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IB Chemistry TOPIC #10: ATOMIC THEORY

IB Chemistry TOPIC #10: ATOMIC THEORY. Atomic Structure. Atomic Structure. Atoms are very small ~ 10 -10 meters All atoms are made up of three sub-atomic particles: protons, neutrons and electrons. The protons and neutrons form a small positively charged nucleus

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IB Chemistry TOPIC #10: ATOMIC THEORY

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  1. IB Chemistry TOPIC #10: ATOMIC THEORY

  2. Atomic Structure

  3. Atomic Structure • Atoms are very small ~ 10-10 meters • All atoms are made up of three sub-atomic particles: protons, neutrons and electrons • The protons and neutrons form a small positively charged nucleus • The electrons are in energy levels outside the nucleus

  4. Atomic Structure • The actual values of the masses and charges of the sub-atomic particles are shown below: • A meaningful way to consider the masses of the sub-atomic particles is to use relative masses

  5. Atomic Structure - Definitions • Atomic number (Z) is the number of protons in the nucleus of an atom. The number of protons equals the number of electrons in a neutral atom • N.B. No. of protons always equals the no. of electrons in any neutral atom of an element. • Mass number (A) is the sum of the number of protons and the number of neutrons in the nucleus of an atom. • So how can you work out the number of neutrons in an atom? No. of neutrons = Mass number – atomic number

  6. Atomic Structure - Example • So how can you work out the number of neutrons in an atom? • Example No. of neutrons = Mass number – atomic number No. of neutron = Mass No. – Atomic No. = 23 – 11 = 12

  7. Atomic Structure - Questions • What are the three sub atomic particles that make up the atom? • Draw a representation of the atom and labelling the sub-atomic particles. • Draw a table to show the relative masses and charges of the sub-atomic particles. • State the atomic number, mass number and number of neutrons of: a) carbon, b) oxygen and c) selenium. • Which neutral element contains 11 electrons and 12 neutrons?

  8. Atomic Structure - Questions 5. Copy and complete the following table:

  9. Summary Slide • All atomic masses are relative to the mass of carbon-12. • Eg one hydrogen atom weighs 1/12 the mass of a carbon-12 atom.

  10. 35 17 37 17 Cl Cl Isotopes • Isotopes are atoms of the same element with the same atomic number, but different mass numbers, i.e. they have different numbers of neutrons. Each atom of chlorine contains the following: 17 protons 17 electrons 18 neutrons 17 protons 17 electrons 20 neutrons The isotopes of chlorine are often referred to aschlorine-35andchlorine-37

  11. Isotopes • Isotopes of an element have the same chemical properties because they have the same number of electrons. When a chemical reaction takes place, it is the electrons that are involved in the reactions. • However isotopes of an element have the slightly different physical properties because they have different numbers of neutrons, hence different masses. • The isotopes of an element with fewer neutrons will have: • Lower masses • faster rate of diffusion • Lower densities • lower melting and boiling points

  12. Isotopes - Questions • Explain what isotopes using hydrogen as an example. • One isotope of the element chlorine, contains 20 neutrons. Which other element also contains 20 neutrons? • State the number of protons, electrons and neutrons in: a) one atom of carbon-12 b) one atom of carbon-14 c) one atom of uranium-235 d) one atom of uranium-238

  13. Mass Spectrometer • The mass spectrometer is an instrument used: • To measure the relative masses of isotopes • To find the relative abundance of the isotopes in a sample of an element When charged particles pass through a magnetic field, the particles are deflected by the magnetic field, and the amount of deflection depends upon the mass/charge ratio of the charged particle.

  14. Mass Spectrometer – 5 Stages • Once the sample of an element has been placed in the mass spectrometer, it undergoes five stages. • Vaporization – the sample has to be in gaseous form. If the sample is a solid or liquid, a heater is used to vaporise some of the sample. X (s) X (g) or X (l)  X (g)

  15. Mass Spectrometer – 5 Stages • Ionization – sample is bombarded by a stream of high-energy electrons from an electron gun, which ‘knock’ an electron from an atom. This produces a positive ion: X (g) X +(g) + e- • Acceleration – an electric field is used to accelerate the positive ions towards the magnetic field. The accelerated ions are focused and passed through a slit: this produces a narrow beam of ions.

