Kin 290 chapter 2, week 2
Kin 290 chapter 2, week 2 Kin 290
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This 16 page Class Notes was uploaded by Leonard Carey on Monday April 11, 2016. The Class Notes belongs to Kin 290 at 1 MDSS-SGSLM-Langley AFB Advanced Education in General Dentistry 12 Months taught by Dr. Satern in Spring 2016. Since its upload, it has received 14 views. For similar materials see Anatomy & Physiology in Kinesiology at 1 MDSS-SGSLM-Langley AFB Advanced Education in General Dentistry 12 Months.
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Date Created: 04/11/16
Chapter 2 Chapter 2 – Part A Chemistry Comes Alive Why This Matters • Understanding chemistry and biochemistry helps to determine the most effective solutions to use to treat dehydration and fluid loss Chemistry and Physiological Reactions • Body is made up of many chemicals • Chemistry underlies all physiological reactions: • Movement, digestion, pumping of heart, nervous system • Chemistry can be broken down into: – Basic chemistry – Biochemistry Part 1 – Basic Chemistry 2.1 Matter and Energy Matter • Matter is anything that has mass and occupies (takes up) space • Matter can be seen, smelled, and/or felt • Weight is mass plus the effects of gravity Matter • States of matter • Matter can exist in three possible states: • Solid: definite shape and volume (example: bones, hair, organs) • Liquid: changeable shape; definite volume (example: Blood) • Gas: changeable shape and volume (example: oxygen, carbon dioxide) Energy • Energy is the capacity (ability) to do work or put matter into motion • Energy does not have mass, nor does it take up space • The greater the work done, the more energy it uses up Energy (cont.) • Kinetic versus potential energy – Energy exists in two possible forms: • Kinetic – energy in action (kin is prefix for movement) • Potential – stored (inactive) energy (get the energy from food) – Energy can be transformed from potential to kinetic energy • Stored energy can be released, resulting in action © 2016 Pearson Education, Inc. 1 Chapter 2 Energy (cont.) • Forms of energy • Chemical energy • Stored in bonds of chemical substances (when atoms move together to form a new substance) • Electrical energy • Results from movement of charged particles (ions) • Mechanical energy • Directly involved in moving matter (potential and kinetic energy) • Radiant or electromagnetic energy • Travels in waves (example: heat, visible light, ultraviolet light, and X rays) Energy (cont.) • Energy form conversions • Energy may be converted from one form to another • Example: turning on a lamp converts electrical energy to light energy • Energy conversion is inefficient • Some energy is “lost” as heat, which can be partly unusable energy 2.2 Atoms and Elements • All matter is composed of elements • Elements are substances that cannot be broken down into simpler substances by ordinary chemical methods • Four elements make up 96% of body: • Carbon, oxygen, hydrogen, and nitrogen • 9 elements make up 3.9% of body • 11 elements make up <0.01% • Periodic table lists all known elements 2.2 Atoms and Elements • All elements are made up of atoms, which are: • Unique building blocks for each element • Smallest particles of an element with properties of that element • What give each element its particular physical & chemical properties 2.2 Atoms and Elements (Table 2.1 – p. 26) • Atomic symbol • One or twoletter chemical shorthand for each element • Example: “O” for oxygen, “C” for carbon • Some symbols come from Latin names: “Na” (natrium) is sodium; “K” (kalium) is potassium Structure of Atoms © 2016 Pearson Education, Inc. 2 Chapter 2 • Atoms are composed of three subatomic particles: • Protons • Carry a positive charge (+) • Weigh an arbitrary 1 atomic mass unit (1 amu) • Neutrons • Have no electrical charge (0) • Also weigh 1 amu • Electrons • Carry a negative charge () • Are so tiny they have virtually no weight (0 amu) Structure of Atoms (cont.) • Number of positive protons is balanced by number of negative electrons, so atoms are electrically neutral • Protons