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BIOCHEMISTRY WATER Chapter 3: Water and Life Water • Cells are 70-90% water • Three-fourths of Earth’s surface is covered by water • Water is the biological medium for life on Earth • Water must be present for life (as we know it) Hydrogen bonds between water molecules Water’s polarity • Polar – opposite ends of molecule have opposite charges – Opposing charges due to oxygen’s electronegativity • Oxygen has partial negative • Hydrogens have partial positive • Water forms H bonds (opposite charges attract) – ~15% water molecules in our body are bonded to four partners Four Properties of Water Allowing for Life 1. Cohesion/Adhesion 2. Moderation of temperature 3. Insulation of bodies of water by floating ice 4. The solvent of life Trees Walking on water Cohesion • Cohesion - H bonding keeps water molecules close together • Makes water a “structured” liquid • Adhesion – the clinging of one substance to another – Ex. water sticks to sides of a glass • Both cohesion and adhesion help water move up from roots to plants to leaves • Surface tension – measure of how difficult it is to stretch or break the surface of a liquid – Water has a high surface tension due to H bonds – almost like a thin film on the surface. Moderation of Temperature • Heat – measure amount of total kinetic energy (cal, kcal, or J) – 1 cal = amount of heat needed to raise the temp of 1g of water 1°C – 1000 cal = 1 kcal – 1 cal = 4.184 J • Temperature – measures the intensity of heat due to average kinetic energy of molecules (°C) • Specific heat – amount of heat that must be absorbed or lost for 1 g to change its temp by 1 °C – Specific heat of water = 1 cal/g/°C – Compared to most substances, water’s specific heat is quite high. – This means water will changes its temp slower when it absorbs or gives off heat. – Why? H bonds need heat to break and heat is released when bonds are formed. – High specific heat allows large bodies of water to absorb lots of heat in summer without raising temp too high. In winter, gradual cooling of water helps warm the air. – High specific heat helps stabilize ocean temp to better support marine life. • Heat of vaporization – the quantity of heat a liquid must absorb for 1 g to be converted to gaseous state • Water has a high heat of vaporization – 580 cal of heat needed to evaporate 1 g water at 25°C because H bonds must be broken before molecules can change to gas. – Evaporative cooling – as liquid evaporates, the surface cools because the hottest molecules leave as gas and cooler molecules are left behind • Why sweat? Evaporative cooling in progress! • Why do we sweat more on humid days? Insulation of Bodies of Water by Floating Ice • Water is more dense than ice. At 4°C water is at its most dense state, then as it cools to O°C, the molecules freeze. The H bonds keep the water molecules slightly apart (like a lattice) so air pockets form within ice. • Lakes, oceans etc. would freeze solid if ice was more dense than water. During the summer, only upper few inches of ocean would thaw making life as we know it impossible (not to mention ice skating and hockey! ) Ice Caps/Glaciers Melting • Global warming causing ice to melt. – Loss of habitat – Less sunlight reflected into space so increases temp The structure of ice Ice, water, and steam Floating ice and the fitness of the environment Floating ice and the fitness of the environment: ice fishing Ice floats and frozen benzene sinks The Solvent of Life • Solvent - dissolving agent in a solution • Solute – the substance that is dissolved • Aqueous solution – water is the solvent • Water can dissolve many substances, but obviously not all! – Hydrophilic – likes water – Hydrophobic – repel water (nonpolar and nonionic) A crystal of table salt dissolving in water A water-soluble protein • Concentrations – Molecular mass – sum of mass of atoms (daltons) • Molecular mass of CO2 = 12 + 16(2) = 44 daltons – 1 mole = 6.02 x 1023 molecules – 1 mole CO2 = 44 grams – Molarity (M) = mole/liter Moles Chemical reaction: hydrogen bond shift pH • H+ (protons) occasionally move from one water molecules to another (disassociation). • If water loses a H+ then it becomes OH- (hydroxide ion). • If water gains a H+ then it becomes H3O+ (hydronium ion). • In pure water, the OH- and H+ concentrations are equal. • Acid – increases H+ concentrations • Base – decreases H+ concentrations • pH = -log [H+] – For neutral water: • pH = -log 10-7 = -(-7) = 7 – pH decreases as H+ increases – A pH of 3 vs. pH of 6 is a 1000 fold difference (10 fold for each step) The pH of some aqueous solutions Buffers • Buffers – minimize changes in pH by being able to release or take in H+ • Buffers keeps blood between 