1. E LECTRONS IN A TOMS Chapter 5

2. L IGHT AND Q UANTIZED E NERGY Nuclear atom and unanswered questions Scientists found Rutherford’s nuclear atomic model fundamentally incomplete Did not explain how electrons are arranged Did not address why negatively charged electrons are not pulled into the atom’s positively charged nucleus Wave nature of light Electromagnetic radiation – form of energy that exhibits wavelike behaviors Visible light Microwaves X-rays Radio/TV waves

3. Wavelength (λ) – shortest distance between equivalent points on a continuous wave Frequency (ν) – number of waves that pass a given point per second Measured in hertz (Hz) Ex: 652 Hz – 652 waves/s Amplitude – wave’s height from the origin to a crest, or from the origin to a trough All electromagnetic waves travel at a speed of 3.00x10 8 m/s in a vacuum Speed of light is represented by “c” c = λ ν Wavelength and frequency are inversely related Sunlight passing through a prism is separated into a continuous spectrum of colors

4. Particle nature of light The quantum concept – Max Planck concluded that matter can gain or lose energy only in small, specific amounts called quanta Quantum – minimum amount of energy that can be gained or lost by an atom E quantum = hv where “E” is energy, “h” is Planck’s constant, and “v” is velocity Planck’s constant = 6.626x10 -34 J·s Planck’s theory: for a given frequency, v, matter can emit or absorb energy only in whole-number multiples of hv (1hv, 2hv, 3hv, etc.) Analogous to child building a wall with wooden blocks

5. The photoelectric effect – electrons (photoelectrons) are emitted from a metal’s surface when light of a certain frequency shines on the surface Albert Einstein proposed light has both wavelike and particlelike characteristics Photon – particle of electromagnetic radiation with no mass that carries a quantum of energy Photon’s energy depends on its frequency E photon = hv Atomic emission spectra – set of frequencies of the electromagnetic waves emitted by atoms of the element Each element’s atomic emission spectrum is unique and can be used to determine if that element is part of an unknown compound Neon - light is produced by passing electricity through a tube filled w/ neon gas Neon’s atomic emission spectrum consists of several individual lines of color, not a continuous range of colors as seen in the visible spectrum

6. Q UANTUM T HEORY AND THE A TOM Bohr model of the atom – Neils Bohr proposed that elements’ atomic emission spectra are discontinuous Energy states of hydrogen Ground state – lowest allowable energy state of an atom Excited state – when an atom gains energy The smaller the electron’s orbit the lower the atom’s energy state Hydrogen’s line spectrum - when in the excited state the electron ca drop from the higher-energy orbit to a lower-energy orbit and the atom emits a photon Fig 5-10

7. Quantum mechanical model of the atom – Louis de Broglie accounted for fixed energy levels of Bohr’s model Electrons behave as waves Only half-wavelengths are possible on a guitar b/c the string is fixed at both ends Only whole numbers of wavelengths are allowed in a circular orbit of fixed radius Fig 5-11 Heisenberg Uncertainty Principle – states it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time Impossible to measure an object w/o disturbing it Tried to measure electrons w/ light but b/c a photon has about the same energy as an electron, the interaction changes the electron’s position

8. Quantum mechanical model of the atom – electrons are treated as waves Atomic orbital - 3-dimensional region around nucleus that describes the electron’s probable location Fig 5-13 Hydrogen’s atomic orbitals Principal quantum numbers – indicate the relative sizes and energies of atomic orbitals As “n” increases, the orbital becomes larger, the electron spends more time farther from the nucleus, and the atom’s energy level increases Principal energy levels - atom’s major energy levels (specified as “n”) Energy sublevels Principal energy level 1 consists of a single sublevel; principal energy level 2 consists of 2 sublevels, etc.

9. Sublevels are labeled s, p, d, or f according to the shapes of the atom’s orbitals Each orbital may contain at most 2 electrons All “s” orbitals are spherical All “p” orbitals are dumbbell shaped Not all “d” or “f” orbitals have the same shape Fig 5-15 and 5-16 Table 5-2

10. E LECTRON C ONFIGURATIONS Ground-state electron configurations Electrons tend to assume the arrangement that gives the atom the lowest possible energy The aufbau principle – states that each electron occupies the lowest orbital available All orbitals related to an energy sublevel are of equal energy Energy sublevels within a principle energy level have different energies (Fig 5-17) In order of increasing energy, the sequence of energy sublevels within a principal energy level is s, p, d. and f Orbitals related to energy sublevels within one principal energy level can overlap orbitals related to energy sublevels within another principal level

11. The Pauli exclusion principal – states that a maximum of 2 electrons may occupy a single atomic orbital, but only if the electrons have opposite spins Hund’s rule – states that single electrons w/ the same spin must occupy each equal-energy orbital before additional electrons w/ opposite spins can occupy the same orbitals Orbital diagrams and electron configuration notations Orbital diagram - boxes with zero, one, or two arrows represent orbitals Electron configuration notation – designates the principal energy level and energy sublevel associated w/ each of the atom’s orbitals and includes a superscript representing the number of electrons in the orbital Noble-gas notation

12. Valence electrons – electrons in the atom’s outermost orbitals Electron-dot structure – consists of the element’s symbol, which represents the atomic nucleus and inner- level electrons, surrounded by dots representing the atom’s valence electrons (Lewis dot structure)