Trends of the Periodic Table

9/5/08

1. Atomic Radii• atomic radii (atom size) is determined by

measuring the distance between the nuclei of 2 identical atoms that are chemically bonded together, then dividing this distance by 2. • measured in pm (picometers, or 10-12 m).

• across a period, atomic radii tends to decrease from left to right.• this is due mostly from the increase in proton number. • as electrons are added to the s and p sublevels, they

are pulled closer and closer to the highly positive nucleus, which results in a decrease in atomic radii.

• down any group on the table, the atomic radii will increase. • this is due to an increase in total number of particles

(protons, neutrons, and especially electrons) in the atom. The number of energy levels needed increases, so the overall size increases.

• the electrons that are added to these larger elements are being added in orbitals that lie farther and farther away from the nucleus, and thus, father away from any positive charges that would bring them closer inside and decrease the radii.

• thus, the overall trend line for atomic radii looks like the figure to the right.

• the green box in the bottom left corner of the table indicates the largest element on the table.

• note that fluorine(blue) is one of the smallest elements.

• let’s take a look at the actual values for atomic radii........

De c r e a s e s

TO THE NEXT SLIDE!

Increases

2. Ionization Energy• ionization energy is the energy required to remove one

electron from an atom, forcing the creation of a positive ion. • for example, if you were to remove the one valence

electron that sodium (Na) has in its outermost s and p orbitals, you would create an Na+ ion.

• any process that refers to the creation of an ion is called ionization.

• IE is measured in kilojoules per mole (kJ/mol), which is a unit that represents energy required per amount of substance.

• across a period from left to right, IE will generally increase. • this is due mostly to the number of valence electrons

that can be found in their outermost s and p orbitals. • the lower the number, the easier it is to lose them.

• therefore, the alkali metals will lose electrons very easily (only one to lose) and have low IE, and.......

• therefore, the alkali metals will lose electrons very easily (only one to lose) and have low IE, and.......

• the noble gases (with 8 valence e-) will hold onto those electrons for dear life, and have very high IE.

• this explains why the alkali metals are very reactive with most substances, and noble gases are the most unreactive elements.

• the reason that lower valence electron number equals lower IE is because of a lower number of protons are present in the nucleus on the left side of the table when compared with the right side.

• therefore, there is less positive charge in the nucleus to hold those electrons in towards the nucleus, and they are more easily lost for elements on the left side of the table.

• down a group, IE will tend to decrease.

• just as atomic radii was affected by the electrons being added to s and p orbitals that were farther and farther away from the nucleus, IE is affected in the same way, leading to an easier removal for electrons that lie in the outermost, most distant s and p orbitals.

• also, you must realize that electrons that lie in between the nucleus and those outermost s and p electrons will tend to shield the effect of the nuclear charge from those outermost electrons.

• thus, the overall trend line for IE looks like the figure to the right.

• the green box in the top right corner of the table indicates that you will find the elements with the

largest IE there. • note that fluorine (blue) is the element that has the

highest IE (other than the noble gases). • let’s take a look at the actual values for IE........

Decreases

I n c r e a s e s

3. Ionic Radii• ionic radii (ion size) is the new size of an atom after it

has either gained or lost one or more electrons, forming an ion.

• there are two kinds of ions: cations and anions.• positive ions are known as cations.

• the formation of a cation (by the loss of one or more electrons) always leads to a decrease in atomic radius because the removal of the highest-energy-level electrons results in a smaller electron cloud.

• also, the remaining electrons are drawn closer to the nucleus by its unbalanced positive charge.

• negative ions are known as anions.• the formation of an anion (by the addition of one or

more electrons) always leads to an increase in atomic radius.

• negative ions are known as anions.• the formation of an anion (by the addition of one or

more electrons) always leads to an increase in atomic radius.

• this is because the total positive charge of the nucleus remains unchanged when an electron is added to an atom or an ion. • so the electrons are not drawn to the nucleus as

strongly as they were before the addition of the extra electron.

• the electron cloud also spreads out because of greater repulsion between the increased number of electrons.

• within each period of the periodic table, the metals at the left tend to form cations and the nonmetals at the upper right tend to form anions.

• within each period of the periodic table, the metals at the left tend to form cations and the nonmetals at the upper right tend to form anions. • cationic radii decrease across a period because the

electron cloud shrinks due to the increasing nuclear charge acting on the electrons in the same main energy level.

• anionic radii decrease across each period for the elements in the upper p-block. The reasons for this trend are the same as the reasons that cationic radii decrease from left to right.

• as they are in atoms, the outer electrons in both cations and anions are in higher energy levels as one reads down a group. • therefore, just as there is a gradual increase of atomic

radii down a group, there is also a gradual increase of ionic radii.

• thus, the overall trend line for ionic radii looks like the figure to the right. • the green box in the bottom left corner of the table indicates where you will find the element with the largest ionic radii. • note that fluorine (blue) is an element with a very

small anionic radii. Aluminum (red) is an element with a very small cationic radii.

Increase

D e c r e a s e

4. Electronegativity• electronegativity is a measure of the ability of an atom in

a chemical compound to attract electrons. • remember that valence electrons are the only

electrons involved in bonding. • some elements are better able than others to attract

the valence electrons involved in the bonding process.

• an uneven concentration of charge in a compound affects the chemical properties of that compound. • if electrons within a compound are being pulled

unevenly in one direction, the molecule is said to be “ionic” or “polar” (depending on what elements are involved).

• if there is an equal distribution of electrons, and thus, charge, on a molecule, then it is said to be “non-polar.”

• Linus Pauling, one of America’s greatest chemical minds, created an electronegativity chart to relate elements to each other. • he assigned fluorine, the most electronegative

element, a value of 4.0. • values were assigned to all the other elements

in relation to this value.

• across a period, electronegativity values increase from left to right., and down a group, they tend to decrease. • thus the trend line for electroneg- ativity looks like the figure to the right. • the elements with the largest electronegativities are located in the top right corner of the table.

• notice fluorine is the element with the most electronegativity.

Pauling

I n c r e a s eDecrease

• Let’s take a look at the actual values for electronegativity........