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IB CHEMISTRY HL 1 UNIT 3 PERIODICITY

IB CHEMISTRY HL 1 UNIT 3 PERIODICITY. 11th IB t grade opics 3 and 13. 3.1 The Periodic Table. 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number. The history of the periodic table will not be assessed.

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IB CHEMISTRY HL 1 UNIT 3 PERIODICITY

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  1. IB CHEMISTRY HL 1UNIT 3 PERIODICITY 11th IB t grade opics 3 and 13

  2. 3.1 The Periodic Table • 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number. The history of the periodic table will not be assessed. • 3.1.2 Distinguish between the terms group and period. • 3.1.3 Apply the relationship between the electron arrangement of elements and their position in the periodic table up to Z = 20. • 3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for and element and its position in the periodic table.

  3. The Periodic Table of Elements

  4. The Periodic Table • The columns are called groups. The group number gives the number of electrons in the valence shell. • The rows are called periods and these are labeled 1-7. The period number gives the number of occupied electron shells. • In the IB data booklet, the representative groups in the Periodic Table are numbered from1 to 7 and the last column is labeled as “0”.

  5. The Periodic Table • We can use the electron configuration to split up the valence electrons into sub-levels. Example: C is [He]2s22p2. Note that valence electrons are in the same main energy level.

  6. 3.2 Physical Properties • 3.2.1 Define the first ionization energy and electronegativity. • 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals (Li  Cs) and the halogens (F  I). Explanation for the first four trends should be given in terms of the balance between the attraction of the nucleus for the electrons and the repulsion between electrons. Explanations based on effective nuclear charge are not required. • 3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies and electronegativities for elements across period 3. • 3.2.4 Compare the relative electronegativity value of two or more elements based on their positions in the periodic table.

  7. Effective Nuclear Charge • In any atom the nucleus exerts an attractive force on the electrons. • Across a period the number of protons in the nucleus steadily increases. The effective charge increases with the nuclear charge as there is no change in the number of inner electrons. • The effective nuclear charge experienced by an atom’s outer electrons increases with the group number of the element. • It increases across a period but remains approximately the same down a group.

  8. Core Radius (pm) Z Zeff Na 11 10 1 186 12 10 2 160 Mg Al 13 10 3 143 Si 14 10 4 132 Effective nuclear charge (Zeff) is the “positive charge” felt by an electron. Zeff = Z - s 0 < s < Z (s = shielding constant) Zeff Z – number of inner or core electrons

  9. Atomic Radius • The electron cloud does not have a sharp boundary so atomic radius is usually measured as half the distance between two neighboring nuclei:

  10. Atomic Radii covalent radius metallic radius

  11. Trends in Atomic Radii • Atomic radii increase down a group. • Atomic radii decrease across a period. • Going down a group there are more electron shells so the atomic and ionic radii increase. The effective nuclear charge remains about constant. • Across period attraction between the nucleus and the outer electrons increases as the nuclear charge increases so electrons are pulled in more and atomic and ionic radii decrease.

  12. Trends in Atomic and Ionic Radii

  13. Trends in Ionic Radii • Positive ions are smaller than their parent atoms. To form a positive ions the outer shell is lost ex. Na is 2, 8, 1 whereas Na+ is 2, 8. • Negative ions are larger than their parent atoms. To form a negative ions electrons are added in the outer shell ex. Cl is 2, 8, 7 and Cl- is 2, 8, 8. There is increased electron-electron repulsion in the outer shell so they move farther apart and increase the radius of the outer shell.

  14. Trends in Ionic Radii • The ionic radii decrease from groups 1 to 4 for POSITIVE ions. The ions Na+, Mg2+, Al3+ and Si4+ all have the same electron arrangement: 2, 8. The decrease in ionic radius is due to the increase in nuclear charge with atomic number across the period. The increased attraction between the nucleus and the electrons pulls the outer shell closer to the nucleus. • The ionic radii decrease from groups 4 to 7 for the NEGATIVE ions. The ions Si4-, P3-, S2- and Cl- have the same electron arrangement 2, 8, 8. The decrease in ionic radius is due to the increase in nuclear charge across the period.

  15. Trends in Ionic Radii • The positive ions are smaller than the negative ions, as the former have only two occupied electron shells and the latter have three. This explains the big difference between the ionic radii of the Si4+ and Si4- ions and the discontinuity in the middle of the table. • The ionic radii increase down a group as the number of electron shells increases.

  16. Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed.

