Chemistry 130 Test 4 Study Guide

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Chemistry 130 Test 4 Study Guide
2014-04-27 19:55:43
Chemistry test study guide
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  1. Define viscosity, viscous.
    A fluid's resistance to flow.

    Something that is very viscous will resist flowing.

    "Pancake syrup is more viscous than water."
  2. Define surface tension.
    A tendency for a liquid to form itself into a ball with nothing else around it.

    Raindrops tend to be spherical.
  3. Define evaporation.
    A physical change of a liquid becoming a vapor or gas; opposite of condensation.
  4. Define condensation.
    The opposite of evaporation; a physical change of a vapor forming back into a liquid.
  5. Define vapor pressure.
    The pressure that a vapor of a liquid exerts when it fills the maximum amount available in a container.
  6. Define dynamic equilibrium.
    Like a two-way street.  The reactants form the products, but then as soon as the products form the reactants are formed again.  This goes back and forth.
  7. Define melting point.
    The temperature at which a solid becomes a liquid or melts.
  8. Define boiling point.
    The temperature at which a liquid begins to vaporize and become a gas.
  9. Define normal boiling point.
    The boiling point at 760 tor (1 atmosphere). This is the average boiling point at sea level.
  10. Define freezing.
    The temperature at which a substance becomes a solid.
  11. Define melting.
    The temperature at which a solid becomes a liquid.
  12. Define sublimation.
    Changing directly from a solid to a vapor, bypassing the liquid completely.

    Frost can do this, without melting it gradually disappears.

    The opposite of deposition.
  13. Define deposition.
    Changing directly from a vapor to a solid, bypassing the liquid completely.  

    The opposite of sublimation.
  14. Define exothermic reactions.
    The reaction gives off heat to the surroundings.

    Think of exo as "exit".  Heat is exiting the reaction.

    Striking a match is an exothermic reaction.

    The opposite of endothermic reaction.
  15. Define endothermic reaction.
    The reaction absorbs heat from its surrounding.

    The reaction mixture tends to cool off while the reaction is in progress.

    An example of this is a cold pack.  Break the seal between the two compartments, water mixes with ammonium nitrate, drawing heat from the outside.  The pack gets cold.
  16. Describe the heating curve for water.
  17. Describe London forces/dispersion forces/dipole-induced dipole interactions.
    Two molecules passing near each other, they are both non-polar.  

    However, they make each other polar because the electron clouds tend to repel each other. 

    Other molecules nearby those that are passing close, see what's happening, and they develop a dipole moment where they become polar temporarily.
  18. List the different kinds of intermolecular forces and their strengths releative to each other, starting with weakest first.
    • 1) London/dispersion forces
    • 2) Dipole-induced dipole
    • 3) Dipole-dipole interactions (occur only between polar molecules)
    • 4) Hydrogen bonding (only three elements capable of this: N2, O2, F2)

    Hydrogen bonding is generally the strongest out of these.
  19. With intermolecular forces, all else being equal, the stronger the intermolecular forces the _______ the melting point and the _______ the boiling point.
    higher; higher
  20. All intermolecular forces are much weaker than _______ and _______.
    covalent bonds; ionic bonds
  21. How do you identify a solid as an ionic solid?
    You can tell this from the formula given.

    Opposite sides of the periodic table, Potassium Chloride for example, Potassium is an alkaline metal on the extreme left, and Chlorine is on the far right, so this will be ionic all the time.

    As a solid, it does not conduct electricity.

    Ionic solids often dissolve in water.

    When one is dissolved in water, the solution does conduct electricity.
  22. Ionic solids generally have _______ _______ melting points.  What is an example?
    very high


    NaCl -- 801oC melting point
  23. How do you identify a solid as a molecular solid?
    These are generally gases, liquids, or low-melting solids.

    Sometimes dissolves in water, but the solution does not conduct electricity.


    • * Dry ice
    • * Water
  24. Neither _______ solids nor _______ solids conduct electricity.
    ionic; molecular
  25. How do you identify a solid as a metallic solid?
    Regular metals: copper, iron

    Usually they have melting points.

