Chpt. 3: Arrangement of Electrons in The Atom

Remember from Crookes experiment: when cathode rays (electrons) struck glass at far end of tube they caused it to glow it was the energy of the electrons that caused this glow

Bohr’s Study of Spectra (1920) through his study of spectra Niels Bohr provided an insight into the arrangement of electrons inside the atom Bohr’s Theory of the structure of the atom was based on the line spectra produced when electrons in an atom are excited by heat or high voltage currents

Understanding ‘Spectra’ Do you remember the spectrum of white light??? When white light is passed through a prism it is broken up into a array of colours – continuous spectrum from red to violet

Emission Line Spectrum of Excited Hydrogen Gas (Bohr’s experiment!!!) Light from a hydrogen discharge tube (connected to power supply) passed through a prism results in a series of narrow lines

Spectrum produced is known as a emission line spectrum because: - it consists of lines (rather than being continuous ) - light is emitted when an electric current is passed through *hydrogen

Line Spectra of other elements *Note: each element has its own unique emission line spectrum e.g. like a fingerprint

Demonstration Viewing of emission spectra of elements using a spectroscope or a spectrometer *Note: Emission spectrometer instrument used to analyse materials for the presence of certain elements based on their line emission spectrum

Mandatory Experiment 1 To carry out flame tests with salts of lithium, sodium, potassium, barium, strontium and copper *Note: may also be phrased as test to identify a metal

Bohr realised that any model of the atom must explain two things: 1. why emission spectra of the elements are line spectra 2. why the emission spectrum of each element is unique to that element

Bohr’s Theory of Atomic Structure Bohr suggested a way to relate the emission line spectrum of hydrogen to the structure of an atom of hydrogen and it may be summarised as follows:

1.Electrons revolve around the nucleus in fixed paths called orbits 2.Electrons in any one orbit have a fixed amount of energy – orbits also known as *energy levels 3.As long as an electron is in any one particular energy level, it neither gains nor looses energy 4.When atoms absorb energy, electrons jump from a lower energy level to a higher energy level 5.Energy is lost when an electron falls from a higher energy level to a lower energy level * The energy of the electron in a particular orbit (energy level) is fixed (quantised) at a definite value

Orbit Of Energy E 2 Orbit of energy E 1 Energy absorbed when e - jumps from n = 1 to n = 2 energy level n = 2 n = 1 Energy emitted when e - falls from n = 2 to n = 1 energy level Hydrogen Atom

An energy level: is defined as the fixed energy value that an electron in an atom may have. *Note: each energy level has a number, n, called the Principal Quantum Number

How are the lines of the emission line spectrum formed??? (Bohrs explanation for the existence of energy levels in atoms)

n=1 n=2 n=3 n=4 Spectrum VisibleVisible Excited State Excited State unstable and drops back down Energy released as a photon Frequency proportional to energy drop Excited State But only as far as n = 2 this time IR UV

Electron normally in Ground State (lowest energy level n = 1) Energy supplied [ as heat or electricity] to atom and the electron absorbs a certain amount of it Electron jumps to higher energy level Now in Excited State Unstable therefore only a temporary state Drops back to a lower level Formation of lines of emission line spectrum

Definite amount of energy that was absorbed to make the jump up is now released as a photon Definite amount of energy emitted (photon) appears as a line of particular colour in the spectrum. Each colour corresponds to a particular wavelength or frequency Frequency depends on difference in energy levels : E 2 - E 1 = hf (E = hf) h is Plank’s Constant, f is frequency of light, E is energy of light emitted from atom

When electron falls to n = 1 level gives UV Range (Lyman Series) n = 2 level gives Visible Range (Balmer Series) n = 3,4 or 5 levels gives IR Range (Paschen Series) In conclusion since only definite amounts of energy are emitted, this implies that electrons can occupy only definite energy levels. Thus, energy levels must exist in the atom

Why has each element its own unique emisison line spectrum??? According to Bohr: Each element has a unique emission line spectrum because each element has a different number of electrons, and each element has its own arrangement of these electrons in different energy levels. As a result, the different electron transitions give rise to different emission spectra

Atomic Absorption Spectrometry Atoms can also absorb light If white light is passed through a gaseous sample of an element, the light which comes out has certain wavelengths missing i.e. dark lines will be seen in the spectrum Spectrum known as atomic absorption spectrum Technique used by chemists to detect the presence of certain elements and to measure the concentration of these elements e.g. Analysis of water for metals like lead and mercury

Absorption Spectrum Vs. Emission Spectrum

Energy Sublevels Advances in spectrometer technology resulted in a better understanding of energy levels e.g. Sodium emission spectrum – single yellow line actually two yellow lines very close together Scientists proposed that each main energy level, except the first, was made up of a number of sublevels all of which were close in energy Scientists discovered that number of sublevels in each energy level corresponds to the value of n for the main energy level i.e. n = 2 2 sublevels n = 33 sublevels n = 44 sublevels

Sublevel of lowest energy – s Next highest energy level – p Next highest energy level – d Highest energy level - f Main energy levelSublevel Energy

The Bohr model could explain the emission spectrum of hydrogen and other 1 electron species e.g. He +, Li 2+ however, it could not accurately explain the emission spectra of many electron atoms. Our current view of the atom is based on QUANTUM MECHANICS which builds upon the ideas of Bohr.

De Broglie – ‘Wave – Particle Duality’ (1924) Louis De Broglie suggested that all moving particles (including electrons) which have the properties of a particle also have a wave motion associated with them If electron has a wave motion then Bohr’s picture of it travelling in a precise path at a precise distance from the nucleus cannot be true

Heisenberg – ‘Heisenberg’s Uncertainty Principle’ (1925) Werner Heisenberg used mathematics to confirm De Broglie’s theory. He came up with a very famous principle ‘Heisenberg’s Uncertainty principle’: ‘ it is impossible to measure at the same time both the velocity and the position of an electron’

Heisenberg Vs. Bohr Bohr’s model - electrons moving with certain speed in orbits of fixed distances from the nucleus. Heisenberg disputed this model stating that you cannot measure the speed of the electron and its distance from the nucleus at the same time. * Electron is not travelling in a fixed path therefore can now only talk about the probability of finding it at a particular position inside the atom.

Atomic Orbitals An *orbital: is a region in space within which there is a high probability of finding an electron *Note: do not mix up orbital with Bohr’s term orbit

Schrödinger (1926) Erwin Schrödinger devised WAVE EQUATION to work out probability of finding an electron in any particular sublevel in an atom These equations were used to predict the shape of the s, p, d,& f orbitals

Types of Orbitals s Orbitald Orbitalp Orbital

All s orbitals are spherical and differ only in size

p Orbitals This is a p sublevel with 3 orbitals these are called: x, y & z p orbitals are dumb-bell shaped and only differ in their orientation in space

d orbitals = complex shapes (5 orbitals) f orbitals = complex shapes * Note: A sublevel is a subdivision of the main energy level made up of one or more orbitals of the same energy level.

Arrangement of Electrons in Atoms Electrons in atoms are arranged as LEVELS (n) SUBLEVELS ORBITALS ELECTRONS