Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

LC Chemistry > Electronic Structure of Atoms > Flashcards

Electronic Structure of Atoms Flashcards

You may prefer our related Brainscape-certified flashcards:
AP® Chemistry flashcards Periodic Table flashcards Chemistry 101 flashcards Organic Chemistry 101 flashcards
1
Q

Energy Level

A

Fixed energy value that an electron in an atom may have

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

Ground State

A

Lowest energy state (in 1s orbital)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

Excited State

A

Higher energy state

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Sublevel

A

Subdivision of a main energy level and consists of one or more orbitals of the same energy

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Orbital

A

A region in space within which there is a high probability of finding an electron

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Heisenberg’s Uncertainty Principle

A

It is not possible to ascertain both the position and the momentum of an electron in an atom simultaneously

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

The atomic radius of an atom

A

Defined as: Half the distance between the nuclei of two atoms of the same element that are joined together by a single covalent bond

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

The first ionisation energy of an atom

A

Minimum energy required to completely remove the most loosely bound electron from one mole of neutral gaseous atom in the ground sate

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Aufbau Principle

A

When building up the electronic configuration of an atom in its ground state, the electrons may occupy the lowest available energy level

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Hund’s Rule of Maximum

A

States that when two or more orbitals of equal energy are available, the electrons occupy then singly first before filling them in pairs

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Pauli Exclusion Principle

A

No more than two electrons may occupy am orbital and they must have opposite spins

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Niel’s Bohr

A
  • Investigated location of electrons
  • Looked at emission line spectrum (of elements)
  • White light could be split to produce continuous spectrum
  • Repeated using light from hydrogen discharge bulb
  • No longer see continuous spectrum but narrow lines of emission spectrum
  • Repeated w/ other elements
  • Different line spectrum to hydrogen, but same as itself when repeated
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Bohr’s theory

A
  1. Electrons revolve around nucleus in fixed paths called energy levels
  2. Energy levels are represented by the letter n (lowest: n = 1, then n=2, etc)
  3. Each energy level has fixed amount of energy
  4. An electron in an energy level possess the same amount of energy as the energy level
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Electrons in excited state

A
  1. Electrons normally exist in ground state (occupy lowest available energy level)
  2. If electron is given energy (from electricity/heat) they can jump from lower energy levels to higher ones
  3. When this happens, electron is said to be in its excited state
  4. Amount of energy absorbed by electron as it jumps equals difference in energy between levels
  5. Electrons in excited state are unstable + fall back down to ground state
  6. As they fall back down, emit fixed amounts of light energy (seen as colours)
  7. Amount of light energy emitted is equal to amount of energy absorbed
  8. Light emitted appears as specific line on spectrum
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Formula for amount of energy absorbed by electron as it jumps from energy levels

A

E₂ - E₁ = hf

h = Planck's Constant
f = frequency
E₂ = excited state
E₁ = ground state
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

Spectrum series

A

Paschen series
Balmer series
Lyman series

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

Paschen series

A

invisible, infrared

n = 3

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

Balmer series

A

visible

  • seen as coloured lines in line emission spectrum
  • correspond to transitions of electrons from higher energy levels to second energy level (n = 2)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

Lyman series

A

invisible, ultraviolet

n = 1

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

Atomic absorption in spectrometry

A
  1. White light passed through gaseous sample of element, then through prism
  2. Spectrum of light that comes out has certain wavelengths of light missing
  3. Absorption spectrum of element is exact oppposite of its emission spectrum
  4. Electrons in ground state absorb some energies of light they emit
  5. Emission + absorption spectrums can be used to identify specific elements
  6. Uses: Atomic Absorption Spectrometer used to detect metals in water
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Sub-levels

A
  1. More sophisticated spectrometers developed to look at emission line spectra
  2. Noticed some lines actually made of two/more lines close together on spectrum, eg. ELS for sodium consists of two yellow lines not one
  3. Could not have resulted from electrons dropping to 2 diff. energy levels (lines would be further apart)
  4. Proposed each energy level (Excluding first) is made up of a number of sub-levels close in energy
  5. Discovered number of sub-levels an energy level has is same as value for n for that sublevel
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

number of sublevels in each energy level

A 
n=1 has 1 sublevel
n=2 has 2 sublevels
n=3 has 3 sublevels
n=4 has 4 sublevels
etc
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

s, p, d configurations

A

1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 4d¹⁰ 4f¹⁴

check hardback

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

s orbitals

A

spherical

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

p orbitals

A

dubbell shaped,
3 separate orbitals (each contain 2 electrons)
all tree have same energy

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

Max no. of electrons that can be accommodated in an orbital

A

2

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

Wave nature of the electron

A

Discovered that the electron travels in a wave motion rather than in fixed paths as proposed by Bohr