    They will conduct electricity, unlike ionic and molecular solids.
  26. How do you identify a solid as a network covalent solid (covalent atomic solids)?
    • * Things like Diamond
    • * Very high melting point
    • * Does not conduct electricity
    • * Does not dissolve in anything
  27. In a dry ionic solid, the ions are _______ in place, they _______ move.  Free movement is _______ to conduct electricity.
    locked; cannot; required
  28. List the states of matter and their strength of intermolecular forces relative to energy.
    • Gas (weak)
    • Liquid (moderate)
    • Solid (strong)
  29. In general, the rate of vaporization increases with...
    • Increasing surface area
    • Increasing temperature
    • Decreasing strength of intermolecular forces
  30. Define volatile and nonvolatile.
    Liquids that evaporate easily are termed volatile, while those that do not vaporize easily are termed nonvolatile.

    Rubbing alcohol, for example, is more volatile than water.

    Motor oil at room temperature is virtually nonvolatile.
  31. Vapor pressure increases with...
    • Increasing temperature
    • Decreasing strength of intermolecular forces.
  32. Give an example of dynamic equilibrium.
    In a closed, partially empty water bottle, more evaporation happens than condensation because there are so few gaseous water molecule sin the space above the water.

    However, as the number of gaseous water molecules increases, the rate of condensation also increases.  

    At the point where the rates of condensation and evaporation become equal, dynamic equilibrium is reached and the number of gaseous water molecules above the liquid remains constant.
  33. Define solvent.
    Present in the largest amount (majority component) of a solution.
  34. Define solute.
    Present in the smallest amount (minority component) of a solution.
  35. Define true solution.
    Anything dissolved in a liquid or a gas, either way, it will be transparent.  The contents dissolve 100% in the base solvent.  The solution may be colored, but it is transparent nonetheless.
  36. Define suspension.
    Components that are not transparent, such as muddy water.  Contents are cloudy.  The opposite of true solution.
  37. The solubility of gases tend to _______ with increasing temperature.
  38. The solubility of solids tend to _______ with increasing temperature.
  39. Define unsaturated.
    More solute can be dissolved in the solution than what is there already.
  40. Define saturated.
    No more solute can be dissolved in the solution than what is there already.
  41. Define supersaturated.
    A solution that is more concentrated than a saturated solution.  It is above the normal limit of saturation.

    Supersaturated solutions are inherently unstable.
  42. Define miscible.
    Two liquids that form a homogeneous mixture when added together.

    If they dissolve in each other at all proportions, they are said to be miscible.
  43. Define immiscible.
    Two liquids that do not dissolve in each other are immiscible (not miscible, the opposite of miscible).

    "Vegetable oil is immiscible with water."
  44. Define strong electrolyte.
    Compounds that exist in water solutions only as ions.

    Ordinary ionic compounds that dissolve in water are always strong electrolytes.

    Strong electrolytes are good conductors of electricity.
  45. Define nonelectrolytes.
    Compounds that dissolve as molecules, not as ions.

    Alcohol, sugar, yeast would be nonelectrolytes.  They do not form any ions in a solution.
  46. If you _______ the temperature and sufrace area, the rate at which liquid evaporates will increase.
  47. If you _______ the temperature and surface area, the rate at which liquid evaporates will decrease.
  48. Describe the effect of pressure on the solubility of a gas.
    If you increase the pressure of a gas sitting above a liquid, more of the gas will dissolve.
  49. Describe the effect of pressure on the boiling point of a liquid.
    If you increase the pressure, the boiling point will also increase.

    This is the basis of a pressure cooker.

    If you decrease the pressure, the boiling point will decrease.

    This is why water boils at less than 100oC in Show Low (higher altitude, less pressure).
  50. A mixture of 25g of NaCl (molar mass = 58.44 g/mol) in 90.0mL of water (90.0mL weighs 90.0 grams).  What is the % weight/volume?
    • .              mass
    • % w/v = ---------- x 100
    •                  mL

    •                  25g
    • % w/v = ---------- x 100 = 27.8 %
    •               90.0mL
  51. A mixture of 25g of NaCl (molar mass = 58.44 g/mol) in 90.0mL of water (90.0mL weighs 90.0 grams).  What is the % weight/weight?
    • .                   mass
    • % w/w = ----------------------- x 100
    •              mass + liquid mass

