Development of Periodic Table Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped.

Properties of Metal, Nonmetals, and Metalloids

Metals Compounds formed between metals and nonmetals tend to be ionic. Metal oxides tend to be basic.

Nonmetals Substances containing only nonmetals are molecular compounds. Most nonmetal oxides are acidic.

Metalloids Metalloids have some characteristics of metals and some of nonmetals. For instance, silicon looks shiny, but is brittle and a fairly poor conductor.

Noble gases (Group 18) All atoms of the noble gases have their outer s and p orbitals filled. We will see later that these atoms require very large amounts of energy to form ions, so much in fact, that they are difficult to alter chemically and as such are inert (unreactive) and do not tend form ions.

Alkali metals (Group 1) & Group 1 atoms have an electronic structure [Noble gas] ns1. This means that they tend to lose the s electron when they from an ion, leaving behind an inert noble gas type structure. This explains why Group 1 elements tend to only form 1+ ions.

Alkali Metals Alkali metals (except Li) react with oxygen to form peroxides. K, Rb, and Cs also form superoxides: K + O2  KO2 They produce bright colors when placed in a flame.

Group 2 (Alkaline earth) atoms have an electronic structure [Noble gas] ns2. This means that they tend to lose the two s electrons when they from an ion, leaving behind an inert noble gas type structure. This explains why Group 2 elements tend to only form 2+ ions. A similar argument can be applied to group 3 atoms and their simple ions.

Alkaline Earth Metals Beryllium does not react with water, and magnesium reacts only with steam, but the other alkaline earth metals react readily with water. Reactivity tends to increase as you go down the group.

Chalcogens -Group 16 atoms have an electronic structure [Noble gas] ns2 np4. This means that they tend to gain two p electrons when they from an ion, to reach an inert noble gas type structure with a charge of 2-. Halogens-Group 17 atoms have an electronic structure [Noble gas] ns2 np5. This means that they tend to gain one p electron when they from an ion, to reach an inert noble gas type structure with a charge of 1-.

Group VIIA: Halogens They are diatomic in their elemental form. Therefore, they tend to oxidize other elements easily (cause other elements to lose their electrons). They react directly with metals to form metal halides. Chlorine is added to water supplies to serve as a disinfectant.

Periodic Trends In this chapter, we will rationalize observed trends in Sizes of atoms and ions (atomic and ionic radius) Ionization energy. Electron affinity.

What Is the Size of an Atom? The (bonding) atomic radius is defined as one-half of the distance between covalently bonded nuclei. Knowing the arrangement of atoms in a metallic crystal allows determination of the radii of metals.

Sizes of Atoms The bonding atomic radius tends to — Decrease from left to right across a row (due to increasing Zeff). — Increase from top to bottom of a column (due to the increasing value of n).

Sizes of Atoms The bonding atomic radius tends to — Decrease from left to right across a row (due to increasing Zeff). — Increase from top to bottom of a column (due to the increasing value of n).

As a period is traversed from left to right the atomic size decreases As a period is traversed from left to right the atomic size decreases. This is because the nuclear charge increases (greater positive charge, extra protons) and the subsequent electrons enter the same shell experiencing no extra shielding from inner electrons and are therefore pulled in more tightly.

Effective Nuclear Charge In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. The nuclear charge that an electron experiences depends on both factors.

Effective Nuclear Charge The effective nuclear charge, Zeff, is found this way: Zeff = Z − S where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons.

Group trend atomic radius- Variation of atomic size down a group: As a group is descended the atomic size increases. This is because that although there is again an increase in nuclear charge (Z eff ) (greater positive charge, extra protons) the outer electrons enter new shells much further away from the nucleus ( n is increasing) and experience more shielding. As a result the atomic size increases since the greater the number of shells in an atom the larger the atom.

Sizes of Ions Ionic size depends upon The nuclear charge. The number of electrons. The orbitals in which electrons reside.

Sizes of Ions Cations are smaller than their parent atoms: The outermost electron is removed and repulsions between electrons are reduced.

