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Bettel Plus Chang Chemistry > Quantum Theory and the Electronic Structure of Atoms > Flashcards

Quantum Theory and the Electronic Structure of Atoms Flashcards

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1
Q

Max Planck

A

-discovered that atoms and molecules emit energy only in certain discrete quantities (QUANTA)

-Quantum theory

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2
Q

Classical physics

A

-energy is continuous and that any amount of energy could be released in a radiation process

-assumed that atoms and molecules could emit or absorb any arbitrary amount of energy

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3
Q

Wave

A

-vibrating disturbance by which energy is transmitted
-periodic: wave form repeats itself at regular intervals

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4
Q

Wavelength, λ

A

-distance between identical points on successive waves

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5
Q

Frequence, f

A

-number of waves that pass through a particular point in one second

1 Hz = 1 cycle/s

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6
Q

Amplitude

A

-vertical distance from the midline of a wave to the peak or trough

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7
Q

Wave speed, v

A

-depends on the type of wave and the nature of the medium through which the eave is travelling

v = λ.f

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8
Q

James Clerk Maxwell (1873)

A

-visible light consists of electromagnetic waves

•Electromagnetic wave- has electric field component and magnetic field; these components have same frequency, wavelength, speed but travel in mutually perpendicular planes
-speed of electromagnetic waves: speed of light

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9
Q

Electromagnetic radiation

A

-emission and transmission of energy in the form of electromagnetic waves

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10
Q

Gamma rays

A

-electromagnetic radiation
-shortest wavelength; highest frequency

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11
Q

Radio waves

A

-electromagnetic radiation
-longest wavelength; lowest frequency

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12
Q

Visible light

A

-electromagnetic radiation
-wavelength: 400 nm (violet) - 700 nm (red)

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13
Q

⬆️ Frequency ⬆️ Energy

A
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14
Q

Solids are heated ➡️ emit electromagnetic radiation

A

Electric heater: dull red
Tungsten light bulb: bright white light

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15
Q

Quantum

A

-smallest quantity of energy that can be emitted/ absorbed in the form of electromagnetic radiation

-energy of single quantum:
E =h.f

-(Quantum theory): energy is always emitted in multiples of hf (i.e. hf, 2hf, 3hf but not 3.98hf, 1.67hf)

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16
Q

Albert Einstein (1905)

A

used quantum theory for:

•PHOTOELECTRIC EFFECT
-electrons are ejected from the surface of certain metals exposed to light of at least a certain minimum frequency (threshold frequency)
-no. of electrons ejected proportional to intensity (brightness) of light
-below threshold frequency, no electrons ejected no matter how intense the light
-electrons in metal are held in attractive forces; removing requires light of sufficiently high frequency (~high energy)

Photons- a beam of light is a stream of particles; each photon contains energy
E = h.f

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17
Q

Photoelectric effect formula

A

hf = KE + BE

KE: Kinetic energy of ejected electron
KE = 1/2 .me. v^2
me: mass of electron
BE: Binding energy of electron in metal

f < fthres : no electron ejected
f = fthres : only knock the electrons loose
f > fthres : electrons knock loose with kinetic energy

-more intense beam of light, larger number of protons: more electrons ejected
-higher frequency of light, greater kinetic energy of emitted electrons

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18
Q

Isaac Newton

A

-sunlight is composed of various color components that can be recombined to produce white light

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19
Q

Emission spectra

A

-either continuous or line spectra of radiation emitted by substance
-seen by energizing sample material with thermal energy or some other form of energy
-every element has a UNIQUE emission spectrum

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20
Q

Line spectra

A

-light emission only at specific wavelength
-emission spectra of atoms in gas phase: do not show continuous spread of wavelengths from red to violet

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21
Q

Niels Bohr (1913)

A

-emission spectrum of hydrogen atom
-each orbit has a particular energy associated with; energies associated with electron motion in the permitted orbits must be fixed in value or QUANTIZED
-emission of radiation by energized hydrogen atom: associated with electron dropping from higher-energy orbit to a lower one, giving up a quantum of energy (photon) in the form of light

En = -R(1/n^2)
En: energy of electron; n: principal quantum number (1,2,3,4 etc)

Convention: (-) energy of electron in atom is lower than energy of free electron (far from nucleus)
•Free electron: E∞=0
•Electron gets closer to nucleus (n decreases): En = more negative

