I II III

Ch. 6 - The Periodic Table & Periodic Law

I. Development of the Modern Periodic Table(p. 174 - 181)

A. MendeleevDmitri Mendeleev (1869, Russian)

Organized elements by increasing atomic mass

Elements with similar properties were grouped together

There were some discrepancies

A. Mendeleev Deduced elements existed, but were

undiscovered elements, their properties could be predicted

B. MoseleyHenry Moseley (1913, British)

Organized elements by increasing atomic number

Resolved discrepancies in Mendeleev’s arrangement

This is the way the periodic table is arranged today!

C. Modern Periodic Table

1 2 3 4 5 6 7

Group (Family)Period

1. Groups/FamiliesVertical columns of periodic tableEach group contains elements with similar

chemical & physical properties (same amount of valence electrons in each column)

2 numbering systems exist: Groups # 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) Numbered 1 to 18 from left to right

2. PeriodsHorizontal rows of periodic table

Periods are numbered top to bottom from 1 to 7

Elements in same period have similarities in energy levels, but not properties

Main Group ElementsTransition MetalsInner Transition Metals

3. Blocks

3. Blocks

1 2 3 4 5 6 7

Lanthanides - part of period 6

Actinides - part of period 7

Overall Configuration

I II III

II. Classification of theElements(pages 182-186)

Ch. 6 - The Periodic Table

A. Metallic Character

1 2 3 4 5 6 7

MetalsNonmetalsMetalloids

1. MetalsGood conductors of heat and electricityFound in Groups 1 & 2, middle of table in

3-12 and some on right side of tableHave luster, are ductile and malleableMetallic properties increase as you go

from left to right across a period

a. Alkali MetalsGroup 1(IA)1 Valence electronVery reactive, form metal oxides

(ex: Li2O)Electron configuration

ns1

Lowest melting pointsForm 1+ ion: Cations

Examples: Li, Na, K

b. Alkaline Earth MetalsGroup 2 (IIA)2 valence electronsReactive (not as reactive as alkali metals) form

metal oxides (ex: MgO)Electron Configuration

ns2

Form 2+ ionsCations

Examples: Be, Mg, Ca, etc

c. Transition MetalsGroups 3 – 12 (IB – VIIIB) Reactive (not as reactive as Groups 1 or 2), can

be free elements Highest melting pointsElectron Configuration

ns2(n-1)dx where x is column in d-blockForm variable valence state ionsAlways form Cations

Examples: Co, Fe, Pt, etc

3. MetalloidsSometimes called semiconductorsForm the “stairstep” between metals and

nonmetalsHave properties of both metals and

nonmetalsExamples: B, Si, Sb, Te, As, Ge, Po, At

2. NonmetalsNot good conductorsUsually brittle solids or gases (1 liquid Br)Found on right side of periodic table –

AND hydrogenHydrogen is it’s own group, reacts rapidly

with oxygen & other elements (has 1 valence electron)

Nonmetal Groups/FamiliesBoron Group: IIIA typically 3 valence

electrons, also mix of metalloids and metalsCarbon Group: IVA typically 4 valence

electrons, also has metal and metalloidsNitrogen Group: VA typically 5 valence

electrons, also has metals & metalloidsOxygen Group: VIA typically 6 valence

electrons, also contains metalloids

a. HalogensGroup 17 (VIIA)Very reactiveElectron configuration

ns2np5

Form 1- ions – 1 electron short of noble gas configuration

Typically form salts (NaCl)Anions

Examples: F, Cl, Br, etc

b. Noble GasesGroup 18 (VIIIA)Unreactive, inert, “noble”, stableElectron configuration

ns2np6 full energy level Have an octet or 8 valence e-

Have a 0 charge, no ionsHelium is stable with 1s2, a duetExamples: He, Ne, Ar, Kr, etc

I II III

III. Periodic Trends(p. 187-194)

Ch. 6 - The Periodic Table

0

50

100

150

200

250

0 5 10 15 20Atomic Number

Atom

ic R

adiu

s (p

m)

Periodic LawWhen elements are arranged in order of

increasing atomic #, elements with similar chemical and physical properties appear at regular intervals.