  16. Mass Spectrometer – 5 Stages • Deflection – The accelerated ions are deflected into the magnetic field. The amount of deflection is greater when: • the mass of the positive ion is less • the charge on the positive ion is greater • the velocity of the positive ion is less • the strength of the magnetic field is greater

  17. Mass Spectrometer • If all the ions are travelling at the same velocity and carry the same charge, the amount of deflection in a given magnetic field depends upon the mass of the ion. • For a given magnetic field, only ions with a particular relative mass (m) to charge (z) ration – the m/z value – are deflected sufficiently to reach the detector.

  18. Mass Spectrometer • Detection – ions that reach the detector cause electrons to be released in an ion-current detector • The number of electrons released, hence the current produced is proportional to the number of ions striking the detector. • The detector is linked to an amplifier and then to a recorder: this converts the current into a peak which is shown in the mass spectrum.

  19. Atomic Structure – Mass Spectrometer • Name the five stages which the sample undergoes in the mass spectrometer and make brief notes of what you remember under each stage. • Complete Exercise 4, 5 and 6 in the handbook. Any incomplete work to be completed and handed in for next session.

  20. Atomic Structure – Mass Spectrometer • Isotopes of boron Ar of boron = (11 x 18.7) + (10 x 81.3) (18.7 + 81.3) = 205.7 + 813 100 = 1018.7 = 10.2 100

  21. Mass Spectrometer – Questions • A mass spec chart for a sample of neon shows that it contains: • 90.9% 20Ne • 0.17% 21Ne • 8.93% 22Ne Calculate the relative atomic mass of neon You must show all your working!

  22. (90.9 x 20) + (0.17 x 21) + (8.93 x 22) • 100 Mass Spectrometer – Questions • 90.9% 20Ne • 0.17% 21Ne • 8.93% 22Ne Ar= 20.18

  23. 52.3 23.6 22.6 1.5 m/e 204 206 207 208 Mass Spectrometer – Questions Calculate the relative atomic mass of lead You must show all your working!

  24. (1.5 x 204) + (23.6 x 206) + (22.6 x 207)+(52.3 x 208) • 100 • 306 + 4861.6 + 4678.2 + 10878.4 • 100 • 20724.2 • 100 Mass Spectrometer – Questions • 1.5% 204Pb • 23.6% 206Pb • 22.6% 207Pb • 52.3% 208Pb Ar= 207.24

  25. Energy Levels • Electrons go in shells or energy levels. The energy levels are called principle energy levels, 1 to 7. • The energy levels contain sub-levels. These sub-levels are assigned the letters, s, p, d, f

  26. Energy Levels • Each type of sub-level can hold a different maximum number of electron.

  27. Energy Levels • The energy of the sub-levels increases from s to p to d to f. The electrons fill up the lower energy sub-levels first. Looking at this table can you work out in what order the electrons fill the sub-levels?

  28. Energy Levels • Let’s take a look at the Periodic Table to see how this fits in.

  29. Energy level Number of electrons Sub-level Electronic Structure • So how do you write it? 1s2 Example For magnesium: 1s2, 2s2, 2p6, 3s2

  30. Electronic Structure • The electronic structure follows a pattern – the order of filling the sub-levels is 1s, 2s, 2p, 3s, 3p… • After this there is a break in the pattern, as that the 4s fills before 3d. • Taking a look at the table below can you work out why this is? • This is because the 4s • sub-level is of • lower energy than the • 3d sub-level.

  31. Electronic Structure • The order in this the energy levels are filled is called the Aufbau Principle. • Example (Sodium – 2, 8, 1)

  32. Electronic Structure • There are two exceptions to the Aufbau principle. • The electronic structures of chromium and copper do not follow the pattern – they are anomalous. • Chromium – 1s2, 2s2, 2p6, 3s2, 3p6, 3d5, 4s1 • Copper – 1s2, 2s2, 2p6, 3s2. 3p6, 3d10, 4s1 • Write the electronic configuration for the following elements: • hydrogen c) oxygen e) copper • carbon d) aluminium f) fluorine

  33. Electronic Structure – of ions • When an atom loses or gains electrons to form an ion, the electronic structure changes: • Positive ions: formed by the loss of e- 1s2 2s2 2p6 3s1 1s2 2s2 2p6 Na+ ion Na atom • Negative ions: formed by the gain of e- 1s2 2s2 2p4 1s2 2s2 2p5 O atom O- ion

