Skip to main content
Make the most out of any wiki by using our free sister product,
Blendspace by TES
, to create interactive lessons and presentations!
Pages and Files
Chapter 1 The Study of Change
Chapter 2 Atoms, Molecules, Ions
Chapter 3 Mass Relationships
Chapter 4 Aqueous Reactions
Chapter 5 Gases
Chapter 6 Thermochem
Chapter 7 Electronic Structure
Chapter 8 Periodic Relationships
Chapter 9 Chemical Bonding I
Chapter 10 Chemical Bonding II
Chapter 11 I.M.F. s
Chapter 12 Solutions
Chapter 13 Kinetics
Chapter 14 Equilibrium
Chapter 15 Acids and Bases
Chapter 16 Acid Base Equilibria
Chapter 17 Atmosphere
Chapter 18 Entropy, Free Energy
Chapter 19 Electrochem
Chapter 20 Metallurgy
Chapter 21 Nonmetals
Chapter 22 Transition Metals
Chapter 23 Nuclear Chemistry
Chapter 24 Organic Chemistry
Chapter 25 Polymers
Chapter 11 Intermolecular Forces (Liquids and Solids)
11.1. The Kinetic Molecular Theory and Liquids and Solids
cited from media-2.web.britannica.com
11.2. Intermolecular Forces
11.3. Properties of Liquids
11.4. Crystal Structure
11.5 X-Ray Diffraction by Crystals
11.6 Types of Crystals
11.7. Amorphous Solids
11.8 .Phase Changes
11.9. Phase Changes
This chapter will examine the fundamental properties of Liquids and Solids. Amongst many other topics, this chapter will introduce the reader to many basic principles of matter, particularly liquids and solids. It will define characteristic properties of all three states of matter (11.1), intermolecular forces such as London Dispersion, Dipole-Dipole, etc. (11.2). It will go in depth on properties of liquids (11.3), and crystal structures and types (11.4 and 11.6). By the end of this study, the reader should also be familiar with the difference between normal and amorphous solids and will study in in-depth analysis of phase changes and how to interpret diagrams of these changes (11.8 and 11.9). Now, before we get started, let's get a few things straight. I'm going to be completely honest. This isn't Mass Relations, or the analysis of the Periodic Table. I guarantee your tenth grade chemistry class didn't hit this stuff. It's quite involved but that DOESN'T mean it has to be hard. In fact, it's quite manageable if you focus. I guarantee you that if you can read and follow this text; you will do GREAT on the test. But this isn't a chapter you can pretend to understand. So don't try it. Chapter 11 can be your best friend...or your worst enemy.
11.1 The Kinetic Molecular Theory of Liquids and Solids
A Brief Overview of Gases
In gases, the distance between molecules (in comparison with their diameters) is extremely great. At normal pressures and temperatures, there is no appreciable interaction between the molecules. Since most of the space in gases is empty (not occupied by molecules), gases can be easily compressed. Also, because gases lack strong forces between molecules, they easily expand to fill the volume of the container they are occupying. Due to the amount of empty space in gases, they have very low densities under normal conditions.
Liquids and Solids
The main difference between the two condensed states of matter (liquids and solids) and the gaseous state is the distance between molecules. In liquids, there is very little empty space and the molecules are very close together. Due to this, they are much denser than gasses in normal conditions and also more difficult to compress. Also, they have definite volumes because molecules in liquids do not break away from their attractive forces (liquids are held together by one or more of these attractive forces). The molecules can move past one another freely, allowing the liquid to both flow and take the shape of its container.
Molecules are held together rigidly with no freedom of motion. There is less empty space in a solid than in a liquid and in many solids; molecules are arranged in regular configurations in three dimensions. Solids are almost incompressible and have both a definite shape and volume. For most substances (not including water), densities as a solid is greater than densities as a liquid. In many situations, more than one state of a substance can coexist, such as ice floating in water. The different states of a substance present in a system are known as
Want some reinforcement? Do you respond well to AWESOME, interactive animations? Click Here:
States of Matter
11.2 Intermolecular Forces
are attractive forces between molecules. They are responsible for the non-ideal behavior of gases and exert even more influence in condensed states of matter (liquids and solids). As the temperature of a gas drops, the kinetic energy of its molecules decreases until they break away from the attraction of neighboring molecules. The molecules then aggregate to form small drops of liquid. This is known as condensation.
are forces that hold atoms together in a molecule. They stabilize individual molecules while intermolecular forces are responsible for the bulk properties of matter. Intermolecular forces are much weaker than intramolecular forces. For example, it takes much less energy to evaporate a liquid than to break the bonds in the molecules of a liquid.
The boiling point often reflects the strength of intermolecular forces operating among the molecule. At the boiling point, there must be enough energy supplied to overcome the attractive forces among the molecules before the substance can enter the vapor stage. The more energy it takes to separate molecules of a substance, the stronger the intermolecular forces and the higher the boiling point of the substance must be. Basically, the stronger the intermolecular forces, the higher the boiling point.
van der Waals forces
- named after the Dutch physicist Johannes van der Waal, these are the different types of molecular forces. They are:
are not considered van der Waals forces but are intermolecular forces treated as their own categories.
Molecules with permanent dipole moments align with opposite polarities for maximum attractive interaction as shown above.
attractive forces between polar molecules, that is, between molecules that possess dipole moments. Their origin is electrostatic (which can be explained in Coulomb's law in
) and the greater the dipole moment, the greater the force. In liquids, polar molecules are not held as rigidly as in solids, but still align in a way that will maximize attractive interaction. The diagram to the right shows dipole-dipole forces in a solid. (For polar molecules and dipole moments, see
Coulomb's law can also explain ion-dipole forces. These forces attract an ion (either a cation or anion) and a polar molecule to each other. The strength of this interaction depends on the charge and size of the ion and on the magnitude of the dipole moment and size of the molecule.
On average, cations are more concentrated because they are usually smaller than anions. Therefore, if a cation and anion have the same magnitude, a cation will interact more strongly with dipoles than an anion!
Ion-Dipole Forces: The larger the charge and size of the ion and the larger the magnitude of the dipole moment and size of the molecule, the greater the strength of the interaction!
Dipole induced dipole and ion-induced dipole interactions:
If an ion or a polar molecule is placed near an atom or a nonpolar molecule, what happens? The electron distribution of the atom (or nonpolar molecule) is distorted by the force exerted from the ion or the polar molecule, resulting in a kind of dipole. The dipole in the nonpolar atom or molecule is said to be an induced dipole because the separation of the positive and negative charges in the atom (or nonpolar molecule) is due to the proximity of an ion or polar molecule. The attractive interaction between an ion and the induced dipole is an
ion-induced dipole interaction
. If the attraction is between a polar molecule and induced dipole, it's called a
dipole-induced dipole interaction
. The likelihood of a dipole moment being induced depends on: charge on the ion, strength of the dipole, and polarizability.
is the ease with which the electron distribution in an atom or molecule can be distorted. The greater the number of electrons and more diffuse the electron cloud in the atom, the greater its polarizability. A diffuse cloud is an electron cloud that is spread over an appreciable volume, so that the electrons are not held tightly by the nucleus. Polarizabiltity allows gases containing atoms or nonpolar molecules (like Helium) to condense.
