General Chemistry 2024.pptx

Clinical Biochemistry Kristina Hultén Ph.D. 1 2 3 4 5 6 Fundamental for all Biological Disciplines Atoms Chemistry Molecules Molecular biology...

Clinical Biochemistry Kristina Hultén Ph.D. 1 2 3 4 5 6 Fundamental for all Biological Disciplines Atoms Chemistry Molecules Molecular biology Biochemistry Cell biology Cells Pharmacology Physiology Pathology Tissues Organs Organ systems 7  Course content General Chemistry Organic Chemistry Acid-Base Chemistry The Cell Proteins, Enzymes, Vitamins, Term 1 Cell signaling DNA Metabolic Pathways Disorders Fed and Fasted State Term 2 8  Term 1 Interim Exam (quiz 1): Multiple choice + one bonus question (free write short answer). 50 questions, 1h  Term exam (quiz 2): Multiple choice + one bonus question (free write short answer). 50 questions, 1h  Term 2 Interim Exam (quiz 3): Multiple choice + one bonus question (free write short answer). 50 questions, 1h  Final exam: Comprehensive exam of all course material. Multiple choice + one bonus question (free write short answer). 100 questions, 2h Grading: Quiz 1-3: 60% of grade, Final exam: 40% of grade You need a B (≥70%) to pass this course. Blackboard: Lectures, syllabus, study questions Contact: [email protected] 9 will start at the very beginning… Amino acidsproteins Isotopes Nucleotides  DNA, RNA, Major and minor signal molecules Atoms: essential C,N, O, H are the Different Protons, elements main elements in elements neutrons, e- Biomolecules molecules Glycerol and FA form lipids Ions Carbohydrates 10 then continue through the metabolic pat Lippincott: Biochemistry 11  Fuel Oxidative Fuel Storage & Pathways Mobilization Pathway Cellular respiration Energy for work (ATP) Energy stored in CO2 H 20 biomolecules heat Energy for waste disposal Detoxification or Biosynthetic Waste 12 to EXAMPLES we will come across late Biochemistry/Meta bolism - Sodium (or other ion) imbalances - Acid/base disturbances - Vitamin deficiencies - Protein defects eg. CF, hemoglobinopathies - Amino-acid metabolic disorders - Diabetes, lipidemia, hypercholesteremia, obesity, metabolic syndrome 13 Marathon runners: These guys are on the last mile of their run. What do you expect is going on biochemically speaking at this point? (Yes- they look great despite the long run! but apart from that?) Wiki commons accessed 6/18/2018 14 15  These kids have been in the pool for a long time now. They have been under the water more than above water, in fact. What could be happening biochemically speaking in this situation? /www.bing.com/images/ Creative Commons accessed 06/18/2019 16 Principles in General Chemistry Kristina Hultén Ph.D. 17 Syllabus detailed objectives: Concepts of General Chemistry, Organic Chemistry and Acid-Base To understand the following basic chemistry terms: SI units, scientific notation, atomic number and atomic weight, electron configuration To discriminate between the following atomic bonds: covalent bond, polar covalent bond, ionic bond, carbon double-bond formation, saturated and unsaturated bonds and van der Waals forces To explain chemical reactions such as proton transfer, hydrolysis, condensation, substitution, addition, group transfer To differentiate between oxidation and reduction reactions To define hydrophilic and hydrophobic, lipophilic, 18 lipophobic Outlin e 1. Review the structure of atoms, bonds formation, solutions 2. Explain the concept of reduction-oxidation and other pertinent reactions 3. Introduce the concept of concentration measurements (molarity, normality, osmolality, tonicity) 4. Laboratory values 19 Atomic theory All matter is composed of atoms Atoms differ between elements In chemical compounds atoms are combined in specific ratios Chemical reactions change the combinations of atoms 20 Key features of an element Atoms - protons, neutrons, electrons Differences in outer electron shell population  differences in electron bonding All elements have a half-life; all are radioactive 21 Wikimedia.org 22  Most common elements of the human body: Oxygen 61% 46% in the Earth’s crust Carbon 23% Hydrogen 10% 0.1% in the Earth’s crust Nitrogen 