Unit 3: Electrons and the Periodic Table Periodic Table Trends Unit 3: Electrons and the Periodic Table

What patterns do you notice? Periodic Law When elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals. What patterns do you notice? Atomic #s 3, 11, 19 are all alkali metals

Atomic Radius Atomic Radius size of atom Atomic Radius © 1998 LOGAL Atomic Radius Average distance in an atom between the nucleus and the outermost electron

Atomic Radius Atomic Radius Increases to the LEFT and DOWN Smallest Fr biggest

Atomic Size Trend Atomic Size increases down a group Why larger going down? Adding more energy levels. Atomic Size decreases across a period Why smaller across? Increased nuclear charge (more protons) without additional energy levels pulls e- in closer. Greater Coulombic Attraction

Atomic Size & Radius Examples The closer you are to Francium, the larger you will be! Which is larger: Rb or Li Rb N or Ne N

Ionization Energy Ionization energy is the amount of energy needed to remove an electron. M + energy  M+1 + e- Electrons that are close to the nucleus are hard to remove because they are under a strong force of attraction

Ionization Energy Trend Ionization Energy Increases across a period Why? Valence electrons experience a greater nuclear force because they are closer to the nucleus. Smaller atoms have higher Ionization energy. Ionization Energy Decreases down a group. Why? Valence electrons removed are farther from the nucleus because they are in higher energy levels. Bigger atoms have lower Ionization energy.

Ionization Energy Trends Why opposite of atomic radius? In small atoms, e- are close to the nucleus where the attraction is stronger Small atoms have High IE Big Atoms have Low IE

Ionization Energy Lowest as you go DOWN and to the LEFT High IE Low IE Fr Low IE Bottom left elements (Metals) WANT to lose an electron to become more stable.

Which would have a higher Ionization energy, Sodium or Chlorine? Chlorine has higher IE. Chlorine is smaller and has a higher nuclear charge (more protons) = stronger hold on electron = higher energy to take it away. Also, remember – Na wants to lose an electron (it is a metal) and Cl wants to gain an electron (non-metal)

E. Ionization Energy First Ionization Energy He Ne Ar Li Na K

Ionization Energy Successive Ionization Energies Large jump in I.E. occurs when a CORE e- is removed. Mg 1st I.E. 736 kJ 2nd I.E. 1,445 kJ Core e- 3rd I.E. 7,730 kJ

Ionization Energy Successive Ionization Energies Large jump in I.E. occurs when a CORE e- is removed. Al 1st I.E. 577 kJ 2nd I.E. 1,815 kJ 3rd I.E. 2,740 kJ Core e- 4th I.E. 11,600 kJ

Electronegativity The ability of an atom to attract an electron. The smaller the atom, the more electronegative it is because of a greater nuclear force.

Electronegativity Trends Electronegativity Increases across a period. Why? Non-metals such as F, O and N want more electrons to complete their valence shell. Smaller atoms have greater nuclear charge and thus, more force to attract electrons. Exception: Noble gases are not included because they generally do not want to gain electrons. They are already stable.

Electronegativity Trends Electronegativity Decreases Down a Group Why? Atomic size increases and valence electrons are farther from the nucleus. More energy levels increases shielding. So the pull from the positive nuclear charge is less. In General: Non-Metals have high Electronegativities Metals have low Electronegativities

Electronegativity Trends Highest as you go UP and to the RIGHT towards Fluorine F Remember- Noble gases not included in this trend!

Ionic Radius Na  Na+ Ionic Radius Cations (+ ions) the ionic radius is smaller than the original atom. Why? There is an increased attraction for the fewer electrons that remain. Na  Na+

Ionic Radius For Anions (– ions) the ionic radius is larger than the original atom. Why? The nuclear attraction is less for an increased number of electrons. Extra electrons repel each other and spread out – larger!) © 2002 Prentice-Hall, Inc. Cl  Cl-1

Practice Which atom is larger H or He? Which atom has a greater ionization energy, Ca or Sr? Which atom is more electronegative, F or Cl? Hydrogen – Smaller nuclear charge Ca – smaller, less shielding, lower effective nuclear charge Fluorine – Smaller, less shielding with less energy levels, so easier to attract electron

Ca – lower nuclear charge Examples Which atom has the larger radius? Be or Ba Ca or Br Ba –more energy levels Ca – lower nuclear charge

Examples Which atom has the higher 1st I.E.? N or Bi Ba or Ne N Ne

S or S2- Al or Al3+ S2- Al Examples Which particle has the larger radius? S or S2- Al or Al3+ S2- Al

Alkali Metal Reactivity