SS1_2023_Bis2A_Facciotti_Reading_04

Learning Objectives Associated with SS1_2023_Bis2a_Facciotti_Reading_04 GC.21 Define and correctly use the law of conservation of mass. GC.25 Explain the role of water in condensation reactions and hydrolytic reactions. GC.16 Discuss the concept of reaction reversibility and its relationship to equilibrium. GC.17 Describe the equilibrium state.  GC.18 Relate the magnitude of the equilibrium constant to the equilibrium position. GC.19 Explain the effect of changing the concentration of a reactant or product on a system initially at equilibrium - apply Le Chatelier's principle. GC.35 Apply the concept of chemical equilibrium and the equilibrium constant to describe the progress of a chemical reaction, initially out of equilibrium, towards equilibrium, and finally at equilibrium in terms of “forward” and “reverse” reaction rates and concentrations of chemical reactants. GC.27 Define and correctly use pH and understand the relationship between pH and the hydrogen ion concentration. GC.20 Use the definition of pH to determine the difference in [H^+] concentration between two aqueous solutions.

 

Characteristic Chemical Reactions

Chemical reactions occur when two or more atoms bond together to form molecules or when bonded atoms break apart. We call the substances that "go in" to a chemical reaction reactants (by convention, we usually list these on the left side of a chemical equation), and the substances found that "come out" of the reaction products (by convention, we usually list these on the right side of a chemical equation). An arrow drawn between reactants and products typically shows the direction of the chemical reaction. By convention, for one-way reactions (a.k.a. unidirectional), reactants are listed on the left and products on the right of the single-headed arrow. However, you should be able to identify reactants and products of unidirectional reactions that are written in any orientation (e.g. right-to-left; top-to-bottom, diagonal right-to-left, around a circular arrow, etc.) by using the arrow to orient yourself.

In chemical reactions, the atoms and elements present in the reactant(s) must all also be present in the product(s). Similarly, there can be nothing present in the products that was not present in the reactants. This is because chemical reactions are governed by the law of conservation of mass, which states that matter cannot be created nor destroyed in a chemical reaction. This means that when you examine a chemical reaction, you must try to account for everything that goes in AND make sure that you can find it all in the stuff that comes out!

Just as you can express mathematical calculations in equations such as 2 + 7 = 9, you can use chemical equations to show how reactants become products. By convention, chemical equations are typically read or written from left to right. Reactants on the left are separated from products on the right by a single- or double-headed arrow indicating the direction in which the chemical reaction proceeds. For example, the chemical reaction in which one atom of nitrogen and three atoms of hydrogen produce ammonia would be written as:

\[\ce{N + 3H→NH_3}.\]

Correspondingly, the breakdown of ammonia into its components would be written as:

\[\ce{NH3→N + 3H.}\]

Note that in either direction, you find 1 N and 3 Hs on both sides of the equation.

 

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Possible NB Discussion Point

In General Biology courses, it is important to appreciate the law of the conservation of mass in the context of biological processes.  In chemistry, you will take a quantitative approach to this topic, learning to balance equations, and making sure that the total number of atoms and the total charge does not change.  In General Biology, we take a more qualitative approach to the topic. Do you think this leads to confusion? Should we place more emphasis on balancing equations in General Biology?

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Reversibility

While all chemical reactions can technically proceed in both directions, some reactions tend to favor one direction over the other. Depending on the degree to which a reaction spontaneously proceed in either both or one direction a different name can be given to characterize the reactions reversibility. Some chemical reactions, such as the one shown above, proceed mostly in one direction with the "reverse" direction happening on such long time scales or with such low probability that, for practical purposes, we ignore the "reverse" reaction. These unidirectional reactions are also called irreversible reactions and are depicted with a single-headed (unidirectional) arrow. By contrast, reversible reactions are those that can readily proceed in either direction. Reversible reactions are usually depicted by a chemical equation with a double-headed arrow pointing toward both the reactants and the products. In practice, you will find a continuum of chemical reactions; some proceed mostly in one direction and nearly never reverse, while others change direction easily depending on various factors like the relative concentrations of reactants and products. These terms are just ways of describing reactions with different equilibrium points.

