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PowerPoint® Lecture Slides prepared by Barbara Heard, Atlantic Cape Community College CHAPTER 2 Chemistry Comes Alive: Modified by Dr. Par Mohammadian © Annie Leibovitz/Contact Press Images © 2013 Pearson Education, Inc. Overview • Part 1: Basic Chemistry – Matter – Energy – Atoms & Elements – Chemical Bonds – Functional Groups • Part 2: Biochemistry – Bio-macromolecules: • Carbohydrates • Lipids • Proteins – Enzymes • Nucleic acids Body Chemistry A general understanding of chemistry is necessary for understanding human physiology. Why? Name specific examples! – Physiological processes are based on chemical interactions. Matter • Matter—anything that has mass & occupies space • 3 states of matter – Solid—definite shape and volume – Liquid—changeable shape; definite volume – Gas—changeable shape and volume © 2013 Pearson Education, Inc. Energy • Definition: – Capacity to do work or put matter into motion • Types of energy – Kinetic—energy in action Name an example! – Potential—stored (inactive) energy • Energy can be transferred from potential to kinetic energy Forms of Energy Name an example for each! • Chemical energy – Stored in bonds of chemical substances • Electrical energy – Results from movement of charged particles • Mechanical energy – Directly involved in moving matter • Radiant or electromagnetic energy – Travels in waves (e.g., visible light, UV light, and x-rays) Composition of Matter: Elements • Elements Examples? – Matter is composed of elements – Elements cannot be broken into simpler substances by ordinary chemical methods • Atoms – Building blocks for each element © 2013 Pearson Education, Inc. Major Elements of the Human Body • Four elements make up 96.1% of body mass Element Carbon Hydrogen Oxygen Nitrogen © 2013 Pearson Education, Inc. Atomic symbol C H O N Lesser Elements of the Human Body 9 elements make up 3.9% of body mass Element Calcium Phosphorus Potassium Sulfur Sodium Chlorine Magnesium Iodine Iron © 2013 Pearson Education, Inc. Atomic symbol Ca P K S Na Cl Mg I Fe Atoms – Subatomic particles • An atom is the smallest unit of an element. It has: – A nucleus with positively charged protons and uncharged neutrons – Orbiting electrons with negative charges – An atomic mass equal to the number of protons plus the number of neutrons – An atomic number equal to the number of protons Electron Orbitals • Orbitals (or shells) are energy levels that surround the nucleus of an atom. • Electrons fill the shells, starting with the one closest to the nucleus. – The first shell holds 2 electrons. – Each shell thereafter holds 8 electrons. (Non-biological elements fill distant shells that hold more than 8.) – Atoms are most stable when the outer shell is filled. Electrons in unfilled outer shells participate in bonding; they are called valence electrons. Figure 2.1 Two models of the structure of an atom. Nucleus Nucleus Helium atom Helium atom 2 protons (p+) 2 neutrons (n0) 2 electrons (e–) 2 protons (p+) 2 neutrons (n0) 2 electrons (e–) Planetary model Proton Neutron Electron © 2013 Pearson Education, Inc. Electron cloud Orbital model Identifying Elements • Different elements contain different numbers of subatomic particles – Hydrogen has 1 proton, 0 neutrons, and 1 electron – Lithium has 3 protons, 4 neutrons, and 3 electrons © 2013 Pearson Education, Inc. Identifying Elements: • Atomic number = # of protons in nucleus – Written as subscript to left of atomic symbol • Ex. 3Li • Mass number – Total # of protons and neutrons in nucleus • Total mass of atom – Written as superscript to left of atomic symbol • Ex. 7Li © 2013 Pearson Education, Inc. Ions and Isotopes • An atom that loses or gains electrons becomes an ion: – Cation; charge: + – Anion; charge: • An atom that loses or gains protons becomes a different element. • An atom that gains or loses neutrons becomes an isotope of the same element. Radioisotopes • Valuable tools for biological research and medicine – Share same chemistry as their stable isotopes – Most used for diagnosis • All damage living tissue – Some used to destroy localized cancers – Radon from uranium decay causes lung cancer © 2013 Pearson Education, Inc. Free Radicals •Free radicals: Unstable; have at least one unpaired electron; caused in our body by natural (sun) and/or manufactured •Sources (microwave, TV); e.g peroxide ion .O2•Can contribute