Chemistry 130 Test 3 Study Guide

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Chemistry 130 Test 3 Study Guide
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  1. Compounds containing which ions are mostly soluble?
    • Li+ (lithium)
    • Na+ (sodium)
    • K+ (potassium)
    • NH4+  (ammonium)
    • NO3- (nitrate)
    • C2H3O2- (acetate)
    • Cl- (chlorine)
    • Br- (bromine)
    • I- (iodine)
    • SO42-  (sulfate)
  2. Cl-, Br-, I- are normally soluble.  However, the compound they are a part of becomes insoluble when paired with which ions?
    • Ag+
    • Hg22+
    • Pb2+
  3. Compounds containing SO42- are usually soluble, except when it is paired with which ions?
    • Sr2+ (strontium)
    • Ba2+ (barium)
    • Pb2+ (lead)
    • Ca2+ (calcium)
  4. Compounds containing which ions are mostly insoluble?
    • OH- (hydroxide)
    • S2- (sulfur)
    • CO32- (carbonate)
    • PO43- (phosphate)
  5. Compounds containing OH- and S2- ions are usually insoluble.  They become soluble though when these ions are paired with which others?
    • Li+
    • Na+
    • K+
    • NH4+
  6. Compounds containing CO32- and PO43- ions are usually insoluble.  They become soluble though when these ions are paired with which others?
    • Li+
    • Na+
    • K+
    • NH4+
  7. S2- is a special type of ion.  Compounds containing S2- ions are usually insoluble, except when paired with which others?
    The usual suspects:

    • Li(lithium)
    • Na+ (sodium)
    • K+ (potassium)
    • NH4+ (ammonium)

    but also:

    • Ca2+ (calcium)
    • Sr2+ (strontium)
    • Ba2+ (barium)
    • Pb2+ (lead)
  8. What are the "usual suspects" (ions) that when paired with, will turn an insoluble compound into a soluble one?
    • Li+
    • Na+
    • K+
    • NH4+
  9. Li+, Na+ K+, NH4+ are ions that make compounds soluble.  What are some exceptions?
    None.
  10. NO3- (nitrate) and C2H3O2- (acetate) are ions that make compounds soluble.  What are some exceptions?
    None.
  11. Define precipitation reaction.
    A reaction that forms a solid.

    The solid is known as the precipitate.
  12. What is the key to predicting a precipitation reaction?
    Only insoluble compounds form precipitates.
  13. Define molecular equation.
    An equation showing the complete neutral formulas for every compound in the reaction.

    example:

    AgNO3(aq) + NaCl(aq) --> AgCl(s) + NaNO3(aq)
  14. Define complete ionic equation.
    Equations showing the reactants and products as they are actually present in the solution.

    example:

    Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) --> AgCl(s) + Na+(aq) + NO3-(aq)
  15. Define spectator ions.
    Ions that do not participate in the reaction of a chemical equation.

    example:

    Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) --> AgCl(s) + Na+(aq) + NO3-(aq)
  16. Define net ionic equations.
    Equations that only show the species that actually participate in the reaction.

    example:

    Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) --> AgCl(s) + Na+(aq) + NO3-(aq)

    would be written as:

    Ag+(aq) + Cl-(aq) --> AgCl(s)

    (spectactor ions are omitted).
  17. Define solution.
    A homogeneous mixture of two or more substances.
  18. Define aqueous solution.
    A homogeneous mixture of a substance with water.
  19. Define solvent.
    The majority component of a solution.
  20. Define solute.
    The minority component of a solution.
  21. Define precipitate.
    An insoluble product formed through the reaction of two solutions containing soluble compounds.

    (the solid that is formed)
  22. Define strong electrolyte.
    A substance whose aqueous solutions are good conductors of electricity.
  23. Define soluble.
    Dissolves in solution.
  24. Define solubility.
    The amount of a compound, usually in grams, that will dissolve in a certain amount of solvent.
  25. Define insoluble.
    Not soluble in water.  

