VCU CHEM 403: Biochemistry Exam 1 Notes - Roesser PDF

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Professor: Roesser...

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Lecture 1 - 1/13 - Introduction and start of review for Chapter 1 Know prefixes:

Free amino acids do NOT come together to create polypeptides: Figure 1: Know that this picture is in the textbook but prof said it was wrong.

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Why?: Because this is an endergonic (positive delta G) reaction. ● Requires energy to build the polypeptide ● Entropy: The randomness - Entropy value has a HUGE decrease when building the polypeptide - which is unfavorable. - “Two amino acids have more entropy than one polypeptide” ● Enthalpy: tells how much heat and work was added or removed from the substance. - “NH2 is a pretty good nucleophile” - “Carboxylic acid is a bad electrophile because it has a negative charge” - How does the cell get around this?: They couple this unfavorable reaction with a favorable reaction What actually happens with amino acids: - It takes amino acids and small RNAs called tRNAs, and covalently links the two together to make aminoacyl tRNA doing this requires ENERGY - SO it COUPLES the unfavorable reaction with one that is favorable and RELEASES energy aka it hydrolyses ATP to AMP and two phosphates, which releases energy. Coupling these two together is what allows this reaction to move forward Figure 2: EDIT: It's supposed to amino-acetyl tRNA - ATP in the cycle is regenerated by oxidizing food 1

Figure 3: 2 amino acyl tRNA reacting to form an ester linkage which is a very good electrophile. This means that they will react exergonically aka spontaneously

Lecture 2 - 1/15 - Finishing Chapter 1/Beginning Chapter 2: Water Strong vs Weak forces: - Strong forces: chemical bonds, not easy to break a) Covalent bonds - sharing a pair of electrons. Energy of stabilization is 180 KJ/mol and higher b) Ionic bonds - he didn't talk about these but know its an example I guess. Transfer of electrons - Weak forces - 4 Types a) van der Waals interactions - Induced electrical interactions (aka dipoles) when e- (electron) clouds are overlapping because atoms are about ≤ 0.3 nm apart from each other. - They’re attractive interactions are very weak, we’re talking 0.4 KJ/mol to 4.0 KJ/mol for energy of stabilization b) Hydrogen bonds - unequal sharing of a Hydrogen atom (aka a proton) between 2 heteroatoms - Definition of heteroatoms: something that is highly electronegative a) Ex: O,N,S,Cl,F - We will be talking mainly about the underlined. Cl and F aren't really present in nature Figure 1: Hydrogen bond between H and N (denoted by the dotted line). Hydrogen bonds are longer and more stable than covalent bonds (H bonds energy of stabilization is 10-20 KJ/mol)

c) Ionic interactions - Opposite charges attract while like charges repel each other Figure 2: Example he showed in class. Glycine and Lysine have opposite charges. (I put a bond so you know they’re attracted to each other but lines typically mean covalent bonds.) 2

- This attraction could help hold together the tertiary structure of the protein d) Hydrophobic interactions - Based on water being a ubiquitous solvent in nature (meaning found everywhere) and is a polar solvent - If you put a nonpolar solute with a polar solvent (water) the solute tends to aggregate - Def of aggregate: Solute minimizes the contact between the polar solvent - This is the lowest energy, most stable situation a) The molecules in the middle of the aggregate are NOT reacting with water at all’ b) Think of hydrophobic tails being protected Figure 3: Example of aggregation

“Life depends on weak forces. Double stranded DNA is held together by weak forces. Two strands are not held together covalently. This is good because during replication the strands need to be separated, and with weka forces is fairly easy to do so” Enzymes are essential for life because they catalyze important reaction to make them useful for the cell 1) Most of the time enzymes are proteins - we’ll see later that is not always the situation 2) Enzymes are very powerful catalysts 3) They are very specific than we can be in a chemistry lab. This is a huge advantage 4) Require specific tertiary (3°) structure. These structures are held together by weak forces 5) Most enzymes are inactive above 60 °C, because the heat disrupts the 3° structure - Not always true, sometimes there are proteins that are active at 100°C when necessary (extremophiles) - Since life requires weak forces that limits the conditions that life is capable of existing in 3 classes of organisms 1) Eubacteria 2) Archaea bacteria 3