Let's use Helium, a nonpolar molecule, atom as an example for this particular term. The electrons are constantly moving at some distance from the nucleus. It is very possible that, due to the configuration of the electrons, the atom has a dipole moment known as an instantaneous dipole because it lasts only for a fraction of a second. In the next second, it may have a new instantaneous dipole in a different location. However, because all these instantaneous dipoles cancel each other out the atom has no dipole moment. In a collection of atoms, like Helium, an instantaneous dipole of one Helium molecule can induce a dipole in the neighboring Helium atoms and than the next second, create a different instantaneous dipole in a different Helium atom. This type of interaction creates
, which are attractive forces that arise as a result of temporary dipoles induced in atoms or molecules. At very low temperatures, when atomic speeds are reduced, dispersion forces are strong enough to hold some atoms, like gaseous Helium, together, causing the gas to condense!
These are also known as London forces. Dispersion forces will increase with molar mass because molecules with larger molar masses more electrons.
Dispersion Forces increase in strength with the number of electrons. A larger molar mass will also imply a bigger atom with an electron distribution that is more easily disturbed since the outer electrons are held less tightly to the nucleus than smaller atoms. Also, the more electrons, the higher the melting point. Dispersion forces are comparable to and sometimes greater than dipole-dipole forces between polar molecules.
For the most part, the boiling points of similar compounds containing elements in the same periodic group will increase with molar mass, due to the increase in dispersion forces for molecules with more electrons. However, Hydrogen compounds of the elements in Groups 5A, 6A, and 7A do not follow this trend. Examples: HCl and HI. In each of these periods, the lightest compounds such as NH
O and HF have the highest boiling points! Why? This must mean that intermolecular attractions in these compounds are stronger than other molecules in the same group. This strong bond is known as a Hydrogen Bond, which is a special type of dipole-dipole interaction between hydrogen atoms in a polar bond (like N-O or O-H from the compounds above) and an electronegative O, N, or F atom. An example of how this interaction is written is shown below.
X and Y represent the elements Oxygen, Nitrogen, and/or Fluorine
The X-H represents one molecule (or part of one) and the Y is part
of another molecule. The dotted line represents Hydrogen bonds.
The three atoms usually lie in a straight line but bond angles can be
as much as thirty degrees! Important: O, N, and F atoms possess at
least one lone pair capable of interacting with a Hydrogen atom in
A hydrogen bond has a large average energy for a dipole-dipole interaction (up to 40 kJ/mol) and has a powerful effect on the structures of compounds. How can someone determine the strength of a Hydrogen bond? The strength is determined by the interaction between the hydrogen nucleus and the lone pair electrons of the electronegative atom. The higher the electronegativity, the stronger the bond! For example HF would have stronger bonds than H
0 because Fluorine is the most electronegative element.
Remember! All the forces discussed are attractive in nature. HOWEVER, molecules also exert repulsive forces on one another. Electrons of one molecule will repel electrons of another and the same thing will happen between the nuclei. The magnitude of these forces rise as the distance between molecules in liquids and solids decrease. This explains why condensed states are harder to condense FURTHER. They are already in close contact and will resist further compression.
11.3 Properties of Liquids
Surface Tension: Molecules within a liquid are pulled by intermolecular forces in all directions, with no tendency to any one direction. HOWEVER, molecules on the surface of a liquid are pulled downward and sideways by other molecules. They, unlike other molecules, are not pulled upward. Due to this, the surface of the liquid will tighten like an elastic film. Can we measure the elastic force in the surface area of a liquid? YES. When a water molecule (polar) is dropped on a nonpolar surface, like a wax tablecloth, the water will form a small bead to minimize its surface are. Surface Tension is the amount of energy required to stretch or increase the surface area of a liquid by one unit area.
Let's look at...
Liquids with strong intermolecular forces will have high surface tensions.
CAN YOU FIGURE THIS ONE OUT
- Water has a greater surface tension that most other liquids. WHY?
HINT: Last lesson we learned about intermolecular bonds...and there are particularly strong ones between Hydrogen and Nitrogen, Fluorine, and Oxygen!
: Pay attention this one is a little tricky.
To the right are capillary tubes. The blue stuff is water spontaneously rising in the tubes. How does this work? A thin film of water adheres to the wall of the glass tubes. The water's surface tension will cause this film to contract and as it does, it pulls the water up the tube. Let's break it down even more. Two types of forces will bring about capillary action. One is
. Cohesion is the intermolecular attraction between molecules. The second is
, which is an attraction between unlike molecule. Let's get back to the capillary tubes. Water on the sides of a capillary tube would be a prime example of adhesion. If adhesion is stronger than cohesion, than the water in the tube will be pulled upward. This process will continue until the weight of the water balances out the adhesive forces. However, this is not a universal trait in liquids. In some, like Mercury, cohesion is greater than adhesion, thus resulting in a depression or lowering, at the Mercury level.
: A measure of a fluid's resistance to flow. The greater the viscosity, the more slowly the liquid flows.
Let's look at...
! The viscosity of a liquid will INCREASE when the temperature DECREASES. Also, liquids with STRONGER intermolecular forces have HIGHER viscosities.
Water will have a high viscosity because of its ability to form strong hydrogen bonds. HOWEVER, Glycerol has one of the highest viscosities. WHY?
Like water, Glycerol can form hydrogen bonds. Each molecule has three hydroxide groups that can participate in hydrogen bonding with OTHER glycerol molecules. Also, because of glycerol's shape, the molecules tend to get entangled rather than slip past one another. These interactions contribute to its high viscosity.
Glycerol: A clear, odorless, syrupy liquid used to make explosives, ink, and lubricants. This COULD be an AP question!
Glycerol's shape will cause it to entangle with other glycerol molecules rather than slide past
The Structure and Properties of Water:
Water is important. Very important. In fact, it's SO important; a whole section of this lesson is dedicated to it. You think you know water...don't you. But you don't.
: All life processes involve water. Easy enough.
: Water is an excellent solvent for many ionic compounds, as well as other substances capable of forming hydrogen bonds with water.
: Water has a high specific heat. WHY? In order to raise the temperature of water (increase the average kinetic energy of water molecules), we must first break the intermolecular hydrogen bonds. Thus, water can absorb substantial amounts of energy while its temperature only slightly rises.
