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07/06/2014

07/06/2014. A is for ATOM. +. Evidence for nuclear atom using alpha particle scattering. Next. The Nuclear Atom. ATOMS. 07/06/2014. From the ancient Greeks through to the 19 th centaury, there has been the question ‘ what is matter is made from ?’

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07/06/2014

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  1. 07/06/2014 A is for ATOM + Evidence for nuclear atom using alpha particle scattering. Next

  2. The Nuclear Atom ATOMS 07/06/2014 From the ancient Greeks through to the 19th centaury, there has been the question ‘what is matter is made from?’ The idea of atoms was first proposed by Greek Philosopher Democritus in 530 B.C. The concept was that matter could only be split in half and half again until indivisible units were reached. History of the Atom - Democritus Back Next

  3. The Nuclear Atom 07/06/2014 In 1808, John Dalton (a teacher!) proposed the modern ATOMIC THEORY. It simply states that all elements are made up of atoms and an element is only made up from one type of atom. Dalton's view of tiny indivisible spheres remained unchallenged until the end of the 19th Century. History of the Atom – John Dalton Dalton’s Atom Back Next

  4. The Nuclear Atom 07/06/2014 In 1897, Joseph (JJ) Thomson (British Physicist) was experimenting with electrical currents through gases. The cathode rays he produced could be deflected or moved when in electromagnetic fields. Cathode rays were made up of tiny negatively charged particles – ELECTRONS. Discovering the Electron Back Next

  5. The Nuclear Atom 07/06/2014 From his evidence, Thomson proposed that atoms were made up of just tiny electrons. He accounted for the neutrality of atoms by the stating the electrons existed in a ‘soup of positive charge’. Sometimes referred to the plum-pudding model. Thomson’s Model of the Atom Back Next

  6. The Nuclear Atom 07/06/2014 Around the same time, Henri Becquerel discovered that some unstable elements gave off smaller particles – RADIOACTIVITY. Therefore atoms must be divisible and made up of smaller parts – SUBATOMIC PARTICLES. Marie and Pierre Curie and Ernest Rutherford confirmed this. Radioactivity Back Next

  7. DIG – The Dating Game 07/06/2014 Alpha Particle: Positively charge. Large in comparison. Essentially a Helium Nucleus (as proved by Rutherford)  Radiation Beta Particle: Negatively charged. Light. (later to be shown as electrons)   Gamma Rays: Neutrally charged. No mass – Energy. Back Next

  8. The Nuclear Atom 07/06/2014 Rutherford and his colleagues bombarded a thin foil of gold with a beam of alpha particles and then onto a fluorescent screen. Rutherford’s Experiment Small amounts were deflected. Fluorescent Screen 99.9% passed straight through unaffected. Thin Gold Foil Back Next

  9. The Nuclear Atom 07/06/2014 Why were alpha particles scattered? To explain back scattering Rutherford proposed the Nuclear Model of the Atom. Alpha Particle Scattering + Back Next

  10. The Nuclear Atom 07/06/2014 Alpha particles are positive. High speed alpha particle bullet travels through atom. The electrons have little effect since they are very light and the electrons in the pudding model are very spread out. Very little deflection. Does not support observations. Alpha Particle Scattering Plum-Pudding Model Back Next

  11. The Nuclear Atom 07/06/2014 What is observed is that alpha particles in some instances are strongly deflected. With electrons practically dismissed, the only electrostatic force available could be a positive charge somewhere within the atom. Alpha Particle Scattering Alpha Source An atom Back Next

  12. The Nuclear Atom 07/06/2014 Hesuggested that all of the atom’s positive charge, together with most of its mass, is concentrated in the centre. Alpha particles which travel close to the nucleus are strongly deflected. The degree of deflection depends on how close it approaches. Rutherford’s Model + Back Next

  13. The Nuclear Atom 07/06/2014 The nucleus must be very small in comparison to the atom. This will account for the vast majority making it through unaffected. Rutherford’s Model Back Next

  14. The Nuclear Atom 07/06/2014 In summary, He envisioned an atom that had a positively charged nucleus in the centre. The atom was mostly empty space. An he deemed it reasonable that electrons orbit this nucleus like planets orbit the Sun. Rutherford’s Nuclear Model of an Atom + Nuclear Model of an Atom Back Next

  15. The Nuclear Atom 07/06/2014 The model appeared flawless and convinced most of the scientific community. Rutherford and his colleagues (Hans Geiger and Ernest Marsden) were able to precisely predict the effects of: Alpha particle energyThickness of sampleDifferent metals However there was a problem… Rutherford’s Nuclear Model of an Atom + Nuclear Model of an Atom Back Next

  16. The Nuclear Atom 07/06/2014 As the electrons move in circles, they would lose energy. Losing energy would slow them down. Therefore they would be pulled into the positively charged nucleus. It has been calculated that a Rutherford atom would only exist for about 1 billionth of a second! The answer lies within QUANTUM MECHANICS – when things get really small!! The Problem… Back Next

  17. Rutherford’s model could not explain: • Why the electrons did not lose energy as they orbited. • What held the protons together in the nucleus. • The origins of emission spectra of gases could not be explained. +

  18. The Nuclear Atom 07/06/2014 The number of protons in a nucleus did not match the atomic weight of the atom. Therefore a third neutrally charged particle must exist! These he named NEUTRONS. James Chadwick’s Protons Alpha Radiation Neutron Released Beryllium Foil Back Next

  19. Line Spectra Gases absorb certain frequencies of light. Each gas absorbs a unique combination of frequencies – each frequency corresponding to a unique colour. So each gas has a unique set colours which is known at its “line spectra” – because they are unique they can be used to identify a gas – similar to fingerprints.

