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Chem - Gr11 > Structure of Atom > Flashcards

Structure of Atom Flashcards

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1
Q

Electromagnetic Radiation

A

Charged particles under acceleration producing alternating electrical and magnetic fields

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2
Q

Isotopes vs Isobars vs Isotones

A

Isotopes: Atoms of the same element which have the same atomic numbers but different mass numbers
Isobars: Atoms of different elements having the same mass numbers but different atomic numbers
Isotones: Atoms having same number of neutrons but different mass numbers

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3
Q

Relation between c, mu, lambda

A

c = (mu)(lambda)

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4
Q

Wavenumber

A

Number of wavelengths per unit length = 1 / lambda

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5
Q

Give the electromagnetic spectrum (in increasing wavelength)

A

gamma rays < X rays < UV < visible < IR < Microwave < FM < AM < long radio

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6
Q

Give the electromagnetic spectrum (in increasing frequency)

A

Long radio > AM > FM > Microwave > IR > visible > UV > X rays > gamma rays

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7
Q

Black Body Radiation Explanation

A
  • When solids are heated they emit radiation over a wide range of wavelengths
  • When an iron rod is heated, it turns to dull red and then progressively more and more red as temp increases
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8
Q

What is black body radiation?

A

The ideal black body, which emits and absorbs radiations of all frequencies, is called a black body and the radiation emitted by such a body is black body radiation

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9
Q

Relation between E, h, mu

A

E = h (mu)

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10
Q

Photoelectric effect + results

A

Electrons are ejected from certain metals when exposed to a beam of light
* Ejected as soon as light strikes the surface
* No. of electrons proportional to intensity
* Existence of minimum frequency

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11
Q

Relation between intensity and no. of electrons ejected

A

Directly proportional

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12
Q

Threshold frequency

A

Frequency below which photoelectric effect is not observed

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13
Q

Relation between frequency and KE

A

Directly proportional (Greater energy by photon = greater KE - greater frequency of photon)

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14
Q

Photoelectric effect equation

A

h(mu) = h(mu0) +1/2 mv^2

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15
Q

Longest wavelength

A

Red colour

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16
Q

Spectrum

A

Ray of white light is spread out into a series of coloured bands

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17
Q

Emission spectrum

A

Spectrum of radiation emitted by a substance that has absorbed energy

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18
Q

Absorption spectrum

A

When light is passed through unexcited atomic hydrogen and transmitted light is lacking in intensity

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19
Q

Wavenumber formula (Rydberg’s)

A

(mu)_ = 109,677 (1/n1^2 - 1/n2^2) cm^-1

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20
Q

Angular momentum formula

A

mvr = nh/2pi

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21
Q

All light spectrum correspondance

A

Lyman - Ultraviolet (Till n=1)
Balmer - Visible (Till n = 2)
Paschen - Infrared (Till n=3)

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22
Q

Formula of radius of bohr orbit

A

r = a(n^2)/Z
a = 52.9 pm

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23
Q

Convert:
1) picometer
2) nanometer
3) Armstrong
4) Micrometer

A

1) 10^-12m
2) 10^-9m
3) 10^-10m
4) 10^-6m

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24
Q

Ionized hydrogen atom

A

When an electron is taken free from the influence of nucleus, the energy is taken as zero. The electron in this situation is associated with the stationary state of Principal Quantum Number = n = infinity

25
Q

Formula of E {Rydberg’s}

A

E = -Rh(Z^2/n^2)
Rh = 2.18 * 10^-18J

26
Q

Magnitude of velocity of electron

A

Increases with positive charge on nucleus
Decreases with increase in principal quantum number

27
Q

For which species can Bohr’s theory by applied

A
  • Isoelectronic species having one electron
  • H, He+, Li2+, Be3+, etc.
28
Q

What happens when n1 > nf

A

Delta E is negative and energy is released

29
Q

What does the brightness / intensity of spectral depend upon?

A

Number of photons of same wavelength or frequency absorbed or emitted

30
Q

Relation between wavelength and momentum (dual behavior)

A

lambda = h / mv = h / p

31
Q

Heisenberg Uncertainty Principle

A

Impossible to determine simultaneously, the exact position and exact momentum of an electron

32
Q

Heisenberg Uncertainty Principle Formula

A

delta x * delta mv >= h / 4 pi

33
Q

Principal quantum number

A
  • Determines size and to large extent the energy of the orbit
  • Identifies the shell
  • Ranges from 1 to n
34
Q

Number of allowed orbitals

A

n^2

35
Q

With increase in ‘n’ value what increases?

A
  • Energy
  • Size of orbital
36
Q

Azimuthal quantum number

A

Defines the 3D shape of the orbital

37
Q

What do ‘l’ values range from

A

0 to (n-1)

38
Q

How many subshells are there?

A

n total

39
Q

Magnetic orbital quantum number

A

Gives information about spatial orientation of the orbit with respect to standard set of co-ordinate axis

40
Q

How many possible values of magnetic quantum number

A

2l + 1

41
Q

Electron spin

A

Can have two orientations relative to chosen axis {Spin of the electron}

42
Q

nodes

A

Refers to the area / region where probability density function reduces to zero

43
Q

Nodes:
1) Total
2) Radial
3) Angular

A

1) n-1
2) n - l -1
3) l

44
Q

Draw 2px, 2py, 2pz

A

-

45
Q

Draw dxy, dyz, dxz, d(x^2-y^2), dz^2

A

-

46
Q

Degenerate

A

Orbitals having the same energy

47
Q

Effective nuclear charge

A

The attractive positive charge of nuclear protons acting on valence electrons

48
Q

Formula of Zeff

A

Zeff = Z - S
Z = atomic number
S = number of shielding electrons

49
Q

Relation between l and Zeff and why?

A

Inversely proportional
Increase in azimuthal causes s electron to be more tightly bound to the nucleus than p electron which in turn will be more tightly bound than the d electron

50
Q

Energy of orbit formula (using n and l values)

A

n+l

51
Q

Aufbau Principle

A

In the ground state of the atoms, the orbitals are filled in order of increasing energies

52
Q

Pauli Exclusion Principle

A
  • No two electrons in an atom can have the same set of four quantum numbers
  • Only two electrons may exist in the same orbital and these electrons must have opposite spin
53
Q

Maximum number of electrons in shell with principal quantum number n

A

2n^2

54
Q

Hund’s rule of maximum multiplicity

A

Pairing of electrons in the orbitals belonging to the same subshell does not take place until each orbital belonging to that subshell has got one electron each

55
Q

Valence Electrons

A

Electrons in the completely filled shells are known as core electrons and electrons added to electronic shell with highest principal quantum number

56
Q

Config of copper and chromium

A

Chromium: 1s2 2s2 2p6 3s2 3p6 4s1 3d5
Copper: 1s2 2s2 2p6 3s2 3p6 4s1 3d10

57
Q

Explain copper and chromium’s electronic config

A
  • Fully filled and half-filled orbitals have extra stability
  • Hence, they take on that config
58
Q

Causes of stability of completely filled and half filled sub-shells

A
  • Symmetrical distribution of electrons
  • Exchange energy - Maximum and hence max stability
59
Q

Fe2+ configuration

A

1s2 2s2 2p6 3s2 3p6 3d6
We have to remove from the valence electron shell first

Chem - Gr11 (8 decks)

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