Periodic Properties of the Elements

1. Periodic Properties of the Elements Chapter 7

2. 7.1 Development of the Periodic Table • 1st developed by Dmitri Mendeleev (Russia) & Lothar Meyer (Germany) on the basis of the similarity in chemical and physical properties • Mendeleev … • started by organizing elements by increasing mass. • Recognized a repetition of pattern. • Placed elements by same column  same properties • Predicted correctly about the existence of new elements • Henry Moseley • established that each element has a unique atomic number, which added more order to the periodic table • Identified the atomic number with the # of protons in the nucleus of the atom & the # of electrons in the atom.

3. 7.2 Electron Shells and the Sizes of Atoms • Atoms aren’t hard spheres with well-defined shells of electrons • The edges of atoms are a bit “fuzzy” • The quantum mechanical model of the atom supports the notion of electron shells: certain distances from the nucleus at which there is a higher likelihood of finding an electron

4. Atomic Sizes • The size of an atom can be gauged by its bonding atomic radius, based on measurements of the distances separating atoms in their chemical combinations with other atoms • Measure the atomic radius from the center of the nucleus to the outermost electron. • Atom size increases going down a group. • Atomic size decreases going left to right across the period.

5. 7.3 Ionization Energy • Ionization energy – the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion • 1st ionization energy (I1) – The energy needed to remove the first electron from a neutral atom, forming a cation • 2nd ionization energy (I2) – the energy needed to remove the second electron • The greater the ionization energy, the harder it is to remove an electron

6. 7.3 Ionization Energy • HIGH ionization energy means the atom hold onto the electron tightly and a lot of energy is need to pull it off • LOW ionization energy means the atom holds onto the electron loosely so breaking it apart doesn’t require much energy

7. 7.3 Ionization Energy Periodic Trends in Ionization Energies • Ionization energy decreases as you move down a group. • Ionization energy increases as you move from left to right on the periodic table. • Representative elements show a larger range of values of I1 than do the transition metal elements

8. Ionization Energy 3-D

9. 7.4 Electron Affinities • Electron affinity – the energy change that occurs when an electron is added to as gaseous atom • A negative electron affinity = the anion is stable • A positive electron affinity = the anion is higher in energy than are the separated atom and electron. The anion is not stable and will not form

10. 7.4 Electron Affinities • If the electron affinity is negative, the atom releases energy. • Normally, non-metals have a more negative electron affinity than metals. The exception is the noble gases.

11. 7.4 Electron Affinities • Election affinities become more negative as we proceed from left to right • Halogens have the most negative electron affinities • The electron affinities of the noble gases are all positive since the added electron would have to occupy a new, higher-energy subshell • Electron affinity doesn’t change greatly as we move down a group. Electron affinity should become more positive (less energy released).

12. 7.5 Metals, Nonmetals, and Metalloids

13. Pg. 239 --Table 7.3 Characteristic Properties of Metals and Nonmetals

14. 7.5 Metals, Nonmetals, and Metalloids • Metallic Character - The tendency of an element to exhibit properties of metals • Metallic character generally increases going down a column and decreases going from left to right across a period

15. Metals • Metals conduct heat & electricity • They are malleable & ductile • Solids at room temp. except mercury(Hg) (it’s liquid) • Melt at very high temps • Have low ionization energies & are consequently oxidized (lose electrons) when they undergo chemical reaction. • Many transition metals have the ability to form more than one positive ion.

16. Chemical Reactions with Metals • metal oxide + water  metal hydroxide • Most metal oxides are known as basic oxides • Ex: Na2O (s) + H2O(l)  2NaOH (aq) • metal oxide + acid  salt + water • Ex: MgO (s) + 2HCl (aq)  MgCl2 (aq) + H20 (l)

17. Nonmetals • Not lustrous & generally are poor conductors of heat and electricity • Non-metals commonly gain enough electrons to fill their outer p sub-shell completely, giving a noble gas electron configuration. • Molecular substances - Compounds composed entirely of nonmetals • Ex: oxides, halides, and hydrides • Melting points are gen. lower than those of metals

18. Chemical Reactions with Nonmetals • Nonmetal oxide + water → acid • Most nonmetal oxides are acidic oxides • CO2 (g) + H2O (l)  H2CO3 (aq) • Nonmetal oxide + base  salt + water • CO2 (g) + 2NaOH (aq)  Na2CO3 (aq) + H2O (l)

19. Metalloids (aka Semi-metals) • Have properties that are intermediate between those of metals and nonmetals

20. 7.6 Group Trends for the Active Metal Group 1A: The Alkali Metals Characteristics • Soft metallic solids • Silvery • metallic luster • high thermal and electrical conductivities • Low densities and melting points • Most active metals • Exist in nature only as compounds

21. 7.6 Group Trends for the Active Metal Group 2A: Alkaline Earth Metals • Solids with typical metallic properties • Harder, more dense, and melt at higher temperatures when compared to alkali metals • Very reactive towards nonmetals, but not as reactive as alkali metals • Both alkali and alkaline earth metals react with hydrogen to form ionic substances that contain the hydride ion, H-

22. 7.7 Group Trends for Selected Metals Hydrogen • Hydrogen is a nonmetal with properties that are distinct from any of the groups of the periodic table • It forms molecular compounds with other nonmetals, such as oxygen and the halogens

23. 7.7 Group Trends for Selected Metals Group 6A: The Oxygen Group • Most important element in group 6A • Exists in several allotropic forms (different forms of the same element in the same state) • Oxygen is encountered in two molecular forms, O2 (common form) and O3 (aka ozone) • Oxygen has a strong tendency to gain electrons from other elements, thus oxidizing them • In combination with metals, oxygen is usually found as the oxide ion, O2-, although salts of the peroxide ion, O22-, and superoxide ion, O2-, are sometimes formed

24. 7.7 Group Trends for Selected Metals Sulfur!! • 2nd more important element in group 6A • Also exists in several allotropic forms • Elemental sulfur is more commonly found as S8 molecules • In combination with metals, it is more often found as the sulfide ion, S2-

25. 7.7 Group Trends for Selected Metals • Nonmetals that exist as diatomic molecules • There melting and boiling points increase as you go down the column • Have the most negative electron affinities of the elements • Their chemistry is dominated by a tendency to form 1- ions, especially in reactions with metals

26. 7.7 Group Trends for Selected Metals • Group 8A: The Noble Gases aka inert gases • Nonmetals that exist as monoatomic gases • Very unreactive since they have completely filled s and p subshells. Have the complete octet • Have large 1st ionization energies • Only the heaviest noble gases are known to form compounds, and they do so only with very active nonmetals, like fluorine