Matter is anything that takes up space and has a resting mass. Matter is ultimately made up of atoms. Either an assemblage atoms or ions of a single element or atoms of different elements bonded into molecules. From this point on we will discuss that which makes up matter as particles, whether they be atoms or molecules. The bonds between and among these particles will be referred to as intermolecular bonds. These bonds may be covalent or ionic. Sometimes the bonds are even weaker and are referred to as just intermolecular forces. The intermolecular force that holds water together as a solid or liquid is due to hydrogen bonding between adjacent molecules. The slightly positively charged hydrogen atom of a water molecule is attracted to the slightly negatively charged oxygen of an adjacent water molecule. More about this will be discussed later as we consider the unique structure of the water molecule, the fact that is polar, and how this results in some amazing properties.
Students will most readily identify any solid as matter. The definition will become more difficult to accept and apply when considering gases, liquids, and especially living organisms. Essentially, by definition the students themselves are matter. Luckily though, when teaching about matter we stick with a very concrete substance called water. Water is a great example for teaching about matter and phase changes for many reasons. Water is easily accessible of course, but consider also that the three main phases of interest, solid, liquid, and gas, are normal states for water on this planet. Students will have experience water in all three phases many times throughout their lives. There are not many substances that have three phases in such a perfect temperature range such that they exist naturally. Water is thus the ideal matter to use when discussing phase and phase change considering also the fact that water is crucial to life. This allows the teacher the opportunity to have a much larger discussion about why water is so crucial to life.
Phases of Matter
Matter may exist in five different phases as a result of the energy of the particles that make up the matter. The phases that are necessary to focus on are solid, liquid, and gas. The other two phases are plasma and Bose-Einstein condensate. Neither of these two states is necessary to discuss in a freshman level science class. The big idea here is that students gain and apply a conceptual understanding of matter and what causes phase change. Solids, liquids, and gases differ among each other by amount of energy that the particles have and the resulting impact this has on the attractive forces between particles within the matter. The amount of energy that the particles have will consequently impact their amount of movement. Simply put, as the amount of energy that any given particle has increases, the movement of that particle increases as well as its ability to overcome attractive forces from other particles.
Solid, the state of matter wherein the energy needed to overcome intermolecular forces or bonds is greater than the energy of the individual particles. Solids exist when molecules are attracted to each other and do not have sufficient energy to overcome that force of attraction. This results in the properties characteristic of a solid, the fixed shape and relatively high density. Although water is an exception, the solid state of any given matter is denser than the liquid state and gaseous state. With water the solid state is slightly less dense the liquid state, a phenomenon that will be addressed later in this unit. The particles that make up a solid are “locked” in place by the bonds among them, thus the fixed shape of a solid.
There are two types of solids, crystalline and amorphous. Crystalline solids are made up of particles that have a highly regular arrangement, that is to say that there is order to the arrangement. This order is called a crystalline lattice and it is made of smaller units called cell units. A cell unit is the smallest repeating unit of the lattice. Imagine if you had 1 cm cubes and you arranged those cubes into a 10cm cube. The 1cm cubes would be the cell units and the 10cm cube would be the lattice. The crystalline structure will be further discussed, but first amorphous solids will be considered.
Amorphous solids are solids whose particles have no order. The structure is referred as being disordered. The most common example of an amorphous solid is glass. Glass is produced by melting sand. Other elements may be added to acquire different characteristics like color or to increase the melting temperature. Lead can be added to make the glass easier to cut, for example lead crystal. Considering the disorder of the particles in an amorphous solid some argue that these solids can be classified as liquids. Though this argument is beyond the scope of this unit, it is useful to consider because of the frequency that a student will state that glass is a liquid. Lastly, a normally crystalline solid, ice, can be produced as an amorphous solid. If liquid water is frozen fast enough the particles do not have time to arrange themselves into the crystalline lattice structure and an amorphous solid will be produced.