    •                   25g
    • % w/w = ------------- x 100 = 21.7%
    •               25g + 90g
  52. A mixture of 25g of NaCl (molar mass = 58.44 g/mol) in 90.0mL of water (90.0mL weighs 90.0 grams).  What is the molarity?
    • .
    •                (    1 mole    )
    • (25g NaCl)----------------- = .428 moles NaCl
    •                ( 58.44 g/mol)

    90.0 mL = 0.090 L

    • 0.428 moles
    • --------------- = 4.76 M
    •    0.090 L
  53. Need # of grams of HCl to be dissolved in 200mL of water to make a 0.50 M solution. HCl molar mass = 36.5 g/mol.
    • .            ( 0.50 moles) ( 36.5 g  )
    • (0.200 L) --------------- ------------ = 3.65g
    • .            (      1 L      ) ( 1 mole )
  54. If 50 mL of a 2.0 M solution of NaCl is diluted to 175 mL with water, what is the concentration of the diluted solution?
    • M1V1 = M2V2
    • (2.0 M)(50 mL) = M2(175 mL)

    M2 = 0.571 M

    The answer checks out, because since it's being diluted, it must equal LESS than 2.0 M, which is what it was to start out with.

    Units do not matter with dilution equation, as long as they are the same on both sides.
  55. Define qualitative aspects of colligative properties.
    Colligative properties are those properties of a solution that do not depend on what is dissolved in what; it only depends on how many particles are present.  Not what kind they are.

    Particles being molecules, ions, atoms, whatever they might be.

    It only concerns the number.
  56. Describe qualitative aspects of colligative properties concerning vapor pressure.
    If you were to measure the vapor pressure of water sitting above a salt water solution, it would be lower than over a fresh water solution at the same temperature.
  57. Describe qualitative aspects of colligative properties concerning boiling point elevation, and freezing points depression.
    If you dissolve something in water, the boiling point will be higher than that of pure water, and the freezing point will be lower than that of pure water.
  58. Describe qualitative aspects of colligative properties concerning osmosis.
    Osmosis is the passage of solvent molecules through a semi-permeable membrane.
  59. Describe the general properties of acids and bases.
    Acids taste sour, bases taste bitter.

    Arrhenius definition of acids and bases:

    "acid is a substance that when dissolved in water increases the concentration of hydrogen ion, H+(aq)."

    "A base is a substance that when added to water increases the concentration of hydroxide ion, OH-(aq)."

    Bronsted-Lowry definitions of acids/bases:

    "An acid is defined as being able to donate or lose an electron; a proton donor."

    "A base is defined as having the ability to gain or accept an electron, a Hydrogen ion H+ acceptor."
  60. What is the definition of pH?
    pH is the negative logarithm of a hydrogen ion concentration.
  61. What is the conjugate base of a given acid and the conjugate acid of a given base?  What is an example of each?
    The removal of a Hydrogen ion H+ to create the base.

    •                  (remove H+)
    •      HNO3 ------------------->   NO3-
    • (nitric acid)                (conjugate base)

    •              (add H+)
    •   NH3 ------------------->   NH4+
    • (base)                   (conjugate acid)

    Both of these examples can go the other way (be reversed).
  62. Name the 6 strong acids in the book on page 500, and list 3 of them that are the most common.
    • hydrochloric acid (HCl)
    • hydrobromic acid (HBr)
    • hydroiodic acid (HI)
    • nitric acid (HNO3)
    • perchloric acid (HClO4)
    • sulfuric acid (H2SO4)

    This color represents the most common.
  63. Name the 6 strong bases in the book on page 503, and list 2 of them that are the most common.
    • lithium hydroxide (LiOH)
    • sodium hdroxide (NaOH)
    • potassium hydroxide (KOH)
    • strontium hydroxide (Sr(OH)2)
    • calcium hydroxide (Ca(OH)2)
    • barium hydroxide (Ba(OH)2)
  64. What does strong acid and strong base mean?
    When dissolved in water, it ionizes completely.

    In the case of nitric acid HNO3, you get hydrogen ions and nitrate ions only, there are no HNO3 molecules anymore.