Use of vocabulary

You may have realized that the terms "reactants" and "products" are relative to the direction of the reaction. If you have a reaction that is reversible, though, the products of running the reaction in one direction become the reactants of the reverse. You can label the same compound with two different terms. That can be a bit confusing. So, what is one to do in such cases? The answer is that if you want to use the terms "reactants" and "products", you must be clear about the direction of reaction that you are referring to - even for when discussing reversible reactions. The choice of terms, "reactants" or "products" that you use will communicate to others the directionality of the reaction that you are considering.

Let's look at an example of a reversible reaction in biology and discuss an important extension of these core ideas that arises in a biological system. In human blood, excess hydrogen ions (H⁺) bind to bicarbonate ions (HCO₃^-), forming an equilibrium state with carbonic acid (H₂CO₃). This reaction is readily reversible. If carbonic acid were added to this system, some of it would be converted to bicarbonate and hydrogen ions as the chemical system sought out equilibrium.

\[\ce{HCO_3^−+ H^+\rightleftharpoons H_2CO_3}\]

The example above examines and "idealized" chemical systems as it might occur in a test-tube. In biological systems, however, equilibrium for a single reaction is rarely reached as it might be in the test-tube. In biological systems, reactions do not occur in isolation. Rather, the concentrations of the reactants and/or products are constantly changing, often with a product of one reaction being a reactant for another reaction. These linked reactions form what are known as biochemical pathways. The immediate example below illustrates this point. While the reaction between the bicarbonate/proton and carbonic acid is highly reversible, it turns out that, physiologically, this reaction is usually "pulled" toward the formation of carbonic acid. Why? As shown below, carbonic acid becomes a reactant for another biochemical reaction—the conversion of carbonic acid to CO₂ and H₂O. This conversion reduces the concentration of H₂CO_3, thus pulling the reaction between bicarbonate and H⁺ to the right. Moreover, a third, unidirectional reaction, the removal of CO₂ and H₂O from the system, also pulls the reaction further to the right. These kinds of reactions are important contributors to maintaining the H⁺ homeostasis of our blood.

\[ \ce{HCO_3^- + H^+ \rightleftharpoons H_2CO_3 \rightleftharpoons CO_2 + H_20 \rightarrow} \text{ waste}\]

The reaction involving the synthesis of carbonic acid is actually linked to its breakdown into \(CO_2\) and \(H_2O\). These products are then removed from the system/body when they are exhaled. Together, the breakdown of carbonic acid and the act of exhaling the products pull the first reaction to the right.

Synthesis reactions

Many macromolecules are made from smaller subunits, or building blocks, called monomers. Monomers covalently link to form larger molecules known as polymers. Often, the synthesis of polymers from monomers will also produce water molecules as products of the reaction. This type of reaction is known as dehydration synthesis or condensation reaction.

Figure 1. In the dehydration synthesis reaction depicted above, two molecules of glucose are linked together to form the disaccharide maltose. In the process, a water molecule is formed. Attribution: Marc T. Facciotti (original work)

 

Interactive Figure 1. The molecules of glucose and maltose depicted as 3D interactive molecules.

Glucose	Maltose

 

In a dehydration synthesis reaction (Figure 1), the hydrogen of one monomer combines with the hydroxyl group of another monomer, releasing a molecule of water. At the same time, the monomers share electrons and form covalent bonds. As additional monomers join, this chain of repeating monomers forms a polymer. Different types of monomers can combine in many configurations, giving rise to a diverse group of macromolecules. Even one kind of monomer can combine in a variety of ways to form several different polymers; for example, glucose monomers are the constituents of starch, glycogen, and cellulose.

In the carbohydrate monomer example above, the polymer is formed by a dehydration reaction; this type of reaction is also used to add amino acids to a growing peptide chain and nucleotides to the growing DNA or RNA polymer. Visit the modules on Amino Acids, Lipids, and Nucleic Acids to see if you can identify the water molecules that are removed when a monomer is added to the growing polymer.

