Chapter 4 & 5 Chemistry

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Chapter 4 & 5 Chemistry
2014-03-09 16:59:25
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  1. Dual Wave-Particle Nature
    In the early twentieth century, light was determined to have a dual wave-particle nature. Meaning light is both a wave and a particle.
  2. Rutherford
    The major shortcoming of Rutherford's model was that it was incomplete. It did not explain how the atom's negatively charged electrons are distrubuted in the space surronding its positively charged nucleus.
  3. Electromagnetic Radiation
    A form of energy that exhibits wavelike behavior as it travels through space. Examples: X rays, ultraviolet and infared lights, microwaves, and radiowaves.
  4. Electromagnetic Spectrum
    Together, all the forms of electromagnetic radiation form the electromagnetic spectrum.
  5. Electromaganetic Radiation Moves
    All forms of electromagnetic radiation moce at a constant speed of 3.0x108 m/s.
  6. Wavelength ( λ-Lambda )
    Is the distance between corresponding points on adjacent waves. The unit for wavelength is a distance unit.
  7. Frequency ( v-nu )
    Is the defined as the number of waves that pass a given point in a specific time, usually one second.
  8. Photoelectric Effect
    Refers to the emission of electrons from a metal when light shines on a metal.
  9. Quantum of Energy
    Is the minimum quantity of energy that can be lost or gained by an atom.
  10. Planck's Constant
    6.626x10-34J. s
  11. Albert Einstein
    In 1905, Einstein expanded on Planck's theory by introducing the radical idea that electromagnetic radiation has a dual wave particle nature. Although light has many wave like features it can be thought of as a stream of particles.
  12. Photons
    Is a particle of elctomagnetic radiation having zero mass and carrying a quantum of energy. The energy of a particluar photon depends on the frequency of radiation.
  13. Ground State
    The lowest energy state of an atom is its ground state.
  14. Excited State
    A state in which an atom has a higher potential energy than it has in its ground state is an excited state.
  15. Photons are given off when...
    An excited atom returns back to geound state, it gives off energy it gained in the form of elctromagnetic radiation. We see photons in the form of a color.
  16. Line-Emission Spectrum
    When a narrow beam of the emitted light was shined through a prism, it was seperated into four specific colors of the visible spectrum. These four bands are known as the line emission spectrum.
  17. Continous Spectrum
    The emission of a continous range of frequencies of electromagnetic radiation is continous spectrum.
  18. Emission
    When the electron falls to a lower energy level, a photon is emitted, the process is called emission.
  19. Absorption
    Energy must be added to an atom in order to move an electron from a lower energy level to a higher enegry level. This is called absorption.
  20. Bohr's Model of the Hydrogen Atom
    He solved the puzzle of the hydrogen-atom spectrum in 1913. His model linked the atom's electron to photon emission. According to the model, the electron can circle the nuclues in allowed paths or orbits. When the atom is in one of these orbits the atoms has definite energy.
  21. Quantum Theory
    Quantom theory was developed to explain observations such as the photoelectric effect and the line emission spectrum of hydrogen. It describes mathematically the wave properties of electrons and other small particles.
  22. Louis de Broglie
    Questioned whether electrons have a dual particle nature just like light.
  23. Diffraction
    The bending of a wave as it passes by the edge of an object or through a small opening.
  24. Interference
    Occurs when waves overlap.
  25. Heisenburg Uncertanity Principle
    It is impossible to determine simultaneously both the position and velocity of an electron or any other particle.
  26. Electrons do not...
    They do not travel around the nucleus in orbitals instead they exist in regions called orbitals.
  27. Orbital
    Is a three dimensional region around the nucleus that indiates the probable location of an electron.
  28. Quantum Number
    Specify the properties of atomic orbitals and the properties of electrons in orbitals.
  29. Principle Quantum
    Symbolized by n, indicates the main energy level occupied by the electron.
  30. Angular Momentum Quantum Number
    Symbolized by l, indicates the shape of the orbital.
  31. Magnetic Quantum Number
    Symbolized by m, indicates the orientation of an orbital around the nucleus.
