Honors unit 3-4 periodic trends

Honors unit 3-4 periodic trends Chap 8

Sizes of Atoms • Radius of a free atom / ion = distance from nucleus to outermost electron shell. • The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

Effective Nuclear Charge (Zeff) • In a many-electron atom, electrons are both attracted to the nucleus (p#) and repelled by other electrons (core e#). • Shielding effect: Electrons within level and from previous levels block the effects of the (+) nucleus • Zeff is the actual attraction experienced by the valence electron

Atomic Size or Ionic Size Influenced by three factors: 1. Energy Level (n number) • Higher energy level is further away. 2. Effective nuclear charge Zeff= Z - S • Z is the atomic number and S is the number of core electrons. • More Zeffcharge pulls electrons in closer. 3. Outermost electron number • More outermost e- repel each other.

Table of Atomic Radii

Atomic Radius Trend • Radius decreases across a period • Increased effective nuclear charge due to decreased shielding • Radius increases down a group • Each row on the periodic table adds a “shell” or energy level to the atom • Radius of transition metals • increase in size down the Group • atomic radii of transition metals roughly the same size across the same period

N or F, N is further left • N or F • C or Ge • N or Al • Al or Ge? opposing trends • N or F • C or Ge, Ge is further down • N or F • C or Ge • N or Al, Al is further down & left Example 8.5 – Choose the Larger Atom in Each Pair

Sizes of Ions • Ionic size depends upon: • Energy level (n#) • Effective nuclear charge Zeff. • Outermost e# : repulsion to each other.

Sizes of Ions • Cations are smaller than their parent atoms. • The outermost electron is removed and shell# are less.

Sizes of Ions • Anions are larger than their parent atoms. • Electrons are added and repulsions are increased.

Sizes of Ions • Ions increase in size as you go down a column. • Due to increasing value of n.

Sizes of Ions • In an isoelectronic series, ions have the same number of electrons. • Isoelectronic ionic size decreases with an increasing Zeff. • Who are isoelectronic? Li+, O2-, Mg2+

Trends in Ionic Radius • Ions in same group have same charge • Ion size increases down the group • more valence shells (n), larger • Cations smaller than neutral atom; Anions bigger than neutral atom • Cations smaller than anions in the same period • except Rb+1 & Cs+1 bigger or same size as F-1 and O-2 • The radii of isoelectronic ions decrease as the Z increases.

Homework • Atomic radius: Page 359 55, 56, 59, 60, 61 • Ion radius: Page 359 63 a and b, 64 a, b and c, 68, 69, 70

Ionization Energy (IE) • Definition: the energy required to remove an electron from an atom, starting from the valence e-. • first ionization energy = energy to remove electron from neutral atom • Bigger atomic size means less attraction to the e-, therefore less IE needed to remove the e- • Tends to increase across a period • As radius decreases across a period, the electron you are removing is closer to the nucleus and harder to remove • Tends to decrease down a group • Outer electrons are farther from the nucleus and easier to remove

Al or S, Al is further left • Al or S • As or Sb, Sb is further down • Al or S • As or Sb • N or Si, Si is further down & left • Al or S • As or Sb • N or Si • O or Cl? opposing trends Example 8.8 – Choose the Atom in Each Pair with the Higher First Ionization Energy Tro, Chemistry: A Molecular Approach

Trends in Electron Affinity/Electronegativity • Definition: A measure of the ability of an atom in a chemical compound to attract electrons • Electronegativity tends to increase across a period • As radius decreases, electrons get closer to the bonding atom’s nucleus • Electronegativity tends to decrease down a group • As radius increases, electrons are farther from the bonding atom’s nucleus (more shells)

Periodic Trend:Electronegativity

Summary of Periodic Trends

Homework • Page 359 72, 73, 74, 77

Group trends

Metallic Character • Metals • conduct heat and electricity • most oxides basic and ionic • form cations in solution • lose electrons in reactions – oxidized - + charged • Nonmetals • electrical and thermal insulators • most oxides are acidic and molecular • form anions and polyatomic anions • gain electrons in reactions – reduced - - charged • metallic character increases left • metallic character increase down

Tro, Chemistry: A Molecular Approach

Sn or Te • P or Sb • Ge or In, In is further down & left • Sn or Te • P or Sb • Ge or In • S or Br? opposing trends • Sn or Te, Sn is further left • Sn or Te • P or Sb, Sb is further down Example 8.9 – Choose the More Metallic Element in Each Pair Tro, Chemistry: A Molecular Approach

Trends in the Alkali Metals • S block, having only 1 valence e-, 1+ charge when forming ion. • atomic radius increases down the column  electron affinity/ electron negetivity decreases down the column ionization energy decreases down the column • very low ionization energies • good reducing agents, easy to oxidize • very reactive, not found uncombined in nature • react with nonmetals to form salts • Replace H in water and form strong base AOH

2A + 2H2O  2AOH + H2

Trends in the Halogens • P block, having only 7 valence e-, 1- charge when forming monatomic ion. • atomic radius increases down the column  electron affinity decreases down the column ionization energy decreases down the column • Very high electron affinities • good oxidizing agents, easy to reduce • very reactive, not found uncombined in nature • reactivity increases down the column • react with hydrogen to form HX acids • melting point and boiling point increases down the column

Trends in the Noble Gases • atomic radius increases down the column • Full-shell electron configuration • Very high ionization energy. IE decreases down the column • very unreactive • only found in monatomic form in nature • Also called “inert” gas • Very low boiling points • all gases at room temperature

Homework • Page 360 79, 81, 85, 87

Honors unit 3-4 periodic trends