Electrons in Atoms Download

Chapter 5
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
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 3.00x108
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
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
 Equantum = 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
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
 Ephoton = hv
Atomic emission spectra – set of frequencies of the
electromagnetic waves emitted by atoms of the
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
Bohr model of the atom – Neils Bohr proposed that
elements’ atomic emission spectra are
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
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
Quantum mechanical model of the atom – Louis de
Broglie accounted for fixed energy levels of Bohr’s
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
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
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.
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
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
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
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
Valence electrons – electrons in the atom’s
outermost orbitals
Electron-dot structure – consists of the element’s
symbol, which represents the atomic nucleus and innerlevel electrons, surrounded by dots representing the
atom’s valence electrons (Lewis dot structure)