Figure 2. This depicts, using words, (decorated with functional groups colored in red) a generic dehydration synthesis/condensation reaction. Attribution: Marc T. Facciotti (original work)

Hydrolysis reactions

Polymers are broken down into monomers in a reaction known as hydrolysis. A hydrolysis reaction includes a water molecule as a reactant (Figure 3). During these reactions, a polymer can be broken into two components: one product carries a hydrogen ion (H⁺) from the water, while the second product carries the water's remaining hydroxide (OH^–).

Figure 3. In the hydrolysis reaction shown here, the disaccharide maltose is broken down to form two glucose monomers with the addition of a water molecule. Note that this reaction is the reverse of the synthesis reaction shown in Figure 1 above. Attribution: Marc T. Facciotti (original work)

 

Figure 4. This depicts using words (decorated with functional groups colored in red) a generic hydrolysis reaction. Attribution: Marc T. Facciotti (original work)

Dehydration synthesis and hydrolysis reactions are catalyzed, or “sped up,” by specific enzymes. Note that both dehydration synthesis and hydrolysis reactions involve the making and breaking of bonds between the reactants—a reorganization of the bonds between the atoms in the reactants. In biological systems (our bodies included), food in the form of molecular polymers is hydrolyzed into smaller molecules by water via enzyme-catalyzed reactions in the digestive system. This allows for the smaller nutrients to be absorbed and reused for a variety of purposes. In the cell, monomers derived from food may then be reassembled into larger polymers that serve new functions.

Helpful links:

Visit this site to see visual representations of dehydration synthesis and hydrolysis.
Example of Hydrolysis with Enzyme Action is shown in this 3 minute video entitled: Hydrolysis of Sucrose by Sucrase.

Exchange/transfer reactions

We will also encounter reactions termed exchange reactions. In these types of reactions, "parts" of molecules are transferred between one another—bonds are broken to release a part of a molecule and bonds are formed between the released part and another molecule. These enzyme-catalyzed reactions are usually reasonably complex multi-step chemical processes.

Figure 5. An exchange reaction in which both synthesis and hydrolysis can occur, chemical bonds are both formed and broken, is depicted using a word analogy.

 

 

Chemical equilibrium—Part 1: forward and reverse reactions

Understanding the concept of chemical equilibrium is critical to following several of the discussions that we have in BIS2A and indeed throughout biology and the sciences. It is difficult to completely describe the concept of chemical equilibrium without reference to the energy of a system, but for the sake of simplicity, let’s try anyway and reserve the discussion of energy for another chapter. Let us, rather, begin developing our understanding of equilibrium by considering the reversible reaction below:

Hypothetical reaction #1: A hypothetical reaction involving compounds A, B and D. If we read this from left to right, we would say that A and B come together to form a larger compound: D. Reading the reaction from right to left, we would say that compound D breaks down into smaller compounds: A and B.

We first need to define what is meant by a “reversible reaction.” The term “reversible” simply means that a reaction can proceed in both directions. That is, the things on the left side of the reaction equation can react together to become the things on the right of the equation, AND the things on the right of the equation can also react together to become the things on the left side of the equation. Reactions that only proceed in one direction are called irreversible reactions.

To start our discussion of equilibrium, we begin by considering a reaction that we posit is readily reversible. In this case, it is the reaction depicted above: the imaginary formation of compound D from compounds A and B. Since it is a reversible reaction, we could also call it the decomposition of D into A and B. Let us, however, imagine an experiment in which we watch the reaction proceed from a starting point where only A and B are present.

