Introduction into Biochemistry and Acids & Bases .pdf

Join the course’s Teams Introduction into Biochemistry  Course information Recommended textbooks Marks' Basic Medical Biochemistry: A Clinical Approach 5th Edition, by Michael Lieberman (Author), Alisa Peet MD (Author), 2018 Biochemistry 8th edition by Mary Campbell (Author) a...

Join the course’s Teams Introduction into Biochemistry  Course information Recommended textbooks Marks' Basic Medical Biochemistry: A Clinical Approach 5th Edition, by Michael Lieberman (Author), Alisa Peet MD (Author), 2018 Biochemistry 8th edition by Mary Campbell (Author) and Shawn Farrell (Author) Online: https://themedicalbiochemistrypage.org/ Instructors Prof. Mamoun Ahram Dr. Diala Abu Hassan  Outline Introduction Acids and bases, pH, and buffers Macromolecules Carbohydrates, lipids, and amino acids, peptides, and proteins Protein structure-function relationship part I: fibrous proteins: collagen, elastin, and keratins part II: globular proteins (plasma proteins, myoglobin, hemoglobin, and immunoglobulins) part III: Regulation of hemoglobin Enzymes structural features and classification, kinetics, mechanisms of regulation, cofactors Protein purification and analysis Biochemistry & chemical composition of living organisms  Biochemistry = understanding life Know the chemical structures of biological molecules Understand the biological function of these molecules Understand the interaction and organization of different molecules within individual cells and whole biological systems Understand bioenergetics (the study of energy flow in cells) Biochemistry in medicine: explains all disciplines diagnose and monitor diseases design drugs (new antibiotics, chemotherapy agents) understand the molecular bases of diseases  Chemical elements in living creatures The human body is composed mainly of ~30 elements. Four primary elements: carbon, hydrogen, oxygen, and nitrogen (96.5% of an organism's weight) Then, calcium and phosphorus (that’s 98.5%). Others exist in trace amounts but are essential, elements (mostly metals).  Important terms Hydrogen Electronegativity Oxygen Carbon Covalent bonds Nitrogen Polar vs. non-polar covalent bonds Single vs. multiple Non-covalent interactions Electrostatic interactions Hydrogen bonds (donor and acceptor) Van der Waals interactions Hydrophobic interactions Hydrophobic versus hydrophilic molecules Nucleophile vs electrophile  Important properties of bonds Bond strength (amount of energy that must be supplied to break a bond) Bond length: the distance between two nuclei Bond orientation: bond angles determining the overall geometry of atoms The three-dimensional structures of molecules are specified by the bond angles and bond lengths for each covalent linkage. Polarity of covalent bonds Covalent bonds in which the electrons are shared unequally in this way are known as polar covalent bonds. The bonds are known as “dipoles”. Oxygen and nitrogen atoms are electronegative Oxygen and hydrogen Nitrogen and hydrogen Not carbon and hydrogen Both water and CO2 contain polar bonds, but only water is a polar molecule.  What are non-covalent interactions? They are reversible and relatively weak. Electrostatic interactions (charge-charge interactions): They are formed between two charged particles. These forces are quite strong in the absence of water. Hydrogen bonds A hydrogen atom is partly shared between two relatively electronegative atoms (a donor and an acceptor). van der Waals interactions Unequal distribution of electronic charge around an atom changes with time. The strength of the attraction is affected by distance.  Hydrophobic interactions Self-association of nonpolar compounds in an aqueous environment Minimize unfavorable interactions between nonpolar groups and water  Properties of noncovalent interactions Reversible Relatively weak Molecules interact and bind specifically. Noncovalent forces significantly contribute to the structure, stability, and functional competence of macromolecules in living cells. Can be either attractive or repulsive Involve interactions both within the biomolecule and between it and the water of the surrounding environment 13  Carbon The road to diversity and stability Properties of carbon (1) It can form four bonds, which can be single, double, or triple bonds. Each bond is very stable. strength of bonds: triple > double > Single) They link C atoms together in chains and rings. These serve as a backbones.  Properties of carbon (2) Carbon bonds have angles giving molecules three-dimensional structures. In a carbon backbone, some carbon atoms rotate around a single covalent bond producing molecules of different shapes. The electronegativity of carbon is between other atoms. It can form polar and non-polar molecules. Pure carbon is not water soluble, but when Nonpolar carbon forms covalent bonds with other elements like O or N, the molecule that makes carbon compounds is soluble. Water  Properties of water (1) Water is a polar molecule as a whole because of: the different electronegativities between Hydrogen and oxygen It is angular. Water is highly cohesive. Water molecules produce a network. Water is an excellent solvent because It is small, and it weakens electrostatic forces and hydrogen bonding between polar molecules. Note Dipole-dipole interaction Dipole-charge interaction Properties of water (3) It is reactive because it is a nucleophile. A nucleophile is an electron-rich molecule that is attracted to positively- charged or electron-deficient species (electrophiles). Properties of water (4) Water molecules are ionized to become a positively-charged hydronium ion (or proton), and a hydroxide ion: Note: H3O+ = H +  Types of acids and bases Arrhenius acids and bases Acid: a substance that produces H+ when dissolved in water H+ Reacts with water-producing hydronium ion (H3O+). Base: a substance that produces OH- when dissolved in water.  