and neutrons are found in a centrally located nucleus; electrons orbit around the nucleus • Chemists devise models of how subatomic particles are put together • Planetary model • Orbital model Structure of Atoms (cont.) (Figure 2.1 – p. 25) • Planetary model: simplified and outdated because it incorrectly depicts electrons in orbits, fixed circular paths – Still useful for illustrations • Orbital model: current model used that depicts orbitals, probable regions where an electron is most likely to be located (rather than fixed orbits) • Shading in regions of greatest electron density results in an electron cloud around nucleus • Useful for predicting chemical behavior of atoms Identifying Elements (Figure 2.2 – p. 27) • Different elements contain different numbers of subatomic particles • Hydrogen has 1 proton, 0 neutrons, and 1 electron • Helium has 2 protons, 2 neutrons, and 2 electrons • Lithium has 3 protons, 4 neutrons, and 3 electrons • Identifying facts about an element include its atomic number, mass number, isotopes, and atomic weight Identifying Elements (cont.) • Atomic number • Number of protons in nucleus • Written as subscript to left of atomic symbol © 2016 Pearson Education, Inc. 3 Chapter 2 • Example: Li3 • Mass number • Total number of protons and neutrons in nucleus • Total mass of atom • Written as superscr7pt to left of atomic symbol • Example: Li Identifying Elements (cont.) (Figure 2.3 – p. 27) • Isotopes – Structural variations of same element – Atoms contain same number of protons but differ in the number of neutrons they contain • Atomic numbers are same, but mass numbers different • Atomic weight • Average of mass numbers of all isotope forms of an ato 2.3 Combining Matter Molecules and Compounds • Most atoms chemically combine with other atoms to form molecules and compounds • Molecule: general term for 2 or more atoms bonded together • Compound: specific molecule that has 2 or more different kinds of atoms bonded together • Example: C H 6 12 6 • Molecules with only one type of atom (H or O2) are 2ust called molecules Mixtures (Figure 2.4 – p. 29) • Most matter exists as mixtures: two or more components that are physically intermixed • Three basic types of mixtures • Solutions • Colloids • Suspensions Mixtures (cont.) • Solutions • Are homogeneous mixtures, meaning particles are evenly distributed throughout • Solvent: substance present in greatest amount • Usually a liquid, such as water • Solute(s): substance dissolved in solvent • Present in smaller amounts • Example: blood sugar – glucose is solute, and blood (plasma) is solvent Mixtures (cont.) © 2016 Pearson Education, Inc. 4 Chapter 2 • Solutions (cont.) • True solutions are usually transparent • Example: air (gas solution), salt solution, sugar solution • Most solutions in body are true solutions of gases, liquids, or solids dissolved in water Difference Between Mixtures and Compounds • Three main differences: • Unlike compounds, mixtures do not involve chemical bonding between components • Mixtures can be separated by physical means, such as straining or filtering; compounds can be separated only by breaking their chemical bonds • Mixtures can be heterogeneous or homogeneous; compounds are only homogeneous 2.4 Chemical Bonds • Chemical bonds are “energy relationships” between electrons of reacting atoms • Chemical bonds are not actual physical structures • Electrons are the subatomic particles that are involved in all chemical reactions • They determine whether a chemical reaction will take place and if so, what type of chemical bond is formed Role of Electrons in Chemical Bonding • Electrons can occupy areas around nucleus called electron shells • Each shell contains electrons that have a certain amount of kinetic and potential energy, so shells are also referred to as energy levels • Depending on its size, an atom can have up to 7 electron shells • Shells can hold only a specific number of electrons; the shell closest to nucleus is filled first • Shell 1 can hold only 2 electrons • Shell 2 holds a maximum of 8 electrons • Shell 3 holds a maximum of 18 electrons Role of Electrons in Chemical Bonding (cont.) • Outermost electron shell is called valence shell • Electrons in valence shell have the most potential energy because they are farthest from nucleus • These are electrons that are involved in chemical reactions Role of Electrons in Chemical Bonding (cont.) (Figure 2.5 – p. 31) • Octet rule (rule of eights) • Atoms desire 8 electrons in their valence shell • Exceptions: smaller atoms (examples: H and He) want only 2 electrons in shell 1 • Desire to have 8 electrons is driving force behind chemical reactions • Noble gases already have full 8 valence electrons (or 2 for He) so are not © 2016 Pearson Education, Inc. 5 Chapter 2 chemically reactive • Most atoms do not have full valence shells • Atoms will gain, lose, or share electrons (form bonds) with other atoms to achieve stability of 8 electrons in valence shell Types of Chemical Bonds • Three major types of chemical bonds • Ionic bonds • Covalent bonds • Hydrogen bonds Types of Chemical Bonds (cont.) • Ionic bonds • Ions are atoms that have gained or lost electrons and become charged • Number of protons does not equal number of electrons Types of Chemical Bonds (cont.) (Figure 2.6 – p. 32) • Ionic bonds involve the transfer of valence shell electrons from one atom to another, resulting in 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 Types of Chemical Bonds (cont.) (Figure 2.6 – p. 32) • Most ionic compounds are salts • When dry, salts form crystals instead of individual molecules • Example is NaCl (sodium chloride) Types of Chemical Bonds (cont.) (Figure 2.7 – p. 33) • Covalent bonds • Covalent bonds are formed by sharing of two or more valence shell electrons between two atoms • Sharing of 2 electrons results in a single bond • Sharing of 4 electrons is a double bond • Sharing of 6 electrons is a triple bond • Allows each atom to fill its valence shell at least part of the time • Two types of covalent bonds: • Polar and nonpolar covalent bonds Types of Chemical Bonds (cont.) (Figure 2.8 – p. 34) © 2016 Pearson Education, Inc. 6 Chapter 2 • Covalent bonds (cont.) • Nonpolar covalent bonds • Equal sharing of electrons between atoms • Results in electrically balanced, nonpolar molecules such as CO 2 Types of Chemical Bonds (cont.) • Polar covalent bonds • Unequal sharing of electrons between 2 atoms • Results in electrically polar molecules • Atoms have different electronattracting abilities, leading to unequal sharing • Atoms with greater electronattracting ability are electronegative, and those with less are electropositive Types of Chemical Bonds (cont.) (Figure 2.9 – p. 34) • Polar covalent bonds (cont.) • H O is a polar molecule 2 • Oxygen is more electronegative, so it exerts a greater pull on shared electrons, giving it a partial negative charge and giving H a partial positive charge • Having two different charges is referred to as dipole 2.5 Chemical Reactions Chemical Equations • Chemical reactions occur when chemical bonds are formed, rearranged, or broken • These reactions can be written in symbolic forms called chemical equations • Chemical equations contain: • Reactants: substances entering into reaction together • Product(s): resulting chemical end products • Amounts of reactants and products are shown in balanced equations Chemical Equations • Compounds are represented as molecular formulas • Example: H O o2 C H O 6 12 6 • Subscript indicates atoms joined by bonds • Prefix denotes number of unjoined atoms or molecules Chemical Equations (cont.) • Compounds are represented as molecular formulas • Example: H O o2 C H O 6or12 6r CH 2 4 • In chemical equations, subscripts indicate how many atoms are joined by bonds, whereas prefix means number of unjoined atoms (example: 4H) Reactants Product H + H H2 (Hydrogen gas) 4H + 1C CH4 (Methane) © 2016 Pearson Education, Inc. 7 Chapter 2 Types of Chemical Reactions (Figure 2.11 – p. 36) • Three main types of chemical reactions: 1. Synthesis (combination) reactions involve atoms or molecules combining to form larger, more complex molecule • Used in anabolic (building) processes A + B AB Types of Chemical Reactions (cont.) 2. Decomposition reactions involve breakdown of a molecule into smaller molecules or its constituent atoms (reverse of synthesis reactions) • Involve catabolic (bondbreaking) reactions AB A + B Types of Chemical Reactions (cont.) 3. Exchange reactions, also called displacement reactions, involve both synthesis and decomposition • Bonds are both made and broken AB + C AC + B and AB + CD AD + CB Types of Chemical Reactions (cont.) • In living systems, these reactions are also referred to as reductionoxidation or redox reactions • Atoms are reduced when they gain electrons and oxidized when they lose electrons • Example: C H O + 6O → 6CO + 6H O + ATP 6 12 6 2 2 2 • In this example, glucose is oxidized, and oxygen molecule is reduced Energy Flow in Chemical Reactions • All chemical reactions are either exergonic or endergonic • Exergonic reactions result in a net release of energy (give off energy) • Products have less potential energy than reactants • Catabolic and oxidative reactions • Endergonic reactions result in a net absorption of energy (use up energy) • Products have more potential energy than reactants • Anabolic reactions Reversibility of Chemical Reactions • All chemical reactions are theoretically reversible © 2016 Pearson Education, Inc. 8 Chapter 2 A + B ←→ AB • Chemical equilibrium occurs if neither a forward nor a reverse reaction is dominant • Many biological reactions are not very reversible • Energy requirements to go backward are too high, or products have been removed Rate of Chemical Reactions • The speed of chemical reactions can be affected by: • Temperature: increased temperatures usually increase rate of reaction • Concentration of reactants: increased concentrations usually increase rate • Particle size: smaller particles usually increase rate Rate of Chemical Reactions • Catalysts • Catalysts increase the rate of reaction without being chemically changed or becoming part of the product • Enzymes are biological catalysts Chapter 2 – Part B Chemistry Comes Alive Part 2 – Biochemistry Biochemistry is the study of chemical composition and reactions of living matter All chemicals either organic or inorganic – Inorganic compounds Water, salts, and many acids and bases Do not contain carbon – Organic compounds Carbohydrates, fats, proteins, and nucleic acids Contain carbon, are usually large, and are covalently bonded Both equally essential for life 2.6 Inorganic Compounds Water Most abundant inorganic compound – Accounts for 60%–80% of the volume of living cells Most important inorganic compound because of its properties – High heat capacity – High heat of vaporization – Polar solvent properties – Reactivity – Cushioning © 2016 Pearson Education, Inc. 9 Chapter 2 Water High heat capacity – Ability to absorb and release heat with little temperature change – Prevents sudden changes in temperature High heat of vaporization – Evaporation requires large amounts of heat – Useful cooling mechanism Water (cont.) (Figure 2.12 – p. 39) Polar solvent properties – Dissolves and dissociates ionic substances – Forms hydration (water) layers around large charged molecules Example: proteins – Body’s major transport medium Water (cont.) Reactivity – Necessary part of hydrolysis and dehydration synthesis reactions Cushioning – Protects certain organs from physical trauma Example: cerebrospinal fluid cushions nervous system organs Salts Salts are ionic compounds that dissociate into separate ions in water – Separate into cations (positively charged molecules) and anions (negatively charged) Not including H and OH ions Salts (cont.) Salts (cont.) – All ions are called electrolytes because they can conduct electrical currents in solution – Ions play specialized roles in body functions Example: sodium, potassium, calcium, and iron – Ionic balance is vital for homeostasis – Common salts in body NaCl, CaCO , 3Cl, calcium phosphates Acids and Bases Acids and bases are both electrolytes – Ionize and dissociate in water © 2016 Pearson Education, Inc. 10 Chapter 2 Acids + – Are proton donors: they release hydrogen