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 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 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 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

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. Metals Non-Metals Do not have a luster; various colors
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 Table 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 generally 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