Human Anatomy & Physiology Chapter 2 Lecture PDF

HUMAN ANATOMY & PHYSIOLOGY Second Edition Chapter 02 The Chemistry of Life PowerPoint® Lectures created by Suzanne Pundt, University of Texas at Tyler...

HUMAN ANATOMY & PHYSIOLOGY Second Edition Chapter 02 The Chemistry of Life PowerPoint® Lectures created by Suzanne Pundt, University of Texas at Tyler Copyright © 2019, 2016 Pearson Education, Inc. All Rights Reserved  Review 1) What are atoms and atomic structures 2) What are mass number & atomic number 3) What is an element, isotope, radioisotope 4) When atoms are stable and when they react with other atoms 5) What is a chemical bond 6) Definition of three major types of chemical bonds, which one is the strongest 7) Why atoms are electrically neutral. What are ions, anions and cations 8) Definition of molecule and compound 9) Definition of major types of chemical reactions 10) Difference between organic and inorganic compounds. What are major organic and inorganic compounds Review 11) Properties of water 12) Deference between hydrophilic and hydrophobic compounds 13) What are electrolytes 14) What is pH, acidic PH, Basic PH, Neutral PH 15) What are acids, bases, buffers. 16) What are main organic molecules 17) The structures and functions of carbohydrates. 18) The structures and functions of lipids 19) The structures and functions of proteins. 20) What are enzymes, why they are important in metabolism 21) What are Exergonic (Exothermic) and Endergonic (Endothermic) Reactions 22) The structures and functions of nucleic acids 23) The structures and functions of high-energy compounds. Useful Links: https://www.youtube.com/watch?v=o-3I1JGW-Ck what is an atom? https://www.youtube.com/watch?v=QqjcCvzWwww ionic & covalent bonding https://www.youtube.com/watch?v=xdedxfhcpWo How water dissolves salt https://www.youtube.com/watch?v=Xeuyc55LqiY Acids bases, PH https://www.youtube.com/watch?v=rIvEvwViJGk PH, buffers http://www.youtube.com/watch?v=NJyAme5GVF8 How buffers work http://www.youtube.com/watch?v=H8WJ2KENlK0 Bilogical molecules https://www.youtube.com/watch?v=qBRFIMcxZNM What is a Protein? https://www.youtube.com/watch?v=r1ryDVgx0zw How enzymes work? http://www.youtube.com/watch?v=qy8dk5iS1f0 DNA structure http://www.youtube.com/watch?v=nhAcHuzfZfk THE HUMAN GENOME MUSIC PROJECT - CHROMOSOME 1 https://www.youtube.com/watch?v=FgzzQmOZzZk How life began MODULE 2.1 ATOMS AND ELEMENTS © 2016 Pearson Education, Inc.  CHEMISTRY AND PHYSIOLOGICAL REACTIONS Chemicals are the raw materials that make up ▪ our bodies, ▪ the bodies of other organisms, ▪ the physical environment. Chemistry – study of matter and its interactions ▪ underlies all physiological reactions: o Movement, digestion, pumping of heart, nervous system Chemistry can be broken down into: ▪ Basic chemistry & Biochemistry © 2016 Pearson Education, Inc.  MATTER What is the matter ? Matter – anything that has mass and occupies space; ▪ Matter can be seen, smelled, and/or felt ▪ Mass actual amount of matter ▪ Weight is ▪ mass plus the effects of gravity Atom – smallest unit of matter that retains original properties Gravity makes a 1 kilogram mass exert about 9.8 newtons of force © 2016 Pearson Education, Inc. States of matter Matter can exist in three possible states: ▪ Solid: definite shape and volume, high density (Bones) ▪ Liquid: changeable shape; definite volume, Medium density (Blood) ▪ Gas: changeable shape and volume, low density (the air we breath)  ENERGY The release and use of energy by living systems gives us the elusive quality we call it life Energy does not have mass, does not take up space Is the capacity to do work or put matter into motion or fuel chemical reactions Matter is substance, energy is the mover of substance. ▪ It measured by its affect on the matter The greater the work done, the more energy it uses up ENERGY AND CHEMICAL REACTIONS Two general forms of energy: ▪ Potential energy is stored; can be released to do work at some later time ▪ Kinetic energy is potential energy that has been released or set in motion to perform work; ▪ All atoms have kinetic energy as they are in constant motion; the faster they move the greater that energy Solid Fluid Gas © 2016 Pearson Education, Inc. Cautious !!!!!!!!!!!! Energy can be transformed from potential to kinetic energy Stored energy can be released, resulting in action  ENERGY Energy is found in different forms in the human body; ▪ Chemical energy o Stored in bonds of chemical substances ▪ Electrical energy o Results from movement of charged particles ▪ Mechanical energy o Directly involved in moving matter ▪ Radiant or electromagnetic energy o Travels in waves (example: heat, visible light, ultraviolet light, and X rays) The nerve system uses electrical current called nerve impulses to transmit message from one part of the body to another  ENERGY Energy form conversions ▪ Energy may be converted from one form to another o Example: turning on a lamp converts electrical energy to light energy ▪ Energy conversion is inefficient o Some energy is “lost” as heat, which can be partly unusable energy o Heat is a type of kinetic energy and is the random motion of atoms & molecules  ATOMS AND ATOMIC STRUCTURE Atom – smallest unit of matter that retains original properties Made up of even smaller structures called subatomic particles Figure 2.1 Structure of a representative atom. © 2016 Pearson Education, Inc.  