2-3% Calcium 1% 4% in the Earth’s crust Phosphorous 1% McMurry et al. 23 Atoms: composition Atomic number: number of protons Mass number: sum of protons and neutrons Atomic weight (atomic mass units, amu) Mass number (sum of protons + neutrons) 6 3 Li Element Symbol Atomic number ( #protons) McMurry et al. 24  Gains or loses Neutron Gains or loses Proton Gains or loses Electron Becomes a different Becomes an isotope Becomes an ion element ( same element) (same element) 25 Electro ns Gain or loss changes the electrical charge of the atom Molecule configuration, communication Energy can be captured and transferred to other molecules by electrons on certain molecules 26 Questio n Two atoms are walking down the street. One atom says: “Uh oh, I think I just lost and electron” The other atom asks: “Are you sure?” Yeah, I am ________________ 27 28  Electron Configuration and the Molecular Orbital Theory Electrons are most likely to be found within a molecular orbital (MO) in a molecule, depending on its energy level. 1. Electrons are arranged in shells and subshells 2. The first shells are called 1(1 subshell, 2 orbits), 2 (2s [1 subshell, 2 orbits], 2p [3 subshells-6 orbits]) so 1,2,3 etc. or or K,L,M… 3. Orbitals are filled from the innermost shell and outwards. Each subshell receives one electron before the subshells fill up. 4. Electrons with greater energy levels will be found in orbitals further away from the nucleus. 5. Hydrogen has on electron in 1 orbit (incomplete)-wants one interaction to complete outer shell. Wikimedia.org29 orbitals fill from the innermost shell and outwards to higher ener depending on the # of electrons per specific element Wikimedia.org K (1s) L (2n) M (3n) Shell Element 1s 2s 2p 3s 3p 3d (2) (2) (6) (2) (6) (10) Subshell Hydrogen ↑ e- (1) Carbon (6) ↑↓ ↑↓ ↑ ↑ Nitrogen ↑↓ ↑↓ ↑ ↑ ↑ (7) Sodium ↑↓ ↑↓ ↑↓↑↓↑↓ ↑ 30  TIME-OUT Creative commons 1.Atom composition and charges 2.Electron configuration 3.Valence electrons and how they matter 31 Atomic Bonds 32  Electron bonding capacity between elements A molecule is formed by two or more atoms linked by a chemical bond. The total number of interactions depends on the valence electrons (ve-; outer shell electrons). Bonding capacity can be deduced from the periodic table Hydrogen (H), Potassium (K), Sodium (Na): 1 ve- Carbon (C): 4 ve- Nitrogen (N): 5 ve- Oxygen (O), Sulphur (S): 6 ve- The type of bond will depend on the electronegativity of the elements 33 Electronegativity Electronegativity is a measure of the attraction of an atom for the electrons in a chemical bond The higher the electronegativity of an atom, the greater its attraction for bonding electrons. Oxygen (O): 3.5 Unequal sharing Hydrogen (H): 2.1 Equal sharing Carbon (C): 2.5 Nitrogen (N): 3.04 34  http://dl.clackamas.edu/ch104-07/electron.htm The electronegativity generally increases as you go from left to right across the periodic table. It decreases as you go down the periodic table. 35 Covalent vs. polar covalent vs. ionic bond? If the electronegativity difference between atoms is < than 2, the bond is likely covalent. If the difference is less than ca. 0.4, the bond is most likely equal sharing of electrons if the difference is between ca. 0.4-2, the bond will likely be covalent with unequal sharing. For atoms that have an electronegativity difference > 2, the bond is usually ionic. 36 Covalent bond Two similar atoms share (2) electrons Same claim of electrons – Equal sharing H H H H Two hydrogen atoms A hydrogen molecule McMurry et al. 37 Polar covalent Two dissimilar atoms are joined One atom has a stronger pull - Unequal sharing δ+ δ- H-Cl McMurry et 38 al. ttps://chemistrytalk.org/ionic-vs-covalent-bonds/ Coordinate (dated) covalent bond: One atom provides both electrons in a shared pair. 39 Ionic bond An atom that completely gains or loses an electron becomes an ion Transfer of one or more electrons from one atom to another and attraction between these two ions will lead to ionic bond formation Ionic bonds are weak in water 40  Carbon double bond formation The double bond between the carbon atoms involves two pairs of shared electrons. The two pairs aren't equal. H H C-C H H Sigma pi (between C nuclei) pi: e- move above and below the plane of the molecule within (pink) orb 41  Saturated/unsaturated bonds Alkane Alkene Alkyne Saturated Unsaturat Unsaturat ed ed Eg. propane Eg. propene Eg. propyne H H H H H H H-C-C-C-H H-C-C=C-H H-C-C ≡ C-H H H H H H H Double bonds are unsaturated- we will talk about cis and trans unsaturated fatty acids in Term 2. For now, you need to differentiate between saturated and unsaturated bonds and the examples above. 