Periodic Properties of the Elements

Periodic Properties of the Elements

Chapter 7 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. Atoms arent 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 1) 2)

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. Ionization energy the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion

1st ionization energy (I ) The energy 1 needed to remove the first electron from a neutral atom, forming a cation 2nd ionization energy (I ) the energy 2 needed to remove the second electron The greater the ionization energy, the harder it is to remove an electron

HIGH ionization energy means the atom holds 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 doesnt require much 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

Electron affinity the energy change that occurs when an electron is added to a gaseous atom A negative electron affinity means the anion is stable A positive electron affinity means the anion is higher in energy than are the

separated atom and electron. The anion is not stable and will not form 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.

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 doesnt change greatly as we

move down a group. Electron affinity should become more positive (less energy released). 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 Most metal oxides are ionic solids that are basic

Poor conductors of heat and electricity 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 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 conduct heat & electricity They are malleable & ductile

Solids at room temp. except mercury(Hg) (its 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. 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)

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 generally lower than those

of metals 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)

Have properties that are intermediate between those of metals and nonmetals 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 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-

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 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 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- 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

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

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