ATOMS AND ATOMIC STRUCTURE Subatomic particles exist in 3 forms: ▪ Protons (p+) – found in central core of atom (atomic nucleus); positively charged ▪ Neutrons (n0) – found in atomic nucleus; slightly larger than protons; no charge. ▪ Electrons (e-) – found outside atomic nucleus; negatively charged Atoms are electrically neutral – ▪ They have no charge ▪ Number of protons and electrons are equal, cancelling each other’s charge; © 2016 Pearson Education, Inc.  ATOMS AND ATOMIC STRUCTURE Electron shells – ▪ Regions surrounding atomic nucleus where electrons exist ▪ Each can hold a certain number of electrons: o 1st shell (closest to nucleus) can hold 2 electrons o 2nd shell can hold 8 electrons o 3rd shell can hold 18 electrons but “satisfied” with 8 ▪ Some atoms may have more than 3 shells © 2016 Pearson Education, Inc.  ELEMENTS IN THE PERIODIC TABLE AND THE HUMAN BODY Element – substance that cannot be broken down into simpler substance by chemical means Atomic number defines every element Atomic number ; Number of protons that an atom has in its nucleus is its atomic number ▪ Each element is made of atoms with same number of protons © 2016 Pearson Education, Inc.  ELEMENTS IN THE PERIODIC TABLE AND THE HUMAN BODY The periodic table of elements lists elements by their increasing atomic numbers: ▪ Organizes elements into groups with certain properties ▪ Each element is represented by a chemical symbol o Based on The first letter or two of the element’s name » C for Calcium The Element’s Latin or German’s name » Na the Latin word for Natrium © 2016 Pearson Education, Inc.  ELEMENTS IN THE PERIODIC TABLE AND THE HUMAN BODY The human body is made up of ; 4 major elements(96%) o Hydrogen o Oxygen o Carbon o Nitrogen 7 mineral elements 13 trace elements © 2016 Pearson Education, Inc.  THE PERIODIC TABLE Figure 2.2 Elements in the human body and their positions in the periodic table. © 2016 Pearson Education, Inc. COMMON ELEMENTS COMPOSING THE HUMAN BODY © 2016 Pearson Education, Inc. COMMON ELEMENTS COMPOSING THE HUMAN BODY (CONTINUED) © 2016 Pearson Education, Inc. COMMON ELEMENTS COMPOSING THE HUMAN BODY (CONTINUED) © 2016 Pearson Education, Inc.  ISOTOPES AND RADIOACTIVITY Mass number – equal to sum of all protons and neutrons found in atomic nucleus Isotope – atom with same atomic number (same number of protons), but different mass number (different number of neutrons) Radioisotopes – unstable isotopes; (unstable nucleous) high energy or radiation released by radioactive decay; allows isotope to assume a more stable form ▪ Decay rate is expressed as half-life o The emission can brake molecules apart, destroy or damage cells. © 2016 Pearson Education, Inc. Hydrogen: is simplest atom Hydrogen atom contains one proton one electron (seldom contain neutron) All other atoms have both neutrons and protons in their nuclei Electron shell Hydrogen-2, Hydrogen-3, Hydrogen-1 deuterium tritium mass number: 1 mass number: 2 mass number: 3 A typical hydrogen A deuterium (2H) A tritium (3H) nucleus nucleus contains a nucleus contains a contains a pair of proton and no neutrons. proton and a neutron. neutrons in addition to the proton. THE STRUCTURE OF HYDROGEN ATOMS  RADIOISOTOPES Radioisotopes are used in biological research and medicine ▪ Share same chemistry as their stable isotopes so will be taken up by body o Can then be used for diagnosis of disease All radioactivity can damage living tissue ▪ Some types can be used to destroy localized cancers ▪ Some types cause cancer o Radon from uranium decay causes lung cancer  NUCLEAR MEDICINE Common applications of radioisotopes: Radiotracers – injected into patient and detected by camera; image analyzed by computer; shows size, shape, and activity of organs and cells Cancer radiation therapy – radiation damages structure of cancer cells; interferes with functions Treatment of thyroid disorders – high doses of iodine-131 treat overactive or cancerous thyroid tissue; radioisotope accumulates and damages cells © 2016 Pearson Education, Inc.  Radioactive tracers are frequently used in medical diagnosis. Sophisticated imaging instruments are used to detect them. – An imaging instrument that uses positron-emission tomography (PET) detects the location of injected radioactive materials. – PET is useful for diagnosing heart disorders and cancer and in brain research. In Alzheimer patient’s brain accumulate deposits of protein called beta-amyloid Protein molecule (PIB) contain radioactive isotope bind to beta-amyloid Healthy person Alzheimer’s patient  Dangers associated with using radioactive substances. – Uncontrolled exposure can cause damage to some molecules in a living cell, especially DNA. – Chemical bonds are broken by the emitted energy, which causes abnormal bonds to form.  