    Same thing with sodium hydroxide NaOH, when dissolved in water you get sodium ions and Hydroxide ions.  There are no NaOH molecules anymore.
  65. A strong acid or a strong base is a _______ electrolyte, a weak acid or a weak base is a _______ electrolyte.
    strong; weak
  66. What does weak acid and weak base mean?
    When dissolved in water, you get some ionization, but for the most part they exist as molecules.
  67. What is the neutral pH?
  68. If the pH of a solution is less than 7, what is it?  If it is greater than 7, what is it?
    acidic; basic
  69. In terms of pH, a Hydrogen ion concentration of 1 x 10-7 is what?
  70. In terms of pH, a Hydrogen ion concentration of 1 x 10-4 is what?
  71. In terms of pH, a Hydrogen ion concentration of 1 x 10-10 is what?
  72. What is the pH of:   1 x 10-9 = [ H+ ]
    1 x 10-9 = [ H+ ]

    pH = 9
  73. Define buffer solution.
    A solution that resists changes in pH.

    If you have a buffer solution to a particular pH and add a strong acid to it, the pH will decrease but only very slightly.  If you add a strong base to it, the pH will increase, but only very slightly.

    If you do the same in a solution with no buffering capacity at all, add a small amount of a strong base to it and the pH goes through the roof.  If you add a small amount of a strong acid the pH will drop like a rock.
  74. What causes acid rain?  What are the effects of it?
    Most common causes is burning high Sulfur coal.  Sulfur Dioxide SO2 forms that reacts with water and Oxygen to form Sulfuric acid H2SO4.  

    It ends up in lakes and streams when it rains so the pH is lowered.  

    Living things cannot live in really acidic water.
  75. Describe some factors that affect the rate of a reaction.
    Temperature.  Raising the temperature raises the rate.

    If you raise the temperature by 10oC, you double the reaction rate.
  76. Define activation energy.
    The energy required to start a reaction.

    A match will not spontaneously burst into flame, you need friction to start it.
  77. Define collision theory.
    Two molecules that are going to react together, have to hit each other hard enough to break bonds.
  78. Define catalyst.
    Increase the rate of a reaction without being used up.
  79. Define equilibrium.
    Reactants form the products, but then the products form the reactants again, goes back and forth.
  80. How do you write an equilibrium constant expression?
    K is a constant expression that would have to be given.

    • A + B = 2C
    •             |
    •   K = [C]2
  81. Use LeChatelier's Principle to determine how the concentrations of the reactants or products change when one of the reactants or products is added or removed, when the temperature changes, or when a catalyst is added.
    If you add more A, then the amount of C will increase.  If you take away A or B, the amount of C will decrease.  C will decompose to replace what is lost.

    Exothermic reaction, heat is a product.  If you heat it up, you'll make more A and B and less C.  If you cool it (remove heat) you'll make more C.

    • A + B = 2C + heat
    •             |
    •   K = [C]2
    •       -------
    •       [A][B]
  82. Define alpha particle.
    Helium nucleus, emitted by atoms that are heavier than they'd like to be, so it's a way of shedding mass.
  83. Define beta particle.
    Fast moving electrons that originate when a neutron decomposes to form a proton which stays in the nucleus.  The electron is the beta particle that comes flying out at high speeds.
  84. Define positron.
    Positively charged electrons.
  85. Define gamma rays.
    Form of light we cannot see.
  86. Define ionizing power.
    Alpha particles have the highest ionizing power.  Gamma radiation is relatively low by comparison.
  87. Define penetrating power.
    Alpha particles cannot penetrate.  They have no penetrating power.  Gamma rays have high penetrating power.

    An alpha emitter in your lungs is the most dangerous situation of all.  You have living tissue right there on the surface.
  88. How do you protect yourself from the harmful effects of nuclear radiation?
    • Get away from the source
    • Use proper shielding
  89. Define half-life.
    The time that it takes for half the sample to decompose.

    If a sample has 1,000,000 atoms and the half life is one day, at this time tomorrow you'll have 500,000.  The next day at the same time there will be 250,000.  The day after that, 125,000, and so on.
  90. Define fission.
    Take a large atom and break it into smaller ones.

    Opposite of fusion.

    Take a Uranium atom and it breaks up to form maybe a Barium atom and a Krypton atom, something like that, two lighter atoms.
  91. Define fusion.
    Take two or more small atoms and combining them into a larger atom.

    Opposite of fission.

    Hydrogen atoms in the Sun are fusing them together to form Helium nuclei.  This is the power source for the Sun.
  92. Define acute radiation damage.
    A large dose of radiation all at once over a short period of time.  It yields increased cancer risk.