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

4s and 3d sublevel

A
  1. According to Aufbau principle, when building up electronic config., electrons occupy lowest available energy level
  2. Since 4s sublevel is lower in energy than 3d sublevel, 4s sublevel is always filled before 3d sublevel
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
29
Q

Chromium & Copper

A

Electronic configuration of chromium + copper are not what one might expect

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
30
Q

Chromium configuration

A

Cr = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
31
Q

Why chromium config. is different

A
  1. It has been found that sublevels exactly half filled or completely filled have extra stability
  2. One of the electrons in the 4s² orbital ‘flips over’ to one of the d orbitals to give two sblevels which are exactly half filled
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
32
Q

Copper configuration

A

Cu = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰

33
Q

Why copper config. is different

A
  1. One of the electrons in 4s² orbital ‘flips over’ to one of the d orbitals
  2. Completely filled 3d sublevel + half filled 4s sublevel have extra stability associated with them
34
Q

Limitations of Bohr’s theory

A
  1. Worked perfectly to explain emission spectrum of hydrogen. Atoms with more than one electron -> failed to account for many of lines in emission spectra of these atoms
  2. Did not take into account the fact that electron had wave motion.
  3. Heisenberg’s Uncertainty Principle was in conflict with Bohr’s theory.
  4. Did not take into account the presence of sublevels.
35
Q

Hundi’s Rule of Maximum Multiplicity

A

States that when two or more orbitals of equal energy are available, the electrons occupy them singly before filling them in pairs

36
Q

Pauli Exclusion Prinicple

A

States that no more than two electrons may occupy an orbital and they must have opposite spin

37
Q

purpose of Milikan’s oil drop experiments

A

measure magniude of charge on electron

38
Q

why smth might have a higher boiling point

A

hydrogen bonding stronger than other intermolecular forces in other molecules

39
Q

why smth might have a lower boiling point

A
  • smaller molecules
  • fewer electrons
  • smaller degree of intermolecular forces
40
Q

electronegativity

A

relative attraction of an atom for the shared pairs of electrons in a covalent bond

41
Q

explain how line emission spectra occur

A
  • electrons restricted to energy levels
  • jump to higher levels when they absorb energy
  • fall back emitting energy as light
  • energy difference between levels gives specific frequency/wavelength of light in spectrum
42
Q

what evidence do line emission spectra provide for the existence of energy levels in atoms?

A

only fixed frequencies are emitted, therefore electrons must be restricted to certain energy values

43
Q

why is it possible for the line emission spectra to be used to distinguish between diff elements?

Why do diff elements have unique atomic spectra?

A

diff elements have diff characteristic spectra/diff distribution of energy levels

44
Q

how Bohr used line emission spectra to explain the existence of energy levels in atoms

A
  • electrons in ground state
  • fixed energies absorbed, jump to higher energy levels (excited state)
  • excited state unstable, fall back to lower levels, emitting energy as light
  • energy diff between levels gives specific freq of light in spectrum
45
Q

example of a radioactive isotope and one common use of it

A

carbon-14 - dating of ancient remains

cobalt-60 - cancer treatment

americium-241 - smoke alarms

46
Q

series of coloured lines in line emission spectrum of hydrogen corresponding to transitions of electrons from higher enregy levels to the second energy level

A

Balmer series

47
Q

describe the flame test

A
  • dip platinum wire in HCl to clean, clean if it does not change flame colour
  • dip in HCl
  • dip in sample of salt, sticks to wire
  • place salt in/over Bunsen flame
  • note colour of flame observed
  • if question specifies a salt, write the colour the flame should turn into
48
Q

reagents used in brown ring test for nitrate ion

A

iron(II) sulfate (FeSO₄)

concentrated sulphuric acid (H₂SO₄)

49
Q

distinguish between ground state and excited state for the electron in a hydrogen atom

A

ground: in lowest energy state
excited: higher energy state

50
Q

explain how the expression E2 - E1 = hf links the occurance of the visible lines in hydrogen spectrum to energy levels in a hydrogen atom

A

E2 - E1: energy difference between higher and level 2 / energy emitted when electron falls from higher to level 2

f: frequency of line in spectrum.

