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

Week 5

This week in AP Chemistry we learned many of the characteristics of the VSEPR (Valence Shell Electron Pair Repulsion) Theory. This theory helps to explain the reasons why molecules take on the shapes that they do, especially simple central-surrounding atom molecules with covalent bonds such as CH4 and H2O. This theory states that the shapes of these molecules depend strongly on the interactions of the valence electrons of the central atom. It classifies the electron pairs into two categories: bonded and lone pairs. The properties of these electron pairs determine the three scientific classifications for molecules of this type: molecular class, electron domain geometry and molecular domain geometry.

Molecules are described in various ways, including the following. Molecular classes, such as AB2E3 are very concise summaries of the shape of a molecule and can be used to determine the electron domain geometry and molecular domain geometry for that particular molecule. In this notation an A is used to represent the central atom, the B is used to represent a bonded electron pair and the subscript determines the number of these pairs, and the E is used to represent a lone pair of electrons and its subscript shows the number of these pairs. For this notation, the subscripts of B and E should always sum to the number of electron pairs of the central atom. Electron domain geometry is the overall shape of the central atom's electrons, making no significant distinction between bonded and lone pairs. Molecules with central atoms with 2,3,4,5 and 6 electron pairs are classified as linear, trigonal planar, tetrahedral, trigonal bipyramidal and octahedral, respectively. The molecular domain geometry of a molecule is a sub set of its electron domain geometry and are determined by the number of lone pairs within the original structure. A site I used to review these different shapes is listed below.

Two models we used in class in order to show
the interactions of the valence electrons

Throughout middle school and into high school, I have always seen H2O in its bent molecular domain geometry while CO2 was in a linear form, and I was puzzled why this happened. I was very curious to see whether there would be a simple set of rules such as I have found in VESPR or a complicated set of strange rule-breakers requiring large amounts of memorization. I am glad that there is a simplistic, although occasionally difficult, way to determine the shape that a molecule will form, knowing nothing aside from what the molecule consists of.

A significant point was made this week that in VSEPR models the lone pairs of the central atom are significantly large than bonded pairs, as a result of the forces on them. This means that the lone pairs of the molecule will repel other pairs farther than the bonded pairs. In turn, the molecule retains its overall molecular domain geometry, while the angles involved in these shapes become much more complicated, straying from clean numbers such as 90, 120, 180 and 109.5. Instead, without further calculations, you may only estimate that it will be slightly more or less than these original numbers. I hope that sometime in the future we will be able to calculate these angles more precisely.

A good summary of VSEPR characteristics:
http://misterguch.brinkster.net/VSEPR.html

Week 4

This past week in AP Chemistry, we started out with new material on the Lewis structures of atoms. Using POGILs, we learned about the characteristics of the covalent bonds that make up many molecules in the real world. A covalent bond forms between any two separate atoms that share a pair of electrons in order to follow the octet rule. Bond order is the number of pairs of electrons shared between two atoms in a particular bond, specified as single, double or triple and depicted in a Lewis structure as one, two or three lines, respectively. Bond order can be in a theoretical form based on the information from the Lewis structure in which it will have a whole number value, or an experimentally calculated value with any number of decimal places. The bond energy/strength of a particular bond is the energy required to sever the covalent bond and is most often measured in kilojoules per mole (kJ/mol). Calculated bond length is the distance of the bond, being the distance between the nuclei of each atom involved, typically measured in picometers.

With these POGILs focusing on bond order and strength we also learned about the relationships between each of the characteristics of covalent bonds. As the bond order of a bond increases, the bond energy also increases because more electrons are being shared and must be broken apart. Similarly, as the bond order between the same two atoms increases, the bond length decreases because of the increases attractive forces between the two. Conversely, as bond length increases between  two atoms, the bond energy decreases due to the effects of coulomb's  law relating distance to electrical force. Additionally, bond length is strongly influenced by the relative size of the atomic radii of the bonded atoms (see link below for more detail and available practice.

