Unit 1 - Chapter 3

Question 1

Max Planck

Answer 1

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

Question 2

Double slit experiment

Answer 2

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

Question 3

Photoelectric effect

Answer 3

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

Question 4

Photoelectric effect experiments

Answer 4

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

Question 5

Einstein’s explanation of the photoelectric effect

Answer 5

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

Question 6

Wave-particle duality

Answer 6

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

Question 7

Continuous spectrum

Answer 7

when white light is passed through a prism, it produces a continuous spectrum
continuous spectrum = rainbow of colours

Question 8

Dark line (absorption) spectrum

Answer 8

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

Question 9

Bright line (emission) spectrum

Answer 9

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

Question 10

Comparing dark and bright line spectra

Answer 10

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

Question 11

Niels Bohr

Answer 11

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)

Question 12

Bohr’s explanation of the dark line spectrum

Answer 12

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

Question 13

Bohr’s explanation of the bright line spectrum

Answer 13

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

Question 14

Louis de Broglie

Answer 14

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

Question 15

Werner Heisenberg

Answer 15

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

Question 16

Erwin Schrodinger

Answer 16

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)

Question 17

Max Born

Answer 17

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

Question 18

Wolfgang Pauli

Answer 18

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

Question 19

Paul Dirac

Answer 19

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

Question 20

Structure of the quantum atom

Answer 20

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

Question 21

Quantum numbers

Answer 21

principle quantum number = main energy level
secondary quantum number = sub-level
magnetic quantum number = orbital
spin quantum number = up or down spin

Question 22

Restrictions on quantum numbers

Answer 22

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)

Question 23

Sublevels

Answer 23

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)

Question 24

1st energy level

Answer 24

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

Question 25

2nd energy level

Answer 25

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

Question 26

3rd energy level

Answer 26

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

Question 27

Pauli exclusion principle

Answer 27

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

Question 28

Aufbau principle

Answer 28

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

Question 29

Hund’s rule

Answer 29

when there are several orbitals of the same energy within a sublevel, electrons will half-fill each orbital before pairing up

Question 30

Energy level and orbital overlap

Answer 30

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

Question 31

Unexplained electron configurations

Answer 31

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

Question 32

Electron configurations and ionization energies

Answer 32

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

Question 33

Ionic charges in transition metals

Answer 33

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

Question 34

Explaining magnetic properties

Answer 34

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

Question 35

Paramagnetic substances

Answer 35

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

Question 36

Ferromagnetic substances

Answer 36

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