  20. Bohr’s Atom Bohr - brought the concept of quantization into atomic theory. Electrons could only move in certain specific orbits corresponding to specific amounts of energy. These ENERGY LEVELS radiated out from the nucleus with higher energies being further away. Electrons do not radiate energy in these orbits. Energy is gained or lost when they move between orbits. This model enabled Bohr to explain the hydrogen spectrum.

  21. Atomic Spectra A glass prism can be used to generate a colour ………………….. If this the light generated by a hot (glowing) gas is viewed through a prism specific colour lines are seen as AN ……………… ………….SPECTRUM. If light is shone through a cold sample of the same gas, the same specific colour lines are absent and appear as an ……………………….. LINE SPECTRUM.

  22. Atomic Spectra A glass prisim can be used to generate a colour spectrum. If this the light generated by a hot (glowing) gas is viewed through a prism specific colour lines are seen as AN EMISSION LINE SPECTRUM. If light is shone through a cold sample of the same gas, the same specific colour lines are absent and appear as an ABSORPTION LINE SPECTRUM.

  23. Absorption & Emission spectrum ………………………… • In absorption spectrum radiation is again absorbed by electrons being …………… to higher energy levels. • The same frequencies (colours) are again emitted when the excited electrons ……………………. to the ground state in an ………………………spectrum. ……………………… - - - -

  24. Absorption & Emission spectrum EMITTED LIGHT • In absorption spectrum radiation is again absorbed by electrons being excited to higher energy levels. • The same frequencies (colours) are again emitted when the excited electrons drop down to the ground state in an emission spectrum. ABSORBED LIGHT ABSORBED LIGHT - - - -

  25. Emission Spectrum Excited electrons dropping down from unstable energy levels……………………..in the form of light. The frequency (colour) of the radiation is directly related to the ……………………….. between the energy levels. Since each element has its own ……………………series of energy levels, each element also has its own unique series of …………………………… lines. The line spectrum can therefore be used to …………………each element much like a fingerprint.

  26. Emission Spectrum Excited electrons dropping down from unstable energy levelsradiate energy in the form of light. The frequency (colour) of the radiation is directly related to the energy gap between the energy levels. Since each element has its own unique series of energy levels, each element also has its own unique series of emission/absorption lines. The line spectrum can therefore be used to identify each element much like a fingerprint.

  27. THE NEUTRAL ATOM • The atom consists of a _____________________ ______________________________ surrounded by a __________________________. • Atomic NumberZ: ___________________ in the Nucleus = _________________ in a ______ atom. • Mass numberA - Number of ______ + ________ Notation ___________ Number (bigger) ZAX symbol ______________ Number (smaller)

  28. THE NEUTRAL ATOM • The atom consists of a nucleus containing protons and neutrons surrounded by a cloud of electrons. • Atomic NumberZ - Number of protons in the Nucleus = number of electrons in a neutral atom. • Mass numberA - Number of protons + neutrons. Notation Mass Number (bigger) ZAX symbol Atomic Number (smaller)

  29. Relative Masses • Relative atomic(Ar): The mass of the atomrelative to ________________________________________. (Number of times heavier than…) Eg: O - 16one atom of oxygen is ________________ than 1/12 of the mass of a C12 atom, Formula mass (Mr) - The _______________________ of the atoms in a molecule. Water H2Oone molecule of water has a relative mass of _____________________________ that is the molecular or formula mass of water. Mr(H2O) = 18 (Times heavier than…)

  30. Relative Masses • Relative atomic mass (Ar): The averagemass of an atom of an element relative to 1/12 of the mass of a C12 atom. (Number of times heavier than…) Eg: O - 16one atom of oxygen is 16 times heavier than 1/12 of the mass of a C12 atom. • Formula mass (Mr) - The sum of all the atomic masses of the atoms in a molecule. Water H2Oone molecule of water has a relative mass of (2x(1)+16) = 18 - that is the molecular or formula mass of water. Mr(H2O) = 18 (Times heavier than…)

  31. Relative Masses - examples Calculate the Formula masses of: • O2(oxygen gas) Mr(O2) = • Cl2 (chlorine gas) • NaCl (sodium chloride - table salt) • CaCO3(calcium carbonate) • (NH4)2Cr2O7(ammonium dichromate)

