IIIIII Ch. 5 - The Periodic Table C. Johannesson

Mendeleev zDmitri Mendeleev (1869, Russian) yOrganized elements by increasing atomic mass. yElements with similar properties were grouped together. xRepeating patterns are referred to as periodic. yPredicted properties of undiscovered elements..

Moseley zHenry Mosely (1913, British) y Organized elements by increasing atomic number. yResolved discrepancies in Mendeleev’s arrangement. yThe Periodic Law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers.

The Periodic Table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group. The Modern Periodic Table Visual Concept

IIIIII II. Organization of the Elements Ch. 5 - The Periodic Table C. Johannesson

Metallic Character zMetals zNonmetals zMetalloids

Blocks zMain Group Elements zTransition Metals zInner Transition Metals

Periodic Patterns C. Johannesson © 1998 by Harcourt Brace & Company s p d (n-1) f (n-2) yPeriod of an element can be determined from elements electron configuration yThe name of each block (s,p,d &f) determined by what sublevel is being filled

Periodic Patterns zPeriod # yenergy level (subtract for d & f) zA/B Group # ytotal # of valence e - zColumn within sublevel block y# of e - in sublevel

Periodic Patterns zExample - Hydrogen s-block1st Period 1s 1 1st column of s-block

Sample Problem B zAn element has the electron configuration [Kr]4d 5 5s 2. Without looking at the periodic table, identify the period, block, and group in which this element is located. Then, consult the periodic table to identify this element and the others in its group. Periods and Blocks of the Periodic Table,

IIIIII C. Johannesson III. Periodic Trends Ch. 5 - The Periodic Table

Trends zGenerally electron configuration of atom’s highest occupied energy level govern’s atom’s properties zVertical groups share similar chemical properties zHorizontal rows/periods, length is determined by number of e- that can occupy sublevels being filled

zAtomic Radius ysize of atom Increases to the LEFT and DOWN Atomic Radius

zWhy larger going down? yHigher energy levels have larger orbitals yShielding - core e - block the attraction between the nucleus and the valence e - zWhy smaller to the right? yIncreased positive nuclear charge without additional shielding pulls e - in tighter Atomic Radius

Ionization Energy zElectron can be removed from an atom if enough energy is supplied zIon- an atom or group of bonded atoms that has a charge +/- zIonization- process that results in formation of an ion zIonization energy (IE) energy required to remove one electron from a neutral atom yCan compare the ease with which atoms give up electrons © 1998 LOGAL

zFirst Ionization Energy yIncreases UP and to the RIGHT Ionization Energy

z Ease of e- loss is major reason for high reactivity zWhy opposite of atomic radius? yIn small atoms, e - are close to the nucleus where the attraction is stronger zWhy small jumps within each group? yStable e - configurations don’t want to lose e - Ionization Energy

C. Johannesson zIonic Radius yCations (+) xlose e - xsmaller © 2002 Prentice-Hall, Inc. yAnions (–) xgain e - xlarger Ionic Radius

yVisual ConceptVisual Concept Ion

zMelting/Boiling Point yHighest in the middle of a period. yIncrease as you go down Melting/Boiling Point

zWhich atom has the larger radius? yBeorBa yCaorBr Ba Ca Examples

zWhich atom has the higher 1st I.E.? yNorBi yBaorNe N Ne Examples

zWhich atom has the higher melting/boiling point? yLiorC yCrorKr C Cr Examples

zWhich particle has the larger radius? ySorS 2- yAlorAl 3+ S 2- Al Examples

IIIIII C. Johannesson III. Periodic Trends Continued Ch. 5 - The Periodic Table

Electron Affinity The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity. Electron affinity generally increases across periods. Increasing nuclear charge along the same sublevel attracts electrons more strongly Electron affinity generally decreases down groups. The larger an atom’s electron cloud is, the farther away its outer electrons are from its nucleus.

yVisual ConceptVisual Concept Electron Affinity

Electronegativity zMeasure of the ability of an atom in chemical compound to attract e- yThere is an un even concentration of charge in a compound yEffects chemical properties yFluorine most electronegative element yGroups 1 & 2 least electronegative elements yTends to increase across periods decrease or remain the same down

yVisual ConceptVisual Concept Electronegativity

Example zOf the elements gallium, Ga, bromine, Br, and calcium, Ca, which has the highest electronegativity? Explain your answer in terms of periodic trends. zBromine should have the highest electronegativity because electronegativity increases across the periods.

Valence Electrons Chemical compounds form because electrons are lost, gained, or shared between atoms. The electrons that interact in this manner are those in the highest energy levels. The electrons available to be lost, gained, or shared in the formation of chemical compounds are referred to as valence electrons. Valence electrons are often located in incompletely filled main-energy levels. example: the electron lost from the 3s sublevel of Na to form Na + is a valence electron.

yVisual ConceptVisual Concept Valence Electrons