to diseases (cancer) and aging! •Antioxidants (e.g. Vit C and E) prevent damage by releasing electrons without becoming free radicals! Combining Matter: Molecules and Compounds • Most atoms chemically combined with other atoms to form molecules and compounds – Molecule • Two or more atoms bonded together (e.g., H2 or C6H12O6) • Smallest particle of a compound with specific characteristics of the compound – Compound • Two or more different kinds of atoms bonded together (e.g., C6H12O6 , but not H2) © 2013 Pearson Education, Inc. Mixtures • Most matter exists as mixtures – Two or more components physically intermixed • Three types of mixtures – Solutions – Colloids – Suspensions © 2013 Pearson Education, Inc. Types of Mixtures: Solutions • Homogeneous mixtures • Most are true solutions in body – Gases, liquids, or solids dissolved in water – Usually transparent, e.g., atmospheric air or saline solution • Solvent – Substance present in greatest amount – Usually a liquid; usually water • Solute(s) – Present in smaller amounts • Ex. If glucose is dissolved in blood, glucose is solute; blood is solvent © 2013 Pearson Education, Inc. Concentration of True Solutions • Can be expressed as – Percent (%) of solute in total solution (assumed to be water) Prepare 0.9% NaCl solution! • Parts solute per 100 parts solvent – Milligrams per deciliter (mg/dl) – Molarity, or moles per liter (M) • 1 mole of an element or compound = Its atomic or molecular weight (sum of atomic weights) in grams • 1 mole of any substance contains 6.02 1023 molecules of that substance (Avogadro’s number) Use the example in the book (pg 30) to explain! © 2013 Pearson Education, Inc. Colloids and Suspensions • Colloids (AKA emulsions) – Heterogeneous mixtures, e.g., cytosol – Large solute particles do not settle out • Suspensions – Heterogeneous mixtures, e.g., blood – Large, visible solutes settle out © 2013 Pearson Education, Inc. Figure 2.4 The three basic types of mixtures. Solution Colloid Suspension Solute particles are very tiny, do not settle out or scatter light. Solute particles are larger than in a solution and scatter light; do not settle out. Solute particles are very large, settle out, and may scatter light. Solute particles Solute particles Solute particles Example Example Example Mineral water Jello Blood Plasma Settled red blood cells Unsettled Settled © 2013 Pearson Education, Inc. Mixtures versus Compounds • Mixtures – No chemical bonding between components – Can be separated by physical means, such as straining or filtering – Heterogeneous or homogeneous • Compounds – Chemical bonding between components – Can be separated only by breaking bonds – All are homogeneous © 2013 Pearson Education, Inc. Chemical Bonds • A molecule forms when electrons of several atoms interact to form chemical bonds. – The number of bonds an atom can form is determined by the number of valence electrons. • Hydrogen has one electron; it needs one more to fill the inner shell so that it can form one bond. • Carbon has 6 electrons; 2 fill the inner shell and 4 are in the next shell. It needs 4 more electrons so that it can form 4 bonds. Figure 2.5a Chemically inert and reactive elements. Chemically inert elements Outermost energy level (valence shell) complete 2e Helium (He) (2p+; 2n0; 2e–) © 2013 Pearson Education, Inc. 8e 2e Neon (Ne) (10p+; 10n0; 10e–) Figure 2.5b Chemically inert and reactive elements. Chemically reactive elements • Valence shell not full • Tend to gain, lose, or share electrons (form bonds) with other atoms to achieve stability Outermost energy level (valence shell) incomplete 1e Hydrogen (H) (1p+; 0n0; 1e–) 6e 2e Oxygen (O) (8p+; 8n0; 8e–) © 2013 Pearson Education, Inc. 4e 2e Carbon (C) (6p+; 6n0; 6e–) 1e 8e 2e Sodium (Na) (11p+; 12n0; 11e–) Types of Chemical Bonds • Three major types – Ionic bonds – Covalent bonds – Hydrogen bonds © 2013 Pearson Education, Inc. Ionic Bonds • Ions – Atom gains or loses electrons and becomes charged • # Protons ≠ # Electrons • Transfer of valence shell electrons from one atom to another forms ions – One becomes an anion (negative charge) • Atom that gained one or more electrons – One becomes a cation (positive charge) • Atom that lost one or more electrons • Attraction of opposite charges results in an ionic bond © 2013 Pearson Education, Inc. Figure 2.6a–b Formation of an ionic bond. + Sodium atom (Na) (11p+; 12n0; 