Moles / Light Flashcards

1
Q

1 mole

A
  • 6.022 x 1023 atoms
  • As many particles as there are in 12.0 g of 12C
  • Avagadro’s number
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2
Q

Molar conversions

A
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3
Q

A mole of a compound (covalent, ionic)

A
  • Covalent: Avogadro’s number of molecules
    • 1 mole CO2 = 6.022 x 1023 CO2 molecules
  • Ionic: Avogadro’s number of formula units
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4
Q

How many moles of C, H, O in C6H12O2?

A
  • C) 6 moles
  • H) 12 moles
  • O) 6 moles
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5
Q

How to find molar mass

A
  • Add for each element:
    • Number of moles in the element * atomic mass (from periodic table, g/mole)
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6
Q

Electromagnetic radiation

A
  • Generated by moving electrons
  • Can transport energy without a medium (unlike sound which requires matter to move through)
  • Displays wave properties
    • Diffraction
    • Interference
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7
Q

Electromagnetic spectrum

A
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8
Q

Features of waves

A
  • Wavelength: the distance between two identical points on two asjacent waves
    • Symbol: lambda λ
    • Unit: m, nm
      • 109 nm in m
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9
Q

Frequency

A
  • The amount of crests that pass a point in one second
  • Symbol: f, nu, ν
  • Unit: s-1, Hz
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10
Q

Relationship between wavelength and frequency

A
  • All EM waves travel at speed of light
  • c = 3.0 x 108 m/s
  • c = λv
  • Long wavelength = low frequency
  • Short wavelength = high frequency
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11
Q

Max Planck

A
  • Solids emit radiation as heated
  • Study relationships between wavelength and intensity of radiation emitted and temperature
  • Energy is quantizes
  • Lower frequency = intense light
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12
Q

Quantization of energy

A
  • Energy is released or absorbed in quanta
  • Energy radiation is directly proportional to frequency (specific allowable energies)
  • Quantized - values are restricted to certain quantities
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13
Q

Energy equation (Planck’s constant)

A
  • E = hv
  • h = Planck’s constant = 6.626 x 10-34 J S
  • v = frequency
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14
Q

Wavelength vs energy vs frequency (red vs violet)

A
  • RED: Large wavelength = low frequency = low energy
  • VIOLET: Small wavelength = high frequency = high energy
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15
Q

Photoelectric effect

A
  • When light is shone on a metal, electrons are ejected from the metal
  • The ejected electrons have a ceratin amount of kinetic energy
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16
Q

Classical physics prediction

A
  • Higher amplitude (brighter) = more energy
  • Bright red = eject electrons, dim red = no electrons
  • Bright blue = eject electrons, dim blue = no electrons
  • Brighter light = higher energy of electrons popped off
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17
Q

Blue vs red light actually observed

A
  • Bright red = no electrons, dim red = no electrons
  • Bright blue = eject electrons, dim blue = eject electrons (but fewer)
  • Suggests that the ability of light to eject electrons depends on its frequency, not brightness
  • Planck - energy and frequency directly related
  • Einstein uses this to explain the photoelectric effect
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18
Q

Einstein explains photoelectric effect

A
  • Light as particles, each with its own quantum (specific amount) of energy, depending on its frequency
  • Particles called photons
  • Only those photons with the minimum amount of energy needed to eject electrons from metal will do so
19
Q

Photon

A
  • A particle of electromagnetic radiation having zero mass and carrying a quantum of energy
  • Light has a dual nature
20
Q

Classical view of atom

A
  • An electron traveled about the nucleus in an orbit
  • Any orbit should be possible for any electron
  • Did not know why the electrons did not fall into the nucleus
21
Q

Neils Bhor

A
  • Built a simple model of an atom
  • Based on an understanding of the sharp line emission spectrum of excited atoms
22
Q

Line emission spectra

A
  • When atoms or molecules absorb energy, that energy is often released as light energy (fireworks, etc)
  • When that light is passed through a prism. a pattern is seen that is unique to that type of atom or molecule — pattern called an emission spectrum
    • Non-continuous
    • Can be used to identify the material
      • Flame tests
23
Q

Line emission spectra of excited atoms

A
  • Excited atoms emit light of only certain wavelengths
  • The wavelengths of emitted light depend on the element
24
Q

Ground state

A

Electron in lowest energy level possible

25
Atomic spectra and Bhor
* e- can only exist in certain **discrete orbits** of specific energy (quantized energy states) * e- have a **specific** amount of energy that is equal to the PE of their location * PE of e- is determined by its distance from the nucleus * Closer to nucleus = lower PE * Each type of photon has an energy equal to the energy lost by specific jumps in energy levels * Nobel prize 1922 * **Problem**: only works for hydrogen --\> quantum/wave mechanics
26
Balmer series
* The series of visible lines in the hydrogen atom spectrum
27
L de Broglie (+ equations)
* Proposed (1924) that all moving objects have wave properties * For light: E = mc2 * E = hν = hc / λ * mc = h / λ * For particles: c = velocity * (mass)(velocity) = h / λ * mv = h / λ * λ = h / mv
28
Quantum / wave mechanics (+ equation)
* Using deBroglie's equation **λ = h/mv** can calulcate the wavelength for moving objects * λ = wavlength (meters) * v = velocity (m/sec) * m = mass (kg) * h = Planck's constant
29
Uncertainty principle
* Problem of defining nature of electrons in atoms * Solved by W. **Heisenberg** * Cannot simultaneously define the **position** and **momentum** (m\*v) of an electron * We define e- energy exactly but accept limitation that we do not know exact position * m = kg * v = m/s
30
Double slit experiment
* Demonstration that light and matter can display characteristics of both classically defined waves and particles * Electrons act like waves and particles * Interference pattern
31
E Schrodinger
* Applied idea of e- behaving as a wave to the problem of electrons in atoms * Developed the **wave equation** * Solution gives set of math expressions called **wave functions** * Each wave function describes an allowed energy state of an e- which corresponds to an **_orbital_** - the region of space within which an electron is found * Area of probability, not exact location
32
Orbital
* For an electron with a given energy, can describe a region in the atom of high probability of finding it * Many of the properties of atoms are related to the energies of the electrons
33
s orbital
* 1 direction spherical
34
p orbital
* 3 directions * Three p orbitals lie 90 degrees apart in space
35
d orbital
* 5 directions
36
f orbital
* 7 directions
37
g orbital
* 9 directions * Hypothetical
38
Principal quantum number
* Symbol: n * The main energy level occupied by the electron * Values of n are positive integers and the value of n=1 has the lowest energy
39
Sublevels and relative energy
* 4 sublevels for the elements discovered to date * Sublevel indicates the shape of the orbital * First energy level, n = 1 has 1 sublevel * Fourth energy level, n = 4 has 4 sublevels * Levels split because negative e- next to each other * Sublevels broken up into orbitals
40
The Aufbau Rule
* The **lowest** energy levels are filled first
41
The Pauli Exclusion Principle
* Orbitals can have **max of 2 electrons** * e- in same orbital must have **opposite** spins
42
Hund's Rule
* When orbitals of identical energy (same shape) are available, e- enter those orbitals singly before any spin pairing takes place
43
Energy emitted by e- energy level change
E in J
44
Wavelength emitted by e- energy level change