IIIIII Ch. 6 - The Periodic Table & Periodic Law I. Development of the Modern Periodic Table (p )

A. Mendeleev zDmitri Mendeleev (1869, Russian) yOrganized elements by increasing atomic mass yElements with similar properties were grouped together yThere were some discrepancies

A. Mendeleev zDmitri Mendeleev (1869, Russian) yPredicted properties of undiscovered elements

B. Moseley zHenry Moseley (1913, British) yOrganized elements by increasing atomic number yResolved discrepancies in Mendeleev’s arrangement yThis is the way the periodic table is arranged today!

C. Modern Periodic Table zGroup (Family) zPeriod

1. Groups/Families zVertical columns of periodic table zNumbered 1 to 18 from left to right zEach group contains elements with similar chemical properties

2. Periods zHorizontal rows of periodic table zPeriods are numbered top to bottom from 1 to 7 zElements in same period have similarities in energy levels, but not properties

zMain Group Elements zTransition Metals zInner Transition Metals 3. Blocks

Lanthanides - part of period 6 Actinides - part of period 7 Overall Configuration

IIIIII II. Classification of the Elements (pages ) Ch. 6 - The Periodic Table

1. Metals zGood conductors of heat and electricity zFound in Groups 1 & 2, middle of table in 3-12 and some on right side of table zHave luster, are ductile and malleable

a. Alkali Metals zGroup 1 z1 Valence electron zVery reactive zElectron configuration yns 1 zForm 1 + ions zCations yExamples: Li, Na, K

b. Alkaline Earth Metals zGroup 2 zReactive (not as reactive as alkali metals) zElectron Configuration yns 2 zForm 2 + ions zCations yExamples: Be, Mg, Ca, etc

c. Transition Metals zGroups zReactive (not as reactive as Groups 1 or 2), can be free elements zElectron Configuration yns 2 (n-1)d x where x is column in d-block zForm variable valence state ions zCations yExamples: Co, Fe, Pt, etc

2. Nonmetals zNot good conductors zFound on right side of periodic table – AND hydrogen zUsually brittle solids or gases

a. Halogens zGroup 17 (7A) zVery reactive zElectron configuration yns 2 np 5 zForm 1 - ions – 1 electron short of noble gas configuration zAnions yExamples: F, Cl, Br, etc

b. Noble Gases zGroup 18 zUnreactive, inert, “noble”, stable zElectron configuration yns 2 np 6 full energy level zHave a 0 charge, no ions zExamples: He, Ne, Ar, Kr, etc

3. Metalloids zSometimes called semiconductors zForm the “stairstep” between metals and nonmetals zHave properties of both metals and nonmetals zExamples: B, Si, Sb, Te, As, Ge, Po, At

C. Valence Electrons zoutermost s & p orbital electrons zStable octet - filled s & p orbitals (8 e - ) in one energy level zGroup #A = # of valence e - (except He) 1A 2A 3A 4A 5A 6A 7A 8A

C. Valence Electrons zValence electrons = yelectrons in outermost orbitals (highest principle energy level) zYou can use the Periodic Table to determine the number of valence electrons zEach group has the same number of valence electrons 1A 2A 3A 4A 5A 6A 7A 8A

IIIIII III. Periodic Trends (p ) Ch. 6 - The Periodic Table

A. Periodic Law zWhen elements are arranged in order of increasing atomic #, elements with similar chemical and physical properties appear at regular intervals.

B. Chemical Reactivity zFamilies ySimilar valence e - within a group result in similar chemical properties

zAtomic Radius ysize of atom © 1998 LOGAL zIonization Energy yEnergy required to remove an e - from a neutral atom © 1998 LOGAL zElectronegativity C. Other Properties

Shielding Effect zThere is a Nuclear charge experienced by the outer (valence) electron(s) in a multi-electron atom is due to the difference between the charge on the nucleus and the charge of the core electrons (inner electron shells). z-Results in the reduction of attractive force between the positive nucleus and the outermost electrons due to “shielding effect” of the inner electron shells(core electrons). zPeriodic Trend, 1. Shielding effect increases down a group. 2. Shielding effect remains constant across a period.

zAtomic Radius = ½ the distance between two identical bonded atoms 1. Atomic Radius

zAtomic Radius yIncreases to the LEFT and DOWN 1. 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 nuclear charge without additional shielding pulls e - in tighter 1. Atomic Radius

zThe minimum energy required to remove an electron from the ground state of an isolated gaseous atom or ion. zThe ease with which an atom loses an e -. zFirst Ionization Energy = Energy required to remove one e - from a neutral atom. 2. Ionization Energy K Na Li Ar Ne He

zFirst Ionization Energy yIncreases UP and to the RIGHT 2. Ionization Energy

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 - 2. Ionization Energy

zSuccessive Ionization Energies yMg1st I.E.736 kJ 2nd I.E.1,445 kJ Core e - 3rd I.E.7,730 kJ yLarge jump in I.E. occurs when a CORE e - is removed. 2. Ionization Energy

yAl1st I.E.577 kJ 2nd I.E.1,815 kJ 3rd I.E.2,740 kJ Core e - 4th I.E.11,600 kJ zSuccessive Ionization Energies yLarge jump in I.E. occurs when a CORE e - is removed. 2. Ionization Energy

Electron Affinity zMost atoms can attract e - to form negatively charged ions zThe energy change that occurs when an e - is added to a gaseous atom or ion. zThe ease with which an atom gains an e -. zFor most atoms, the energy released when an e - is added. (in kJ/mol) zPeriodic Trend 1. Electron affinity slightly decreases down a group. 2. Electron affinity generally tends to increase across a period.

3. Electronegativity zThe measure of the ability of an atom in a chemical compound to attract electrons zGiven a value between 0 and 4, 4 being the highest zTendency for an atom to attract e - closer to itself when forming a chemical bond with another atom.

zWhy increase as you move right? yMore valence electrons, need less to fill outer shell zWhy increase as you move up? ySmaller electron cloud, more pull by + nucleus 3. Electronegativity

Ionic Radius zThe size atoms become when losing or gaining electrons. zPositive Ions – Metal - Atoms that lose e - and form positive ions become smaller. zThe lost e - is a valence e - and the atom may lose a shell.The repulsion between the remaining e - is lessened and allows the effective positive nuclear charge to pull the remaining e - closer.

zNegative Ions – Nonmetal - Atoms that gain e - and form negative ions become larger. zThe repulsion between the added e - and existing e - is increased and the effective positive nuclear charge cannot hold onto the e - tightly.

zWhich atom has the larger radius? yBe orBa yCa orBr Examples

zWhich atom has the higher 1st I.E.? yNorBi yBa orNe Examples

zWhich element has the higher electronegativity? yCl or F yBe or Ca Examples

More Practice zAnswer questions on page 189 and on page 194.