Sizes of Ions Anions are larger than their parent atoms” Electrons are added and repulsions between electrons are increased.

Sizes of Ions Ions increase in size as you go down a column: This increase in size is due to the increasing value of n.

Sizes of Ions In an isoelectronic series, ions have the same number of electrons. Ionic size decreases with an increasing nuclear charge.

Arrange the following species in order of increasing size. Ar, K+, Ca2+, S2-, Cl

Ionization Energy The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion. The first ionization energy is that energy required to remove the first electron. The second ionization energy is that energy required to remove the second electron, etc.

Ionization energies are often measured in units of kJ mol-1 Ionization energies are often measured in units of kJ mol-1. They have positive values indicating that energy must be “put in” in order to remove electrons. It is endothermic.

The magnitude of the ionization energy is determined by the attraction of the positive nucleus for the negative electrons that are being removed. The attraction is dependent upon two factors; 1. The nuclear charge (how many protons are present). 2. The shielding effect of the inner electrons (the extent to which inner electrons protect the outer electrons from the nuclear charge). 3. These two things together are the Zeff

Group Trends in First Ionization Energies As one goes down a groupcolumn, less energy is required to remove the first electron. For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.

Period Trends in First Ionization Energies Generally, as one goes across a period, it gets harder to remove an electron. As you go from left to right, Zeff increases. No extra shielding, as electrons are being removed from the same n

Ionization Energy It requires more energy to remove each successive electron. Slightly less electron repulsion, electrons move closer to nucleus, increasing their attraction to the nucleus

Ionization Energy Look at Na The large "jump" in the data between the 1st and 2nd ionization energies shows that it is relatively easy to remove the 1st electron but extremely difficult to remove the 2nd (the ionization energy increases approximately x10). It can be assumed that the 2nd electron is in a new shell, is closer to the nucleus, is more difficult to remove (higher ionization energy) and therefore the 1st electron was in the outside shell on its' own, hence group I.

The slightly odd behavior of boron & aluminum and oxygen & sulfur can be explained thus; Boron & Aluminum: Elements in group 13 have the valence (outer) electronic configuration ns2 np1. The outer np1 electron is in a p orbital that has a slightly higher energy than the corresponding s orbital, is slightly further away form the nucleus and experiences slightly more shielding from the full ns2 sub-shell in addition to the inner complete shells. As a result these elements exhibit slightly lower first ionization energies than would otherwise be expected. Oxygen & Sulfur: Elements in group 16 have the valence (outer) electronic configuration ns2 np4. The electron that is being removed in an ionization process is paired with another in one of the p orbitals and as a result experiences repulsion. This repulsion means it is more easily lost, and as a result these elements exhibit slightly lower first ionization energies than would otherwise be expected

1. Consider the first and second electron affinities of oxygen 1. Consider the first and second electron affinities of oxygen. Write two equations to represent these processes. In terms of energy, explain the differences in these processes

Give a brief but complete account of the changes in ionization energy as the following transitions in the periodic table are made. (i) Passing from magnesium to strontium (ii) Passing from sodium to argon

Define the term first ionization energy and state the two factors which influence its magnitude.

(i) What is the most likely formula of R's oxide? Consider the following successive ionization energies (kJ mol-1) of element R. 1st 2nd 3rd 4th 5th 6th 737 1450 7732 10540 13360 17995 (i) What is the most likely formula of R's oxide? (ii) Why is the value for the 2nd ionization energy greater than the 1st?

Electron Affinity Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom: Cl + e−  Cl−

Trends in Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row.

Trends in Electron Affinity There are again, however, two discontinuities in this trend.

Trends in Electron Affinity The first occurs between Groups IA and IIA. The added electron must go in a p orbital, not an s orbital. The electron is farther from the nucleus and feels repulsion from the s electrons.

Trends in Electron Affinity The second discontinuity occurs between Groups IVA and VA. Group VA has no empty orbitals. The extra electron must go into an already occupied orbital, creating repulsion.

Metals versus Nonmetals Differences between metals and nonmetals tend to revolve around these properties.