Chem 102a - Test 1

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Chem 102a - Test 1
2011-09-27 14:56:57
chem chemistry

Chapters 1-5 Chemistry, A Molecular Approach
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  1. kinetic energy
    energy associated with motion
  2. potential energy
    energy associated with position
  3. total energy
    kinetic energy + potential energy
  4. thermal energy
    • -energy associated with temperature
    • -often converted from kinetic energy
  5. mass
    the measure of a quantity of matter within an object
  6. weight
    • the gravitational pull on an object
    • W=mg
  7. temperature
    measure of the the amount of average kinetic energy - the energy due to motion - of the atoms or molecules that compose the matter.

    • thermal energy transfers from hot to cold
    • 0 K is absolute zero -- molecular motion virtually stops
  8. intensive property
    porperty that is independent of the amount of substance. ex: density.
  9. extensive property
    property that is dependent on amount of substance. ex: mass.
  10. significant figure rules
  11. significant figures in calculations
    • multiplication or division: result carries the same number of significant figures as the factor with the fewest figures.
    • addition or subtraction: result carries the same number of decimal places as the quantity with the fewest decimal places.
    • rounding: only round final answer, not steps
  12. accuracy
    how close the measured value is to the actual value
  13. precision
    how close a series of measurements are to one another
  14. general problem solving strategy
  15. elements
    about 91 different naturally occuring elements
  16. law of definite proportions
    all samples of a given compound, regardless of their source or how they are prepared, have the same proportions as their constituent elements.
  17. law of multiple proportions
    when two elements (A and B) form two different compounds, the masses of element B that combine with 1g of element A can be expressed as a ratio of small whole numbers.
  18. atomic theory
    • each element is composed of tiny indestructible particles called atoms.
    • all atoms of a given element have the same mass and other properties which distinguish them.
    • atoms combine in simple, whole-number ratios to form compounds.
    • atoms of one element cannot change into atoms of another element. they can only change the way they are bound together through chemical reactions.
  19. discovery of electron
    • J.J. Thompson
    • by cathode ray tube
  20. electrical charge
    • the fundamental property of some of the particales that compose atoms, which results in attractive & repulsive forces (called electrostatic forces) between the particles.
    • proton=positive, electron=negative, neutron=neutral
  21. electric field
    the area around a charged particle where forces exist
  22. radioactivity
    emission of small energetic particles from the core of certain unstable atoms.
  23. rutherford's gold foil experiment
  24. nuclear theory
    • most of the atom's mass and all of it's positive charge are contained in its core called the nucleus (over 99.9%).
    • most of the volume of the atom is empty space, throughout which tiny, negatively charged electrons are dispersed.
    • there are as many electrons outside the nucleus as there are protons inside the nucleus. ie. the atom is electrically neutral.
  25. proton qualities
    • symbol: p+
    • charge: 1+
    • mass (amu): 1.007
    • mass (g): 1.673 x 10-24
    • location: nucleus
  26. neutron qualities
    • symbol: n0
    • charge: 0
    • mass (amu): 1.009
    • mass (g): 1.675 x 10-24
    • location: nucleus
  27. electron qualities
    • symbol: e-
    • charge: 1-
    • mass (amu): 5.486 x 10-4
    • mass (g): 9.109 x 10-28
    • location: outside the nucleus
  28. atomic number
    • represents the number of protons in an atom's nucleus.
    • unique to each element
  29. isotope
    atoms with the same number of protons but a different number of neutrons
  30. natural abundance
    the relative amount of each different isotope in a naturally occurring sample of a given element.
  31. mass number
    • the sum of the number of protons and neutrons
    • A=p+n
  32. ion
    an atom or molecule in which the total number of electrons is not equal to the total number of protons, giving it a net positive or negative electrical charge.
  33. cation
    • positively charged ion
    • typically metals
  34. anion
    • negatively charged ion
    • typically nonmetals
  35. periodic law
    • when elements are arranged in order of increasing mass certain sets of elements with recurring properties arrange periodically.
    • Medeleev
  36. periodic metals
    • lower left side
    • good conductors of heat and electricity
    • malleability (can be pounded into sheets)
    • ductility (can be drawn into wires)
    • shiny
  37. periodic nonmetals
    • upper right side
    • poor conductors of heat and electricity
    • varied properties
  38. periodic metalloids
    • lie along the zigzag line which divides metals and nonmetals.
    • mixed properties
    • semiconductors
  39. noble gases
    • group 8A
    • mostly unreactive nonmetals
    • full valence shell
  40. alkali metals
    • group 1A
    • reactive metals
  41. alkaline earth metals
    • group 2A
    • fairly nonreactive metals
  42. halogens
    • group 7A
    • highly reactive nonmetals
  43. atomic mass
    • represents the average mass of the isotopes that compose the element.
    • weighted according to the natural abundance of each isotope.
  44. mole (mol)
    • the standard scientific unit for dealing with atoms in macroscopic quantities.
    • which is defined arbitrarily as the amount of a substance with as many
    • atoms or other units as there are in 12 grams of the carbon isotope C-12.
  45. avogadro's number
    the number of atoms in a mole is called avogadro's number, the value of which is approximately 6.022 × 1023 particles.
  46. molar mass
    an element's molar mass in grams per mole is numerically equal to the element's atomic mass in atomic mass units.