    Does not dissolve in water.
  26. Define miscible.
    Forming a homogeneous mixture when added together.

    example:

    "sorbitol is miscible with glycerol"
  27. Define acid-base reaction.
    Reactions that form water upon mixing of an acid and a base.
  28. Define gas evolution reaction.
    Reactions that evolve a gas.
  29. True or False: Many gas evolution reactions also happen to be acid-base reactions.
    True
  30. Define neutralization reaction.
    Another name for acid-base reaction, which is a reaction that forms water when an acid and a base are mixed.
  31. Acid-base reactions generally form water and a _______ called a _______, that usually remains dissolved in the solution.
    ionic compound; salt
  32. What is the net ionic equation for many acid-base reactions?
    H+(aq) + OH-(aq) ---> H20(l)
  33. What intermediate products are formed when sulfides react in a gas evolution reaction?
    None.  Sulfides do not result in an intermediate product.
  34. What gas is evolved when sulfides react in a gas evolution reaction?
    H2S
  35. What intermediate product do carbonates and bicarbonates yield when they react in a gas evolution reaction?
    H2CO3
  36. Which gas is evolved from carbonates and bicarbonates reacting in a gas evolution reaction?
    CO2
  37. What intermediate product do sulfites and bisulfites yield when they react in a gas evolution reaction?
    H2SO3
  38. Which gas is evolved from sulfites and bisulfites reacting in a gas evolution reaction?
    SO2
  39. What intermediate product does ammonium yield when reacting in a gas evolution reaction?
    NH4OH
  40. Which gas is evolved from ammonium reacting in a gas evolution reaction?
    NH3
  41. What does H2CO3(aq) decompose into?
    H2O(l) and CO2(g)
  42. Define oxidation-reduction reaction a.k.a. redox reaction.
    Reactions involving the transfer of electrons.
  43. Many redox reactions involve the reaction of a substance with _______.
    oxygen
  44. True or False: redox reactions must involve oxygen.
    False.
  45. What is a fundamental definition of oxidation?
    The loss of electrons.
  46. What is a fundamental definition of reduction?
    The gain of electrons.
  47. Oxidation and reduction must occur _______.
    together.

    If one substance loses electrons (oxidation), then another substance must gain electrons (reduction).
  48. Redox reactions are those in which...
    • A substance reacts with elemental oxygen
    • A metal reacts with a nonmetal
    • More generally, one substance transfers electrons to another substance.
  49. Combustion reactions are a type of _______.
    Redox reaction
  50. Which element is being oxidized and which is being reduced:  Sn2+ + Pb4+ --> Sn4+ + Pb2+
    Tin (Sn) is being oxidized because it is losing electrons, and Lead (Pb) is being reduced, because it is gaining electrons.
  51. A combustion reaction is a reaction with _______ involved.
    Oxygen
  52. Combustion reactions will always yield products of _______ and _______.
    CO2 and Water (H2O)
  53. Compounds containing carbon and hydrogen, or carbon hydrogen and oxygen, always form _______ and _______ upon combustion.
    carbon dioxide (CO2) and water (H2O).
  54. In an incomplete combustion of a hydrocarbon, _______ and _______ are products.
    Soot (carbon, C) and carbon monoxide (CO).
  55. Name the type of reaction: A + B ---> AB
    synthesis or combination reaction
  56. Name the type of reaction: AB ---> A + B
    decomposition reaction
  57. Name the type of reaction: A + BC ---> AC + B
    displacement reaction (or single-displacement reaction).
  58. Name the type of reaction: AB + CD ---> AD + CB
    double-displacement
  59. What happens in a synthesis or combination reaction?
    Simpler substances combine to form more complex substances.

    example:

    2Na(s) + Cl2(g) ---> 2NaCl
  60. What happens in a decomposition reaction?
    A complex substance decomposes to form simpler substances.  The simpler substances may be elements or compounds.

    example:

    • 2H2O(l) ---> 2H2(g) + O2(g)
    •                     (elements)


    • CaCO3(s) ---> CaO(s) + CO2(g)
    •                       (compounds)
  61. What happens in a displacement (single-displacement) reaction?
    One element displaces another in a compound.

    example:

    Zn(s) + CuCl2(aq) ---> ZnCl2(aq) + Cu(s)
  62. What happens in a double-displacement reaction?
    Two elements or groups of elements in two different compounds exchange places to form two new compounds.

    example:

    AgNO3(aq) + NaCl(aq) ---> AgCl(s) + NaNO3(aq)
  63. Define stoichiometry.
    It allows us to predict the amounts of products that form in a chemical reaction based on the amounts of reactants.

