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Study Guide for Chemistry Honors Midterm

Courtesy of Illana Ben-Ezra

  1. Chapters 1 and 2
  2. Mass, weight, matter definitions
  3. Mass – Measurement or quantity of matter object has
  4. matter – objects that take up space and have mass
  5. Weight- measure of gravitational force acting on object
  6. Hypothesis, theory, law
  7. Hypothesis – testable statement or prediction
  8. Theory – broad explanation of experiments
  9. Law – concise statement
  10. Independent, dependent and controlled variables
  11. Dependant Variable – changes to affect Independent Variable
  12. Independent Variable – reacts to dependant Variable
  13. Control – remains same
  14. SI-length, mass, temperature, etc.
  15. Length – meter, m
  16. Mass – gram, g
  17. Temperature – Kelvin, K / Celsius, C
  18. Amount of substance – mole, mol
  19. Volume – Liter, L / cubic meter, m3
  20. Density – gram/cubic meter (g/cm3) , gram/milliliter (g/m/l)
  21. SI prefixes and what they stand for
  22. Giga- G, 10 to ninth
  23. Mega – M, 10 to sixth
  24. Kilo – K, 10 to third
  25. Hector – h, 10 to second
  26. Deca – da, ten to first
  27. base
  28. Deci – d, 10 to neg first
  29. Centi – c, 10 to neg second
  30. Milli- m, 10 to neg third
  31. Micro - │┴ (really connected), ten to neg sixth
  32. Nano – n, 10 to neg ninth
  33. Pico – p, 10 to neg twelfth
  34. Converting from one metric prefix to another
  35. King Henry Doesn't [Usually] (unit)Drink Chocolate Milk
  36. each step is ten times or one-tenth as much as the step on either side
  37. move from one prefix to the next by moving the decimal point one place to the left or right, filling in, as necessary, with zeroes
  38. Scientific notation
  39. Working with either huge or tiny numbers
  40. To multiply add exponents and multiply #
  41. Significant figures