3) Eukaryotes “How many genes does it take to make a living organism?” - When thinking about this question it depends on what is considered living Figure 4: Organism complexity chart drawn in class Organism

Comment he made

# Genes

Virus

Not living because it has to take over a living cell to replicate

12

Mycobacteria

Once cause what you call consumption - cellular living organism They can only procreate in human lung tissue (tuberculosis), so they have a limited environment in which they could thrive in

525

E. coli

Able to live in many different environments

4,400

Yeast

Single cellular eukaryotes

6,600

C. Elegans

Multicellular organism - Small worms

19,000

Arabidopsis

Plant

27,000

Puffer fish

N/A

38,000

Lizard

N/A

70,000

Up until now it seems that the more complex an organism gets the more genes it has. This is because it has been one gene per one protein Human

Most complex but yet we don’t have that many genes. Why?

23,000

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Why: Evolution stopped creating more genes for more complex creatures, in the mammalian lineages. a) Most of our genes we create multiple proteins usually done by alternative RNA splicing Chapter 2: Chemistry of Water 2 take home messages for water 1) Water is a very unusual, unique solvent (He reviewed this twice, so remember this) - Water has structure. They like to hydrogen bond with another water molecule - Differences are caused by polarity of H2O - It not liner there is a 104° angle - this makes it polar - Hydrogen bond create higher boiling point and higher melting point - Has a high dielectric constant a) Definition of dielectric constant: Measurement of solvents ability to separate out ions b) Think when you put salt in water and it dissolves - Has a high surface tension - Liquid form is more dense than solid form - water is more dense than ice (ice floats on water) - There is a density maximum (4°C), which is when it is the most dense Lecture 3 - 1/17 - Chapter 2: Water More on water - Water wants to create hydrogen bonds a) “No matter if it’s liquid or gas hydrogen bonds are constantly being made and/or broken” b) Gaseous phase, molecules don't interact - this requires energy and a high temperature 4

Polar molecules in water - When we put molecules into water, hydration spheres are formed around the solute if they’re soluble a) This why water has a high dielectric constant b) Also why most solutes are very soluble in water Figure 1: Hydration spheres dissociating NaCl

Nonpolar molecules in water - When you put a nonpolar solutes in water they tend to aggregate, this is the basis of hydrophobic interaction (one of the weak forces we talked about) a) The point here is polar molecules interacting with nonpolar molecules is a high energy situation aggregation minimizes the contact between the nonpolar molecule and water - This happens because molecules in the middle of the aggregate are not contacting water at all, so it’s a lower energy situation Amphipathic/Amphiphilic molecules in water - These molecules have two parts of the molecule - part of the molecule is polar and the other part is nonpolar - A good example is fatty acids - one end has a carboxylic acid (negative charge, polar) the other hydrocarbon tail (very nonpolar) a) These will automatically form Micelles in water - Micell: 3D bubbles - on the outside there is a carboxylic acid group, which interacts well with water - on the inside there is a hydrocarbon tails, which are very nonpolar - This is the lowest energy interaction, this happens naturally - This is a hydrophobic interaction b) Micelles are why soaps work - If you washed your hands with just pure water you get rid of the polar dirt, but not the nonpolar stuff - The soap is micelles, which is part of nonpolar and polar. So the nonpolar stuff is able to get broken up with these Density maximum Figure 2: Most solvents vs water densities at different temperatures

