W1-Chapter 1 biochemistry ASASIpintar SEM1.pdf

Week1: Biochemistry 2: Atoms & Molecules 3: The Chemistry of Water 4: Carbon: The Basis of Molecular Diversity 5: Biological Macromolecules & Lipids Dr Ikhwan Zakaria. 2425  Matter Takes up space and has mass Exists as elements (pure form) and in chemical combination...

Week1: Biochemistry 2: Atoms & Molecules 3: The Chemistry of Water 4: Carbon: The Basis of Molecular Diversity 5: Biological Macromolecules & Lipids Dr Ikhwan Zakaria. 2425  Matter Takes up space and has mass Exists as elements (pure form) and in chemical combinations called compounds  Elements Can’t be broken down into simpler substances by chemical reaction Composed of atoms  The Elements of Life Essential elements in living things include carbon C, hydrogen H, oxygen O, and nitrogen N making up 96% of an organism A few other elements make up the remaining 4% of living matter  Trace Elements Trace elements are required by an organism in only minute quantities Minerals such as Fe and Zn are trace elements  Deficiencies If there is a deficiency of an essential element, disease results (b) Iodine deficiency (a) Nitrogen deficiency (Goiter)  Compounds Are substances consisting of two or more elements combined in a fixed ratio Have characteristics different from those of their elements  Properties of Matter An element’s properties depend on the structure of its atoms Each element consists of a certain kind of atom that is different from those of other elements An atom is the smallest unit of matter that still retains the properties of an element Sub atomic particles:  Energy Levels Are represented by electron shells Third energy level (shell) Second energy level (shell) Energy absorbed First energy level (shell) Energy lost Atomic nucleus (b) An electron can move from one level to another only if the energy it gains or loses is exactly equal to the difference in energy between the two levels. Arrows indicate some of the step-wise changes in potential energy that are possible.  Atomic Number & Atomic Mass Atoms of the various elements differ in their number of subatomic particles The number of protons in the nucleus = atomic number The number of protons + neutrons = atomic mass Neutral atoms have equal numbers of protons & electrons (+ and – charges) Electron distribution diagram Electron distribution diagram for the first 18 elements  Isotopes Same number of protons, different number of neutrons  Isotopes May be radioactive, spontaneously giving off particles and energy  radioactive decay May be used to date fossils (C-14) or as medical tracers (I-131)  CHEMICAL BONDS Covalent Ionic Hydrogen Van der Waals Polar Non-polar  Covalent Bonds Sharing of at least a pair of valence electrons –Electrons in the outermost shell Examples: CH4, H2, O2  Covalent Bonding in 4 molecules Fill in the Blanks ___hydrogen atoms share ___ pair of electrons, forming a ____ bond ___oxygen atoms share ___ pair of electrons, forming a ____ bond ___hydrogen atoms and ____ oxygen atom(s) are joined by ____ bonds, forming a _______ of water ___hydrogen atoms can satisfy the valence of one ______ atom, forming methane.  Covalent Bonding Electronegativity – Is the attraction of a particular kind of atom for the electrons in a covalent bond The more electronegative an atom – The more strongly it pulls shared electrons toward itself Electronegativity  Covalent Bonding In a nonpolar covalent bond – The atoms have similar electronegativities – Share the electron equally  Covalent Bonding In a polar covalent bond – The atoms have different electronegativities – Share the electrons unequally Because oxygen (O) is more electronegative than hydrogen (H), shared electrons are pulled more toward oxygen. d– This results in a partial negative charge on the oxygen and a partial positive O charge on the hydrogens. H H d+ d+ H2O  Balanced Symmetrical Unbalanced asymmetrical  Polar bond vs. Polar molecule Bond polarity – Difference in electronegativities Molecule polarity – Sum of all bond polarities in the molecule polar nonpolar  Ionic Bonds Electron transfer between two atoms –Anion (-ve), Cation (+ve) Metals+Non-metals = ionic compound –Ionic compounds are often called salts, which may form crystals Electron transfer and ionic bonding http://eiyecaieyre.deviantart.com/art/Science-homework-comic- 68583801  Hydrogen Bonds Between positively charged H atom and the strongly electronegative O /N/F of another molecule  intermolecular