Periodic Properties of the Elements Chapter 7

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.

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

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.

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

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

7.3 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

Ionization Energy 3-D

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

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.

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).

7.5 Metals, Nonmetals, and Metalloids

Metals Non-Metals Have a shiny luster; various colors, although most are silvery Do not have a luster; various colors Solids are malleable and ductile Solids are usually brittle; some are hard, and some are soft Good conductors of heat and electricity Poor conductors of heat and electricity Most metal oxides are ionic solids that are basic Most non-metallic oxides are molecular substances that form acidic solutions Tend for form cations in aqueous solutions Tend to form anions or oxyanions in aqueous solution Pg. 239 --Table 7.3 Characteristic Properties of Metals and Nonmetals

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

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.

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)

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

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)

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

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

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-

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

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

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-

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

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