Week 12

This past week in AP chemistry we were introduced to the concept of equilibrium in chemical reactions. We were first told that most of the information we had learned about chemical reactions going to completion was not accurate in reality, instead, every reaction has a proportion that it will stay at once it is reached under certain conditions. This is the basis of equilibrium. Using information of temperature, volume, concentration or pressure, and constants, you can often calculate more useful information. After becoming familiar with the topic, you are also able find how a change in any characteristic of a system will affect the particles contained, as well as how you may make such a change.

Reactions under identical conditions will approach the same reaction quotient of equilibrium. A reaction quotient is equal to the product of the concentrations of the products, each to their stoichiometric coefficients, divided by the product of the concentrations of the reactants, each to their stoichiometric coefficients. This is called the Kp of a reaction under these conditions. This value is always constant for a particular reaction under constant temperature. Temperature is the only value that can affect the Kp of a reaction, while many others still affect the system as a whole. In the same way, you are able to find the Qp of a system. This value tells you the current place on the equilibrium scale and with it you can find which way the reaction will shift.

Even in equilibrium, reactions never stop. Reactants and products are continuously being converted between each other. Only in equilibrium are the rates of each direction equal, resulting in a constant concentration for each group. This is why adding or removing any reactant or product to a system results in a significant change in the proportions of each. Adding reactant will make a reaction shift to the right, toward the products, while adding product will make a reaction shift back to the left, toward the products. Removing either of these two does just the opposite. In the case of adding one of the reactants, you are able to push the reaction further, decreasing the remaining amount of the other reactants, while increasing the amount of the one you are adding. Chemical engineers can use this technique to save money, pushing more of an expensive reactant through with an inexpensive one (see link below for practice).

An  illustration of manipulating conditions to obtain desirable products

In class we were given a challenge, requiring us to use our prior knowledge in combination with the topics we have just learned. We were given a reaction and a yield and told to find the reactant quotient of the reaction and draw a diagram. I found this exercise very useful in improving my ability to draw together larger concepts to find useful and more realistic values in the real world.

Practice with manipulating reactions through conditions
http://www.files.chem.vt.edu/chem-ed/courses/equil/intro/lechatel.html

Week 11

These past weeks in AP Chemistry we have worked on the unit of gas laws. These laws, formed long ago, define the interactions between the many properties of any system of gas. Our lesson included both theoretical calculations as well as more complicated real-world calculations, with far more accurate results. As usual, we were first given the information through online lectures, allowing us to pause and repeat sections, followed by a simple quiz to test our overall understanding of the information and for our teacher to see the points in the lecture that did not properly explain themselves. We followed these up in class with ConcepTests, which are taken in the open, each student holding up the number of fingers corresponding to their answer, for the teacher to find and solve the areas in which multiple students are still having the most difficulty. I personally found this very useful because, although I understand the main concept areas already, it helps me to see which spots are most problematic for me and my peers in regards to simple mistakes due to negligence.

We used an online PHeT simulator to initially introduce us to ideal gas situations and interactions. We filled out a pre-made Google Doc spreadsheet in order to find the relationships on our own, before learning about them. This was very useful way to offer the students a much more realistic look into the way science works outside of the classroom: data is gathered in as high volume as possible, and then pulled in together and analyzed as a whole. Through this process we were able to easily be guided into finding Boyle's, Charles', and Guy Lussac's laws of ideal gas interactions of various properties. We were later guided into finding the Combined Gas Law, stating that the product of pressure and volume divided by the temperature of a gas will always equal a constant. We continued by adding n (number of moles) into the mix. Through further data collection from the simulator you could find that the number of moles was directly proportional to pressure or volume, while it was inversely proportional to temperature, given in the well known equation PV = nRT, where R is a constant coefficient. For further practice see the link below.

Classmate holds balloon of oxygen gas in liquid
 nitrogen changing the volume

After learning all about ideal gas situations, we were told that this was not exactly how the world worked. We had previously assumed in our calculations that all particles had volumes of zero and there were no interactions between the different molecules in the container, which is not the case due to intermolecular forces (covered in an earlier blog). In reality, substances contain dipole-dipole, dipole-induced dipole, London dispersion and hydrogen bonding interactions, in different combinations. These in turn attract all particles within a system toward each other, decreasing the overall pressure of the system. Similarly, every particle has a small but significant volume, which increases the overall volume of the system. We learned that in order to calculate more accurate values you must find the constant for each molecule and incorporate it into the equation.

