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Topic 3: Periodicity

Understand the arrangement of elements based on atomic number, group, and period. Learn about physical properties like ionization energy and atomic size. Explore trends in the periodic table and how they affect element properties.

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Topic 3: Periodicity

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  1. Topic 3: Periodicity 3.1 The periodic table 3.1.1      Describe the arrangement of elements in the periodic table in order of increasing atomic number 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 highest occupied energy level for an element and its position in the periodic table.

  2. Groups: vertical columns (18) • Have similar properties because have same number of electrons in outer shell • Periods: horizontal row (7) • Family Names: • Group 1: alkali metals • Group 2: alkaline earth metals • Group 17: halogens • Group 18: noble gases • Group 3-12: Transition metals • Groups 1,2, 13-18: representative elements

  3. 3.2 Physical properties3.2.1      Define the terms first ionization energy and electronegativity3.2.2      Describe and explain the trends in atomic radii, ionic radii, first ionization energy, electronegativities and melting points for alkali metals (Li  Cs) and the halogens (F  I).3.2.3      Describe and explain the trends in atomic radii, ionic radii, first ionization energy, and electronegativities for elements across period3.2.4      Compare the relative electronegative values of two or more elements based on their position on the periodic table.

  4. Atomic Size • The electron cloud doesn’t have a definite edge. • They get around this by measuring more than 1 atom at a time. • Summary: it is the volume that an atom takes up • http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/atomic4.swf

  5. Group trends H • As we go down a group (each atom has another energy level) the atoms get bigger, because more protons and neutrons in the nucleus Li Na K Rb

  6. Periodic Trends atomic radius decreases as you go from left to right across a period. • Why? Stronger attractive forces in atoms (as you go from left to right) between the opposite charges in the nucleus and electron cloud cause the atom to be 'sucked' together a little tighter. Remember filling up same energy level, little shielding occurring. Na Mg Al Si P S Cl Ar

  7. Ionic Size • Cations form by losing electrons. • Cations are smaller than the atom they come from. • Metals form cations. • Cations of representative elements have noble gas configuration.

  8. Ionic size • Anions form by gaining electrons. • Anions are bigger than the atom they come from. • Nonmetals form anions. • Anions of representative elements have noble gas configuration.

  9. Periodic Trends • Metals losing from outer energy level, more protons than electrons so more pull, causing it to be a smaller species. • Non metals gaining electrons in its outer energy level, but there are less protons than electrons in the nucleus, so there is less pull on the protons, so found further out making it larger. N-3 B+3 O-2 F-1 Li+1 C+4 Be+2

  10. Size of Isoelectronic ions • Positive ions have more protons so they are smaller. N-3 O-2 F-1 Ne Na+1 Al+3 Mg+2

  11. Electronegativity

  12. Electronegativity • The tendency for an atom to attract electrons to itself when it is chemically combined with another element. • How fair it shares. • Big electronegativity means it pulls the electron toward it. • Atoms with large negative electron affinity have larger electronegativity.

  13. Group Trend • The further down a group the farther the electron is away and the more electrons an atom has. • So as you go from fluorine to chlorine to bromine and so on down the periodic table, the electrons are further away from the nucleus and better shielded from the nuclear charge and thus not as attracted to the nucleus. For that reason the electronegativity decreases as you go down the periodic table.

  14. Period Trend • Electronegativity increases from left to right across a period • When the nuclear charge increases, so will the attraction that the atom has for electrons in its outermost energy level and that means the electronegativity will increase

  15. Period trend Electronegativity increases as you go from left to right across a period. • Why? Elements on the left of the period table have 1 -2 valence electrons and would rather give those few valence electrons away (to achieve the octet in a lower energy level) than grab another atom's electrons. As a result, they have low electronegativity. Elements on the right side of the period table only need a few electrons to complete the octet, so they have strong desire to grab another atom's electrons.

  16. Group Trend electronegativity decreases as you go down a group. • Why? Elements near the top of the period table have few electrons to begin with; every electron is a big deal. They have a stronger desire to acquire more electrons. Elements near the bottom of the chart have so many electrons that loosing or acquiring an electron is not as big a deal. • This is due to the shielding affect where electrons in lower energy levels shield the positive charge of the nucleus from outer electrons resulting in those outer electrons not being as tightly bound to the atom.

  17. Shielding • Shielded slightly from the pull of the nucleus by the electrons that are in the closer orbitals. • Look at this analogy to help understand

  18. Melting Points of Group 1

  19. Metallic bonding • Collective bond, not a single bond • Strong force of electromagnetic attraction between delocalized electrons (move freely). • This is sometimes described as "an array of positive ions in a sea of electrons

  20. Why does the melting point decrease going down the alkali metals family? • Atoms are larger and their outer electrons are held farther away from the positive nucleus. • The force of attraction between the metal ions and the sea of electrons thus gets weaker down the group. • Melting points decrease as less heat energy is needed to overcome this weakening force of attraction.