All together, many of my ideas have changed in regard to the covalent bonds that make up many of the molecules I am so familiar with. Previously, I had never considered that there may be a relationship between the properties of individual atoms and how they form bonds with each other. Furthermore, I had not thought that many of the characteristics of these bonds would allow you to predict how to draw its Lewis structure or predict its many molecular abilities and how the Lewis structure or its properties could be used to find the characteristics of these bonds. In summary, nearly all pieces of information related to Lewis structures can be used to find many other characteristics of the situation.

The reaction of nitric acid and brass

In class this week we began our second experiment, trying to determine the percent of copper by mass of a piece of brass, an alloy of copper and zinc. The way that we intend to calculate this value is through measuring the absorbance of light by the solution of the brass and nitric acid we added as well as the absorbance of a solution of known concentration of copper to find a calibration curve to use to find the original concentration of copper.



Extensive overview of bond characteristics and practice:
http://s-owl.cengage.com/ebooks/vining_owlbook_prototype/ebook/ch8/Sect8-3-a.html

Working on the experiment

Week 3

This past week in AP Chemistry, we began with some extended applications of stoichiometry to real world situations. We were asked to calculate the empirical formula of a compound consisting of only carbon, hydrogen and oxygen with a known mass, given only that it had reacted with an unknown amount of oxygen gas in a combustion reaction forming certain masses of carbon dioxide and water. Up to this point we had only learned how to calculate empirical formulas given the mass percentage of each element directly, through the use of molar masses and mole ratios.

This new challenge required no additional knowledge of stoichiometry, but significantly more thought and work. One had to retrieve previous information from other units such as the law of chemical equations stating that the number of atoms of each element as well as the total mass must be equivalent on the reactant and product sides of every balanced chemical reaction. Using what we had learned in this lesson, we could solve for the mass of each of the three elements through the mass percentages of the elements in the products. Subtracting the total mass of the reactants by the given compound gave the mass of the oxygen gas that had been burned. You were left with the mass of each element in the compound being examined; the same setup as all other problems this unit.

Challenging problems such as these are greatly beneficial to our learning process in chemistry. They help students connect multiple ideas, often from different content areas together to answer a more complex problem than those they have encountered in the past. These are the best way to assess and strengthen a students true understanding of the material at hand because they must pull the information together by feel, navigating on their own, rather than simply following a step-by-step process they were given for one specific set of problems.

After being assessed on stoichiometry, we moved on to the Lewis structures of atoms (as depicted in the images below). Developed by scientist G. N. Lewis, this notation is commonly used to show the structure of molecules through the covalent bonds they form. Covalent bonds form between atoms that share electrons in order to complete their octet and become more stable like the noble gases nearest them, aside from hydrogen which only forms a duet with a single bonding pair. The octet rule is mainly a result of the eight electron spaces in their valence electron shell. The valence electrons of any atom are those in the s and p orbitals of their highest energy level. These particular electrons are the only reason that the nonmetal elements form compounds, because nonmetal elements will not form positive ions (aside from ammonium). For the best practice drawing Lewis structures, see the link below.

Lewis structure practice:
http://www.chem.purdue.edu/vsepr/practice.html

Lewis structure of HCN (left) and other common
molecules (right)

Week 2

This past week in AP Chemistry we focused on stoichiometry, the area of chemistry dealing with the ratios of atoms in chemical reactions, and the many applications it has in the real world. To help the class better understand the topic of stoichiometry, we "white-boarded" problems from various worksheets in class. This consists of each table group being assigned a single problem to work through quickly, but thoroughly and presenting their work to the class.

This approach to students doing practice problems offers multiple advantages over the standard way of having each student work through every problem on their own. First, a significant amount of time is saved by doing this simply because each student does less work. One might jump to the conclusion that saving this time would only sacrifice the students' understanding of the material, but this is not the case. This is because of the presentation of work. While students present their work, others are encouraged to double-check their classmates work, which requires thinking through the process of the problem and getting the required practice, while excluding the manual calculation of numbers which students gain little from. Working in groups also increases the efficiency of getting students through parts they may get stuck on or make a mistake. Instead of spending time trying to remember something that is not there, or looking for a mistake, students are quickly aided by their peers to get them through while reminding them what area they may need to review. Lastly, students have an added positive motivation to focus on their own problem as to not make a mistake for the rest of class to see.

