Arrangement of Electrons in the Atom.docx

Arrangement of Electrons in the Atom Neils Bohr, a young Danish scientist, provided an insight into the arrangement of electrons in the atom. Bohr’s Study of Spectra - When white light is passed through a glass prism, the white light is broken up into an array of colours, a continuous spectrum. For...

Arrangement of Electrons in the Atom Neils Bohr, a young Danish scientist, provided an insight into the arrangement of electrons in the atom. Bohr’s Study of Spectra - When white light is passed through a glass prism, the white light is broken up into an array of colours, a continuous spectrum. For example, a rainbow. -Bohr passed the light from a Hydrogen Gas Discharge Tube through a prism. The spectrum observed consisted of some narrow, coloured lines. He called this an emission (line) spectrum (1) - Emission spectrums play a very important role in modern chemistry. Each element has its own emission spectrum. This can be used in the identification of elements. - A spectrometer is used to measure the wavelength of each band of light in a spectrum. - A spectroscope is used to view spectrums but cannot measure wavelengths. Mandatory Experiment Number 1 To carry out flame tests with salts of lithium, sodium, potassium, barium, strontium, and copper 1. Using a damp wooden splint, hold a sample of lithium salt in the blue flame of a Bunsen burner. - Repeat the procedure using the other salts. - Record the colour of the flame for each salt. Metal present Colour of flame Lithium Crimson Potassium Lilac Barium Green Strontium Red Copper Blue green Sodium Yellow The Bohr Theory Bohr surmised that an electron in an atom can only have a fixed amount of energy. This fixed amount of energy is called a quantum of energy. - Electrons move around the nucleus of in fixed paths called orbits (energy levels). Electrons in any one orbit have a fixed amount of energy. - Energy levels are represented by the letter ‘n’. The lowest energy level is n = 1. - The energy of the electron in a particular orbit is quantised. - If an electron is in any one energy level it neither gains nor loses energy. - Atoms normally exist in the ground state. This means that electrons have the lowest possible amount of energy because they occupy the lowest available energy levels. - When energy is provided to an atom in its ground state, some of this energy is absorbed and the electrons jump from lower energy levels to higher energy levels. They are now in an excited state. - The energy absorbed is equal to the difference in energy between the ground state and the excited state. **Definitions: An energy level is the fixed energy value that al electron in an atom may have. The ground state of an atom is one in which the electrons occupy the lowest available energy levels. The excited state of an atom is one in which the electrons occupy higher energy levels than those available in the ground state.** - Electrons in the excited state are unstable so they fall back down to the ground state after a short period of time - As the electron falls, the excess energy is given off as a photon of light which has a definite amount of energy. - The frequency of light emitted depends on the difference in energy between the two energy levels. E2 – E1 = hf where E2 is the excited state and E1 is the ground state. h is Planck’s constant and f is the frequency of emitted light. - Each frequency of light emitted appears as a coloured line on the spectrum. - The separate lines obtained in the spectrum show that the electron can only have specific values of energy. - The emission of visible light is due to electrons falling to the n=2 energy level. This is known as the Balmer Series of emissions lines. Atomic Absorption Spectrometry (AAS) Scientists studying spectra have found that atoms can absorb light. - If light is passed through a gaseous sample of an element, it is found that the light that emerges has certain wavelengths missing. This spectrum is called an Atomic Absorption Spectrum. - These wavelengths are absorbed by the element and correspond exactly to the wavelengths that would be detected if an emission spectrum of the element were recorded. - The atoms in the ground state absorb the same radiations as they emit in the excited state. (2) Uses of AAS The device used to measure absorption spectrums is called an Atomic Absorption Spectrometer. - Used in the analysis of water for metals such as lead, mercury, and cadmium. - Used for measuring the concentration of lead in blood samples. - Used in forensic science to test for gunshot residue on samples like skin and clothing. An Atomic Absorption Spectrometer works based on two principles, 1. Atoms of an element in the ground state absorb light of a particular wavelength that is characteristic of that element. 2. The amount of light absorbed is directly proportional to the concentration of the element present in the sample. Energy Sublevels Every energy level (except n=1) consists of a number of sublevels ** Definition: A sublevel is a subdivision of a main energy level and consists of one or more orbitals of the same energy** When looking at emission spectra, what may appear to be one line in the spectrum, is actually a number of lines very close together. Each energy level, except the first, are made up of a number of sublevels all of which are close in energy. The number of sublevels is the same as the value of n. For example, n = 2 main energy level has two sublevels, n = 3 has three sublevels etc. (3) Wave Nature of the Electron In 1924 a French scientist called Louis de Broglie suggested that all moving particles have a wave motion associated with them. It was discovered that electrons had a wave motion associated with them, this is referred to as wave-particle duality. De Broglie’s theory was treated mathematically by a German physicist called Werner Heisenberg. Heisenberg’s Uncertainty Principle states that it is impossible to measure the velocity and speed of an electron simultaneously This led to limitations surrounding Bohr’s theory. Limitation of Bohr’s theory - Worked perfectly for hydrogen but in atoms with more than one electron it failed to account for many of the lines in the emission spectra of these atoms. - Did not consider the wave motion of the electron. - Could not explain the splitting of certain lines in emission spectra and did not consider the presence of sublevels. Atomic Orbitals **Definition: An orbital is a region in space within which there is a high probability of finding an electron.** An ‘s’ orbital is spherical in shape. a ‘p’ orbital is shaped like a dumbbell Sublevel Number of Orbitals Shape Number of Electrons per orbital Total number of electrons in sublevel s 1 Spherical 2 2 p 3 Dumbbell 2 6 d 5 N/A 2 10 Exam Questions 2014 – HL – Section B – Question 4 4. Answer eight of the following items (a), (b), (c), etc. (a) What colour is observed in a flame test on a salt of (i) barium, - Green (ii) lithium? Crimson (b) Describe the structure of Thomson’s ‘plum pudding’ model of the atom. Mass of positively-charged material with electrons embedded in it (d) The scientist pictured on the right is Werner Heisenberg. State the famous principle, published in 1927, which bears his name. Heisenberg’s Uncertainty principle - Position and velocity of an electron cannot be known simultaneously 5. (a) Name the scientist whose work on energy levels in the hydrogen atom is depicted in the Google doodle reproduced on the right. Neils Bohr Distinguish between the terms energy level and atomic orbital. Energy level - Fixed energy of an electron in an atom Atomic Orbital - Region in space where there is a high probability of fi nding an electron Write the electron configuration (s, p) of an atom of silicon showing the distribution of electrons in atomic orbitals in the ground state. 1s2 2s2 2px2 2py2 2pz2 3s2 3px1 3py 1 Hence, state how many (i) main energy levels, 3 (ii) atomic orbitals, 8 are occupied in the silicon atom in its ground state. 2018 – Hl – Section B – Question 4 (a) What was the purpose of Millikan’s ‘oil drop’ experiments of 1908 to 1913?  To measure the magnitude of the charge on an electron (c) Write the electron configuration (s, p, etc) for an iron atom. 1s2 2s2 2p6 3s2 3p6 4s2 3d6 References Assignmentpoint.com Khanacademy.com Slideplayer.com