Unit 1 - Chapter 3 Flashcards
Max Planck
proposed that energy is quantized (non-continuous)
this means that energy must be absorbed or released in discrete bundles that he called quanta
quanta = particles can only vibrate at defined frequencies, multiples of some fundamental frequency
Double slit experiment
performed by Thomas Young
when monochromatic light was passed through a screen with two slits, it produced an interference pattern (constructive and destructive interference)
this means that light behaves like a wave
Photoelectric effect
performed by Hertz and Lenard
electrons are emitted from certain metals as a result of absorbing energy from electromagnetic radiation (light)
these electrons are called photoelectrons
Photoelectric effect experiments
experiment 1 = a metal plate is exposed to red light, no photoelectrons detected, increase of intensity has no effect
experiment 2 = same metal plate is exposed to UV light, photoelectrons detected, increase of intensity causes photoelectrons to be detected at a faster rate
Einstein’s explanation of the photoelectric effect
he suggested it could only be explained if light has particle-like properties and travels in quantized packets called photons
only photons with appropriate energy knock electrons from metal
increasing the intensity increases the number of photons, but not the energy of the photons
Wave-particle duality
some experiments show that light has wave-like properties, while others show that it has particle-like properties
wave-particle duality = light is simultaneously both a wave and a particle
Continuous spectrum
when white light is passed through a prism, it produces a continuous spectrum
continuous spectrum = rainbow of colours
Dark line (absorption) spectrum
when white light is passed through a gaseous sample of an element, and then passed through a prism, it produces a pattern of dark lines on the continuous spectrum
this represents the specific frequencies of light that have been absorbed by the atoms of the element
Bright line (emission) spectrum
when a gaseous sample of an element is energized, and the light passes through a prism, it produces a series of light bands separated by regions of black
this represents the only energies of light that the excited atoms of an element can emit
Comparing dark and bright line spectra
the frequencies of light missing from the dark line spectrum are identical to the frequencies of light present in the bright line spectrum
the spectra are specific to each element
Niels Bohr
Bohr saw atomic spectra as evidence that the energy of electrons is quantized
he introduced the first quantum model of the atoms, concluding that electrons are confined to discrete energy levels
in unenergized atoms, the electrons would be found in their ground state (lowest possible energy level)
Bohr’s explanation of the dark line spectrum
electrons can be energized and excited to higher energy levels if they gain the exact amount of energy to make the transition
the missing frequencies of light in a dark line spectrum correspond to the specific quanta of energy needed for that transition
these frequencies are equal to the difference in energy between the two levels
Bohr’s explanation of the bright line spectrum
excited electrons eventually will return to lower energy levels by losing energy
this energy is released as a photon of light of a specific frequency
the distinct bands of light in the bright line spectrum correspond to the energy difference between the two levels
the energy differences between levels gets lesser in the higher levels
Louis de Broglie
he determined the wavelength of any moving particle given its mass and speed, as he reasoned that particles may also have wave properties like light
applying this idea to the atom, he discovered that an electron behaves like a standing wave bound to the nucleus, and there is only a defined number of electron wavelengths due to quantized energy
Werner Heisenberg
Heisenberg Uncertainty Principle = it is impossible to know simultaneously, with exact precision, the position and momentum of a particle
applying this to the structure of an atom, it is impossible to know with any degree of certainty where or how an electron moves in an atom
the uncertainty is about the size of the atom itself
Erwin Schrodinger
Schrodinger used conventional wave equations to develop a probabilistic quantum model of the atom
the solutions to this equation provides us with the electron’s wave function, which defines the probability of finding an electron in any given location
the region of probability of finding an electron is an orbital
these equations assigned an electron a set of three quantum numbers (energy level, sublevel, orbital)
Max Born
used Schrodinger’s equations to develop probability functions that could produce a plot of probability densities
this 3D volume of space in which there is a 90% probability of finding an electron represents the electron’s orbital
Wolfgang Pauli
proposed a fourth quantum number which relates to a property of the electron called spin which is responsible for an electron’s weak magnetic field
Paul Dirac
