CHEMISTRY

M2: UNIT M READING NOTES: Heat and Phases of MatterPage 1 of 18

M-1 Reading

Energy is quite a different “kettle of fish” than the molecules and atoms which have been studied so far this year!

Where molecules and atoms are matter, Energy is not. Where molecules and atoms have volume, Energy does not. In fact, Energy is defined as the ability to do work or produce heat. There is nothing to it that we can hold in our hands—it is an action or a potential to do action.

And in fact, there are two very specific forms:

  • Potential—stored energy
  • This is the energy stored in bonds
  • Kinetic—energy of motion
  • Energy released by bonds which makes molecules move faster
  • Energy absorbed by bonds which makes molecules move slower

These energies can be found in many different types, and are always conserved in reactions. For example, chemical energy (from a coal burning power plant) can be converted into electrical energy which is used in your home. There is sometimes an amount of energy which is converted to heat (usually seen as a loss), but essentially, there is a law of conservation of energy which says that energy can be neither created nor destroyed.

In the body, chemical (potential) energy is converted to mechanical (kinetic) energy as we “burn” fuel we need to move think, breathe, etc. Additionally, there is a heat release associated with the breaking of chemical bonds heats up our body to 37ºC (98.6ºF). This allows all of our very necessary processes that are temperature dependent to continue.

In the metric system, the calorie (cal) is the unit of energy. It was defined as the amount of energy to raise one gram of water by 1ºC. Later, the SI unit of the joule (J) was developed.

While these units are very important in chemical laboratory processes, some other units are important in nutrition—surely you have heard of the Calorie (big “C”) before! This is a version of the calorie, but it is 1000 times larger—it is also called the kilocalorie. The same thing can be said for the joule, in that the kilojoule (kJ) is needed to represent nutritional data.

The following table relates some very important unit relationships.

Equivalence Statements
1 calorie = 4.184 joule / 1 joule = 0.2390 calorie
1000 calorie = 1 Calorie / 1000 joule = 1 kJ
1 Calorie = 1 kilocalorie
  • As has been previously discussed, chemical bonds store potential energy.
  • As the chemical bonds are broken (often requiring energy) or formed (often releasing energy) energy is transferred back and forth between the surrounding environment and the chemical bonds.
  • At this point it is helpful to introduce a couple of vocabulary terms specific to discussions of energy transfer.
  • System: This is the reaction that you are studying. If you were investigating an ice cube melting. The melting ice cube is your system. The system is defined at the beginning of the investigation.
  • Surroundings: This is everything besides the system you are studying. The air, table, etc. around the melting ice cube are all part of the surroundings.
  • Chemical reactions can be classified by the net energy change between the system (the reaction) and the surroundings.
  • Reactions that show a net absorption of energy are said to be endothermic.
  • Reactions that show a net release of energy are said to be exothermic.
  • Exothermic Reactions
  • These reactions release energy to the surroundings.
  • The energy released is absorbed by the molecules in the surroundings, increasing the kinetic energy of the molecules and causing an increase in temperature.
  • For example, look at the following reaction.
  • Fe2O3 + 2 Al → 2 Fe + Al2O3 + 851 kJ
  • This is an exothermic reaction.
  • Heat is RELEASED and is written as a PRODUCT in the reaction.
  • This reaction is called a thermite reaction and the energy is released all at once in avery dramatic display.
  • The change in energy of a reaction can be graphed. These graphs are called energy diagrams.
  • Below is an example of an energy diagram for an exothermic reaction. All exothermic reactions have the same shaped energy diagrams.

  • Endothermic Reactions
  • These reactions absorb energy from the surroundings.
  • Since energy is absorbed from the surroundings, the molecules in the surrounds lose kinetic energy causing a decrease in temperature.
  • The energy absorbed from the surroundings is used to break and rearrange the chemical bonds.
  • Look at the following endothermic reaction:
  • NH4SCN(s) + Ba(OH)2· 8H2O(s) + heat → Ba(SCN)2(aq) + 2NH3(aq) + 10H2O(l)
  • Notice that heat is ABSORBED and is written as a REACTANT.
  • The reaction absorbs so much heat that it causes water to freeze.
  • All endothermic reactions have similar shaped energy diagrams displaying a net absorption of energy by the system

M-2 Reading

  • You have previously discussed the heat transferred in intramolecular processes. In other words, the heat released or absorbed when chemical bonds are broken and formed.
  • There are forces that also exist between neighboring molecules.
  • These forces are not nearly as strong as the intramolecular forces (ionic and covalent bonds).
  • However, these intermolecular forces strongly influence the physical properties of substances (boiling point, freezing point, specific heat, etc.)