11e–) Chlorine atom (Cl) (17p+; 18n0; 17e–) Sodium gains stability by losing one electron, and chlorine becomes stable by gaining one electron. © 2013 Pearson Education, Inc. Sodium ion (Na+) — Chloride ion (Cl–) Sodium chloride (NaCl) After electron transfer, the oppositely charged ions formed attract each other. Figure 2.6c Formation of an ionic bond. • Most ionic compounds are salts – When dry salts form crystals instead of individual molecules – Example is NaCl (sodium chloride) Cl– Na+ Large numbers of Na+ and Cl– ions associate to form salt (NaCl) crystals. © 2013 Pearson Education, Inc. Covalent Bonds • Formed by sharing of two or more valence shell electrons • Allows each atom to fill its valence shell at least part of the time Reacting atoms Resulting molecules or + Nitrogen atom Nitrogen atom Formation of a triple covalent bond: Two nitrogen atoms share three electron pairs. © 2013 Pearson Education, Inc. Structural formula shows triple bond. Molecule of nitrogen gas (N2) Figure 2.7a Formation of covalent bonds. Reacting atoms Resulting molecules + or Structural formula shows single bonds. Carbon atom Hydrogen atoms Formation of four single covalent bonds: Carbon shares four electron pairs with four hydrogen atoms. Reacting atoms Molecule of methane gas (CH4) Resulting molecules + Oxygen atom Oxygen atom Formation of a double covalent bond: Two oxygen atoms share two electron pairs. © 2013 Pearson Education, Inc. or Structural formula shows double bond. Molecule of oxygen gas (O2) Nonpolar Covalent Bonds • Electrons shared equally • Produces electrically balanced, nonpolar molecules such as CO2 Carbon dioxide (CO2) molecules are linear and symmetrical. They are nonpolar. © 2013 Pearson Education, Inc. Polar Covalent Bonds • Unequal sharing of electrons produces polar (AKA dipole) molecules such as H2O • Small atoms with six or seven valence shell electrons are electronegative, e.g., oxygen • Most atoms with one or two valence shell electrons are electropositive, e.g., Na – + + V-shaped water (H2O) molecules have two poles of charge—a slightly more negative oxygen end (–) and a slightly more positive hydrogen end (+). Hydrogen Bonds • + – – Not true bond Hydrogen bond – Common between dipoles (indicated by such as water dotted line) + – – + – + + Attractive force between electropositive hydrogen of one molecule and an electronegative atom of another molecule – Also act as intramolecular bonds, holding a large molecule in a threedimensional shape + – The slightly positive ends (+) of the water molecules become aligned with the slightly negative ends (–) of other water molecules. Functional Groups Classes of molecules are named after their functional group. Chemical Reactions (reactants) A+B -> AB (product) Synthesis reactions Smaller particles are bonded together to form larger, more complex molecules. Example Amino acids are joined together to form a protein molecule. Decomposition reactions Bonds are broken in larger molecules, resulting in smaller, less complex molecules. Exchange reactions Bonds are both made and broken (also called displacement reactions). Example Glycogen is broken down to release glucose units. Example ATP transfers its terminal phosphate group to glucose to form glucosephosphate. + Amino acid molecules Glycogen Adenosine triphosphate (ATP)Glucose Protein molecule + Glucose molecules Adenosine diphosphate (ADP) Glucosephosphate Energy Flow in Chemical Reactions • All chemical reactions are either exergonic or endergonic – Exergonic reactions—net release of energy • Products have less potential energy than reactants – Endergonic reactions—net absorption of energy • Products have more potential energy than reactants © 2013 Pearson Education, Inc. Reversibility of Chemical Reactions • All chemical reactions are theoretically reversible – A + B AB – AB A + B • Chemical equilibrium occurs if neither a forward nor reverse reaction is dominant © 2013 Pearson Education, Inc. Rate of Chemical Reactions • Affected by – Temperature Rate – Concentration of reactant Rate – Particle size Rate – Catalysts: Rate without being chemically changed or part of product • Enzymes are biological catalysts © 2013 Pearson Education, Inc. Part 2: Biochemistry Classes of Compounds • Inorganic compounds • *Water, **salts, and many acids and bases • Do not contain carbon • Organic compounds • Carbohydrates, fats, proteins, and nucleic