Example #1: Left-balanced reaction

Hypothetical reaction #1: time course

Concentration	t=0	t=1	t=5	t=10	t=15	t=20	t=25	t=30	t=35	t=40
[A]	100	90	80	70	65	62	60	60	60	60
[B]	100	90	80	70	65	62	60	60	60	60
[C]	0	10	20	30	45	38	40	40	40	40

At time t = 0 (before the reaction starts), the reaction has 100 concentration units of compounds A and B and zero units of compound D. We now allow the reaction to proceed and observe the individual concentrations of the three compounds over time (t=1, 5, 10, 15, 20, 25, 30, 35, and 40 time units). As A and B react, D forms. In fact, one can see D forming from t=0 all the way to t=25. After that time, however, the concentrations of A, B and D stop changing. Once the reaction reaches the point where the concentrations of the components stop changing, we say that the reaction has reached equilibrium. Notice that the concentrations of A, B, and D are not equal at equilibrium. In fact, the reaction seems left balanced so that there is more A and B than D.

Note: Common student misconception warning

Many students fall victim to the misconception that the concentrations of a reaction’s reactants and products must be equal at equilibrium. Given that the term equilibrium sounds a lot like the word “equal,” this is not surprising. But as the experiment above tries to illustrate, this is NOT correct!

Example #2: right-balanced reaction

We can examine a second hypothetical reaction, the synthesis of compound \(\ce{J}\) from the compounds \(\ce{E}\) and \(\ce{F}\).

\[ \ce{E +F <=> J} \nonumber\]

Hypothetical reaction #2: A hypothetical reaction involving compounds E, F and J. If we read this from left to right, we would say that E and F come together to form a larger compound: J. Reading the reaction from right to left, we would say that compound J breaks down into smaller compounds: E and F.

The structure of hypothetical reaction #2 looks identical to that of hypothetical reaction #1, which we considered above—two things come together to make one bigger thing. We just need to assume, in this case, that E, F, and J have different properties from A, B, and D. Let’s imagine a similar experiment to the one described above and examine this data:

Hypothetical reaction #2: time course

In this case, the reaction also reaches equilibrium. This time, however, equilibrium occurs at around t=30. After that point, the concentrations of E, F, and J do not change. Note again that the concentrations of \(\ce{E}\), \(\ce{F}\), and \(\ce{J}\) are not equal at equilibrium. In contrast to hypothetical reaction #1 (the ABD reaction), this time the concentration of J, the thing on the right side of the arrows, is at a higher concentration than E and F. We say that, for this reaction, equilibrium lies to the right.

Four more points need to be made at this juncture.

• Point 1: Whether equilibrium for a reaction lies to the left or the right will be a function of the properties of the components of the reaction and the environmental conditions that the reaction is taking place in (e.g., temperature, pressure, etc.).
• Point 2: We can also talk about equilibrium using concepts of energy, and we will do this soon, just not yet.
• Point 3: While hypothetical reactions #1 and #2 appear to reach a point where the reaction has “stopped,” you should imagine that reactions are still happening even after equilibrium has been reached. At equilibrium the “forward” and “reverse” reactions are just happening at the same rate. That is, in example #2, at equilibrium J is forming from E and F at the same rate that it is breaking down into E and F. This explains how the concentrations of the compounds aren’t changing despite the fact that the reactions are still happening.
• Point 4: From this description of equilibrium, we can define something we call the equilibrium constant. Typically, the constant is represented by an uppercase K and may be written as K_eq. In terms of concentrations, K_eq is written as the mathematical product of the reaction product concentrations (stuff on the right) divided by the mathematical product of the reactant concentrations (stuff on the left). For example, K_eq,1 = [D]/[A][B], and K_eq,2 = [J]/[E][F]. The square brackets "[]" indicate the “concentration of” whatever is inside the bracket.

 

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Possible NB Discussion Point

The following statement is true: A chemical equilibrium can be established starting with equal concentrations of reactants and products. Can you think of and describe other starting conditions for which a chemical equilibrium can also be established?  Are there any starting conditions for which a chemical equilibrium can NOT be established?

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The Role of Acid/Base Chemistry in General Biology

We have learned that the behavior of chemical functional groups depends on the composition, order, and properties of their constituent atoms. We will see that pH, a measure of the hydrogen ion concentration of a solution, can alter the chemical properties of some key biological functional groups in ways that change how they interact with other molecules and thus their biological role.