42 Non-covalent bonds and other weak forces (Secondary bonding) Hydrogen bonding Van der Waals forces Hydrophobic attractions -In nature molecules will form weak interactions with other molecules -Multiple weak bonds or forces can cause strong interactions 43  Hydrogen bonds δ+H 2δ- O H H O δ+ H H O H H O H Lippincott: Biochemistry 44 Examples of hydrogen bonds Hydrogen bonding is responsible for ammonia's high solubility in water. Multiple hydroxyl groups provide many opportunities for hydrogen bonding leading to the high viscosities of glycerine and sugar syrup Many organic (carboxylic) acids form hydrogen-bonded dimers in the solid state. Adapted from: http://www.chem1.com/acad/webtext/states/water.html; Creative Commons 45  Hydrogen bonds that stabilize molecules DNA –hydrogen bonds Proteins-hydrogen bonds between bases between atoms on side chains DNA lectures Protein lectures Figures from Creative commons 46  TIME-OUT 1.Discuss the bonds - Covalent - Polar covalent - ionic - Saturated/unsaturated - Hydrogen bond 2.What is electronegativity? 47 48 Since the body consists mainly of water, all particles are dispersed in water-so let’s consider the properties of water.... 49  Water Each H share a pair of electrons with O Negative pole O has two unshared pairs of δ- electrons 4 pairs total δ- O more electronegative O than H H δ+ Polar covalent Up to 4 hydrogen bonds possible for one water H Positive pole molecule (two delta positive and two delta negative charges-each charge can form a hydrogen bond δ+ with a delta charge at another molecule) 50  Hydrogen bonds explain the following water characteristics: δ+H 2δ- O H H O Ice floats (more stableδ+ H bonds) H Strong surface tension O H High specific heat H O Solvent properties Water density decreases both when water is heated and when it freezes 51  Water density depends on temperature, as temperature will affect the movement and the stability of bonds. Water molecules move liquid gradually less when the temperature decreases. At 4°C – water has the greatest density. They are organized but still relatively close together Below freezing point, the molecules slow down, resulting solid in more steady hydrogen Wiki Commons bonds, more space between 52 Water: “The molecule of life” The most polar of all common liquids Hydrogen bonding: the electrostatic attraction between molecules of water resulting from the polarity Self-ionization 2 H2O  H3O+ + OH- Can act as both acid and base 53 Water content in humans Differs by age and by tissue. Usually represents 50-60% of the body weight in adults, 75% of the body weight in children Assuming 60% of the total body weight, most is found intracellularly and about 1/3 is found extracellularly (plasma, interstitial) Dissolves and transports compounds in the blood (ions, CHO, lipids, proteins, nucleic acids Medium for moving molecules into and throughout cellular compartments Separates charged molecules, dissipates heat and participates in chemical reactions 54 Distribution of water in the body 25 L 15 L Intracellul Extracellul ar fluid ar fluid 10 L Interstitia l 5L Blood 55 Three important reactions in biochemistry 1. Oxidation-reduction (Redox) reactions 2. Group transfer reactions 3. Hydrolysis reactions 56 1. Redox reaction Electrons are transferred from one atom to another Oxidation: Loss of one or more electrons Reduction: Gain of one or more Oxidation electrons A A+ + electron(s) Reduction 57  Oxidation Is Loss of electrons - Oxidizing agent causes oxidation by gaining those e-. It causes oxidation by undergoing reduction (oxidation number decreases) - Which is the oxidizing agent in the reaction below? -Which is oxidized? Reduction Is Gain of electrons - Reducing agent causes reduction by losing those e—it causes reduction by undergoing oxidation - Which is the reducing agent in the reaction below? Feis2+reduced? - Which + Cu2+  Fe3+ + Cu+ 58  OIL RIG Oxidation Is Loss of electrons Reduction Is Gain of electrons Fe2+ + Cu2+  Fe3+ + Cu+ Fe3+ + Cu+  Fe2+ + Cu2+ 59 One molecule will cause the oxidation by accepting electrons and one molecule will cause the reduction by giving up electrons: Oxidizing agent (molecule) Causes oxidation Gains one or more electrons Undergoes reduction Becomes more negative (less positive) Reducing agent (molecule) Causes reduction Loses one or more electrons Undergoes oxidation Oxidation number of atom increases 60  TIME-OUT 1. Identify the oxidizing agent (-it accepts electrons and is reduced) 2. Identify the reducing agent (-it loses electrons and is oxidized) Fe2+ + Cu2+  Fe3+ + Cu+ Oxidation Is Loss of electrons Reduction Is Gain of electrons 61  - We will come across many redox reactions in biochemistry - For purposes of organic redox reactions in aqueous solutions, it is best to consider atomic hydrogen as a proton fused to an electron. H = H+ + e- - One molecule we will come across in organic redox reactions is NAD+ /NADH - Which is the oxidized form? Which is the reduced form? NAD+ + H+ + 2 e-  NADH iscussed in this course: Redox potentials (E´0) describe the likelihood of a reaction going to the left or the right 62  Oxidation-Reduction in Biochemistry: Kreb’s Cycle: The Electron Transport Chain: Candeias, LP: http://iriaxp.iri.tudelft.nl/~scwww/candeias/bio-et/redoxf.html G. Parslow, University of Melbourne 63 2. Group transfer reaction: Example of a group transfer: A phosphoryl transfer reaction From phosphoryl group donor molecule (designated 'R2' in the figure below) to a phosphoryl group acceptor (designated 'R1'). Transfer of a single phosphate group from R2 to R1: 64 Transfer of phosphate group to alcohol (monophosphate) ATP ADP 65 3. Hydrolysis Hydrolysis (a water molecule is added and a bond is broken) …. More to consider: Condensation (bond formed, water released) Proton transfer: a hydrogen ion will move from one chemical to another (acid-base reactions). Substitution: atom or functional group in a chemical compound is replaced by another atom or functional group Addition: chemical reaction in which two or more molecules react to one larger molecule. Hydrophilic vs. hydrophobic characteristics Amphipathic molecule: molecule with both 66 What type of reaction is this? http://chemwiki.ucdavis.edu/Organic_Chemistry/Organic_Chemistry_With_a_Biological_Emphasis/Chapter_12%3A_Acyl_substitution_reactions/ Section_12.1%3A_Introduction_to_carboxylic_acid_derivatives_and_the_nucleophilic_acyl_substitution_reaction An exam question could look 67 68 Lecture recap:  Atom composition  Electronegativity, bonds  Water  Important organic reactions  If we dissolve particles (molecules in water we can get different concentrations of solutes.  Our body required stable concentrations of key ions to maintain homeostasis.  Concentration, Osmolality, Tonicity 69 Concentra tion Mole: A mole is the amount of a substance whose weight (g) equals the molecular weight (6.023x 1023 molecules) Molarity: moles of solute per litre of solution Normality (acids/bases): moles of H+ ions per litre of solution (acid) moles of OH- ions per litre of solution (base) - We use normality for acids and bases, because we want to reflect the number of hydrogen or hydroxide ions in solution as this defines the pH - Acids and bases can of course also be described in molarity (moles/L) 70 Concentration: Acids/Bases Example: Molarity Normality Molarity Normality HCl (H+) 1M 1N NaOH 1M 1N 0.2M 0.2N 0.2M 0.2N H3PO4(3H+) 1M 3N Al(OH)3 1M 3N 0.2M 0.6N 0.2M 0.6N Example: Peak stomach acid output has gastric acid concentration approaching 1N. [Gastric] enzymes are highly sensitive to acid concentration 71 Homeostasis (“similar condition”) Organism in homeostasis External change Internal Malfunction Internal Change Compensation to Restore Homeostasis Failure Success Disease/illness Wellness 72 Marathon runners Wiki commons accessed 6/18/2018 73 Fluid and Electrolyte Homeostasis Volume Osmolarity Concentration of individual ions pH of body fluids Integration among multiple systems (respiratory, cardiovascular, renal and behavioral) 74 Distribution of water in the body 25 L 15 L Intracellul Extracellul ar fluid ar fluid 10 L Interstitia l 5L Blood 75 Important ions of the body Cations: Anions: H+ Hydrogen Cl- Chloride K+ Potassium HCO3- Na+ Sodium Bicarbonate Ca2+ Calcium HPO42- Mg2+ Magnesium Phosphate SO42- Sulfate 76  Important ions/electrolytes Plasma and interstitial Intracellular: fluid: Cations: Cations: Most: K+, Mg2+, Most: Na+ Some: Ca2+ Some: Ca2+, Mg2+ Anions: Anions: HPO42- Most: Cl- , proteins, Proteins (anions) Some HPO42- , SO42- Some SO 2-, Cl-, HCO - 4 3 HCO3- This is not a complete list! 77 The blood contains/transports: Water WBC and platelets RBC Proteins (albumin, ab) Salts (ions) Nutrients, hormones Gases, waste materials Source: KnuteKnudsen CC-BY 3.0 Unported via Wikimedia Comm 78 Blood Analytes Proteins –hydration, nutrition and liver function -enzymes, albumin Electrolytes – “lytes”: Na+, K+, Cl- , HCO3- -hydration, acid-base balance, pH, osmolality and heart and muscle contraction Minerals- E.g. calcium, phosphorous, iron Serum Creatinine BUN (Blood Urea Nitrogen) Uric Acid 79 EXAMPLE: Laboratory values Don’t memorize for this course! Molecule Normal Plasma Concentration Na+ 135-145 mEq/L K+ 3.5-5.1 mEq/L Cl- 98-106 mEq/L HCO3- 22-29 mmol/L BUN 7-18 mg/dL Creatinine 0.6-1.2 mg/dL Glucose 70-115 mg/dL 80 Concentration: Ions Ions can be described in mEq/L An equivalent is the number of moles of an ion in a solution, multiplied by the valence of that ion m=milli =1/1000 Example: 1 mmole/L of Na+ =1 mEq 1 mmole/L of Ca2+ = 2 mEq 1mmole/L of Cl- = 1 mEq 81 82 smolarity 1) Molarity solution and Osmolality –concentration of one molecule in 2) Osmolarity and osmolality- combined concentration of all molecules in solution 3) Some molecules are in higher concentration they contribute more to the osmolarity (“osmotically active”) Osmolarity: mOsm/L Same concept but different units Osmolality: mOsm/kg In a diluted solution (mainly water) osmolarity and osmolality are virtually the same Clinically the values are reported in milliosmoles per liter (mOsm/L). 83 Osmolality is used to define the balance across cell membranes All physiological fluids can be tested for osmolality. Plasma (serum) and urine are most typically measured in the clinical setting. Hyponatremia (unexpectedly low serum sodium level) and hypernatremia, diabetes Kidney function – renal concentrating ability Reported in milliosmoles per liter. 84 Osmolality Serum Osmolality Calculation: 2 x Na+ + BUN / 2.8 + glucose / 18 Normal range in serum: 285-295 mOsm/L Normal range in urine: 800-1300 mOsm/L Hypo/hypertonic: 2 x Na+ + glucose / 18 85 Cell membrane permeability Semi-permeable Lipid bi-layer with external polar “heads” SMALL HYDROPHOBIC MOLECULES O , CO SMALL UNCHARGED 2 2 POLAR MOLECULES H2O, glycerol, ethanol LARGER UNCHARGED POLAR MOLECULES Amino acids, glucose, nucleotides IONS 86 Distribution of solutes in the body fluid compartments Extracellular environmentIntracellular environment Cl- Cl- Na+ Cl- ATP pump K+ K+ K+ Na+ Na+ K+ K+ Na+ Na+ X X Proteins*, phosphates*, glucose-6-phosphat X Charged molecules are trapped within the cell Cell membrane-semi impermeable 87 Water can move in and out of the cell! Consider the concentration balance (osmolality on inside vs outside of the cell) Water movement in or out of the cell can attempt to even out imbalances due to, for example, dehydration. Diffusion: Movement of a particle from a higher to a lower concentration within a defined space or, spreading out more widely/evenly Osmosis: Diffusion of water to even out [ion -K+ and Na+] concentrations (from 88  Osmotic pressure Osmotic pressure: The pressure required to stop water from moving from one compartment to another (stop the osmosis). Equals the pressure created by water moving across a membrane  Depends on the balance between intracellular and extracellular fluid and the balance between the number of solutes (where will water flow?)  