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Inorganic Chemistry > Atomic Structure > Flashcards

Atomic Structure Flashcards

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1
Q

John Dalton

A

In 1808, suggested that all matter was made up of tiny spheres that were able to bounce around with perfect elasticity and called them ATOMS

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2
Q

Mendelev and Meyer

A

In 1869, independently, they proposed the periodic table of the elements with its tendencies and components.

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3
Q

Balmer

A

In 1885, showed that energies of visible light emitted from hydrogen atom follow the equation:

E = Rh(1/2^2 - 1/nh^2)

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4
Q

Rh

A

Rydberg constant for hydrogen: 1.0097x10^7 1/m or 2.179x10-18 J

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5
Q

Joseph John Thompson

A

In 1898, found that atoms could sometimes eject a far smaller negative particle which he called an ELECTRON

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6
Q

What happened in 1904?

A

Thompson developed the Plum Pudding Model.

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7
Q

What is the Plum Pudding Model?

A

Developed by Thompson in 1904, this model stated the idea that an atom was made up of electrons scattered unevenly within an elastic sphere surrounded by a soup of positive charge to balance the electron’s charge, like plums surrounded by pudding.

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8
Q

Ernest Rutherford

A

In 1910, he proposed a more detailed model of the atom that included a central nucleus. He suggested that the positive charge was all in a central nucleus holding the electrons in place by electrical attraction.

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9
Q

Niels Bohr

A

In 1913, he refined Rutherford’s idea by adding that the electrons were in orbits. With each orbit only able to contain a set number of electrons.

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10
Q

What are the three major parts of an atom?

A

Proton, neutron, electron

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11
Q

T/F: Each atom has different amount of protons, neutrons and electrons.

A

True

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12
Q

Describe the proton (4)

A
  1. Positively charged particles found in the atomic nucleus
  2. Protons were discovered by Ernest Rutherford.
  3. Experiments done in the late 60’s and early 70’s showed that protons are made from other particles called quarks.
  4. Protons are made from two “up” quarks and one “down” quark.
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13
Q

Describe the neutron (4)

A
  1. Uncharged particles found in the atomic nucleus.
  2. Neutrons were discovered by James Chadwick in 1932.
  3. Experiments done in the late 60’s and early 70’s showed that neutrons are made from other particles called quarks.
  4. Neutrons are made from one “up” quark and two “down” quarks.
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14
Q

Describe the electron (4)

A
  1. Negatively charged particles that surround the atom’s nucleus.
  2. Electrons were discovered by J.J. Thompson in 1897.
  3. Electrons determine properties of the atom.
  4. Chemical reactions involve sharing or exchanging electrons.
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15
Q

Describe the nucleus

A
  1. Central part of an atom.
  2. Composed of protons and neutrons.
  3. Contains most of an atom’s mass.
  4. The nucleus was discovered by Ernest Rutherford in 1911.
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16
Q

Electron mass

A

1/1836 mass of H atom

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17
Q

Describe the quark (6)

A
  1. Believed to be one of the basic building blocks of matter.
  2. Quarks were first discovered in experiments done in the late 60’s and early 70’s.
  3. Three families of quarks are known to exist. Each family contains two quarks.
  4. The first family consists of up and down quarks (the quarks that join together to form protons and neutrons)
  5. The second family consists of strange and charm quarks (only exist at high energies).
  6. The third family consists of top and bottom quarks (only exist at very high energies).
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18
Q

Describe the isotope

A

atoms that have the same number of protons but different number of neutrons

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19
Q

Give an example of an isotope

A

Hydrogen has three isotopes.

  1. Protium: 1 proton
  2. Deuterium: 1 proton + 1 neutron
  3. Tritium: 1 proton + 2 neutrons
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20
Q

He²₄

A

The superscript (2): is the atomic number, the number of protons in an atom.

The subscript (4): is the atomic mass, the number of protons and neutrons in an atom.

*In neutral species (without charge) the number of protons = number of electrons.

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21
Q

The three quantum numbers

A

n, l and ml

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22
Q

The quantum number ‘n’

A

principal quantum number (number of the period)

n > 0

23
Q

The quantum number ‘l’

A

orbital angular momentum quantum number (s, p, d, f) goes from 0 to n-1. Specifies the shape of the sub-level.
l < n

24
Q

The quantum number ‘ml’

A

magnetic (azimuthal) quantum number (orientation)

|ml| ≤ l, it goes from -l to +l

25
Q

Energy levels formula

A

En = -Eo/n^2

26
Q

Radial wave functions

A

provides the distance and depends on quantic numbers ‘n’ and ‘l’

In the formula ‘ao’ is Bohr’s radius.