ions (H ), bare protons (have no electrons) in solution + – Example: HCl → H + Cl – Important acids HCl (hydrochloric acid), HC H O (acetic acid, abbreviated HAc), and H CO 2 3 2 2 3 (carbonic acid) Acids and Bases (cont.) Bases + – Are proton acceptors: they pick up H ions in solution Example: NaOH → Na + OH + – – – When a base dissolves in solution, it releases a hydroxyl ion (OH ) – Important bases – Bicarbonate ion (HCO ) a3d ammonia (NH ) 3 Acids and Bases (cont.) pH: Acidbase concentration + – pH scale is measurement of concentration of hydrogen ions [H ] in a solution – The more hydrogen ions in a solution, the more acidic that solution is + – pH is negative logarithm of [H ] in moles per liter that ranges from 0–14 – pH scale is logarithmic, so each pH unit represents a 10fold difference Example: a pH 5 solution is 10 times more acidic than a pH 6 solution Acids and Bases (cont.) (Figure 2.13 – p. 40) pH: Acidbase concentration (cont.) – Acidic solutions have high [H ] but low pH Acidic pH range is 0–6.99 – Neutral solutions have equal numbers of H and OH ions – All neutral solutions are pH 7 Pure water is pH neutral + –7 – pH of pure water pH 7: [H ] 10 m – Alkaline (basic) solutions have low [H ] but high pH Alkaline pH range is 7.01–14 Acids and Bases (cont.) Neutralization – Neutralization reaction: acids and bases are mixed together Displacement reactions occur, forming water and a salt NaOH + HCl → NaCl + H O 2 © 2016 Pearson Education, Inc. 11 Chapter 2 Acids and Bases (cont.) Buffers + + – Acidity involves only free H in solution, not H bound to anions – Buffers resist abrupt and large swings in pH Can release hydrogen ions if pH rises Can bind hydrogen ions if pH falls – Convert strong acids or bases (completely dissociated) into weak ones (slightly dissociated) Carbonic acid–bicarbonate system (important buffer system of blood): 2.7 Organic Compounds: Synthesis and Hydrolysis Organic molecules contain carbon – Exceptions: CO an2 CO, which are inorganic Carbon is electroneutral – Shares electrons; never gains or loses them – Forms four covalent bonds with other elements – Carbon is unique to living systems Major organic compounds: carbohydrates, lipids, proteins, and nucleic acids 2.7 Organic Compounds: Synthesis and Hydrolysis (Figure 2.14 – p. 42) Many are polymers – Chains of similar units called monomers (building blocks) Synthesized by dehydration synthesis Broken down by hydrolysis reactions 2.8 Carbohydrates Carbohydrates include sugars and starches Contain C, H, and O – Hydrogen and oxygen are in 2:1 ratio Three classes – Monosaccharides: one single sugar Monomers: smallest unit of carbohydrate – Disaccharides: two sugars – Polysaccharides: many sugars Polymers are made up of monomers of monosaccharides 2.8 Carbohydrates (Figure 2.15 – p. 43) Monosaccharides – Simple sugars containing three to seven carbon atoms – (CH O)2 —ngeneral formula n number of carbon atoms – Monomers of carbohydrates © 2016 Pearson Education, Inc. 12 Chapter 2 – Important monosaccharides Pentose sugars – Ribose and deoxyribose Hexose sugars – Glucose (blood sugar) Carbohydrates (cont.) Disaccharides – Double sugars – Too large to pass through cell membranes – Important disaccharides Sucrose, maltose, lactose – Formed by dehydration synthesis of two monosaccharides glucose + fructose → sucrose + water Carbohydrates (cont.) (Figure 2.15 – p. 44) Polysaccharides – Polymers of monosaccharides Formed by dehydration synthesis of many monomers – Important polysaccharides Starch: carbohydrate storage form used by plants Glycogen: carbohydrate storage form used by animals – Not very soluble 2.9 Lipids Contain C, H, O, but less than in carbohydrates, and sometimes contain P Insoluble in water Main types: – Triglycerides or neutral fats – Phospholipids – Steroids – Eicosanoids Lipids (cont.) Triglycerides or neutral fats – Called fats when solid and oils when liquid – Composed of three fatty acids bonded to a glycerol molecule – Main functions Energy storage Insulation Protection © 2016 Pearson Education, Inc. 13 Chapter 2 Lipids (cont.) Triglycerides can