Electronic Structure of Atoms Electrons inhabit regions of space known as orbitals. Heisenberg Uncertainty Principle NOTE Impossible to define with absolute precision, at the same time, both the position and the momentum of an electron. As we view the Bohr model the circles are NOT orbits. The electrons are NOT moving around the nucleus along the circles. Instead, the circles represent energy levels. The electrons on the circle closest to the nucleus have the lowest energy. The eight electrons on the next circle have a higher energy, etc. Electrons occupy energy levels that are outside the nucleus Different levels have different energy values Level 1 (N=1) can hold 2 electrons Level 2 (N=2) can hold 8 electrons Level 3 (N=3) can hold 18 electrons Level 4 (N=4) can hold 32 electrons Hydrogen's electron - the 1s orbital In the case of hydrogen, the electron can be found anywhere within a spherical space surrounding the nucleus. The diagram shows a cross-section through this spherical space. Each orbital has a name. The orbital occupied by the hydrogen electron is called a 1s orbital. The "1" represents the fact that the orbital is in the energy level closest to the nucleus. The "s" tells you about the shape of the orbital. s orbitals are spherically symmetric around the nucleus Emissions & light Absorption by electrons If we blast an atom with energy (under particular conditions) then the atom will emit light Remember flame tests
Element Colour Barium Yellow Green Copper Blue Green Lithium Deep Red Potassium Lilac Sodium Yellow Strontium Red We use this to our advantage in street lights A sodium Discharge Tube emits yellow light when subjected to a high voltage Analysing light emission A spectroscope can analyse light emitted by an element White Light is a mixture of visible light of all wavelengths In a spectroscope the light is bent to different extents Light from the highest energy levels are bent most The result is a spectrum of colours A continuous band Hydrogen emits a line spectrum Emission spectra are characteristic of an element We can use absorption spectroscopy to measure the amount of a sample of a particular element PRACTICAL DEMONSTRATION
1. Convert elemental sample into flame 2. Pass light of a suitable wavelength through it 3. Amount of light absorbed depends on the amount of the element in the sample Line Spectra and energy levels Niels Bohr 1913 The energy of the hydrogen electron is restricted to certain values and these values can be quantified When an electron moves from a higher energy level to a lower one a definite amount of energy is emitted E2 E1 = hf The amount of energy emitted is equal to the difference in the 2 energy levels E2 E1 = hf Where h = Planks constant (6.63 x 10-34 Js) F = frequency of light emitted Important If an atoms absorbs an amount of energy equal to the energy difference between two levels then it will move back up to the higher energy level Each energy level is assigned an integer n Called the principal quantum number Values correspond to energy levels with 1 being lowest Lyman Series Under normal circumstances the electron is in n=1 (ground state If the electron receives enough energy then it moves to n=2 (excited state) Electrons in this state are unstable and will eventually return to n=1 (Ground state) with emission of energy equal to the difference between n=1 an n=2 As a result a line appears in spectrum in the ultraviolet region. This line is a member of the Lyman Series Balmer Series Hydrogen electrons in an excited state may not drop back directly to a ground state but may in fact fall into another lower energy level n=3 to n=2 or n=4 to n=2 When this happens a line in the visible range is detected These are called Balmer Lines Paschen Series Hydrogen electrons in an excited state may not drop back to energy level 2 either Paschen lines are found in the infrared area of the spectrum and these result from electrons dropping from an excited state back down to n=3 NOTE The line spectrum for each element is unique but the theory remains constant An electron receives energy and moves from ground state to excited state On the return journey it will emit energy (light) of a value equal to the difference between the energy level it reached and the energy level it falls to Orbitals 1923 Louis de Broglie says electrons behave like waves and particles The location of the electron within an atom was undefined until Heisenberg (1927)
Heisenbergs uncertainty Principle it is impossible to define with absolute precision, at the same time, both the position and the momentum of an electron. This uncertainty lead to the location of an electron being computated mathematically using probability Edwin Schrodinger developed a mathematical probability equation in 1926 Using a number of different mathematical functions it is found that a boundary may be drawn about an atoms nucleus as to where there is a high probability of finding an electron at any given point in time This region is called an atomic orbital There are a number of different types of orbital but all can hold 2 electrons N=1 has one orbital Its spherical in shape and can hold 2 electrons called the 1S orbital N=2 has four orbitals 2s also spherical holding 2e - 2p x, 2p y, 2p z also each holding 2e - All 3 are equal in energy but higher than 2s All 3 are dumb-bell shaped N=3 has nine orbitals 3s is spherical holding 2e - 3p x, 3p y, 3p z also each holding 2e - All 3 are equal in energy but higher than 3s All 3 are dumb-bell shaped There are five 3d orbitals equal in energy but higher than 3p Each hold 2 electrons and have a complicated shape N=4 has 16 orbitals 4s is spherical 3 x 4p 5 x 4d 7 x 4 f orbitals Electronic Configurations A group of orbitals that all have the same energy is called an energy sub level 2p sub level is made up of 2p x, 2p y, 2p z There are five 3d sublevels
Remember if they are on the same sub level they all contain the same energy but may have different capacities Aufbau Principle Electrons will occupy the lowest energy sublevel available Atomic radii The atomic Radius of an element is half the distance between the nuclei of two atoms of the elements that are joined together by a single covalent bond Values for atomic radii are given in nanometres (nm) Values determined by x-rays Size of atomic radii depends on 1. Nuclear charge 2. Screening effect of inner electrons 3. Number of energy levels occupied Note: atomic radius decreases on moving across a period and increases on going down a group Ionisation Energies Defining first ionisation energy
The first ionisation energy is the energy required to remove the most loosely held electron from an isolated atom of the element in its ground state In an atom with lots of electrons the energy required to remove an electron will depend on where it is (the energy level) Measured in kilojoules per mole The second ionisation energy refers to the removal of a second electron from the now positive ion Values generally increase on moving left to right across a period Exception is Beryllium and Nitrogen which are abnormally high The increase in value is due to the nuclear charge and the decreasing atomic radii which makes it increasingly more difficult to remove the electron Why the Beryllium aberration Due to the electronic configuration 1s 2 2s 2 vs. 1s 2 2s 2 2p x 1 (Be) (B) The presence of a lone electron in the 2p shell means it is easier to remove than the more stable Beryllium electron...the same is true for nitrogen The second ionisation energy refers to the removal of a second electron from the now positive ion in a mole of these ions and the third ionisation energy refers to the removal of a third electron from the remaining positive ion after the second has been removed