    Destroys lining of the stomach, greater problems digesting and absorbing nutrients from food, hair falls out.

    We are normally exposed to about 1/3 of a rem per year.
  93. Describe the different levels of rem and what they do.
    500 rems all at once, the person dies, at least 50/50 chance of dying, gets extremely sick if nothing else.

    The human body is exposed normally to 1/3 of a rem per year.

    100 rems is increased cancer risk.

    Can be used to diagnose disease, because radioactive iodine can be adsorbed by the thyroid.  Rapidly dividing cells take up lots of radioactive phosphorous.
  94. How is radioactivity detected?
    • A film-badge dosimeter contains photographic film in a case that is pinned to clothing.  The more exposed the film, the greater the exposure to radioactivity.
    • In a Geiger-Muller counter, radiation produces ionized argon gas, which gives off an electrical signal.
    • In a scintillation counter, the radioactivity produces visible or ultraviolet light when it passes through compounds such as NaI or CsI.
  95. Define organic chemistry.
    The study of carbon-containing compounds and their reactions.

    The basis for living organisms.

    Organic compounds are prevalent in odors and fragrances.  They are common in foods, drugs, petroleum products and pesticides.
  96. Define hydrocarbon.
    Compounds containing only Carbon and Hydrogen.
  97. Hydrocarbons can be classified into four different types.  What are they?
    • Alkanes CnH2n+2
    • Alkenes CnH2n
    • Alkynes CnH2n-2
    • Aromatic hydrocarbons
  98. Classify each type of hydrocarbon into saturated or unsaturated.  Why are they classified this way?
    • Alkanes are saturated
    • Alkenes are unsaturated
    • Alkynes are unsaturated
    • Aromatic hydrocarbons are unsaturated

    Alkenes have at least one double bond between carbon atoms, and Alkynes have one triple bond.  As a result they have fewer Hydrogen atoms.  This means they are not loaded to full capacity, unlike Alkanes.
  99. Define saturated hydrocarbon.
    A hydrocarbon that is loaded to capacity, such as Alkanes.


    Methane CH4

    •        H
    •        |
    • H -- C -- H
    •        |
    •        H
  100. Define structural formula.
    A formula that shows not only the number and type of each atom in a molecule, but the structure as well.


    Methane CH4

    •        H
    •        |
    • H -- C -- H
    •        |
    •        H
  101. Define condensed structural formula.
    A shorthand way to write a structural formula in which you eliminate many or all of the bonds, and group like-atoms together.

    Structural formula for propane:

    •        H    H    H
    •        |     |     |
    • H -- C -- C -- C -- H
    •        |     |     |
    •        H    H    H

    condensed structural formula for propane:

  102. Define normal alkanes, n-alkanes.
    Alkanes composed of carbon atoms bonded in a straight chain without any branching.

    The n-alkanes with three or more carbon atoms have the general structure:


  103. How do you name an alkane?
    If it's normal, put an n- on the front.

    • 1) methane
    • 2) ethane
    • 3) propane
    • 4) n-butane
    • 5) n-pentane
    • 6) n-hexane
    • 7) n-heptane
    • 8) n-octane
    • 9) n-nonane
    • 10) n-decane
  104. Define isomer.
    Often involving branching, molecules with the same molecular formula but different structures.  Because of their different structures, they have different properties, so they are different compounds.


    butane has 2 isomers

    n-butane: C4H10

    •       H     H    H    H
    •        |     |     |     |
    • H -- C -- C -- C -- C -- H
    •        |     |     |     |
    •       H     H    H    H

    isobutane: C4H10

    •        H    H    H
    •        |     |     |
    • H -- C -- C -- C -- H
    •        |     |     |
    •       H     |    H
    •              |
    •        H-- C -- H
    •              |
    •             H
  105. What are the prefixes for base names of alkane chains?
    • Carbon atoms/Prefix
    • 1/meth-
    • 2/eth-
    • 3/prop-
    • 4/but-
    • 5/pent-
    • 6/hex-
    • 7/hept-
    • 8/oct-
    • 9/non-
    • 10/dec-
  106. What are the common alkyl groups?
  107. Define alkenes.
    Hydrocarbons containing at least one double bond between carbon atoms.
  108. Define alkynes.
    Hydrocarbons containing at least one triple bond between carbon atoms.
  109. Describe how to name alkenes and alkynes.
    They are named the same way as alkanes with the following exceptions:

    1) The base chain is the longest continuous carbon chain that contains the double or triple bond

    2) The base name has the ending -ene for alkenes and -yne for alkynes.