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Density maximum and lakes - Why? a) The lake cools down from the top, as the water cools it becomes slightly more dense than all the other water b) This cycle continues until we reach 4°C - after this point the cold less dense water stays floating at the surface, and freezes c) Ice begins to form at the surface - its less dense and stays at surface d) Ice is a good insulator - this is why the lake doesn’t freeze all the way through Colligative properties of water - Definition: When you put solute into solvent it changes the solvents properties a) It doesn’t matter about the identity of the solute, it matters about the concentration of solute you put in b) Examples: freezing point depression, boiling point elevation, osmotic pressure Osmotic pressure is when the cell wants to even itself out Figure 3: Imagine a lipid bilayer with a bunch of Cl ions that can't get through the other side - water wants to get through the other side

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If there is too much osmotic pressure the cell can lice Conclusion of first unique property: polarity Ionization of water - If you have pure water to a very small extent it’ll disassociate to these free ions:

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There aren't many free protons free floating in solution because once H+ groups find water it bonds with it:

Hydrogen ion concentration in solution aka pH - We talk about this but it doesn’t actually mean that there are free floating protons Equilibrium constant concentrationof products - Equilibrium constant is always concentrationof reactants - Concentration is referred to with [X], so I’ll be using that from now on - also keep in mind that M (molar) is mol/L Equation 1: Equilibrium constant for water +¿ −¿ ¿ OH ¿ ¿ H ¿ ¿ K eq =¿ - You can’t predict this K value, you must do an experiment to find it out - It has been found that in pure water: a) [H+] = 1× 10−7 M b) [OH-] = 1× 10−7 M c) These have to be equal to each other because you can't get on without the other d) [H2O] = 55.5 M - If you put these numbers back into the equation Keq = 1.8 ×10−16 M a) What is a very low number like that mean? Which predominates reactants or products?: Reactants predominates, and it tells you that almost all of the water is in the H2O form - A much more useful equation is to ignore the [H2O] to get Kw (water autoionization constant) +¿ −¿ ¿ OH ¿ H ¿ K w =¿ −14 a) Kw will always equal 1× 10 M b) If you add H+ this number still stays, because when you add H+ it’ll rebond with OHc) If you know one you should be able to find the other pH +¿¿ H - pH is defined as −log ¿ - The higher the pH the lower the number (1 is very acidic, while 14 is very basic) pOH

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pOH is defined as

−¿ ¿ OH −log ¿ ALWAYS : pH + pOH =14

Strong vs weak electrolytes - Strong electrolytes: substances that completely dissociates into ions in water a) Examples: Strong acids (HCl), ionic compounds (NaCl, KBr), and strong bases (NaOH) - Weak electrolytes: substance that partially dissociates or doesn't dissociates at all in water a) Examples: Sugars (dissolves but does not form ions), alcohols (ethanol - dissolves but does not form ions), weak acids (acetic acid - CH3COOH), and weak bases - Why do we call them electrolytes? a) If you put an anode and a cathode attached to a light bulb, your circuit will light up - aka it conducts electricity - Pure H2O = light bulb does not light up - H2O with NaCl = light bulb does light up - If he tells you have 1 Molar of a strong electrolyte (HCl) - What is the [H+] in solution? - 1 Molar - because of complete dissociation - If you have 10 Molar of NaOH - What is the hydroxide (OH-) concentration? - 10 Molar - because of complete dissociation - If you were to have a 10 Molar solution of acetic acid - What is the [H+] in solution? −¿ +¿ ¿ H ¿ ¿ C H 3 CO O ¿ ¿ K a =¿ - Experimentally determined Ka value is 1.7 ×10−5 M - Which means at equilibrium almost 2 out of 20,000 molecules are going to be in the ionic form while the rest are in the acetic acid form a) Equilibrium leans towards the reactants side Lecture 4 - 1/22 - Chapter 2: Continuing of strong and weak electrolytes, and practice problems Weak electrolytes require Henderson Hasselbalch Equation −¿ ¿ +¿+ A : A weak electrolyte in its acid forms dissociates to H+ and the conjugate base A¿ HA ⇔H Figure 1: Henderson Hasselbalch Equation and using it to find the pH equation