acids • Contain carbon, usually large, and are covalently bonded • Both equally essential for life *part of hydrolysis and dehydration synthesis reactions **Common salts in body: NaCl, CaCO3, KCl, calcium phosphates; Ions (electrolytes) conduct electrical currents in solution Dehydration synthesis Monomers are joined by removal of OH from one monomer and removal of H from the other at the site of bond formation. + Monomer 2 Monomer 1 Monomers linked by covalent bond Hydrolysis Monomers are released by the addition of a water molecule, adding OH to one monomer and H to the other. Monomer 1 + Monomers linked by covalent bond Example reactions Dehydration synthesis of sucrose and its breakdown by hydrolysis Water is released + Water is consumed Glucose Fructose Sucrose Monomer 2 Acids, Bases, and pH • Sometimes a solution has more H+ ions than OH- ions ready to release: acid, pH < 7. • Acids are proton donors – Release H+ (a bare proton) in solution – HCl H+ + Cl– • Sometimes a solution has more OH- ions than H+ ions ready to release: base (alkaline), pH > above 7. Often called a proton acceptor (Take up H+ from solution) • NaOH Na+ + OH– OH– accepts an available proton (H+) OH– + H+ H2O Some Important Acids and Bases in Body • Important acids – HCl, HC2H3O2 (HAc), and H2CO3 • Important bases – Bicarbonate ion (HCO3–) and ammonia (NH3) © 2013 Pearson Education, Inc. Concentration (moles/liter) [OH−] [H+] pH pH = - log [H+] • Pure water has a concentration of 10-7, so the pH is 7. A pH 6 solution actually has 10 times the number of H+ ions. H+ 100 10−14 14 1M Sodium hydroxide (pH=14) 10−1 10−13 13 Oven cleaner, lye (pH=13.5) 10−2 10−12 12 10−3 10−11 11 10−4 10−10 10 10−5 10−9 9 10−6 10−8 8 10−7 10−7 7 Neutral 10−8 10−6 6 10−9 10−5 5 10−10 10−4 4 10−11 10−3 3 10−12 10−2 2 10−13 10−1 1 10−14 100 0 Increasingly basic From 0 to 14: • 0: the strongest acid • 14: strongest base. Examples Household ammonia (pH=10.5–11.5) Household bleach (pH=9.5) Egg white (pH=8) Blood (pH=7.4) Milk (pH=6.3–6.6) Increasingly acidic pH Scale Black coffee (pH=5) Wine (pH=2.5–3.5) Lemon juice; gastric juice (pH=2) 1M Hydrochloric acid (pH=0) pH: Acid-base Concentration – Relative free [H+] of a solution measured on pH scale – As free [H+] increases, acidity increases • [OH–] decreases as [H+] increases • pH decreases – As free [H+] decreases alkalinity increases • [OH–] increases as [H+] decreases • pH increases © 2013 Pearson Education, Inc. Neutralization • Results from mixing acids and bases – Displacement reactions occur forming water and a salt – Neutralization reaction • Joining of H+ and OH– to form water neutralizes solution Example? © 2013 Pearson Education, Inc. Buffers • Buffers stabilize H+ concentration in a solution. – In blood, two molecules stabilize pH: bicarbonate ion (HCO3-) and carbonic acid (H2CO3). HCO3- + H+ ↔ H2CO3 • If blood falls below pH 7.35, the condition is called acidosis. • If blood rises above pH 7.45, the condition is called alkalosis. Buffers • Buffers resist abrupt and large swings in pH – Release hydrogen ions if pH rises – Bind hydrogen ions if pH falls • Convert strong (completely dissociated) acids or bases into weak (slightly dissociated) ones • Carbonic acid-bicarbonate system (important buffer system of blood): © 2013 Pearson Education, Inc. Carbohydrates • Sugars and starches Organic compounds: • Polymers • Contain C, H, and O [(CH20)n] • Three classes – Monosaccharides – one sugar – Disaccharides – two sugars – Polysaccharides – many sugars • Functions of carbohydrates – Major source of cellular fuel (e.g., glucose) – Structural molecules (e.g., ribose sugar in RNA) Carbohydrates • Monosaccharide: simple sugar, one carbon ring – Examples: glucose, fructose, galactose • Disaccharide: two monosaccharides joined by a covalent bond – Examples: sucrose, maltose, lactose • Polysaccharide: several monosaccharides joined together – Example: starch (composed of thousands of glucose molecules) Carbohydrates • Glycogen: another polysaccharide formed to store sugar in a cell – Glycogen does not pull in water via osmosis as simple sugars do. • Cellulose: a polysaccharide made by plants – Cellulose is not digestible by humans. Lipids • Contain C, H, O (less than in carbohydrates), and sometimes P • Insoluble in water • Main types: – Triglycerides or neutral fats – Phospholipids – Steroids – Eicosanoids Figure 2.16a Lipids. Triglyceride