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

I II III

Ch. 6 - The Periodic Table & Periodic Law

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

A. Mendeleev

Dmitri 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. Moseley

Henry 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/Families

Vertical 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. Periods

Horizontal 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. Metals

Good 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 periodForm bases in water

a. Alkali Metals

Group 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 Metals

Group 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 Metals

Groups 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. Metalloids

Sometimes called semiconductorsForm the “stairstep” between metals and

nonmetalsHave properties of both metals and

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

2. Nonmetals

Not 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)

Form acids in water

Nonmetal Groups/Families

Boron Group: IIIA typically 3 valence electrons, also mix of metalloids and metals

Carbon Group: IVA typically 4 valence electrons, also has metal and metalloids

Nitrogen Group: VA typically 5 valence electrons, also has metals & metalloids

Oxygen Group: VIA typically 6 valence electrons, also contains metalloids

a. Halogens

Group 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 Gases

Group 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

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 Electrons

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

energy level

1A

2A 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)

1A

2A 3A 4A 5A 6A 7A

8A

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

Ato

mic

Ra

diu

s (

pm

)

Periodic Law

When 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

Ra

diu

s (

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 Effect

There 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

Successive Ionization Energies

Mg 1st I.E. 736 kJ

2nd I.E. 1,445 kJ

Core e- 3rd I.E. 7,730 kJ

Large jump in I.E. occurs when a CORE e- is removed.

The greater the IE the more difficult it is to remove an electrons

2. Ionization Energy

Al 1st I.E. 577 kJ

2nd I.E. 1,815 kJ

3rd I.E. 2,740 kJ

Core e- 4th I.E. 11,600 kJ

Successive Ionization Energies

Large jump in I.E. occurs when a CORE e- is removed.

2. Ionization Energy

Electron Affinity

Electron Affinity

The greater the attraction between a given atom and an added e-, the more negative the atom’s EA. Halogens’ ns2p5 have the most negative EA.

Noble Gases have EA > 0; as do Be, Mg, & N because e- have to enter previously unoccupied, higher energy orbitals, an unfavorable energy state.

Periodic Trend

1. Electron affinity slightly increases up a group.

2. Electron affinity generally tends to increase across a period.

Electron Affinity

Electron 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. Electronegativity

The 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

Ionic Radius

The size atoms become when losing or gaining electrons.

Positive Ions – Metal - Atoms that lose e- and form positive ions become smaller.

The 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.

Negative Ions – Nonmetal - Atoms that gain e- and form negative ions become larger.

The repulsion between the added e- and existing e- is increased and the effective positive nuclear charge cannot hold onto the e- tightly.

Periodic Trend

1. Ionic Radius increases down a group.

2. Ionic radius tends to gradually decrease across a period for the positive ions, then beginning in group VA or VIA the much larger negative ions also gradually decreases

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