Practice with Gas Laws:
http://www.sciencegeek.net/Chemistry/taters/Unit5GasLaws.htm

Week 10 (Trimester 2)

These past few weeks in AP chemistry we have been introduced to the many aspects of thermochemistry and their applications to the real world. Inside and outside of class we have utilized various tools to understand these new concepts from every possible angle. Making connections on our own. These tools included videos, worksheets and class discussions, all of which offered new ways of thinking about ideas and the ways that different properties interact. As a result of learning this new material on thermochemistry, I now believe that it is one of the best connections from the microscopic world of atoms and molecules, and the macroscopic world of "us."

We first cleared all confusion on the ideas of heat and temperature. Heat is the transfer of kinetic energy between the particles of different systems, while temperature is a quantified measure of the speed of the particles in a particular system. We were then introduced to some key concepts behind thermodynamic calculations. The specific heat capacity (Cp) of a substance is the energy required to raise a particular mass of the substance a particular degree of temperature, typically J g-1 K-1. Enthalpy is the total energy within a system, and is generally considered useless on its own. Instead, scientists use the change in enthalpy (ΔH) of a system when calculating certain values. This value is special in that it includes the energy change from temperature as well as from a change in the state of the substance. Entropy is a measure of the number of separate microstates that a substance may be found in. In other words, it measures the number of different forms the atoms or molecules of a system may line up at various points in time. These values are each known as state functions, meaning that the final value is independent of the path taken for it to be achieved.

Enthalpy as a state function

Using the properties of thermodynamics that we had learned, we began to use certain equations in order to calculate the various hypothetical temperature changes within different experiments. The first equation we used multiplied the mass of the system by the specific heat by the change in temperature to find the energy change in the system, assuming the state did not change. Given any three variables of this equation, the last could be solved for with algebra. We then added the heat of fusion and vaporization, in energy per mass unit, to find the total enthalpy change of any system, with state changes included. We then moved on to Hess's law, stating that the enthalpy change required to form a compound is equal to the sum of the products minus the sum of the reactants (see link below). The next equation we used was Gibbs free energy, for calculating the ability of a reaction to do work. This reaction states that the ability of a reaction to do work is equal to the change in enthalpy minus the temperature multiplied by the change in entropy. I had previously never imagined that I would be able to accurately calculate these types of properties of reactions with so few variables and reference materials.

Combustion of magnesium in class

Practice with Hess' law: http://chemistry.about.com/od/workedchemistryproblems/a/Hesss-Law-Example-Problem.htm

Week 9

This week in AP chemistry we delved into greater detail on the electrostatic forces between atoms and their effects on the different characteristics of the substance they make up. Each of these separate classifications has its own particular set of rules determining which of its properties are affected by its particular structure. The separate classifications are atomic bonding structure, bonds formed, boiling/ melting point and vapor pressure.

The way in which a substance is made up overall will typically have the largest magnitude in determining the properties of that substance. First, with the lowest melting and boiling points in most cases, covalently bonded compounds are molecules made up of nonmetal elements who share their electrons in order to complete their octets. Such as carbon dioxide, nitrogen trifluoride and sulfur dichloride, these elements have the weakest total molecular forces, resulting in smaller quantities of energy required to pull them further apart. Next, substances with ionic bonds generally have much higher melting and boiling points than compounds with covalent bonds, and for this reason they are more commonly encountered as solids at room temperature, such as sodium chloride, better known as table salt, a combination of sodium cations and chlorine anions. The final group with the highest melting and boiling points, are often solids that can not be melted at all. These substances are known as network solids. They are composed of elements with covalent bonds, but no particular distinctions of individual molecules. Instead, these substances are composed of  "macromolecules," a single lattice structure of a very large quantity of atoms, all with covalent bonds between those next to it. Common examples of this include diamond and graphite, both structures of only carbon atoms bonded together in one piece, and quartz, composed of a large network of silicon dioxide. Their molecular bond strength comes from the continuous network of bonds within the entire substance, and that melting them would require breaking all of the covalent bonding (See figures below).

