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Wavelength and Electrons!

Wavelength and Electrons!

Wavelength ( λ) - the distances from peak to peak (nm, pm, cm, m) 

Frequency (v) - the number of peaks per second (/s, s-, Hz)

Light - form of energy also known as electromagnetic radiation that travels by waves and particles

Amplitude - the height of the wave from the origin to peak or crest

Brightness or intensity - amplitude adjusts this and it's not the type of light, it is how much light

Vacuum - no matter exists (empty space)

Quantum mechanics - description of the motion and interaction of subatomic particles

Particle - exists in a single place at a particular moment in time, probability of finding a particle is 100% in a particular position and 0% everywhere else

Wave - we know its pattern (wavelength) but like a ripple, it's got a high probability of being in many places

Momentum - mass * velocity (speed + direction) 

We only see a small section of the electromagnetic spectrum (visible light).


C = λv:

  • Constant (c) = 3.00 * 10^8 m/sec
  • Wavelength and frequency have an inverse relationship 

E = hv:

  • Constant (h) = planck's constant or 6.626 * 10^-34 J*s
  • Energy and frequency have a direct relationship. Longer wavelengths have lower frequencies and lower energies. Shorter wavelengths have higher frequencies and higher energies.
  • Violet has a shorter wavelength, higher frequency, and more energy. Red has a longer wavelength, lower frequency, and less energy. As frequency increases, energy increases, and wavelength decreases. All wavelengths have the same speed.


De Broglie's Atom:

  • The larger the radius, the higher the electron energy.
  • The larger the transition between energy levels, the higher the photon energy.
  • The electron can exist in only discrete energies; it is quantized. If the electron could exist at any energy level, we would see the entire visible spectrum.
  • In De Broglie's model of the atom, the electron can exist as a standing wave at specific energies.
  • The electron can only absorb the exact amount of energy that will allow it to move from one standing wave to another.
  • If the electron absorbs energy that’s not enough to get it to another standing wave, the electron would not be in phase with itself and would interfere with itself.


Heisenberg Uncertainty Principle:

  • Large objects that have a high momentum have a very short wavelength.
  • Fast objects, heavy objects have a lot of momentum even if they’re moving slowly.
  • Small objects, like atoms and electrons, have wavelengths we can measure so we can know its momentum.
  • The small object doesn’t have a position.
  • If we have a pure wave, we can measure its wavelength and thus its momentum but we don’t know its pattern.
  • Although we know the particle’s position we don’t know its momentum.
  • To get a particle’s wavelength, we superimpose its different wavelengths and see where the peaks overlap.
  • We can then find an area with a clear wavelength in one small section or area.
  • Thus, we now have a quantum object, something that has both wavelengths and particle nature.


Interference:

  • Waves including electromagnetic waves interact with each other in a characteristic way called interference.
  • They either cancel each other out or build each other up.
  • Ex. if two waves of equal amplitude are in phase when they interact (when they align with overlapping crests) a wave with twice the amplitude results, which is called constructive interference.
  • If two waves are completely out of phase when they interact (when they align so that the crest from one overlaps with the trough of the other, the waves cancel by destructive interference.
  • Diffraction: when a wave encounters an obstacle or a slit that is comparable in size to its wavelength, it bends or diffracts around it.
  • Interference from two slits: when a beam of light passes through two small slits, the two resulting waves interfere with each other. Whether the interference is constructive or destructive at any given point depends on the path lengths traveled by the waves. The resulting interference pattern appears as a series of bright and dark lines on a screen.


The Particle Nature of Light:

