PERIODICITY

Development of the Periodic Table Mendeleev developed periodic table to group elements in terms of chemical properties. Alkali metals develop +1 charge, alkaline earth metals + 2 Nonmetals usually develop negative charge (1 for halides, 2 for group 6A, etc.) Blank spots where elements should be were observed. Discovery of elements with correct properties.

Periodic Properties Periodic law = elements arranged by atomic number gives physical and chemical properties varying periodically. We will study the following periodic trends: Atomic radii Ionization energy Electron affinity

TRENDS IN THE PERIODIC TABLE - USE OF WHAT WE HAVE JUST LEARNED A)Atomic radius - a number of physical properties of elements are related to the size of an atom, but with our probability picture where does an atom end?

The atomic radius is ½ the distance between the 2 nuclei of the adjacent atoms.

Atomic Radius Fig. 8.17 Atomic Radii for Main Group Elements (s,p) Atomic radii actually decrease across a row in the periodic table. Due to an increase in the effective nuclear charge. (The more protons you have, the harder they pull the e- to them.) Within each group (vertical column), the atomic radius tends to increase with the period number.

Atomic Radius 2 If positively charged the radius decreases. (lost e-) If the charge is negative the radius increases (gained e-). When substances have the same number of electrons (isoelectronic), the radius will depend upon which has the largest number of protons.

C) Atomic radius, in general, decreases as we move from left to right in a row of the periodic table. D) Atomic radius increases from top to bottom in a family or group. E) These 2 trends are the result of 3 influences on size.

As the number of the principal energy level “n” increases, the size increases, they extend further from the nucleus and the covalent radius increases. (bigger outer level, higher floor in the motel—the larger the radius) 2) As nuclear charge (number of protons) increases across a row, the positive charge on the nucleus increases to electrons are pulled closer to the nucleus…radius get smaller.

3) The shielding effect is The attraction for electrons in the outermost shell by the nucleus is shielded by electrons in lower energy levels. *As you gain electrons it becomes harder to pull in the farther ones. a) The smaller size of atoms going across a row can be attributed to minimum shielding.

Electrons in the same shell are attracted more strongly as the nuclear charge (# of protons) increases, because the shielding effect remains the same. If the shielding effect remains the same, the Effective Nuclear Charge increases. The ENC is the positive charge that an electron experiences from the nucleus and is equal to the nuclear charge minus the number of shielding electrons. For example: Li has 3protons in the nucleus, 2e in the 1s orbital (shielding) and 1e in the 2s orbital. ENC = 3 - 2 = 1. The outermost electron "feels" a net attraction by the inside of +1.

Ionization Energy/Electron Affinity Ionization Energy = energy necessary to remove an electron---is always endothermic and positive +. M(g) + h  M+ + e. b) Electron Affinity = energy change upon the addition of an electron can be either endothermic or exothermic depending on the element (for a gaseous atom) A(g) + e A-(g) *An exothermic Energy = - value

IONIZATION ENERGY Ionization energy, Ei: minimum energy required to remove an electron from the ground state of atom (molecule) in the gas phase. M(g) + h  M+ + e. Sign of the ionization energy is always positive, for example, it requires energy for ionization to occur.

Ionization Energy: Periodic table Fig. 8.18 Ionization Energy vs atomic #

IONIZATION ENERGY A(g) + energy  A+ + 1e A(g) + energy A+ + 1e H = + kJ/mol 2) THIS IS A VERY IMPORTANT CONCEPT because the chemical properties of any atom are determined by the configuration of an atom's valence electrons, those electrons in the outermost shell.

6) The trend across from left to right is accounted for by a) the increasing nuclear charge.

WHY? IONIZATION ENERGY The electrons in the outermost shell are more strongly bound to the nucleus due to the increasing effective nuclear charge. as we go across a row of the periodic table energy is larger nuclear charge becomes larger as the number of protons in the nucleus of the atom becomes larger.

IONIZATION ENERGY b) With an electron already in the orbital there is repulsion between the two in the same orbital and it comes out with less energy input. c) The trend from top to bottom of a column shows a decrease in the FIE which corresponds to an increase in the atomic radius. 9) The 2nd, 3rd, and 4th ionization energies are those required to remove the 2nd, 3rd, and 4th electrons.

Electron Affinity 1) Electron affinity is the energy change which occurs when an electron is accepted by an atom in the gaseous state. A(g) + e A-(g) 2) In contrast to ionization energy, what do we observe on the following graph of EA's?

Electron Affinity The greater the negative value of the electron affinity, the greater the tendency of an atom to accept an electron. d) A +value indicates that energy must be absorbed for an atom to gain an electron. e) left to right on the periodic chart, general increasing tendency to form negative ions. However, there are more exceptions than with Ionization Energy.

Electron Affinity

Electron affinities generally become smaller as we go down a column of the periodic table for two reasons. First, the electron being added to the atom is placed in larger orbitals, where it spends less time near the nucleus of the atom. Second, the number of electrons on an atom increases as we go down a column, so the force of repulsion between the electron being added and the electrons already present on a neutral atom becomes larger.

Electron Affinity

Electron Affinity

ELECTRON AFFINITY Electron Affinity, Eea, is the energy change that occurs when an isolated atom in the gas phase gains an electron. E.g. Cl + e  Cl Eea = 348.6 kJ/mol Energy is often released during the process. Magnitude of released energy indicates the tendency of the atom to gain an electron. From the data in the table the halogens clearly have a strong tendency to become negatively charged Inert gases and group I & II elements have a very small Eea.

c) What should you be able to do as a result of this??? I should be able to give you a list of elements and you should be able to put them in order of size from smallest to largest by just looking at their positions on the chart. You should be able to tell me the reasons why they are smaller or larger.