Types of acids and bases The Brønsted-Lowry acid: any substance (proton donor) able to give a hydrogen ion (H+-a proton) to another molecule. Monoprotic acid: HCl, HNO3, CH3COOH Diprotic acid: H2SO4 Triprotic acid: H3PO3 Brønsted-Lowry base: any substance that accepts a proton (H +) from an acid. NaOH, NH3, KOH Water = amphoteric Substances that can act as an acid in one reaction and as a base in another are called amphoteric substances. Example: water With ammonia (NH3), water acts as an acid because it donates a proton (hydrogen ion) to ammonia. NH3 + H2O NH4+ + OH– With hydrochloric acid, water acts as a base. HCl+ H2O → H3O+ + Cl- Ampho = ‘both’ or ‘dual’ Join the course’s Teams Acid/base strength Acids differ in their ability to release protons. Strong acids dissociate 100%. Bases differ in their ability to accept protons. Strong bases have a strong affinity for protons. For multi-protic acids (H2SO4, H3PO4), each proton is donated at different strengths. Rule The stronger the acid, the weaker the conjugate base. Strong vs. weak acids Strong acids and bases are one-way reactions HCl → H+ + Cl- NaOH → Na+ + OH- Weak acids and bases do not ionize completely HC2H3O2 H+ + C2H3O2- NH3 + H2O NH4+ + OH- Equilibrium constant and Acid dissociation constant Acid/base solutions are at constant equilibrium. We can write equilibrium constant (Keq) for such reactions HA H++ A- Note: H3O+ = H + The value of the Ka indicates the direction of the reaction. When Ka is greater than 1 the product side is favored. When Ka is less than 1 the reactants are favored. What is pKa?  Molarity of solutions Solutions can be expressed in terms of its concentration or molarity. Moles of a solution are the amount in grams in relation to its molecular weight (MW or a.m.u.). moles = grams / MW A molar solution is where the number of grams equal to its molecular weight (moles) in 1 liter of solution. M = moles / volume (L) Since (mol = grams / MW), you can calculate the grams of a chemical you need to dissolve in a known volume (L) of water to obtain a certain concentration (M) using the following formula: grams = M x volume (L) x MW Acids and bases can also be expressed in terms of their normality (N) or equivalence (Eq).  Exercise How many grams do you need to make 5M NaCl solution in 100 ml (MW 58.4)? grams = 58.4 x 5 M x 0.1 liter = 29.29 g  Equivalents When it comes to acids, bases and ions, it is useful to think of them as equivalents. An equivalent is the amount of moles of hydrogen ions that an acid can donate. or a base can accept. A 1 g-Eq of any ion is defined as the molar mass of the ion divided by the ionic charge.  Examples For acids: 1 mole HCl = 1 mole [H+] = 1 equivalent 1 mole H2SO4 = 2 moles [H+] = 2 equivalents 1 eq of H2SO4 = ½ mol (because 1 mole gives two moles of H+ ions) Remember: One equivalent For ions: of any acid neutralizes one One equivalent of Na+ = 23.1 g equivalent of any base. One equivalent of Cl- = 35.5 g One equivalent of Mg2+ = (24.3)/2 = 12.15 g Molarity and equivalents Equivalents = n x M x volume (L) One equivalent of any acid neutralizes one equivalent of base. Based on the equation above, since x eq of an acid is neutralized by the same x eq of a base, then (n x M x vol) of an acid is neutralized by (n x M x vol) of a base. Problem 1 Note that each one produces 1 mole of the ions (H+ or OH-), so 1M of HCl is equal to 1M of NaOH. Eq of base = Eq of acid n x M1 x Vol1 = n x M2 x Vol2 1 x M1 x 12= 1 x 0.12 x 22.4 M1 = (0.12 x 22.4) / 12 M1 = 0.224 M  Problem 2 Note that 1 mole of HNO3 produces 1 mole of H+, but 1 mole of Ba(OH)2 produces 2 moles of OH-. In other words, the n is different. Also, remember that Equivalents = n x M x volume (L), where n is the number of charges or the number of H + (or OH-) the acid or base can produce or accept. Titration means that we an acid to a base slowly. At one point during titration, the acid and the base neutralize or cancel each other. In other words, “to titrate” means “to neutralize”. At the point of neutralization, the concentration of H+ is equal to the concentration of OH-. The best way to calculate how much acid is needed to neutralize a base (or the opposite) is to calculate the equivalents. Eq of acid = Eq of base N x M1 x Vol1 = n x M2 x Vol2 1 x 0.085 x Vol = 2 x 0.12 x 15 Vol = (2 x 0.12 x 15) / 1 x 0.085 Vol = 42.35 mL  Ionization of water Water dissociates into hydronium (H3O+) and hydroxyl (OH-) ions. For simplicity, we refer to the hydronium ion as a hydrogen ion (H+) and write the reaction equilibrium as:  Equilibrium constant The equilibrium constant Keq of the dissociation of water is: The equilibrium constant for water ionization under standard conditions is 1.8 x 10-16 M.  Kw Since there are 55.6 moles of water in 1 liter, the product of the hydrogen and hydroxide ion concentrations results in a value of 1 x 10-14 for: This constant, Kw, is called the ion product for water  + [H ] and - [OH ] For pure water, there are equal concentrations of [H+] and [OH-], each with a value of 1 x 10-7 M. Since Kw is a fixed value, the concentrations of [H+] and [OH-] are inversely changing. If the concentration of H+ is high, then the concentration of OH- must be low, and vice versa. For example, if [H+] = 10-2 M, then [OH-] = 10-12 M