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Unit 2-8 Periodic trends

Chapter 7. Unit 2-8 Periodic trends. Development of Periodic Table. Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped. Development of Periodic Table.

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Unit 2-8 Periodic trends

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  1. Chapter 7 Unit 2-8 Periodic trends

  2. Development of Periodic Table Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped.

  3. Development of Periodic Table Mendeleev, for instance, predicted the discovery of germanium (which he called eka-silicon) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon.

  4. Periodic Trends • In this chapter, we will rationalize observed trends in • Sizes of atoms and ions. • Ionization energy. • Electron affinity.

  5. Effective Nuclear Charge (Zeff) • In a many-electron atom, electrons are both attracted to the nucleus (p#) and repelled by other electrons (core e#). • Shielding effect: Electrons within level and from previous levels block the effects of the (+) nucleus • The nuclear charge that an electron experiences depends on both factors.

  6. Effective Nuclear Charge (Zeff) The effective nuclear charge, Zeff, the actual attraction experienced by the valence electron, is found this way: Zeff = Z−S where Z is the atomic number and S is a screening constant, usually close to the number of core electrons.

  7. Sizes of Atoms • Radius of a free atom / ion = distance from nucleus to outermost electron shell. • The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

  8. Atomic Size • Influenced by three factors, influences go down as the listed sequence: 1. Energy Level (n number) • Higher energy level is further away. 2. Effective nuclear charge : Zeff = Z - S • More charge pulls electrons in closer. 3. outermost e- amount. • Electrons repel each other. More e-, the bigger size.

  9. Sizes of Atoms Bonding atomic radius tends to… …decrease from left to right across a row due to increasing Zeff. but transitional metals in the same row have similar radius. …increase from top to bottom of a column due to increasing value of n

  10. Sizes of Ions • Ionic size depends upon: • Energy level (n#) • Effective nuclear charge Zeff. • Outermost e#.

  11. Sizes of Ions • Cations are smaller than their parent atoms. • The outermost electron is removed and shell# are less.

  12. Sizes of Ions • Anions are larger than their parent atoms. • Electrons are added and repulsions are increased.

  13. Sizes of Ions • Ions increase in size as you go down a column. • Due to increasing value of n.

  14. Sizes of Ions • In an isoelectronic series, ions have the same number of electrons. • Isoelectronic ionic size decreases with an increasing Zeff. • Who are isoelectronic? Li+, O2-, Mg2+ • Who is the biggest in radius in this isoelectronic series?

  15. Ionization Energy • Amount of energy required to remove an electron from the ground state of a gaseous atom or ion. • Resulting in increasing in + charge. • First(1st) ionization energy is that energy required to remove first electron. • Second(2nd) ionization energy is that energy required to remove second electron, etc.

  16. Ionization Energy • It requires more energy to remove each successive electron. • When all valence electrons have been removed, the ionization energy takes a quantum leap.

  17. Trends in First Ionization Energies • As one goes down a column, less energy is required to remove the first electron. • For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.

  18. Trends in First Ionization Energies • Generally, as one goes across a row, it gets harder to remove an electron. • As you go from left to right, Zeff increases.

  19. Trends in First Ionization Energies However, there are two apparent discontinuities in this trend.

  20. Trends in First Ionization Energies • The first occurs between Groups IIA and IIIA. • Electron removed from p-orbital rather than s-orbital • Electron farther from nucleus • Small amount of repulsion by s electrons.

  21. Trends in First Ionization Energies • The second occurs between Groups VA and VIA. • Electron removed comes from doubly occupied orbital. • Repulsion from other electron in orbital helps in its removal.

  22. e e C H Electronegativity Electronegativity is a measure of an atom’s attraction for the shared pair of electrons in a bond Which atom would have a greater attraction for the electrons in this bond and why?