  32. Spin Quantum Number
    +1/2    -1/2 Indicate the two fundamental spin states of an electron in an orbital.
  33. Electron Configuration
    The arrangment of electrons in an atom.
  34. Aufbau Principle
    An electron occupies the lowest energy orbital that can reieve it.
  35. Pauli Exclusion Principle
    If two electrons occupy the same orbital then they must have opposite spin states.
  36. Hund's Rule
    Electrons are spread out in singulary occupied orbitals before a second electron is added.
  37. Noble Gas
    The Group 18 elements (helium, neon, argon, kyrpton, xenon, & radon).
  38. Noble Gas Configuration
    Refers to an outer main energy level occupied, in most cases, by eight electrons.
  39. Periodic Law
    The physical and chemical properties of the elements are periodic functions of their atomic numbers.
  40. Periodic Table
    Is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group.
  41. Ar
    Found by Sir William Ramsey & John William Strutt
  42. Lanthanide
    Are the 14 elements with atomic numbers from 58 (cerium, Ce) to 71 (lutetium, Lu).
  43. Actinide
    Are the 14 elements with atomic numbers from 90 (thorium, Th) to 103 (lawrencium, Lr).
  44. Alkali Metals
    The elements of group 1 of the periodic table (lithium, sodium, potassium, rubidium, cesium, and francium) are known as the Alkali metals.
  45. Alkaline-Earth Metals
    The elements of group 2 of the periodic table (beryllium, magnesium, calcium, barium, and radium) are called the alkaline-earth metals.
  46. Transition Metals
    The d-block elements are metals with typical metallic properties and are often referred to as transition metals.
  47. Main-Group Elements
    The p-block elements together with the s-block elements are called the main group elements.
  48. Halogens
    The elements of group 17 (flourine, chlorine, bromine, iodine, and astatine) are known as halogens. The halogens are the most reactive nonmetals They form salts.
  49. Atomic Radius
    • 1/2 the distance between a nuclei of identical atoms that are bonded together. 
    • Across the Period: Decrease. The positive charge of the nuclues increase and everything moves closer to the nucleus decreasing its size.
    • Down the Group: Increase. Theres more electrons so theres more repulsion.
  50. Ion
    Is an atom or group of bonded atoms that has a positive or negative charge.
  51. Ionization
    Any process that results in the formation of an ion is refered to as an ion.
  52. Ionization Energy
    • The energy required to remove one electron from a neutral atom of an element.
    • Across the Period: Increase. This increase is caused by increasing nuclear charge, the more electrons in your final shell the more energy it tkaes to get electrons.
    • Down the Group: Decrease. The electons get farther away from the nucleus so it is easier to take away their electrons.
  53. Electron Affinity
    The energy change that occurs when an electron is aquired by a neutral atom is called the atom's electron affinity.
  54. Cation
    A positive ion is known as a cation.
  55. Anion
    A negative ion is known as a anion.
  56. Valence Electrons
    The electrons avaliable to be lost, gained, or shared in formation of chemical compounds are referred to as valence electrons.
  57. Electronegativity
    • Is a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound.
    • Across the Period: Increase. The number of valence electrons increases.
    • Down the Group: Decrease. The number of electron shells increase but not the # of valence electrons so there is less attratcive force between the nucleus and valence e
  58. Ionic Radius: Metals
    Metals beome cations because they lose electrons. The size of the ionic radius decreases as it loses electrons. (Less electrons, less repulsion, less space taken up)
  59. Ionic Radius: Nonmetals
    Nonmetals become anion because they gain electrons. The size of the ionic radius increases as gains electrons. (More electrons, more repulsion, more space taken up)
  60. Mendeleev
    "Father of the Periodic Table;" 1st person to arrange known elements into an orgainzed table.Arranged them by atomic mass w/exceptions, left gaps for undiscovered elements Sc, Ga, Ge, & He.
  61. Moseley
    After the discovery of subatomic particles, Mooseley rearranged the periodic table by atomic number.