Chemistry in Living Systems. By Dr. Carmen Rexach Physiology Mt SAC Biology Department

Transcription

1 Chemistry in Living Systems By Dr. Carmen Rexach Physiology Mt SAC Biology Department

2 Matter and Energy Definitions Types of energy Kinetic vs. potential Forms of energy Chemical Ex: ATP Electrical Ex: Action potential of an neuron Mechanical Ex: Action of muscles in moving body Radiant (Electrochemical) Ex: Light energy

3 Atoms Smallest unit of matter composed of subatomic particles protons, neutrons, electrons atomic mass = #protons + # neutrons atomic number = # protons Elements = every specific type of atom

4 Atomic composition of the body 112 elements, 92 naturally occurring 24/112 are essential 7 essential minerals Ca, P, K, S, Na, Cl, Mg 13 essential trace elements Fe, I, Cu, Zn, Mn, Co, Cr, Se, Mo, F, Sn, Si, V 99% of the body s atoms Carbon Hydrogen Oxygen Nitrogen

5 Ions Atom that has gained or lost one or more electrons = electrolytes anions vs. cations H, minerals, trace elements

6 Isotopes Elements with a different number of neutrons Radioisotopes Instability and disintegration of atomic nucleus Occurs in heavier isotopes Half-life = amount of time required for ½ of radioactivity to be lost as isotope disintegrates Applications Radioactive tracers Radiation therapy for cancer

7 Thyroid Scan Measures health of thyroid by detecting radioisotope iodine-131 taken up by thyroid gland normal thyroid enlarged cancerous

8 Free radicals Atoms containing a single unpaired electron in the outermost orbital Can bond to either fill orbital or create free radicals in the donor molecule Formed by specific enzymes Highly reactive Begins process of pathogen destruction by white blood cells

9

10 Valence Electrons Shells surround electrons: First shell can contain only 2 electrons. Second shell can contain 8 electrons. If more than 2 electrons, must occupy shells more distant. Valence electrons: Electrons in the outer most orbital that participate in chemical reactions (if orbit incomplete). Form chemical bonds.

11 Chemical bonds Chemical bonds: Interaction of valence electrons between 2 or more atoms. Number of bonds determined by number of electrons needed to complete outermost shell.

12 Chemical bonds link atoms together to produce molecules Covalent Polar covalent Ionic Hydrogen van der Waals forces strongest weakest

13 Covalent bonds Atoms share electrons Most prevalent in organic compounds Non-polar = no charge H H

14 Polar covalent bonds δ+ δ+ H H O δ- Unequal sharing of electrons Result = positive and negative poles H 2 O Makes hydrogen bonds possible Oxygen, nitrogen, phosphorous have tendency to pull electrons towards themselves.

15 Ionic bonds Complete transfer of one or more electrons from one atom to another form ions when dissociate Na Cl Na + Cl -

16 Hydrogen bonds Electrical attraction between hydrogen in one polarized bond and oxygen or nitrogen on another or in the same molecule Surface tension, cohesion

17 Van der Waals forces Very weak attractions between nonpolar regions of molecules important in protein structure and structure of other large molecules Caused by hydrophobic interactions polar groups turn outward toward aqueous solution nonpolar groups turn inward towards each other

18 Solutions Solutes and solvents Solubility Hydrophilic Hydrophobic Amphipathic

19 Solubility Glucose, amino acids, are H 2 0 soluble. Hydration spheres form around atoms of oxygen, nitrogen, phosphorous. Charged complex ions and their cations form hydration spheres Hydrophilic molecules Molecules composed of nonpolar covalent bonds are not H 2 0 soluble. Cannot form hydration spheres. Hydrophobic molecules

20 Hydration spheres and glucose Hydrophilic interaction

21 Lipids in water Hydrophobic interaction

22 Solutions Concentration Amount of solute present per unit volume of solution (g/l) Moles = amount of a compound in grams equal to its molecular weight Ex) 180g of glucose in 1 L of solution = 1 Molar solution of glucose (1 mol/l) Small volumes in human body mmol/l = mol/l; μmol/l = mol/l; nmol/l = mol/l, etc.

23 Acids, bases, and ph Acids = proton donors Bases = proton acceptors ph = -log [ H + ] ph scale 0 (most acidic) to 14 (most basic) more free H + in solution = lower ph = more acidic!!!

24

25 Buffers System of molecules and ions that act to prevent changes in [H + ]. Stabilizes ph of a solution. In blood: H C0 2 H 2 C0 3 H + + HC0-3 Reaction can proceed in either direction (depending upon the concentration of molecules and ions).

