Transcript Metathesis Problems (and Some Solutions) Identified Through

Chapter 10: Modern atomic theory
Chemistry 1020: Interpretive chemistry
Andy Aspaas, Instructor
Rutherford’s atom
• Recall Rutherford’s atomic theory
– Positively charged nucleus
– Surrounded by negatively charged electrons
• Unanswered questions
– How are electrons arranged
– How do they move?
Electromagnetic radiation
• Electromagnetic radiation: energy transmitted by
waves, “radiant energy”
• Wavelength: distance between peaks of these
waves
• Different forms of electromagnetic radiation have
different wavelengths
Electromagnetic spectrum
Types of electromagnetic radiation
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Radio waves: low frequency and energy
Microwaves
Infrared
Visible
Ultraviolet
X-rays
Gamma rays: high frequency and energy
Energy and electromagnetic radiation
• The shorter the wavelength, the higher the energy
transmitted
– Blue light: shorter wavelength: higher frequency:
higher energy
– Red light: longer wavelength: lower frequency:
lower energy
Wave calculations
• Velocity = c = speed of light
• 2.997925 x 103 m/s
• All types of light energy travel at same speed
• Amplitude = A = height of wave, brightness of light
• Wavelength = = distance between peaks
• Frequency =  = number of waves that pass a point
in a given amount of time
– Generally measured in Hertz (Hz)
– 1 Hz = 1 wave/sec = 1 sec-1
• c = x 
Planck’s nuclear theory
• Light energy behaves as particles in certain
situations
• Each particle of light (a photon) has a certain fixed
amount of energy
– Energy of photon is directly proportional to
frequency of the light
– Higher frequency = more energy in photon
Atomic emission spectra
• Atoms that gain extra energy will release that energy in the
form of light
• Light is given off in very specific wavelengths
• Different atoms give off different characteristic wavelengths
of light when excited
– Line spectrum: shows wavelengths of light that are
emitted
– Only certain wavelengths are given off, so only specific
amounts of energy can be absorbed or given off for any
one type of atom
– Atoms are “quantized” - only specific energy levels
Bohr’s model of the atom
• Explains line spectrum of hydrogen
• Energy of atom is related to distance of electron from nucleus
– Electrons can “jump” to different possible orbits around
nucleus
– Gain in energy: electron jump to higher quantum level “excited state”
– Lines in spectrum correspond to difference in energy
levels
• Ground state: minimum energy level
• But, only explains hydrogen atom behavior
– Plus, electrons do not have simple circular orbits
Wave mechanical model of the atom
• Electrons can be treated as waves (in the same way
that light can also be treated as particles)
• Mathematics can calculate the probability densities
of finding an electron in a particular region of the
atom
– Schrödinger equation - cannot predict location of
any one particle, only probability of it being a
certain place
Orbitals
• Solutions to wave equations give regions of high
probability for finding electrons
– Called orbitals
– 90% probability of finding an electron
– 3-dimensional shape
Orbitals and energy levels
• Principal energy level (n) = how much energy the
electrons in the orbital have
– Higher values mean higher energy and farther
average distance from nucleus
• Each principal energy level has n sublevels
– Different shape and energy
– Named s, p, d, f
• Each sublevel has 1 or more orbitals
– s = 1 orbital, p = 3, d = 5, f = 7
Pauli exclusion principle
• No orbital may have more than 2 electrons
• Electrons in same orbital must have opposite spin
– s holds 2 electrons
– p holds 6 electrons
– d holds 10 electrons
– f holds 14 electrons
Electron configurations
• Hydrogen electron configuration: 1s1
– Superscript indicates number of electrons in
orbital
• Helium: 1s2
• Follow the periodic table: row number = principal
energy level (first number in electron configuration)
• Column and section determine which sublevel (s, p,
d, f) is filled
Valence electrons
• Valence electrons: only those in outermost energy
level - determine most of an atom’s reactivity
properties
• Can indicate Na electron configuration as
– 1s22s22p63s1 or [Ne]3s1 (using nearest Noble gas
with smaller atomic number than the atom)
Atomic properties and the periodic table
• Ionization energy: energy required to remove an electron
from an atom
– Decreases down a group (less energy required to remove
electron)
– Increases across a period (more energy required to
remove an electron)
• Atomic size
– Increases down group to account for greater mass
– But decreases across period because more electrons
mean more attraction to the nucleus