: The converse of the third is also true. Water can give off great amounts of heat while its temperature only slightly decreases!
! Due to the third and fourth characteristics of water, bodies such as oceans and lakes can moderate the climate by absorbing heat in the summer and giving off heat in the winter with only small changes in the temperature of the body of water.
: Water is less dense as a solid than a liquid. This is why ice floats. What makes this SOOOOO cool is that almost every other substances density is greater as a solid than a liquid. Let's analyze why. This is the electronic structure of a water molecule:
As you can see, there are two lone pairs on the oxygen atom. Many compounds can form intermolecular hydrogen bonds. However, unlike many other compounds that form these bonds, the oxygen atom can form TWO hydrogen bonds, the same as the number of its lone pairs. Therefore, water molecules are joined together in three dimensional networks where each oxygen atom is bonded to four other hydrogen atoms in a tetrahedral-esque structure. Two of the bonds are covalent (in the water molecule) and two are hydrogen bonded. Because the equality between hydrogen atoms and lone pairs is not in other molecules capable of hydrogen bonds, they can form chains and rings but not three dimensional structures.
The highly ordered three dimensional structure of ice prevents the molecules from getting too close to one another. Let's analyze ice melting for a second: At its melting point, water molecules have enough kinetic energy to pull free of the hydrogen bonds. These molecules will become trapped in the three dimensional structure which begins to breakdown into smaller clusters. As a result,
there are more molecules per unit volume in liquid water than in ice. Since density=mass/volume, water's density is greater than ice. With further heating, more molecules are released from the intermolecular hydrogen bonding which leads to an increase in density with an increase in temperature. HOWEVER, don't forget that water expands when heated so the density is slightly decreased. The trapping of free water molecules and thermal expansion act in different directions.
From zero to four degrees Celsius, the trapping of water molecules prevails and water becomes progressively denser. However, after four degrees, thermal expansion predominates and the density of water will decrease with increasing temperature. Whoah! That's a LOT of information about water you probably didn't know about! I told you your mind would be fried. You must be thirsty after all of that. Get a glass of water. It's refreshing!
11.4 Crystal Structure
Solids can be divided into two categories: crystalline and amorphous
possesses rigid and long range order; its atoms, molecules, or ions occupy specific positions.
PSYCHE. That's section 11.7 you're not ready yet.
Back to Crystalline Solids...the arrangement of particles in a crystalline solid is such that the net attractive intermolecular forces are at their maximum. Forces responsible for crystal structures can be ionic, covalent, van der Waal, hydrogen, or a combination of them! Let's DELVE into this a little bit deeper.
The basic repeating structural unit of a crystalline structure.
To the right is a unit cell. Each sphere represents an atom, ion, or molecule and is called a lattice point. However, not ev
ery crystals lattice point contains these particles but rather has a group of them arranged AROUND the lattice point. However, we will assume, for the sake of understanding, that each lattice point is occupied by an atom, which is the case with most metals. All crystalline structures can be defined by one of seven types of unit cells. Any of the unit cells, when repeated in all three dimensions, forms the lattice structure of a crystalline solid.
SEVEN TYPES OF UNIT CELLS:
simple cubic cell
, all sides are equal length and all angles are ninety degrees
, C is larger than A and all angles are again, equal to ninety degrees.
, the length, width, and height are all different lengths but all angles are ninety degrees.
, all the sides are equal and all the angles are equal. However, the angles cannot be right angles.
, length, width, and height are different lengths. both B and Y are ninety degrees. However, A is not a right angle.
, length, width, and height are all different. Also, the three angles, A, B, and Y are all unequal and not ninety degrees.
, the length of all sides A are equal and unequal to C. Angles are comprised of ninety and one hundred and twenty degrees.
Packing Spheres: To understand the geometric requirements for crystal formation, let's consider the different ways of packing a number of identical spheres to form an ordered three-dimensional stucture. The way the spheres are arranged in layer determines the type of unit cell. It SOUNDS difficult to do but TRUST me it's not as hard as you think. As with everything, let's start w
ith the simplest example, the simple cubic cell. Let's think of the diagram below as a bird's eye view of the spheres. A three dimensional structure can be generated by placing a layer above and below this layer in a way that spheres in one layer are directly above and below another layer. This can be done many times to generate many layers, as is the case of a crystal. Let's focus on the middle sphere. It is in direct contact with four spheres in its own layer and one above and below. How many is this? Let's count it up: 1 + 1 + 4 = 6. Therefore, each sphere in this arrangement will have a
of six. What the heck is that?! A coordination number can be defined as the number of atoms or ions surrounding an atom or ion in a crystal lattice. Its value gives a measure of how tightly the spheres are packed together. Let's look at
The LARGER the coordination number, the closer the spheres are to each other. A SIMPLE cubic cell is SIMPLE enough. But, with everything else in chemistry, there's always a way to make easy stuff a little more complicated. So what are we waiting for?!
Okay, pay attention because this part is a little more involved conceptually. You need to be at one hundred percent so turn the TV off and really pay attention. The other types of cubic cells are body-centered and face-centered. These are simply other types of cubic cells. Let's look at some pictures to help us out.
Simple Cubic Cell
Face Centered Cubic
How do they differ from the simple cubic cell? Well they're not as simple. DUH. A body centered cubic arrangement differs from a simple cell in that the second layer of spheres fits into the depressions of the first layer, the third layer fits into the second layer's depressions, etc. etc. The coordination number is 8 because each sphere is connected to four spheres above and four spheres below it. In a face centered cubic cell, there are spheres at the center of each of the six faces of the cube, in addition to the eight corners of the cube. Below, are some clearer representations of the placement of spheres within simple, body, and face centered cubic cells. The dots represent the location of the spheres. While this is clearer, the diagrams to the right are more accurate.
Remember, every unit cell in a crystalline solid is adjacent to other unit cells. Therefore, most of a shell's atoms are shared bit neighboring cells. For example, in all types of cubic cells, the corner atom belongs, not just to one unit cell, but is shared by eight! A face-centered atom (the one in the MIDDLE of each face of the cube) is shared by two unit cells. Let's use a simple cubic cell for this next example. It's a little tricky but not TOO hard to follow so pay attention. Let's pretend, for a second, that rather than eight unit cells sharing one corner atom, each unit cell is greedy and decides to take 1/8 of the atom (or sphere) to itself. Now, there are eight corner atoms being shared so 1/8 multiplied by 8 shared corner spheres equals a TOTAL of only one sphere. This means that, essentially, there is only one complete sphere within a simple cubic cell. GET IT?! Now, in a body centered cubic, the same arithmetic happens but in addition to the eight corner spheres, there's always one unshared sphere in the middle of each cube. SO, there's the one sphere in the middle PLUS the one sphere being made up of the 1/8 portions of corner spheres, meaning that a body centered cubic will have a total of two spheres. A face centered cubic will have a total of FOUR complete spheres can you figure out how? Like the rest, one total sphere is made up from the corner spheres. However, each of the six faces of the cube have a sphere in the middle of them being shared by ONE other cube thus, splitting it in half. 6/2 = 3. Three plus one equals four! This is very hard to grasp mentally so, here are the pictures:
Imagine this as a simple cubic cell with neighbors on all sides sharing corner spheres. As you can see, each sphere on the corners is only 1/8 of a whole sphere.