0

50

100

150

200

250

0 5 10 15 20

Ato

mic

Rad

ius

(pm

)

Atomic Number

Atomic Radius size of atom

© 1998 LOGALIonization Energy

Energy required to remove an e- from a neutral atom

© 1998 LOGAL

Electronegativity

Properties of Atoms

Shielding EffectThere 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). As atoms add more protons the nuclear charge increases Atoms are also adding more e- which are attracted to the p+

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).

Periodic Trend, 1. Shielding effect increases down a group.2. Shielding effect remains constant across a period.

Atomic Radius = ½ the distance between two identical bonded atoms

1. Atomic Radius

1 2 3 4 5 6 7

Atomic Radius Increases to the LEFT and DOWN

1. Atomic Radius

Why larger going down? Higher energy levels have larger orbitals Shielding - core e- block the attraction between the

nucleus and the valence e-

Why smaller to the right? Increased nuclear charge(total charge of protons in

nucleus) without additional shielding pulls e- in tighter

1. Atomic Radius

The minimum energy required to remove an electron from the ground state of an isolated gaseous atom or ion. The ease with which an atom loses an e-.First Ionization Energy (IE1) = Energy required to remove one e- from a neutral atom.

Na(g) + IE1 (energy) → Na+(g) + e- ; +∆H (positive)Second Ionization Energy (IE2) = energy needed to remove a second electron, and so forth

Na+(g) + IE2 (energy) → Na2+ (g) + e- ; +∆H (positive)

2. Ionization Energy

1 2 3 4 5 6 7

First Ionization Energy Increases UP and to the RIGHT

2. Ionization Energy

Why does it increase up a group? The closer the e- are to the nucleus the more difficult it is to remove them Decreased shielding effect increases the positive nuclear charge

Why does it increase across a period? Atomic radius decreases

Positive nuclear charge increases pulling e- closer to the nucleus

2. Ionization Energy

Electron Affinity

Electron AffinityPeriodic Trend 1. Electron affinity slightly increases up a group.2. Electron affinity generally tends to increase across a period.Electron affinity increases up a groupdecreases the atomic radius taking the electrons closer to the nucleus’

positive attraction. less shielding effect increases the positive nuclear charge (+) as additional

shells are added and e- are held on tighter.Electron affinity increases across a period

atomic radius decreases effective positive nuclear charge increases steadily and the e- are

drawn closer to the nucleus making it easier to add e-

Electron AffinityElectron affinity increases up a groupdecreases the atomic radius taking the electrons closer

to the nucleus’ positive attraction. decreasing shielding effect increases the effective

positive nuclear charge (+) as additional shells are added and e- are held on tighter.

Electron affinity increases across a period atomic radius decreases effective positive nuclear charge increases steadily

and the e- are drawn closer to the nucleus making it easier to add e- to unfilled sublevels.

3. ElectronegativityThe measure of the ability of an atom in a chemical compound to

attract electrons Given a value between 0 and 4, 4 being the highestTendency for an atom to attract e- closer to itself when forming a

chemical bond with another atom.

1 2 3 4 5 6 7

Why increase as you move right? More valence electrons, need less to fill outer shell Increased nuclear charge

Why increase as you move up? Smaller electron cloud, more pull by + nucleus

3. Electronegativity

Which atom has the larger radius?

Be or

or Br

Examples

Ba

Ca

Which atom has the higher 1st I.E.?

or Bi

Ba or

Examples

N

Ne

Which element has the higher electronegativity?

Cl or

or Ca

Examples

F

Be

B. Chemical ReactivityMetals Period - reactivity decreases as you go from left to right across a period.

Group - reactivity increases as you go down a groupReact to form bases when combined with water Non-metals Period - reactivity increases as you go from the left to the right across a period.

Group - reactivity decreases as you go down the group. React to form acids when combined with water

C. Valence ElectronsValence Electrons

e- in the outermost s & p energy levels Stable octet: filled s & p orbitals (8e-) in one

energy level

1A2A 3A 4A 5A 6A 7A

8A

C. Valence ElectronsYou can use the Periodic Table to determine

the number of valence electronsEach group has the same number of valence

electrons Group #A = # of valence e- (except He)

1A2A 3A 4A 5A 6A 7A

8A