be constructed of: – Saturated fatty acids All carbons are linked via single covalent bonds, resulting in a molecule with the maximum number of H atoms (saturated with H) Solid at room temperature (Example: animal fats, butter) Lipids (cont.) (Figure 2.16 – p. 45) – Unsaturated fatty acids One or more carbons are linked via double bonds, resulting in reduced H atoms (unsaturated) Liquid at room temperature (Example: plant oils, such as olive oil) Trans fats – modified oils; unhealthy Omega3 fatty acids – “heart healthy” Lipids (cont.) Phospholipids – Modified triglycerides Glycerol and two fatty acids plus a phosphoruscontaining group – “Head” and “tail” regions have different properties Head is a polar region and is attracted to water Tails are nonpolar and are repelled by water – Important in cell membrane structure Lipids (cont.) Steroids – Consist of four interlocking ring structures – Common steroids: cholesterol, vitamin D, steroid hormones, and bile salts – Most important steroid is cholesterol Is building block for vitamin D, steroid synthesis, and bile salt synthesis Important in cell plasma membrane structure 2.10 Proteins Comprise 20–30% of cell mass Have most varied functions of any molecules – Structural, chemical (enzymes), contraction (muscles) Contain C, H, O, N, and sometimes S and P Polymers of amino acid monomers held together by peptide bonds Shape and function due to four structural levels Amino Acids and Peptide Bonds (Figure 2.17 – p. 48) © 2016 Pearson Education, Inc. 14 Chapter 2 All proteins are made from 20 types of amino acids – Joined by covalent bonds called peptide bonds – Contain both an amine group and acid group – Can act as either acid or base – Differ by which of 20 different “R groups” is present Enzymes and Enzyme Activity (Figure 2.19 – p. 51) Enzymes: globular proteins that act as biological catalysts – Catalysts regulate and increase speed of chemical reactions without getting used up in the process – Lower the energy needed to initiate a chemical reaction Leads to an increase in the speed of a reaction Allows for millions of reactions per minute! 2.11 Nucleic Acids Nucleic acids, composed of C, H, O, N, and P, are the largest molecules in the body Nucleic acid polymers are made up of monomers called nucleotides – Composed of nitrogen base, a pentose sugar, and a phosphate group Two major classes: – Deoxyribonucleic acid (DNA) – Ribonucleic acid (RNA) 2.11 Nucleic Acids DNA holds the genetic blueprint for the synthesis of all proteins – Doublestranded helical molecule (double helix) located in cell nucleus – Nucleotides contain a deoxyribose sugar, phosphate group, and one of four nitrogen bases: Purines: adenine (A), guanine (G) Pyrimidines: cytosine (C) and thymine (T) 2.11 Nucleic Acids (Figure 2.21 – p. 53) DNA holds the genetic blueprint for the synthesis of all proteins (cont.) – Bonding of nitrogen base from strand to opposite strand is very specific Follows complementary basepairing rules: – A always pairs with T – G always pairs with C 2.11 Nucleic Acids RNA links DNA to protein synthesis and is slightly different from DNA – Singlestranded linear molecule is active mostly outside nucleus – Contains a ribose sugar (not deoxyribose) – Thymine is replaced with uracil © 2016 Pearson Education, Inc. 15 Chapter 2 – Three varieties of RNA carry out the DNA orders for protein synthesis Messenger RNA (mRNA), transfer RNA (tRNA), and ribosomal RNA (rRNA) 2.12 ATP (Figure 2.22 – p. 55) Chemical energy released when glucose is broken down is captured in ATP (adenosine triphosphate) ATP directly powers chemical reactions in cells – Offers immediate, usable energy needed by body cells Structure of ATP – Adeninecontaining RNA nucleotide with two additional phosphate groups 2.12 ATP (Figure 2.23 – p. 55) Terminal phosphate group of ATP can be transferred to other compounds that can use energy stored in phosphate bond to do work – Loss of phosphate group converts ATP to ADP – Loss of second phosphate group converts ADP to AMP © 2016 Pearson Education, Inc. 16
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