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22
Q

Ground state/ ground level

A

n=1
-lowest energy state of a system

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23
Q

Excited state/ excited level

A

n = 2, 3, 4…
-higher energy than in ground state, less tightly held by nucleus

⬆️ n (farther from nucleus) ⬇️ stability

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24
Q

Rydberg’s formula

A

∆E = hf = R(1/ni^2 - 1/nf^2)

ni > nf : photon is emitted (-∆E: energy released)
ni < nf : (+∆E: energy absorbed)

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25
Q

Various series in atomic hydrogen emission spectrum

A

SERIES. nf. ni. Spectrum Region
Lyman. 1. 2,3,4… Ultraviolet
Balmer. 2. 3,4,5… Visible, UV
Paschen. 3. 4,5,6… Infrared
Brackett. 4. 5,6,7… Infrared

26
Q

Josef Fraunhofer

A

-studied emission spectrum of the sun and noticed certain dark lines at specific wavelength
-for atoms, emission and absorption of light occur at the same wavelength but differ in appearance.
°Colored lines: emission
°Dark lines: absorption
-matching the absorption lines with emission spectra of known elements: deduce the elements present in stars

27
Q

Pierre Janssen

A

-new dark line in solar emission spectrum: Helium

28
Q

William Ramsay

A

-discovery of Helium on Earth (in mineral of Uranium)
-only source of Helium: alpha particles emitted during nuclear decay

29
Q

Louis de Broglie

A

-if light waves can behave like a stream of particles (photons), particles such as electrons can possess wave properties
-electron bound to nucleus behave like standing/stationary wave
-greater frequency, shorter wavelength, greater number of nodes

•Node: amplitude of the wave at these points is zero
•Standing wave: generated by plucking (i.e. in guitar string)

-waves can behave like particles and particles can exhibit wavelike properties
λ = h/(m.v)
λ: wavelength
h: Planck’s constant
m: mass of moving particle
v: velocity

30
Q

Laser

A

-light amplification by stimulated emission of radiation
-special type of emission: involves atoms or molecules
-Ruby laser: first laser

31
Q

Clinton Davisson, Lester Germer, G.P. Thomson

A

-electrons do indeed possess wavelike properties

32
Q

Niels Bohr’s theory limitation

A
  1. Did not account for emission spectra of atoms containing more than 1 electron (i.e. Li and He)
  2. Did not explain why extra lines appear in hydrogen emission spectrum when magnetic field is applied
33
Q

Electron microscopy

A

-application of wavelike properties of electrons
-produces images of objects that cannot be seen with naked eye or light microscopes (law of optics: impossible to form image of an object that is smaller than half the wavelength of light used for observation)

34
Q

Werner Heisenberg

A

-Heisenberg uncertainty principle: it is impossible to know simultaneously both the momentum p (derined as mass times velocity) and the position of a particle with certainty

∆x∆p => h/(4π)

∆x∆p: uncertainty in mesuring position and momentum

-making measurement of momentum more precise (smaller ∆p), measurement of position of particles less precise (larger ∆x)
-electron does not orbit the nucleus in a well defined path (contradicts Bohr’s theory)

35
Q

Erwin Schrondiger (1926)

A

-formulated equarion that describes the behavior and energies of submicroscopic particles in general

-incorporates both particle behavior (mass) and wave behavior (wave function) which depends on the location in space of the system

-probability of finding electron in certain region of space is PROPORTIONAL to square of wave function

-start of QUANTUM MECHANICS field/ wave mechanics

36
Q

Electron density

A

-probability that an electron will be found in a particular region of an atom

37
Q

Atomic orbital

A

-wave function of an electron in an atom
-has characteristic energy and characteristic distribution of electron density

38
Q

Quantum numbers

A

-describe the distribution of electrons in hydrogen and other atoms:
•Principal quantum number
•Angular momentum quantum number
•Magnetic quantum number

°Spin quantum number
-describes the behavior of a specific electron and completes the description of electrons in atoms

39
Q

Principal quantum number, n

A

(n = 1,2,3…)
-larger n: greater average distance of an electron in the orbital from the nucleus; larger and less stable orbital

40
Q

Angular momentum quantum number, l

A

-Given value of n:
l = 0 to (n-1)

-shape of orbitals; Sharp, Diffuse, Principal, Fundamental … alpha arrange na

s orbital: l=0
p orbital: l=1
d orbital: l=2
f orbital: l=3
g orbital: l=4
h orbital: l=5

