegative, so the hydrogen is left almost bare and attracts a neighboring
lone pair very strongly.
Boiling breaks intermolecular forces, not covalent bonds. Boiling water gives
steam, which is still $\mathrm{H_2O}$. Breaking the $\mathrm{O\text{-}H}$ bonds
would need about fifteen times more energy.
Size can beat type. Iodine is nonpolar and solid at room temperature, because
its enormous electron cloud makes its dispersion forces stronger than the
dipole-dipole forces in many small polar molecules.
Sort It: Strongest Force Present
$\mathrm{CO_2}$ is the one to check. Its bonds are polar and the molecule is
linear, so the dipoles cancel and only dispersion remains. Bond polarity does not
determine the intermolecular force; molecular polarity does.
Three absorb energy (melting, vaporization, sublimation), because they are
pulling particles apart against attraction.
Three release energy (freezing, condensation, deposition), because attractions
are forming.
Sublimation goes straight from solid to gas. Dry ice does it at ordinary
pressure, and so does ice in a freezer over time, which is why old ice cubes shrink
and food gets freezer burn.
Each pair is exactly reversible in energy. Melting ice absorbs $334$ J/g and
freezing water releases $334$ J/g. Same magnitude, opposite sign, which is Hess's
law applied to a physical change.
Reading a Heating Curve
Plot temperature against energy added, heating a solid steadily all the way to
gas. The line is not straight, and the flat parts are the interesting bit.
Sloped sections: temperature rises. Energy is going into kinetic energy, so
the particles move faster and the thermometer climbs. Here $q = mc\Delta T$
applies.
Flat sections: temperature does NOT rise. All the energy is going into
breaking attractions rather than into speeding particles up. Here
$q = m\Delta H_{\text{fusion}}$ or $q = m\Delta H_{\text{vaporization}}$, and
$\Delta T$ is zero.
This is the point of the whole graph. Ice at $0°C$ and water at $0°C$ are at
the same temperature and have very different energy contents, and the difference is
the energy that went into separating the molecules.
The vaporization plateau is far longer than the melting one.
Nearly seven times longer, because melting only loosens particles while boiling
separates them completely. That is also why a steam burn is so much worse than a
boiling-water burn: condensing steam dumps $2260$ J/g into your skin before the
temperature starts falling at all.
Sort It: Rising or Flat
Reading a heating curve, left to right.
Example 1: Reading the Curve
Here is every step, worked all the way through.
Heating $10$ g of ice from $-10°C$ to steam at $110°C$.
How many sloped sections are there on the curve? 3
How many flat sections? 2
During melting, what happens to the temperature? Type one word nothing
Which flat section is longer, melting or boiling? Type one word boiling
Which equation applies on a FLAT section? Type the letter q equals what m delta H
Worked Example 1: Energy for Melting
Melting $50.0$ g of ice at $0°C$. $\Delta H_{\text{fus}} = 334$ J/g.
Worked Example 2: Comparing Boiling Points
$\mathrm{H_2O}$ boils at $100°C$; $\mathrm{H_2S}$, which is heavier, boils at
$-60°C$.
Vapor Pressure and Boiling
Liquids evaporate below their boiling point, because molecular speeds vary and the
fastest molecules at the surface escape.
Vapor pressure is the pressure that escaped vapor exerts above the liquid,
and it rises with temperature as more molecules have enough energy to leave.
That is the definition, and it explains why boiling point depends on where you
are. On a mountain the air pressure is lower, so less vapor pressure is needed
and water boils cooler: around $70°C$ at the top of Everest, which is why food
cannot be cooked properly there. A pressure cooker does the reverse, raising the
external pressure so water reaches about $120°C$ before boiling, and food cooks
much faster.
Volatile substances have high vapor pressure and evaporate readily, which
means weak intermolecular forces. Gasoline you can smell; motor oil you cannot.
Evaporation cools what is left behind, because the fastest molecules leave and
the average speed of the remainder drops. That is how sweating works.
Worked Example 3: Boiling Point and Pressure
Common Mistakes
Mistake 1: Thinking boiling breaks covalent bonds
Steam is still $\mathrm{H_2O}$. Only the forces BETWEEN molecules break.
Mistake 2: Using $mc\Delta T$ on a flat section
$\Delta T$ is zero there. Use $m\Delta H$.
Mistake 3: Expecting temperature to rise during melting
It cannot until all the solid is gone.
Mistake 4: Assuming heavier means higher boiling point
$\mathrm{H_2O}$ beats $\mathrm{H_2S}$ by $160$ degrees.
Mistake 5: Reading bond polarity as molecular polarity
$\mathrm{CO_2}$ has polar bonds and only dispersion forces.
Where This Goes Next
Water is anomalous in almost every way, and every anomaly comes from hydrogen
bonding plus the bent shape.
Solid water floats, which almost nothing else does. Hydrogen bonds lock ice
into an open hexagonal lattice that is less dense than the liquid, so lakes freeze
from the top down. If ice sank, bodies of water would freeze solid from the bottom
and nothing in them would survive a winter.
Its specific heat is unusually high, so oceans buffer the planet's temperature
and organisms made largely of water resist sudden thermal change.
Its surface tension is unusually high, which is why insects walk on ponds and
why water climbs to the tops of tall trees against gravity.
It dissolves an extraordinary range of substances, which is what makes it the
medium for essentially all biochemistry.
Every one of those follows from two lone pairs bending the molecule and leaving
the hydrogens exposed. A linear water molecule would be nonpolar, would not
hydrogen bond, and would have none of these properties.
Beyond ordinary conditions, phase diagrams map out where each state exists at any
temperature and pressure. Two features are worth knowing: the triple point,
where all three states coexist, and the critical point, beyond which liquid
and gas become indistinguishable. Above it a supercritical fluid flows like a
gas and dissolves like a liquid, and supercritical $\mathrm{CO_2}$ is used to
decaffeinate coffee without any chemical solvent at all.
Check Your Understanding
Five of them, no instructions this time.
Summary
Every state is energy against attraction, and heating shifts the balance.
Dispersion, then dipole-dipole, then hydrogen bonding, and hydrogen bonding
needs H on N, O or F. Size can beat type.
Boiling breaks intermolecular forces, not covalent bonds.
Flat sections of a heating curve use $m\Delta H$; sloped sections use
$mc\Delta T$. Boiling's plateau is far longer than melting's.
Boiling happens when vapor pressure equals external pressure, which is why
altitude and pressure cookers change it.