    mass of 1 mol=molar mass
  47. chemical bonds
    the result of interactions between charged particles.

    can be strong bonds such as ionic or covalent or weaks such as hydrogen bonding.
  48. ionic bonds
    • occur between metals and nometals
    • involve the transfer of electrons from one atom to another. in this type the electron is not shared at all but transferred to the outer atomic orbital of the atom with the vacancy. one atom will have a net neg charge and the other a net positive charge.
  49. covalent bonds
    • bonds occurring between two or more nonmetals
    • involve the sharing of electrons between two atoms. ie. the electrons are attracted to the nuclei of both atoms so they are "shared" by both. low to nonexistent electronegative difference.
    • most organic compounds
  50. empirical formula
    simplest whole number ratio of atoms of each element present in a compound.

    standard for ionic compunds and macromolecules. different compunds can have the same empirical formula.

    ex: glucose = CH2O
  51. molecular formula
    identifies the number of each type of atom in a molecule. specific to one molecule of a certain compound.

    ex: glucose = C6H12O6
  52. structural formula
    shows the structure of the molecule using lines to represent covalent bonds.

    ex: glucose
  53. ball and stick model
    represents atoms as balls and bonds as sticks

    ex: glucose
  54. space filling model
    atoms fill the space to more closely spresent our best estimates of how a molecule might appear if scaled to a visible size.

    ex: glucose
  55. atomic elements
    • exist in nature with single atoms as their basic units,
    • most elements fall into this category.
  56. molecular elements
    • elements which exist as molecules.
    • most are diatomic, ex: O2, H2
  57. molecular compounds
    generally composed of two or more covalently bonded nonmetals.

    • you can recognize molecular compounds because the first element in the compound name is a nonmetal. some molecular compounds contain hydrogen, but if you see a compound which starts with "H", you can assume it is an acid and not a molecular
    • compound. compounds consisting only of carbon with hydrogen are called hydrocarbons. hydrocarbons have their own special nomenclature, so they are treated differently from other molecular compounds.
  58. ionic compunds
    • composed of cations and anions bound together by ionic bonds.
    • basic unit is the formula unit.
    • they are held together by the electrostatic force between oppositely charged bodies. Ionic compounds have a high melting and boiling point, and they are hard and very brittle. They are also called salts .
  59. formula unit
    • the lowest whole number ratio of ions represented in an ionic compound.
    • ex: NaCl and K2O
    • Ionic compounds do not exist as individual molecules; a formula unit is used to indicate the lowest reduced ratio of ions in the compound.
  60. polyatomic ions
    an ion composed of two or more atoms.
  61. acid
    • molecular compounds that release hydrogen ions (H+) when dissolved in water.
    • react with bases and metals
    • sour taste
    • pHs of acids are less than 7 because the concentration of hydronium ions is greater than 10−7 moles per liter.
    • can be binary or oxyacid
  62. binary acids
    composed of hydrogen and a nonmetal

    [hydro] + [base name of nonmetal + ic] + [acid]

    ex: HCl, HF. HI
  63. oxyacid (or oxoacid)
    contain hydrogen and an oxyanion

    • -ate ending: [base name oxyanion + ic] + [acid]
    • -ite ending: [base name oxyanion + ous] + [acid]

    ex: sulfuric, nitric, phosphoric, chlorous
  64. formula mass
    the sum of the atomic weights of the atoms in the empirical formula of the compound.

    • ex: the molecular formula for glucose is C6H12O6, so the empirical formula is CH2O.
    • the formula mass of glucose is (12)+2(1)+16 = 30 amu.