  17. Isoelectronic: have the same number of electrons, and hence the same ground-state electron configuration Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne] O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne] Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne

  18. Ionization Energy • The first ionization energy of an element is the energy required to remove one mole of electrons from one mole of gaseous atoms. • Ionization energies increase across a period. Number of protons increases across period 3 so effective nuclear charge increases and ionization energy increase with it. • Ionization energies decrease down a group. Down a group electrons are further from nucleus so ionization energy decreases.

  19. Increasing First Ionization Energy Increasing First Ionization Energy General Trends in First Ionization Energies

  20. Trends in First Ionization Energies

  21. Trends in Ionization Energies • There are some small exceptions to the increasing trend across a period: • Ionization energy for a p sub-shell is lower than for an s sub-shell. This is because p orbitals are slightly higher in energy than s orbitals (in the same period). • There is also a decrease from the 5th element to the sixth as the p sub-shells start to be doubly filled. • It is easier to remove the 6th electron as it is repelled by its partner whereas the 5th electron is not paired so it takes more energy to remove it.

  22. Trends in Ionization Energies • Down a group ionization energy decreases as the outer electron is further from the pull of the nucleus. • Successive ionization energies for one element increase (but not smoothly) due to increased effective nuclear charge. • When electrons are removed from a new subshell there is a further increase in ionization energy.

  23. Electronegativity • Electronegativity is the ability of an atom to attract electrons in a covalent bond. • Electronegativity is related to ionization energy but is specific to BONDING electrons. • Electronegativity increases from left to right across a period owing to the increase in nuclear charge, resulting in an increased attraction between the nucleus and the bond electrons. • Electronegativity decreases down a group. The bond electrons are furthest from the nucleus and so there is reduced attraction.

  24. Electronegativity • Maximum value is 4.0 which Fluorine has. • Minimum value is 0.7 which Francium has.

  25. The Electronegativities of Common Elements

  26. Melting Points

  27. Melting Points • Melting points of alkali metals (group 1) decrease down the group as the metallic bonds weaken – valence electrons are further from the nucleus so the attraction between the delocalized electrons and the positive ions decreases. • Melting points of halogens (group 7) increase down a group as van der Waals forces increase with molar mass. The halogens all exist as diatomic molecules in their standard elemental form.

  28. Melting Points • Melting points will increase with stronger bonding and intermolecular forces. It is a measure of the difference in forces between the solid and liquid states. • Boiling point is a measure of the absolute size of these forces.

  29. Melting points across a period • Across a period the bonding changes from metallic (strong) to giant covalent (very strong) to van der Waals forces between molecules (weak). • Melting points generally increase across a period and reach a maximum at group 4. • The melting points increase and then decrease accordingly with the changes in strength of bonding. • The bonding changes from metallic (Na, Mg and Al) to giant covalent (Si) to weak van der Waals forces between molecules (P4, S8 and Cl2) and single atoms (Ar). All the period 3 elements are solids at room temperature except chlorine and argon.

  30. Summary of Trends across Period 3

  31. 3.3 Chemical Properties • 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group. The following reactions should be covered: Alkali metals (Li, Na and K) with water; Alkali metals (Li, Na and K) with halogens (Cl2, Br2, I2); Halogens (Cl2, Br2, I2) with halide ions (Cl-, Br-, I-).

  32. Chemical Properties • Chemical properties of an element are largely dependent on the number of electrons in the outer shell. • This means that groups tend to have similar chemical properties - they react in a similar way.

  33. The noble gases, group 0 • These are the least reactive elements. • They are monatomic – exist as single atoms. • They are colorless gases. • They have complete outer shells of electrons so have the highest ionization energies for each period. • Other elements tend to react to attain the electron configuration of the noble gases. • Compounds of xenon, krypton and argon have been made but it requires special conditions to create these.

  34. The Alkali Metals, group 1 • Physical properties: • Good conductors of electricity due to delocalized valence electrons • Low densities • Soft • Grey shiny surfaces when freshly cut • Chemical properties: • Very reactive due to single valence electron that is lost easily • Always form 1+ ions and combine easily with non-metals such as oxygen and halogens. • Ex. 2Na(s) + Cl2(g)  2NaCl(s)

  35. The Alkali Metals, group 1 • Reactivity increases down the group as the valence electron is further from the attraction of the nucleus and ionization energy decreases. • All alkali metals react vigorously with water to form a metal hydroxide solution (basic) and hydrogen gas • 2Na(s) + 2H2O(l)  2Na+(aq) + 2OH-(aq) + H2(g) • All alkali metals tarnish quickly in air so they lose their shiny surface. They are stored under oil to prevent this.

