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electron configuration Flashcards

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1
Q

define relative atomic mass (bp)

A

ratio of average mass of an atom to 1/12 the mass of C-12 isotope OR average mass of an atom on a scale where one atom of C-12 has a mass of 12/sum of the weighted average mass of isotopes of an element compared to C-12

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2
Q

define isotope (bp)

A

atom of same element/same number of protons but with different mass number/number of neutrons

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3
Q

define limit of convergence [1]

A

the frequency (/wavelength) at which the spectral lines converge – from this the ionisation energy can be calculated

(as the energy levels are closer together at high energy and short wavelength, the emission lines merge at the limit of convergence)

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4
Q

define ionisation energy

A

energy (kJmol^-1) required to completely remove one mole of electron from one mole of gaseous atom

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5
Q

give the equation of ionisation

A

X (g) →X+ (g) + e-

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6
Q

is ionisation exothermic or endothermic, and why?

A

endothermic

because energy is required to break the force of attraction between the electron and the central positive nucleus (study mind)

take in energy to break the bond

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7
Q

excited states meaning

A

at higher energy state

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7
Q

how are emission spectra produced?

A

by atoms emitting photons when electrons in excited states return to lower energy levels

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8
Q

how to calculate frequency

A

f = c / λ

c = speed of light

λ = wavelength (m)

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9
Q

how to convert m to nm?

A

1nm = 1x10^-9m

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10
Q

how to calculate relative atomic mass?

A

(isotope abundance x isotope mass number)/isotope abundance [total]

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10
Q

outline the Bohr’s model deduced from the hydrogen line emission spectrum [2] (WS)

A

electrons are in specific energy levels and energy levels get closer together at higher energies

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11
Q

state one limitation of the Bohr model [1] (WS)

A

it cannot be applied to many electron atoms / does not predict the intensity of different lines

it does not take into account the interactions between external fields

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12
Q

how the first ionisation energy of an atom can be determined from its emission spectrum [2] (WS)

A

the line referring to the highest frequency of electromagnetic waves emitted is used to determine the first ionisation energy. the first ionisation energy of an atom can be calculated usiing the frequency of the limit of convergence, using E = hf & f = c/wavelength

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13
Q

explain the general increase in trend in the first ionisation energies of the period 3 elements, Na to Ar

A
  1. nuclear charge increases, increasing number of protons
  2. valence electrons are in the same energy level - similar shielding effect as the number of inner shell electrons are the same – valence electron experience stronger attractive force from the nucleus and gets pulled in
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14
Q

why the first ionisation energy of nitrogen is greater than oxygen

A

electron config of nitrogen: electrons fill up the 2p sub level and are unpaired

electron config of oxygen: one paired electron and 2 unpaired electrons in the 2p sub level, the paired electron causes there to be more electron-electron repulsion = easier to remove the valence electron and there is lower ionisation energy

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15
Q

what happens to wavelength when frequency increases?

A

wavelength decreases

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16
Q

what happens to energy if frequency increases?

A

frequency increases

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17
Q

relationship between frequency, wavelength & energy

A

wavelength decrease = frequency increase = energy increase

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18
Q

how to calculate ionisation energy?

A

E = hf

h = planck’s constat

f = frequency

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19
Q

convert 700nm to m

A

700 x (1 x 10^-9) = 7 x 10^-7

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20
Q

explain what is limit of convergence

A

energy levels are closer together at high energy / high frequency / short wavelength