  32. Relative Masses - examples Calculate the Formula masses of: • O2(oxygen gas) Mr (O2) = 2x16 = 32 • Cl2 (chlorine gas) Mr (Cl2) = 2x35.5 = 71.0 • NaCl (sodium chloride - table salt) Mr (NaCl) = 23+35.5 = 58.5 • CaCO3(calcium carbonate) Mr (CaCO3) = 40.1+12+(3x16) = 100.1 • (NH4)2Cr2O7(ammonium dichromate) Mr ((NH4)2Cr2O7 ) = 2(14+4)+2(52)+7(16) =252

  33. e- e- e- e- e- e- e- e- e- e- e- e- Isotopes The two atoms below both belong to carbon but they are not identical – can you spot what is different? Isotopes Atoms of the same element which have different numbers of neutrons. Others – Boron 10 & 11, Hydrogen 1 & 2, Chlorine 35 & 37. Write notation and work out numbers of neutrons. 613C 612C

  34. Isotopes • Isotopes - Atoms of the same element which have different numbers of neutrons. Eg: 613C & 612C • 37Cl (25%) & 35Cl (75%) - ratio 1:3 Av Ar(Cl) = (37x25)+(35x75) = 35.50 100 Or Av Ar(Cl) = (37x1)+(35x3) = 35.50 4 Relative atomic mass is (actually) the average mass of an atom of an elementrelative to1/12 of the mass of a carbon-twelve atom.

  35. Bohr’s Atom - problems • Only explain hydrogen spectrum. • Could not explain molecules (bonding of atoms) - formation or properties. • Why fixed orbits and no energy radiation in orbits. • At variance with Heisenberg’s uncertainty principle. Heisenberg: Not possible to know both the position and velocity of an electron at the same time with the same amount of accuracy.

  36. A Wave Model • De Broglie - light - ‘matter wave’ - theory. • Davisson & Germer - electron diffraction -proof of ‘matter waves’. • Shroedinger, Heisenberg et. al. - wave & quantum mechanical model. • Orbits ---> Orbitals - standing electron waves - a region or space defining the standing wave pattern. Orbital: - a region where there is a high probability of finding an electron.

  37. Ionisation Energy (Ei) The ENERGY REQUIRED to REMOVE AN ELECTRON completely from an atom. Sodium ion Sodium atom Electronic structure Na: 1s22s22p63s1 Na+: 1s22s22p6 FIRST ionisation energy (Ei1): Energy required to remove OUTERMOST electron. M  M+ + 1e- SECOND ionisation energy (Ei2): Energy required to remove SECOND OUTERMOST electron. M+ M2+ + 1e-

  38. Successive Ionization Energies Analyse this graph in light of your knowledge of atomic & electronic structure. This graph can be used to provide ‘evidence’ for some of the features of modern atomic theory. What can be inferred about the electronic structure of the atom?

  39. Successive Ionization Energies FIRST ionisation energy (Ei1): Energy required to remove OUTERMOST electron. M  M+ + 1e- SECOND ionisation energy (Ei2): Energy required to remove SECOND OUTERMOST electron. M+ M2+ + 1e- Hard to remove close to nucleus Inner level Second energy level Outer (valence) level Easy to remove far from the nucleus This graph provides EVIDENCE for energy levels. What can be inferred about the electronic structure of the atom?

  40. Patterns in Energies FIRST ionisation energy (Ei1): Energy required to remove OUTERMOST electron. M  M+ + 1e- What does this graph tell us about the electronic structure of the atom?

  41. Patterns in Energies FIRST ionisation energy (Ei1): Energy required to remove OUTERMOST electron. M  M+ + 1e-

  42. First Ionisation Energies Which of the following electrons would be easier to remove? H’s electron would be removed. 1s 1 proton H

  43. Electron Structure Bohr Orbits - energy levels N = 4 N = 3 N = 2 N = 1 3s 2s

  44. Electron Structure Energy sub levels and orbitals Bohr Orbits - energy levels N = 4 N = 3 N = 2 N = 1 4p (3d orbitals) 4s 3p orbitals 3s 2p orbitals 2s 1s orbital

  45. Electron distribution Energy levels (1,2,3 etc) are divided up into sub-levels (s, p, d, f) each of which has a specific number of orbitals (s-1, p-3, d-5). Pauli: • Electrons occupy lowest vacant energy levels. • Max two electrons per orbital • spin paring occurs when two electrons sharing the same orbital.  N=2 second energy level N S N N=1 first energy level S

  46. 4s 3s 2s 1s 4s 3s 2s 1s 3p 2p 3p 2p Electron Structure

  47. 4s 3s 2s 1s 4s 3s 2s 1s 3p 2p 3p 2p Electron Structure

  48. 4s 3s 2s 1s 4s 3s 2s 1s 3p 2p 3p 2p Electron Structure

  49. 4s 3s 2s 1s 3p 2p Electron Structure Afbau Diagrams 1s He 1s2 H 1s1

  50. 4s 3s 2s 1s 4s 3s 2s 1s 3p 2p 3p 2p Electron Structure Be 1s2 2s2 Li 1s2 2s1

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