    It also allows us to predict how much of the reactants are necessary to form a given amount of product, or how much of one reactant is required to completely react with another reactant.
  64. Define limiting reactant.
    The reactant that makes the least amount of product.

    The reactant that is completely consumed in a chemical reaction.
  65. Define theoretical yield.
    The amount of product that can be made in a chemical reaction based on the amount of limiting reactant.
  66. Define actual yield.
    The amount of product actually produced in a chemical reaction.
  67. Define percent yield.
    The percentage of the theoretical yield that was actually attained.
  68. How do you calculate percent yield?
    • .                 actual yield
    • % yield = --------------------  x 100%
    •                theoretical yield
  69. When calculating how much product can be made from each reactant, the _______ amount determines the _______.
    smallest; limiting reactant
  70. What is the theoretical yield in pancakes if you have 3 cups of flour, and 1 cup of flour equates to 5 pancakes?
    Theoretical yield = 15 pancakes

    •                      5 pancakes
    • 3 cups flour = -------------- = 15 pancakes
    •                      1 cup flour
  71. Define wavelength.
    The distance between adjacent wave crests.

    It's symbol is:  λ  (lambda)
  72. Define frequency.
    The number of cycles or crests that pass through a stationary point in one second.

    It's symbol is:  v
  73. Define amplitude.
    Measures the strength of waves in a wavelength by how high the crests and how low the troughs are.
  74. Define photon.
    A particle of light; a single packet of light energy.
  75. In a photon, the shorter the wavelength, the _______ the energy of the photon.
    greater
  76. _______ is a form of electromagnetic radiation.
    Light
  77. Frequency and energy of electromagnetic radiation are _______ related to its wavelength.
    inversely
  78. The speed of electromagnetic radiation (light) in a vacuum is _______.  It does not depend on _______ or _______ of the electromagnetic radiation.
    constant; wavelength; frequency
  79. In the electromagnetic spectrum, where are radio waves located?
    Far left, everything greater than 10-1.
  80. In the electromagnetic spectrum, where are microwaves located?
    Left of visible spectrum, between radio waves and infrared radiation.

    10-3 < x < 10-1
  81. In the electromagnetic spectrum, where is infrared radiation located?
    Just left of the visible spectrum, between microwaves and the visible spectrum.

    10-6.5 < x < 10-3
  82. In the electromagnetic spectrum, where is ultraviolet radiation located?
    Just to the right of visible spectrum, between the visible spectrum and x-rays.

    10-9 < x < 10-6.5
  83. In the electromagnetic spectrum, where are x-rays located?
    To the right of the visible spectrum, between ultraviolet radiation and gamma rays.

    10-11 < x < 10-9
  84. In the electromagnetic spectrum, where are gamma rays located?
    Far right of everything, x-rays are to its left.

    x < 10
  85. In the electromagnetic spectrum, where is the visible spectrum located?
    Sliver between infrared radiation and ultraviolet radiation.

    10-6.6 < x < 10-6.3
  86. Define quantized.
    Restrict the number of possible values of (a quantity) or states of (a system) so that certain variables can assume only certain discrete magnitudes.

    A fixed amount of possible values.
  87. The Bohr model: Electrons exist in _______ orbits at specific, fixed _______ and specific, fixed _______ from the nucleus.
    quantized; energies; distances
  88. The Bohr model: When _______ is put into an atom, electrons are excited to _______ orbits.
    energy; higher-energy
  89. The Bohr model: When an electron _______ (or _______) from a _______ orbit to a _______ orbit, the atom emits light.
    relaxes; falls; higher-energy; lower-energy
  90. The Bohr model: The energy (and therefore the _______) of the emitted light corresponds to the energy _______ between the two orbits in the transition.  Since these energies are fixed and discrete, the energy (and therefore the _______) of the emitted light is fixed and discrete.
    wavelength; difference; wavelength
  91. The Bohr model: Describe the difference between an electron relaxing from n = 3 to n = 2 and one relaxing from n = 2 to n = 1.
    The higher energy levels are more closely spaced than the lower ones, so the difference in energy between n = 2 and n = 1 is greater than the difference in energy between n = 3 and n = 2.

