chapter 6 the periodic table & periodic law. section 6.1 development of the modern periodic...

Chapter 6The Periodic Table

&Periodic Law

Section 6.1Development of the Modern

Periodic Table

John Newlands• In 1864, noticed

when the elements were arranged in order of increasing atomic mass, their properties repeated every eight elements. – THE LAW OF

OCTAVES

Meyer & Mendeleev• In 1869, published almost identical versions

with the elements in order of increasing atomic mass and in columns with similar properties.

Mendeleev

• Mendeleev is given more credit than Meyer BECAUSE: – He published his table first– He better demonstrated his table

• Suggested some of the previously measured masses were incorrect

• Left blanks for not yet discovered elements

Mendeleev’s Predicted Properties of Ge

“eka Silicon” and Its Actual Properties

Table 8.1

PropertyPredicted Properties

eka Silicon Actual Properties Ge

atomic massappearancedensitymolar volumespecific heat capacityoxide formulaoxide densitysulfide formula and solubility

chloride formula (boiling point)

chloride density

72amugray metal5.5g/cm3

13cm3/mol0.31J/g*KEO2

4.7g/cm3

ES2; insoluble in H2O; soluble in aqueous (NH4)2SECl4; (<1000C)

1.9g/cm3

72.61amugray metal5.32g/cm3

13.65cm3/mol0.32J/g*KGeO2

4.23g/cm3

GeS2;insoluble H2O; soluble aqu (NH4)2SGeCl4; (840C)

1.844g/cm3

Was Mendeleev psychic????

• periodic law: when arranged by atomic # elements exhibit a periodic recurrence of similar properties– Quantum-mechanical model of atom

explains organization of table

Development of Periodic Table

Mosley• In 1913, using X-rays, he discovered a

unique number of protons in the nuclei of atoms for each element.

• Today the elements are arranged in order of increasing atomic number

• PERIODIC LAW– There is a periodic repetition of chemical

and physical properties of the elements when they are arranged in order of increasing atomic number

Arrangement of the Periodic Table

• Groups/Families– 18 vertical columns (↑↓)– Two Labeling Systems

1. Number-and-letter system- A through 8A columns (representative elements)- 1B through 8B short columns (transition elements)

2. Number system- 1-18

• Periods– 7 horizontal rows (↔)

PERIODS

GROUPS/FAMILIES

Arrangements of the Periodic Table

Metals• Shiny• Good conductors of heat and electricity• Malleable & Ductile• Generally Solid at room temperature

Group 1Alkali Metals

Group 2Alkaline Earth Metals

Groups 3-12Transition Metals

Lanthanide & Actinide GroupsInner Transition Metals

Nonmetals & Metalloids• Nonmetals

– Dull– Generally gases or brittle

solids at room temperature– Poor conductors of heat and

electricity

• Metalloids– Elements with physical and

chemical properties of both metals and nonmetals

– Rest on the “stair-step”

BB

SiSi

AsAs

TeTe

AtAt

GeGe

SbSb

PoPo

←←MetalsMetals

Nonmetals →Nonmetals →

Section 6.2Classification of Elements

Element PlacementWhy are elements put into groups/families together?

Because they have similar chemical properties

Why do elements have similar chemical properties?

Because they have the same number of valence electrons

Group 1 – Alkali Metals

Period 2 Lithium 1s22s1 [He]2s1

Period 3 Sodium 1s22s22p63s1 [Ne]3s1

Period 4 Potassium 1s22s22p63s23p64s1[Ne]4s1

ALL ELEMENTS IN GROUP 1 (ALKALI METALS) HAVE ONE VALENCE ELECTRON

Recurring pattern in e- configuration is basis for periodic behavior.

• Main group, group # = valence e- count– Valence e- responsible for chemistry

• Elements in same group behave similarly

Dot Diagrams for Representative Elements

Figure 8.12 A periodic table of partial ground-state electron configurations

Representative Elements• s-block elements

– Groups 1&2, hydrogen & helium– Valence electrons occupy outermost s sublevels only

• p-block elements– Groups 13-18 (except helium)– Valence electrons include a full outermost s sublevel and a

filled or partially filled p sublevel

Period number is equal to the principle energy level where the valence electrons are located

Transition Elements• d-block elements

– Groups 3-12– Valence electrons include a full outermost s

sublevel and a filled or partially filled d sublevel

The period number minus 1 equals the principle energy level where the valence electrons are located

Inner transition metals• f-block elements

– Lanthanide & Actinide Groups– Full or partially full outermost s sublevel, and full or

partially full outermost f sublevel

The period number minus 2 equals the principle energy level where the valence electrons are located.