δ- δ- δ+ δ+ δ+ δ+ δ- δ+ δ+  Van der Waals Interactions Between molecules  intermolecular Occur when transiently positive and negative regions of molecules attract each other 31 THE CHEMISTRY OF WATER  1. Cohesive behavior Cohesion:  hydrogen bonds high surface tension, difficult to stretch/break surface of liquid Help transport of water upwards in plants Adhesion: water adhere to unlike molecules Water – plant cell-wall (hydrogen bond) Help counter the downward pull of gravity Surface tension, – Related to cohesion – Ease to stretch/break surface of liquid  2. Moderation of Temperature High Specific Heat, – the amount of energy needed to increase the temperature of 1 gram of a substance by 1 ℃ (changes temp slowly)  1 cal/gram℃ – Water changes temperature less than other liquids – Due to hydrogen bonding, heat absorbed to break H-bonds Heat released to form H-bonds 1 cal heat  small ∆ in water’s temperature  Water can moderate air temperature in coastal areas – High Heat of Vaporization, the amount of heat energy needed to change 1g liquid to a gas (or evaporate) Hottest liquid molecules (greatest kinetic energy) leave surface as gas Hydrogen bonds need to be broken Evaporative cooling –Stability of temperature in lakes and ponds –Dissipation of body heat by sweating – High Heat of Fusion, the energy needed to convert a substance from a liquid to a solid (freezing, in the case of water)  3. Floating of ice on liquid water >4°C  water is liquid 0−4°C  water begins to freeze 0°C  H2O molecules form stable H-bonds Solid water is less dense than liquid water,  ice floats. Allows life to exist under frozen surfaces of lakes and sea  4. A universal solvent Hydration shell As the universal solvent, highly polar compound, dissolves polar substances and ionic compounds – Hydrophilic  ionic and polar – Hydrophobic  non- ionic and non-polar  Water as a solvent Figure 3.9 A water soluble protein This oxygen is attracted to a slight positive charge on the lysozyme molecule. d– d+ This hydrogen is attracted to a slight negative charge on the lysozyme molecule. (a) Lysozyme molecule (b) Lysozyme molecule (purple) (c) Ionic and polar regions on the protein’s in a non-aqueous in an aqueous environment surface attract water molecules. environment such as tears or saliva  pH (power of Hydrogen) pH scale runs between 0 and 14 and measures the relative acidity and alkalinity of aqueous solutions The pH of a solution is defined by the negative logarithm of H+ concentration, written as pH = –log [H+] The pH scale and pH values of some aqueous solutions.  pH and buffers Buffers: Materials that have both acid and base properties Resist change in pH Absorb excess H+ or donate H+ E.g. Bicarbonate ion HCO3- H+ donor H+ acceptor H+ ion (acid) pH rise (base) H2CO3 ⇌ HCO3- + H+ pH drop Carbonic acid Bicarbonate ion Blood pH imbalance CARBON CHEMISTRY  Carbon chemistry Carbon is the backbone of biological molecules (macromolecules) All living organisms are made up of chemicals based mostly on the element carbon  Carbon Chemistry Organic chemistry is the study of carbon compounds Carbon atoms can form diverse molecules by bonding to four other atoms Carbon compounds range from simple molecules to complex ones Carbon has four valence electrons and may form single, double, triple, or quadruple bonds 46  The bonding versatility of carbon allows it to form many diverse molecules, including carbon skeletons Figure 4.3 The shapes of three simple organic molecules.  The electron configuration of carbon gives it covalent compatibility with many different elements Fig 4.4 Valences of the major elements of organic molecules.  Hydrocarbons Hydrocarbons are molecules consisting of only carbon and hydrogen Hydrocarbons are found in many of a cell’s organic molecules  Isomers Isomers are molecules with the same molecular formula but different structures and properties Three types of isomers are – Structural Variation in colavent arrangement – Geometric/cis-trans Variation in arrangement about a double bond – Enantiomers/optical isomers Variation in spatial arrangement around an asymmetric carbon Molecules are mirror image 3 types of isomers  Enantiomers are important in the pharmaceutical industry  Chemical Groups Chemical groups that are directly involved in CH3 OH chemical reactions are Estradiol called functional groups HO They are the chemically reactive Female lion groups of atoms within an organic CH3 OH molecule CH3 Give