Quartz

Many characteristics of substances are further determined by the specific types of bonding that occur within the molecule. Especially present in covalent bonding, the presence of extra van der Waal forces each increase melting and boiling points. Typically the weakest, London Dispersion Forces (LDFs) are present in all substances, although to varying degrees. The strength of this force depends on the polarizability of the molecule, which increases with the size of the molecule, factoring in the number of elements and their sizes, which is ultimately the area with which they may form temporary dipoles with those next to it. Next, dipole interactions increase melting and boiling as a result of increasing the attraction between the positive and negative sides of separate atoms and requiring more energy to pull them apart. Most significant when present, hydrogen bonding can greatly increase the melting and boiling points of these compounds. The particular reasoning behind this trend was not explicitly described to us, but I have concluded that this may only occur between hydrogen and the elements nitrogen, oxygen and fluorine because only these elements have greater electronegativity values than hydrogen, making it negative and the hydrogen positive, and more likely to bond with the typically negative outside of another molecule.

For an additional recap of molecular trends, try the link below:
http://www.masterorganicchemistry.com/2010/10/25/3-trends-that-affect-boiling-points/

Week 8

This week in AP chemistry we were introduced to the correlation of intermolecular forces and characteristics of different substances. The new vocabulary we were given included distinguishing between intermolecular forces and intramolecular forces. Intermolecular forces are those that form between separate compounds, including London Dispersion Forces, hydrogen bonds, dipole- dipole bonds, dipole- induced dipole bonds and ion- dipole bonds. Intramolecular bonds are those which form inside of the compound, holding it together, including molecular and ionic bonds. The combination of these bonds is the largest contributing factor in the melting and boiling points of any compound.

The intermolecular forces, most prevalent in covalently bonded compounds have a significant influence on the melting and boiling points of the compound. The weakest of these forces is London Dispersion Forces, in which a temporary dipole in one compound due to an increased or decreased electron density creates a temporary dipole in the compound adjacent to it, followed by a weak attractive force between the two partially positive and partially negative ends (See figure below). This particular force is found in all compounds, while magnitude can vary greatly due to polarizability. Polarizability is a compounds capacity to form temporary induced dipole- induced dipole bonds, determined majorly by its contact or surface area with the compounds around it.

London Dispersion Forces between two atoms

Another force of similar strength is the dipole- induced dipole force. This bond forms between a molecule with a significant dipole and a non-polar molecule in which the electron density is influenced by the dipole of the other, forming a temporary dipole along with attractive forces. Next and the most simple of this group, dipole-dipole interactions form between the positive dipole end of one molecule and the negative dipole end of another, these interactions are significantly stronger as a result of permanent partial charges belonging to both molecule. The last of the van der Waal forces is hydrogen bonding, typically stronger than all of the rest. These bonds form in substance with hydrogen atoms bonded to nitrogen, oxygen or fluorine. Only these three elements have electronegativity values greater than that of  hydrogen and are able to form these bonds. They are the main reason why water takes its rigid hexagonal rings shape as ice, with a greater volume and lesser density than liquid water (See link below for practice with these).

Intramolecular forces within particular substances have a larger impact melting points and boiling points than intermolecular forces. Most importantly, ionic compounds with large charges can have incredibly high melting points and boiling points. It has been observed that these characteristics increase along with bond strength and opposite bond length as a result of atomic radius. Increasing the charge of the ions also has a great impact on these characteristics, such as in the case of sodium chloride have lower melting and boiling points than magnesium sulfide. Covalent bonds also have much more strength than van der Waal forces and increasing the molecular weight of a molecule increases its melting and boiling points, such as adding additional carbon layers to alkanes.

Molecular Forces and Melting Points worksheet with useful graphs:
http://www.dublinschools.net/Downloads/Key-boiling%20points%20and%20IMF.pdf

Week 7

This week in AP Chemistry we reviewed and were tested on covalent bonding. Some of the topic areas in this unit were the Lewis structures formed by these bonds, the electron domain geometry and molecular domain geometry around the central atoms or any other particular atoms and the hybridization of s- and p-orbitals along with the very recent debate over the hybridization of the d-block orbitals and their effects. In order to review for the test we utilized a new tool on the class Moodle called a TaskChain, a series of quizzes in which you must achieve a 90 percent or higher score, given partial credit for correct second and third answers, in order to move on to the next in the series. This simple quiz offered me three significant advantages. First, it offered me more peace-of-mind, knowing that I was able to pass all nine with at least an A grade on my first attempt, even when clicking mistakes were accounted for. Second, this way of reviewing was perfect for the teacher to show me the types of questions that would be on the test and what they would generally look like. Third, these quizzes were extensive and covered nearly all material, reminding me of any relationships or topic areas that may have slipped my mind over the course of the few weeks spent since the previous test.

