Periodic table

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3 Periodic trends in the properties of atoms One of the most fundamental principles of chemistry is the periodic law, states that, the chemical and physical properties of elements are a periodic function of atomic number. This is, of course, the principle behind the structure of the periodic table. Elements within a given vertical group resemble one another chemically because chemical properties repeat themselves at regular intervals of 2, 8, 18, or 32 elements. We will discuss how, atomic radius, ionization energy, electronegativity vary horizontally and vertically in the periodic table.

Period Group

5 I- Atomic radius We can’t speak strictly about the size of an atom as the electron cloud surrounding the nucleus does not have a sharp boundary. However, a quantity, called the atomic radius can be defined and measured assuming a spherical atom and the atomic radii can be taken as one half the distance of closest approach between atoms in an elemental substance. Examples, For copper atoms in metallic copper, the atomic radius is found to be nm, while in case of chlorine, the arrangement of atoms gives a radius of nm.

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7 The atomic radii of the main group elements, 1.Decreases across a period from left to right in the periodic table. 2.Increases down a group in the periodic table.

8 Explanation Consider the first increase in radius observed as we move down the table for example in the the alkali metals group. All these elements have a single s electron outside a filled level or filled p sublevel. Electrons in these inner levels are much closer to the nucleus than the outer s electron and hence effectively shield it from the positive charge of the nucleus. To a first approximation, each inner electron cancels the charge of one proton in the nucleus, so the outer s electron is attracted by a net positive charge +1. In this it has properties of an electron in the hydrogen atom. Because the average distance of the electron from the hydrogen nucleus increases with the principal quantum number n, the radius increases moving from Li ( 2s electron) to Na (3s electron) and so on down the group.

9 Explanation (cont.) The decrease in atomic radius across the periodic table can be explained in a similar manner. Consider the third period, where electrons are being added to the third principal energy level. The added electrons should be relatively poor shields for each other because they are all at about the same distance from the nucleus. Only ten core electrons in the inner filled levels (n=1, n=2) are expected to shield the outer electrons from the nucleus. This means that the charge felt by an outer electron, called the effective nuclear charge, should increase steadily with atomic number as we move across the period. As effective nuclear charge increases, with increase of atomic number, the outermost electrons are pulled in more tightly and atomic radius decreases.

10 II- Ionic radius The radii of cations and anions derived from atoms of a group have the same trend as observed for atoms, i.e. It increases moving down a group and decreases going through a period from left to right. Comparing the radii of cations and anions with those of the atoms they are derived: Positive ions are smaller than the metal atoms from which they are formed. The Na + ion has a radius of nm (95 pico meter), only a little more than half that of the Na atom, nm. Negative ions are larger than the nonmetals from which they are formed. The radius of the Cl - ion, nm, is nearly twice that of the Cl atom, 0.099nm.

11 The difference in radii between atoms and ions can be explained quite simply. A cation is smaller than the corresponding metal atom because the excess of protons in the ion draws the outer electrons in closer to the nucleus. In contrast, an extra electron in an anion adds to the repulsion between outer electrons, making a negative ion larger than the corresponding nonmetal atom.

12 Example Using only the periodic table arrange the following sets of atoms and ions in order of increasing size. Mg, Al, Ca.b) S, Cl, S 2- c) Fe, Fe 2+, Fe 3+ a) Compare the other atoms with Mg. Al to the right is smaller than Mg. Ca below Mg, is larger. The order is Al<Mg<Ca b) Compare with S atom. Cl to the right is smaller. The S 2- anion is larger than the S atom The predicted order is Cl<S<S 2-. c) Compare with Fe 2+ ion, the Fe atom with no charge, is larger. The Fe 3+ ion with a 3+ is smaller. The predicted order is Fe 3+ < Fe 2+ < Fe.

13 III- Ionization energy (IE) Ionization energy is a measure of how difficult it is to remove an electron from a gaseous atom. Energy must be absorbed to bring about ionization, so ionization energies are always positive quantities. The first ionization energy is the energy change for the removal of the outer most electron from a gaseous atom to form a +1 ion. M (g) M + (g) + e  E1 = first ionization energy The more difficult it is to remove electrons, the larger the ionization energy.

14 Notice that Ionization energy (IE) increases across the periodic table from left to right and decreases moving down the periodic table. Comparing the trends of ionic radii and ionization energy it is clear that there is an inverse correlation between them. The smaller the atom, the more tightly its electrons are held to the positively charged nucleus and the more difficult they are to remove thus the higher the ionization energy. Conversely, in a large atom such as that of group one metal, the electron is relatively far from the nucleus, so the lower the energy that has to be supplied to remove it from the atom.

15 III- Ionization energy (IE) (cont.) There are some exceptions to the observed trend which can be explained in terms of electron repulsions, Examples: The decrease in the IE in going from Be to B reflects the fact that the electrons in filled 2s orbital provide some shielding for electrons in the 2p orbital from the nuclear charge. The decrease in IE in going from N to O reflects the extra electron repulsions in doubly occupied oxygen 2p orbitals.

16 III- Ionization energy (IE) (cont.) Example Consider atoms with following electron configurations 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 Which atom has the largest first ionization energy, explain your choice. The atom with the largest value of IE is 1s 2 2s 2 2p 6 (this is a neon atom), because this element is found at the right end of period 2. As for the other two configurations including 3s electrons they will be of lower IE.

17 IV- Electron affinity, (some times referred to as) Electronegativity Electron affinity is the energy change associated with the addition of an electron to a gaseous atom X (g) + e X - (g) Electronegativity increases moving from left to right in the periodic table and it decreases moving down a group. Electronegativity defines the strength of an atom to attract shared electrons within a covalent bond. The result is a polarized bond with partial negative or positive charges on the respective atom. Partial charges are indicated by the symbol  + or  -.

18 Electronegativity (cont.) Consider the hydrogen fluoride molecule H-F :  + H-F  - F is more electronegative than H that is why it tends to have a  - charge while H will acquire  + charge. This is called bond polarity. Decreasing electronegativity Increasing electronegativity

19 The relationship between electronegativity and bond type is shown below  For identical atoms ( an electronegativity difference of zero), the electrons in the bond are shared equally and no polarity develops.  When two atoms with very different, electronegativities interact, electron transfer can occur, to form ions that make up an ionic substance.  Intermediate cases give polar covalent bonds with unequal electron sharing.

20 Example Order the following bonds according to polarity H-H, O-H, Cl-H, S-H and F-H The polarity of the bond increases as the difference in electronegativity ( obtained from the table) increases. (2.1) (2.1) (2.5)(2.1)(3.0)(2.1)(3.5)(2.1)(4.0)(2.1) Thus the order is H-H <S-H<Cl-H <O-H<F-H Covalent bond Polarity increases Polar covalent bond