Brainscape's Knowledge GenomeTM

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


See full index

IB Chem HL > S1.3 (Electron Configuration) > Flashcards

S1.3 (Electron Configuration) Flashcards

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

Emission Spectra

A

Range of frequencies or wavelengths of electromagnetic radiation emitted during an electron transition from a higher energy to a lower energy level shown by lines converging at higher energy

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

Electron Transition

A

Movement of an electron between energy levels in an atom

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

electromagnetic Spectrum

A

Radiowaves<Microwaves<Infra red<Visible Light<UV<X Ray< Gama Ray

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

highest energy & frequency wave

A

UV

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

continuous spectrum

A

all the freuencies across a range of electromagnetic radiation (rainbow)

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

Emission line Spectra (How it works)

A
  • energy is applied to a sample of vaporized hydrogen
  • only shows specific wavelengths: colored lines on black sheet of paper
    *
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Why electrons prodcue colored lines

A
  1. electron absorbs energy
  2. n=2 -> n=3
  3. electron is excited & unstable
  4. electron releases energy
  5. n=3 -> n=2
  6. emitted energy is shown as wave
  7. line is observed
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

What is special about the lines on an emission spectra sheet

A

they converge at high frequencies

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

Hydrogen Emission Spectrum

A

The emission line spectrum of hydrogen (proves that electrons stay at specific energy levels)

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

Hydrogen Emission Spectrum

UV

A

higher -> n=1

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

Hydrogen Emission Spectrum

Visible Light

A

higher -> n=2

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

Hydrogen Emission Spectrum

IR

A

higher -> n=3

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

meaning of n=∞

A

point where the electron has left the atom thus having an infinite amount of energy

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Heisenbergs uncertainty principle

A

it is impossible to know the postion and momentum of an electron at once

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

principle quantum numbers

A

shows main energy levels of an electron (n)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Orbital

A

A region of space with the highest possibility of finding an electron

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

How many electrons can each/any orbital hold

A

2

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

S orbital

A
  • spehrical
  • lowest in energy
  • size increases as shell number increases
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

Total number of orbitals in each energy level

A
  • n=1; 1
  • n=2; 4
  • n=3; 9
  • n=4; 16
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

P Orbital

A
  • dumbell shaped
  • has 3 degenerate orbitals (Py, Px, Pz)
  • total of 6 e-
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

D Orbital

A
  • double dumbel shaped
  • 5 orbitals
  • can hold a total of 10 e-
21
Q

F orbital

A
  • 7 orbitals
  • total of 14 e-
22
Q

Total number of electrons in each energy level

A
  • n=1; 2
  • n=2; 8
  • n=3; 18
  • n=4; 32
23
Q

Sublevels in each energy level

A
  • n=1; 1s
  • n=2; 2s,2p
  • n=3; 3s,3p,3d
  • n=4; 4s,4p,4d,4f
24
Q

Aufbau Principle

A

Electrons fill the lowest energy levels first and fill orbitals in order of increasing energy

25
Q

Hunds Rule

A

When degenerate orbitals are available, electrons will fill those orbitals singly and with parallel spins before pairing up

25
Q

Pauli Exclusion Principle

A

Two electrons in the same orbitals must have opposite spins

25
Q

Aufpau Principles filling order

A

1s>2s>2p>3s>3p>4s>3d>4p>5s>4d>5p>6s

26
Q

Noble Gas Notation example

Calcium (20 e-)

A

(Ar)4s2

26
Q

Shortened Spectroscopic Notation/ Noble Gas notation

A

factors out the part of the notation which overlaps with one of the noble gases

26
Q

Spectroscopic Notation

Exceptions

A
  • copper
  • chromium

they first fill the d orbitals singly rather having them half full

27
Q

Orbital box Notation

A

shows electrons and their pairing + spins

28
Q

Eelectron configurations of ions

A
  • filling= lowest first
  • removing= highest first
  • except: 4s first in & first out
29
Q

Ionisation Energy

A

The amount of energy required to remove one mole of electrons from one mole of gaseous atoms

30
Q

Why ionisation energy increases across a period

A
  • nuclear charge increases as protons increase
  • this means stronger attraction between nucleus and electrons
  • energy required to remove the outer electrons increases
30
Q

Trend in ionisation energies

A
  • ionistaion energy increases across a periods
  • ionisation energy decreases going down a group
31
Q

Why ionisation energy decreases down a group

A
  • number of shells increase, increasing the distance from nucleus and electrons
  • attraction between electrons and nucleus decreases => easier to remove electrons
  • &shielding from electrons in shells decreases attraction

=> energy required to remove an electron decreases

32
Q

Convergence limit

A

en electron is removed from the attraction of the nucleus (ionised)

33
Q

While knwoing the frequency

ionisation Energy calculations (kJ/mol)

A
  1. use E=hf to calculate energy in Joules
  2. then divide by 1000 for kJ
  3. then times by avogadros constant (L) for K+kJ per mole
34
Q

Without knowing the frequency

Ionisation energy calculations (kJ/mol)

A
  1. f=c/λ and E=Lhf (per mole) so subtitude f into E=Lhf
  2. E=Lhc/λ
  3. to get in in kJ = E= LHC/1000λ
34
Q

Succesive Ionisation Energies

A

process of removing successive electrons from an atom or positive ion

34
Q

Important about the second ionisation energy

A

It is always significantly larger than the first as you’re trying to remove an electron from a positive ion meaning the attraction to the nucleus is greater

(the more you need to remove the larger the ionisation energy)

35
Q

Deducing valence electrons based on ionisation energy

A

the ionisation energy difference which is the largest means that was the last electron in the shell that was being removed

36
Q

why is there a general increasing value for ionization energy across a period

A

electrons are being added to the same energy level, but the amount of protons increases
=> this means the attraction between them gets stronger

37
Q

Things to look out for when comparing ionization energies (especially graph)

A
  • trends across period (ionization increases)
  • trend across groups (ionization increases down)
  • DRAW ELECTRON CONFIGUARTION and speak about energy levels
38
Q

What happens to ionisation energy when electrons are paired vs unpaired

A
  • if all are unapired = harder to remove 0 ionisation energy increases
  • if some are paired & unpaired= paired ones are easier to remove due to repulsion = decreases ionisation energy
39
Q

Evidence that electrons exist in fixed energy levels

A
  • emission spectra is discrete lines (if continuous: electrons would be anywherre)
  • the lines converge showing different energy levels as the lines are succesively converging
  • the differneces in ionization energies shows that different energies are needed to remove from different levels
40
Q

limitations of hydrogen emission spectra

A
  • does not represent sub levels/orbitals
  • only applies to hydrogen atom/atom of one electron
  • does not explain why only certain energy levels allowed
  • the atom is considered to ber isolated
  • does not consider the number of electrons thta can fit into a energy level
  • does not consider the probability of finding an electron (orbitals)
41
Q

emission spectrum described

A
  • series of lines
  • only at specific wavelengths
  • converge at high frequences
42
Q

relationship of emission spectrum tp hydrogen atom (outlined)

A
  • electron transition between high energy level to lower energy level
  • DRAW DIAGRAM
43
Q

converting nm to m

A

multiply by 10xe9

44
Q

things to consider when speaking about ionisation enrgy trends

A
  • shielding => more shielding = less attaraction
  • which orbital/subshell or main energy level it is taking e- out of
  • attraction to nucleus (is it already psotive or negatively charged?)

IB Chem HL (8 decks)

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