As you can see, each corner contains one eighth of a sphere IN ADDITION TO an unshared atom in the middle of the cube.
With a face centered cubic, each corner contains an eigth of an atom and each face contains half an atom!
Closest Packing: Clearly, there is more empty space in a simple cubic and body centered cubic cell than a face-centered cubic cell. Like everything in Chemistry, efficiency is the key. Closest packing is the most efficient arrangement of spheres. Let's use the image below to explain closest packing a little more detailed. In the first image in the top left, let's focus on the enclosed sphere (the one in
the middle). It has six immediate neighbors in this layer, which is the first layer. In the second layer (the picture to its immediate right) spheres are packed into the depressions between the spheres in the first layer so that all the spheres are as close together as possible! At this point, as you can see, there are two variations of the third layer that are possible. In the first scenario (the picture in the bottom left of the diagram), the spheres will fit into the depressions of the second layer so that each third layer sphere will be directly over a first layer sphere. Let's refer to the first layer as Layer A and the second layer as B. In this first scenario, because all the third layer spheres are in the exact same positions as first layer spheres, this would also be referred to as Layer A. However, the third layer spheres may fit into depressions that lie directly over depressions in the first layer, thus making a completely unique layer known as Layer C in the bottom right.
For more help on close packed structures CLICK HERE!
The figures to the direct right show hexagonal and cubic close packed structures. In a hexagonal close packed structure, every other layer of spheres occupies the same positions. For example, the spheres in the first layer are in the same position as spheres in the third layer. For this reason, it is known as an ABA arrangement. In a cubic close packed structure (which corresponds to the face centered cubic structure already described); the spheres in every fourth layer occupy the same arrangement, thus giving it an ABC arrangement. A continuation of the hexagonal close packed structure would result in an ABABABA...
Hexagonal close packed structure
Cubic Close Packed Structure
nd in a cubic close packed structure, an ABCABC...
structure would form. In both structures, each sphere has a coordination number of 12. How? Each sphere is in contact with six spheres in its own layer, three in the layer above, and three in the layer below. Both these structures represent the MOST efficient way of packing ideal spheres in a unit cell. There is no way to increase the coordination number to higher than 12. Hmmm...what kind of elements has these types of structure? I'm glad you asked that question. Many metals and noble gases, which are monatomic, form crystals with hexagonal or cubic close packed structures. For example, magnesium, titanium, and zinc crystallize with their atoms in a hexagonal close packed structure and aluminum, nickel, and silver crystallize with their atoms in a cubic close packed structure. This is good to remember, as a section of every AP test is devoted to minor facts about specific elements. ALL solid noble gases will have a cubic close packed structure save helium, which will hake a hexagonal close packed structure. REMEMBER THAT! Why would different crystal structures form in a series of related substances like noble gases? It all has to do with the relative stability of a particular crystal structure, which is governed by intermolecular forces. For example, magnesium has a hexagonal close packed structure because this arrangement of its atoms will result in the greatest stability of the solid. Now, the moment we've all been waiting for....CALCULATIONS. I know the below calculations LOOK difficult but they are really easy to understand. The below figure summarizes the relationship between side length and the atomic radius of one of the spheres in the cube. The atomic radius of a sphere can be calculated if the density of the crystal is known.
CONCLUSION: This is by far the most difficult lesson in the chapter and, hey, it wasn't that bad right? We learned all about the basics of a crystalline solid, the seven basic types of unit cells, packing spheres, closest packing, and the A to Z of simple cubic, body centered, and face centered cubic cells! WOW! If you can understand this stuff, the rest of the chapter is a PIECE OF CAKE. I suggest studying the pictures to help understand the concepts because the diagrams within this lesson are very useful and its good to know what each crystal structure LOOKS like and what close packing looks like. Trust me when I say it's all downhill from here!
11.5 X-Ray Diffraction by Crystals
Almost everything known about crystal structures has been learned from X-Ray diffraction studies. What the heck does that mean?! X-Ray diffraction refers to the scattering of X rays by the units of a crystalline solid. The scattering patterns produced are used to deduce the arrangement of particles in the solid lattice. In
, the interference phenomenon associated with waves was discussed. HISTORY LESSON! In 1912, Max von Lauer, a German Physicist, correctly suggested that, because the wavelength of X-rays is comparable in magnitude to the distances between lattice points in a crystal, the lattice should be able to diffract x-rays! In 1914, he received a Nobel Prize for his discovery of X-ray diffraction.
The figure to the right shows a typical X-Ray diffraction setup. A bean of X-Rays is directed at a mounted crystal. Atoms in the crystal absorb some of the incoming radiation and then reemit it. This process is called the scattering of X-Rays. To understand how a diffraction pattern may be generated, let's consider the scattering of X Rays by atoms into two parallel planes. Initially, the two incident rays are in phase with each other. What does that mean? It means that their maxima and minima occur in the same position. Let's look at a diagram of these rays to help us better understand.
The diagram to the right shows the reflection of X Rays. Imagine on each incident and diffracted (reflected) x-ray lines there is a sine curve starting at the origin, B. The upper wave is scattered, or reflected, by an atom in the first layer, while the lower wave is scattered by an atom in the second layer. In order for these two scattered waves to be in phase again, the extra distance traveled by the lower wave must be an integral multiple of the wavelength of the X-Ray. That means the following:
AB + BC = 2dsin
= nλ where
= 1, 2, 3...
is the angle between the X-Rays and the plane of the crystal.
= the distance between adjacent planes. Basically, the distance between B and B'. HISTORY LESSON! This equation is known as the Bragg Equation after William H. Bragg and Sir William L. Bragg, both English physicists. William worked mainly in X-Ray crystallography and shared the Nobel Prize with his son Sir William in 1915 for Physics. Sir William formulated the fundamental equation for X-Ray diffraction. Physics in Chemistry...and who said sciences didn't overlap?!?!
The X-Ray diffraction technique offers the most accurate method for determining bond lengths and bond angles in molecules in the solid state. Since X-Rays are scattered by electrons, chemists can construct an electron density contour map from the diffraction patterns by using a complex mathematical procedure. This contour map will tell us the relative densities of electrons at various locations in a molecule. The densities will reach at maximum at the center of each atom. In this manner, we can determine the positions of the nuclei and hence the geometric parameters of the molecule.
If you're like me at all you’re saying to yourself Holy Ravioli wha-wha-what?!?!?! So...this site will go much further into Bragg's equation and what every variable stands for and a more in depth history.
11.6 Types of Crystals
I know what your wondering right now. You're saying "Wow this is all so interesting! But how are properties of crystals, such as density, hardness, and melting points of crystals determined?" If you're asking yourself this, my reply is that these properties of crystals are determined by the kinds of forces that holds the particles together. There are four types of crystals: Ionic, Covalent, Molecular, and Metallic.
Ionic Crystals have two important characteristics...
1). They are composed of charged species.
2). Anions and Cations are quite different in size.
Knowing the radii of the ions is also helpful in understanding the structure and stability of these compounds. There is no way to measure the radii of an individual ion, but sometimes you can come up with a reasonable estimate. For example, we know the radius of I- in KI is 216 pm, we can determine the radius of the K+ ion in the compound and, from that, the radius of Cl- in KCl and so on and so on. Let us consider the NaCl crystal, which has a face centered cubic lattice. The figure below shows that the edge length of the unit cell of NaCl is twice the sum of the radii of Cl- and Na+. Using the chart of ionic radii in
we know the atomic radii of Na+ and Cl-. So we calculate the edge length of the unit cell to be 2(95 + 181) pm = 552 pm. But WAIT! The diagram below shows the length to add up to 564 pm. This means is that the radius of an ion will vary slightly from one compound to another. Oh Well!
Fluorites and Zincblendes and Cesium Oh My!
Let's talk examples. You know, just to help clear the air of discrepancies and make sure our understanding stretches farther than that of face centered Sodium Chloride! Let's start with CsCl. A Cs+ ion is much larger than a Na+ ion so CsCl will have a simple cubic lattice. ZnS has the zincblende structure, which is based on a face centered cubic lattice. If the sulfur ions (with a 2- charge) occupy the lattice points, the Zn ions (with a 2+ charge) are located one fourth of the distance along each body diagonal. Other ionic compounds that have the zincblende structure include CuCl, BeS, CdS, and HgS. Then there are fluorite structures. The Ca ions (with a 2+ charge) occupy the lattice points and each F- ion is tetrahedrally surrounded by four of the Calcium ions. Other compounds with a Fluorite structure are SrF2, BaF2, BaCl2, and PbF2. I know what you’re thinking and no, zincblendes do NOT need to contain zinc and fluorites do NOT need fluorine. It's the structure that warrants the name not the elements. This small section is not a super important lesson, as you will probably only see a question of two on the AP test so don't slave on memorizing every little thing about these structures. It's just good to know they exist and be somewhat familiar!
Most ionic crystals have high melting points, an indication of the strong cohesive forces holding the ions together. A measure of the stability of ionic crystals is the lattice energy
The higher the lattice energy, the more stable the compound. These solids do not conduct electricity because the ions are in a fixed position. However, when it is melted down into the molten state of dissolved in water, thus freeing the ions and allowing them to be more mobile, the resulting liquid WILL conduct electricity!
In a Covalent Crystal, all the atoms are held together in three dimensional networks entirely by covalent bonds. Two GREAT examples of this are the two allotropes of carbon, diamond and graphite. In diamond, each carbon atom is bonded to four other atoms. The strong covalent bonds in three dimensions contribute to the unique hardness of diamonds and it's excessively high melting point of 3550 degrees Celsius! What about graphite? Carbon is arranged in six-member rings and each atom is covalently bonded to three other atoms. The remaining unhybridized 2p orbital is used in pi bonding. Each layer of graphite has a delocalized molecular orbital that is present in benzene. Because electrons can move about freely in this delocalized orbital, graphite is a good conductor of electricity in directions along the planes of carbon atoms. The layers are held together by week van der Waals forces. The covalent bonds will also explain graphite's hardness. However, because the layers can slide over one another in graphite (as their held together, not covalently, but by van der Waal forces), it makes a good lubricant and is slippery to the touch. It's used in pencils and ribbons for printers and typewriters.
Another covalent crystal is Quartz (SiO2). The arrangement of the silicon atoms is similar to the carbon atoms in diamond. however, there is an Oxygen atom between each silicon atom. Since Si and O have different polarities, the Si-O bond is polar. Nevertheless, quartz is similar to diamond in that it is very hard and has a high melting point of 1620 degrees Celsius.
structure of diamond
structure of graphite
In a molecular crystal, the lattice points are occupied by molecules, and the attractive forces between them are either van der Waals forces or hydrogen bonding. Sulfur dioxide is a good example of a molecular crystal. The predominant attractive forces are dipole-dipole. Intermolecular hydrogen bonding is responsible for keeping the three dimensional lattice of ice as well. Other examples are I2, P4, and S8.
Besides ice, molecules in molecular crystals are packed as closely as their size and shape will allow. However, molecular crystals are broken easier than covalent because van der Waals forces and hydrogen bonding are weaker than covalent and ionic bonding! Most molecular crystals will melt at temperatures below 100 degrees Celsius.
The structure of these are the simplest because every lattice point in a crystal is occupied by an atom of the same metal. They are generally body centered cubic, face centered cubic, or hexagonal close packed. As a result, they are usually very dense. In a metal, electrons are delocalized throughout the entire crystal. The great cohesive force resulting from delocalization is responsible for the strength of metals and also explains their conductivity to both heat and electricity.
11.7 Amorphous Solids
Solids are most stable in crystalline form. But that DOESN'T mean their all in that form. If a solid is formed rapidly, its atoms or molecules won't have time to align and may be locked in a position other than that of a regular crystal. An example of a solid being formed too quickly would be when a liquid is cooled too rapidly. This is an amorphous solid. Its scientific definition is
a solid that lacks a regular three dimensional arrangement of atoms.
An example: Glass.
What do we know about Glass? It's fragile. It's used to make windows, it's see through, it's made from sand, and very nice looking, expensive things are made from it, which you will, at one point in your life, have the experience of breaking! But that's not all! It's one of the most valuable and versatile materials. It's also one of the oldest, dating back to 1000BC. By DEFINITION,
glass is an optically transparent fusion product of inorganic materials that has cooled to a rigid state without crystallizing
Glass is formed by mixing molten silicon dioxide (its main component), sodium oxide, boron oxide, and other transition metals to attain certain colors and properties. (This mixing is why glass is known as a fusion product.) There's a twist to glass. One that is as surprising as Dumbledore's death and finding out Malcolm Crowe was a ghost. Glass behaves more like a liquid than a solid! X-ray diffraction studies have shown it lacks long range periodic order.
Guess how many types of glass there are. 10? Higher. 75? Much higher. 10,472? You've gone too far. That's right there are 800 different types of glass! Wow! The color of glass is due to the presence of metal ions (as oxides).
Contains Iron (III) Oxide (Fe2O3) or copper (II) oxide (CuO)
Contains Uranium (IV) oxide (UO2)
Cobalt (II) oxide (CoO) or Copper (II) oxides (CuO)
: Small particles of gold and copper
Did you notice that most of the ions here are transition metals? Yeah...it's like that. Note that on the AP test it's good to know that Transition metals generally are linked with colors. Every test will probably have one or two questions on colors.
11.8 Phase Changes
I have one thing to say about this lesson. It’s easy. No big deal. If you got through crystal structures you’re looking at a cake walk. Let's get started. What's a phase change?
It is a transformation from one phase to another
. Whoa that was SO difficult to figure out.
Characteristics of Phase Changes:
1). Occurs when heat is added or removed
2). A PHYSICAL change characterized by a change in molecular order. Molecules in a solid have the greatest order and those in gases have the greatest randomness.
Think of a liquid as the bowl of porridge. It is not fixed in a rigid lattice like a solid and lack's the freedom of a gas. It's smack in the middle. The molecules are in constant motion. Liquids are denser than gases. Therefore, the collision rate of molecules is higher in liquid phases than gas phases. I know exactly what you’re thinking. What happens when molecules of a liquid have sufficient energy to escape from the surface? Well, a phase change known as evaporation/vaporization occurs. It is the process in which a liquid is transformed into a gas. The higher the temperature, the greater the kinetic energy, and hence the MORE molecules that leave the liquid! The hotter the water, the more molecules that escape!
What is that diagram below? It'll help explain vapor pressure. Before evaporation, the mercury in either side of the U-Shaped tube is at equal levels. Once the molecules begin to leave the liquid, a vapor pressure is formed. Evaporation is not indefinite, and eventually, the mercury level will stabilize and no change will be seen any longer. At first, molecules are only moving from the liquid to empty space in evaporation. However, as the concentration of the vapor increases, some molecules will condense and return to the liquid phase. This is known as condensation:
the change from the gas phase to the liquid phase. This occurs when a molecule strikes the liquid surface and becomes trapped by intermolecular forces in the liquid.
The rate of evaporation at a certain temperature is constant. However, condensation will increase with the increasing concentration of vapor molecules.
Dynamic equilibrium is when the rate of the forward reaction is exactly balanced by the rate of the reverse reaction
. Basically, this state is reached when the rates of condensation and evaporation are equal. The
equilibrium vapor pressure is the vapor pressure measured when a dynamic equilibrium exists between condensation and evaporation.
Equilibrium vapor pressure is the MAXIMUM vapor pressure of a liquid at a certain temperature.
Molar Heat of Vaporization and Boiling Point
Molar Heat of Vaporization:
A measure of the strength of intermolecular forces in a liquid. This is also defined as the energy in kilojoules required vaporizing 1 mole of a liquid. The molar heat of vaporization is directly related to the strength of liquids intermolecular forces. The stronger the intermolecular attraction, the more energy it takes to free the molecules from the liquid to the solid phase. Therefore, if a liquid has a low vapor pressure, it will have a high molar heat of vaporization!
The equation above is known as the Clausius-Clapeyron. It shows the relationship between the vapor pressure (P) and the temperature (T). Ln is the natural logarithm of pressure and R represents the gas constant, 8.314 J/K x mol and C is a constant. It is a linear equation that translates into y=mx + b. What does this equation do for us? By measuring the vapor pressure of a liquid at different temperatures and plotting lnP versus 1/T, we determine the slope, which will inevitably equal the negative heat of vaporization over the gas constant, R. If we know the values of the heat of vaporization and the pressure of a liquid, we can use this equation to calculate the vapor pressure at different temperatures! How? By using this rather long equation!
A good way to demonstrate the molar heat of vaporization is by rubbing an alcohol such as ethanol or rubbing alcohol on your hands. These have a lower heat of vaporization than water so, by rubbing them on your hands, the kinetic energy will increase and molecules will evaporate. This loss of heat will make your hands feel cool. This is the same basic principal as perspiration.
The temperature at which the vapor pressure of a liquid is equal to the external pressure. This is also the temperature at which a liquid boils...duh. At the boiling point, bubbles form. When they form, the liquid that occupied that space is pushed aside and the level of the liquid will begin to rise!
A bubble will experience pressure both in and on it. The atmospheric pressure and hydrostatic pressure (pressure due to the presence of a liquid) exert pressure ON the bubble whereas the vapor pressure of the liquid exerts pressure INSIDE the bubble. When the vapor pressure is equal to the external pressure, the bubbles will rise to the surface and burst. If the vapor pressure is lower in the bubble than the external pressure, the bubble will collapse before it reaches the surface. Let's draw a conclusion...the boiling point of a liquid is LARGELY dependant on the external pressure. Example: At 1 atm, water boils at 100 degrees Celsius. At .5 atm it boils at only 82 degrees Celsius.
The higher the heat of vaporization, the higher the boiling point will be. Both are determined by the strength of the intermolecular forces. The weaker the forces, the lower the boiling point and heat of vaporization. Simple Enough! Weak dispersion forces will produce low heat of vaporization and boiling points whereas dipole-dipole forces will attribute to the opposite.
DID YOU KNOW?!?! Of all liquids, Mercury has the highest boiling point and heat of vaporization. Can you guess why? It's strong metallic bonding!
Critical Temperature and Pressure
The opposite of evaporation is condensation. The liquefaction of gas baby! How does this occur? There's two ways.
First: cooling a gas will decrease its kinetic energy so the molecules will eventually aggregate and form small drops of liquids.
Second: Apply Pressure. Compression reduces the distance between molecules so they're held together by mutual attraction.
Third: Apply both,
1. A substance in the gas phase WILL NOT liquefy, no matter how great the pressure, if the temperature is above the critical point.
2. The highest temperature a substance can appear as a liquid.
3. Above the critical temperature, there is no distinction between liquid and gas
The minimum pressure that must be applied to bring about liquefaction at the critical temperature.
What exactly causes the existence of the critical temperature? The intermolecular forces below this temperature are strong enough to hold molecules together. However, above it, the molecules become so energetic, that some break from their bonds. Again, high intermolecular forces will result in high critical temperatures.
