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Atomic Structure - Questions

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.

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Atomic Structure - Questions

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  1. 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?

  2. 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

  3. 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

  4. 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

  5. Isotopes – H/W • Complete Exercise 1, 2, and 3 in the handbook for next session. • Task: Find out the uses of isotopes in as much detail as possible. • N.B. Please make sure you understand and write in your own words – DO NOT COPY out of a text-book.

  6. 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.

  7. Mass Spectrometer – 5 Stages • Once the sample of an element has been placed in the mass spectrometer, it undergoes five stages. • Vaporisation – 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)

  8. Mass Spectrometer – 5 Stages • Ionisation – 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.

  9. 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

  10. 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.

  11. 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.

  12. Atomic Structure – Mass Spec • 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. • Card Sort Activity ???

  13. Atomic Structure – Mass Spec • Isotopes of boron Ar of boron = (10 x 18.7) + (11 x 81.3) (18.7 + 81.3) = 187 + 894.3 100 = 1081.3 = 10.8 100

  14. Mass Spectrometer – Questions • Complete Exercise 7  14

  15. Energy Levels • Electrons go in shells or energy levels. The energy levels are called principle energy levels, 1 to 4. • The energy levels contain sub-levels. These sub-levels are assigned the letters, s, p, d, f

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

  17. 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?

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

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

  20. 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.

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

  22. 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

  23. 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 2p6 O atom O- ion

  24. 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 3s1 3p6 3d2 1s2 2s2 2p6 3s1 3p6 3d24s2 Ti atom Ti2+ ion • Chromium (Cr) - loss of 3 e- 1s2 2s2 2p6 3s1 3p6 3d3 1s2 2s2 2p6 3s1 3p6 3d54s1 Cr atom Cr3+ ion

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

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

  27. 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:

  28. 1s 2s Orbitals • The orbitals in different sub-levels have different shapes: • s orbitals • p orbitals

  29. 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, sp2  

  30. 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 reduced the mutual repulsion between the paired electrons. Electronic structure of carbon, 1s2, 2s2, 2p2

  31. 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

  32. 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   

  33. 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

  34. 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

  35. 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

  36. 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

  37. Ionisation Energy • Ionisation of an atom involves the loss of an electron to form a positive ion. • The first ionisation energy is defined as the energy required to remove one electron from a gaseous electron. • The first ionisation energy of an atom can be represented by the following general equation: • X(g) X+ + e-ΔH +ve • Since all ionisations requires energy, they are endothermic processes and have a positive enthalpy change (ΔH) value.

  38. Ionisation Energy • The value of the first ionisation 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)

  39. + + Ionisation 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 ionisation energy Greater ionisation energy

  40. + + Ionisation 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 ionisation energy Smaller ionisation energy

  41. Ionisation energy - Questions • Write an equation to represent the first ionisation of: a) aluminium b) lithium c) sodium

  42. + + + + Trends across a Period • Going across a period, the size of the 1st ionisation energy shows a general increase. • This is because the electron comes from the same energy level, but the size of the nuclear charge increases. Going across a Period

  43. Trends across a Period (2 exceptions) • The first ionisation of Al is less than that of Mg, despite the increase in the nuclear charge. • The reason for this is that the outer electron removed from Al is in a higher sub-level: the electron removed from Al is a 3p electron, whereas that removed from Mg is a 3s.

  44. Sulphur Phosphorus 3p     3p    3s   3s   Trends across a Period (2 exceptions) • The first ionisation energy of S is less than that of P, despite the increase in the nuclear charge. • In both cases the electron removed is from the 3p sub-level. However the 3p electron removed from S is a paired electron, whereas the 3p electron removed from P is an unpaired electron. • When the electrons are paired the extra mutual repulsion results in less energy being required to remove an electron, hence a reduction in the ionisation energy.

  45. Trends across a Period - Questions • There is a break in this general trend going across a Period. • Look at the table below and point out where the break in the the trend is and try to give an explanation. Clue: which sub-level (s, p, d or f is the outer electron in?

  46. He Ne F Ar N Cl C P Be H O Mg Ca S Si B Li Na Al K Trends across a Period - Questions • Now take a look at the graph below: • Explain what the graph shows in as much detail as possible • There is one other break in the general pattern going across a Period. What is it and explain why that is.

  47. + + + + Trends down a Group Ionisation energy decreases going down a Group. Going down a Group in the Periodic Table, the electron removed during the first ionisation is from a higher energy level and hence it is further from the nucleus. The nuclear charge also increases, but the effect of the increased nuclear charge is reduced by the inner electrons which shield the outer electrons. Down the Group

  48. Ionisation energy - Questions • Explain why sodium has a higher first ionisation energy than potassium. • Explain why the first ionisation energy of boron is less than that of beryllium. • Why does helium have the highest first ionisation energy of all the elements? • Complete Tasks

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