For example, depending on the pH, some functional groups on the amino acid that make up proteins can exist in different chemical states. We will learn that the chemical state of these functional groups can have a profound effect on the shape of the protein or on its ability to carry out chemical reactions. As we move through the course, we will see many examples of this type of chemistry in different contexts. 

In pure water, hydrogen ions are spontaneously generated by the dissociation (ionization) of a small percentage of water molecules into equal numbers of hydrogen (H⁺) ions and hydroxide (OH^-) ions. The OH^- that result from the ionization of water departs into the sea of water molecules interacting with other molecules through polar interactions, while the now "free" (unbonded) H⁺ ions produced by the ionization associates with water molecules (line two of the figure below) to create a new molecule called a hydronium ion, H₃O⁺.  At some point, the hydroxide ion from line 1 in the figure below will rejoin with a proton and reform another water molecule.  This process of dissociation and re-association between hydroxide and hydrogen ions happens continuously at equilibrium. 

While most H⁺ ions in solution really exist as H₃O⁺ ions, we usually represent the H₃O⁺ in figures or equations more simply as H⁺. Why?  Because it is easier. Just remember that since nearly all chemistry in biology happens in water that when you see H⁺ referred to in the text, figures, or in equations, it usually represents H₃O⁺.

Figure 1: Water spontaneously dissociates into a proton and hydroxyl group. The proton will combine with a water molecule forming a hydronium ion.
Attribution: Marc T. Facciotti

 

While some paradoxes to this rule can be found in the chemistry of concentrated solutions, in General Biology, it is convenient to formally define pH as:

\[ pH = -\log_{10} [H^+]\]

In the equation above, the square brackets surrounding [H⁺] indicate concentration. If necessary, try a math review at wiki-logarithm or kahn-logarithm. Also see: definition-concentration or wiki-concentration.  The pH of a solution is therefore a measure of the concentration of hydrogen ions in a solution (or the number of hydronium ions).  

The pH is typically reported on a logarithmic pH scale that ranges from 0 to 14 (Figure 2). We define pH=7.0 as neutral. We call anything with a pH below 7.0 acidic and any reported pH above 7.0 alkaline or basic. Extremes in pH in either direction from 7.0 are often considered inhospitable to life, although examples exist to the contrary. pH levels in the human body usually range between 6.8 and 7.4, except in the stomach where the pH is more acidic, typically between 1 and 2. Some microbial species like Sulfolobus acidocaldarius thrive in hyper acidic environments (pH < 3) while others like Natronomonas pharaonis have been found living in lakes with pH > 11. These organisms are classified as "extremophiles" for their abilities to thrive in extreme environments. Proteins from these organisms are sometimes used in industrial processes where their ability to withstand environmental stress is a valued property.  

Figure 2: The pH scale ranging from acidic to basic with various biological compounds or substances that exist at that particular pH. Attribution: Marc T. Facciotti

 

For Additional Information

Watch this video for an expanded explanation of pH and its relationship to [H+] and the logarithmic scale.

 

Let's work out an example to see how the pH scale works.

For reference: 1 mole (mol) of a substance (which can be atoms, molecules, ions, etc.), is defined as being equal to 6.02 x 10²³ particles of the substance. Therefore, 1 mole of water is equal to 6.02 x 10²³ water molecules.

Mathematically this can be written as:
1 mol = 6.02x10²³ particles in a substance
1 mol H₂O = 6.02x10²³ water molecules

Example: The concentration of hydrogen ions dissociating from pure water is approximately 1 × 10^-7 moles H⁺ ions per liter of water. The pH is calculated as the negative of the base 10 logarithm of this unit of concentration. The log₁₀ of 1×10^-7 is -7.0, and the negative of this number yields a pH of 7.0 (neutral pH).

Mathematically this can be represented as:
 pH = -log₁₀[H⁺]
pH = -log₁₀[1×10^-7] 
pH = 7.0 (neutral pH)

The figure below provides another way to visualize the inverse relationship between proton and hydroxide ion concentrations by graphically illustrating how proton concentration decreases as pH increases while the hydroxide ion concentration simultaneously increases. 