7 and 7.8 • The equation below shifts right to decrease pH and left to increase pH (bicarbonate buffer) H2CO3 Carbonic acid HCO3- + bicarbonate H+ proton Acid Rain • Acid precipitation – below 5.6 • Caused primarily by increased levels of sulfur and nitrogen oxides released from the burning of fossil fuels • Acid precip can damage lakes, streams, and soil. • Acid can make harmful heavy metals more soluble in water. The effects of acid precipitation on a forest Pulp mill Acid rain damage to statuary, 1908 & 1968 CARBON AND THE MOLECULAR DIVERSITY OF LIFE CHAPTER 4 ISOMERS • Compounds with the same chemical formula but different structures Functional Groups See diagram of functional groups Three types of isomers Structural isomers Functional Groups of Organic Compounds A comparison of functional groups of female (estradiol) and male (testosterone) sex hormones THE STRUCTURE AND FUNCTION OF LARGE BIOLOGICAL MOLECULES CHAPTER 5 Macromolecules • Polymer – long molecule consisting of many similar building blocks connected by covalent bonds • Monomer – the building blocks SYNTHESIS AND BREAKDOWN • Dehydration synthesis (condensation reaction) – removal of water to join 2 compounds (anabolic) • Hydrolysis – addition of water to break a bond between 2 compounds (catabolic) The synthesis and breakdown of polymers CARBOHYDRATES • Monomers are sugars. • Used for energy. • Types of carbohydrates – Monosaccharides – one sugar • Examples: glucose, fructose, and galactose – Disaccharides – two sugars joined • Examples: lactose, sucrose and maltose • Joined by glycosidic linkage via dehydration synthesis – Polysaccharides – many sugars joined • Examples: starch, glycogen, chitin (exoskeletons of insects), and cellulose (fiber) The structure and classification of some monosaccharides Linear and ring forms of glucose Examples of disaccharide synthesis Storage polysaccharides Starch and cellulose structures Starch and cellulose structures The arrangement of cellulose in plant cell walls Cellulose digestion: termite and Trichonympha Cellulose digestion: cow Chitin, a structural polysaccharide: exoskeleton and surgical thread LIPIDS • Little or no affinity for water (hydrophobic) • No monomers • Many uses including insulation, store energy, hormones • Examples: Fat, phospholipids, and steroids • Fats – composed of glycerol (an alcohol) and fatty acids – Saturated – no double bonds in carbon chain – Unsaturated – at least one double bond in carbon chain • Phospholipids – make up plasma membrane • Steroids – consist of 4 carbon rings; examples include cholesterol, vitamin D, estrogen, and testosterone Examples of saturated and unsaturated fats and fatty acids Saturated and unsaturated fats and fatty acids: butter and oil The structure of a phospholipid Two structures formed by self-assembly of phospholipids in aqueous environments The synthesis and structure of a fat, or triacylglycerol Cholesterol, a steroid An Overview of Protein Functions PROTEIN • Monomers are amino acids. • Many uses including enzymes, antibodies, receptors, structural, hormones, and transport • Protein – consists of one or more polypeptides with specific 3-D structure – Polypeptide = polymer of amino acids – Peptide = bond between 2 amino acids • There are 20 different amino acids differing only by the R group The 20 amino acids of proteins: nonpolar The 20 amino acids of proteins: polar and electrically charged Making a polypeptide chain • Four levels of protein structure – Primary – sequences of amino acids – Secondary – hydrogen bonds cause coils and folds • Beta pleated sheet and alpha helix – Tertiary – irregular contortions due to various weak bonds: • Hydrophobic interactions • Disulfide bridges • Ionic bonds • Van der Waals interactions (weak attractions from transient partial charges) – Quaternary – two or more polypeptide chains aggregated into one functional macromolecule • Examples: collagen and hemoglobin The primary structure of a protein A single amino acid substitution in a protein causes sickle-cell disease Sickled cells The secondary structure of a protein Examples of interactions contributing to the tertiary structure of a protein The quaternary structure of proteins Review: the four levels of protein structure Protein Structure • The function of a protein depends on its 4 levels. • Denaturation – the “unraveling of a protein” so that the 2nd, 3rd, and 4th level structures are gone and all that is left is the primary sequence – Loss of structure = loss of function – Causes of denaturation - High temp, high or low pH Denaturation and renaturation of a protein Spider silk: a structural protein Silk drawn from the spinnerets at the rear of a spider NUCLEIC ACIDS • Monomers are nucleotides. • Nucleotides – are made of sugar, phosphate, and a N-base • Examples: DNA, ATP, and RNA • We will discuss these in great detail later in the semester! The components of nucleic acids The DNA double helix and its replication AN INTRODUCTION TO METABOLISM CHAPTER 8 Transformations between kinetic and potential energy Kinetic and potential energy: dam Kinetic and potential energy: cheetah at rest and running BIOENERGETICS • Bioenergetics – the study of how energy flows through living organisms • Catabolic – reactions or pathways where a larger molecule is broken down into smaller molecules • Anabolic - reactions or pathways where smaller molecules are joined to build a larger molecule THERMODYNAMICS • First Law of Thermodynamics – energy can be transferred and transformed, but not created nor destroyed • Second Law of Thermodynamics – every energy transfer or transformation makes the universe more disordered (have more entropy) – Entropy – measure of disorder or randomness – Most energy transformations involve at least some energy be changed to heat – Heat is the lowest form of energy – Biological order has increased over time – Second law requires only that processes increase the entropy of the universe – Organisms may decrease entropy but entire universe must increase entropy Order as a characteristic of life • Free energy (G) – energy available to do work when temperature is uniform throughout system G = H – TS • H = system’s total energy (enthalpy) • T = temp in Kelvin (° C + 273) • S = entropy (measure of disorder) ∆ G = G final state – G starting state ∆ G = ∆ H - T∆ S • For a spontaneous reaction: – ∆ G = negative • So must: give up energy (decrease H) and/or give up order (increase S) ∆ G = 0 at equilibrium • Free energy increases if move away from equilibrium and decreases if move toward equilibrium The relationship of free energy to stability, work capacity, and spontaneous change • Exergonic – ∆ G = negative – Spontaneous – Net release of energy • Endergonic – ∆ G = positive – NOT spontaneous – Stores free energy in molecules Energy changes in exergonic and endergonic reactions • Cells at equilibrium are dead! • Cells can keep disequilibrium by having products of one reaction not accumulate but instead become reactants of another reaction • Energy coupling – an exergoinc reaction drives an endergoinc reaction (or a reaction that increase entropy is paired with one that decreases entropy) Disequilibrium and work in closed and open systems ATP = ADENOSINE TRIPHOSPHATE ATP + H2O ADP + Pi ∆ G = -7.3kcal/mol • ADP = adenosine diphosphate • Normally the phosphate is bonded to an intermediate compound which is then considered phosphorylated • The reverse reaction is endergonic and requires +7.3 kcal/mol to make ATP from ADP The structure and hydrolysis of ATP Energy coupling by phosphate transfer The ATP cycle ENZYMES • Catalytic proteins or enzymes – change the rate of reaction without being consumed by the reaction • Activation energy (EA) – energy needed to start a reaction – Energy needed to contort the reactants so the bonds can change • Enzymes lower activation energy by enabling reactants to absorb enough energy to reach transition state at moderate temps • Enzymes are substrate specific • Substrate – the reactant on which an enzyme works • Active site – area on enzyme where substrate fits • Induced fit – model of enzyme activity Energy profile of an exergonic reaction The effect of an enzyme on activation FThe induced fit between an enzyme and its substrate The catalytic cycle of an enzyme • Effects of pH and Temp – Optimal temperatures and pH ranges exist for enzymes – High temperatures can denature an enzyme – Changes in pH can also denature enzymes Environmental factors affecting enzyme activity • Cofactors – non-protein helpers that bind to active site or substrate (zinc, iron) – Coenzymes – cofactors that are organic (vitamins) • Enzyme Inhibitors – reduce enzyme activity – Competitive inhibitors – block substrate from entering active site • Reversible • Overcome by adding more substrate – Noncompetitive inhibitors – bind to another part of enzyme thereby changing the enzyme’s shape making it inactive • Irreversible • Examples: DDT, sarin gas, and penicillin Inhibition of enzyme activity • Allosteric regulation – Allosteric sites – receptors on enzymes (not the active site) that may either inhibit or stimulate enzyme activity • Cooperativity – an enzyme with multiple subunits where binding to one active site causes shape changes to rest of subunits which in turn activates those subunits • Feedback inhibition – End product of a pathway acts as a inhibitor of an enzyme within the pathway Allosteric regulation of enzyme activity Cooperativity Feedback inhibition