formation Three fatty acid chains are bound to glycerol by dehydration synthesis. + Glycerol © 2013 Pearson Education, Inc. + 3 fatty acid chains Triglyceride, or neutral fat 3 water molecules Saturated and Unsaturated Fats • If every carbon on the fatty acid chain shares a single electron, the fatty acid is saturated. • If there are double bonds between carbons, the fatty acid is unsaturated. • Trans fats – modified oils – unhealthy • Omega-3 fatty acids – “heart healthy” Ketone Bodies • Hydrolysis of triglycerides forms free fatty acids in the blood. These can be used for energy or converted into ketone bodies by the liver. – Strict low-carbohydrate diets and uncontrolled diabetes can result in elevated ketone levels, called ketosis. – Ketone levels lower pH ->can cause ketoacidosis, which can lead to coma and death. Phospholipids • Modified triglycerides: – Glycerol + two fatty acids and a phosphorus (P) containing group • “Head” and “tail” regions have different properties • Important in cell membrane structure “Typical” structure of a phospholipid molecule Two fatty acid chains and a phosphorus-containing group are attached to the glycerol backbone. Example Phosphatidylcholine Polar “head” Nonpolar “tail” (schematic phospholipid) Phosphorus-containing group (polar “head”) © 2013 Pearson Education, Inc. Glycerol backbone 2 fatty acid chains (nonpolar “tail”) Steroids • A steroid is structurally very different from a triglyceride but nonpolar, so considered a lipid. – 3 six-carbon rings + 1 five-carbon ring + functional groups • Cholesterol is a steroid used (1) as a precursor to steroid hormones, such as testosterone, estrogen, and aldosterone, and (2) to make molecules such as vitamin D and bile salts. Eicosanoids • Many different ones • Derived from a fatty acid (arachidonic acid) in cell membranes • Most important eicosanoid – Prostaglandins • Role in blood clotting, control of blood pressure, inflammation, and labor contractions © 2013 Pearson Education, Inc. Proteins • Contain C, H, O, N, and sometimes S and P • Proteins are polymers • Amino acids (20 types) are the monomers in proteins – Joined by covalent bonds called peptide bonds – Contain amine group and acid group – Can act as either acid or base – All identical except for “R group” (in green on figure) © 2013 Pearson Education, Inc. Amino Acids • An amino acid has an amino group, a carboxyl group, and a functional group. – The functional group is what differentiates the 20 amino acids. Amine group Acid group Generalized structure of all amino acids. © 2013 Pearson Education, Inc. Glycine is the simplest amino acid. Aspartic acid (an acidic amino acid) has an acid group (—COOH) in the R group. Lysine (a basic amino acid) has an amine group (—NH2) in the R group. Cysteine (a basic amino acid) has a sulfhydryl (—SH) group in the R group, which suggests that this amino acid is likely to participate in intramolecular bonding. Figure 2.18 Amino acids are linked together by peptide bonds. Dehydration synthesis: The acid group of one amino acid is bonded to the amine group of the next, with loss of a water molecule. Peptide bond + Amino acid Dipeptide Amino acid Hydrolysis: Peptide bonds linking amino acids together are broken when water is added to the bond. © 2013 Pearson Education, Inc. Protein Structure • A chain of amino acids is called a polypeptide chain. – primary structure: The chain varies in length from 3 to 4,500 amino acids. • Weak hydrogen bonds may form between neighboring amino acids. – This may form an alpha helix or a beta fold: secondary structure. Protein Structure • Attraction to amino acids further away produces bends and folds, creating a specific 3D shape: tertiary structure – This structure dictates function. – Since weak bonds hold tertiary structure together, a protein is easily denatured (unfolded) by changes in pH or temperature. Protein Structure • Some functional proteins are composed of multiple polypeptide chains covalently bonded together. – This is called the quaternary structure of the protein. – Examples are the hemoglobin in blood and the hormone insulin. Protein Functions • Structural: collagen fibers in connective tissues; keratin in skin • Enzymes: assist every chemical process in the body • Antibodies: part of the immune system • Receptors: receive communication from other cells for regulation of cell activity • Carriers: across