Section 6.3Periodic Trends

Atomic Radius• Half the distance between two nuclei of

identical atoms that are chemically bonded together

• Down the group– atomic radius increases

• Across the period– atomic radius decreases

Atomic Radius Decreases

Ato

mic

Radiu

s In

crease

s

Figure 8.16

Periodicity of atomic radius

Practice Atomic Radius• Which has the larger atomic radii of the following?

B or Al Na or Mg F or Cl

• Which has the smaller atomic radii of the following?

H or He K or Cs N or Ne

• Circle the one with the largest atomic radius and underline the one with the smallest.

C, Si, Ge V, Cr, W N, Mg, Ca

Ionization Energy• The amount of energy required to remove

an electron from the atom (how tightly an atom holds on to its electrons)

• Down a group – ionization energy decreases

• Across a period– ionization energy increases

• Ionization Energy (IE)– Energy required for complete removal of 1

mole of e- from 1 mole of atoms• Atoms w/ low IE form cations (lose e-)• Atoms w/ high IE form anions (gain e-)

Trends in Atomic Properties

Na(g) Na+(g) + e- I1

Na+(g) Na2+(g) + e- I2

I1 < I2 < I3

Figure 8.18

8.4 Trends in Atomic Properties

• Greater IE, more difficult to remove e-

– Positive values, energy into atom• Larger atoms easier to ionize

Ionization Energy Increases

Ioniz

ati

on E

nerg

y D

ecr

ease

s

Figure 8.17

Periodicity of first ionization energy (IE1)

Practice Ionization Energy• Which has the greater ionization energy?

Ne or Ar N or O Sc or Ti

• Which has the smaller ionization energy?

Al, Si, P K, Rb, Sr Be, Mg, Ca

Ionic Radius• Octet Rule

– Atoms tend to gain, lose, or share electrons in order to achieve a full outer energy level (typically 8 are needed)

• Ion– An atom that has an overall charge due to the

gaining or losing of electrons

Figure 8.25

Main-group ions and noble gas configurations

Ionic Radius Comparisons• Metals have LOW ionization and

electron affinity– They lose electrons to form positively

charged ions– Positive charged ions are smaller than the

original atom

• Nonmetals have HIGH ionization energy and electron affinity– They gain electrons to form negatively

charged ions– Negatively charged ions are larger than

the original atom

Figure 8.29

• Cation smaller than parent– e- removed, other

e- feel greater Zeff

• Anion larger than parent– e- added, e-/e-

repulsions occupy more space

Trends in Properties of Monatomic ions

Ionic Radius Increases

Ionic

Radiu

s In

crease

s

FOR IONIC RADIUS… MUST FOLLOW METAL/NON-METAL RULES

Ionic Radius Practice• Which is the smaller of the two?

Lithium ion or Lithium atomChlorine ion or Chlorine atom

• Underline the following that will form a positively charged ion and circle the ones that will form a negatively charged ion.

Mg F Al CuBr N S K

• How will the radius of each of the above change when an ion is formed?

Mg F Al CuBr N S K

Electronegativity• The ability of an atom to attract electrons in a

chemical bond.

• Down the group– Electronegativity values decrease

• Across the period– Electronegativity values increase

*Noble gases are the exception to this rule.

Electronegativity Increases

Ele

ctro

negati

vit

y D

ecr

ease

s

Electronegativity Practice• Which has the greater electronegativity

value?

B or N Si or Sn Cr or W

• Which has the smaller electronegativity value?

Rb, Sr, Y Ga, In, Sn As, Se, S

CUMULATIVE REVIEW• Which has the smallest atomic radius between

Ga, In, & Tl? – Which has the highes ionization energy?

• Which is the smallest: an atom of sodium, an ion of sodium, or an atom of potassium?

• Which has the greatest electron affinity between zinc, arsenic, or bromine? Which has the lowest ionization energy?

Figure 8.21

8.4 Trends in Atomic Properties

Electronegativity Increases Atomic Radius Decreases

Ionic Radius Increases Electron Affinity Increases

Ionization Energy Increases

Ele

ctro

negati

vit

y D

ecr

ease

s A

tom

ic R

adiu

s In

crease

s

Ionic

Radiu

s In

crease

s

Ele

ctro

n A

ffinit

y D

ecr

ease

s

Ioniz

ati

on E

nerg

y D

ecr

ease

s