Unit 3: The Unit 3: The Periodic TablePeriodic Table

Chapter Chapter 66

Lesson1: Development of the Periodic Table

1) History of the Periodic Table – By the end of the 1700’s scientists had identified only 30 elements (ex. Cu, Ag, Au, H2, N2, O2, C). By the mid 1800’s by using spectroscopy additional elements were identified by using their line spectra and about 65 elements had been identified.

A. J.W. Dobereiner: 1829Organized the elements

into groupsgroups with similar propertiesproperties.He called these groups triadstriads. The middle element is often the average of the other two.

Ex) Cl – 35.5 Br – 79.9 I – 126.9CaAvg SrBa

Cl + I Avg.2

Triads on the Periodic Table

B. J.A.R. Newlands: 1867• Law of OctavesOctaves.

He said properties repeated every 88thth element. There were 62 known elements at the time. He was also a musician.

C. Dimitri Mendeleev: 18691. Organized the 1st

periodic table according to increasing atomic massatomic mass and put elements with similar properties in the same columncolumn.

2. He arranged some elements out of atomic mass orderorder to keep them together with other elements with similar propertiesproperties. He also left three blanksblanks in his table and correctly predictedpredicted the properties of these 3 unidentified elements that were later identified and matched his predictions.

D. Moseley: 1915Each element has a certain amountof positive charge in the nucleuswhich are called protonsprotons.1. Moseley reorganized the

periodic table by Atomic Atomic NumberNumber.

2. The Periodic Law: When elements are arranged in order of increasing atomic atomic numbernumber, their physical and chemical propertiesproperties show a periodicperiodic pattern.

Glenn Seaborg “Seaborgium”

Sg #106• Born in 1912 in Michigan, Seaborg proposed reorganizing the Periodic Table one last time as a young chemist working on the Manhattan Atomic Bomb Project during WWII by pulling the “f-block” elements out to the bottom of the table. He was the principle or co-discoverer of 10 transuranium elements. He was awarded the Noble prize in 1951 anddied in 1999.

6-2: Reading the Periodic Table

A. Information in each square:

B. Organizing the Squares1. Vertical Columns – groups & families2. Horizontal Rows - periods

13Al

Aluminum26.9815

Atomic # (protons)

SymbolName

Atomic MassWeighted average of all an element’s isotopes

Mass # - protons & neutrons

particular isotope of aluminum like

Al-26 or Al-27

Parts of the Periodic Table

E. Metal, Nonmetals and Metalloids (Semimetals):

1. Metals

2. Nonmetals

3. Metalloids - Properties of both metals & nonmetals

- good conductors of heat & electricity- most solids at room temperature (except Hg)- high luster (shiny)- ductile (can be drawn into thin wire)- malleable (bends without breaking)- high melting points- high densities- react with acids

- brittle (shatters when struck)- low luster (dull)- neither ductile nor malleable- nonreactive with acids- nonconductors

(Semimetals)

C. Electron Configuration & Families

1. Valence electrons – outermostoutermost electrons responsible for bonding.

2. Elements in the same groupgroup have the same number of valence electrons.

Carbon has 4 valence electrons

Atomic Families:

Alkali Metals*Alkaline Earth Metals*Nobel Gases*Halogens*Oxygen FamilyCarbon FamilyNitrogen FamilyBoron FamilyTransition Metals*

*Valance electrons?*Physical Properties?*Chemical Properties?

http://www.learner.org/interactives/periodic/groups.html

6-3: Trends in the Periodic Table

A. Atomic Radius1. The distance from the center of the the center of the

nucleus to the nucleus to the outermost outermost electronelectron.2. Atoms get largerlarger going down a group

and smallersmaller going across a period.

Ex) Na is larger than Mg Na is smaller than K

Atomic Radii of the Representative Elements

Atomic Radii vs Atomic Number

Ionic Size1. When atoms gaingain

electrons, they become (-) and get larger.

Positive Ion Size1. When atoms

loselose electrons, they become (+) and get smallersmaller.

2. Ions get largerlarger as you go down a group.

Relative Sizes of Positive &

Negative Ions

The sodium ion lost an electron, and therefore the positive

protons in the nucleus exert a stronger pull on the remaining negative electrons, shrinking

the orbitals. Thus positive ions are smaller than their atoms.

The chloride ion gained an electron, and therefore the fewer positive protons in the nucleus

exert a weaker pull on the extra negative electrons, increasing the size of the orbitals. Thus negative ions are larger than

their atoms.

Electron Attraction in a Bond &

Ion Size

C. Ionization Energy:1. The energy needed to removeremove one

electron from an atom.2. Elements that do not want to lose

their electrons have highhigh ionization energies.

3. Elements that easily lose electrons have lowlow ionization energies.

4. I.E. decreasesdecreases down a group (opposite of atomic radius)

5. I.E. increasesincreases across a period. (opposite of atomic radius)

Ex) Na IE smaller than Mg Na IE larger than K

Ionization Energy of the 1st 20 Elements

Ionization Energy vs. Atomic Number

Sublevels by the Periodic Table

Which element would have the lowest ionization energy?Which element would have the lowest ionization energy?

Which element would have the highest ionization energy?Which element would have the highest ionization energy?

Will the Lithium ion lose Will the Lithium ion lose any more electrons?any more electrons?

D. Successive Ionization Energies:

1. Energy required to removeremove electrons beyond the 1st electron.2. Ionization energies will increaseincrease for every electron removed.3. Na [Ne]3s1 Na• 1st = ____ kJ 2nd = ____ kJ4. Mg [Ne]3s2 Mg: 1st = ____ kJ 2nd = ____ kJ 3rd = ____kJ5. Al [Ne]3s23p1 Al: 1st = ____kJ 2nd = ____kJ 3rd = ____kJ 4th = ___kJ73

8

4560

496 145

0773057

71814

2742

11,600

E. Electron Affinity:

1. Energy change that occurs when an atom gainsgains an electron.

2. A highly negativenegative electron affinity attracts electrons. (nonmetals)

3. A positivepositive electron affinity does not attract electrons. (metals)

Electron Affinity

decreases

F. Electronegativity:

1. Reflects an atoms ability to attractattract electrons in a chemical bond.

2. E.N. decreasesdecreases going down a group3. E.N. increasesincreases going across a period.4. Examples: NaCl and H2

G. Octet Rule:1. An atom will tend to loselose, gaingain or

shareshare electrons in order to acquire a full set of valence electrons.

2. “Octet” = 8 = s2p6 configuration

H. Oxidation Number:

The charge on an ion when it gains orloses electrons to acquire a stable

octet.

Which of the following would have the largest?

• Atomic Radii?• Ionization Energy?• Electron Affinity?

Element DElement C Element A

Which of the following would have the smallest?

• Atomic Radii?• Ionization Energy?• Electron Affinity?

Element BElement D Element D

Electron Shielding