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  1. Atom Structure Cook IB Chemistry II

  2. The Atom • First Periodic Tables were much different than the one today • First Periodic table had 100 elements • Elements • Simplest form of matter and can’t be broken down into simpler componets • Atoms • Smallest unit of an element • There are 92 elements that occur naturally

  3. Discovery of the atom • J. J. Thompson • Discovered that different metals produce a stream of negatively charged particles when a high voltage is applied across 2 electrodes= electrons • These were the same regardless of the metal. • Atoms have no charge • Plum pudding model: Negative ions scattered

  4. Rutherford’s Model • Fired alpha particles at a piece of gold foil. So, he hypothesized that these particles should pass straight through or get struck in the positive sponge.

  5. The atom • Protons and neutrons are present in the nucleus of an atom. Electrons are in orbits or energy levels around the nucleus. • The relative masses and relative charges of the sub atomic particles are:

  6. Atomic Number • Atomic Number (Z)=number of protons. It is the fundamental characteristic of and element • Mass Number(A)=number of protons + neutrons • Isotopes: • Are atoms with the same atomic number, different mass number or the same number of protons, but different number of neutrons. • Number of Protons=Z • Number of Electrons=Z-q • Number of neutrons=A-Z

  7. Dalton’s Atomic Theory • All matter is composed of tiny indiviable particles called atoms • Atoms can’t be created or destroyed • Atoms of the same element are alike in every way • Atoms of different elements are different • Atoms can combine together in small number to form molecules. • What do atoms look like. Kind of like hard spheres

  8. Isotopes • • Isotopes differ in physical properties that depend on mass such as: • density, rate of diffusion, etc. This difference is very significant for the isotopes of hydrogen as deuterium has the twice the mass of the more abundant . As isotopes have the same electron arrangement they have the same chemical properties. • • Examples of the uses of radioisotopes: C-14 in radiocarbon dating, • CO-60 in radiotherapy and • I-131 and I-125 as medical tracers.

  9. Isotopes • Show the same chemical properties, as a difference in the number of neutrons makes no difference to how they react and so they occupy the same place in the periodic table • Chlorine exist as 2 isotopes: • 35Cl and 37Cl The average relative atomic mass of the isotopes is not 36 but 35.45. • 35Cl is the more abundant isotope, in a sample of 100 chlorine atoms, there are 75 atoms of 35Cl and 25 atoms of 37Cl. • How would you calculate the relative atomic mass?

  10. TO WORK IT OUT • To work it out first calculate the total mass of the hundred atoms. • (75 x 35) + (25 x 37) = 3550 • = 35.5 • The 2 isotopes are both atoms of chlorine with 17 protons and 17 electrons • 35Cl; number of neutrons: 35-17=18 • 37Cl; number of neutrons: 37-17=20

  11. Uses of Radioisotopes • The stability of a nucleus depends on the balance between the number of protons and neutrons. When a nucleus contains either too many or too few neutrons, it is: • Radioactive • And will change to a more stable nucleus by giving out radiation • There are several different forms of radiation based on ionization and penetration abilities: • Alpha • Beta • Gamma

  12. Uses • Carbon-14 dating • The most stable isotope of carbon is 12C: • Has 6 protons and 6 neutrons. • 14C has • 8 neutrons , which is too many to be stable. It can reduce the neutron to proton ratio when a neutron changes to a proton and an electron. • The protons stays in the nucleus but the electron is ejected from the atom as beta particles. • 146C  147N + o-1e

  13. Carbon 14 • The relative abundances of carbon-14 present in living plants is constant as the carbon continually replenishes from carbon present in CO2 in the atmosphere. • When organisms die no carbon 14 is absorbed and the levels carbon 14 fall due to decay. • As this process occurs at a regular rate, it can be used to date carbon containing materials. • The rate of decay is measured in half life • This is the time taken for half the atoms to decay • The carbon-14 to carbon-12 ratio falls by 50% every 5730 years after the death of an organism • This is what archeologist use to date objects.

  14. Colbalt 60 • Radiotherapy • Radiation Therapy • Is the treatment of cancer and other diseases with ionizing raditation. • Cancerous cells are abnormal cells which divide at rapid rates to produce tumors that invade surrounding tissue. • The treatment damages the genetic material inside the cell by knocking off electrons and making it impossible for the cell to grow • This therapy damages both cancer and normal cells, the normal cells are able to recover if the treatment is carefully controlled.

  15. Radiation Therpy • Can treat localized • Solid tumors • Cancers of • Skin • Tongue • Larynx • Brain • Breast • Unterine cervix • Blood • Leukemia • Colbalt 60 is commonly used as it emits very penetrating gamma radiation when their protons and neutrons change their positions in the nucleus.