emission lines merge at the limit of convergence

the frequency is used to determine the IE

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21
Q

X(g) + ______ → X+ (g) + e-

A

energy

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22
Q

convert PHz to Hz

A

1PHz = 1 x 10^15Hz

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23
how to calculate ionisation energy a hydrogen atom (g) , given the frequency of 3.2PHz
1. convert PHz to Hz 3.2 x 1 x 10^15 2. E = hf 3.2 x 10^15 x 6.63 x 10^-34 = 2.1 x 10^-18 J 3. "atom" = molar ionisation energy of hydrogen E x avogadro constant 2.1 x 10^-18 x 6.02 x 10^23 = 1.3 x 10^6 Jmol^-1 = 1.3 x 10^3 kJmol^-1
24
convert J to kJ
1J = 0.001 kJ divide 1000
25
convert Jmol^-1 to kJmol^-1
1J = 0.001 kJ divide 1000
26
why do group 1 elements have the lowest value?
electrons put into a higher energy level, *which is further from nucleus*, the attractive force is weaker
27
explain the trend: IE increases across a period
nuclear charge increase (with increasing number of proton) all valence electrons in same energy level similar screening/shielding effect (number of inner shell electrons are the same) valence electron experience stronger attractive force from the nucleus and gets pulled in
28
UV (energy + wavelength)
- high energy - short wavelength
29
infrared (IR) (energy + frequency)
- low energy - low frequency
30
effective nuclear charge
net positive charge experience by an electron in a multi-electron atom
31
define shielding effect
*reduction in the effective nuclear charge on the electron cloud* due to a difference in the attraction forces of the electrons of the nucleus. alternate names: screening effect, atomic shielding
32
what is the evidence of main energy level?
group 1 elements have lowest value + IE increases across a period (generally)
33
what is the evidence for sub level (e.g. s, p.. etc)
drop in IE from group 2 to group 13 drop in IE from 15 to group 16
33
s orbital
s orbital for every energy level probability of finding electrons spherical shaped each s orbital can hold 2 electrons 1s, 2s, 3s etc
34
p orbital
start at second energy level dumbbell shaped equal energy, 3 diff directions (Pz, Px, Py) each p orbital can hold 2 electrons
35
d orbital
start at 3rd energy level 5 diff shape dxy, dyz, dxz, dx2-y2, and dz2
36
f sublevel
start at 4th energy level 7 diff shape
37
aufbau principle
electrons placed in lowest energy orbitals first difficulties: overlapping of orbitals of diff energies
38
pauli exclusion principle
at most 2 electrons per orbital with opposite spins
39
hund's rule
orbitals of equal energy are singly filled first, before doubly filled
40
sequence with which the orbitals fill with electrons
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s etc
41
why is there a fall in IE from Be to B?
p orbital is a higher orbital, and the electron is further from nucleus (2p)
42
why is there a fall in IE from N to O?
electron pairing N: electrons fill up p sub-levels and are unpaired O: paired electrons in p sublevel - more electron-electron repulsion - easier to remove valence electrons - hence lower IE
43
give the equation for 2nd ionisation of Mg
Mg+(g) → Mg2+ (g) + e- from monocharge cation in gaseous state
44
give the equation of 3rd ionisation of Mg
Mg2+ (g) → Mg3+ (g) + e-
45
define second ionisation energy
energy (kJmol^-1) required to completely remove one mole of electron from one mole of monocharge cation in gaseous state
46
successive ionisation energy what are the (general) factors affecting the interaction?
distance between protons and electron charge
47
[Mg] why does the IE increase when removing the 2nd electron?
– Remove electron from positively charge ion – Less electron-electron repulsion – Attract closer to the nucleus → higher IE
48
[Mg] why is there a sudden rise in IE when removing 3rd electron?
- Lower energy level – Electron closer to nucleus – Sudden large rise in IE
49
[Mg] why is there a gradual increase in IE when removing 3rd-10th electron?
– Gradual increase in IE – As in same energy level – increasing positive charge on ion attracts electrons more strongly – decreasing repulsion between electrons
50
[Mg] what causes the difference (increase) in IE between 10th and 11th IE?
* 10th is removing electron from electronic arrangement 2,1 while 11th ionization energy is removing electron from electronic arrangement 2; * 11th electron removed is much closer to the nucleus / 11th electron removed from a (much) lower energy level/shell;
51
why is this equation not for ionisation energy? Br2(g) + 2e- → Br2 2- (g)
not removing one mole of electron → electron should be on RHS
52
why is this equation not for ionisation energy? Br2 (g) → Br2 + (g) + e-
Br2 (g) is a molecule, not atom
53
fill in the blanks when _____ return to a lower energy level (from _____), a _____ will be ______.
when electron return to a lower energy level (from higher energy level), a photon will be emitted.
54
tip about relative atomic mass, it should be between:
range between heaviest and lightest isotope
54
when predicting mass spectrum, the length of lines should also consider:
probability e.g. 79Br-81Br have 2/4 probability, so 2x length of other 2 lines
55
at higher emission there is ____ lines
denser also lower frequency
56
x and y axis of mass spectrum (graph)
y = abundance x = m/Z (values given next to the nuclear symbol + numbers added up tgt considering probability)
57
define isotope (cky)
