Periodic Trends and Ionization September 19th 2011 Periodic Trends and Ionization SCH4U1 Mr. Dvorsky

By now you should know..... -the periodic table is arranged by atomic number, Z increases from left to right and from top to bottom. -the number of valence electrons also increases from left to right, but down a group they remain the same. [see chart on board] –explains why elements in same group have similar properties since they have the same number of valence electrons to react with.

Factors Affecting the Properties of the Elements

The properties of the elements are determined by the force of attraction between the nucleus and outermost electrons. This force is itself determined by two factors: the nuclear charge and more importantly the distance between the nucleus and the outer electrons

Effective nuclear charge: see explanation on board.

Atomic Radius The distance from the centre of the atom to the outmost electrons. –atomic radius indicates the size of an atom Technically speaking, atoms don’t have a fixed radius like a ball, so the radius is found by finding the distance between the nuclei of two touching atoms and dividing that distance by half. For the scope of this class, the definition above shall suffice.

Atomic Radius There is an increase in atomic radius down a group because successively more orbitals are added There is a decrease in atomic radius from left to right on the periodic table. -more electrons are added to the outer shell -the effective nuclear charge of the nucleus becomes stronger. -therefore the nucleus and outer shells are pulled more strongly together and the radius decreases.

Ionization Energy Definition: -It is the minimum amount of energy required to remove a valence electron from a gas phase atom. -Since energy must be absorbed to remove an electron, ionization energies are always positive quantities. -The electron that is the least tightly bound will be removed first. –it will be one of the valence electrons.

Ionization Energy The removal of the first electron is known as the first ionization energy, if a second is removed, this is the second ionization energy, and so on. A(g)  A+(g) + e- = first ionization energy A+(g)  A2+(g) + e- = second ionization energy

Ionization Energy The energy required for each successive ionization increases because the net positive charge of the nucleus felt by the remaining electrons is increasing. As a result, the remaining electrons are held more tightly to the nucleus.

Ionization Energy Across the rows in the periodic table, ionization energy correlates with Z*. Down the groups, ionization energy decreases. The reason for this is that as atoms get larger, the distances between the nucleus and valence electrons is larger = less attraction and easier removal.

show periodicity (that means it varies in a repetitive way on the periodic table) e.g. Look at Li to Ne and compare with Na to Ar

Ionization Energy There are factors that affect ionization energy. Remember that ionization energy is a a measure of the energy needed to pull a particular electron away from the attraction of the nucleus. A high value of ionization energy shows a high attraction between the electron and the nucleus.

Ionization Energy -See explanation on board for anomalies in trend.

Ionization Energy – Transition Series Trend

Electron Affinity -it is the change in energy when a neutral atom (in gas phase) gains an electron A(g) + e-  A-(g) -similar to ionization energy, the addition of the first electron is known as the first electron affinity, and the addition of a second is known as the second electron affinity.

Electron Affinity -Electron affinity can be energetically favoured depending on the atom in question and how many electrons are being added. e.g. F(g) + e-  F-(g) is favoured more than O(g) + e-  O-(g)

Electronegativity Electronegativity is the ability of an atom to draw electron density to itself in a bond in a compound. The closer to F on the periodic Table, the more Electronegative.

Electronegativty Because atomic size decreases across the rows in the periodic table, and smaller more compact atoms attract electrons more than larger atoms, the electronegativity trend is to increase across the rows. Conversely, as atomic size increases down the groups, the larger atoms will be less electronegative, so the electronegativity trend is to decrease down the groups.