Intermolecular Force / Found in Molecules
that are… / Notes
London Dispersion Force
/ Nonpolar
(CH4, C2H6) /
  • The strength of dispersion forces increases with increasing molar mass.
  • Caused by the formation of an “instantaneous” dipole in neighboring molecules.

Dipole-dipole Force
/ Polar
(NF3, CO) /
  • Polar molecules have a dipole. One side of the molecule attracts more electrons and has a net negative charge.
  • The oppositely “charged” sides of neighboring molecules attract one another.
  • This does not form a bond but makes the molecules “stick” together affecting any process that requires separating neighboring molecules.

Hydrogen bonding / Polar where H is bonded to N, O, F
(H2O, NH3, HF) /
  • A special case of dipole-dipole force.
  • The strongest of the intermolecular forces.
  • Has the greatest affect on physical properties.
  • Compare the boiling points of H2O (100°C) and CH4 (-161°C). Both have approximately the same molar mass.

Changes of phase require a transfer of energy.

  • Increasing the distance between neighboring molecules (melting and boiling) requires the addition of energy.
  • Decreasing the distance between neighboring molecules (condensation and freezing) requires the release of energy.
  • Heating curves show the change in temperature that occurs in a substance as the substance changes state.
  • The focus will be on the heating curve of water. However, all heating curves have the same basic parts, these parts just occur at different temperatures.

Point on Graph / Phase(s) / Equation for heat (q) / Specific heat and enthalpy values (water)
A / Solid / m x Cp x ∆T / 2.11 J/g*C
B / Solid & Liquid / m x Hfus / 335.5 J/g
C / Liquid / m x Cp x ∆T / 4.18 J/g*C
D / Liquid & Gas / m x Hvap / 2260 J/g
E / Gas / m x Cp x ∆T / 2.00 J/g*C
  • Notice that the parts of the heating curve where a phase change is occurring are flat. These sections of the heating curve show no change in temperature. Therefore, there is no change in kinetic energy. All of the energy is going to increasing the distance (potential energy) between neighboring molecules and overcoming the intermolecular forces.
  • Notice that the plateau separating the liquid and gas is much larger than the plateau separating the solid and liquid.
  • The amount of heat that is required to undergo a phase change is additive.
  • Each part of the curve requires an equation with a different constant. The amount of substance is assumed to remain constant.
  • Look at the following example: How much heat is released when 111 grams of water vapor is cooled from 110°C to water at 100°C?
  • Find the points on the curve.
  • Find the equations and constants that you need to use.
  • Calculate the heat needed for each step in the heating curve.
  • Add the heat values together to get a total.

M-3 Reading

  • As previously discussed, an exothermic reaction has heat released (on the product side) and endothermic reactions have heat absorbed (energy is on the reactant side)
  • One can also describe this heat flow (enthalpy), ∆H, for a system as a positive or negative quantity.
  • Exothermic: ∆H is negative (-) when heat flows out of the system into the surroundings; increasing temperature
  • Endothermic: ∆H is positive (+) when heat flows into the system from the surroundings; decreasing temperature
  • The following equations represent this constant

Fe2O3 + 2 Al → 2 Fe + Al2O3 ∆H = - 851 kJ

This reaction is exothermic. Ex. Combustion of methane gas

  • NH4NO3(s) → NH4+(aq) + NO3-(aq) ∆H = +264 kJ

This reaction is endothermic. Ex. Melting of ice

  • How do we measure this heat flow?
  • Calorimeter : an apparatus measuring heat flow in a reaction containing water and/or other materials of known heat capacity. The walls are insulated so no exchange of heat with the air outside occurs
  • Water is very resistant to temperature change. Recall that the energy required for temperature change in water is 4.184 J/g∙ºC. This explains why air temperature in late spring could be 90ºF, but water temperature will remain below 60ºF.
  • The calorimeter is the surroundings we construct to measure the temperature change that results from this transferred heat.
  • The heat flow for a reaction system is equal in magnitude but opposite in sign to that for the calorimeter

