Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
WATER Chapter 3: Water and the Fitness of the Environment 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) Figure 3.1 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 2. Moderation of temperature 3. Insulation of bodies of water by floating ice 4. The solvent of life Figure 3.2x Trees Figure 3.3 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 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 less 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 – 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 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! ) Figure 3.5 The structure of ice (Layer 2) Figure 3.5x1 Ice, water, and steam Figure 3.6 Floating ice and the fitness of the environment Figure 3.6x1 Floating ice and the fitness of the environment: ice fishing Figure 3.6x2 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) Figure 3.7 A crystal of table salt dissolving in water Figure 3.8 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 Figure 3.x2 Moles Unnumbered Figure (page 47) 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 • For pure (neutral) water at 25°C: – [OH-] [H+] = 10-14 – [H+] = 10-7 – [OH-] = 10-7 • If enough acid is added to increase the [H+] to 10-4, then the [OH-] will decrease by an equivalent amount or 10-10 • Because concentrations can vary by factors of 100 trillion, scientists use a log scale. • pH = -log [H+] – For neutral water pH = -log 10-7 = -(-7) = 7 – A pH of 3 vs. ph of 6 is a 1000 fold difference (10 fold for each step) Figure 3.9 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- + H+ bicarbonate 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. Figure 3.10 The effects of acid precipitation on a forest Figure 3.10x1 Pulp mill Figure 3.10x2 Acid rain damage to statuary, 1908 & 1968 CARBON AND THE MOLECULAR DIVERSITY OF LIFE CHAPTER 4 Figure 3.10x2 Acid rain damage to statuary, 1908 & 1968 ISOMERS Compounds with the same chemical formula but different structures FUNCTIONAL GROUPS See diagram of functional groups Figure 4.6 Three types of isomers Figure 4.6ax Structural isomers Table 4.1 Functional Groups of Organic Compounds Figure 4.8 A comparison of functional groups of female (estradiol) and male (testosterone) sex hormones Figure 4.8x1 Estrone and testosterone THE STRUCTURE AND FUNCTION OF MACROMOLECULES CHAPTER 5 Macromolecules – long molecule consisting of many similar building blocks connected by covalent bonds Monomer – the building blocks Polymer SYNTHESIS AND BREAKDOWN Dehydration synthesis (condensation reaction) – removal of water to join 2 compounds Hydrolysis – addition of water to break a bond between 2 compounds Figure 5.2 The synthesis and breakdown of polymers CARBOHYDRATES Monosaccharides Examples: glucose, fructose, and galactose One sugar Disaccharides Examples: Lactose, sucrose and maltose Two sugars Joined by glycosidic linkage via dehydration synthesis Polysaccharides Examples: starch, glycogen, and cellulose Many sugars Figure 5.3 The structure and classification of some monosaccharides Figure 5.4 Linear and ring forms of glucose Figure 5.5 Examples of disaccharide synthesis Figure 5.6 Storage polysaccharides Figure 5.7a Starch and cellulose structures Figure 5.7b,c Starch and cellulose structures Figure 5.8 The arrangement of cellulose in plant cell walls Figure 5.x1 Cellulose digestion: termite and Trichonympha Figure 5.x2 Cellulose digestion: cow Figure 5.9 Chitin, a structural polysaccharide: exoskeleton and surgical thread LIPIDS Little or no affinity for water (hydrophobic) 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 Figure 5.11 Examples of saturated and unsaturated fats and fatty acids Figure 5.11x Saturated and unsaturated fats and fatty acids: butter and oil Figure 5.12 The structure of a phospholipid Figure 5.13 Two structures formed by self-assembly of phospholipids in aqueous environments Figure 5.10 The synthesis and structure of a fat, or triacylglycerol Figure 5.14 Cholesterol, a steroid Figure 5.14x Cholesterol Table 5.1 An Overview of Protein Functions PROTEIN Polypeptide – polymer of amino acids Protein – consists of one or more polypeptides with specific 3-D structure There are 20 different amino acids differing only by the R group Figure 5.15 The 20 amino acids of proteins: nonpolar Figure 5.15 The 20 amino acids of proteins: polar and electrically charged Figure 5.16 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 Quaternary – two or more polypeptide chains aggregated into one functional macromolecule Examples: collagen and hemoglobin Figure 5.18 The primary structure of a protein Figure 5.19 A single amino acid substitution in a protein causes sickle-cell disease Figure 5.19x Sickled cells Figure 5.20 The secondary structure of a protein Figure 5.22 Examples of interactions contributing to the tertiary structure of a protein Figure 5.23 The quaternary structure of proteins Figure 5.24 Review: the four levels of protein structure Figure 5.25 Denaturation and renaturation of a protein Figure 5.21 Spider silk: a structural protein Figure 5.21x Silk drawn from the spinnerets at the rear of a spider NUCLEIC ACIDS Examples: DNA and RNA Nucleotides – building blocks of nucleic acids (made of sugar, phosphate, and a N-base) We will discuss these in great detail later in the semester! Figure 5.29 The components of nucleic acids Figure 5.30 The DNA double helix and its replication AN INTRODUCTION TO METABOLISM CHAPTER 8 Figure 6.2 Transformations between kinetic and potential energy Figure 6.2x1 Kinetic and potential energy: dam Figure 6.2x2 Kinetic and potential energy: cheetah at rest and running 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 Bioenergetics THERMODYNAMICS 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) First 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 Figure 6.4 Order as a characteristic of life energy (G) – energy available to do work when temperature is uniform throughout system Free G = H – TS = system’s total energy (enthalpy) T = temp in Kelvin (° C + 273) S = entropy H ∆ 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 Figure 6.5 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 Figure 6.6 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 Figure 6.7 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 Figure 6.8 The structure and hydrolysis of ATP Figure 6.9 Energy coupling by phosphate transfer Figure 6.10 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 Figure 6.12 Energy profile of an exergonic reaction The effect of an enzyme on activation Figure 6.14 The induced fit between an enzyme and its substrate Figure 6.15 The catalytic cycle of an enzyme Effects of pH and Temp Optimal temperatures and pH ranges exist for enzymes Figure 6.16 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 Figure 6.17 Inhibition of enzyme activity Allosteric regulation Allosteric sites – receptors on enzymes (not the active site) that may either inhibit or stimulate enzyme activity Feedback inhibition End product of a pathway acts as a inhibitor of an enzyme within the pathway 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 Figure 6.18 Allosteric regulation of enzyme activity Figure 6.19 Feedback inhibition Figure 6.20 Cooperativity