MODULE 2.2 MATTER COMBINED: MIXTURES AND CHEMICAL BONDS © 2016 Pearson Education, Inc.  MATTER COMBINED Matter can be combined physically to form a Mixture. Mixture – atoms of two or more elements physically intermixed without changing chemical nature of atoms themselves There are 3 basic types of mixtures: ▪ Suspensions, ▪ Colloids, ▪ Solutions © 2016 Pearson Education, Inc.  MIXTURES Suspension – mixture containing two or more components with Large, unevenly distributed particles; Will settle out when left undisturbed Figure 2.3a The three types of mixtures. © 2016 Pearson Education, Inc.  MIXTURES Colloids – two or more components with o Small, evenly distributed particles o Will not settle out Figure 2.3b The three types of mixtures. © 2016 Pearson Education, Inc.  MIXTURES Solutions – two or more components with ▪ Extremely small, evenly distributed particles; ▪ Will not settle out; contain a solute dissolved in a solvent (no chemical change): usually translucent o Solute – substance that is dissolved, (solid, liquid, gas) o Solvent – substance that dissolves solute Water is the most important solvent in the body o We can separate the solute from solvent by physical means (such as Evaporation) Figure 2.3c The three types of mixtures. © 2016 Pearson Education, Inc.  CHEMICAL BONDS Matter can be combined chemically when atoms are combined by chemical bonds. A chemical bond is not a physical structure but rather an energy relationship or attractive force between atoms ▪ Molecule – formed by chemical bonding between two or more atoms of same element ▪ Compound – formed when two or more atoms from different elements combine by chemical bonding CH4 © 2016 Pearson Education, Inc.  CHEMICAL BONDS Macromolecules – very large molecules composed of many atoms O2 N2 CH4 Molecular formulas – represent molecules symbolically with letters and numbers; show kinds and numbers of atoms in a molecule Table 2.1 Electron Sharing in Covalent Bonds. © 2016 Pearson Education, Inc.  CHEMICAL BONDS Chemical bonds are formed when valence electrons (in outermost valence shell) of atoms interact Valence electrons determine how an atom interacts with other atoms and whether it will form bonds with a specific atom ▪ The octet rule states that an atom is most stable when it has 8 electrons in its valence shell (as in CO2) © 2016 Pearson Education, Inc.  THE DUET RULE ▪ The duet rule (for atoms with 5 or fewer electrons) states that an atom is most stable when its valence electron shell holds 2 electrons Figure 2.5 Formation of a covalent bond. © 2016 Pearson Education, Inc.  CHEMICAL BONDS A chemical bond can change an elements property Sodium Chloride Sodium chloride © 2016 Pearson Education, Inc.  CHEMICAL BONDS Ionic bond Covalent bonds © 2016 Pearson Education, Inc. Remember :Atoms are electrically neutral. ▪ Number of protons and electrons balance each other ION: Is an atom in which the total number of electrons is not equal to the total number of protons, giving the atom a net positive or negative electrical charge  IONS AND IONIC BONDS Ionic bond – formed when electrons are transferred from a metal atom to a nonmetal atom; ▪ results in formation of ions: cations and anions (Figure 2.4) ▪ Cation – positively charged ion; forms when metal loses one or more electrons ▪ Anion – negatively charged ion; forms when nonmetal gains one or more electrons The attraction between opposite charges bonds ions to one another forming a compound called a salt © 2016 Pearson Education, Inc.  IONIC BONDS Figure 2.4 Formation of an ionic bond. © 2016 Pearson Education, Inc. What is salt?  COVALENT BONDS Covalent bonds – strongest bond; form when two or more nonmetals share electrons (Figures 2.5, 2.6; Table 2.1) Two atoms can share one (single bond), two (double bond), or three (triple bond) electron pairs: Table 2.1 Electron Sharing in Covalent Bonds. © 2016 Pearson Education, Inc.  COVALENT BONDS Electronegativity; All elements have protons that attract electrons; property known as electronegativity The more electronegative an element the more strongly it attracts electrons, pulling them away from less electronegative elements © 2016 Pearson Education, Inc.  NONPOLAR COVALENT BONDS Nonpolar covalent bonds result when two nonmetals in a molecule with similar or identical electronegativities pull with equal force; therefore share electrons equally (Figure 2.6a) © 2016 Pearson Education, Inc.  