i. All nonzero digits sig

  1. 2300 – SF are 2 and 3
  2. 1234.56 - 6 SF

ii. Zeros b/w SF, sig

1. 105 – 3 SF

iii. Trailing zeros in # with decimal are sig

iv. Trailing zeros in # with no decimal are not sig

v. leading Zeros never sig even if act as placeholder

  1. Adding SF
  2. # of SF after decimal in final sum/difference is determinded by lowest # of SF after decimal in any of orignals
  3. Multiplying/ Dividing SF
  4. # of SF in final product determined by smallest SF in any of original #s
  5. Rounding
  6. If digit to immediate right of last SF is less than 5, don’t change last SF
  7. If digit to immediate right of last SF greater than 5, round last SF up
  8. If digit to immediate right of last SF = to 5 and followed by nonzero  round last SF up
  9. If digit to immediate right of last SF = 5 and is not followed by nonzero digit  look at last SF, if odd round up, if even dont
  10. Accuracy and precision
  11. Precision – how close together or repeatable results are
  12. Accuracy – how close measurement is to actual measure
  13. Simple conversions
  1. Chapter 3
  2. States of matter
  3. Solid – definite shape and volume, particles tightly linked together
  4. Liquid – definite volume not shape, particles not as tightly linked
  5. Gas – no definite volume or shape, particles move rapidly b/c little to no attraction b/w particles
  6. Chemical and physical changes
  7. Chem – change in matter in which one or more new substances is produced, difficult to reverse
  8. cooking, burning, rusting, flaming, flammability, corrosiveness, acidity, ect..
  9. Phys –change in matter in which no new substance is produced, may involve change of state and can usually be reversed
  10. melting, freezing ect..
  11. Chapter 4
  12. Atomic theory
  13. Atoms can be destroyed but not through chem. Reactions
  14. Elements have isotopes
  15. Men and what they discovered (Thomson, Rutherford, etc.)
  16. J.J. Thomspon and atom
  17. named electrons
  18. “plum pudding”, electrons randomly mixed in positive ball
  19. described ration of electron charge to mass
  20. e/m = 1.76 x 10^8
  21. Milikan’s oil drop experiment
  22. atoms are neutral b/c protons = electrons
  23. Rutherford’s gold foil experiment
  24. fired alpha particle into foil, most went straight through, some slightly defected, very few deflected back
  25. proved atom mostly empty space w/ most mass in center (nucleus)
  26. Chadwick – discovered neutrons
  27. Atomic number, atomic mass, number of protons, neutrons and electrons
  28. Atomic # - # of protons (also # of electrons in norm atom), abbreviated as Z
  29. Atomic Mass
  30. avg mass of all naturally occurring non radioactive isotopes of element
  31. # of protons + # of neutrons
  32. Calculating average atomic mass
  33. (mass x percentage) +(mass x percentage) +(mass x percentage)
  34. Percentage should be decimal
  35. Chapter 5
  36. Calculating wavelength, frequency, and energy in photons of light
  37. C=vλ
  38. v = frequency (cycles per sec = hertz (Hz)), # of cycles that pass given point in sec
  39. to find frequency, divide C by λ
  40. λ = wavelength (distance from crest to crest or trough to trough)
  41. to find, divide C by v
  42. C= speed of light (3.00 x 10^8)
  43. E=nhv
  44. n = # of photons
  45. photon – packets of energy that make up light
  46. h = constant, 6.626 x 10^ -34
  47. E = energy of photon(s)
  48. Frequency and wavelength are proportional, if one increases other decreases
  49. frequency up, wavelength down
  50. Atomic emission series of light (Lyman, Balmer, etc.)
  51. When electron absorb energy can move to higher energy level, will drop back down to original level emitting the photon as light, when this light goes through prism it is shown in different colored lines called atomic spectrum, each element had own atomic spectrum
  52. lyman line – electron drops from higher level to lvl n=1, ultra violet
  53. balmer – to n=2, infa red
  54. paschen- n=3, infa red
  55. Brackett – n=4, infa red
  56. Pfund - n=5, infa red
  57. Principle quantum numbers (n)
  58. Indicates main energy level occupied by electron
  59. As n becomes bigger, increase in energy levels therefore further from nucleus therefore more energy
  60. 7 diff energy levels
  61. n= # of level
  62. further away from nucleus = more energy
  63. the # indicates the specific shell that electron belongs to which corresponds to period on table
  64. Azimuthal quantum number (l)
  65. Indicates angular movement and shape of cloud
  66. # of orbitals (l) in each principle energy level
  67. n=1, has 1 orbital
  68. s orbital
  69. s orbital is round
  70. has 1 sublevel, each sublevel holds 2 electrons
  71. n=2, has 2 orbital
  72. s and p orbitals
  73. p is dumbbell shaped
  74. p has 3 sublevels, each sublevel holds 2 electrons
  75. n=3, has 3 orbitals
  76. s, p, and d
  77. d has 5 sublevels, each sublevel holds 2 electrons
  78. n=4, has 4 orbital
  79. s, p, d, and f
  80. f has 7 sublevels, each sublevel holds 2 electrons
  81. magnetic quantum number (ml)
  82. indirect orientation of orbital around nucleus
  83. depends on l
  84. value b/w 1 and -1
  85. gives the 3D orientation of each orbital
  86. spin quantum number (ms)
  87. indicates spin of electrons in orbital
  88. if orbital sublevel has 2 electrons, spin opposite way

1. electrons always spin opposite

  1. Electron configuration including exceptions
  2. Aufbau principle
  3. each electron occupies lowest energy orbital available
  4. # is energy level
  5. S, p, d, f are orbitals
  1. Electron Configuration- ways electrons arrange around nucleus
  2. s is lowest level, but if filled move to next lowest, p, then d, then f, follow arrows in Aufbau diagram
  3. Pauli Exclusion Principle – no 2 electrons can have same 4 quantum #s
  4. Even if in same energy level, orbital, and sublevel, spin will be different b/c electrons cant spin same way
  5. Hund’s rule – when filling orbitals with more that one sub level, ½ fill each sublevel (arrows up), then go back and pair electrons (arrows down)
  6. 6C = 1s^2 2s^2 2p^2
  7. 1s^2 – 1 up 1 down (1st do up then down)
  8. 2s^2 – 1 up 1 down
  9. 2p^2 – 2 up b/c p has room for 3 and 3 down electrons, but 1st have to fill up electrons before go to down
  10. 8O = 1s^2 2s^2 2p^4
  11. Same as above except
  12. 2p^4 – 3 up 1 down, b/c filled ups so then go back to downs

6. exceptions

a. when s^2 and d^4 becomes s^1 d^5  s loses electron to d b/c more stable if ½ full