  34. Electronic Structure – of transition metals • With the transition metals it is the 4s electrons that are lost first when they form ions: • Titanium (Ti) - loss of 2 e- 1s2 2s2 2p6 3s2 3p6 3d2 1s2 2s2 2p6 3s2 3p6 3d24s2 Ti atom Ti2+ ion • Chromium (Cr) - loss of 3 e- 1s2 2s2 2p6 3s2 3p6 3d3 1s2 2s2 2p6 3s2 3p6 3d54s1 Cr atom Cr3+ ion

  35. Electronic Structure - Questions • Give the full electronic structure of the following positve ions: a) Mg2+ b) Ca2+ c) Al3+ • Give the full electronic structure of the negative ions: a) Cl- b) Br- c) P3-

  36. Electronic Structure - Questions • Copy and complete the following table:

  37. Orbitals • The energy sub levels are made up of orbitals, each which can hold a maximum of 2 electrons. • Different sub-levels have different number of orbitals:

  38. 1s 2s Orbitals • The orbitals in different sub-levels have different shapes: • s orbitals • p orbitals

  39. 2p 2s 1s Orbitals • Within a sub-level, the electrons occupy orbitals as unpaired electrons rather than paired electrons. (This is known as Hund’s Rule). • We use boxes to represent orbitals:     Electronic structure of carbon, 1s2, 2s2, 2p2  

  40. 2p 2s   1s     Orbitals • The arrows represent the electrons in the orbitals. • The direction of arrows indiactes the spin of the electron. • Paired electrons will have opposite spin, as this reduces the mutual repulsion between the paired electrons. Electronic structure of carbon, 1s2, 2s2, 2p2

  41. 2p 2s 1s Orbitals • Using boxes to represent orbitals, give the full electronic structure of the following atoms: a) lithium b) fluorine c) potassium d) nitrogen e) oxygen

  42. 2p 2s 1s Orbitals • Using boxes to represent orbitals, give the full electronic structure of the following atoms: a) lithium b) fluorine c) potassium d) nitrogen e) oxygen Electronic structure of lithium: 1s2, 2s1   

  43. 2p     2s     1s Orbitals • Using boxes to represent orbitals, give the full electronic structure of the following atoms: a) lithium b) fluorine c) potassium d) nitrogen e) oxygen Electronic structure of fluorine: 1s2, 2s2, 2p5

  44. 4s  3p       3s   2p       2s   1s   Orbitals • Using boxes to represent orbitals, give the full electronic structure of the following atoms: a) lithium b) fluorine c) potassium d) nitrogen e) oxygen Electronic structure of potassium: 1s2, 2s2, 2p6, 3s2, 3p6, 4s1

  45. 2p    2s   1s   Orbitals • Using boxes to represent orbitals, give the full electronic structure of the following atoms: a) lithium b) fluorine c) potassium d) nitrogen e) oxygen Electronic structure of nitrogen: 1s2, 2s2, 2p3

  46. 2p     2s     1s Orbitals • Using boxes to represent orbitals, give the full electronic structure of the following atoms: a) lithium b) fluorine c) potassium d) nitrogen e) oxygen Electronic structure of oxygen: 1s2, 2s2, 2p4

  47. Ionization Energy • Ionization of an atom involves the loss of an electron to form a positive ion. • The first ionization energy is defined as the energy required to remove one mole of electrons from one mole of atoms of a gaseous element. • The first ionization energy of an atom can be represented by the following general equation: X(g) X+ + e-ΔH > 0 • Since all ionizations require energy, they are endothermic processes and have a positive enthalpy change (ΔH) value.

  48. Ionization Energy • The value of the first ionization energy depends upon two main factors: • The size of the nuclear charge • The energy of the electron that has been removed(this depends upon its distance from the nucleus)

  49. + + Ionization Energy • As the size of the nuclear charge increases the force of the attraction between the negatively charged electrons and the positively charged nucleus increases. Small nuclear charge  Large nuclear charge    Small force of attraction  Large force of attraction  Smaller ionization energy Greater ionization energy

  50. + + Ionization energy • As the energy of the electron increases, the electron is farther away from the nucleus. As a result the force of attraction between the nucleus and the electron decreases. Electrons further away from positive nucleus  Electrons closer to positive nucleus  Large force of attraction  Small forceof attraction  Greater ionization energy Smaller ionization energy

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