 

 

   Figure 3:  A graphical representation of acidity and basicity.  This figure illustrates the relationship between H⁺ and OH^- concentrations on the pH scale.  At low pH values H⁺ ions are plentiful.  As the pH increases the relative abundance of OH^- ions increase while H⁺ abundance decreases.

                                                                                                            Attribution: Mary O. Aina

The inverse relationship between pH and the concentration of protons confuses many students - take the time to convince yourself that you "get it." One way could be to predict whether different pH values are acidic or basic and then do the calculations to make sure. Start by trying these practice questions.

Knowledge Check Quiz

 

 

 

Acids and Bases

Acids and bases are molecules that can influence the pH of a solution. In General Biology it is often convenient to use the Brønsted-Lowry definition of acids and bases.  Using this formalism we define:

Acids = molecules that can donate a proton to another molecule (including water to form a hydronium ion)
Bases = molecules that can accept a proton from another molecule (including hydronium ions)

When protons from acidic molecules dissociate from their "parent" they increase the H⁺ concentration and thereby lower the pH of the solution. By contrast, when a base absorbs a "free" proton from a solution onto the "parent" molecule, the decrease in proton concentration in solution results in a shift to higher pH values.

Generically we can represent acids and bases as follows:

 

Figure 4:  Generic Acids and Bases.  This figure shows the behavior of Brønsted-Lowry acids and bases. The acid (A in a light purple circle) starts in a protonated form bound to an H⁺ ion, drawn as a red H. The acid deprotonates, shedding its H⁺ into solution or to another molecule. Meanwhile the base (B in a light green circle) begins deprotonated and absorbs a proton (red H⁺⁾ from solution or other molecule.  
Attribution: Marc T. Facciotti

 

In the figure above, the molecule A^- - the deprotonated form of the acid AH - can also be referred to as the conjugate base of the acid AH. Likewise the molecule BH⁺ - the protonated form of the base B - can be referred to as the conjugate acid of the base B.  

We call acids that completely dissociate into A^- and H⁺ ions at equilibrium strong acids. These reactions are characterized by an equilibrium position that lies far to the right (favoring product formation) and their chemical equations are often drawn with a single arrow separating reactants and products. By contrast, acids that do NOT completely dissociate into A^- and H⁺ ions at equilibrium are called weak acids. Depending on the pH, it is common to find both protonated and deprotonated forms of the weak acid (or both the acid and it's conjugate base) in solution at the same time. The chemical equations representing these reactions are therefore usually depicted with double arrows, indicating that the protonation/deprotonation of A^-/AH, respectively, is reversible.  

Two important examples of weak acids/bases in biology are the carboxyl and amino functional groups. At physiological pH values (around pH = 7) the carboxyl group tends to behave as an acid by donating it's proton to solution or other molecules. Under the same conditions, the amino group tends to act as a base, absorbing protons from solution or other molecules. As we will soon see, these and other protonation/deprotonation reactions play key roles in many biological processes.  

 

Figure 5: The carboxylic acid group acts as an acid by releasing a proton. This can increase the number of protons in solution and thus decrease the pH. The amino group acts as a base by accepting hydrogen ions, which can decrease the number of hydrogen ions in solutions, thus increasing the pH.
Attribution: Marc T. Facciotti (original work)

                                                                                                                                                                                                                                                                                                                                                                                            

   

Additional pH resources

Here are some additional links on pH and pKa to help learn the material. Note that there is an additional module devoted to pKa.

ChemLibreText Links

• Determining and calculating pH
• pH and pKa 
  • This has an expanded discussion about pH that is a bit more detailed than how we present the concept.