  16. Iodine 31 • Radioisotopes have the same chemical properties as any other atom of the same element, and so they play the same role in the body • Their positions, unlike other isotopes can be monitored by detecting radiation levels making them suitable

  17. Radioactivity • Unstable atomic nuclei will spontaneously decompose to form nuclei with a higher stability. • The decomposition process is called radioactivity. • The energy and particles which are released during the decomposition process are called radiation. • When unstable nuclei decompose in nature, the process is referred to as natural radioactivity. When the unstable nuclei are prepared in the laboratory, the decomposition is called induced radioactivity

  18. Type of Radiation • Alpha Particles: • Emitted by nuclei with too many protons to be stable • Composed of 2 protons and 2 neutrons • Beta Particles • Emitted by nuclei with too many neutrons, are electron which have been ejected from the nucleus by neutron decay • Gamma Particles • Are form of electromagnetic radiation

  19. Mass Spectrometry • Mass spectrometry (MS) is an analytical technique that produces spectra (singular spectrum) of the masses of the atoms or molecules comprising a sample of material. • The spectra are used to determine the elemental or isotopic signature of a sample, the masses of particles and of molecules, and to elucidate the chemical structures of molecules, • such as peptides • and other chemical compounds. • Mass spectrometry works by ionizing chemical compounds to generate charged molecules or molecule fragments and measuring their mass-to-charge ratios.[1]

  20. Mass Spectrometry • Vaporization: • The sample is turned into a gas using an electrical heater • Ionization • The gas particles are bombarded with high-energy electrons which knock electrons which ionize them. Electrons are knocked off the particles leaving positive ions. • Acceleration • Positive ions are attracted to negatively charged plates. • The positive ions are accelerated by an electric field. • Deflection: • The positive ions paths are altered with a magnetic field at right angles of each other. The amount of deflection is proportional to the charge mass ratio. Ions with smaller mass are deflected more than heavier ones. Lighter ions have less momentum and are deflected more than heavier ions. For a given field, only ions with a particular mass/charge ratio will make it to the detector. • Detection • The magnetic field strength is slowly increased. This changes the mass charge ration of ions that can reach the detector. A mass spectrum is produced. Mass charge ratio is detected and a signal is sent to a recorder

  21. Relative Atomic mass • ) of an element is the average mass of an atom according to the relative abundances of its isotopes, on a scale where the mass of one atom of is 12 • For example for 3517Cl which has two isotopes (75 %) and 3717Cl(25 %).

  22. Electromagnetic • This type of radiation comes in different forms of different energy • All electromagnetic waves travel at the same speed (c) • These waves can be distinguished by their different wavelengths (λ)\ • Different colors of visible light have different wavelengths • Red light has a lower wavelength than blue • The number of waves that which pass a particular point in 1 sec is called: • Frequency (f)

  23. Parts of a wave • Wavelength • Practice formulas: • λ= units are meters (m) • F= units are Hz or • v=f x λ units are m/s

  24. Visible light • Forms only a small part of the elctromagnetic spectrum • Infrared waves have longer wavelengths than red light and ultraviolet waves have shorter wavelengths than violet.

  25. Electromagnetic spectrum

  26. Visible light spectrum

  27. Line Spectra • When white light is passed through hydrogen gas, an absorption spectrum is produced. This line spectrum with some colors of the continuous spectrum missing • See diagram on page 51 • Evidence of Bohr model • Hydrogen atoms absorb and emit energy. This picture of the atom was considered with the electrons orbiting the nucleus in a circular energy level. Niels Bohr proposed that an electron moves into orbit or higher energy level further away from the nucleus when an atom absorbed energy. • This is called the : Excited state • This is produced • It is unstable • Electrons soon fall back to lowest state=Ground State • The energy the electron gives out as it falls back into lower levels is called • Electromagnetic Radiation

  28. Photon • This energy is called • Packet of energy=photon • Photons are released for each electron transistion. • The energy of the photon of light emitted is equal to the energy change in the atom • ∆Eelectron=Ephoton • It is also related to the frequency of the radiation by planck’s equation • ∆Eelectron=hf • Planck’s Constant =h=6.63x10-34Js • You will use this equation to calculate the wavelength to break bonds. Page 7 of chemistry data booklet

