Arrangement of electrons in atom

Neils Bohr Investigated the location of electrons in the atom

Neils Bohr Looked at the emission line spectrum of elements When light passes through a prism it splits to give a continuous spectrum

When light passes through a prism it splits to give a spectrum

Bohr repeated this experiment using light from a hydrogen discharge bulb hydrogen discharge bulb = glass tube filled with hydrogen gas and an electric current passed through Instead of seeing a continuous spectrum he saw a series of narrow lines

Bohr repeated it with other elements Saw that each element had a different emission line spectrum Therefore, you can identify an element by looking at its line spectrum

Bohr concluded that the different arrangement of electrons in each atom must explain why each element has a different emission line spectrum

Spectrometer

Uses of Emission Spectra Yellow street lights = sodium Advertising signs = Neon

Uses of Emission Spectra Fireworks!

Mandatory Experiment 1.1 Flame Tests

Bohr’s Theory Electrons revolve around the nucleus in fixed paths called energy levels (or orbits/shells) Energy levels are represented by the letter n Lowest: n=1, n=2 n=3 etc n=1 n=2 n=3

Bohr’s Theory Each energy level has a fixed amount of energy An electron in an energy level posses the same amount of energy as the energy level An energy level is the fixed energy value that an electron in an atom may have

n=1 n=2 n=3

Atoms normally exist in the ground state (they occupy the lowest available energy level) If an electron is given energy (from electricity/heat) they can jump from the lower energy levels to higher ones. n=1 n=2

When this happens an electron is said to be in its excited state The excited state of an atom is when the electrons occupy higher energy levels than those available in the ground state n=1 n=2

The amount of energy absorbed by the electron as it ‘jumps’ from the lower to the higher energy level equals the difference in energy between the levels Orbit of energy E1 Orbit of energy E2 n=1 n=2

Electrons in the excited state are unstable and they fall back down to the ground state As they fall back down they emit fixed amounts of light energy (seen as colours) The amount of light energy emitted is equal to the amount of energy absorbed n=1 n=2

The energy corresponds to a certain wavelength of light The amount of light energy emitted is equal to the amount of energy absorbed The energy corresponds to a certain wavelength of light E=hf n=1 n=2

Excited State Spectrum Excited State Excited State UV IR Ground State Excited State unstable and drops back down Excited State UV But only as far as n = 2 this time n=2 Vi s ible Energy released as a photon Frequency proportional to energy drop IR n=1 Ground State

The light emitted appears as a specific line on the spectrum

The frequency of the light emitted depends on the difference in energy between the two energy levels E2 - E1 = hf Frequency of light emitted Energy of higher energy level Planck’s constant Energy of lower energy level

Bohr Theory- Key points The electron in a hydrogen atom occupies fixed energy levels in its ground state, electrons occupy the lowest available energy level The elctron can move (jump/become excited) to a higher energy level if it receives a certain amount of energy The energy (photon) absorbed must exactly equal the energy difference between the ground state (lower level) and excited state (higher level) according to E2-E1= hf The excited state is unstable and the electron falls back to a lower level It emits the excess energy in the form of a photon of light (hf) with definite frequency (wavelength) according to E2 – E1 = hf

Bohr Theory- Key points Frequency depends on difference in energy levels When an electron falls to: n=1 level gives UV range (Paschen Series) n=2 level gives Visible Range (Balmer Series) n=3 level gives IR Range (Lyman Series)

Exam Q How does the lines on emission spectrum provide evidence for the existence of energy levels? The lines on the emission spectrum are produced when an electron falls from the excited state to the ground state The line is the light energy produced when it falls The light energy is the difference between the energy level, so shows the existence of energy levels

Atomic Absorption Spectrometry

White light is passed through a gaseous sample of the element and then through a prism The spectrum of light that comes out has certain wavelengths of light missing The absorption spectrum of an element is the exact opposite of its emission spectrum

Principle behind AAS based on 2 facts: Electrons in the ground state absorb the same energies of light that they emit The amount of light absorbed is directly proportional to the concentration of the element present in the sample Emission and absorption spectrums can be used to identify specific elements Uses: Atomic Absorption Spectrometer used to detect metals in water

Sub-levels In many cases, it appeared that some lines were actually made of two or more lines close together on the spectrum E.g: ELS for sodium consists of two yellow lines not one Could not have resulted from electrons dropping to 2 different energy levels (lines would be much further apart)

Sub-levels Proposed that each energy level (excluding the first) is made up of a number of sub-levels which are close in energy Discovered that the number of sub-levels an energy level has is the same as the value for n for that sub-level n = 1; 1 sub-level n = 3; 3 sub-levels n = 2; 2 sub-levels n = 4; 4 sub-levels

Sub-level of lowest energy is s The next lowest energy is p Followed by d and f.

Energy sublevels n=1 1s n=2 2s 2p n=3 3s 3p 3d n=4 4s 4p 4d 4f n=5 5s 5p 5d 5f n=6 6s 6p 6d 6f n=7 7s 7p 7d 7f s=sharp p=principal d=diffuse f=fundamental

Bohr theory and emission spectra https://www. youtube. com/watch Absorption v emission spectra https://www.youtube.com/watch?v=1uPyq63aRvg

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

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

Heisenberg’s Uncertainty Principle: It is impossible to measure both the position and velocity of an electron at the same time

Bohr’s Theory did not take into account the fact that the electron had a wave motion Electrons do not revolve around the nucleus in fixed paths as proposed by Bohr -> Heisenberg’s Uncertainty Principle was in conflict with Bohr’s Theory. Therefore, we can only talk about the probability of finding an electron in a region in space They exist in orbitals

Schrodinger Worked out the probability of finding an electron in any sublevel Schrodinger’s equations allowed for the calculation of the SHAPES of these orbitals

An Orbital: An orbital is a region in space where there is a high probability of finding an electron A Sublevel: Is a subdivision of a main energy level and consists of one or more orbitals of the same energy

S orbitals are spherical An Orbital: An orbital is a region in space where there is a high probability of finding an electron S orbitals are spherical

An Orbital: An orbital is a region in space where there is a high probability of finding an electron P orbitals are ’dumbbell” shaped Consist of three separate orbitals (each containing 2 electrons) on the three axis All three have the same energy

Energy sublevels https://www.youtube.com/watch?v=9E3QaRxqXZc

Limitations to Bohr’s Theory Bohr’s theory only worked to explain the emission spectrum of Hydrogen Bohr’s theory did not take into account the fact that an electron had a wave motion Heisenberg’s Uncertainty Principle was in conflict with Bohr’s theory Bohr’s theory could not explain the splitting of certain lines in emission spectra and did not take into account the presence of sublevels

N.B: Do not confuse orbit and orbital Key Words Energy Level/Orbit Emission spectrum Ground state Excited state Atomic Emission Spectrometry Atomic Absorption Spectrometry Energy Sublevels Wave nature Orbital N.B: Do not confuse orbit and orbital