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Chapter 2 The Chemical Basis of Life

Chapter 2 The Chemical Basis of Life. I. Elements: Substances that can not be broken down into simpler substances by chemical reactions. There are 92 naturally occurring elements: Oxygen, carbon, nitrogen, calcium, sodium, etc. Life requires about 25 of the 92 elements Chemical Symbols:

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Chapter 2 The Chemical Basis of Life

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  1. Chapter 2 The Chemical Basis of Life

  2. I. Elements: • Substances that can not be broken down into simpler substances by chemical reactions. • There are 92 naturally occurring elements: Oxygen, carbon, nitrogen, calcium, sodium, etc. • Life requires about 25 of the 92 elements • Chemical Symbols: • Abbreviations for the name of each element. • Usually one or two letters of the English or Latin name of the element • First letter upper case, second letter lower case. Example: Helium (He), sodium (Na), potassium (K), gold (Au).

  3. Main Elements: Over 98% of an organism’s mass is made up of six elements. • Oxygen (O): 65% body mass • Cellular respiration, component of water, and most organic compounds. • Carbon (C): 18% of body mass. • Backbone of all organic compounds. • Hydrogen (H): 10% of body mass. • Component of water and most organic compounds. • Nitrogen (N): 3% of body mass. • Component of proteins and nucleic acids (DNA/RNA) • Calcium (Ca): 1.5% of body mass. • Bones, teeth, clotting, muscle and nerve function. • Phosphorus (P): 1% of body mass • Bones, nucleic acids, energy transfer (ATP).

  4. Minor Elements: Found in low amounts. Between 1% and 0.01%. • Potassium (K): Main positive ion inside cells. • Nerve and muscle function. • Sulfur (S): Component of most proteins. • Sodium (Na): Main positive ion outside cells. • Fluid balance, nerve function. • Chlorine (Cl): Main negative ion outside cells. • Fluid balance. • Magnesium (Mg): Component of many enzymes and chlorophyll.

  5. Trace elements: Less than 0.01% of mass: • Boron(B) • Chromium(Cr) • Cobalt(Co) • Copper(Cu) • Iron(Fe) • Fluorine (F) • Iodine (I) • Manganese (Mn) • Molybdenum (Mo) • Selenium (Se) • Silicon (Si) • Tin (Sn) • Vanadium (V) • Zinc (Zn)

  6. II. Structure & Properties of Atoms Atoms: Smallest particle of an element that retains its chemical properties. Made up of three main subatomic particles. ParticleLocationMassCharge Proton (p+) In nucleus 1 +1 Neutron (no) In nucleus 1 0 Electron (e-) Outside nucleus 0* -1 * Mass is negligible for our purposes.

  7. Atomic Particles: Protons, Neutrons, and Electrons Helium Atom Carbon Atom

  8. Structure and Properties of Atoms 1. Atomic number = # protons • The number of protons is unique for each element • Each element has a fixed number of protons in its nucleus. This number will never change for a given element. • Written as a subscript to left of element symbol. Examples: 6C, 8O, 16S, 20Ca • Because atoms are electrically neutral (no charge), the number of electronsandprotons are always the same. • In the periodic table elements are organized by increasing atomic number.

  9. Structure and Properties of Atoms: 2. Mass number = # protons + # neutrons • Gives the mass of a specific atom. • Written as a superscript to the left of the element symbol. Examples: 12C, 16O, 32S, 40Ca. • The number of protons for an element is always the same, but the number of neutrons may vary. • The number of neutrons can be determined by: # neutrons = Mass number - Atomic number

  10. Structure and Properties of Atoms: 3. Isotopes: Variant forms of the same element. • Isotopes have different numbers of neutrons and therefore different masses. • Isotopes have the same numbers of protons and electrons. • Example: In nature there are three forms or isotopes of carbon (6C): • 12C: About 99% of atoms. Have 6 p+, 6 no, and 6 e-. • 13C: About 1% of atoms. Have 6 p+, 7 no, and 6 e-. • 14C: Found in tiny quantities. Have 6 p+, 8 no, and 6 e-. Radioactive form (unstable). Used for dating fossils.

  11. Electrons determine how an atom can bond with other atoms A. Energy levels: Electrons occupy different energy levels around the nucleus. Level (Shell) Electron Capacity 12 (Closest to nucleus, lowest energy) 28 38 (If valence shell, 18 otherwise) 4, 5, & 6 18 B. Electron configuration:Arrangement of electrons in orbitals around nucleus of atom. C. Valence Electrons: Number of electrons in outer energy shell of an atom.