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+¿ +¿ +¿ −¿ ¿ −¿ −¿ A ¿ ¿ A A ¿ ¿ ¿ {rearrange is that H+ is by itself} {take the log} {multiply by -1} [ HA] ¿ H ¿ K a [ HA ] ¿ ¿ H = ¿ ¿ ¿ H =log(K a)+log ¿ K a=¿ ¿ lo g ¿ +¿ −¿ ¿ A ¿ [ HA ] ¿ H ¿ =−log (K a)− log ¿ −log ¿ +¿¿ H : the equation for pH −log ¿ −log(K a) : the equation for pKa value −¿¿ A ¿ {flip the sign of the log} {Replace values with pH and pKa} [ HA ] ¿ pH= p K a−log¿ −¿ ¿ A ¿ - this is the final Henderson Hasselbalch Equation! ¿ ¿ pH= p K a +log ¿ Example problems Ex 1) What is the pH after 100 mL of 2 M NaOH is added to 100 mL of acetic acid and H 2O is added to make the final volume of 0.5 L? (pKa of acetic acid = 4.76) Steps: 1) Determine which one is in excess mol )=0.2 moles NaOH :(0.1 L× 2 L mol )=0.1 moles Acet ic acid :(0.1 L× 1 L - Here NaOH is in excess, so we won’t have to use the Hasselbalch equation 2) Determine amount of [OH-] −¿ ¿ 0.2 moles−0.1 moles=0.1 moles of O H 3) Determine M for [OH-] −¿ ¿ OH 0.1moles =0. 2 M ¿ 0.5 L 4) Find pOH 9

−¿ ¿ O H ⇒−log[ 0.2]=0.69 pOH =−log ¿ 5) Subtract from 14 14 −0.69=13.3 pH Ex 2) What is the pH when 100 mL of 1 M NaOH is added to 200 mL of 2 M acetic acid? (pKa of acetic acid = 4.76) Steps: 1) Determine which one is in excess mol )=0.1 mol NaOH :(0.1 L× 1 L mol )=0.4 mol C H 3 COOH=(0.2 L× 2 L - Here CH3COOH is in excess, meaning we have to use the Hasselbalch equation 2) Determine amount of [OH-] +¿ ¿ −¿+ H ¿ C H 3 COOH ⇔C H 3 O O −¿ ⇔H 2 O Every OH molecule is going to react to create water + ¿ + OH ¿ shifts from left to right ¿ H

0.4 CH3COOH and 0 CH3COO-0.1 OH Add 0.1 OH This converts acid form to conjugate base form

0.3 moles Acid form

+0.1 OH 0.1 moles Conjugate base

3) Plug into the Henderson Hasselbalch equation −¿¿ A ¿ ¿ pH = pKa+ log ¿ - Notice the (⅓) it’s a ratio, it doesn’t matter what the final volume is. If you put this ratio then you will get the right number every time. Titration Curves - How titrations work: you add a known amount of strong base while monitoring pH, so you can graph the pH as you add base - Figure 1: If you do the titration correctly then you should have a graph similar to the one below

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● In the middle of the figure you have half of the OH- equivalence the pKa of the weak acid Figure 2: Example of a titration curve that relevant. - Phosphoric acid has 3 H+ to give away. Be able to distinguish different parts of this graph, it’ll become important in future classes