developed a new version of the wave equation that included the fourth quantum number
an electron can have one of two possible spins = up or down, resulting in two opposing magnetic fields
result = two spin paired electrons (of opposite spin) can occupy the same orbital
Structure of the quantum atom
electrons exist as wave functions and are restricted to fixed energy levels
within an energy level there exists a defined number of sublevels, which represent wave functions of different energy
the location of an electron within a sublevel is defined by an orbital
each orbital can hold one set of spin paired electrons
Quantum numbers
principle quantum number = main energy level
secondary quantum number = sub-level
magnetic quantum number = orbital
spin quantum number = up or down spin
Restrictions on quantum numbers
principle = (n) = any integer from 1 to infinity secondary = (l) = any integer from 0 to n-1 magnetic = (ml) = any integer from -l to +l spin = (ms) = up spin (½), down spin (-½)
Sublevels
0 = s sublevel, 1 = p sublevel, 2 = d sublevel, 3 = f sublevel
each energy level introduces a fundamental wave function (sublevel) and its associated orbitals
subsequent energy levels have a higher energy harmonic of the previous sublevels plus one new fundamental wave function (new sublevel)
1st energy level
the electrons in the 1s orbital occupy a spherical region of probability around the nucleus where the densest region is the region most likely to find an electron
the probability of finding the electrons decreases as you move away from the nucleus
2nd energy level
the 2s orbital has a node (a region of zero probability) separating two spherical regions of probability
the 2p orbital has a node on the nucleus and opposite algebraic signs on the two sides of the axis, creating three dumbbell shaped regions
3rd energy level
as the wave functions become more complex (from s to p to d to f), so do the shapes of the orbitals associated with them
the d sublevel consists of five 3d orbitals
Pauli exclusion principle
no two electrons in an atom may have the same four quantum numbers, as a result, electrons sharing in the same orbital must also have opposite spin
Aufbau principle
as electrons are added to an atom, they will seek the lowest possible energy, they will fill lowest energy levels and sublevels before higher ones are filled
Hund’s rule
when there are several orbitals of the same energy within a sublevel, electrons will half-fill each orbital before pairing up
Energy level and orbital overlap
as energy levels become more complex, higher energy orbitals of a lower energy level are higher in energy than the lower energy orbitals of a higher energy level
this first occurs with the 3rd energy level where the energy of 3d is higher than that of 4s
Unexplained electron configurations
electrons are easily moved between higher level s and d sublevels because, as we move along the periodic table, more protons are added to the nucleus, increasing the effective nuclear charge
this charge pulls the d electrons closer to the nucleus which results in a lower energy d sublevel
filled and half-filled sublevels result in greater electron stability than do partially filled sublevels
Electron configurations and ionization energies
an element with a higher atomic number may have a lower ionization energy than expected if there is a single valence electron in a higher sublevel
this is because the electron is more weakly attracted to the nucleus than an electron in a lower energy sublevel
it may also have a lower ionization energy than expected if the valence electrons are sharing an orbital, because there is already a repulsion force between the two electrons, making it easier to remove one of them
Ionic charges in transition metals
electrons in the highest energy level are, on average, further from the nucleus, and therefore easiest to remove
metal atoms will lose 1, 2, or at most 3 electrons to form stable ions, however in groups 4-15, they need to lose more than 3 to form a noble gas configuration
transition metals form stable ions by losing their outermost s electrons resulting in ions with a charge of 2+ or 1+
multivalent metals may also lose one of their d electrons to achieve a 3+ charge
Explaining magnetic properties
magnetic properties of substances are based on the magnetic field generated by the electron’s spin
when two electrons are spin-paired, their magnetic fields cancel out
when an atom has several unpaired electrons, the magnetic fields cause the atom itself to act like a tiny magnet
Paramagnetic substances
a substance that is weakly attracted to a magnetic field when placed between magnetic poles
the atoms and their magnetic poles are arranged randomly throughout the substance
magnetic properties are not maintained without the magnetic field
Ferromagnetic substances
a substance that is strongly attracted to a magnetic field and can retain its magnetic properties when removed from magnetic field
eg. iron, cobalt, nickel
atoms are arranged in clusters called domains
all of the atoms within a domain have their magnetic poles oriented in the same direction creating a much stronger magnetic field than an individual atom