different atoms of the same element, with the same number of protons (and electrons) but different number of neutrons
58
isotope same electron configuration = ?
same chemical properties
58
isotope different number of neutrons = ?
different masses and hence *slightly different physical properties* e.g. density, melting point, boiling point
59
what is continuous spectrum?
consist of radiation at every wavelength (all energy levels are possible)
60
what is line spectrum?
a line spectrum has only lines of specific wavelengths a line spectrum has only (lines of) sharp/discrete/*specific* colors/*wavelengths*/*frequencies* specific wavelengths / specific frequencies
61
what does the line emission spectrum of hydrogen provide evidence for?
existence of electrons in discrete *energy levels*, which *converge* at higher energies
62
characteristics of hydrogen emission spectrum & explanation based on Bohr's model 1
1) radiation is emitted at specific wavelengths (not continuous) explanation: electrons in discrete/specific/certain/different shells/energy levels
63
characteristics of hydrogen emission spectrum & explanation based on Bohr's model 2
converge at low wavelength / high energy explanation: energy levels get closer together at higher energies
64
limitations of Bohr's model
- only applies to atoms with one electron / hydrogen - does not consider probability of finding electron at different positions - does not take into account the interactions between atoms / molecules / external fields
65
peschen series (infra-red)
down to n=3
66
balmer series (visible)
down to n=2
67
lymann series (UV)
down to n=1
68
which transition will give you the highest energy transmission a) n=2 -> n=infinity b) n=infinity -> n=2 3) n=1 -> n=2 d) n=2 -> n=1
D
69
n=1 -> n=2 what does this show?
excitation requires energy
70
is gap from n=2 to n=1 larger or smaller than gap from n=infinity to n=2, and why?
n=2 to n=1 is LARGER energy levels converge at higher energy levels
71
true or false: He (helium) atoms give emissions, which is a line spectrum
TRUE
72
true or false: n=2 -> n=3 is an IR emission
FALSE excitation
73
true or false: emission spectra can be used to identify different elements
TRUE elements give emissions at own specific wavelengths
74
why do the emission lines converge?
energy diff decrease when energy increase
75
n=1 is which shell?
outermost
76
predict the absorption spectrum of hydrogen
opposite of continuous spectrum energy of emissions -> energy diff of energy levels amount absorbed = amount emitted if energy transition process is reversed coloured background with dark colored lines at the same wavelength as the emission spectrum
77
exceptions to aufbau principle
chromium Cr and copper Cu
78
electron configuration of chromium
[Ar] 4s1 3d5 4s become 3d
79
electron configuration of S2-
1s2 2s2 2p6 3s2 3p4
79
electron configuration of Na+
1s2 2s2 2p6
80
electron configuration of Zn+
[Ar] 4s1 3d10 energy lost from 4s before lost from 3d, bc 4s is the outermost energy level (main energy level 4) 4s electron is removed 4s removed before 3d in transition elements (now 4s have higher energy)
81
electron configuration of Cu+
1s2 2s2 2p6 3s2 3p6 3d10
82
energy levels for an electron in a hydrogen atom are () near the nucleus
farther apart near the nucleus
83
how many occupied main electron energy levels are there in an atom of sodium? (electron arrangement 2,8,1)
3 1s2 2s2 2p6 3s1 1,2,3
84
how many occupied p orbitals are there in an atom of sodium?
3 because Px, Py, Pz 1s2 2s2 2p6 3s1
85
UV light (in relation to emission spectrum) is only for?
hydrogen
86
state a physical property that is different for isotopes of an element
mass/density/melting point:boiling boiling for gases: rate of effusion/diffusion NO MASS NUMBER
87
describe the emission spectrum of hydrogen. outline how this spectrum is related to the energy levels in the hydrogen atoms. (bp 21)
series of lines/lines electron transfer/transition between higher energy level to lower energy level/electron transitions into first energy levels causes UV series/transition into second energy level causes visible series/transition into third energy levels causes infrared series convergence at higher frequency/energy/short wavelength
88
electromagnetic radiation in order of increase wavelength (shortest first) (=highest energy first)
y-rays/gamma ray → x-rays → UV → visible (VIBGYOR) → infrared → microwaves → radiowaves
89
colors in visible region
VIBGYOR violet indigo blue green yellow orange red
90
acronym for electromagnetic wave (starting with shortest wavelength)
Ronald McDonald Is Very Unusually Xtra Great! RMIVUXG
91
diff between continuous spectrum vs line spectrum (bp)
continuous: has all colors/wavelengths/frequencies line: has only (lines of) sharp/discrete/specific colors/wavelengths/frequencies
92
explain the large increase between 2 ionisation energies (10/11) (bp)
10th electron comes from 2nd energy level/n=2 AND 11th electron comes from 1st energy level/n=1; electron in 1st energy level closer to nucleus; electron in 1st energy level not shielded by inner electrons/exposed to greater effective nuclear charge
93
explain the general increase in successive ionisation energies of the element (bp)
successive electrons (are more difficult to remove because each is) taken from more positive charged ion [increased nuclear charge?] increased electrostatic attraction
94
which species possesses only 2 unpaired electrons? a. Zn b. Mg c. Ti 2+ d. Fe 2+
c. Ti 2+

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