After the covalent bonding test, on Mole Day, we were assigned to read a passage about polarity and its significance in paintball and then write an essay on what polarity is, including facts, definitions and quotations. Polarity was introduced to us in the previous unit, with dipole moments, being the partial charges on each side of a molecule pointing in a specific direction with a specific magnitude measured in Debye. This article carried this information over as well as introducing the significance of polar molecules in chemistry. The most common of these molecules, water, lines up end-to-end, positive-to-negative because the liquid molecules are allowed to slip passed each other and opposite charges attract (See link below for polarity determination help). The contact between the positive and negative ends of the molecules form polar bonds. the most common type of this bond is the hydrogen bond, between a hydrogen atom and another oppositely charged atom. This concept relates to the covalent bonding we just finished in that atoms are brought together because of charges dealing with electrons. This most significant difference in these different types of bonds is that polar bonds do not share the electrons in any way and are much weaker as a result.

A paintball match beginning

Polar bonds are what allow other polar substances to dissolve in water. The significance of this was shown in the article we read, which mentioned how scientists who were developing new types of paintballs looked for a substitute for the old paint that would be water soluble and therefor much easier to wash for the players. Applying this change to the game of paintball made its popularity skyrocket, showing how little innovations involving chemistry can improve all the different things that we do in our lives.

Toward the end of the week we were briefly introduced to the next subject of ionic bonds and their properties that can be readily determined from given information. The most important principle was that the shorter the length of an ionic bond, the more energy is released. From this you can take that the atomic radii of the elements bonded are connected to the energy of the bond. Then, knowing the law of the conservation of energy, it is logical to conclude our next idea, that boiling points increase as bond length decrease. This is because more energy was released when the bond was formed, and as a result more energy is required to pull them apart when it enters its gaseous state.

Polar molecule determination summary:
http://users.stlcc.edu/gkrishnan/polar.html

Week 6

This week in AP Chemistry we examined the hybridization theory held by many modern scientists. We were reminded early on that it was very important that we remember that this particular conjecture is a theory and there is currently an ongoing debate over many of the specifics of these processes, with data to support all sides of the debate. In class, we took extra care to focus on the areas with less debate, that scientists are much more sure of. In the case of hybridization, scientists are most sure of the hybridization of molecules with two to four electron domains around the central atom. The hybridization for the molecules is sp,sp2 and sp3 (numbers should be superscript), for molecules with two, three and four electron domains, respectively. These names mean that the hybridized orbitals were formed from the combination of an s-orbital and that particular amount of p-orbitals. For molecules with five or six electron domains, it had been believed, until a few years ago, that d-orbitals were involved with these hybridization, but today it is more widely accepted that molecules with these geometries do not hybridize at all.

Many students initially struggled with the idea of hybridization. For me in particular, it seems that this was mainly a result of not being able to see the significance of hybridization, and how the orbitals combined to make these many different shapes, struggling to draw the connection from the various-shaped orbitals we learned about over the summer and the uniformity of the ones presented to us now. The biggest help came when it was simply broken down into the relationships between hybridization, electron domain geometry and the number of electron domains. Simply, if you know any of this information, there is only one possibility for each of the others for that molecule. To check or solve, all you need to do is add the superscripts to get the number of electron domains to find the electron domain geometry, in any order.

Along with the introduction of hybridization, this week we analyzed some relatively ordinary molecules in a very advanced way with the use of the WebMO program and a supercomputer from Hope College. This program allowed us to find many important details about each molecule with incredible accuracy simply through entering the structure of the molecule (atoms involved and bonds). These details include all bond angles, dipole moments, individual partial charges on each atom, and manipulable space filling diagrams to show polarity (pictured below). In class we all filled out a chart with this information after building certain molecules, on for each electron and molecular domain configuration. I noticed that all bond angles followed our rules for these geometries, the standard angles for those without unshared pairs and less than the standard angles for those with unshared pairs repelling the bonded pairs.

For an overview of electron and molecular domain geometries with example molecules go to the link below:
http://www.sparknotes.com/testprep/books/sat2/chemistry/chapter4section8.rhtml

NSF (Thiazyl Fluoride) space filling diagram from WebMO