Each line produced due to electrons falling from particular higher level to particular lower level
h is Planck’s constant

the expression indicates that the energy difference (E2 - E1) is proportional to the frequency (f)

51
Q

max no. of electrons that can be accommodated in a p-orbital

A

2

52
Q

how melting point of either the crude product or recrystallised product of benzoic acid could have been measured

A
  • sample on heating block
  • thermometer in metling block
  • block heated
  • note temp range over which sample melts
53
Q

two ways you could conclude from melting point measurements that recrystallised product was purer than the crude product

A
  • higher melting point
  • melting point closer to correct values in table
  • sharper (narrower) range
54
Q

properties of cathode rays

A
  • negatively charged (attracted to anode)
  • negligible mass
  • straight-line motion
  • cause fluorescence
  • high-speed
55
Q

why some alpha particles deflected at large as they passed thru gold foil

A

repelled when passing near nucleus

56
Q

why some alpha particles reflected back along original paths when passed thru gold foil and why only a small number did this

A
  • collided with nucelus (positive (+) core/centre)

- nucleus very small

57
Q

new structure of atom proposed by rutehrford

A

dot as nucleus in centre, surrounded by a circle with a dot (electron) on it

58
Q

explain, in terms of atomic structure, why diff flame colours are observed in flame tests using salts of diff metals

A
  • metal atoms of diff elements have diff sets of energy levels/diff electron arrangement
  • emit diff frequencies of light (have difff characteristic spectra)
59
Q

main energy levels involved in electron transition that gives rise to first (red) line of Balmer series in emission spectrum of hydrogen atom

A

3 and 2

60
Q

when counting how many orbitals are occupied

A

take into account that some have multiple orbitals, eg.
p has px, py, pz.
if there is one electron in px and one in py, then it only occupies px and py, so that’s 2 orbitals occupied in the p.

61
Q

why is it difficult to specify the absolute boundary of an atom?

A

orbitals do not have defined volumes / Heisenberg’s uncertainty principle (state) / electrons have wave nature

62
Q

successive ionisation energy values for electrons in carbon are: 1086, 4620, 6223, 37831, 47277. How does this provide evidence for number of electrons in a carbon atom and the number of electrons in each main energy level?

A

1: six values, six electrons
2: gradual increase for first four values - four electrons in outer main energy level. Last two values very high - two electrons in inner main energy level

63
Q

Why is the first ionisation energy of oxygen lower than nitrogen despite increase in values across second period?

A

nitrogen is relatively stable, has a half-filled 2p sublevel and has a more stable electron configuration

64
Q

How an electron can become excited

A

add energy (eg. heat/electricity)

65
Q

explain the origin of the series of visible lines in emission spectrum of hydrogen

A
  • excited electron falls back from n = 3, 4, etc
  • to n = 2
  • energy lost is emitted as light
66
Q

why is there no yellow line in hydrogen emission spectrum?

A

no corresponding electron transition

67
Q

distinguish between a 2p orbital and a 2p sublevel

A

2p sublevel consists of three 2p orbitals of equal energy

68
Q

describe how electrons are arranged in orbitals of the highest occupied sub-level of nitrogen in its ground state

A

just write electronic config split up into orbitals (eg. 2px¹ 2py² 2pz³) or use the arrow diagrams

69
Q

when accounting for trends in ionisation energies

A

explain why it increases/decreases and also explain any exceptions

70
Q

experimental evidences for existence of energy levels in atoms

A
  • line emission spectra of elements

- successive ionisation energies of elements

71
Q

bond energy

A

average energy required to break a bond and to separate the atoms

72
Q

what do the electron config of the series of elements from scandium to zinc have in common?

A

electrons occupying 3d sublevel all end in 3dₓ
/
all have electrons in 3d

73
Q

explain, in terms of structures of the atoms, the trend in reactivity down the group 1 of the periodic table

A
  • reactivity increases
  • increase in atomic radius
  • effective nuclear charge is the same, screening efect of inner shells cancels increase in nuclear charge
  • outermost electrons less tightly held by nucleus
74
Q

explain how the emission spectrum of x gives evidence for energy levels

A

describe electrons becoming excited and falling back down

75
Q

difference between electron in 2s and 3s orbital

A

energy of 2s less than 3s / nuclear attraction for 2s electron greater than 3s

76
Q

“How many orbitals are in x” vs “how many sublevels are in x” vs “how many energy levels are occupied in x”

A

energy levels: 1, 2, 3, etc
sub-levels: 1s, 2s, 2p, 3s, 3p, 4s, 3d etc
orbitals: 1s, 2s, 2px, 2py, 2pz, 3s, etc…

77
Q

state the max number of electrons that can be in a p orbital

A

2

78
Q

when asked for an element

A

write the full name, dont just write the symbol

eg. Write Helium (He), rather than just He

79
Q

explain why second ionisation energy of Sodium significantly higher than first, while its not significantly higher in Neon

A

Sodium: first electron removed from outer shell, whereas second electron removed from second/inner shell

Neon: second electron from same sublevel, both removed from 2/same shell

LC Chemistry (26 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions