Section 6.3 - Periodic Trends

1. Section 6.3 - Periodic Trends

2. Objectives • Compare period and group trends of several properties. • Relate period and group trends in atomic radii to electron configuration.

3. Periodic Trends • When the properties of an element change in a predictable way, we call it a “trend”. • In the periodic table, there are trends observed within a group (from top to bottom) and across a period (from left to right).

4. Periodic Trends • Atomic radius is defined as half the distance between the nuclei of identical atoms that are chemically bonded together. • Atomic radius is a measure of atomic size

5. Trends in Atomic Radius WITHIN A GROUP

6. WHY does radius increase WITHIN A GROUP? 1. As the number of the energy level of the valence electrons increases, the size of the energy level (and its orbitals) increases because there is more space to occupy. Therefore, the size of the atom increases. 2. Valence electrons in higher energy levels are further from the nucleus & feel the pull of the positively charged nucleus less and less. 3. Outer energy level electrons are also shielded from increasing positive nuclear charge by electrons in the inner energy levels.

7. Trends in Atomic Radius ACROSS A PERIOD • Across a period, atomic radius decreases • Increasing atomic number means an increase in nuclear charge across the period • Since the energy level remains the same, the valence electrons do not move further away. (There is NO shielding.) Therefore, the increase in positive charge in the nucleus pulls on the increasing number of electrons with equal force. • The electrons are pulled to the nucleus and radius decreases. • Prac. Probs. pg. 189 #16-18

8. Trend in Ionic Radius • An ION is an atom or a bonded group of atoms that has a positive or negative charge • When atoms lose electrons, they form positively charged ions (the number of protons will be greater than the number of electrons) • Positive ions have empty orbitals so the ion will always be smaller than the atom. • In addition, the remaining electrons experience less repulsion so they can get closer to each other and to the nucleus.

9. Ionic Radius Trends • When an atom gains electrons, it becomes negatively charged (more electrons than protons). • The ion will be larger than the atom because 1) The pull on each electron will be smaller 2) Increased electron repulsion causes an increase in radius

10. Ionic Radius Trend Across a Period • Negative ions are always larger than positive ions. • As charge on positive ions increases, ionic radius decreases. • As charge on negative ions decreases, ionic radius decreases.

11. Ionic Radius Trend Within a Group • As atomic number increases (top to bottom), ionic radius increases for both positive & negative ions. • This is because there is an increase in energy levels down a group.

12. Trends in Ionization Energy • Ionization energy is the energy required to remove an electron from a gaseous form of that atom. • Think of it as an indication of how strongly an atom’s nucleus holds onto its valence electrons - a high value means the atom has a strong hold on its electrons - they are not likely to form positive ions!

13. Types of Ionization energy • First ionization energy: energy required to remove the first valence electron • Second ionization energy: energy required to remove the second valence electron from a +1 ion • Third ionization energy?

14. First Ionization Energy Trends

15. First Ionization Energy Trends • Ionization Energy INCREASES across a period. An increasing nuclear charge produces an increased hold on valence electrons. • Ionization Energy DECREASES within a group (top to bottom). Valence electrons are further away from the nuclear positive charge and thus easier to remove.

16. Periodic Trends Decreasing Ionization Energy

17. Ionization Energy Trends • Open your books to page 192. • In Table 5, you will see that the energy required for each successive ionization always increases. • For each element there is an ionization for which the required energy jumps dramatically.

18. Ionization Energy Trends • Find this ionization for Boron. • This means a boron atom can “easily” lose the first, second, and third valence electrons but it is extremely hard to remove the 4th. Therefore, very unlikely that it will lose the fourth electron.

19. Ionization Energy Trends • Boron has 3 valence electrons and will “easily” form a +3 ion. (It will NOT form a +4 ion!) • The ionization at which the large jump in energy occurs isrelated to the atom’s number of valence electrons.

20. The Octet Rule • Sodium atom (Na) 1s22s22p63s1 • Sodium ion (Na+) 1s22s22p6 • The sodium ion has the same electron configuration as neon, a noble gas. • Filled s and p orbitals of the same energy level are unusually stable. • Octet Rule – atoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electrons.

21. The Octet Rule • Useful for determining types of ions likely to form • Left side of table (METALS) - will LOSE electrons. (Will form positive ions.) • Right side of table (NONMETALS) - will GAIN electrons to acquire an octet. (Will form negative ions.)

22. Trends in Electronegativity • The electronegativity of an element indicates the relative ability of its atoms to attract electrons in a chemical bond. • Noble gases have essentially NO electronegativity. • EN is expressed in terms of a numerical value of 4.0 or less; see pg. 194

23. Trends in Electronegativity