∆H reaction = - ∆H calorimeter

  • So heat lost or gained from chemical bonds in a reaction can be calculated by heat gained or lost by the water in the calorimeter.
  • Magnitude of Heat Flow
  • Just a reminder the SI unit of heat is the joule (J).
  • ∆H of water can be calculated by the equation:

∆H = m * cp * ∆T

where m is the mass (g of H2O) and ∆T (of H2O) is the change in temperature. The value of c is the water’s heat transfer. This is called specific heat.

  • Specific Heat (cp) : Defined as the amount of heat required to raise the temperature of one gram 1º and is in units of J/g∙ºC. The specific heat of water is 4.184 J/g∙ºC.
  • Example: When 1.00g of ammonium nitrate, NH4NO3, dissolves in 50.0g of water in a calorimeter, the following reaction occurs:

NH4NO3(s) → NH4+(aq) + NO3-(aq)

and the temperature drops from 25.00 to 23.32ºC.

Assuming all heat is absorbed from the water,

calculate ∆H for the reaction system.

∆H = m * cp * ∆T

∆Hcalorimeter = 50.0g * 4.18 J/g∙ºC * (23.32ºC -25.00 ºC)

∆Hcalorimeter = -351 J so . . . ∆Hreaction = +351 J

  • Thermochemistry
  • The magnitude of ∆H is directly proportional to the amount of reactant or product. For example, amount of heat that must be absorbed to boil a sample of water is directly proportional to its mass.
  • For the following equation a recipe can be written:

H2(g)+ Cl2(g) →2HCl(g) ∆H= -185 kJ

1 mol H21 mol Cl2 2 mol HCl -185 kJ

71.0 g Cl2

and the following question could be asked:

Example: Calculate ∆H for the equation when 1.00g of Cl2 reacts with excess H2.

H2(g)+ Cl2(g) → 2HCl(g) ∆H= -185 kJ

x = -2.61 kJ

  • ∆H for a reaction is equal in magnitude but opposite in sign to ∆H for the reverse reaction.
  • H2O(s) →H2O(l)∆H = +6.00 kJ

H2O(l) →H2O(s)∆H = -6.00 kJ

M-4 Reading

We now turn for just a moment to Section 11-1 to understand the states of matter in terms of kinetic theory. The three states of matter that concern us are solid, liquid, and gas. If you study Figure 3 on page 379, you can see that the least ordered state is a gas. Gas particles are independent of each other and move in straight lines until they collide with each other or with the walls of their container. This means that they take the shape and volume of their container.

Liquid particles are much closer together than gas particles because their motion is more limited than gas particles due to intermolecular attractions. In fact, liquid particles seem to have a vibrational sort of motion about a sliding line of motion. Because of this, liquid particles are not arranged in a definite pattern, and will take the shape of their container. They are therefore more ordered than gases, and at normal atmospheric pressure will be much denser than gases. Because liquid particles are so limited in their motion and are so close together, they will also be relatively incompressible compared to gases, and will transmit pressure equally in all directions. The diffusion of liquid particles is generally slower than the diffusion of gas particles because liquid particles are closer together and therefore experience greater attractive forces than gas particles. Heating of a liquid will increase diffusion because the KE (and the average speed) of the particles increases.

The text defines surface tension as a force which tends to pull adjacent parts of a liquid’s surface together. This acts to make surface area of a liquid sample to be as small as possible. Because a spherical shape has the smallest possible surface area for a given volume, liquids will tend to form spherical drops (see Figure 4 on page 380). Surface tension is a result of intermolecular attractive forces—the greater the attractive force, the greater the surface tension will be. This means that substances that experience hydrogen bonding, like water, will have very strong surface tensions.