  23. Electronegativities are used to evaluate the polarity of the covalent bond and the oxidation states of the atoms in the covalent compound. • Looking across a row or down a group of the periodic table we can see a trend in values. • We can explain these trends by applying the same reasoning used for ionization energies. • EN has the same trend as IE. Electronegativities

  24. IncreasingElectronegativity Looking across a period F C B Li Be N O 2.0 2.5 3.0 3.5 4.0 1.0 1.5 What are the electronegativities of these elements? Across a period electronegativity increases The charge in the nucleus increases across a period. Greater Zeff= Greater attraction for bonding electrons

  25. F DecreasingElectronegativity Cl Br I Looking down a group 4.0 3.0 What are the electronegativities of these halogens? 2.8 2.6 Down a group electronegativity decreases Atoms have a bigger radius (more electron shells) The positive charge of the nucleus is further away from the bonding electrons and is shielded by the extra electron shells.

  26. Electron Affinity Electron affinity = Energy change accompanying addition of electron to gaseous atom: Cl + e− Cl− + heat N + e− + heat N− • Released E (exothermic): negative value • Absorbed E (endothermic): positive

  27. Trends in Electron Affinity • In general, electron affinity becomes more exothermic (-E) as you go from left to right across a row.

  28. Trends in Electron Affinity There are again, however, two discontinuities in this trend.

  29. Trends in Electron Affinity • The first occurs between Groups IA and IIA. • Added electron must go in p-orbital, not s-orbital. • Electron is farther from nucleus and feels repulsion from s-electrons.

  30. Trends in Electron Affinity • The second occurs between Groups IVA and VA. • Group VA has no empty orbitals. • Extra electron must go into occupied orbital, creating repulsion.

  31. Homework • Based on the PES bellow, explain which element requires less 1st ionization energy. Why? • Why C peak y has higher energy than B peak y? • Page 254 7.5, 7.15, 7.17, 7.21, 7.23, 7.25, 7.29, 7.62

  32. Properties of Metal, Nonmetals,and Metalloids

  33. Metals versus Nonmetals Differences between metals and nonmetals tend to revolve around these properties.

  34. Metals versus Nonmetals • Metals tend to form cations when reacting with nonmetal. • Nonmetals tend to form anions when reacting with metal.

  35. Metals Tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.

  36. Metals • Compounds formed between metals and nonmetals tend to be ionic. • Metal oxides tend to be basic. • If soluble, combine H2O to form base. • Neutrolization rxn with acid.

  37. Nonmetals • Dull, brittle substances that are poor conductors of heat and electricity. • Tend to gain electrons in reactions with metals to acquire noble gas configuration.

  38. Nonmetals • Substances containing only nonmetals are molecular compounds. • Most nonmetal oxides are acidic. • Combine with H2O to form oxyacid. If forming weak acid, reversible rxn. • Neutralization rxn with strong base.

  39. Metalloids • Have some characteristics of metals, some of nonmetals. • For instance, silicon looks shiny, but is brittle and semi-conductor. • Most of metalloids are semi-conductors.

  40. Group Trends

  41. Alkali Metals • Soft, metallic solids. • Name comes from Arabic word for ashes.

  42. Alkali Metals • Found only as compounds in nature. • Have low densities and melting points. • Also have low ionization energies.

  43. Alkali Metals Their reactions with water are famously exothermic.

  44. Alkali Metals • Alkali metals (except Li) react with oxygen to form peroxides. • K, Rb, and Cs also form superoxides: K + O2 KO2 • Produce bright colors when placed in flame.

  45. Alkaline Earth Metals • Have higher densities and melting points than alkali metals. • Have low ionization energies, but not as low as alkali metals.

  46. Alkaline Earth Metals • Bedoes not react with water, Mg reacts only with steam, but others react readily with water. • Reactivity tends to increase as go down group.

  47. Group 6A • Oxygen, sulfur, and selenium are nonmetals. • Tellurium is a metalloid. • The radioactive polonium is a metal.

  48. Oxygen • Two allotropes: • O2 • O3, ozone • Three anions (3 oxidation states): • O2−, oxide • O22−, peroxide • O21−, superoxide • Tends to take electrons from other elements (oxidation reagent)

  49. Sulfur • Weaker oxidizing agent than oxygen. • Most stable allotrope is S8, a ringed molecule.

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