Crystalline solids, that which water normally produces, are of more importance to this unit. Crystalline solids fall into three categories, molecular, ionic, and atomic solids. Atomic solids can be further categorized by the nature of the bonds that occur between the points of the cell units.
Molecular crystalline solids are typified by the presence of a discrete molecule at the point of each cell unit and inherently the lattice points as well. They are held together by weak intermolecular forces. If the molecules are non-polar dispersion forces hold them together. With polar molecules it is the dipole-dipole interaction that holds the crystal together. In the case of water, hydrogen bonds are the relatively weak force that allows ice to have a solid form. No matter the composition of a molecular solid the intermolecular forces are weak resulting in a low melting point. Molecular crystalline solids also lack ions and are thus poor electrical conductors.
Ionic crystalline solids are composed of ions bound together by their opposite electric charges. Ions are thus present at the lattice points. This bond is stronger than that of the molecular crystalline solids but not as strong as the covalent bonds found within atomic solids. Ionic crystals may be composed of monatomic ions, polyatomic ions, or a combination of the two. This mix of anions, negatively charged ions, and cations, positively charged ions, creates a solid that is brittle, hard, and has a high melting point. Electrical conduction is not present in the solid state but is strong when in aqueous solution or molten state. Commonly these crystals are composed of Group 1 or 2 metals combined with Group 16 or 17 nonmetals or a nonmetallic polyatomic ion. The typical class room example is sodium chloride (NaCl). The sodium is a Group 1 metal and the chloride is a Group 17 nonmetal. This crystal is formed in a 1:1 ratio of sodium cations and chloride anions. You can see the crystalline structure with even the most basic of classroom microscopes.
Atomic crystals are solids defined by atoms at their lattice points. They fall into three categories, two of which have strong, covalent, bonds that bind them together and the third which is bonded by London dispersal forces. The third type are the crystals formed by Group 8A elements, they have extremely low melting points and are thus not appropriate for the high school classroom. The two covalently bonded atomic crystals are metallic solids and network solids. Metallic solids are composed of metal cations surrounded by loose valence electrons. These electrons are not permanently associated with any given atom but are free to travel. Metallic solids are very effective electrical conductors for this reason. The bonding that occurs because of this structure is called delocalized covalent bonding because of the “free” nature of the valence electrons. An interesting property of some metallic solids is ductility, the ability to stretch into a wire and malleability, capable of being shaped or formed. Copper is a readily available classroom example. The malleable nature and electrical conductivity can be easily demonstrated by students. Network solids have directional covalent bonds between the nonmetal atoms at the lattice points. Because of the directional nature they form large molecules and are thus relatively hard but brittle. The lack of ions in these crystals results in poor electrical conductivity. A classroom example is solids formed from carbon such as diamond or graphite.
Consider modeling solids with students in the classroom by having them attempt to arrange themselves as both a crystalline solid with an organized structure and an amorphous solid with a disorganized arrangement. This will help students to form conceptual images of the two types of solids as they are forced to apply the rules that define each type of solid. With a crystalline solid they should arrange themselves so that they equal spaced from each other and in a fixed position. With an amorphous solid they should be randomly spaced and in fixed position. You can discuss the different types of bonds that will be present among them.
When matter is in a liquid state the molecules are in constant motion, are closely associated with each other, and are constantly forming and breaking bonds between molecules. In the case of water these bonds are hydrogen bonds between the slightly negative oxygen of one molecule and the slightly positive hydrogen of an adjacent molecule. This is due to water’s polar nature and these bonds are breaking and forming trillions of times per second. The constant motion of molecules in a liquid, tempered with the attractive forces between molecules results in the properties of liquids. Liquids have low compressibility due to their high density, though molecules are in constant motion they are closely associated with each other because or intermolecular forces. This close association leaves little room for empty space, thus the low compressibility and high density. Liquids also lack rigidity and will take the shape of the container that they are placed in. This low rigidity of structure is again due to the movement of the molecules that make up the matter and lack of lasting bonds between molecules. Students will have a better understanding of why a liquid acts the way the way that it does if they have a conceptual understanding of what the molecules that make up the liquid are doing.
