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Chapter 2 Chemistry Matter Occupies space and has mass Mass does not depend on gravity, weight does Three states of matter – Solid – definite shape and volume – Liquid – indefinite shape, definite volume – Gas – indefinite shape and volume Energy Capacity to do work (put matter into motion) Energy cannot be created or destroyed Does not take up space, has no mass Measured by its effects on matter Two types of energy – Kinetic energy – the energy of motion – Potential energy – stored energy Forms of Energy Chemical energy – stored in chemical bonds, especially adenosine triphosphate (ATP) Electrical energy – results from movement of charged particles Mechanical energy – directly moves matter Radiant energy – waveform energy, includes light waves, radio waves, UV waves, and X rays Energy Form Conversions Energy is constantly converted to different forms During conversion, some energy is “lost” as heat Atoms and Elements Elements CANNOT be broken down 112 different elements, 92 are natural Main elements in the body = carbon, oxygen, hydrogen, nitrogen Atoms are building blocks of elements Atoms and Elements (cont.) Atoms give each element its special properties – Physical properties – detectable by our senses (color, texture, boiling point) – Chemical properties – interaction of atoms with other atoms (some elements create rust, some can be digested, etc.) Atomic symbol = chemical abbreviation of element (C = carbon, Ca = calcium, Na = sodium) Atomic Structure The atom has a central nucleus, containing protons and neutrons – Protons (p+) – positively charged, have a mass of 1 atomic mass unit (1 amu) – Neutrons (n0) – uncharged, have a mass of 1 amu – Most of the atom’s mass is in the nucleus Atomic Structure (cont.) Electrons orbit the nucleus – Electrons (e-) – negatively charged, only 1/2000 the mass of a proton (0 amu) Atoms are electrically neutral, because number of protons = number of electrons Identification of Elements: Atomic Structure Figure 2.2 Identifying Elements Atomic Number = number of protons Mass Number = sum of masses of protons and neutrons in an element Isotopes = structural variation of an element, contains different amount of neutrons Atomic Weight = average of mass numbers of all isotopes of an element; dependent on abundance of the isotopes Identification of Elements: Isotopes of Hydrogen Figure 2.3 Molecules Molecule = combination of two or more atoms joined by chemical bonds – Compounds = combination of two or more different atoms joined by chemical bonds Molecules are the smallest parts of compounds that display characteristics of that compound Mixtures Mixture = made of two or more mixed components which are not chemically bonded Three types of mixtures are solutions, colloids, and suspensions Solutions Homogenous mixture; may be gas, liquid, or solid Contain solvent (greater amount) and solute (lesser amount) Transparent and do not scatter light Components do not settle out Examples include salt water and sugar water Concentration of Solutions Can be indicated by the percent of solute in total solution Can be indicated by molarity (moles/L) Mole = molecular weight of a compound weighed out in grams Avogadro’s number – one mole of any substance always contains 6.02 X 1023 molecules Concentration of Solutions (cont.) Glucose = C6H12O6 C = 12.011, H = 1.008, O = 15.999 Total atomic weight = 180.156 To make a 1 molar solution of glucose, weight out 180.156 grams of glucose, and add water up to 1 L Colloids (Emulsions) – Heterogeneous mixture – Do not settle out – Milky in appearance – Will scatter light – Can convert from a fluid to solid state – Examples include gelatin and cytosol Suspensions – Heterogeneous mixture – Visible solutes that settle out – Will scatter light – Examples include sand mixed with water and blood Comparison of Mixtures and Compounds No chemical bonding in mixtures Mixtures can be separated by physical means Compounds require chemical means to be separated All compounds are homogenous, mixtures can be homogenous or heterogenous Chemical Bonds Created by interaction of electrons Electrons travel around nucleus of atom in electron shells Atoms can have 7 shells Electron shells are also called energy levels Shell 1 = 2 e-, shell 2 = 8 e-, shell 3 = 18 e- Chemical Bonds (cont.) Shells fill up in order Only the outermost energy level is reactive!! Valence shell = reactive outer energy level of an atom If outer shell is full, that atom is not reactive (inert) If outer shell is not full, the atom will react to become stable Chemical Bonds (cont.) Only 8 electrons bond, no matter how many are in a shell Octet rule = except in shell 1, atoms interact with each other to have eight electrons in their valence shell Chemically Inert and Reactive Elements Figure 2.4a Figure 2.4b Types of Chemical Bonds Ionic – transfer of electrons Covalent – sharing of electrons Hydrogen – weak attractions Ionic Bonds Ion = charged particle formed from gain or loss of electrons Ionic bond = electrons are transferred from one atom to another If an atom accepts electrons, it has a negative charge = anion If an atom loses electrons, it has a positive charge = cation Formation of an Ionic Bond Figure 2.5a Ionic Bonds (cont.) Ionic compounds are usually salts Salts form crystals Figure 2.5b Covalent Bonds Covalent bond = sharing of pairs of electrons between two or more atoms Shared electrons will orbit the whole molecule Single Covalent Bonds Figure 2.7a Double Covalent Bonds Figure 2.7b Triple Covalent Bonds Figure 2.7c Polar and Nonpolar Covalent Bonds Nonpolar molecules are formed from equal electron sharing (symmetrical) Polar molecules are formed from unequal electron sharing (nonsymmetrical) – Electronegative atoms attract electrons – Electropositive atoms donate electrons Comparison of Ionic, Polar Covalent, and Nonpolar Covalent Bonds Figure 2.9 Hydrogen Bonds Formed when hydrogen is covalently linked to one electronegative atom and is attracted to another electronegative atom Hydrogen bonding is common in water, causing surface tension Hydrogen bonding is also found in proteins and DNA Hydrogen Bonds Figure 2.10a Chemical Reactions Occur when chemical bonds are formed, rearranged, or broken Reactions are written as chemical equations – Equations show reactants and products Types of Chemical Reactions Synthesis, or combination – atoms or molecules join to form a more complex molecule Also called anabolic A + B AB Types of Chemical Reactions (cont.) Decomposition reaction = breakdown of a molecule into smaller particles Also called catabolic AB A + B Types of Chemical Reactions (cont.) Exchange reactions = both synthesis and breakdown occur AB + C AC + B AB + CD AD + CB Oxidation-Reduction Reactions Redox reactions are how the body breaks down food for energy Electrons are transferred between reactants Reactant that loses electrons (electron donor) is oxidized Reactant that gains electrons (electron acceptor) is reduced Energy of Reactions Exergonic reactions = reactions that release energy – Products have less energy than reactants – Energy produced can be used Endergonic reactions = reactions that absorb energy – Products have more energy than reactants – Energy is put in to the reaction Reversibility of Reactions All reactions are reversible A + B AB AB A + B Chemical equilibrium = when a reaction proceeds equally in both directions Factors Influencing Rate of Reactions Atoms have to collide to react – Temperature – increasing temperature speeds up reactions – Concentration – increasing reactant amounts will speed up reactions – Particle size – smaller particles move faster, speeds up reactions – Catalysts – increase rate of reactions without being changed themselves Biochemistry Study of chemistry as related to living matter Organic compounds – Contain carbon – Have covalent bonds Inorganic compounds – Waters, salts, acids, and bases – Do not contain carbon Water 1. 2. 60-80% of all cells High heat capacity – can absorb and release large amounts of heat without large temperature change High heat of vaporization – when water evaporates, a lot of heat is released (perspiration) Water (cont.) 3. Polar solvent properties – Universal solvent – Ions will dissociate in water – Water surrounds larger molecules, keeps them from settling out of solution – Major transport medium of the body Water (cont.) 4. Reactivity – Water plays a part in many chemical reactions – Decomposition using water = hydrolysis – Synthesis using water = dehydration synthesis 5. Cushioning – protects internal organs Salts Inorganic compounds Contains cations other than H+ and anions other than OHSalts dissolve in water to create electrolytes Common salts in body = NaCl, calcium phosphates Acids and Bases Dissociate into ions in water Acid – releases hydrogen ions in solution, proton donor (HCl, H2CO3) HCl H+ + Cl – Base – releases hydroxide ions in solution, proton acceptor (NaOH, HCO3-, NH3) NaOH Na+ + OH– pH More H+ free in solution, more acidic More OH- free in solution, more basic (alkaline) Neutral solutions, H+ = OHConcentration of H+ in solution = pH Acid-Base Concentration (pH) Acidic: pH 0–6.99 Basic: pH 7.01–14 Neutral: pH 7.00 pH change of 1 represents 10-fold change in [H+] Figure 2.13 Neutralization Mixture of acids and bases in solution causes formation of water and salt HCl + NaOH NaCl + H2O Buffers Prevent large changes in pH of body fluids Release H+ when pH rises Bind H+ when pH falls pH of a solution due only to free H+ Strong acids and bases = completely dissociate in water Buffers (cont.) Weak acids and bases = do not dissociate completely in water, so less effect on pH of solution Weak acids will bind H+ if acid is added, and will release H+ if base is added Weak bases will bind OH- if base is added, and will bind more H+ if acid is added Organic Compounds Contain carbon Carbon is electroneutral (always shares electrons), and can form 4 bonds Four types of organic compounds – Carbohydrates – Lipids – Proteins – Nucleic acids Carbohydrates Sugars and starches Contain carbon, hydrogen, and oxygen 2:1 ratio of hydrogen and oxygen Larger carbohydrates are less soluble in water Carbohydrates (cont.) Easily available source of fuel Classes of carbohydrates – Monosaccharides – Disaccharides – Polysaccharides Monosaccharides Building blocks of carbohydrates Single rings or chains with 3-7 carbons General formula = (CH2O)n Most important in body are pentoses and hexoses Isomers = same chemical formula, different arrangement of atoms Monosaccharides Figure 2.14a Disaccharides Formed from joining of two monosaccharides by dehydration synthesis Water is released in this reaction Must be broken down by hydrolysis to be absorbed from intestines Disaccharides Figure 2.14b Polysaccharides Long chains of monosaccharides Important storage molecules Two main ones important for humans = starch and glycogen Figure 2.14c Lipids Do not dissolve in water Contain C, H, and O, but the proportion of oxygen in lipids is less than in carbohydrates Also can contain phosphorus Examples: – – – – Neutral fats or triglycerides Phospholipids Steroids Eicosanoids Triglycerides Composed of three fatty acids bonded to a glycerol molecule Fats when solid, oils when liquid Nonpolar molecules Figure 2.15a Triglycerides (cont.) Saturated = fatty acid with single bonds only – Straight hydrocarbon chains – Solid form Unsaturated = fatty acid with double bonds between carbons – Kinks in hydrocarbon chains – Oil form at room temperature Phospholipids Phospholipids – modified triglycerides with two fatty acid groups and a phosphorus group Phosphate group is polar, rest is nonpolar = amphipathic or amphiphilic Figure 2.15b Other Lipids Steroids – flat molecules with four interlocking hydrocarbon rings Eicosanoids – 20-carbon fatty acids found in cell membranes Figure 2.15c Functions of Lipids Triglycerides – insulation Phospholipids – form cell membranes Steroids – part of cell membranes, also used in hormone production Eicosanoids – prostaglandins most important; used for blood clotting and inflammation Proteins Amino acids are building blocks of protein Contain C, H, O, and N Amino group NH2 Carboxyl group COOH Amino acids can be acidic or basic Amino Acids Figure 2.16a–c Protein Macromolecules composed of combinations of 20 types of amino acids bound together with peptide bonds Peptide bond H H R O N C C OH H Amino acid + H H R O N C C OH H Amino acid Dehydration H O 2 synthesis Hydrolysis H H2O H R O H R O N C C N C C H H OH Dipeptide Figure 2.17 Structural Levels of Proteins Primary = amino acid sequence Secondary = twisting and bending creates alpha helices and beta chains Tertiary – folding of secondary structures Quaternary – separate polypeptide chains linked together Structural Levels of Proteins Figure 2.18a–c Structural Levels of Proteins Figure 2.18b,d,e Fibrous and Globular Proteins Fibrous proteins (structural) – Strandlike, some have only secondary structure – Insoluble in water, very stable – Examples: keratin, elastin, and collagen Globular proteins (functional) – Spherical proteins with tertiary and quaternary structures – Water soluble, chemically active – Examples: antibodies, hormones, and enzymes Protein Denaturation Proteins must fold into a specific pattern to function Destruction of hydrogen bonds by temperature or pH changes leads to denaturation Denaturation = unfolding of proteins which causes loss of function Protein Denaturation (cont.) Active sites must be intact for protein to function Figure 2.19a Protein Denaturation Irreversibly denatured proteins cannot refold, and cannot function Figure 2.19b Enzymes Globular proteins; act as catalysts Chemically specific Speed up reactions Usually end in –ase Often have to be activated Lower activation energy of reactions Enzymes Decrease Activation Energy Figure 2.20 Enzyme Action 1. 2. 3. Enzyme binds with substrates Enzyme-substrate complex rearranges to form products Enzyme releases the products; enzyme is unchanged Active site Amino acids + Enzyme (E) Substrates (S) Enzyme-substrate complex (E-S) H2O Free enzyme (E) Peptide bond Internal rearrangements leading to catalysis Dipeptide product (P) Figure 2.21 Nucleic Acids Composed of C, N, H, O, and P Built from nucleotides – Nucleotides have N-containing base, pentose sugar, and a phosphate group 5 bases = adenine (A), guanine (G), cytosine (C), thymine (T), and uracil (U) Nucleic Acids (cont.) Adenine and guanine are purines – Large, have 2 rings Cytosine, thymidine, and uracil are pyrimidines – Smaller, have 1 ring Two major classes of nucleic acids = DNA and RNA Deoxyribonucleic Acid (DNA) Found in nucleus Contains genetic material which is the blueprint for protein production Sugar is deoxyribose Replicates before cell division Forms a double helix Structure of DNA Figure 2.22a Structure of DNA Figure 2.22b Ribonucleic Acid (RNA) Single-stranded Usually located outside of nucleus Uses uracil instead of thymine Important for creation of proteins Sugar is ribose instead of deoxyribose 3 kinds: mRNA, tRNA, and rRNA Adenosine Triphosphate (ATP) Stores energy in phosphate bonds Readily usable by body cells Adenine-containing nucleotide with three phosphate groups Unstable Terminal phosphate group is transferred to other molecules – called phosphorylation Structure of ATP Figure 2.23 Membrane protein Pi P Solute Solute transported (a) Transport work ADP + Pi ATP Relaxed smooth muscle cell Contracted smooth muscle cell (b) Mechanical work Pi X P X Y + Y Reactants Product made (c) Chemical work Figure 2.24