At equilibrium, no water is moving across the membrane. “Osmotically active particle”: A particle/ substance (usually ion) that causes osmosis to occur (high concentration and cannot pass the membrane easily). 89 Which of these do you think could be an important osmotically active particle? a. Na+ b. Glucose c. Vitamin A d. Glycerol 90  Why is this important? If the intracellular fluid increases or decreases dramatically, this can affect intracellular metabolism (hormones, drugs etc.) Hyper osmolality occurs, for example, in untreated type II diabetes where blood glucose levels are very high Tonicity is our final concept for this lecture to describe where water is flowing (in our out of the cell) as aresult of the osmolality and if the cells will swell or shrink as a result. 91 RECAP: One solution: 1) Molarity –concentration of one molecule in solution (moles/L) 2) Osmolarity and osmolality- combined concentration of all molecules in solution (osmoles/L) Two solutions separated by a membrane: Now we have to consider concentration differences across a membrane (focus on molecules that cannot pass). 3) Tonicity (effective osmolality) 92 Tonicity describes concentrations relative to each other in 2 compartments, sharing a semipermeable membrane Isotonic: Same concentration of ions intra- and extracellularly Human blood cells are isotonic in relation to 0.9% NaCl (308 mOsm/l) Hypertonic: Higher concentration of ions in the extracellular fluid than inside of the cell Hypotonic: Lower concentration of ions in the extracellular fluid than inside of 93 Hypertonic solution, i.e. >0.9% NaCl Less water in relation to salt (higher salt concentration relative to the inside of the cell) mOsm/L increased Osmotic pressure will move water out from the cell. Cell will shrivel up. 94 Hypotonic solution, i.e. 145mEq/L hypernatremia) The diuresis improved Measured serum osmolality declined as well, but urine osmolality continued to be elevated. A serum PG level drawn on hospital day 4 measured 986 mg/dL, and the corresponding urine PG level was 3000 105 Question BUN and tonicity We will discuss next lecture together A patient has a plasma [Na+] of 98, a glucose of 100 and a BUN of 9. - What is his osmolality? - Is he hypo-, iso- or hypertonic? - Why? Serum Osmolality = 2 x Na+ + BUN / 2.8 + glucose / 18 Normal range in serum: 285-295 mOsm/L 106 Laboratory values (don’t memorize for this course) Molecule Normal Plasma Concentration Na+ 135-145 mEq/L K+ 3.5-5.1 mEq/L Cl- 98-106 mEq/L HCO3- 22-29 mmol/L BUN 7-18 mg/dL Creatinine 0.6-1.2 mg/dL Glucose 70-115 mg/dL 107 Question osmolality We will discuss next lecture together A patient with diabetic ketoacidosis has an increase in plasma glucose, which cannot move freely across cell membranes in the absence of insulin and therefore causes water to move from the cells to the ECF. What will happen to the cell? What will happen to measured serum [Na+]? 108 Organize the following elements in the order of highest to lowest electronegativity: A. H-hydrogen B. N-nitrogen C. C-carbon D. O-oxygen 109 Electronegativity Electronegativity is a measure of the attraction of an atom for the electrons in a chemical bond The higher the electronegativity of an atom, the greater its attraction for bonding electrons. Oxygen (O): 3.5 Hydrogen (H): 2.1 Equal sharing Unequal sharing Carbon (C): 2.5 Nitrogen (N): 3.04 110 Diabetes- to think about: Explain why it can be useful to measure C-peptide levels in an unconscious patient that is self administering insulin. Figures show insulin production in the pancreatic beta cells Lippincott Clinical Biochemistry 111 Summary Knowledge in general chemistry is necessary in order to understand biochemistry and related sciences. Highlights of important knowledge Self-study Study session Questions? 112 Further reading/watching Youtube Hyponatremia explained clearly Medcram Molarity vs. osmolarity (Khan Academy) http://www.youtube.com/watch?v=o_Bb43 LApog https://www.youtube.com/watch?v=00c3P wSbiys 113