Surface tension is another property of liquids and can be demonstrated easily in the classroom. Surface tension is the liquids resistance to increase surface area due to the intermolecular forces. This helps to explain what water will form droplets when dripped from a pipette or bead up on a waxed surface. Water will form a spherical droplet because the water molecules are attracted to other water molecules. Conceptually one can imagine that the molecules have a high affinity for each other and no affinity for the surrounding environment. This affinity will result in the shape with smallest surface area, a sphere. Consider the example of how fish, such as sardines, will school in the open ocean.
Matter in a gaseous state has no definite volume or shape because the particles that make up the gas have enough energy to overcome the attractive forces between particles. Gases have high compressibility because of the open space between particles. In fact, unless contained, a gas will in theory expand infinitely because the particles are free to move. In the classroom an easy example to help students conceptualize a gas is to light a scented candle in one corner of the room. Eventually the gaseous molecules responsible for the scent will find their way to all corners of the room.
Ultimately it is important that students can appropriately differentiate the states of matter. Differentiation is the result of the intermolecular bonds and particle energy. If a student has a strong working concept of the three states they can begin to understand why shifts between states occur as energy thresholds are met. As students develop their understanding of phase change their ability to apply this understanding increases resulting in a better understanding of the natural world.
A phase change occurs when the particles that make up matter make a transition from one phase to another. Phase change can be explained by discussing particle movement, energy, and intermolecular forces. With water these phase changes happen at 0˚C and 100˚C. Solid water, or ice, must first be heated to 0˚C to melt into a liquid if pressure is held constant at 1 atm. Pressure will be discussed as factor in phase change, but for discussion purposes here after a constant pressure of 1 atm is assumed.
Evaporation and Condensation
The transition of a molecule from the liquid state to the gas state is evaporation or vaporization, conversely condensation is the transition from a gas to a liquid. Molecules in the gas state have higher energy than molecules in the liquid state, this is a fact that students will have memorized. Their understanding of this phenomenon is elicited through the explanation of why the transition occurs. Liquids normally have a vapor pressure, a propensity for molecules within the liquid to escape as gas. This vapor pressure is affected by the temperature of the liquid. If heat is added to a liquid, the kinetic energy of the molecules increases due to the absorption of heat energy. As kinetic energy increases for a molecule, that molecules ability to overcome intermolecular forces increase. This is an endothermic process because energy is being absorbed by the liquid. Thus an increase in temperature correlates with an increase in vapor pressure for liquids. Students can understand the process of evaporation if they understand that molecules have energy and are moving, and that any given molecule in a liquid needs only to have enough energy to overcome the attraction that exists between the molecule and other molecules in the liquid. This is why it is important to work with students to create models to exemplify the states of matter and then apply that understanding to phase change. State transitions in the opposite direction of vaporization require the loss of energy and are thus exothermic. If a gas loses energy it will condense into a liquid. The molecule no longer has the necessary energy to overcome attractive forces of other molecules.
Melting and Solidification
The transition of matter from a solid to a liquid or the reverse can also due to energy. To melt matter enough heat has to be absorbed by the molecules to break free from a fixed position and move about. The molecules do not have enough energy to escape completely and will continue to form and break intermolecular bonds. Solidification occurs when sufficient energy is lost preventing the molecules to move, they are then held in a fixed position as a solid.
It is important to introduce the concept of pressure when discussing phase transitions. With the application or removal of pressure the temperature at which a phase change occurs can be manipulated. More specifically the energy threshold of a molecule to change phase can be changed. If a liquid is subjected to increased pressure the amount of energy required for molecules to change phase increases because they are combatting a force that is restricting their movement. The opposite is also true, if a liquid is exposed to conditions of reduced pressure the molecules can more readily escape the intermolecular forces and will vaporize at a lower temperature.