    The photon emitted when an electron falls from n = 2 to n = 1 therefore carries more energy, corresponding to radiation with a shorter wavelength and higher frequency.
  92. Radiation with a shorter wavelength means what about its frequency?
    Higher frequency.

    Short wavelengths yield higher frequencies.

    Long wavelengths yield lower frequencies.
  93. Quantum Mechanical Orbitals: What does 1s mean?
    1 is the principal quantum number, and s indicates the subshell of the orbital and specifies its shape.

    so n = 1 is the principal quantum number, and s means the orbital that looks like a smaller sphere around the nucleus.
  94. Quantum Mechanical Orbitals: The number of _______ in a given principal shell is equal to the value of n.
    subshells
  95. Quantum Mechanical Orbitals: The shell number directly corresponds to what?
    n = 1, 2, 3, such as 1s, 2s etc.

    1s means n = 1

    2s means n = 2

    and the number signifies the shell number, or principal shell.

    See page 296 and 297 in the book.
  96. Quantum Mechanical Orbitals: with n = 1, how many subshells are there, and what are their letters?
    1

    s
  97. Quantum Mechanical Orbitals: with n = 4, how many subshells are there, and what are their letters?
    4

    s, p, d, f
  98. Quantum Mechanical Orbitals: with n = 3, how many subshells are there, and what are their letters?
    3

    s, p, d
  99. Quantum Mechanical Orbitals: with n = 2, how many subshells are there?
    2

    s, p
  100. Quantum Mechanical Orbitals: the number of subshells in a given principal shell is equal to the value of what?
    It's equal to the value of n.
  101. Define orbital.
    The region of space an electron is most likely to be found.

    The region of space can be visibly shown with a probability map.
  102. Define valence electron.
    Electrons in the outermost principal shell (the principal shell with the highest principal quantum number, n).
  103. Define core electrons.
    Electrons that are not in the outermost shell.

    (Every electron that is not a valence electron).
  104. In the following electron configuration, what is the value of n, what is the number of valence electrons, and what is the number of core electrons?  O: 1s22s22p4
    Highest coefficient number is 2, so n = 2.  There are 2 of them, so adding their exponents together gives the number of valence electrons:

    2 + 4 = 6 valence electrons.

    1s22s22p4

    To find core electrons, add up the exponents of the rest of the configuration, so in this case, 2 core electrons from 1s2.
  105. What are the shapes of d orbitals?
  106. What are the shapes of p orbitals?
  107. What are the shapes of s orbitals?
  108. Define ionization energy.
    The energy required to remove an electron from the atom in the gaseous state.

    For example, the ionization of sodium can be represented with the equation:

    Na + Ionization energy ---> Na+ + 1e-
  109. As you move across a period, or row, to the right in the periodic table, ionization energy _______.
    increases
  110. As you move down a column (or family) in the periodic table, ionization energy _______.
    decreases
  111. Which area of the periodic table has the highest ionization energy?  Which has the lowest?
    Highest is upper right, lowest is lower left.
  112. Define metallic character.
    The properties typical of a metal, especially the tendency to loose electrons in a chemical reactions.  Elements become more metallic as you move from right to left across the periodic table.
  113. As you move across a period, or row, to the right in the periodic table, metallic character _______.
    decreases
  114. As you move down a column, or family, in the periodic table, metallic character _______.
    increases
  115. Which area of the periodic table has the highest metallic character?  Which has the lowest?
    Highest is lower left, lowest is upper right.
  116. Define octet rule.
    In chemical bonding, atoms transfer or share electrons to obtain outer shells with eight electrons.
  117. How is a Lewis structure of an anion usually written?
    Within brackets with the charge in the upper right corner (outside the brackets).

    example:

  118. Explain the differences of bond order, such as the differences between single, double, and triple bonds.
    Single bonds happen when two electrons are shared in a covalent bond.  It has the longest bond length so its the weakest bond.

    Double bonds happen when four electrons are shared in a covalent bond (2 pairs).  It has a shorter bond length than a single bond and therefore it is stronger than a single bond.

    Triple bonds happen when 6 electrons are shared in a covalent bond (3 pairs).  It has a shorter bond length than a double bond and therefore it is stronger than a double bond.

  119. Define resonance structures.
    Representing a molecule with two Lewis structures that form single and double bonds (or double and triple bonds) on alternating atoms in a molecule that are both equally correct, and equally existent in nature.