Lecture 4 - 1/24 - Chapter 1,2,13: Buffers Buffers - Buffers are solutions that resist changes in pH ● Addition of acid or base causes little change in pH - Weak electrolytes make good buffers ● They’re only good if the pH is close to the pKa (+/- 1 unit) ● Ex: Acetic acid pKa = 4.76 in pH range of 3.76-5.76 will work as a buffer Explaining why acetic acid works as a buffer in 3.76-5.76 What happens if you add acid to a solution with a pH of 4.76?: It’ll be pushed to the left (3.76 pH) to establish equilibrium. Acid is going to combine with acetate to make more of the acetic acid - To resist changes when acid is added you need some acetate (conjugate base form) What happens if you add base to a solution with a pH of 4.76?: It’ll be pushed to the right (5.76 pH) to establish equilibrium. OH- reacts with H+ to make water. Acetic acid is going to break down to create OH- and acetate - To resist changes when adding base you need some acetic acid (acid form) To withstand acid/base addition you need a significant amount of both acid form and conjugate base form. This happens when the solution's pKa is close to pH value! Figure 1: Comparing ratios - Note these are logarithmic scales, and that the numbers are for the acetate only not for H+ +¿ ¿ −¿+H ¿ C H 3 COOH ⇒C H 3 CO O

pH

4.76

1

1

3.76

10

1

2.76

100

1

5.76

1

10

11

6.76

1

100

Titration curve - You want to do a pH of the buffer system to be in that plateau area (see figure 2 from previous lecture) Making buffers in lab There are 2 ways to do this (example making an acetic acid buffer [pKa = 4.76]) 1) Make a 0.1 M solution of acetic acid by pouring enough acetic acid into solution to make it that concentration - Initially the pH of this will be very low, so you have to titrate it with a base (NaOH) until you get to the pH of 4.76 - This will happen when you add 0.5 M equivalent of NaOH (so 0.05), - Here you know you have an equal amount of the acid and the conjugate base form. 2) Make it a 1 L solution. Add 0.5 moles of acetic acid and 0.5 moles of sodium acetate (Na-CH3COO+) - Strong electrolyte completely dissociates) - Here you have equal amounts of the acid and conjugate base form so if you measured it correctly you will have a pH of 4.76 Sample buffer problems Ex 1a: If 100 mL of 0.05 M NaOH is mixed with 900 mL of 0.1 M glycine buffer at pH 2.2 what is the final pH? (Note glycine has 2 pKas: pK1 = 2.2 ; pK2 = 9.4) Steps 1) Determine which one is in excess mol )=0.005 moles NaOH :(0.1 L× 0.05 L mol =0.09 moles 0.9 L× 0.1 L Glycin e :¿ - Glycine is in excess use H-H equation 2) Determine what form of glycine we’re using in the Henderson-Hasselbalch equation

- At 2.2 pH you’ll have 0.045 gly+ and 0.045 gly0 3) Determine end amount of glycine

At 2.2 pH you’ll have 0.045 gly+ and 0.045 gly0 -0.005 OH +0.005 OH Add 0.005 OH 0.04 moles 0.05 moles This converts gly+ to gly0 gly+ gly0 4) Plug everything into Henderson-Hasselbalch equation −¿ ¿ A ¿ ¿ pH = pKa+ log ¿ Ex 1b: If 100 mL of 0.05 M NaOH is mixed with 900 mL of H2O. What is the final pH? 12

mol )=0.005 moles L −¿ 0.005 moles ¿ OH = =0.005 M 1L ¿ pOH=−log (0.005 )=2.301 pH=14 −2.301=11.7 - Note the large change - jumped from 7 to 11.7 Biological Thermodynamics Definitions - System: Portion of the universe we’re measuring/concerned with - Surroundings: Everything else outside of the system 3 kinds of systems 1) Isolated systems - No transfer of either energy (E) or matter (m) may be transferred between system and surrounding - Preventing (completely stopping it) energy transfer is nearly impossible, so there is not many of these systems in nature - The only isolated system we have is the universe 2) Closed system - Energy (E) is transferred but matter (m) is not between systems and surrounding 3) Open system - Both energy (E) and matter (m) may be transferred between system and surroundings First Law of Thermodynamics - Definition: Total energy of an isolated system is constant or conserved ● Total energy in the universe is constant ΔE (change in E) is heat transferred at constant volume (V) vs ΔH (enthalpy) is heat transferred at constant pressure (P) - Can equate enthalpy by this equation Δ H =ΔE + PV Standard state - Denoted with ° - example: ΔH° = standard state of ...