    3) The base chain is numbered to give the double or triple bond the lowest possible number.

    4) A number indicating the position of the double or triple bond (lowest possible number) is inserted just before the base name.
  110. Name the different types of hydrocarbon reactions.
    Combustion reactions: the burning of hydrocarbons in the presence of oxygen.  Alkanes, Alkenes, Alkynes all undergo combustion.  The hydrocarbon reacts with oxygen to form carbon dioxide and water.


    Substitution reactions: Only alkanes do this.  One or more hydrogen atoms on an alkane are replaced by one or more other types of atoms, the most common being a halogen substitution.

    CH4(g) + Cl2(g) ----> CH3Cl(g) + HCl(g)

    The general form for halogen substitution reactions is:

    • R--H  +  X2   -->   R--X      +     HX
    • Alkane    Halogen      Alkil halide      Hydrogen halide


    Addition reactions: Alkenes and alkynes do these, in which atoms add across the multiple bond.

    For example, ethene reacts with chlorine gas to form dichloroethane.

    They can also react with Hydrogen in hydrogenation reactions.  In the presence of the appropriate catalyst, propene reacts with hydrogen gas to form propane:

    Hydrogen reactions convert unsaturated hydrocarbons into saturated hydrocarbons, such as partially hydrogenated vegetable oil.
  111. All hydrocarbons undergo _______ reactions.
  112. Only _______ undergo substitution reactions.
  113. Only alkenes and alkynes undergo _______ reactions.
  114. What are the naming conventions for aromatic hydrocarbons?
    Monosubstituted benzenes: benzenes in which only one of the hydrogen atoms has been substituted.  They are often named as derivatives of benzene.

    They have the general form:

    (name of substituent)benzene

    Many monosubstituted benzenes also have common names that can be learned only through familiarity:

    Benzenes with large substituents are named by treating the benzene ring as the substituent.  In these cases, the benzene substituent is called a phenyl group:


    Disubstituted benzenes: benzenes in which two hydrogen atoms have been substituted.  They are numbered, and the substituents are listed alphabetically.  The order of numbering within the ring is then determined by the alphabetical order of the substituents:

    When the two substituents are identical, we use the prefix di- :

    In common use--in place of numbering-- are the prefixes ortho- (1,2 disubstituted), meta- (1,3 disubstituted), and para- (1,4 disubstituted):

  115. What is an aromatic compound?
    Compounds containing benzene rings that have pleasant aromas (as a result benzene rings are also referred to as aromatic rings).

    Examples of these are the pleasant smells of cinnamon, vanilla, and jasmine.  They are all caused by aromatic compounds.
  116. Define functional group.
    A characteristic atom or group of atoms that has been inserted into a hydrocarbon.

    The letter R is often used to represent a hydrocarbon group.  If the letter G represents a functional group, then a generic formula for families of organic compounds is:

    A group of organic compounds with the same functional group forms a family.  The members of the family of alcohols have an ---OH functional group and the general formula R---OH.  Some specific examples of alcohols are methanol and isopropyl alcohol.

  117. A group of organic compounds with the same _______ forms a family.
    functional group
  118. List the common functional groups, their general formulas, examples, names, etc.
  119. Define polymer.
    Long, chainlike molecules composed of repeating units.  

    The individual repeating units are called monomers.
  120. Define addition polymer.
    A polymer in which the monomers link together without eliminating any atoms.

    Made from molecules containing carbon-carbon double bonds.
  121. Define condensation polymer.
    A polymer that eliminate an atom or a small group of atoms during polymerization.
  122. What are some properties of Aromatic Hydrocarbons?
    Physical properties:

    Nonpolar, thus insoluble in water.

    Many have an odor, "aromatic".

    Chemical properties:

    Substitution reactions are most common, ring remains in tact.

    Polycyclic aromatic compounds: benzo[a]pyrene, naphthalene.

    The human body cannot synthesize phenyl rings, although plants can; several vitamins are compounds containing phenyl rings.