27
Q

Normalized spherical harmonics

A

angular function. Depends on two angles and the quantic numbers ‘l’ and ‘ml’

For ‘l’=0, the function is independent of the angle.

28
Q

Radial probability density

A 
is P(r) = r^2 |R(r)|^2
depends only on 'n' and 'l'. It gives information of distance. How far from the nucleus is the electron.
29
Q

Probability distribution functions

A

Only ‘s’ orbitals have the probability to be in the vicinity of the nucleus.

30
Q

Orbital Angular Momentum Quantum Number ‘l’

A

energy levels are degenerate with respect to l (the energy is independent of l)

31
Q

‘l’ values and letters

A 
l =           0   1     2     3     4     5
letter =   s    p    d     f      g     h
32
Q

Atomic states are usually referred to by their values of which quantum numbers?

A

‘n’ and ‘l’
Example:
2p: n = 2, l = 1

33
Q

What are the quantum numbers of a ‘2p’ state?

A 
n = 2 
l = 1
ml = -1, 0, +1 (3 orbitals with different orientations)
34
Q

Intrinsic spin

A

the spinning electron reacts similarly to the orbiting electron in a magnetic field. The magnetic spin quantum number: ‘ms’

35
Q

The magnetic spin quantum number: ‘ms’

A

Two values: -1/2 or +1/2

36
Q

ms = -½

A

Electron’s spin is down, against the applied magnetic field.

37
Q

ms = +½

A

Electron’s spin is up, in favor of the applied magnetic field.

38
Q

According to quantum mechanics, each electron is described by _____ quantum numbers.

A

four:

  1. Principal quantum number (n)
  2. Angular momentum quantum number (l): sublevel and shape of orbital
  3. Magnetic quantum number (ml): orientation of orbital
  4. Spin quantum number (ms)

The first three define the wave function for a particular electron. The fourth refers to the magnetic property of electrons.

39
Q

1s Orbital

A
  • n = 1, l = 0, ml = 0
  • sphere around the nucleus
  • The ‘1’ means the electron is in the orbital closest to the nucleus
  • The ‘s’ (l = 0) tells about the shape.

In this level, penetration occurs.

40
Q

2s Orbital

A
  • n = 2, l = 0, ml = 0
  • Similar to ‘1s’ except the electron is most likely in the region farther from the nucleus.
  • Has a spherical node
41
Q

p Orbitals

A
  • Appear in the second energy level and on.
  • Look like dumbbells (two lobes)
  • They can be oriented in three directions (Px, Py, Pz)
  • Has a nodal plane
42
Q

T/F: 2p orbitals and 3p orbitals have nodal planes, but 3p has also one spherical node.

A

True

43
Q

d Orbitals

A
  • Appear in the third energy level and on.
  • They have four lobes, except one.
  • They can be oriented in 5 directions (ways)
  • Three orbitals reside between the axis (dyz, dxz, dxy)
  • Two reside on the same axis (dx2-y2 and dz2)
  • They have two nodal planes.
44
Q

Nodal Surfaces: Spherical node [R(r)=0]

A

1s 0 2p 0 3d 0
2s 1 3p 1 4d 1
3s 2 4p 2 5d 2

45
Q

Nodal Surfaces: Radial nodes

A

n - l - 1

46
Q

Nodal Surfaces: Angular nodes [Y(θɸ)=0

A

s orbitals = 0
p orbitals = 1 plane
d orbitals = 2 planes

47
Q

Max number of electrons in an energy level

A

Energy level Max # of electrons
1 2
2 8
3 18
4 32
5 50

48
Q

Electron configuration

A

A detailed way of showing the order in which electrons fill in around the nucleus

49
Q

Electron configurations symbols:

1s²

A 
1 = energy level
s = sublevel (s, p, d, f) ('l' quantum number)
2 = # of electrons in the sublevel
50
Q

Three important rules while determining electron configuration

A
  1. Aufbau Principle
  2. Pauli Exclusion Principle
  3. Hund’s Rule
51
Q

Aufbau Principle

A
  • States that electrons occupy energy levels with lowest energy first.
  • For polielectronic atoms. This principle does not apply to hydrogen.
52
Q

Pauli Exclusion Principle

A

• If 2 electrons occupy the same energy level they must have opposite spins. ⇅

53
Q

Hund’s Rule

A
  • Maximum Multiplicity Principle.
  • Electrons that occupy orbitals of the same energy level will have the maximum number of electrons with the same spin.
  • Ex: 2p ↑↑↑

Inorganic Chemistry (5 decks)

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