This week has also changed the way that I think about the numbers and units given to me in any particular problem. Previously, I had considered all given information to be finite, single pieces of information that can help to find other information. I now know that given information used to find the solution to a problem is part of a large web of values relating to the problem that can always be traversed with the right conversion factors and aided by units (see figure below)(see URL below for stoichiometry summary and practice problems). The idea that units should be used as a guide was also reinforced as we did many multiple-step problems in which you could easily get lost in.

I have observed many connections between the material we covered during this last week and the weeks before that. During the summer we learned how to verify that a chemical equation is balanced and change coefficients if necessary to follow the law of conservation of mass and the idea that atoms themselves are not changes in chemical reactions. These same principles are present in the stoichiometry we covered this week. Without following the law of conservation of mass, many of our calculations regarding limiting and excess reagents would be impossible.

Stoichiometry summary and practice:
http://s-owl.cengage.com/ebooks/vining_owlbook_prototype/ebook/ch3/Sect3-3-c.html

The general molar web connecting each value

Week 1

This past week in AP Chemistry we learned about molarity, or concentration of solutions, being the number of moles of a substance in a liter of given solution. This number can be used to calculate the mass of solute in a solution. Given the equation: M1(V1)=M2(V2), we learned how to calculate either the molarity (M) or volume (V) of a solution before (1) or after (2) dilution with the solvent, given the other three values. We used this equation when calculating the concentration of Blue #1 in each solution we made of a known ratio of stock solution and distilled water. Knowing this new concept of molarity and how to calculate it allowed us to develop a linear regression equation that related molar concentration to the absorbance of light by the solution. Absorbance is a measure of the percent of light of a particular wavelength impeded by the solution, related to transmittance. Transmittance is the percent of light of a particular wavelength allowed through a solution.

We were detailed on all of the aspects involved in a lab experiment. I learned that you must always follow the step-by-step process in order to succeed in the testing. This includes completing pre-lab and thinking through all aspects of the experiment before arriving in class on the specified day. During the lab, it was very apparent who had prepared before class and who had not, as seen by the many puzzled faces of those who watched others starting before beginning their own experiments. You must also stay organized and follow the rubric formatting in order for your lab to look professional and presentable. Neglecting this seems to run a higher risk of error. Your responses to the post lab questions

We also explored the ways in which you can identify the molar concentration of a solute in distilled water if there is only one solute, through the application of light. This requires the use of a colorimeter (see figure below) which can accurately measure the absorbance of light of a particular wavelength by the solution being tested. With this data you could find many important details about the solution, which can have many practical details outside of this controlled school experiment. These applications could include the measurement of dissolved minerals in lake or ocean waters of certain areas in which fluctuations could result in catastrophic changes to the life in that area.

We came to know most of the information in class from guides this week. To understand both the principles of light transmittance and absorbance and the instructions for the colorimeter we read through the Vernier Colorimeter manual, which detailed these by describing the units of each term, the use for each and how they relate to each other physically and mathematically. In class we also had discussions on how to calculate the concentration of the stock solution using beer's law and known absorbance and molar extinction coefficient. These calculations were carried throughout the entire experiment and were therefor very important to the success of the experiment as a whole.

Here is a reference for the use of notebooks throughout experimentation:
http://www.ruf.rice.edu/~bioslabs/tools/notebook/notebook.html

This is an overview of the mathematical relationships for light transmittance in solutions:
http://teaching.shu.ac.uk/hwb/chemistry/tutorials/molspec/beers1.htm

Vernier Colorimeter for measuring absorbance (Left) and standard lab notebooks (Right)