Physical Behavior of Matter

Chemistry on the Regents is organized into ten major units. Unit 5 is the one your teacher probably calls the "big one" because it pulls in ideas from atomic structure, the periodic table, and bonding, and turns them into rules about how mixtures, solutions, gases, and phase changes actually behave.

By the end of this page you will be able to classify any sample of matter, separate mixtures the right way, calculate concentrations, work the combined gas law, read solubility and vapor-pressure curves, walk through a heating curve and explain what energy is doing in every segment, and identify hydrogen bonding from a molecule's formula.

Every sample of matter is either a pure substance or a mixture. A pure substance has only one kind of particle. A mixture has two or more, just blended together.

A pure substance has one set of fixed properties (boiling point, density, melting point). A mixture has a range, because the proportions can change. That is the reason "salt water" boils at slightly different temperatures depending on how salty it is.

Regents loves particle diagrams. The rule is simple. Different colored circles mean different elements. Two circles bonded together mean a compound. A bunch of unconnected single circles all the same color mean an element. Toggle the diagram below to see all four cases plus the three phases.

Mixtures can be pulled apart by physical means. You do not need a chemical reaction. Pick the method that exploits whatever property is different between the parts.

A pure substance cannot be separated by physical means. If you can boil, filter, or chromatograph it apart, it was a mixture.

A solution is a homogeneous mixture. The component you have less of (and that gets dissolved) is the solute. The component you have more of (and that does the dissolving) is the solvent. In salt water, salt is the solute and water is the solvent.

Table G on the Reference Tables shows how many grams of a solute will dissolve in 100 g of water at a given temperature. The line is the saturation point. Below the line, unsaturated. On the line, saturated. Above the line, supersaturated.

Most solid solutes get more soluble as temperature goes up. Most gas solutes get less soluble as temperature goes up. That is why a warm soda goes flat faster than a cold one.

"Concentration" is just how much solute is in a given amount of solution. Two ways the Regents asks for it:

Both formulas live on Table T. Use molarity for concentrated solutions (lab work, titrations). Use ppm for very dilute solutions (drinking-water contaminants, trace pollutants).

When you dissolve a solute in a solvent, two things happen to that solvent's phase change temperatures:

Real-world examples. Salting roads in winter lowers the freezing point of water so ice melts. Adding salt to pasta water raises the boiling point a tiny bit (and seasons the pasta). The bigger the concentration, the bigger the shift.

Try each one before opening the answer.

Energy is the ability to do work or cause change. It comes in several forms. The Regents recognizes six.

These three get mixed up constantly. They are not the same.

Heat always flows from hot to cold. Always. It does not flow uphill on its own. When two objects touch, faster particles bang into slower particles, transferring kinetic energy until the temperatures match. That final state is called thermal equilibrium.

Both scales measure temperature. Kelvin starts at absolute zero (the lowest possible temperature, where all particle motion stops). Celsius starts at the freezing point of water. The conversion is on Table T.

Examples. Water freezes at $0^\circ\text{C}$ or 273 K. Water boils at $100^\circ\text{C}$ or 373 K. Room temperature is about $20^\circ\text{C}$ or 293 K.

A reaction or phase change is described by which way energy flows.

A way to remember it. Exo means out, like an exit. Heat exits the system. Endo means in. Heat enters the system.

Three formulas live on Table T. Which one you use depends on whether the substance is changing temperature or changing phase.

Notice the pattern. If the substance is changing temperature, you use $c$ and $\Delta T$. If the substance is changing phase, the temperature stays constant, so you use $H_f$ or $H_v$ instead. Use the right one or your numbers will be way off.

Real gases are messy. To make the math work, chemists invented an idealized model called the ideal gas. An ideal gas particle has three properties:

No real gas obeys these rules perfectly. But a real gas behaves most like an ideal gas at low pressure and high temperature. Why? Low pressure means particles are far apart, so their tiny volume does not matter. High temperature means particles move fast, so weak attractive forces are easy to overcome.

The opposite is also true. At high pressure (squeezed close together) and low temperature (slow moving), real gases deviate the most from ideal behavior, because attractions and particle volume start to matter. Hydrogen and helium come closest to behaving ideally because their particles are small and weakly attracted.

Pressure, volume, and temperature of a fixed amount of gas are linked by one equation on Table T.