When the surface of a liquid is attracted to the surface of a solid, we say that capillary action is occurring. This attraction can actually draw some of the liquid onto the solid surface. Think of a tube—the smaller the diameter of a tube holding a liquid is, the higher that liquid will move up the inside surface of the tube because the attractive force between the solid surface is larger than the pull of gravity downwards. We can see examples of this in the chemistry lab in the observation of a meniscus when a liquid is measured in a graduated cylinder. If you did a chromatography lab in biology last year, capillary action between the paper fibers and the liquid sample helped to separate the different substances in a liquid sample.

We call the process by which a liquid or solid changes to a gas vaporization. When this escape to the gas phase from a liquid surface occurs when the substance is not boiling, we say that the substance has evaporated. How can we explain this in terms of the kinetic theory? Recall that the kinetic-molecular theory of matter states that temperature is directly proportional to the average KE of a sample. This means that some of the particles in a sample have a higher KE, and others will have a slower KE, than the average value (most of the particles, however, would have the average value). There is a minimum KE necessary for a particle to leave the liquid phase and enter the vapor phase, and if a given sample is held at a higher temperature, more of the particles in the liquid phase will have that minimum KE necessary to enter the vapor phase (see Figure 23 on page 401). If that particle is located near the surface of the liquid, it is likely to enter the vapor phase. By the way—a gas and vapor are actually the same phase. We give them different names because a gas is always in the gas phase at room temperature. A particle that is in the gas phase below its boiling point (it can be normally a solid or a liquid), is considered to be a vapor.

For all liquids, there exists a temperature (and pressure) at which the particles are moving slowly enough that they cannot overcome the attractive forces between them, and they can no longer slide past one another. The particles settle into an ordered arrangement that minimizes repulsion forces, and form a solid. This is how freezing occurs. All pure substances will have a definite temperature at which freezing or solidification occurs. We call this the freezing point for a given pressure. Substances with strong intermolecular forces (e.g., polar molecules) will freeze at relatively high temperatures, while substances with weak intermolecular forces (e.g., nonpolar molecules) will tend to freeze at substantially lower temperatures.

SOLIDS

Solid particles appear to occupy a fixed space about which they vibrate because they are closely packed, and travel only a small fraction of their diameters before they collide with a neighboring particle. Crystalline solid particles are arranged in a definite pattern, and will therefore have both a definite shape and a definite volume. The solid state is more ordered still than liquids. They will generally have the highest density and incompressibility compared to the other states of matter, (an exception to this is water, H2O, which actually has a lower solid density than liquid water near its freezing point). Solids also exhibit the lowest rates of particle diffusion. These characteristics are, of course, due to the very close packing of the particles.

We have already stated that solids are made up of particles that vibrate around a fixed position, but what dictates these fixed positions, and how do they affect the properties of solids? The truth is that some solids are more ordered than others. Most true solid substances are crystalline, with regular angles and crystal faces. A crystal is defined as a rigid body in which the particles are arranged in units which form a repeating pattern. The arrangement of these units is determined by the bonds between the particles. Therefore, the bonding in the crystal partially determines the crystal properties. The bonding can be ionic (NaCl), molecular (dry ice—solid CO2), or atomic (diamond):

(A)(B)(C)

The particles in any substance, no matter what its state, will always move faster when they are heated. When a solid is heated, the particles collide more often, and with greater force. They will therefore move farther apart. This in turn reduces any attractive forces existing between forces, and they will start to slide past one another. The ordered arrangement of the solid breaks down (entropy, the tendency of all matter to become disordered, increases), and the solid melts to form a liquid. Once again, all pure crystalline substances will have a definite temperature at which melting occurs. We call this the melting point for a given pressure. Substances with strong intermolecular forces (e.g., polar molecules) will melt at relatively high temperatures, while substances with weak intermolecular forces (e.g., nonpolar molecules) will tend to melt at substantially lower temperatures. Because glass and most plastics are not pure crystalline substances, they will tend to “flow” over a wide range of temperatures. Amorphous substances are therefore sometimes called supercooled liquids, which are substances that retain some liquid properties even at temperatures where they appear to be solid. This is due to the somewhat random arrangement of amorphous solid particles: