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Chapter 2 Chemical Principles

Chapter 2 Chemical Principles. 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 Chemical Principles

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  1. Chapter 2 Chemical Principles

  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 copounds. • 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. 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

  13. 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

  14. 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

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

  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. B. Covalent Bond - Involve 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

  19. 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

  20. Number of covalent bonds: Carbon (4) Nitrogen (3) Oxygen (2) Sulfur (2) Hydrogen (1)

  21. Two Types of Covalent Bonds: Polar and Nonpolar A. 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

  22. Polar and Nonpolar Covalent Bonds B.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, N2, Cl2 C. 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 - A Unique Compound for Life

  25. 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

  26. 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.

  27. Unique properties of water caused by H-bonds • Cohesion:Water molecules stick to each other. • Adhesion:Water molecules stick to many surfaces. • StableTemperature:Water resists changes in temperature. • High heat of vaporization: Water must absorb large amounts of energy (heat) to evaporate. • Expands when it freezes (water denser than ice) • Solvent: Dissolves many substances.

  28. II. Biological Consequences of Water’s Polarity A. Capillary Action:Water tends to rise in narrow tubes. This is caused by two factors: • Cohesion: Molecules of water “stick together” • Adhesion: Water molecules stick to walls of tubes. Examples: Upward movement of water through plant vessels and fluid in blood vessels. B. Surface tension: Difficulty in “stretching or breaking” • At water/air interface, difficult to pull water apart • Causes water to “bead” into tiny balls • Used by some insects who live on the surface of water

  29. C. Temperature Regulation Water has a very 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/degree C • Organisms can absorb a lot of heat without drastic changes in temperature. D. Evaporative Cooling • Vaporization: Transformation from liquid to gas. • Heat of Vaporization: Energy required to convert 1 gram of a liquid -> gas is high (540 calories/gram) • Sweating is a form of evaporative cooling. • Can regulate temperature w/o great water loss.

  30. E. Ice floats on Water: Life can exist in bodies of water Ice floats because liquid water is more dense than ice (solid water). • Water gets more dense as it cools to 4oC. • Water gets less dense (expands) as it cools further to form ice. • Crystalline lattice forms, molecules farther apart Because ice floats, life can survive and thrive in bodies of water, even though the earth has gone through many winters and ice ages

  31. III. Water is the ideal solvent for chemical reactions • Solution: Homogeneous mixture of 2 or more substances. • Examples: Salt water, air, tap water. • Solvent: Dissolving substance of a solution. • Example: Water, alcohol, oil. • Solute: Substance dissolved in the solvent. • Example: NaCl, sugar, carbon dioxide. • Aqueous solution: Water is the solvent. • Solubility: Ability of substance to dissolve in a given solvent.

  32. Solubility of a Solute Depends on its Chemical Nature 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” 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.

  33. 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-

  34. 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 )

  35. 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)

  36. 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+].

  37. pH of Common Solutions

  38. D. Buffers keep pH of solutions relatively constant • Buffer:Substance which prevents sudden large changes in pH when acids or bases are added. • Buffers are biologically important because most of the chemical reactions required for life can only take place within narrow pH ranges. • Example: • Normal blood pH 7.35-7.45. Serious health problems will arise if blood pH is not stable.

  39. CHEMICAL REACTIONS • A chemical change in which substances (reactants) are joined, broken down, or rearranged to form new substances (products). • Involve the making and/or breaking of chemical bonds. • Chemical equations are used to represent chemical reactions. Example: 2H2 + O2 -----------> 2H2O 2Hydrogen Oxygen 2 Water Molecules Molecule Molecules

  40. Organic Compounds

  41. I. Organic Chemistry: Carbon Based Compounds • Organic Compounds: Compounds that contain carbon and are synthesized by cells (except CO and CO2). • Diverse group:Several million organic compounds are known. More are identified daily. • Common:After water,organiccompounds are the most common substances in cells. • Over 98% of thedry weightof living cells is made up oforganiccompounds. • Less than 2% of thedry weightof living cells is made up of inorganic compounds. • Inorganic Compounds: Compounds without carbon.

  42. Organic Compounds are Carbon Based Carbon Has 4 Valence Electrons and Can Form 4 Covalent Bonds

  43. Organic compounds are incredibly diverse Organic molecules can vary dramatically in: • Length (1-100s of C atoms) • Shape (Linear chain, branched, ring) • Type of bonds: • Single • Double • Triple bonds • Other elements that bond to C: • Nitrogen (N) • Oxygen (O) • Hydrogen (H) • Sulfur (S) • Phosphorus (P)

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