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 yDeduced elements existed, but were undiscovered elements, their properties could be predicted

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 zEach group contains elements with similar chemical & physical properties (same amount of valence electrons in each column) z2 numbering systems exist: yGroups # I through VIII with ea. # followed by A or B A groups are Main Group Elements (s&p electrons) B groups are Transition Elements (d electrons) yNumbered 1 to 18 from left to right

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

A. Metallic Character zMetals zNonmetals zMetalloids

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 zMetallic properties increase as you go from left to right across a period

a. Alkali Metals zGroup 1(IA) z1 Valence electron zVery reactive, form metal oxides (ex: Li 2 O) zElectron configuration yns 1 zLowest melting points zForm 1+ ion: Cations yExamples: Li, Na, K

b. Alkaline Earth Metals zGroup 2 (IIA) z2 valence electrons zReactive (not as reactive as alkali metals) form metal oxides (ex: MgO) zElectron Configuration yns 2 zForm 2 + ions zCations yExamples: Be, Mg, Ca, etc

c. Transition Metals zGroups 3 – 12 (IB – VIIIB) zReactive (not as reactive as Groups 1 or 2), can be free elements zHighest melting points zElectron Configuration yns 2 (n-1)d x where x is column in d-block zForm variable valence state ions zAlways form Cations yExamples: Co, Fe, Pt, 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

2. Nonmetals zNot good conductors zUsually brittle solids or gases (1 liquid Br) zFound on right side of periodic table – AND hydrogen zHydrogen is it’s own group, reacts rapidly with oxygen & other elements (has 1 valence electron)

Nonmetal Groups/Families zBoron Group: IIIA typically 3 valence electrons, also mix of metalloids and metals zCarbon Group: IVA typically 4 valence electrons, also has metal and metalloids zNitrogen Group: VA typically 5 valence electrons, also has metals & metalloids zOxygen Group: VIA typically 6 valence electrons, also contains metalloids

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

b. Noble Gases zGroup 18 (VIIIA) zUnreactive, inert, “noble”, stable zElectron configuration yns 2 np 6 full energy level yHave an octet or 8 valence e- zHave a 0 charge, no ions zHelium is stable with 1s 2, a duet zExamples: He, Ne, Ar, Kr, etc

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

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

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

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). yAs atoms add more protons the nuclear charge increases yAtoms are also adding more e - which are attracted to the p + zResults 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(total charge of protons in nucleus) 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 (IE 1 ) = Energy required to remove one e - from a neutral atom. zNa(g) + IE 1 (energy) → Na + (g) + e - ; +∆H (positive) zSecond Ionization Energy (IE 2 ) = energy needed to remove a second electron, and so forth zNa + (g) + IE 2 (energy) → Na 2+ (g) + e - ; +∆H (positive) 2. Ionization Energy

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

zWhy does it increase up a group? yThe closer the e- are to the nucleus the more difficult it is to remove them yDecreased shielding effect increases the positive nuclear charge zWhy does it increase across a period? yAtomic radius decreases yPositive nuclear charge increases pulling e- closer to the nucleus 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) z Cl(g) + e - → Cl — (g) + EA (kJ/mol) ; -∆H (negative)

Electron Affinity zPeriodic Trend 1. Electron affinity slightly increases up a group. 2. Electron affinity generally tends to increase across a period. zElectron affinity increases up a group zdecreases the atomic radius taking the electrons closer to the nucleus’ positive attraction. zless shielding effect increases the positive nuclear charge (+) as additional shells are added and e- are held on tighter. zElectron affinity increases across a period yatomic radius decreases yeffective positive nuclear charge increases steadily and the e - are drawn closer to the nucleus making it easier to add e -

Electron Affinity zElectron affinity increases up a group zdecreases the atomic radius taking the electrons closer to the nucleus’ positive attraction. zdecreasing shielding effect increases the effective positive nuclear charge (+) as additional shells are added and e- are held on tighter. zElectron affinity increases across a period yatomic radius decreases yeffective positive nuclear charge increases steadily and the e - are drawn closer to the nucleus making it easier to add e - to unfilled sublevels.

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 yIncreased nuclear charge zWhy increase as you move up? ySmaller electron cloud, more pull by + nucleus 3. Electronegativity

zWhich atom has the larger radius? yBe or y or Br Examples Ba Ca

zWhich atom has the higher 1st I.E.? yorBi yBa or Examples N Ne

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

B. Chemical Reactivity zMetals zPeriod - reactivity decreases as you go from left to right across a period. Group - reactivity increases as you go down a group zReact to form bases when combined with water zNon-metals zPeriod - reactivity increases as you go from the left to the right across a period. Group - reactivity decreases as you go down the group. zReact to form acids when combined with water

C. Valence Electrons zValence Electrons ye - in the outermost s & p energy levels yStable octet: filled s & p orbitals (8e-) in one energy level 1A 2A 3A 4A 5A 6A 7A 8A

C. Valence Electrons zYou can use the Periodic Table to determine the number of valence electrons zEach group has the same number of valence electrons zGroup #A = # of valence e - (except He) 1A 2A 3A 4A 5A 6A 7A 8A