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Flashcards in Electronic Structure of Atoms Deck (79)
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1

Energy Level

Fixed energy value that an electron in an atom may have

2

Ground State

Lowest energy state (in 1s orbital)

3

Excited State

Higher energy state

4

Sublevel

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

5

Orbital

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

6

Heisenberg's Uncertainty Principle

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

7

The atomic radius of an atom

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

8

The first ionisation energy of an atom

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

9

Aufbau Principle

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

10

Hund's Rule of Maximum

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

11

Pauli Exclusion Principle

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

12

Niel's Bohr

-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

13

Bohr's theory

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

14

Electrons in excited state

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

15

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

E₂ - E₁ = hf

h = Planck's Constant
f = frequency
E₂ = excited state
E₁ = ground state

16

Spectrum series

Paschen series
Balmer series
Lyman series

17

Paschen series

invisible, infrared

n = 3

18

Balmer series

visible

-seen as coloured lines in line emission spectrum
-correspond to transitions of electrons from higher energy levels to second energy level (n = 2)

19

Lyman series

invisible, ultraviolet

n = 1

20

Atomic absorption in spectrometry

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

21

Sub-levels

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

22

number of sublevels in each energy level

n=1 has 1 sublevel
n=2 has 2 sublevels
n=3 has 3 sublevels
n=4 has 4 sublevels
etc

23

s, p, d configurations

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

(check hardback)

24

s orbitals

spherical

25

p orbitals

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

26

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

2

27

Wave nature of the electron

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

28

4s and 3d sublevel

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

29

Chromium & Copper

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

30

Chromium configuration

Cr = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