26 Types of chemical reactions Anabolic reactions endergonic Catabolic reactions exergonic Coupled reactions Exchange reactions Oxidation-reduction reactions Oxidation = giving up electrons Reduction = accepting electrons

27 General characteristics of Molecules that contain carbon and hydrogen. Carbon has 4 electrons in outer shell. Carbon covalently bonds to fill its outer shell with 8 electrons. Tend to bond together to form large molecules = polymers Usually made up of subunits or monomers organic molecules

28 Dehydration synthesis & hydrolysis H OH CH 2 OH O H OH H H H OH OH CH 2 OH O H OH H H OH H OH CH 2 OH H H O OH H OH H H O CH 2 OH O H OH H H OH H OH H H OH OH Dehydration synthesis = removal of one molecule of water for each bond formed Sometimes called condensation Hydrolysis = addition of one water molecule for each bond broken

29 Functional groups Inactive backbone to which more reactive atoms are attached Examples: Carbonyl group: Aldehydes and ketones. Carboxyl group: Organic acids Ex) lactic and acetic acids Hydroxyl group: Alcohol.

30 Stereoisomers Exactly the same atoms arranged in same sequence. Differ in spatial orientation of a functional group. D-isomers: right-handed. L-isomers: left-handed. Ensure enzymes cannot combine with wrong stereoisomer. Enzymes of all cells can combine only with the L-amino acids and D-sugars.

31 Four categories of organic molecules Carbohydrates Lipids Proteins Nucleic acids

32 Carbohydrates: pentoses vs hexoses Pentoses are five carbon sugars Hexoses are six carbon sugars H C H C H C H C CH 2 OH RIBOSE O OH OH OH H C O HO C H H C OH HO C H HO C H CH 2 OH L-GLUCOSE

33 Carbohydrates General formula: (CH 2 O) n Primary function: available energy Monomer: monosaccharides hexoses = glucose, fructose, galactose pentoses = ribose, deoxyribose Dimer: disaccharides sucrose, lactose, maltose Polymer: polysaccharides starch, glycogen

34 Formation of Disaccharides 2 Monosaccharides are covalently bonded via dehydration synthesis to produce disaccharides

35 Lipids Composed predominately of hydrocarbon chains and rings Primary functions: stored energy, insulation, water proofing, hormones Insoluble in polar solvents Major types Fatty acids Eicosinoids Triglycerides Ketone bodies phospholipids Steroids

36 Lipids: fatty acids Fatty acids = long chain of C and H with COOH at one end (usually 16-18C) Saturated vs unsaturated Monounsaturated vs polyunsaturated Prostaglandins and eicosanoids = modified 20 C polyunsaturated fatty acid (arachadonic acid) Important regulators of cell function

37 Prostaglandins and Eicosanoids Diverse group Include the prostaglandins Cramps and labor contractions Blood clotting Inflammation Fever Etc. We will discuss these more with the endocrine system!

38 Lipids: triglycerides Neutral fats and oils Composed of 3 fatty acids + 1 glycerol

39 Hydrolysis of triglycerides in adipose tissue release free fatty acids. Free fatty acids can be converted in the liver to ketone bodies. Ketoacidosis: Increased ketone bodies in the blood which lowers ph. Ketone Bodies

40 Lipids: phospholipids Amphipathic molecules structurally similar to triglycerides major component of plasma membrane Hydrophilic head Hydrophobic tails

41 Phospholipid structure

42 Lipids: steroids Ring structure (3) 6-C rings + (1) 5-C ring cholesterol major precursor for steroid hormones CH 3 CH3 CH 3 CH 3 CH 3 OH A B C D

43 Examples of steroids

44 Proteins Structural and functional composed of amino acids (20) Sensitive to environmental change four levels of structure primary secondary tertiary quatrenary

45 Proteins: amino acid structure R COOH C NH 2 carboxyl amine H

46

47 Proteins: amino acids asparagine arginine histidine

48 Building proteins Dehydration synthesis: Hydrogen from the amino end of one amino acid combines with hydroxyl group of carboxyl end of another amino acid. Covalent bond formed Peptide bond: Bond between two adjacent amino acids. Length of polypeptide chains vary in length. If greater than 100 amino acids, is called a protein.

49 Proteins: primary structure Sequence and number of amino acids ala val ile pro phe met ser thr cys tyr leu gly asn gln glu trp asp lys arg his

50 Proteins: secondary structure Alpha helix Beta-pleated sheet

51 Proteins: tertiary structure 3-D structure which determines the function

52 Proteins: quaternary structure Association of two or more polypeptide chains

53 Four levels of protein structure: summary

54 Nucleic acids DNA: genetic material Double-stranded in cells RNA: protein synthesis Single-stranded in cells three types mrna = messenger RNA trna = transfer RNA rrna = ribosomal RNA

55 Nucleic Acid Structure Composed nucleotides Phosphate Sugar DNA = deoxyribose RNA = ribose Nitrogen base purines adenine (A) guanine (G) pyrimidines cytosine (C) thymine (T) [only in DNA] Uracil (U) [only in RNA]

56 DNA Structure O O P O O 5 CH 2 O R 4 H H H 3 OH H Double helix

57 Base pairing rules

58 RNA Single stranded (variation in viruses) Contains 5-C sugar, ribose Uracil replaces thymine as nitrogen base

59 mrna trna rrna RNA

60 RNA vs. DNA

61 Adenosine Triphosphate (ATP) Predominant form of energy currency in living things Composed of an adenine nucleotide + 2 additional phosphates bound in high energy bonds (7kcal)