Chemistry Video 5 Flashcards
Issac Newton
Worked with prisms, showed that white light contains all the colours of the rainbow. He used a corpuscular model, which says that light is composed of particles
Christiaan Huygens
He used a wave model to describe reflection and refraction.
Thomas Young
Light travelling through narrow slits produced interference patterns that could only be explained using a wave model
James Clerk Maxwell
Developed theory of magnetic radiation and showed that visible light is only a tiny portion of the electromagnetic spectrum
Classical physics in 19th century
Newton’s laws governed particles of matter. Maxwell’s equations governed waves of light. This view was was incorrect. In reality, both matter and light exhibit wave-particle duality.
Newton’s classical mechanics and Maxwell’s classical electrodynamics is limited
Wave
Oscillations that transport energy from one place to another. The molecules stay in place but the kinetic energy travels. Has amplitude, wavelength, crest, trough, frequency
Electromagnetic spectrum
Longer wavelength = decreased frequency
Shorter wavelength = greater frequency = greater energy
Very short wavelengths transmit a lot of energy. Gamma rays have the shortest wavelength.
Very long wavelengths transmit very little energy. Radio waves have the longest wavelength.
Composition of light
Simultaneously both wave-like and particle-like properties. Called wave-particle duality.
Max Planck
Solved the Ultraviolet Catastrophe, which examined the distribution of light from the sun. Realized that the light received did not conform to mathematical expectations, which predicted much more UV light than what is actually received.
Planck introduced concept of quantization of energy, where vibrational energies of atoms cannot possess absolutely any arbitrary value from a spectrum of values. The values must be quantized, AKA possessing a multiple of a discrete unit of energy. This allowed the math to conform to observations of light
Photoelectric effect
Electrons are ejected from metal when irradiated with light above a threshold frequency. Below a certain threshold frequency of light, no matter how intense the light beam was, no electron was ejected. Above the threshold frequency, the electron was always ejected no matter how weak the beam was.
Albert Einstein
Explained photoelectric effect using principles developed by Planck when he worked with black body radiation. Stated that light is quantized. Describes light as stream of particles called photos, whose energy depend on their frequency. E = hv, where v is frequency and h is Planck’s constant.
Referring to photoelectric effect, an electron is ejected if a singular photon of sufficient energy strikes the metal, which is why the intensity of the beam is irrelevant.
Therefore, intensity of light corresponds with number of photons, rather than wave amplitude
Additionally, the greater the brightness, the greater the number of electrons ejected. Brighter light = more photons = higher likelihood that photons will collide with electrons
Photons are quanta, which are discrete fundamental units of energy
Photoelectric effect equation and work function
Equation: Kinetic energy of electron expelled from atom = Energy of the photon - work function
Work function is the minimum energy required to expel an electron from the atom. To expel an energized electron, the kinetic energy of the electron must be greater than 0. So, the energy of the photon must be greater than the work function.
Line spectra
When materials are heated, they radiate some heat off as light. The light will have a range of energy called a continuous spectrum. Some gases will not display a continuous spectrum, and will generate a line spectra. The line spectra shows only specific values for the light emitted, which are represented by certain lines, which each correspond to the specific wavelengths of light emitted.
Johann Balmer
Derived an equation (mathematical model) that related the 4 lines in the visible spectrum of hydrogen. “k” is a constant and “n” is one of four integers.
Johannes Rydberg
There were other lines found in the UV and IR regions for hydrogen. He generalized Balmer’s work to predict all of hydrogen’s emission lines, where “n” are integers and “R” is the Rydberg constant. This mathematical formula was developed to predict the numbers accurately before a conceptual explanation was offered for the mathematical truth
Niels Bohr
Formula from Rydberg was rationalized by him, which helped understand Rydberg’s formula. He set the trajectory of physics towards quantum mechanics
Planetary model
After Rutherford (who discovered the atomic nucleus), the model of the atom consisted of the nucleus and electrons orbiting. Resembles solar system. The simplest system is the H atom with 1 electron orbiting 1 proton.
However, this created problems. The math suggests that electrons should lose energy as they orbit and fall into the nucleus. This means that atoms should be unstable. But in reality, atoms are stable since they exist.