  36. Alkali Metals Stored Under Oil

  37. Alkali Metal + Water • Lithium floats and reacts slowly. It releases hydrogen but keeps its shape. • Sodium reacts with a vigorous release of hydrogen. The heat produced is sufficient to melt the unreacted metal, which moves around on the surface of the water. • Potassium reacts even more vigorously to produce sufficient heat to ignite the hydrogen produced. It produces a lilac colored flame and moves excitedly on the water surface.

  38. The Halogens, group 7 • F, Cl, Br and I are very reactive non-metals in group 7. • All require one electron to complete their valence shell. • All exist as diatomic molecules joined by covalent bonds ex. F2, Cl2, Br2, I2 • Van der Waals forces between the molecules increase down the group with molar mass. • They are all quite electronegative with F being the most electronegative element (smallest atomic radius).

  39. The Halogens, group 7 • Physical properties: They are colored They show a gradual change from gases (F2 and Cl2) to liquid (Br2) to solids (I2 and At2). • Chemical Properties: Very reactive non-metals. Reactivity decreases down the group. They form ionic compounds with metals or covalent compounds with non-metals.

  40. The Halogens, all toxic!

  41. The Halogens • At room temperature, F (pale yellow) and Cl (yellow-green) are gases, Br is a red-brown liquid and I is a solid that forms a black-purple vapor on heating, brown solution in water and purple solution in non-polar solvents. • Gain an electron easily to form Hal- • Ease of gaining an electron (and reactivity) decreases down the group as electrons are further from nucleus. • Slightly soluble in water as non-polar.

  42. Halogen + Alkali Metal • Halogens react easily with alkali metals to form ionic halides. • One electron is transferred from the alkali metal to the halogen so that the alkali metal forms a 1+ ion and the halogen forms a 1- ion. • These oppositely charged ions are strongly attracted to each other and form a strong ionic bond. • The most vigorous reaction will occur between the elements which are furthest apart in the periodic table: francium at the bottom of group 1 and fluorine at the top of group 7.

  43. Reactions of Halogens • Ex. 2Fr(s) + F2(g)  2FrF(s) • The relative reactivity of the halogens can be seen by combining a halogen element with a metal halide: • 2KBr(s) + Cl2(aq)  2KCl(aq) + Br2(aq) • Chlorine is more reactive than bromine so it can displace bromine from the compound. The net ionic equation could also show this: • 2Br-(aq) + Cl2(aq)  2Cl-(aq) + Br2(aq) • The reverse reaction would not occur as bromine is less reactive and cannot displace chlorine.

  44. Halogen + Halide • If the Cl2 is reacted with either the Br- or I- ions then Br2 or I2 will be formed, respectively. • If this is done in aqueous solution then with both Br2 and I2 an orange-brown color will appear from an originally colorless solution. • The halogens can be distinguished more clearly in non-polar solvents where they have the following colors: chlorine is a pale green, bromine is orange and iodine is violet.

  45. CL2, BR2 AND I2 IN CYCLOHEXANE

  46. Silver Halides • Halogens form insoluble salts with silver and lead. • Common test for halide ions is to add nitric acid followed by aqueous silver nitrate. • A precipitate confirms presence of halide: • AgCl is white but darkens in sunlight • AgBr is cream • AgI is pale yellow • AgF is soluble so this test wouldn’t work for F-.

  47. AgI , AgBr, AgCl, AgF

  48. SUMMARY OF Ag+ + Hal-

  49. 13.1 Trends across period 3 • 3.3.2 Discuss the changes in nature from ionic to covalent and from basic to acidic of the oxides across period 3. Equations are required for the reactions of Na2O, MgO, P4O10 and SO3 with water. • 13.1.1 Explain the physical states (under standard conditions) and electrical conductivity (in the molten state) of the chlorides and oxides of the elements in period 3 in terms of their bonding and structure. Include the following oxides: Na2O, MgO, Al2O3, SiO2, P4O6 and P4 O10; and the following chlorides: NaCl, MgCl2, Al2Cl6, SiCl4, PCl3, PCl5 and Cl2.

  50. Group 1 and 2 Oxides are BASIC • Across period 3 the nature of the elements changes. • Na and Mg form cations so they bond with O2- to form ionic oxides. • The oxide ion can bond with H+ ions so they act as bases dissolving in water to give alkaline solutions. • Na2O(s) + H2O(l)  2Na+(aq) + 2OH-(aq) • They will also neutralize acids to produce salt and water. • MgO(s) + 2HCl(aq)  Mg2+(aq) + 2Cl-(aq)

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