41
Q

Shell

A

-collection of orbitals with same value of n

42
Q

Subshell

A

-one or more orbitals with same n and l values
i.e. n=2; l=0,1
Subshells: 2s subshell and 2p subshell

43
Q

Magnetic quantum number, msub.l

A

-orientation of orbital in space
-depends on angular momentum quantum number:
number of msub.l = (2l + 1) orbitals
-l, (-l+1), …, 0, …, (+l+1), +l

•l=0, msub.l=0
•l=1, no. of msub.l=[(2x1)+1]=3
msub.l= -1, 0, 1

44
Q

Electron spin quantum, msub.s

A

-electrons act like tiny magnets
-counterclockwise spin: upward arrow
clockwise spin: downward arrow

msub.d = +1/2, -1/2

45
Q

Otto Stern and Walther Gerlach (1924)

A

-proof of electron spin:
beam of atoms directed through a magnetic field;
in a stream consisting of many atoms, there will be equal distribution of the two kinds of spin, two spots of equal intensity are detected on screen

46
Q

Atomic orbitals

A

-s subshell: 1 orbital (based on number of msub.l)
-p subshell: 3 orbitals
-d subshell: 5 orbitals

47
Q

S orbital

A

-spherical shape but differ in size
⬆️ Size ⬆️ Principal quantum number

48
Q

Boundary surface diagram

A

-encloses about 90% of the total electron density in an orbital
-electron density falls off rapidly as the distance from the nucleus increases

49
Q

P orbitals

A

-start: n=2, l=1, msub.l=-1,0,1
2px, 2py, 2pz
-each p orbital can be thought of as two lobes on opposite sides of nucleus
⬆️ Size ⬆️ Principal quantum number

50
Q

Energies of orbitals

A

-orbitals with same principal quantum number,n, have the same energy
⬆️Principal quantum number ⬆️ energy
-1s: closest to nucleus, most strongly held by nucleus, lowest energy

*Remember!!
Direction of electron spin has no effect on the energy of the electron

51
Q

Electron configuration

A

n, l, msub.l, msub.s: “address” of electrons

-order in which atomic subshells are filled in a many-electron atom

52
Q

Pauli Exclusion Principle
Wolfgang Pauli

A

-determine electron configuration
-no two electrons in an atom can have the same four quantum numbers
-only two electrons may occupy same atomic orbital (each with opposite spin)

53
Q

Diamagnetic substance ⬆️⬇️

A

-electron spins are paired/ ANTIPARALLEL to each other; magnetic effects cancel out
-slightly repelled by magnet
-general rule: atoms with EVEN NUMBER OF ELECTRONS- diamagnetic

54
Q

Paramagnetic substance ⬆️⬆️

A

-electrons have same/ PARALLEL spins, their net magnetic fields would reinforce each other
-general rule: atoms with ODD NUMBER OF ELECTRONS- paramagnetic

55
Q

Measurement of magnetic properties

A

-most direct evidence for specific electron configuration of elements

56
Q

Shielding effect in many-electron atoms

A

-i.e. 2s or 2p electron are partly shielded from attractive forces of nucleus by 1s electrons
⬆️Shielding effect ⬇️Electrostatic attraction
-⬇️Penetrating power ⬆️Angular momentum quantum number
s>p>d>f
⬇️Shielding effect ⬆️Penetrating power

57
Q

Hund’s rule
Frederick Hund

A

-most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins

(A) |⬆️⬇️| | |
(B) |⬆️ |⬇️ | |
(C) |⬆️ |⬆️ | |

Stability A < B < C

58
Q

Determination of maximum number of electrons that an atom can have

A

2n^2

59
Q

Building-Up Principle/
Aufbau Principle

A

-protons are added one by one to the nucleus to build up the elements, electrons are similarly added to the atomic orbitals

60
Q

Noble gas core

A

-shows in brackets the noble gas element that most nesrly precedes the element being considered

61
Q

Transition metals

A

-either gave incompletely filled d susbhells or readily give rise cations that have incompletely filled d subshells

62
Q

Lanthanides/ Rare Earth Series

A

-have incompletely filled 4f subshells or readily give rise to cations that have incompletely filled 4f subshells

Bettel Plus Chang Chemistry (19 decks)

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