Khan Academy Links

• What is pH
• strong acids and bases
• weak acids
• weak acid-base equilibria
• pH and pKa

Simulations

• Acid-base simulation. -------
  • ----------embed this in document---------

Extra pH Links

• https://chemclass-ol.org/logarithms-and-ph
• https://web.mst.edu/~gbert/logs/pH.html
• http://www.biology.arizona.edu/biomath/tutorials/applications/pH.html
• Intro to acids, bases, pH

 

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PRACTICE POST-GUIDE

General Practice

3. Why: The concept of equilibrium underlies many other concepts in biochemistry and is therefore critical to get a good grasp of early. Heck! It’s crucial for understanding the reversible dissociation of water we discussed in the pH lecture/reading. For the many students who struggle with the concept of equilibrium, we’ve found that having them work on developing a mental picture involving MANY molecules often helps them “get it”. One idea that’s sometimes hard to connect with is the idea that chemical reactions are dynamic - reactions happen in time.

Another thing that makes equilibrium hard to visualize is the disconnect between the way we write chemical equations and how we think about equilibrium. Chemical equations are written to symbolically represent a single example of a chemical reaction not describe all individual molecules in a container.

For example, the imaginary equation 3S + Z <—> ZS₃ describes what happens to 3 molecules of S when they meet a molecule of Z; they make on molecule of ZS₃. The concept of equilibrium asks us to imagine that this is just one of many such reactions that can take place in a container. To understand equilibrium we must imagine that there might be 1 billion molecules of S and 1 billion of Z in the container and that the reaction 3S + Z <—> ZS₃ might happen hundreds of thousands of times. In addition, we must be ok with the idea that the reaction can also go “backwards”; sometimes ZS₃ falls apart into Z and three molecules of S.

The last point of confusion to work through is the idea that the forward reaction - in our example 3S + Z —> ZS₃ - might happen at a different rate than the reverse reaction, ZS₃ —> 3S + Z.  The forward reaction might happen faster than the reverse reaction or vice versa depending on the reaction.  That’s a lot to put together all at once.

How to practice: Load the following web simulation <http://billvining.com/mmlib_sims/#gen_14_5>.  This is a simulation of a chemical reaction in which a red molecules converts to blue and the blue molecule can revert back to blue. I recommend setting Temperature to High and Number of Spheres to 10.

The reaction might be written Red <—> Blue.

1. Select Equilibrium Constant K > 1 (recall that K is related to the ratio of Product/Reactant).  Push “Reset”.  At the very start all molecules should be red.  Follow the reaction over time by looking at the dynamic graph on the right.  What is happening over time? Which color molecule is dominant in the box?  Does the reaction ever stop?   What do you think would change if you could start with 1 million red molecules in the box?
2. Next select Equilibrium constant K = 1.  Push “Reset”.  Again, at the beginning you should have all red balls.  What happens as the reaction proceeds?  Which molecule is dominant in time?
3. Finally, select Equilibrium Constant K < 1.  Push “Reset”.  You should start with all red molecules.  Let the reaction go and follow it.  What happens as the reactions proceeds?  What molecule is dominant after some time?

Now for the fun of it, go to <http://billvining.com/mmlib_sims/#gen_14_3>.   This simulation allows you to understand what happens when a reaction has reached equilibrium and it is perturbed by the addition or removal of either a substrate or product.

1. Note the reaction A + B <—> C + D.  Also note the equilibrium constant expression the product or product concentrations divided by the product of reactant concentrations and that the value of K is 4.0.  From the simulation above, and math, can you already predict whether products or reactants will be in greater abundance at equilibrium?
2. Push “Equilibrate”.  This will simulate the reaction going to equilibrium and calculate the equilibrium concentrations of reactants and products. Remember from the simulation above that even when the curves are flat and the concentrations aren’t changing that the reaction is still happening.  There are many A and B molecules colliding and making C and D and the reverse reaction also happening.
3. Here’s the cool thing about this simulation! Pushing Equilibrate the first time simulates the initial reaction reaching equilibrium. Now you can play with adding or subtracting reactants and products from the first equilibrium and seeing how the reaction adapts to the change you made.  Start easy and change one thing. Lets subtract some product D.  Slide the D slider to the left about half way to the zero point.  Before you push equilibrate again, try to predict what will happen to the concentrations of the three other molecules, A, B, and C. Now push Equilibrate. What happened to the amount of reactants A and B?  Why?  What happened to product C?  Why?  Reset the simulation and continue to play with this until you feel like you can reason what’s happening out. Ask your TA or instructor if you are still trying to figure this out.