organic molecules distinctive chemical properties O Testosterone Figure 4.9 Male lion Chemical Groups Chemical Groups Chemical Groups Chemical Groups  Organic Compounds Polymer: most macromolecules are polymers…a single unit repeated many times, hence “poly”mer Monomer: a single unit (“mono”mer) that makes up a polymer Macromolecule: a large organic molecule, made of multiple polymers Four classes of macromolecules (biomolecules) – Carbohydrates – Lipids – Proteins – Nucleic Acids  The Synthesis and Breakdown of Polymers Monomers form larger molecules by – dehydration reaction/synthesis – - H2O molecule Polymers can disassemble by – Hydrolysis – + H2O molecule 60  Carbohydrate C, H, O (CH2O) n, where n is any number from 3 to 8 Serve as fuel and building material 3 classes: – Monosaccharides – Disaccharides – polysaccharides  Monosaccharides Are the simplest sugars Can be used for fuel Can be converted into other organic molecules Can be combined into polymers glucose, galactose, fructose  Monosaccharides Chemical equilibrium between the linear and ring – May be linear structures greatly favors the formation of rings – Can form rings  α- and β-glucose Two forms of glucose, seen below, differ only in a reversal of the H and OH on the first carbon. Even very small changes in the position of certain atoms may dramatically change the molecule  Examples of monosaccharides Sugars vary in the position of the carbonyl group – Length of skeleton – The chemical groups are arranged around an asymmetric carbon  Disaccharides Di (two) sacchar (sugar): a double sugar that consists of two monosaccharides joined by a glycosidic linkage. Chemical formula – C12H22O11 Glycosidic linkage: covalent bond formed by a condensation/dehydration (water is lost) reaction between two sugar monomers –Common Disaccharides: Maltose= Glucose+Glucose (beer) Lactose= Glucose+Galactose (milk) Sucrose= Glucose+Fructose (table sugar/fruit sugar) –All saccharides are storage molecules, they store energy to be used by living system Disaccharide synthesis  Polysaccharides A polysaccharide consists of a series of connected monosaccharides. A polysaccharide is a polymer. Cells hydrolyze storage polysaccharides into sugars as needed. Common storage polysaccharides are starch, glycogen, cellulose, and chitin. Polysaccharides of plants and animals branched  Storage polysaccharide in plants Chloroplast Starch granule Starch – α-glucose polymer that is an energy storage molecule in plants 1 m Amylose Amylopectin (a) Starch: a plant polysaccharide Structural polysaccharide in plants Cellulose – polymer of -glucose molecules. – It serves as a structural molecule in the walls of plant cells. – Cellulose is the major component in wood. Starch and cellulose structure Storage polysaccharide in animals Glycogen – α-glucose polymer – more tightly branched than starch Mitochondria Giycogen granules 0.5 m Glycogen (b) Glycogen: an animal polysaccharide  Structural polysaccharide in animals Chitin is a polymer similar to cellulose, but each - glucose molecule has a nitrogen containing group attached. Chitin serves as a structural molecule in the walls of fungus cells, exoskeletons of insects, and mollusks  Lipids Triglyceride phospholipid steroid (fat)  Fats/Triglyceride They consist of 3 fatty-acids attached to a glycerol molecule. Fatty acids are hydrocarbon chains with a carboxyl group at one end of the chain Hydrophobic  Hydrocarbon chains of fatty acids: HC chains of Saturated fatty acids have a single covalent bond between each pair of carbon atoms, and each carbon has 2 hydrogen bonded to it. The carbon is “saturated” with hydrogen HC chains of Unsaturated fatty acids have a double covalent bond, and each of the two carbons in this bond have only one hydrogen atom bonded to it. Saturated vs. Unsaturated  Phospholipids Have only two fatty acids Have a phosphate group instead of a third fatty acid  Phospholipids Phospholipid structure – Consists of a hydrophilic “head” and hydrophobic “tails” CH2 + N(CH ) 3 3 Choline CH2 O O P O– Phosphate O CH2 CH CH2 Glycerol O O C O C O Fatty acids Hydrophilic head Hydrophobic tails (c) Phospholipid (b) Space-filling model Figure 5.11 (a) Structural formula symbol Lipid bilayer  Steroids Steroids are characterized by a backbone of four linked carbon rings. Examples include cholesterol (a component of cell membranes), hormones, testosterone and estrogen. Cholesterol  Proteins: functions Proteins can be grouped