Periodic Properties of the Elements

1. Periodic Properties of the Elements Chapter 7

2. The Periodic Table • Developed independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev (1870”s). • Found that similar chemical and physical properties recur periodically when the elements are arranged in order of increasing atomic number. • At this time they didn’t know about atomic numbers, so used masses which generally increases as atomic # increases. • Had blanks in the table and used periodicity to guess about the characteristics of the missing elements.

3. Moseley • After Rutherford proposed the nuclear model of the atom, Henry Moseley developed the concept of atomic numbers.

4. Bonding Atomic Radius- based on the distance separating atoms when they are chemically bonded to one another • This radius is shorter than the non bonded radius due to the nuclear attraction between the two atoms. • The bonding atomic radius decreases as you go across the period, and increases as you go down a group. (Rb> F)

5. Bond length • Predict which will be greater, the P-Br bond length in PBr3 or the As-Cl bond length in AsCl3 • P-Br

6. Radial Electron Density • The probability of finding an electron with respect to the nucleus • The 1s subshell in Ar is much closer to the nucleus than the 1s subshell of He. This is because of the Zeff

7. Ionic Radius • Cations are formed when metallic atoms lose valence electrons. • These ions have smaller radii than their parent atoms • Anions are formed when nonmetallic atoms gain electrons • These ions are larger than their parent atoms due to the extra repulsions of another electron

8. Periodic Trends • How easily an electron will be removed from an atom is an important indicator of the chemical nature of that atom. • Ionization energy is the energy required to remove an electron from the ground state a gaseous atom • The greater the ionization energy, the more difficult it is to remove an electron

9. Ionization Energy • Highest energy electron removed first (outermost). • First ionization energy (I1) is that required to remove the first electron. • Second ionization energy (I2) - the second electron • etc. etc.

10. Trends in ionization energy • for Mg • I1 = 735 kJ/mole • I2 = 1445 kJ/mole • I3 = 7730 kJ/mole • The effective nuclear charge increases as you remove electrons. • It takes much more energy to remove a core electron than a valence electron because there is less shielding. • This trend is because the positive nuclear charge that provides attractive forces remains the same, while the number of electrons which provide repulsive forces decreases.

11. Explain this trend • For Al • I1 = 580 kJ/mole • I2 = 1815 kJ/mole • I3 = 2740 kJ/mole • I4 = 11,600 kJ/mole

12. Ionization Trends • Generally from left to right, I1 increases because there is a greater nuclear charge with the same shielding. (Generally, the alkali metals show the lowest ionization energies in a row, and the noble gases the highest. • As you go down a group I1 decreases because electrons are farther away.

13. It is not that simple • Zeff changes as you go across a period, so will I1 • Half filled and filled orbitals are harder to remove electrons from. • here’s what it looks like.

14. First Ionization energy Atomic number

15. First Ionization energy Atomic number

16. First Ionization energy Atomic number

17. Electron Affinities • The energy change that occurs when an electron is added to a gaseous atom to form a negative ion. (A measure of the affinity or attraction for the added electron. • For most atoms the electron affinity is negative because energy is released when an electron is added.

18. Different entities • Remember!!! • Ionization Energy is the desire to lose an electron (+) • Electron Affinity is the desire to gain an electron

19. Electron Affinity Trends • Generally becomes increasingly negative as you go toward the halogens. For the noble gases, the electron affinity is positive, meaning the ion will not form because that would mean that the gas would have to go to a higher energy sub shell which is energetically unfavorable. • Any time the value is zero, the ion will not form. The bigger the negative, the more likely that the ion will form. • Chlorine has the highest electron affinity.

20. Parts of the Periodic Table Metals, Non metals, and Metalloids

21. Metals • Roughly 3/4 of the elements are metals. • Properties of metals include luster, malleability, ductility, god conductors of heat & electricity, form cations in an aqueous solution. • The more an element shows properties of metals, the greater it’s metallic character • (Increases down a row, decreases across a period)

22. More Properties of Metals • All metals are solid except for mercury • Metals tend to have low ionization energies which is why they are oxidized (lose electrons) to form a cation when they undergo a chemical reaction. • When the outermost electrons are lost, the ion achieves a noble gas configuration • Many of the transition metals have the ability to form more than one ion

23. Metal Oxides • Metal-nonmetal compounds are said to be ionic • Oxides are especially important (oxygen is everywhere!) • Most metal oxides are basic, which means they dissolve in water to form bases • Metal oxides can also react with acids to form salt + water

24. Nonmetals • Poor conductors of heat & electricity • Vary in appearance • Have lower melting points than metals • Several exist as diatomic molecules • Tend to gain electrons in a chemical reaction to fill their outer p subshell completely giving a noble gas configuration

25. More properties of Nonmetals • Nonmetal bonded to a nonmetal makes a molecular substance • These molecules tend to be gases, liquids, or low-melting solids. • Non-metal oxides are generally acidic which means they combine with water to form an acid. (This is why carbonated water is acidic) • These acidic nonmetal oxides will combine with a base to produce salt & water

26. Metalloids • Have properties of both metal and nonmetal! • Silicon- looks like a metal, brittle like a nonmetal, semiconductor used in computer chips

27. Group Trends Active Metals

28. The Alkali Metals • Group 1 • Soft metallic solids • Doesn’t include hydrogen- it behaves as a non-metal • Down the group-decrease in IE • Down the group-increase in radius • Decrease in density • Decrease in melting point

29. Alkali Metals • For each row, the alkali has the lowest ionization energy • All very reactive and lose 1 electron to form +1 cation • Exist in nature only as compounds • Electrolysis used • React vigorously with water to produce hydrogen gas and metal hydroxides • Exothermic enough to ignite Hydrogen

30. Also extremely reactive with oxygen • Stored in kerosene • Do not produce a colored solution because no electron to excite

31. Alkaline Earth Metals • All solids • Harder , more dense, melt at higher temps than alkali • Slightly higher ionization energies, thus slightly less reactive (compared to Alkali). • Increasing reactivity as you go down the group that accounts for why Berylium does not react with water but Calcium and everything below it do. • Tend to lose 2 electrons and form +2 cation

32. Alkaline Earth Metals • Highly reactive so usually found in nature as part of a compound

33. Group Trends Nonmetals

34. Hydrogen • Nonmetal that occurs as a colorless diatomic gas • Since there is no shielding, has an extremely high ionization energy • Usually combines with other non-metals to form molecular compounds • Reacts with active metals to form metal hydrides (H is -)

35. Oxygen’s Group • Changing trends as you go down the group • Oxygen usually found in two molecular forms oxygen and ozone (allotropes- different forms of the same element in the same state) • Oxygen makes up 21% of air • Ozone is toxic and smelly • Oxygen usually present as the oxide ion

36. Oxygen’s Group • Sulfur (exists as eight membered rings of sulfur atoms • Most sulfur is found as metal sulfides • Can be burned in oxygen to produce sulfur dioxide (pollutant)

37. The Halogen Family • (Astatine omitted because extremely rare, radioactive and unknown) • All typical non-metals • Melting and boiling point increase as you go down the group • All diatomic • Tend to gain electrons and form -1 anion • Have highly negative electron affinities • Fluorine and Chlorine most reactive

38. The Noble Gases • All non-metal gases • All monoatomic • Rn too highly radioactive to study • Completely filled s and p subshells • Large ionization energies which decrease as you go down the group • Inert gases because thought to be unable to form compounds