  21. Melting Points for halogens

  22. Why does melting point increase going down the halogens? • The halogens are diatomic molecules, so F2, Cl2, Br2, I2 • As the molecules get bigger there are more electrons that can cause more influential intermolecular attractions between molecules. • The stronger the I.A, the more difficult it will be to melt. (more energy needed to break the I.A)

  23. What are these I.A? van der Waals forces (London dispersion): • Electrons are mobile, and although in a diatomic molecule they should be shared equally, it is found that they temporarily move and form slightly positive end and negative end. • Now that one end is + and the other -, there can be intermolecular attractions between the opposite charges of the molecules

  24. van der Waals forces

  25. IB requires knowledge specifically for halogens. Check out this site for more detail. http://www.chemguide.co.uk/inorganic/group7/properties.html

  26. Period 3 melting point trends

  27. Explanation • M.P rise across the 3 metals because of the increasing strength of the metallic bonds. • Silicon has a giant covalent structure just like diamond which makes its structure remarkably strong and therefore takes more energy to break apart.

  28. The atoms in each of these molecules are held together by covalent bonds (except Ar) • They would have weak I.A affecting the amount of energy needed to melt them. • Ar has extremely weak forces of attraction between its atoms, so its easiest to melt.

  29. 3.3 Chemical properties 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group. 3.3.2 Discuss the changes in nature from ionic to covalent and from basic to acidic of the oxides across period 3

  30. Reactivityof alkali metals • Generally group 1 metals become more reactive as you go down a group. • The valence electron of group 1 are found further from the nucleus as you go down the group. • It is easier to remove an electron from francium than from lithium

  31. Alkali metal + water • Li(s) + H2O (l)  LiOH(aq) + H2 (g) (Li + and OH- in solution) • The metal reacts with water to form the hydroxide of the metal (strong base) and bubbles off hydrogen gas. • The larger the alkali metal, the more vigorous the reaction. Sometimes the H2 gas actually lights itself (exothermic reaction, releases heat) causing the H2 to burn.

  32. MUST KNOW! • Na (s) + H2O (l)  NaOH (aq)+ H2(g) • K (s) + H2O (l)  KOH (aq)+ H2(g)

  33. Alkali metals + halogens • 2Na (s) + Cl2(g)  2NaCl (s) • Halogens are good oxidizing agents, which means they cause electrons to be lost from another atom (the reducing agent) • Halogens are 1 electron from stable octet and will try to remove electrons from valence electrons of other metallic atoms.

  34. MUST KNOW! • 2K (s) + Br2(l)  2KBr (s) • 2Li (s) + I2(g)  2LiI (s)

  35. Halogens reacting with halides • Halogens want an electron and even will remove electrons from other soluble salts, we refer to as halides. • When a salt dissolves it forms both of its ions in solution. • Ex: NaCl (aq)  Na+(aq) and Cl- (aq) • So halides are easily available for reactions

  36. Done in aqueous systems • Chlorine is stronger OA (oxidizing agent) than bromine because its found higher on the periodic table, so Cl2 will remove the electron from Br-, making Cl- and Br2 • Cl2 (aq) + 2Br- 2Cl- + Br2 (aq) • Cl2 (aq) + 2I- 2Cl- + I2 (aq) • Br2 (aq) + 2I- 2Br- + I2 (aq)

  37. Properties of Metals • Shiny (lustre) • Good conductors of heat and electricity • Malleable and ductile (change shape and make wires) • Tend to lose electrons • Metal oxides form basic solutions in water (pH greater than 7)

  38. Properties of non-metals • Brittle • Poor conductors of heat and electricity • Tend to gain electrons • Non-metal oxides tend to be basic when dissolved in water (pH less than 7)

  39. Across Period 3: metallic to non-metallic oxides • Basic solution from metallic oxide. Na2O(s) + H2O (l)  2 NaOH (aq) MgO (s) +H2O (l)  Mg(OH)2 (aq) Hydroxides of group 1 and 2 generally considered strong. • Acidic solution from non-metallic oxide. SO3(g) + H2O (l)  H2SO4 (aq) P4O10 (s) + 6H2O (l)  4 H3PO4 (aq) Aqueous hydrogen involved with acidity

  40. Properties of metalloids • Based on chemical and physical properties • Tend to have semi-conductive properties and form amphoteric oxides. • Considered metalloids are: • Boron (B) • Silicon (Si) • Germanium (Ge) • Arsenic (As) • Antimony (Sb) • Tellurium (Te) • Polonium (Po

  41. Amphoteric • Behave as an acid or a base depending upon the reaction it is involved with. • Also called amphiprotic (donate or accept a proton, H+) • Aluminum’s oxide is amphoteric. • Al2O3(s) + 3HCl (aq)→ AlCl3 (aq)+ 3H2O (l) • Reacts with a strong acid to make a to make a salt with water. • Al2O3(s)+ NaOH (aq) → NaAl(OH)4 (aq) • Reacts with a strong base to form sodium aluminate

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