This exercise is intended to help build a mental picture and intuition that can help you master the first five learning goals on the pre-class guide for this class. See if you can see and stretch this model to make it fit each of those learning objectives.

 

Practice for pH

4. Why: Since the [H⁺] is critically important to the function of so many biomolecules, it is important to understand the relationship between hydrogen ion concentration and the more common unit we use to report it, the pH. Although you will not be asked to do any complicated calculations on the exam, you should explore the relationship between pH and the [H⁺] by solving for [H⁺] at different arbitrary pH values. It is important that you understand that pH is a logarithmic scale and that a change of 1 unit on the pH scale corresponds to a factor 10 change in hydrogen ion concentration. You should be able to do easy calculations without a calculator such as determining the hydrogen ion concentration at each whole number pH value.  You should be able to say how much the hydrogen ion concentration has changed when the pH changes by a certain number of whole units.

How to practice: For example, what is the H⁺ concentration at pH=1? pH=2? What is the difference in H⁺ concentration between pH 6 and pH 10?  You can practice this skill yourself by creating many similar questions with different numbers. Make up 10 different questions like this for yourself and convince yourself that you understand the relationship between pH and [H⁺]. Test your friends with a few.

Exercises 4 and 5 are practice for learning goals: “GC.27 Define pH and understand the relationship between pH and the hydrogen ion concentration.”; “GC.20 Use the definition of pH to determine the difference in [H⁺] concentration between two aqueous solutions.”

 

5.  Why: As mentioned in the rationale for question #4 understanding how to work with pH and converting between pH and [H⁺] is important for dealing with a lot of biology that happens at the molecular scale.  Getting a good grasp on this skill and developing some intuition is key.

How to practice: The pH of blood is usually slightly basic with a value of pH 7.365 and standardly rounded up to 7.4. Although you will not be responsible for this type of calculation on the exam, it would be good practice to use a calculator to answer the following questions more precisely.

1. What is the hydrogen ion concentration of blood at pH 7.4?
2. Now compare this hydrogen ion concentration for blood at pH 7.4 to the hydrogen ion concentration at pH 7.0. Which value is larger? By how much?
3. An overall shift of just 0.1 pH units in the blood can have a profound impact on human health. What does this apparently small change in pH mean in terms of the change in the hydrogen ion concentration?
    

PRACTICE EXAM QUESTIONS

Question Q4.1

Q4.1:  A class of antiviral drugs (highlighted by the commercial product Tamiflu) work by binding to a viral enzyme. This interaction inhibits the viral enzyme's function,  thereby lowering viral growth. The figure to the right shows the structure of one such drug and the amino acids on the protein that it interacts with. While we haven’t covered proteins yet, you can assume that they are composed of amino acids (like the ones shown in the figure to the right—this is really an application of identifying functional groups and predicting types of bonds they might form together). The following are possible ways to correctly describe the interaction between the drug and its target:

1. Covalent bonds are formed between atom #2 on the drug and atom #5 on the amino acids.
2. Ionic bonds are formed between atom #3 on the drug and atom #7 on the amino acids.
3. Hydrogen bonds are formed between atoms #3 on the drug and atom #4 on the amino acids.
4. Hydrogen bonds are formed between atoms #3 on the drug and atom #6 on the amino acids.
5. Ionic bonds are formed between atom #1 on the drug and atom #6 on the amino acids.

Question Q4.2

Q4.2  The optimum soil pH for growing strawberries is 6.5, whereas the optimum soil pH for growing blueberries is 4.5. Therefore, blueberries need______________ the number of hydrogen ions than strawberries to grow optimally.

1. 2x more
2. 2x less
3. 10x more
4. 10x less
5. 100x more
6. 100x less
7. It is not possible to answer this question without a calculator.