POLAR COVALENT BONDS Polar covalent bonds form polar molecules when nonmetals with different electronegativities interact resulting in an unequal sharing of electrons (Figure 2.6b) ▪ Atom with higher electronegativity becomes partially negative (δ−) as it pulls shared electrons close to itself ▪ Atom with lower electronegativity becomes partially positive (δ+) as shared electrons are pulled toward other atom Polar molecules with partially positive and partially negative ends are known as dipoles © 2016 Pearson Education, Inc.  HYDROGEN BONDS Hydrogen bonds – weak attractions between partially positive hydrogen of one dipole and partially negative oxygen of another dipole Hydrogen bonds are generally between molecules Figure 2.7a Hydrogen bonding and surface tension between water molecules. © 2016 Pearson Education, Inc.  HYDROGEN BONDS Polar water molecules are more strongly attracted to one another than they are to nonpolar air molecules at surface Hydrogen bonds are responsible for a key property of water—surface tension Surface tension makes rain to fall in drops Important in the structure of proteins Blood sample to form as droplet Figure 2.7b Hydrogen bonding and surface tension between water molecules. © 2016 Pearson Education, Inc. Water's Surface tension created by hydrogenic bonds enable insects to walk on water The legs of water striders are hydrophobic MODULE 2.3 CHEMICAL REACTIONS © 2016 Pearson Education, Inc.  CHEMICAL NOTATION A chemical reaction has occurred every time a chemical bond is formed, broken, or rearranged, or when electrons are transferred between two or more atoms (or molecules) Chemical notation – series of symbols and abbreviations used to demonstrate what occurs in a reaction; the chemical equation (basic form of chemical notation) has two parts: ▪ Reactants on left side of equation are starting ingredients; will undergo reaction ▪ Products on right side of equation are results of chemical reaction o Reversible reactions can proceed in either direction as denoted by two arrows that run in opposite directions (as below) o Irreversible reactions proceed from left to right as denoted by a single arrow CO2 + H2O H2CO3 Reactants (carbon dioxide + water) Product (carbonic acid) © 2016 Pearson Education, Inc. ENERGY AND CHEMICAL REACTIONS Energy, inherent in all chemical bonds, must be invested any time a chemical reaction occurs: Exergonic reactions release excess energy so products have less energy than reactants Endergonic reactions require input of energy from another source; products contain more energy than reactants because energy was invested so reaction could proceed © 2016 Pearson Education, Inc.  HOMEOSTASIS AND TYPES OF CHEMICAL REACTIONS Three fundamental processes occur in the body to maintain homeostasis (breaking down molecules, converting the energy in food to usable form, and building new molecules); carried out by three basic types of chemical reactions: 1. Catabolic reactions (decomposition reactions) 2. Exchange reactions 3. Anabolic reactions (synthesis reactions) © 2016 Pearson Education, Inc.  HOMEOSTASIS AND TYPES OF CHEMICAL REACTIONS Catabolic reactions (decomposition reactions) – when a large substance is broken down into smaller substances General chemical notation for reaction is AB → A + B Usually exergonic because chemical bonds are broken © 2016 Pearson Education, Inc.  HOMEOSTASIS AND TYPES OF CHEMICAL REACTIONS Exchange reactions occur when one or more atoms from reactants are exchanged for one another General chemical notation for reaction is AB + CD → AD + BC HCL + NaOH → H2O + NaCL © 2016 Pearson Education, Inc.  HOMEOSTASIS AND TYPES OF CHEMICAL REACTIONS Anabolic reactions (synthesis reactions) occur when small simple subunits and united by chemical bonds to make large more complex substances General chemical notation for reaction is A + B → AB These reactions are endergonic; fueled by chemical energy © 2016 Pearson Education, Inc.  HOMEOSTASIS AND TYPES OF CHEMICAL REACTIONS Oxidation-reduction reactions (redox reactions) – special kind of exchange reaction; occur when electrons and energy are exchanged instead of atoms ▪ Reactant that loses electrons is oxidized ▪ Reactant that gains electrons is reduced Redox reactions are usually exergonic reactions capable of releasing large amounts of energy © 2016 Pearson Education, Inc. REACTION RATES AND ENZYMES Activation energy (Ea) For a reaction to occur atoms must collide with enough energy overcome the repulsion of their electrons This energy required for all chemical reactions is called the activation energy (Ea) Figure 2.8 Activation energy. © 2016 Pearson Education, Inc.  REACTION RATES AND ENZYMES Activation energy must be supplied so that reactants reach their transition states (i.e., get to the top of the energy “hill”) in order to react and form products (i.e., roll down the hill) Figure 2.8 Activation energy. © 2016 Pearson Education, Inc. REACTION RATES AND ENZYMES Factors increase reaction rate by reducing activation energy or increasing likelihood of strong collisions between reactants: ▪ Concentration ▪ Temperature ▪ Reactant properties ▪ Presence or absence of a catalyst © 2016 Pearson Education, Inc.  