  • Prior to early 1900s and especially after the discovery of the diffraction of light, light was thought to be purely a wave phenomenon. However, discoveries like the photoelectric effect brought a particle nature to light.
  • We thought that energy transferred from the light to an electron in the metal resulted in the dislodgement of electrons. If this was correct, the amplitude of the light would eject electrons.
  • Planck suggested hot objects do not emit EM energy continuously (as would be expected if energy were in form of waves)
  • Instead, Planck suggested objects emit energy in small, specific amounts called quanta.
  • Quantum is the minimum quantity of energy that can be lost or gained by an atom.
  • Planck proposed the following relationship between a quantum of energy and the frequency of radiation. E=hv
  • 1905: Albert Einstein's radical idea: EM radiation has a dual wave-particle nature; while light exhibits many wave-like properties It can also be thought of as a stream of particles.
  • Each particle carries a quantum of energy. Einstein calls these particles photons.
  • A photon is a particle of EM radiation having zero mass and carrying a quantum of energy.
  • The energy of a particular proton depends on the frequency of radiation.
  • Einstein explains that the photo electric effect by proposing that the EM radiation is absorbed by matter only in whole numbers of photons.
  • In order for an electron to be ejected from a metal surface the electrons must be struck by a single photon possessing at least the minimum energy required to knock the electronics. According to the E equals HV the minimum energy corresponds to the minimum frequency. If the frequencies of the photons are below the minimum the electrons remain bound to the metal surface.
  • Electrons and different metals are bound more or less tightly so different metals require different minimum frequencies to exhibit the photoelectric effect.


Wave Nature for Particles:

  • Just as a photo electric effect suggests park on nature of lights certain observations about the islands began to suggest wave nature for particles.
  • Atomic spectra: when an atom absorbs energy in the form of heat, light, or electricity, it often reemits that Energy as late. For example one in electric current is Pastor a tube of neon gas in them it’s a characteristic red light. It shows that the light emitted contains several different wavelengths. We can separate the light committed by a single element into its constituent wave links by passing through a prism.
  • Light is both a wave and a particle.


Democritus:

  • “atmos” - uncuttable

Aristotle:

  • earth, fire, water, air

John Dalton:

  • Atoms is like little indivisible balls, they arrange themselves in different combinations and make different compounds.

JJ Thomson:

  • “plum pudding model”, atoms are not exactly indivisible, they are made of smaller things.

Robert Millikan:

  • oil drop experiment

Earnest Rutherford:

  • most of the atom's mass and it’s entire positive charge I can find in a small core called the nucleus.
  • The positively charged particles are called protons.
  • most of the volume of an atom is empty space.
  • The number of negatively charged electrons dispersed outside the nucleus is the same number as positively charged protons in the nucleus. It explains the overall neutrality of the atom.

Bohr Model:

  • was inadequate, it couldn’t explain the spectrum for any atom except hydrogen.
  • It did not explain chemical behavior.
  • More electrons in atoms —> interfere with each other.

De Broglie:

  • electrons behave as waves, important for only very small objects.

Schrodinger:

  • applies with equations to electrons and atoms.
  • Well-defined energy, probability distribution in space.
  • Probability of where an electron might be.
  • Probability cloud is a wave function and tells us the probability of finding a moving particle at that time and at that point.
  • Atomic orbital is 90% of the electron in the sphere closest to the nucleus but energy is well defined.

James Chadwick:

  • discovery of the neutron


Orbitals:

  • probability regions where electrons are found (Bohr's model)
  • As n increases, orbitals get larger and higher energy.
  • Each n gets one new sub-level.
  • Each energy level and contains N squared total orbitals.
  • Each orbital holds a maximum of two electrons.
  • First energy level has one sub-level and one orbital. S orbital is a spherical shape and has one S orbital in every energy level.
  • Second energy level has two sub levels and four total orbitals. P orbital is paired shaped or peanut-shaped and has 3P orbitals starting at n = 2.
  • Sir energy level has reserve levels and nine total electrons. The D orbital is clover-shaped and has five orbitals and every energy level starting with n equals three. Each new sub-level adds two more orbitals.
  • Forth energy level has four sub-levels and 16 total orbitals. F orbital is flower-shaped and has 7 F orbitals starting with n equals four.
  • Orbitals overlap in space, but they were made of stinking energy and contain different electrons. Note that 2S orbital is larger than 1S and at higher energy but has the same shape. All orbitals exist but only use those needed to hold given amounts of electrons.
  • Electron configuration is the address of the electron.


Aufbau Principle:

  • each electron occupies the lowest energy level.
  • All orbitals within an energy level or at the same, equal energy.

Hund’s Rule:

  • Single electrons with the same spend most occupy equal energy sub orbitals before pairing with opposite spin. In other words first arrow fills up all boxes before the second.

Pauli’s Exclusion Principle:

  • A max of two electrons may occupy a single orbital but only if the electrons have opposite spins. In other words you can’t have two up arrows.

Valance Electrons:

  • The atom's highest energy level electrons.