b. when s^2 d^9 s loses electron to make d full

c. s^2 f^6 and s^2 f^13  s loses one in both to make f full and ½ full

  1. Names of groups of elements (lA, llA, VIIA, VlllA, transition inner transition)
  2. VlllA – noble gases
  3. outermost s and p sublvl filled
  4. very stable
  5. don’t form compounds
  6. inner transition –
  7. fill f sublevel
  8. lanthanide and actinide series
  9. transition metals
  10. valence electrons present in more than one shell, often show several common oxidation states
  1. Chapter 6
  2. Naming compounds (also from 8)
  3. Periodic trends and definitions of each (?)
  4. Atomic radius
  5. L R gets smaller
  6. All added electrons are in same energy lvl
  7. Additional protons have strong electrostatic pull on electrons
  8. Top  Bottom increases
  9. Each energy level makes atom larger
  10. Shielding effect of inner shells block electrostatic attraction of nucleus to outermost electrons
  11. Ionic radius
  12. cations always smaller than atoms from which they were formed b/c increased pos charge but less neg charge so electron cloud drawn closer together
  13. anions gain electrons and are always bigger that their atoms b/c electrons make it puffier
  14. TB increases
  15. LR decreases
  16. Ionization energy
  17. “how much do I want to give you for your electron”
  18. if big jump in amount, electron wont come off
  19. energy required to remove an electron from gaseous atom of ion in its ground state
  20. amount of energy needed to make ion from gas
  21. LR increases
  22. Atomic radius gets smaller as go across making it harder to pull off electrons
  23. TB decreases
  24. Atomic radius larger and shielding effect makes easier to remove outer electrons
  25. electro negativity
  26. “how much do I want your electron”
  27. ability of atom to attract shared electrons
  28. values to help predict type of bonding
  29. TB decreases
  30. b/c of Shielding effect, more of an ability to give b/c shielding effect weakens bonds
  31. LR increases
  32. Left wants to lose b/c have less electrons and right wants to gain b/c need fewer electrons to be happy
  33. Characteristics of periods and groups on periodic table
  34. Metals
  35. high electrical and heat conductivity
  36. high luster
  37. ductile (can be made into wires)
  38. malleable (can be made into sheets)
  39. solid at room temp
  40. Nonmetals
  41. non lustourous
  42. poor conductors of heat and electricity
  43. some are gaseous (Ar, Cl)
  44. some brittle solids (Ca, S)
  45. metalloids
  46. properties of both metals and nonmetals
  47. border stair step
  48. representative series lA-VlllA
  49. phys and chem properties range
  50. outermost s and p levels only partially full
  51. alkali metals
  52. softer than most metals
  53. 1 valence electron so give up 1 electron easily
  54. highly reactive
  55. don’t occur freely in nature
  56. alkaline earth metals
  57. 2 electrons in outer shell, give up 2 easily
  58. highly reactive
  59. not as metalic
  60. Halogens
  61. valence of -1 (7 electrons in outer shell), gain one easily
  62. Basis of modern periodic table
  63. Meyer – periodic table of 56 elements based on atomic weight
  64. Mendeleev/ periodic law
  65. table based on atomic weight
  66. arranged periodically w/ elements having similar properties under each other
  67. mosely
  68. organized periodic table by atomic #
  69. Drawing Lewis dot structure
  70. 1st determine central atom
  71. written in middle
  72. arrange other atoms around it
  73. O rarely bonds to self
  74. H always at end
  75. F, Cl, Br, I occupy end position
  76. Count total # of valence electrons of atoms
  77. Draw single bonds b/w central atom and others
  78. Distribute electrons in pairs to satisfy octet rule
  79. Most electronegative gets first
  80. Form 2x or 3x bonds if cant give all atoms octets
  81. Xtra electrons put on central atom in pairs
  82. Chapter 8
  83. Definition of ionic, covalent (polar/nonpolar) and metallic bonds
  84. Ionic bonds –oppositely charged particles in ionic bond held together by electrostatic force
  85. high melting point b/c strong attraction b/w cations and anions
  86. soluble in water/ conduct electricity b/c of attraction to waters polarity
  87. hard brittle
  88. Covalent- sharing of electrons
  89. polar – uneven
  90. non polar – even
  91. b/w 2 nonmetals
  92. grps 4,5,6 share with each other when 2 atoms from 1 of those present
  93. metallic- outer electrons form cloud shared by all atoms in bond, no atom has more electrons than other, electrons float
  94. good electrical conductors b/c easily move from one place to another
  95. malleable/ductile b/c bond not so strong so when force applied atoms slide over each other
  96. heat conductors b/c atomic vibration (aka heat) can easily be passed from one atom to another
  97. Only first 39 slides of Chapter 9
  98. Sigma and pi bonds
  99. Sigma – single covalent bond
  100. bonded atoms rotate freely w/o losing contact
  101. pi
  102. double bond = 1 sigma and 1 pi
  103. triple = 1 sigma and 2 pi
  104. side to side parallel overlap enables sharing of electrons
  105. Single/double/triple covalent bonds and bond length
  106. Single = 1 pair of electrons
  107. Double = 2
  108. Triple = 3
  109. Bond length –
  110. distance b/w two bonding nuclei
  111. determined by atoms size and amount of atoms shared, the more pairs the closer b/c bond stronger

©2008 YHS Review.com and Ilana Ben-Ezra