  29. Planck’s constant • In 1900, Max Planck was working on the problem of how the radiation an object emits is related to its temperature. • He came up with a formula that agreed very closely with experimental data, but the formula only made sense if he assumed that the energy of a vibrating molecule was quantized--that is, it could only take on certain values. • The energy would have to be proportional to the frequency of vibration, and it seemed to come in little "chunks" of the frequency multiplied by a certain constant. • This constant came to be known as Planck's constant, or h, and it has the value

  30. Hydrogen spectrum • Hydrogen atoms gives out energy when an electron falls from a higher to a lower energy level. • Hydrogen produces visible light when the electron falls to the second energy level (n=2) • The transition from to the first energy level corresponds to a higher energy change and are in the ultraviolet region of the spectrum. • Infrared radiation is produced when an electron falls to the third energy level

  31. Hydrogen Spectrum • Looking at figure 2.13 page 53 HL and 45 SL show how the energy levels inside the atom. • The lines converge at higher energy levels • This is due the energy levels inside the atom are closer together. • When an electron is at its highest energy e=∞, it is no longer in the atom and the atom has been ionized. • The energy needed to remove an electron from the ground state is called • Ionization energy

  32. • The hydrogen spectrum: • Series Region Electron falls to • Lyman UV n = 1 • Balmer Visible n = 2 • Paschen IR n = 3– • • The ionization energy of hydrogen corresponds to the convergence limit of the Lyman series.

  33. Building Atoms using the Bohr Model • Atoms react based the arrangement of sub atomic particles. We can now explore the structure of the atoms beyond hydrogen.

  34. Energy Levels • Each energy level can hold a limited number of electrons. • Ground State • Electrons are placed in the lowest energy level first, and when this becomes complete, electrons move to the second energy level, and so on. • Helium has 42H, has 2 protons, 2 neutrons and 2 electrons. The protons and neutrons from the nucleus and the 2 electrons both occupy the lowest energy level.

  35. Electron arrangment

  36. Ionization of energy • • The first ionization energy is the minimum energy required to remove one mole of electrons from a mole of gaseous atoms to form a mole of univalent cations in the gaseous state. It is the enthalpy change for the reaction: • X (g)  X + (g) + e–. • When an atom becomes ionized it loses an electron or proton. e-

  37. Electron lost when it becomes positive

  38. Electron Configurations • Bohr model has limitations. It doesn’t explain levels after level 3. • More energy is needed to remove electrons at higher ionization energy. • More difficult to remove, so we have: -why we have sublevels See table on page 57 HL.

  39. Atomic Orbitals • we know that the first energy level is made up of 1s sub level. Due to Heisenberg Uncertainty principle we don’t know the position of the electron.. • So we just say its in an orbital • Atomic orbital is a region around the atomic nucleus in which there is a 90% probability of finding electron.

  40. S orbitals S orbitals at either level are spherical/circular 1s and 2s 2s are larger

  41. P orbitals • P sub levels contain 3 p atomic orbitals of equal energy. • They are dumbbell shape and are arranged at right angles

  42. D and F orbitals • D orbitals are made up of 5 sublevels • F orbitals are made up of 7 sublevels • See page 61 HL

  43. Pauli Exclusion Principle • No more than 2 electrons can occupy an one orbital, and if two electrons are in the same orbital they must spin in opposite directions

  44. Energy Levels

  45. Energy Levels • Aufbau Principle: • Orbitals with lower energy are filled before those with higher energy • Hunds Rule • Every orbital in a sub level is singly occupied with electrons of same spin before any one orbital is doubly occupied

  46. Aufbau table

  47. 3d and 4s confusion • Electrons fill low energy orbitals (closer to the nucleus) before they fill higher energy ones. Where there is a choice between orbitals of equal energy, they fill the orbitals singly as far as possible. • The diagram (not to scale) summarises the energies of the orbitals up to the 4p level. •

  48. Alpha Radiation • Alpha radiation is a heavy, very short-range particle and is actually an ejected helium nucleus. Some characteristics of alpha radiation are: • •Most alpha radiation is not able to penetrate human skin. • •Alpha-emitting materials can be harmful to humans if the materials are inhaled, swallowed, or absorbed through open wounds. • •A variety of instruments has been designed to measure alpha radiation. Special training in the use of these instruments is essential for making accurate measurements. • •A thin-window Geiger-Mueller (GM) probe can detect the presence of alpha radiation. • •Instruments cannot detect alpha radiation through even a thin layer of water, dust, paper, or other material, because alpha radiation is not penetrating. • •Alpha radiation travels only a short distance (a few inches) in air, but is not an external hazard. • •Alpha radiation is not able to penetrate clothing. • Examples of some alpha emitters: radium, radon, uranium, thorium.