Bohr Model of the Hydrogen Atom
Used energy quantization. Electrons can inhabit different energy levels and electrons can move between energy levels if it absorbs or emits a photon.
Absorbing photon causes electron to move up to excited state. Emitting photon causes electron to move down to a lower state.
The electron will only emit or absorb photon when moving between energy levels. The energy of the photon equals the change in energy for the electron when moving between energy levels. The energy levels mean that the energy of the electron is quantized (specific values).
Created version of the Rydberg equation, calculating a theoretical value for the Rydberg constant, which agreed closely to the experimentally accepted value. This is strong evidence in favour for the assumptions that Bohr made.
Energy of photon absorbed = Change in energy in Joules unit = (Rydberg constant of 2.17910^(-18))[(1/ni^(2))-(1/nf^(2))]
ni is the initial energy level and nf is the final energy level
Also,
energy = (Planck’s constant of 6.62610^(-34) Jsec)*(frequency)
speed of light of 3.0010^8 m/sec = (wavelength)(frequency)
Louis de Broglie
Proposed that particles can also behave like waves (Not just waves that can behave like particles). All matter must have a wavelength called de Broglie wavelength.
Wavelength = (Planck’s constant)/[(mass}*(velocity)]
As mass increases, wavelength decreases.
Electrons have wave-particle duality.
Werner Heisenberg
Heisenberg uncertainty principle. Limits to ability to measure both position and momentum of a particle. The more we know about location, the less we know about momentum. It is a fundamental quality of matter.
(delta x)*(delta p) greater than or equal to (h bar / 2)
delta x = uncertainty in position
delta y = uncertainty in momentum
h bar = h/(2pi)
Also applies to energy/time pairing. Only applies to small particles.
Schrodinger Equation
Extended the work of de Broglie. Mathematical model that describes wave-like behaviour of electrons. When applied to the H atom, it reproduced many known values. Electrons are represented as 3D wave functions.
(Hamiltonian operator)(wave function psi) = (total energy of particle)(wave function psi)
Hamiltonian operator = set of mathematical operations representing the total energy of the particle
wave function psi = wave function used to describe the probability distribution of the particle
Max Born
Think of electrons as waves of probability density. The square of the magnitude of the wave function describes the probability of an electron existing in a particular space.
Quantum Mechanics
Describes the motion of quantum objects. Different from classical mechanics. Electrons are quantum objects.
Atomic orbitals
Can be assigned sets of quantum numbers. Orbitals are 3D regions of probability of where an electron may be found; based on Schrodinger equation
Principle quantum number, n
Indicates energy level or shell that the electron resides in. Higher n value means higher energy and further away from the nucleus.
Angular momentum quantum number (l)
Value is from 0 to (n-1). Tells us the type of orbital that the electron is in.
l = 0; s orbital, spherical, increase in radius as n increases l = 1; p orbitals, lobes that extend on each of the x, y, z axis l = 2; d orbitals
Magnetic quantum number (ml)
Values from - l to + l. Includes 0.
Spin quantum number (ms)
+1/2 or -1/2. Maximum of 2 electrons can fit in any atomic orbital. Each orbital can hold 2 electrons of opposing spin. Spin up or spin down.
Pauli Exclusion Principle
No 2 electrons in an atom can have the same set of 4 quantum numbers
n = energy level l = type of orbital ml = which orbital of a set ms = spin up or down
Aufbau principle
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p <5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p
Hund’s rule
When looking at a set of degenerate orbitals, each orbital will get one electron before any electrons are paired
Electron configuration
List all types of orbitals that are occupied and a superscript number to indicate number of electrons in each orbital.
n value, orbital type letter, superscript for number of electrons in the subshell
Orbital diagrams
Visual depictions of arrangement of electrons within the orbitals of an atom
Paramagnetic
One or more unpaired electron. Attracted towards a magnetic field. Ex. 1s^2 2s^1
Diamagnetic
All electrons are paired. Weakly repelled by a magnetic field. Ex. 1s^2 2s^2
Exceptions in electron configurations
Ex. Cu: [Ar] 4s^1 3d^10
Ex. Ag, Au
Ex. Cr, Mo
Ex. Pd: [Kr] 4d^10