cell membranes or in blood Protein Denaturation • Denaturation – Globular proteins unfold and lose functional, 3-D shape • Active sites destroyed – Can be cause by decreased pH or increased temperature • Usually reversible if normal conditions restored • Irreversible if changes extreme – e.g., cooking an egg © 2013 Pearson Education, Inc. Enzymes – Proteins that act as biological catalysts • Regulate and increase speed of chemical reactions – Lower the activation energy, increase the speed of a reaction WITHOUT ENZYME WITH ENZYME Less activation energy required Energy Energy Activation energy required Reactants Reactants Product Progress of reaction © 2013 Pearson Education, Inc. Product Progress of reaction Characteristics of Enzymes • Enzymes are specific – Act on specific substrate • Usually end in -ase • Often named for the reaction they catalyze – Hydrolases, oxidases Figure 2.21 Mechanism of enzyme action. Substrates (S) e.g., amino acids + Energy is Water is absorbed; released. bond is formed. Product (P) e.g., dipeptide Peptide bond Active site Enzyme (E) Enzyme-substrate complex (E-S) 1 Substrates bind at active 2 The E-S complex site, temporarily forming an undergoes internal enzyme-substrate complex. rearrangements that form the product. Enzyme (E) 3 The enzyme releases the product of the reaction. Nucleic Acids • Building blocks: Nucleotides – Composed of a five-carbon sugar, a phosphate group, and a nitrogenous base – Nitrogenous bases fall into two categories: • Pyrimidine: a single carbon ring + nitrogen • Purine: 2 carbon rings + nitrogen Deoxyribonucleic Acid (DNA) • The sugar in this molecule is called deoxyribose and can bind to one of four nitrogenous bases: – – – – Guanine Thymine Cytosine Adenine DNA Structure • Deoxyribose bonds with a phosphate group to form a long chain, which serves as the backbone of the molecule. • Each nitrogenous base can form a hydrogen bond with another to result in a double-stranded molecule. – Cytosine can only bind with guanine. – Tyrosine can only bind with adenine. • The two chains of DNA are twisted, forming a double helix. Ribonucleic Acid (RNA) • Similar to DNA except: – Has ribose sugar instead of deoxyribose – Is single-stranded instead of doublestranded – Has uracil instead of thymine Adenosine Triphosphate (ATP) • Chemical energy in glucose captured in this important molecule • Directly powers chemical reactions in cells • Energy form immediately useable by all body cells © 2013 Pearson Education, Inc. Figure 2.23 Structure of ATP (adenosine triphosphate). High-energy phosphate bonds can be hydrolyzed to release energy. Adenine Phosphate groups Ribose Adenosine Adenosine monophosphate (AMP) Adenosine diphosphate (ADP) Adenosine triphosphate (ATP) © 2013 Pearson Education, Inc. Figure 2.24 Three examples of cellular work driven by energy from ATP. Function of ATP Solute game • Phosphorylation – Terminal phosphates are enzymatically transferred to and energize other molecules – Such “primed” molecules perform cellular work (life processes) using the phosphate bond energy + Membrane protein Transport work: ATP phosphorylates transport proteins, activating them to transport solutes (ions, for example) across cell membranes. + Relaxed smooth muscle cell Contracted smooth muscle cell Mechanical work: ATP phosphorylates contractile proteins in muscle cells so the cells can shorten. + Chemical work: ATP phosphorylates key reactants, providing energy to drive energy-absorbing chemical reactions. © 2013 Pearson Education, Inc. Case Studies to be discussed: Ch 2: Critical Thinking and Clinical Application Questions #2: Some antibiotics act by binding to certain essential enzymes in the target bacteria. A) How might these antibiotics influence the chemical reactions controlled by the enzymes? B) What is the anticipated effect on the bacteria? On the person taking the antibiotic prescription? #3: Mrs. R., in a diabetic coma, has just been admitted to N hospital. Her blood pH indicates that she is in severe acidosis, and measures are quickly instituted to bring her pH back to normal. A) Define pH and note the normal pH of blood. B) Why is severe acidosis a problem? #4: Jason (12 yo) was awakened by a loud crash. As he sat up, his fright was revealed by his rapid breathing (hyperventilation). A) What happens during hyperventilation?, B) At this point, was his blood pH rising or falling?