Liquid Solid Equilibrium
Liquid Solid Equilibrium
Freezing and melting is our next study. The freezing and melting points are the temperatures of a substance at which the liquid and solid phases can coexist. Above the freezing point, the substance will solidify. Below, it will liquefy. For example, at 1 atm of pressure, water's melting/freezing point is at 0 degrees Celsius. To the left, there is a normal
heating curve of water. Let's start from the bottom left and work our way up the graph. A solid is gradually heated until it reaches its plateau. at this point, the solid will begin melting. The line is flat at the plateau because heat is being absorbed by the system, but its temperature remains constant. The heat aids molecules overcome the attractive forces in the solid. Once the ice has completely melted, the absorbed heat will now increase water's kinetic energy and the diagonal line will represent the gradual heating of water until it reaches its plateau, at which point it will vaporize. Again, at the plateau, the liquid is changing to a gas while the temperature remains constant. Afterwards, the temperature of the gas will continue to rise. In this particular graph, the plateaus are most likely at 0 and 100 degrees Celsius. A graph like this is almost ALWAYS on the AP test.
This is usually how a heating curve appears on exams. However, it would have no key terms of what's happening at each letter A-E.
Molar Heat of Fusion:
The energy required to melt 1 mole of a solid! The heat o f fusion is usually smaller than the heat of vaporization Molecules in a liquid are closely packed together so there is some energy required to bring abound the rearrangement from solid to liquid. When a liquid evaporates however, molecules will become completely separated, requiring much more energy to OVERCOME the attractive forces. Easy Enough!
Cooling a substance will have the opposite effect as warming it...that makes sense! Therefore, if heat is removed from a gas what will happen? It will lower in temperature. As the gas changes into a liquid, heat is given off because its potential energy is decreasing. Just like heating a substance, during the phase change from a gas to a liquid (condensation), the temperature remains...CONSTANT. Then, as the liquid continues cooling, the temperature drops until it finally freezes. Basically, this process is the same as warming a substance only...the OPPOSITE!
Solid Vapor Equilibrium
Did you know that it is possible for a solid to transform directly into a gas? This process is known as
When gases form directly into a solid, the process is known as deposition. Some substances, like Naphthalene and Iodine, with fairly high vapor pressures, will sublime. The vapor pressure of a solid is generally much lower than that of a liquid because its molecules are held together tighter. The
molar heat of sublimation
is the energy in kilojoules required to sublime 1 mole of a solid. Uh oh...equation time!
This equation proves that the molar heat of sublimation is the sum of the molar heat of fusion and sublimation. This equation is an illustration of Hess's Law
This equation also demonstrates that the heat change for the overall process from a solid changing into a gas whether the solid changes directly into a gas or changes into a liquid first. However, this equation is only valid if these phase changes occur over a constant temperature. If not, it is only an approximation.
11.9 Phase Diagrams
Phase diagrams will usually appear at least ONCE on the AP exam so it's good to know these inside out. Lucky for us, they're very easy to grasp. A
summarizes the conditions at which a substance exists as a solid, liquid, or gas. The best way to study phase diagrams are to simply study examples. Let's start with water!
This diagram is a simple phase diagram of water. As you can see, the diagram is split into three sections, each representing a phase of water. In every phase diagram of every substance, the solid will generally command most of the left side, the gas will generally take over the bottom (but increase as the temperature gats higher), and the liquid will be on top, sandwiched between the two. The lines that separate the phases stipulate the conditions in which each state can exist in equilibrium. For example, the curve between the gas and solid phase shows what the pressure at any given temperature must be for the two to occur at equilibrium and for sublimation to occur. The other curves represent equilibrium states for a solid and liquid to coexist and a liquid and gas to coexist. The point in the middle is known as the
. Whoa what the heck is that? The triple point is the ONLY condition under which all three phases can coexist TOGETHER in equilibrium. In the case of good old water, at .01 degrees Celsius and .006 atm, water reaches its critical point.
Phase diagrams enable us to predict changes in melting and boiling points of a substance as a result of a change in external pressure. For example, the normal boiling point of water is 100 degrees Celsius at 1 atm of pressure. What would happen if the pressure was raised above 1 atm? The boiling point would also rise. However, the freezing point would dip below 0 degrees Celsius. See how much can be determined! I know, the critical point is not described in the section. It's still important to know what it is though. Once the critical point is reached, it will be impossible to distinguish the phase as either liquid or gas. This is known as a super critical fluid!
Review Questions and Calculations!
1. The following compounds are examples of a dipole-dipole interaction, hydrogen bonding, ion dipole, and London Dispersion Forces. Label each of the following a. Hydrochloric Acid (HCl) b. Cl
2 c. HF d. HO e. Cl- mixed with water f. HBr
a. dipole-dipole b. London Dispersion Forces c. Hydrogen bonding d. Hydrogen Bonding e. ion-dipole force f. dipole dipole
2. What's a permanent dipole? What's an instantaneous dipole?
Answers: A permanent dipole occurs in a polar molecule. the positive and negative chemical bonds will not overlap completely and the compound will have a permanent dipole moment. An example of this would be Carbon Dioxide. An instantaneous dipole exists in non-polar molecules. The electrons are constantly orbiting the nucleus and it is possible that, for a fraction of a second, there may be more electrons on one side than the other and, for that split second, the molecule will have a dipole moment.
3. Explain some of the differences between gases and the condensed states.
Answers: Liquids and solids, for the most part, cannot be compressed and have definite volumes. There molecules are also very close together with very little empty space. Gases have no definite volume and there is a great deal of free space between the molecules.
4. Which has a higher boiling point: Krypton or Argon? HI or H2O? PbH4 of CH4?
Krypton, because it is in the same period as Argon but contains a greater molar mass.
H2O because H2O takes part in more intermolecular hydrogen bonds than HI
PbH4 because Pb has a greater molar mass than C and they are in the same period.
5. Define surface tension and viscosity and explain the correlation between intermolecular forces and surface tension.
Surface tension is the amount of energy required to stretch or increase the surface of a liquid by one unit area. Usually, the greater the intermolecular forces in a compound, the greater the surface tension. Viscosity is a measure of a fluid's resistance to flow. The greater the viscosity, the more slowly the liquid flows.
6. What's the difference between cohesion and adhesion?
Cohesion is the intermolecular interaction between like molecules. For example, water molecules interacting in a capillary tube. Adhesion is an interaction between unlike molecules, such as these same water molecules and the glass side of the capillary tube. Both of these forces play a part in capillary action.
7. Which has a greater surface tension? Water or benzene? Why?
Answers: Water, because water undergoes hydrogen bonding, a very strong type of intermolecular force.
8. Explain the difference between body centered, face centered, and simple cubic cells.
Body Centered Cell: The second layer of spheres fit in the depressions of the first and third layer. Coordination number of 8 and contains the equivalent of two complete spheres. One in the middle and one made up by the sum of its eight shared corner spheres.
Face Centered Cubic: Spheres at the center of each face of the cube as well as each of the eight corners. It contains four complete spheres: Three from the six face-centered atoms and one from the eight shared corner spheres.