REACTION RATES AND ENZYMES When reactant concentration increases, more reactant particles are present, increasing chance of successful collisions between reactants Raising the temperature of the reactants increases kinetic energy of their atoms leading to more forceful and effective collisions between reactants © 2016 Pearson Education, Inc.  REACTION RATES AND ENZYMES Both particle size and phase (solid, liquid, or gas) influence reaction rates: ▪ Smaller particles move faster with more energy than larger particles ▪ Reactant particles in the gaseous phase have higher kinetic energy than those in either solid or liquid phase © 2016 Pearson Education, Inc.  REACTION RATES AND ENZYMES Catalyst – Substance that increases reaction rate by lowering activation energy without being consumed or altered in reaction Enzymes – biological catalysts; most are proteins with following properties: ▪ Speed up reactions by lowering the activation energy (Figure 2.9) ▪ Highly specific for individual substrates (substance that can bind to the enzyme’s active site) ▪ Do not alter the reactants or products ▪ Not permanently altered in reactions catalyzed © 2016 Pearson Education, Inc.  REACTION RATES AND ENZYMES Induced-fit mechanism – describes enzyme’s interaction with its substrate(s) ▪ Binding of substrate causes a small shape change that reduces energy of activation ▪ Allows transition state to proceed to final products Figure 2.10 Enzyme-substrate interaction. © 2016 Pearson Education, Inc.  ENZYME DEFICIENCIES Examples of common enzyme deficiencies: Tay-Sachs Disease – deficiency of hexosaminidase; gangliosides accumulate around neurons of brain; death usually by age 3 Severe Combined Immunodeficiency Syndrome (SCIDs) – may be due to adenosine deaminase deficiency; nearly complete absence of immune system; affected patients must live in sterile “bubble” Phenylketonuria – deficiency of phenylalanine hydroxylase; converts phenylalanine into tyrosine; resulting seizures and mental retardation can be prevented by dietary modification © 2016 Pearson Education, Inc.  MODULE 2.4 INORGANIC COMPOUNDS: WATER, ACIDS, BASES, AND SALTS BONDS © 2016 Pearson Education, Inc.  BIOCHEMISTRY Biochemistry – the chemistry of life Inorganic compounds generally do not contain carbon bonded to hydrogen; Include water, acids, bases, CO2 and salts Organic compounds – those that do contain carbon bonded to hydrogen © 2016 Pearson Education, Inc.  ORGANIC AND INORGANIC COMPOUNDS Inorganic Compounds ▪ Carbon dioxide, oxygen, water, and inorganic acids, bases, and salts (CO2, O2, H2O,NaCl, HCl, NaOH) Organic Compounds The chemical compounds of living things are known as organic compounds because of their association with organisms and because they are carbon- containing compounds. Typically associated with H, O and less commonly with nitrogen, phosphorus, sulfur, iron, other elements  Inorganic Compounds Water (H2O) makes up 60–80% of mass of human body and has several key properties vital to our existence (Figure 2.11): 1) High heat capacity – able to absorb heat without significantly changing temperature itself 2) Carries heat with it when it evaporates (when changing from liquid to gas) 3) Cushions and protects body structures because of relatively high density 4) Acts as a lubricant between two adjacent surfaces (reduces friction) 5) Water serves as body’s primary solvent; © 2016 Pearson Education, Inc.  WATER Water serves as body’s primary solvent; often called the universal solvent because so many solutes will dissolve in it entirely or to some degree (Figure 2.11) Water is a polar covalent molecule: ▪ Oxygen pole – partially negative (δ−) ▪ Hydrogen pole – partially positive (δ+) Allows water molecules to interact with certain solutes, surround them, and keep them apart © 2016 Pearson Education, Inc.  WATER Water is only able to dissolve hydrophilic solutes (those with fully or partially charged ends); “like dissolves like”, so water dissolves ionic and polar covalent solutes Figure 2.11a, b The behavior of hydrophilic and hydrophobic molecules in water. © 2016 Pearson Education, Inc.  WATER Solutes that do not have full or partially charged ends are hydrophobic; do not dissolve in water; includes uncharged nonpolar covalent molecules such as oils and fats Figure 2.11c The behavior of hydrophilic and hydrophobic molecules in water. © 2016 Pearson Education, Inc.  ACIDS AND BASES The study of acids and bases is really the study of the hydrogen ion (H+) Water molecules in solution may dissociate (break apart) into positively charged hydrogen ions (H+) and negatively charged hydroxide ions (OH−) Acids and bases are defined according to their behavior with respect to hydrogen ions (next slide) Figure 2.12a The behavior of acids and bases in water. © 2016 Pearson Education, Inc.  ACIDS AND BASES Acid – hydrogen ion or proton donor; number of hydrogen ions increases in water when acid is added ▪ HCl ⇄ H ++Cl- Base (alkali) –proton acceptor;number of hydrogen ions decreases in water when base is added (Figure 2.12c) © 2016 Pearson Education, Inc.  