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Chapter 2

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Chapter 2

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  1. Chapter 2 Biology 25: Human Biology Prof. Gonsalves Los Angeles City College Loosely Based on Mader’s Human Biology,7th edition

  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. 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.

  8. 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

  9. 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.

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

  11. 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

  12. 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

  13. 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

  14. 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

  15. 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+)

  16. 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

  17. 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 --

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

  19. 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

  20. 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

  21. Nonpolar Covalent Bonds: Electrons are Shared Equally

  22. 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.

  23. 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).

  24. 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

  25. 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.

  26. 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.

  27. 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

  28. 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

  29. 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.

  30. 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”

  31. 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.

  32. 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-

  33. 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 )