Wavelength and Electrons!

Wavelength ( λ) - the distances from peak to peak (nm, pm, cm, m) 

Frequency (v) - the number of peaks per second (/s, s-, Hz)

Light - form of energy also known as electromagnetic radiation that travels by waves and particles

Amplitude - the height of the wave from the origin to peak or crest

Brightness or intensity - amplitude adjusts this and it's not the type of light, it is how much light

Vacuum - no matter exists (empty space)

Quantum mechanics - description of the motion and interaction of subatomic particles

Particle - exists in a single place at a particular moment in time, probability of finding a particle is 100% in a particular position and 0% everywhere else

Wave - we know its pattern (wavelength) but like a ripple, it's got a high probability of being in many places

Momentum - mass * velocity (speed + direction) 

We only see a small section of the electromagnetic spectrum (visible light).


C = λv:

  • Constant (c) = 3.00 * 10^8 m/sec
  • Wavelength and frequency have an inverse relationship 

E = hv:

  • Constant (h) = planck's constant or 6.626 * 10^-34 J*s
  • Energy and frequency have a direct relationship. Longer wavelengths have lower frequencies and lower energies. Shorter wavelengths have higher frequencies and higher energies.
  • Violet has a shorter wavelength, higher frequency, and more energy. Red has a longer wavelength, lower frequency, and less energy. As frequency increases, energy increases, and wavelength decreases. All wavelengths have the same speed.


De Broglie's Atom:

  • The larger the radius, the higher the electron energy.
  • The larger the transition between energy levels, the higher the photon energy.
  • The electron can exist in only discrete energies; it is quantized. If the electron could exist at any energy level, we would see the entire visible spectrum.
  • In De Broglie's model of the atom, the electron can exist as a standing wave at specific energies.
  • The electron can only absorb the exact amount of energy that will allow it to move from one standing wave to another.
  • If the electron absorbs energy that’s not enough to get it to another standing wave, the electron would not be in phase with itself and would interfere with itself.


Heisenberg Uncertainty Principle:

  • Large objects that have a high momentum have a very short wavelength.
  • Fast objects, heavy objects have a lot of momentum even if they’re moving slowly.
  • Small objects, like atoms and electrons, have wavelengths we can measure so we can know its momentum.
  • The small object doesn’t have a position.
  • If we have a pure wave, we can measure its wavelength and thus its momentum but we don’t know its pattern.
  • Although we know the particle’s position we don’t know its momentum.
  • To get a particle’s wavelength, we superimpose its different wavelengths and see where the peaks overlap.
  • We can then find an area with a clear wavelength in one small section or area.
  • Thus, we now have a quantum object, something that has both wavelengths and particle nature.


Interference:

  • Waves including electromagnetic waves interact with each other in a characteristic way called interference.
  • They either cancel each other out or build each other up.
  • Ex. if two waves of equal amplitude are in phase when they interact (when they align with overlapping crests) a wave with twice the amplitude results, which is called constructive interference.
  • If two waves are completely out of phase when they interact (when they align so that the crest from one overlaps with the trough of the other, the waves cancel by destructive interference.
  • Diffraction: when a wave encounters an obstacle or a slit that is comparable in size to its wavelength, it bends or diffracts around it.
  • Interference from two slits: when a beam of light passes through two small slits, the two resulting waves interfere with each other. Whether the interference is constructive or destructive at any given point depends on the path lengths traveled by the waves. The resulting interference pattern appears as a series of bright and dark lines on a screen.


The Particle Nature of Light:

  • Prior to early 1900s and especially after the discovery of the diffraction of light, light was thought to be purely a wave phenomenon. However, discoveries like the photoelectric effect brought a particle nature to light.
  • We thought that energy transferred from the light to an electron in the metal resulted in the dislodgement of electrons. If this was correct, the amplitude of the light would eject electrons.
  • Planck suggested hot objects do not emit EM energy continuously (as would be expected if energy were in form of waves)
  • Instead, Planck suggested objects emit energy in small, specific amounts called quanta.
  • Quantum is the minimum quantity of energy that can be lost or gained by an atom.
  • Planck proposed the following relationship between a quantum of energy and the frequency of radiation. E=hv
  • 1905: Albert Einstein's radical idea: EM radiation has a dual wave-particle nature; while light exhibits many wave-like properties It can also be thought of as a stream of particles.
  • Each particle carries a quantum of energy. Einstein calls these particles photons.
  • A photon is a particle of EM radiation having zero mass and carrying a quantum of energy.
  • The energy of a particular proton depends on the frequency of radiation.
  • Einstein explains that the photo electric effect by proposing that the EM radiation is absorbed by matter only in whole numbers of photons.
  • In order for an electron to be ejected from a metal surface the electrons must be struck by a single photon possessing at least the minimum energy required to knock the electronics. According to the E equals HV the minimum energy corresponds to the minimum frequency. If the frequencies of the photons are below the minimum the electrons remain bound to the metal surface.
  • Electrons and different metals are bound more or less tightly so different metals require different minimum frequencies to exhibit the photoelectric effect.


Wave Nature for Particles:

  • Just as a photo electric effect suggests park on nature of lights certain observations about the islands began to suggest wave nature for particles.
  • Atomic spectra: when an atom absorbs energy in the form of heat, light, or electricity, it often reemits that Energy as late. For example one in electric current is Pastor a tube of neon gas in them it’s a characteristic red light. It shows that the light emitted contains several different wavelengths. We can separate the light committed by a single element into its constituent wave links by passing through a prism.
  • Light is both a wave and a particle.


Democritus:

  • “atmos” - uncuttable

Aristotle:

  • earth, fire, water, air

John Dalton:

  • Atoms is like little indivisible balls, they arrange themselves in different combinations and make different compounds.

JJ Thomson:

  • “plum pudding model”, atoms are not exactly indivisible, they are made of smaller things.

Robert Millikan:

  • oil drop experiment

Earnest Rutherford:

  • most of the atom's mass and it’s entire positive charge I can find in a small core called the nucleus.
  • The positively charged particles are called protons.
  • most of the volume of an atom is empty space.
  • The number of negatively charged electrons dispersed outside the nucleus is the same number as positively charged protons in the nucleus. It explains the overall neutrality of the atom.

Bohr Model:

  • was inadequate, it couldn’t explain the spectrum for any atom except hydrogen.
  • It did not explain chemical behavior.
  • More electrons in atoms —> interfere with each other.

De Broglie:

  • electrons behave as waves, important for only very small objects.

Schrodinger:

  • applies with equations to electrons and atoms.
  • Well-defined energy, probability distribution in space.
  • Probability of where an electron might be.
  • Probability cloud is a wave function and tells us the probability of finding a moving particle at that time and at that point.
  • Atomic orbital is 90% of the electron in the sphere closest to the nucleus but energy is well defined.

James Chadwick:

  • discovery of the neutron


Orbitals:

  • probability regions where electrons are found (Bohr's model)
  • As n increases, orbitals get larger and higher energy.
  • Each n gets one new sub-level.
  • Each energy level and contains N squared total orbitals.
  • Each orbital holds a maximum of two electrons.
  • First energy level has one sub-level and one orbital. S orbital is a spherical shape and has one S orbital in every energy level.
  • Second energy level has two sub levels and four total orbitals. P orbital is paired shaped or peanut-shaped and has 3P orbitals starting at n = 2.
  • Sir energy level has reserve levels and nine total electrons. The D orbital is clover-shaped and has five orbitals and every energy level starting with n equals three. Each new sub-level adds two more orbitals.
  • Forth energy level has four sub-levels and 16 total orbitals. F orbital is flower-shaped and has 7 F orbitals starting with n equals four.
  • Orbitals overlap in space, but they were made of stinking energy and contain different electrons. Note that 2S orbital is larger than 1S and at higher energy but has the same shape. All orbitals exist but only use those needed to hold given amounts of electrons.
  • Electron configuration is the address of the electron.


Aufbau Principle:

  • each electron occupies the lowest energy level.
  • All orbitals within an energy level or at the same, equal energy.

Hund’s Rule:

  • Single electrons with the same spend most occupy equal energy sub orbitals before pairing with opposite spin. In other words first arrow fills up all boxes before the second.

Pauli’s Exclusion Principle:

  • A max of two electrons may occupy a single orbital but only if the electrons have opposite spins. In other words you can’t have two up arrows.

Valance Electrons:

  • The atom's highest energy level electrons.