  12. Electron Arrangements of Important Elements of Life 1 Valence electron 4 Valence electrons 5 Valence electrons 6 Valence electrons

  13. III. How Atoms Form Molecules: Chemical Bonds Molecule: Two or more atoms combined chemically. Compound: A substance with two or more elements combined in a fixed ratio. • Water (H2O) • Hydrogen peroxide (H2O2) • Carbon dioxide (CO2) • Carbon monoxide (CO) • Table salt (NaCl) • Atoms are linked by chemical bonds. Chemical Formula: Describes the chemical composition of a molecule of a compound. • Symbols indicate the type of atoms • Subscripts indicate the number of atoms

  14. How Atoms Form Molecules: Chemical Bonds “Octet Rule”: When the outer shell of an atom is not full, i.e.: contains fewer than 8 (or 2) electrons (valence e-), the atom tends to gain, lose, or share electrons to achieve a complete outer shell (8, 2, or 0) electrons. Example: Sodium has 11 electrons, 1 valence electron. Sodium loses its electron, becoming an ion: Na -------> Na+ + 1 e- 1(2), 2(8), 3(1) 1(2), 2(8) Outer shell has 1 e- Outer shell is full Sodium atom Sodium ion

  15. Number of valence electrons determine the chemical behavior of atoms. Element Valence Combining Tendency Electrons Capacity Sodium 1 1 Lose 1 Calcium 2 2 Lose 2 Aluminum 3 3 Lose 3 Carbon 4 4 Share 4 Nitrogen 5 3 Gain 3 Oxygen 6 2 Gain 2 Chlorine 7 1 Gain 1 Neon* 8 0 Stable * Noble gas

  16. How Atoms Form Molecules: Chemical Bonds Atoms can lose, gain, or share electrons to satisfy octet rule (fill outermost shell). Two main types of Chemical Bonds A. Ionic bond: Atoms gain or lose electrons B. Covalent bond: Atoms share electrons

  17. A. Ionic Bond: Atoms gain or lose electrons. Bonds are attractions between ions of opposite charge. Ionic compound: One consisting of ionic bonds. Na + Cl ----------> Na+ Cl- sodium chlorine Table salt (Sodium chloride) Two Types of Ions: Anions: Negatively charged particle (Cl-) Cations: Positively charged particle (Na+)

  18. Ionic Bonding: Sodium Chloride

  19. A Crystal of Sodium Chloride:Ions are Held Together by Ionic Bonds

  20. B. Covalent Bond: Involves the “sharing” of one or more pairs of electrons between atoms. Covalent compound: One consisting of covalent bonds. Example: Methane (CH4): Main component of natural gas. H | H---C---H | H Each line represents on shared pair of electrons. Octet rule is satisfied: Carbon has 8 electrons, Hydrogen has 2 electrons

  21. There may be more than one covalent bond between atoms: 1. Single bond: One electron pair is shared between two atoms. Example: Chlorine (Cl2), water (H2O); methane (CH4) Cl --- Cl 2. Double bond: Two electron pairs share between atoms. Example: Oxygen gas (O2); carbon dioxide (CO2) O=O 3. Triple bond: Three electron pairs shared between two atoms. Example: Nitrogen gas (N2) N = N --

  22. Single and Double Covalent Bonds

  23. Number of covalent bonds formed by important elements: Carbon (4) Nitrogen (3) Oxygen (2) Sulfur (2) Hydrogen (1)

  24. Two Types of Covalent Bonds: Polar and Nonpolar Electronegativity: A measure of an atom’s ability to attract and hold onto a shared pair of electrons. Some atoms such as oxygen or nitrogen have a much higher electronegativity than others, such as carbon and hydrogen. ElementElectronegativity O 3.5 N 3.0 S & C 2.5 P & H 2.1

  25. Polar and Nonpolar Covalent Bonds A.Nonpolar Covalent Bond: When the atoms in a bond have equal or similar attraction for the electrons (electronegativity), they are shared equally. Example: O2, H2, Cl2

  26. Nonpolar Covalent Bonds: Electrons are Shared Equally

  27. Polar and Nonpolar Covalent Bonds B. Polar Covalent Bond: When the atoms in a bond have different electronegativities, the electrons are shared unequally. Electrons are closer to the more electronegative atom creating a polarity or partial charge. Example: H2O Oxygen has a partial negative charge. Hydrogens have partial positive charges.

  28. Polar Covalent Bonds: Electrons are Shared Unequally Creating Partial Charges Water Molecule

  29. Other Bonds: Weak chemical bonds are important in the chemistry of living things. • Hydrogen bonds: Attraction between the partially positive H of one molecule and a partially negative atom of another • Hydrogen bonds are about 20 X easier to break than a normal covalent bond. • Responsible for many properties of water. • Determine 3 dimensional shape of DNA and proteins. • Chemical signaling (molecule to receptor).

  30. Hydrogen Bonds: Weak Attractions between Hydrogen and Partially Negative Atoms Water Molecules

  31. Water: A Unique Compound for Life

  32. Water: The Ideal Compound for Life • Living cells are 70-90% water • Water covers 3/4 of earth’s surface • Water is the ideal solvent for chemical reactions • On earth, water exists as gas, liquid, and solid

  33. I. Polarity of water causes hydrogen bonding • Water molecules are held together by H-bonding • Partially positive H attracted to partially negative O atom. • Individual H bond are weak, but the cumulative effect of many H bonds is very strong. • H bonds only last a fraction of a second, but at any moment most molecules are hydrogen bonded to others.

  34. Hydrogen Bonds in Water are Responsible for Many of its Properties

  35. Unique properties of water caused by H-bonds • Cohesion:Water molecules stick to each other. This causes surface tension. • Film-like surface of water is difficult to break. • Used by some insects that live on water surface. • Water forms beads. • Adhesion:Water sticks to many surfaces. Capillary Action:Water tends to rise in narrow tubes. This is caused bycohesion and adhesion (water molecules stick to walls of tubes). Examples: Upward movement of water through plant vessels and fluid in blood vessels.

  36. Unique properties of water caused by H-bonds • Expands when it freezes. • Ice forms stable H bonds, each molecule is bonded to four neighbors (crystalline lattice). Water does not form stable H bonds. • Ice is less dense than water. • Ice floats on water. • Life can survive in bodies of water, even though the earth has gone through many winters and ice ages

  37. Ice Forms Stable Hydrogen Bonds and is Less Dense than Water

  38. Unique properties of water caused by H-bonds • StableTemperature:Water resists changes in temperature because it has a high specific heat. • Specific Heat: Amount of heat energy needed to raise 1 g of substance 1 degree Celsius • Specific Heat of Water: 1 calorie/gram/oC • High heat of vaporization: Water must absorb large amounts of energy (heat) to evaporate. • Heat of Vaporization of Water: 540 calorie/gram. • Evaporative cooling is used by many organisms to regulate body temperature. • Sweating • Panting

  39. Unique properties of water caused by H-bonds • Universal Solvent: Dissolves many (but not all) substances to form solutions. Solutions are homogeneous mixtures of two or more substances (salt water, air, tap water). All solutions have at least two components: • Solvent: Dissolving substance (water, alcohol, oil). • Aqueous solution: If solvent is water. • Solute: Substance that is dissolved (salt, sugar, CO2). • Water dissolves polar and ionic solutes well. • Water does not dissolve nonpolar solvents well.

  40. A Salt Crystal Dissolving in Water

  41. Solubility of a Solute Depends on its Chemical Nature Solubility: Ability of substance to dissolve in a given solvent. Two Types of Solutes: A. Hydrophilic:“Water loving” dissolve easily in water. • Ionic compounds (e.g. salts) • Polar compounds (molecules with polar regions) • Examples: Compounds with -OH groups (alcohols). • “Like dissolves in like”

  42. Solubility of a Solute Depends on its Chemical Nature Two Types of Solutes: B. Hydrophobic: “Water fearing” do not dissolve in water • Non-polar compounds (lack polar regions) • Examples: Hydrocarbons with only C-H non-polar bonds, oils, gasoline, waxes, fats, etc.

  43. ACIDS, BASES, pH AND BUFFERS A. Acid: A substance that donates protons (H+). • Separate into one or more protons and an anion: HCl (into H2O ) -------> H+ + Cl- H2SO4 (into H2O ) --------> H+ + HSO4- • Acids INCREASE the relative [H+] of a solution. • Water can also dissociate into ions, at low levels: H2O <======> H+ + OH-

  44. B. Base: A substance that accepts protons (H+). • Many bases separate into one or more positive ions (cations) and a hydroxyl group (OH-). • Bases DECREASE the relative [H+] of a solution ( and increases the relative [OH-] ). H2O <======> H+ + OH- DirectlyNH3 + H+ <=------> NH4+ Indirectly NaOH ---------> Na+ + OH- ( H+ + OH- <=====> H2O )

  45. Strong acids and bases: Dissociation is almost complete (99% or more of molecules). HCl (aq) -------------> H+ + Cl- NaOH (aq) -----------> Na+ + OH- (L.T. 1% in this form) (G.T. 99% in dissociated form) • A relatively small amount of a strong acid or base will drastically affect the pH of solution. Weak acids and bases: A small percentage of molecules dissociate at a give time (1% or less) H2CO3 <=====> H+ + HCO3- carbonic acid Bicarbonate ion (G.T. 99% in this form) (L.T. 1% in dissociated form)

  46. C. pH scale: [H+] and [OH-] • pH scale is used to measure how basic or acidic a solution is. • Range of pH scale: 0 through 14. • Neutral solution: pH is 7.[H+ ] = [OH-] • Acidic solution: pH is less than 7. [H+ ] > [OH-] • Basic solution: pH is greater than 7. [H+ ] < [OH-] • As [H+] increases pH decreases (inversely proportional). • Logarithmic scale: Each unit on the pH scale represents a ten-fold change in [H+].

  47. pH of Common Solutions

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