The relationships hidden in this one equation:

• Boyle (T constant): $P_1 V_1 = P_2 V_2$. Squeeze a balloon, pressure goes up.
• Charles (P constant): $\dfrac{V_1}{T_1} = \dfrac{V_2}{T_2}$. Heat a balloon, it expands.
• Gay-Lussac (V constant): $\dfrac{P_1}{T_1} = \dfrac{P_2}{T_2}$. Heat a sealed can, pressure goes up.

All matter is made of tiny particles in constant motion. That idea is the kinetic molecular theory, or KMT. The state of a sample (solid, liquid, gas) depends on how much kinetic energy those particles carry and how strongly the IFAs pull them together.

Water is a famous exception on density. Solid water (ice) is less dense than liquid water, which is why ice floats. Hydrogen bonds force water molecules into an open, hexagonal lattice when they freeze, leaving extra empty space inside the solid.

A fourth state, plasma, exists at extremely high temperatures (stars, lightning, fluorescent bulbs). Particles have so much kinetic energy that electrons strip off and the substance becomes a soup of charged particles. The Regents will sometimes mention plasma but rarely test it in depth.

There are six named phase changes. Three add energy and lift the substance up the phase ladder. Three release energy and drop it down.

Pattern. Going up the phase ladder (solid → liquid → gas) is endothermic and needs heat in. Going down is exothermic and gives heat out.

When ice melts at $0^\circ\text{C}$, you can keep dumping heat in but the thermometer reads $0^\circ\text{C}$ until the last bit melts. That confuses students. Where does the heat go if not into a higher reading?

The heat is going into potential energy, not kinetic. Specifically, it pulls apart the IFAs holding the lattice together. Once the IFAs are broken (the substance has fully melted), heat starts being absorbed as kinetic energy again, and the temperature climbs.

For water, $H_f = 334$ J/g but $H_v = 2260$ J/g. Vaporizing takes almost seven times more energy than melting. Why?

Melting only loosens IFAs enough to let particles slide. Most attractions are still in play. Vaporizing has to break every IFA. The particles have to gain enough kinetic energy to fly apart and stop interacting. That is a much bigger energy bill.

This is why a steam burn (gas condensing on your skin) is far worse than a hot-water burn at the same temperature. Condensing steam dumps 2260 J/g of latent heat into your skin before the water itself even starts to cool.

Both turn liquid into gas. The difference:

• Evaporation happens at the surface, at any temperature. Only the fastest molecules at the top escape. The liquid cools as a result, which is why sweat cools you down.
• Boiling happens throughout the liquid, at the boiling point. Bubbles of vapor form inside the liquid. It happens when the liquid's vapor pressure equals atmospheric pressure (covered later in card 17).

A substance's melting point and freezing point are the same number. Water melts and freezes at $0^\circ\text{C}$. The difference is which direction heat is flowing. Add heat at that temperature and you melt. Remove heat at that temperature and you freeze. Both processes can happen simultaneously at equilibrium (which is exactly what is happening in a glass of ice water).

A heating curve is a graph of temperature versus time as you steadily add heat to a substance. It has five segments. The flat segments are where the trick lives.

This is the most-tested point in the whole unit. During a phase change, KE does not increase. The heat is going into PE. Read that twice. The Regents loves to trip students who say "kinetic energy goes up the whole time."

A cooling curve is the same graph in reverse. Slopes go down. The flat segments are still phase changes, but now the substance is releasing heat to its surroundings (freezing and condensation, both exothermic).

Real heating-curve problems usually ask for the total heat needed to take a substance from a starting state to a final state. The trick is to break the path into segments and use the right formula for each one. Sloped segments get $q = mc\Delta T$. Flat segments get $q = mH_f$ or $q = mH_v$.

If your start or end state stops in the middle of the curve, you only sum the segments your path actually crosses. For example, water at $25^\circ\text{C}$ to steam at $100^\circ\text{C}$ is just steps 3 and 4.

Use water values from Table T (extended): $c_{\text{ice}} = 2.10 \text{ J/g}\cdot^\circ\text{C}$, $c_{\text{water}} = 4.18$, $c_{\text{steam}} = 2.01$, $H_f = 334$ J/g, $H_v = 2260$ J/g.

If a problem hands you a heating-curve graph, you can read off both phase-change temperatures (the y-values of the two flat segments) and even rank $H_f$ vs $H_v$ qualitatively (longer flat segment = bigger heat absorbed).