ACIDS AND BASES pH scale – ranges from 0–14 (Figure 2.13) Simple way of representing hydrogen ion concentration of a solution Literally the negative logarithm of the hydrogen ion concentration: pH = – Log [H+] © 2016 Pearson Education, Inc. CONCEPT BOOST: MAKING SENSE OF THE PH SCALE Why does pH decrease if solution has more hydrogen ions? The smaller the pH number, the bigger its negative log Single-digit changes in negative logarithm (e.g., from 2 to 3) accompanies a 10-fold change in hydrogen ion concentration (e.g., from 0.01 to 0.001) © 2016 Pearson Education, Inc. CONCEPT BOOST: MAKING SENSE OF THE PH SCALE Example: ▪ Solution A has a hydrogen ion concentration of 0.015 M and a pH of 1.82; solution B has a hydrogen ion concentration of 0.0003 M and a pH of 3.52 ▪ The solution with the higher hydrogen ion concentration has the lower −log. For this reason, the more acidic a solution, the lower its pH, and vice-versa © 2016 Pearson Education, Inc.  ACIDS AND BASES When pH = 7 the solution is neutral where the number of hydrogen ions and base ions are equal A solution with pH less than 7 is acidic; hydrogen ions outnumber base ions A solution with pH greater than 7 is basic or alkaline; base ions outnumber hydrogen ions. Figure 2.13 The pH Scale. © 2016 Pearson Education, Inc.  ACIDS AND BASES Buffer – chemical system that resists changes in pH; ▪ prevents large swings in pH when acid or base is added to a solution. A buffer consists of a weak acid and its corresponding anion. ▪ H2CO3⇄H+ + HCO3− ▪ Carbonic acid⇄+Hydrogen ion+HCO3−Bicarbonate ion Blood pH must remain within its narrow range to maintain homeostasis Most body fluids are slightly basic: ▪ Blood pH is 7.35–7.45 ▪ Intracellular pH is 7.2 Figure 2.13 The pH Scale. © 2016 Pearson Education, Inc.  SALTS AND ELECTROLYTES Salt – any metal cation and nonmetal anion held together by ionic bonds Salts can dissolve in water to form cations and anions called electrolytes which are capable of conducting electrical current Figure 2.4 and Figure 2.11a © 2016 Pearson Education, Inc.  MODULE 2.5 ORGANIC COMPOUNDS: CARBOHYDRATES, LIPIDS, PROTEINS, AND NUCLEOTIDES © 2016 Pearson Education, Inc. TABLE 2-3 IMPORTANT FUNCTIONAL GROUPS OF ORGANIC COMPOUNDS  MONOMERS AND POLYMERS Each type of organic compound in body (carbohydrate, lipid, protein, or nucleic acid) consists of polymers built from monomer subunits: Monomers are single subunits that can be combined to build larger structures called polymers by dehydration synthesis (anabolic reaction that links monomers together and makes a molecule of water in process) Hydrolysis is a catabolic reaction that uses water to break up polymers into smaller subunits © 2016 Pearson Education, Inc.  CARBOHYDRATES Carbohydrates, composed of carbon, hydrogen, and oxygen, function primarily as fuel; some limited structural roles Main function: Resource of energy , structural, ▪ Monosaccharides – consist of 3 to 7 a The structural formula of b The structural formula of the ring form, the most the straight-chain form common form of glucose carbons; KEY = Carbon ▪ Monomers: glucose, fructose, = Oxygen = Hydrogen galactose, ribose, and dexoyribose are c most abundant monosaccharides (Figure 2.14) © 2016 Pearson Education, Inc.  CARBOHYDRATES Figure 2.14 Carbohydrates: structure of monosaccharides. © 2016 Pearson Education, Inc.  CARBOHYDRATES ▪ Disaccharides are formed by union of two monosaccharides by dehydration synthesis Figure 2.15 Carbohydrates: formation and breakdown of disaccharides. © 2016 Pearson Education, Inc.  CARBOHYDRATES Polysaccharides consist of many monosaccharides joined to one another by dehydration synthesis reactions (Figure 2.16) Glycogen is the storage polymer of glucose; mostly in skeletal muscle and liver cells Starch: from plants Cellulose: ( in human can not be digested and used as a source of energy) Some polysaccharides are found covalently bound to either proteins or lipids forming glycoproteins and glycolipids; various functions in body © 2016 Pearson Education, Inc.  CARBOHYDRATES Figure 2.16 Carbohydrates: the polysaccharide glycogen. © 2016 Pearson Education, Inc.  LIPIDS Lipids – group of nonpolar hydrophobic molecules composed primarily of carbon and hydrogen; include fats and oils Fatty acids – lipid monomers in some lipids consisting of 4 to 20 carbon atoms; may have none, one, or more double bonds between carbons in hydrocarbon chain (Figure 2.17) © 2016 Pearson Education, Inc.  LIPIDS Saturated fatty acids – Solid at room temperature; have no double bonds between carbon atoms so carbons are “saturated” with maximum number of hydrogen atoms Figure 2.17a Lipids: structure of fatty acids. © 2016 Pearson Education, Inc.  LIPIDS ▪ Monounsaturated fatty acids – generally liquid at room temperature; have one double bond between two carbons in hydrocarbon chain Figure 2.17b Lipids: structure of fatty acids. © 2016 Pearson Education, Inc.  