according to their functions: – Structural proteins: keratin, hair, horns, collagen, connective tissue, spider silk – Storage proteins: casein in milk & ovalbumin in egg whites – Transport proteins: found on cell membranes that transport materials into and out of cell, hemoglobin – Defensive proteins: antibodies to fight infection – Enzymes: regulate the rate of chemical reactions  Proteins: building blocks Amino acids – Are organic molecules possessing both carboxyl and amino groups – Differ in their properties due to differing side chains, called R groups The 20 Amino Acids: Nonpolar side chains The 20 Amino Acids: Polar side chains  The 20 Amino Acids: Electrically charged side chains  Amino acids to polypeptide Dehydration reaction (-H2O) Linking carboxyl group of one amino acid with amino group of the next Peptide bond connects polypeptide backbone  Proteins For a protein to function properly, it has to be folded into a specific structure. Protein folding has four levels of “folding” – Primary – Secondary – Tertiary – Quarternary  Primary Structure The primary structure of a protein describes the linear order of amino acids. Peptide bond  Secondary Structure Secondary structure of a protein is a 3D shape that results from hydrogen bonding – between the amino and carboxyl groups of adjacent amino acids (polypeptide backbone) – Individually, the H-bonds are weak – but repetition over a long region can support a particular shape The bonding can produce a spiral structure called alpha helix Or a folded structure called beta pleated sheet Proteins whose shape is dominated by these 2 patterns, form fibrous proteins  Tertiary Structure Tertiary structure of a protein includes additional 3D shaping. The following factor contribute to tertiary “folding” – Hydrogen bond between R groups of amino acids – Ionic bonding between R groups of amino acids – Hydrophobic interaction – Disulfide bonds, when the sulfur atoms in two cysteine amino acids bond “disulfide bridge”  Quaternary Structure Quaternary structure describes a protein that is assembled from two or more separate polypeptides. – E.g. Transthyretin is made up of 4 polypeptides some proteins only Collagen has 3 helical polypeptides intertwined into a larger triple helix Hemoglobin consists of 4 polypeptide subunits held together by hydrogen bonding, R-groups interactions, and disulfide bonds  What determines protein structure? Physical and chemical conditions of proteins’ environment – Denatured proteins no longer work in their unfolded condition – Proteins may be denatured by changes in pH salt concentration Temperature Environment 100  Nucleic Acids: Role 2 types: – DNA (DeoxyriboNucleic Acid) Inheritable genetic material Can direct synthesis of mRNA – RNA (RiboNucleic Acid) Conveys genetic info for building proteins controls protein synthesis  Nucleic acids: structure Nucleic acids are made up of: – nitrogenous bases Cytosine, Thymine and Uracil and are “Pyrimidines” single-ring bases Adenine and Guanine are Purines double-ring bases – 5-Carbon sugar Deoxyribose Ribose Nitrogenous base + sugar = nucleoside  Nucleoside + phosphate group = nucleotide  Nucleic Acids: Structure of DNA and RNA The two strands of DNA are antiparallel: they run in opposite directions. Form a double helix One is arranged in the 5’3’ direction the other in the 3’5’ direction RNA RNA differs from DNA in the following ways: – The sugar in RNA is ribose, not deoxyribose – The thymine does not occur in RNA, it is replaced with uracil – RNA is a single- stranded molecule and does not form a double helix as DNA does.  Summary I Matter consists of chemical elements in pure form and in combinations called compounds An element’s properties depend on the structure of its atoms The formation and function of molecules depend on chemical bonding between atoms Polar covalent bonds in water molecules result in hydrogen bonding 4 emergent properties of water contributes to Earth’s suitability for life Acidic and basic conditions affect living organisms  Summary II Carbon atoms can form diverse molecules by bonding to four other atoms A few chemical groups are key to molecular function Macromolecules are polymers made from monomers Carbohydrate serve as fuel and building material Lipids are a diverse groups of lipid macromolecules Proteins include a diversity of structures, resulting in a wide range of functions Nucleic acids store, transmit, and help express hereditary information