A cooling curve flips this. The two flat segments now mark condensation (higher temperature) and freezing (lower temperature). Both are exothermic: the substance releases heat to its surroundings.

Inside a molecule, atoms are held together by chemical bonds (covalent or ionic). Between separate molecules, weaker forces still pull them toward each other. These are intermolecular forces of attraction, or IFAs.

IFAs are what hold a liquid together as a liquid. Stronger IFAs mean a higher boiling point, because more energy is needed to pull molecules apart into the gas phase.

A hydrogen bond is the strongest type of IFA. It happens when a molecule contains a hydrogen atom bonded directly to F, O, or N. The exposed proton on hydrogen attracts to a lone pair on a neighboring molecule's F, O, or N.

Hydrogen bonding explains why water has such a high boiling point compared to other small molecules ($\text{H}_2\text{S}$, $\text{H}_2\text{Se}$, etc. are gases at room temperature even though they are heavier). It is also why water has such high surface tension and why ice floats.

If a molecule has H but not bonded to F, O, or N (for example, $\text{CH}_4$, $\text{HCl}$), it does not form hydrogen bonds. Common Regents trap.

Vapor pressure is the pressure exerted by a liquid's vapor in equilibrium with the liquid. Higher temperature means more particles have enough energy to escape, so vapor pressure rises with temperature.

A liquid with weaker IFAs has a higher vapor pressure at any given temperature (its molecules escape more easily). It also has a lower boiling point.

Almost every interesting physical property of water comes from one fact: water is a polar covalent molecule, and a small one. That polarity, combined with water's ability to form hydrogen bonds (covered in Unit 4 and card 16), explains its specific heat, its solubility, and why we use it for cooling almost everything.

Water's specific heat is $c = 4.18 \text{ J/g}\cdot^\circ\text{C}$, one of the highest of any substance. For comparison: aluminum is 0.90, copper is 0.39, iron is 0.45. To raise a gram of water by one degree takes about 5 times as much energy as the same for aluminum.

The reason traces straight back to hydrogen bonds. When you add heat to water, only some of it goes into speeding up the molecules (kinetic energy, which is what the thermometer reads). A big chunk goes into pulling against the hydrogen-bond network, stretching and breaking H-bonds, then letting them re-form. That "wasted" energy is stored in potential energy of the bond network, not in faster motion. So the temperature rises slowly per joule added.

High specific heat makes water the best general-purpose coolant on Earth. Examples:

• Car radiators. Hot engine, cooled by water (mixed with antifreeze to extend the liquid range). Water absorbs huge amounts of heat per degree before getting too hot.
• Computer water cooling. Gaming PCs and data centers loop water through CPU/GPU blocks because air alone is not dense or heat-hungry enough.
• Nuclear and fossil-fuel power plants. All use water as the working fluid: it absorbs reactor heat, boils to steam, drives a turbine, condenses back.
• The ocean. Earth's climate is moderated by oceans. They absorb summer heat slowly and release it in winter slowly, smoothing out temperature swings on coastal areas.
• Sweating. Mammals lose heat through evaporating water on skin. Each gram that evaporates carries away 2260 J (water's $H_v$). Massive cooling per drop.

Polar water molecules can pry apart polar and ionic solutes. Each $\text{Na}^+$ ion in salt water is surrounded by water molecules with their oxygen ends ($\delta^-$) pointed at it. Each $\text{Cl}^-$ is surrounded by water with hydrogen ends ($\delta^+$) pointed at it. The water-ion attractions overcome the ion-ion attractions in the salt lattice. The salt dissolves.

A few more water-property highlights tied to its polar covalent / H-bonding nature:

• $H_v = 2260$ J/g is among the highest of common liquids. Why sweat cools. Why steam burns are so bad.
• $H_f = 334$ J/g is also unusually high. Why ice in a drink keeps the drink cold for so long.
• Ice floats. Hexagonal lattice from H-bonding leaves empty space, lowering density.
• Surface tension. H-bonds at the surface pull inward, creating a skin you can float a paper clip on.
• Universal solvent. Cell biology, blood, oceans, weather, agriculture, every chemistry lab.

These questions all hinge on what is happening at each segment of a heating or cooling curve. Use the walker in card 14 to visualize before answering.

Unit 6 is Equilibrium, the natural sequel to phase changes and saturation. Coming soon.