LIPIDS Polyunsaturated fatty acids – liquid at room temperature; have two or more double bonds between carbons in hydrocarbon chain Figure 2.17c Lipids: structure of fatty acids. © 2016 Pearson Education, Inc.  THE GOOD, THE BAD, AND THE UGLY OF FATTY ACIDS Not all fatty acids were created equally: The Good: Omega – 3 Fats ▪ Found in flaxseed oil and fish oil but cannot be made by humans; must be obtained in diet ▪ Polyunsaturated; positive effects on cardiovascular health The Bad: Saturated Fats ▪ Found in animal fats; also in palm and coconut oils ▪ Overconsumption associated with increased cardiac disease risk © 2016 Pearson Education, Inc.  THE GOOD, THE BAD, AND THE UGLY OF FATTY ACIDS Not all fatty acids were created equally (continued): The Ugly: Trans Fats ▪ Produced by adding H atoms to unsaturated plant oils (“partially hydrogenated oils”) ▪ No safe consumption level; significantly increase risk of heart disease © 2016 Pearson Education, Inc.  LIPIDS Triglyceride – Three fatty acids linked by dehydration synthesis to a modified 3-carbon carbohydrate, glycerol; Storage polymer for fatty acids (also called a neutral fat) Figure 2.18 Lipids: structure and formation of triglycerides. © 2016 Pearson Education, Inc.  LIPIDS Important functions of Triglycerides 1. Energy source 2. Insulation ( Fat deposit under skin) 3. Protection( Fat deposit under certain organs such as kidney) 4. Stored in body as lipid droplets within cells, which accumulate lipid-soluble vitamins, drugs or toxins © 2016 Pearson Education, Inc.  LIPIDS Phospholipids – composed of a glycerol backbone, two fatty acid “tails” and one phosphate “head” in place of third fatty acid (Figure 2.19) A molecule with a polar group (phosphate head) and a nonpolar group (fatty acid tail) is called amphiphilic This amphiphilic nature makes phospholipids vital to the structure of cell membranes Table 2.3 Organic Molecules. © 2016 Pearson Education, Inc.  LIPIDS Figure 2.19 Lipids: structure of phospholipids. © 2016 Pearson Education, Inc.  LIPIDS Steroids – nonpolar and share a four-ring hydrocarbon structure called the steroid nucleus Cholesterol – steroid that forms basis for all other steroids Figure 2.20 Lipids: structure of steroids and Table 2.3 Organic Molecules. © 2016 Pearson Education, Inc. Types of Steroids Cholesterol from animal products, synthesized in body Component of plasma (cell) membranes Estrogens and testosterone Sex hormones Corticosteroids and calcitriol Cholesterol Metabolic regulation, mineral balance Bile salts (Derived from steroids in liver) Required for the normal processing of dietary fats Estrogen Testosterone  PROTEINS Proteins are macromolecules that: ▪ Function as enzymes ▪ Play structural roles ▪ Are involved in movement ▪ Function in the body’s defenses ▪ Can be used as fuel © 2016 Pearson Education, Inc.  Protein Structure Long chains of amino acids (building blocks) Five components of amino acid structure 1. Central carbon atom Amino group Central carbon 2. Hydrogen atom Carboxyl group 3. Amino group (—NH2) 4. Carboxyl group (—COOH) R group (variable side Chain of one or more atoms) 5. Variable side chain or R group  PROTEINS Twenty different amino acids (monomers of all proteins); Can be linked by peptide bonds into polypeptides Figure 2.21a, b Proteins: structure of amino acids. © 2016 Pearson Education, Inc.  PROTEINS Peptides – formed from two or more amino acids linked together by peptide bonds through dehydration synthesis: Figure 2.22 Proteins: formation and breakdown of dipeptides. © 2016 Pearson Education, Inc.  PROTEINS Dipeptides consist of two amino acids, Tripeptides three amino acids, Polypeptides contain 10 or more amino acids Proteins consist of one or more polypeptide chains folded into distinct structures which must be maintained to be functional © 2016 Pearson Education, Inc.  PROTEINS Two basic types of proteins classified according to structure: fibrous and globular Fibrous proteins – long rope-like strands; composed mostly of nonpolar amino acids; link things together and add strength and durability to structures Globular proteins – spherical or globe-like; composed mostly of polar amino acids; function as enzymes, hormones, and other cell messengers Table 2.3 Organic Molecules. © 2016 Pearson Education, Inc.  PROTEINS Complex structure of a complete protein is divided into four levels: Primary structure – amino acid sequence of polypeptide chain Figure 2.23a Levels of protein structure. © 2016 Pearson Education, Inc.  PROTEINS Complex structure of a complete protein (continued): Secondary structure – one or more segments of primary structure folded in specific ways; held together by hydrogen bonds ▪ Alpha helix – coiled spring ▪ Beta-pleated sheet – Venetian blind Figure 2.23b Levels of protein structure. © 2016 Pearson Education, Inc.  