  120. Describe the different atomic arrangements (shape = molecular geometry) of molecules/polyatomic ions containing one central atom as well as the name of their atomic arrangements, degrees of separation, etc.
    • page 338 in the book
  121. Define electronegativity.
    The ability of an element to attract electrons within a covalent bond.
  122. Using the periodic table, how do you determine the elements that have the greatest electronegativity and which ones have the lowest?
  123. When given its shape, how do you determine if a molecule or polyatomic ion is polar or non-polar?
    Visualize each bond as a rope pulling on the central atom.  Is the molecule highly symmetrical?  Do the pulls of the ropes cancel?  If so, there is no net dipole moment and the molecule is non-polar.  If the molecule is asymmetrical and the pulls of the rope do not cancel, the molecule is polar.
  124. Describe Kinetic molecular theory of gases.
    • A gas is a collection of particles in constant motion.
    • No attractions or repulsion between particles; collisions like billiard ball collisions.
    • A lot of space between the particles compared to the size of the particles themselves.
    • The speed of the particles increases with increasing temperature.
  125. Describe Boyle's Law.
    The volume of a gas (V) and its pressure (P) are inversely proportional.  As volume increases, temperature decreases and vice versa.

    •         1
    • V  α  ---
    •         P

    (α is a Greek letter, Alpha, which is used to mean proportional/inversely proportional).

    Boyle's law also assumes that both temperature and the number of particles involved are constant.

    P1V1 = P2V2
  126. Describe Charles' Law.
    The volume (V) of a gas and its Kelvin temperature (T) are directly proportional.

    V α T

    (a is a Greek letter, Alpha, which is used to mean proportional/inversely proportional).

    Charles' Law also assumes constant pressure and a constant amount of gas.

    • V1     V2
    • --- = ---
    • T1     T2
  127. Describe the Combined Gas Law.
    The combined gas law is used when both temperature and pressure are not constant.

    The combined gas law applies only when the amount of gas is constant. 

    The combined gas law encompasses both Boyle's law and Charles' law and can be used in place of them.  If one physical property (P, V, or T) is constant, it will cancel out of your calculations when you use the combined gas law.

    So, pretty much always use the combined gas law.

    • P1V1       P2V2
    • ------ =  ------
    •   T1         T2
  128. Describe Avogadro's Law.
    The volume of a gas (V) and the amount of the gas in moles (n) are directly proportional.

    V α n

    • (a is a Greek letter, Alpha, which is used to mean proportional/inversely proportional).
    • Avogadro's Law assumes constant temperature and pressure.

    • V1     V2
    • --- = ---
    • n1     n2
  129. True or False: In Avogadro's Law when solving for how many moles of a substance were added, you must take into account what the starting mole (n) measurement is, and subtract that from the final result.
    true;

    This is done to find out how many moles were added to the pre-existing amount.
  130. Describe the Ideal Gas Law.
    The law encompassing Boyle's law, Charles' law, and Avogadro's law and combining them into a single expression:

    •       nT
    • V α ---
    •       P

    We can replace the proportional sign with an equal sign by adding R, a proprtionality constant called the ideal gas constant.

    •       Rn
    • V = ----
    •        P

    The value of R is:  

    •                  L*atm
    • R = 0.0821  -------
    •                  mol*K

    R can be expressed in other units, but its numerical value will be different.
  131. What are the only acceptable units of measurement in the Ideal Gas Law?
    • Pressure (P) must be expressed in atmospheres
    • Volume (V) must be expressed in liters.
    • Amount of gas (n) must be expressed in moles.
    • Temperature (T) must be expressed in Kelvins.
  132. Describe Dalton's law of partial pressures.
    The sum of the partial pressures of each of the components in a gas mixture must equal the total pressure.

    Ptot = Pa + Pb + Pc + ...

    tot = total pressure

    Pa, Pb, Pc are partial pressures of the components.
  133. Under which conditions do gas laws not work well, and why?
    Low temperatures and high pressures.

    Weak attractive forces between molecules in the gas and thus Kinetic molecular theory is not entirely true.  

    It falls apart at very low temperatures near the temperature where the gas would condense into a liquid.  

    Very high pressures, gas molecules are crammed together, which means that they no longer interact only through collisions.

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