Simple Cubic: Coordination number of six. Basically one layer of spheres stacked on top of another. Equivalent of one sphere from the sum of eights shared corner spheres.
9. Define: Closest packing, unit cell, and crystalline solid.
Closest packing: the most efficient arrangement of spheres to minimize empty space.
Unit Cell: the basic repeating structural unit of a crystalline solid.
Crystalline Solid: possesses rigid and long range order and its atoms, molecules, or ions occupy specific locations.
10. Iron crystallizes in a body centered cubic unit cell and has a density of 7.784 g/cc. Calculate the atomic radius of iron.
11. What is the coordination number of the following: Simple Cubic Cell, Body Centered Cell, and Face Centered Cell?
12. X rays of wavelength .168 nm strike a crystal and the rays are reflected at 21.2 degrees. Assuming n = 1, calculate the spacing between the atoms that is responsible for this angle of reflection. 1 nm = 1000 pm.
Answers: We need to use the following equation to solve this!
That was easy!
13. The distance between layers in a crystal is 250 pm. X rays are diffracted from these layers at 17 degrees/ Assuming that n = 1, what is the wavelength in nanometers?
Answers: We just use the same equation only now, the wavelength is in question. This is the proper equation:
Our answer is .146 nanometers. Wavelength is generally a small answer!
14. Easy question. This figure to the right is a CsCl unit cell. How many cesium atoms and how many chlorine atoms are
in each unit cell?
Answers: First, Caesium is in the middle. This is a simple cubic cell. A better mental interpretation is shown below. Since each chlorine atom is shared by three other unit cells, each will only count as 1/4. However, there are 8. 8x(1/4) = 2 chlorine atoms and there is one cesium atom, unshared, in the middle!
15. What is an amorphous solid? Give an example!
Answers: An amorphous solid lacks a regular three dimensional arrangement of atoms. The BEST example is glass. Other examples can be boron oxide, sodium oxide, and silicon dioxide.
16. At 19 degrees Celsius, the vapor pressure of ethanol is 40 mmHg. Calculate it's vapor pressure at 28 degrees in mmHg. Heat of vaporization is 39.3 kJ/mol
Answers: Let's define our variables. P1 = 40 mmHg, T1 = 19 degrees Celsius = 292K P2 = ? T2 = 28 degrees Celsius = 301K The given values of the heat of vaporization of ethanol is 39.3. We'll use this equation to solve our problem.
The negative .484 is the sum of the right side of the equation. Our final answer for the new vapor pressure is 24.65 mmHg!
17. Define sublimation and deposition.
Answers: Sublimation is when molecules go directly from the solid to the vapor phase. Deposition is when vapor molecules go right to the solid phase!
18. Below is a phase diagram. Define A, B, C, what is happening at each arrow, and what the triple point and critical point is.
Answers: A is the solid phase, B is the liquid, and C is the gas phase. The arrow from A to B represents the melting point and the arrow from B to A is the freezing point. The arrow from B to C is the boiling point and the arrow from C to B is condensation. From A to C is sublimation and from C to A is deposition. The critical point is where B and C will coexist and there will be no distinguishable difference and the triple point is where A, B, and C can coexist.
19. EASY question: This is a phase diagram of carbon. How many triple points are there in this diagram? What is occuring at these points? What makes this phase diagram unique?
Answers: There are two critical points. At the top one, diamond, liquid carbon, and graphite will coexist. At the bottom, gaseous carbon, graphite, and liquid carbon will coexist. This is unique because carbon can come in two solid forms depending on the pressure and temperature. There's also some overlap between carbon and graphite where either can exist.
20. Last Question: Did you have fun learning?!
Answers: Yes, this was both exciting and informative! Thanks so much Garrett!
Intermolecular Forces and Surface Tensions!
In this experiment, the surface tension of glycerin, rubbing alcohol, and water will be used to determine the strength intermolecular forces. Their will be a variety of experiments done with each liquid to assure accuracy!
3 Erlenmeyer flasks with stoppers
3 petri dishes
3 plastic pipets
100 ml beakers containing:
water, isopropyl alcohol and glycerol (aka glycerin)
Surface tension and vortex.
When a liquid is swirled, a vortex is developed in which the surface level of the center of the liquid is substantially below the surface level of the perimeter.
The greater the surface tension, the longer the vortex will remain after you have stopped swirling the container. Fill one flask half-full with rubbing alcohol, another with water, and yet another with glycerol. (Remember to fill the flasks only half-full).
Stopper the flasks to prevent vapors from polluting the room.
Try to swirl each flask with the sam
vortex in water
e intensity and record the time it takes for the vortex to disappear. Which liquid appears to have greater surface tension and greater intermolecular forces?
Record your answer for Part 1
paper clip on water
Surface tension and droplet shape:
Using an eyedropper or pipet, transfer one drop of each fluid to a sheet of wax paper. The liquid with greater surface tension will maintain a higher profile and will not spread out as much as the one with lower surface tension.
Which liquid appears to have the greater surface tension and greater intermolecular forces?
Record your answer for Part 2.
3. Surface tension and impenetrability:
Liquids with strong intermolecular bonding will be less penetrable than those with weaker intermolecular bonding.
Try to float a paper clip on water, rubbing alcohol, and glycerol by gradually lowering a dry paper clip into each liquid on a cradle fashioned from another paper clip (Figure 3).
It may be best to use a small beaker and some forceps for this procedure.
Which liquid appears to have the greater
surface tension and greater intermolecular forces? Record for Part 3.
Visualization of surface tension:
The surface of a liquid with strong hydrogen bonding will exhibit great tension much like the head of a drum that has been pulled tight.
If a drumstick ruptures the head of a drum, the sides recoil under the tension.
In a similar manner, if a chemical ruptures the surface tension of a fluid, the "skin" of the liquid will recoil away from the point
where the chemical was applied.
Fill one Petri dish with water, another with isopropyl alcohol, a
nd the third with glycerin.
Sprinkle crushed pepper on the surface of both.
The pepper will be more likely to float on the fluid with greater surface tension (Figure 4).
Cover the tip of a paper clip with liquid dish soap and hold over the center of each Petri dish until a drop of soap falls into the liquid.
If the surface of the liquid is under tension, the pepper will recoil towards the sides immediately (see picture).
Which liquid appears to have the greater surface tension and greater intermolecular forces?
Record your conclusion for Part 4.
1. From your observations, list the compounds observed in order from strongest intermolecular force, to weakest.
2. Which liquids do you think would boil the easiest and why?
3. What errors could have possibly changed your results?
This lab was adapted from:
All factual data was taken from Eighth Edition Chemistry by Raymond Chang
All jokes were from Garrett Gallinot!
Closest Image and many other images were revised on paint
All equations made on equation editor
help on how to format text
Turn off "Getting Started"