Section 6.3 - Periodic Trends

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

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).

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

Trends in Atomic Radius WITHIN A GROUP Within periods No additional electrons come between valence electrons and the nucleus. Thus, the valence electrons are not shielded from the increased nuclear charge. The result is than the increased nuclear charge pulls the outermost electrons closer to the nucleus. Within groups Nuclear charge increases and electrons are added to successively higher principal energy levels.

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.

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

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.

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 Within periods, positive ions gradually decrease. Beginning in group 5A or 6A, the size of the much larger negative ions also gradually decreases. Within groups Increase as you go down.

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.

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.

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! To form a positive ion, an electron must be removed from a neutral atom. A high value indicates the atom has a strong hold on its electrons. Atoms with high values are not likely to form positive ions. A low value indicates an atom loses its outer electrons easily. Such atoms are likely to form positive ions.

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?

First Ionization Energy Trends Group 1A metals have low ionization energies. They are likely to form positive ions. First ionization energy is the energy required to remove the 1st electrons. The second ionization energy is the energy required to remove a second electron.

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.

Periodic Trends Decreasing Ionization Energy

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.

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.

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 is related to the atom’s number of valence electrons.

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.

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.)

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

Trends in Electronegativity