PROTEINS Complex structure of a complete protein (continued): Tertiary structure – three-dimensional shape that peptide chain assumes (twists, folds, and coils including secondary structure); stabilized by hydrogen bonding A1 A2 A3 A4 A5 A6 A7 A8 A9 Linear chain of amino acids A1 A2 A3 A4 Hydrogen Hydrogen bond A5 bond A9 A8 A7 A6 A2 A6 A10 A1 A3 A5 A7 A9 OR A11 A12 A13 A14 Alpha-helix Pleated sheet OR Heme units Hemoglobin Keratin or collagen (globular protein) (fibrous protein) © 2016 Pearson Education, Inc.  PROTEINS Complex structure of a complete protein (continued): Quaternary structure – linking together more than one polypeptide chain in a specific arrangement; critical to function of protein as a whole Figure 2.23d Levels of protein structure. © 2016 Pearson Education, Inc.  PROTEINS Figure 2.23 Levels of protein structure. © 2016 Pearson Education, Inc.  PROTEINS Protein denaturation – process of destroying a protein’s shape by heat, pH changes, or exposure to chemicals Disrupts hydrogen bonding and ionic interactions that stabilize structure and function. © 2016 Pearson Education, Inc.  NUCLEOTIDES AND NUCLEIC ACIDS Nucleotides – monomers of nucleic acids; named because of abundance in nuclei of cells; make up genetic material Nucleotide structure: ▪ Nitrogenous base with a hydrocarbon ring structure ▪ Five-carbon pentose sugar, ribose or deoxyribose ▪ Phosphate group Figure 2.24a Structure of nucleotides. © 2016 Pearson Education, Inc.  NUCLEOTIDES AND NUCLEIC ACIDS Two types of nitrogenous bases: purines and pyrimidines Purines – double-ringed molecule; adenine (A) and guanine (G) Pyrimidines – single-ringed molecule; cytosine (C), uracil (U) and thymine (T) Figure 2.24 Structure of nucleotides. © 2016 Pearson Education, Inc.  NUCLEOTIDES AND NUCLEIC ACIDS Adenosine triphosphate (ATP) Adenine attached to ribose and three phosphate groups; Main source of chemical energy in body Synthesized from adenosine diphosphate (ADP) and a phosphate group (Pi) using energy from oxidation of fuels (like glucose) Figure 2.25a Nucleotides: structure and formation of ATP. © 2016 Pearson Education, Inc.  NUCLEOTIDES AND NUCLEIC ACIDS Adenosine triphosphate (continued): Potential energy in this “high-energy” bond can be released as kinetic energy to do work Production of large quantities of ATP requires oxygen; why we breathe air Figure 2.25b Nucleotides: structure and formation of ATP. © 2016 Pearson Education, Inc. ATP is a molecule which carries the most useful form of chemical energy in living systems Food molecules can not be used to energize body activity directly Some of the food energy is captured temporary in the bonds of ATP and the rest is converted to heat  NUCLEOTIDES AND NUCLEIC ACIDS DNA, an extremely large molecule found in nuclei of cells; composed of two long chains that twist around each other to form a double helix DNA contains genes – provide recipe or code for protein synthesis – process of making every protein Figure 2.26a Structure of nucleic acids and Table 2.3 Organic Molecules. © 2016 Pearson Education, Inc.  NUCLEOTIDES AND NUCLEIC ACIDS Other structural features of DNA include: DNA contains: ▪ Pentose sugar deoxyribose (lacks oxygen-containing group of ribose) forms backbone of strand; alternates with phosphate group ▪ Bases: adenine, guanine, cytosine, and thymine Figure 2.26a Structure of nucleic acids and Table 2.3 Organic Molecules. © 2016 Pearson Education, Inc.  NUCLEOTIDES AND NUCLEIC ACIDS Other structural features of DNA include: Double helix strands – held together by hydrogen bonding between the bases of each strand Each base faces the inside of the double helix as strands run in opposite directions. Figure 2.26a Structure of the nucleic acids DNA and RNA. © 2016 Pearson Education, Inc.  NUCLEOTIDES AND NUCLEIC ACIDS Other structural features of DNA (continued): DNA exhibits complementary base pairing; purine A always pairs with pyrimidine T and purine G always pairs with pyrimidine C A = T (where = denotes 2 hydrogen bonds) and C  G (where  denotes 3 hydrogen bonds) Figure 2.26a Structure of the nucleic acids DNA and RNA. © 2016 Pearson Education, Inc.  NUCLEOTIDES AND NUCLEIC ACIDS RNA – single strand of nucleotides; can move between nucleus of cell and cytosol; critical to making proteins RNA contains the pentose sugar ribose RNA contains uracil instead of thymine; still pairs with adenine, (A = U) Figure 2.26b Structure of the nucleic acids DNA and RNA. © 2016 Pearson Education, Inc.  NUCLEOTIDES AND NUCLEIC ACIDS RNA –single strand of nucleotides (continued) RNA copies recipe for specific protein (gene in DNA); process called transcription RNA exits nucleus to protein synthesis location; then directs the making of protein from recipe; process called translation Figure 2.26b Structure of the nucleic acids DNA and RNA. © 2016 Pearson Education, Inc.  NUCLEOTIDES AND NUCLEIC ACIDS Figure 2.26 Structure of the nucleic acids DNA and RNA. © 2016 Pearson Education, Inc.