Periodic Table

Periodic Table

Chapter 6

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What do I know?

• On the back of the blank periodic table write down at least 3 pieces of information you can get from the periodic table.

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A Brief History…

• Joseph Proust• Law of Definite Composition

– elements combine in definite proportions by weight

• The weight of one element that combines with the weight of another element = combining weight

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• Joseph Berzelius • [1807 - 1818 ] • Determined the combining weights of 43

elements with oxygen.• Recognized similarities of certain

elements...– similar metallic properties– similar reactive

properties– Li, Na, K – similar nonmetals– Cl, Br, I,

• “TRIPLETS”

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• Johann Wolfgang Dobereiner • 1829 • mathematician • discovered that combining weight

of middle triplet is the average [or near average]of the combining weights of the other two

• Li, Na, K

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• Jean Stas • 1860 • confirmed Proust theory of definite

composition• established accurate atomic

weight of the known elements

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Was there a relationship between the weight of an element and its properties?

• John A.R. Newlands • 1865 • arranged elements in order of

atomic weight• elements with similar properties

were 7, or multiple of 7 apart• Law of Octaves

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Dimitri Mendeleev 1869

• developed a chart -listed elements by increasing atomic weight

• grouped elements with similar properties in the same row

• Left Gaps where no element fit the pattern.

• Predicted discovery of new elements• Predicted properties of new elements

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Mendeleev’s Table

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.

“The properties of elements are in periodic dependence of their atomic weights.”

Dimitri Mendeleev

Old Periodic Law

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ALTERATIONS and ADDITIONS

• Sir Wm.Ramsay• 1890’s • Discovered Ne, Ar, Kr, Xe• Helium and Radon disc. Previously• New row added to Periodic Table

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• Henry Gwyn-Jeffreys Mosley• 1914-1915 • Number of protons determined

– atomic number - identifies what an element is

• Periodic Table Rearranged– elements arranged by increasing

atomic number– similar elements put in columns

instead of rows

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Modern Periodic Law

• “The properties of elements are in periodic dependence of their atomic numbers.”

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ARRANGEMENT OF THE MODERN PERIODIC TABLE

• A horizonal row on the periodic chart is refered to as either a period, or a series.

• A vertical column on the periodic chart is refered to as either a group, or a family.

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Element Location Electron DotProperties on Chart Notation

• H Grp 1 H . Colorless gas• Li Grp 1 Li . Soft; silver highly reactive• Na Grp 1 Na . Soft; silver highly reactive• K Grp 1 K . Soft; silver highly reactive• Rb Grp 1 Rb . Soft; silver highly reactive• Cs Grp 1 Cs . Soft; silver highly reactive• Fr Grp 1 Fr . Soft; silver • most reactive

metal

Alkali MetalsAlkali Metals

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Element Location Electron DotProperties on Chart Notationc

• Be Grp 2 Be : Reactive metal• Mg Grp 2 Mg : Reactive

metal• Ca

Grp 2 Ca : Reactive metal• Sr

Grp 2 Sr : Reactive metal

• BaGrp 2 Ba : Reactive metal

• RaGrp 2 Ra : Most

reactive metal of group

Alkaline Earth MetalsAlkaline Earth Metals

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What pattern(s) do we see?

• All elements in groups have same electron dot structure.

• Group placement predicts valence.• Groups usually have similar

properties.• Most reactive metals at the bottom

of the group.

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• B Grp 3 B :

nonmetal; black solid

• Al Grp 3 Al:

Metal • Ga

Grp 3 Ga : Metal• In

Grp 3 In : Metal• Tl

Grp 3 Tl : Most reactive metal

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• C Grp 4 C : black→clear

solid

• Si Grp 4 Si :

Metalloid • Ge

Grp 4 Ge : Metal• Sn

Grp 4 Sn : Metal• Pb

Grp 4 Pb : Most reactive

metal

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• N• Grp 5 N : gas; nonmetal

• P • Grp 5 P : nonmetal• As• Grp 5 As : Metalloid• Sb• Grp 5 Sb : Metalloid• Bi• Grp 5 Bi : Metal

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• O• Grp 6 O : gas; nonmetal

reactive• S • Grp 6 S : Nonmetal• Se• Grp 6 Se :

Nonmetal• Te• Grp 6 Te :

Nonmetal• Po• Grp 6 Po : Metal

Chalcogen Family

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• F• Grp 7 : F : gas; most

reactivenonmetal

• Cl • Grp 7 :Cl : gas;

reactive nonmetal

• Br• Grp 7 :Br : liquid; reactive

nonmetal• I• Grp 7 : I : solid; reactive

nonmetal• At• Grp 7 :At : solid; reactive

nonmetalHalogen Family

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• He• Grp 8 He : inert; nonmetal• Ne • Grp8 :Ne : inert; nonmetal • Ar• Grp 8 :Ar : inert;

nonmetal• Kr• Grp 8 :Kr : inert; nonmetal• Xe• Grp 8 :Xe : inert; nonmetal• Rn

Grp 8 :Rn : inert; nonmetal

Noble Gases / Inerts

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What pattern(s) do we see?• All elements in groups have same

electron dot structure.• Group placement predicts valence.• Groups usually have similar

properties – (exception: steps)• Most reactive nonmetals at the top

of the group.• Most reactive metals at the bottom

of the group.

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I spy with my little eye an element with…

• 3 energy levels and 2 valence electrons• Mg• 5 energy levels and 4 valence electrons• Sn• 2 energy levels and 8 valence electrons• Ne• 1 valence electron and 5 energy levels• Rb• 1 valence electron and 7 energy levels• Fr

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• 4 energy levels and 7 valence electrons• Br• 3 energy levels and 5 valence electrons• P• 2 valence electrons and 4 energy levels• Ca• 3 valence electrons and 2 energy levels• B• 8 valence electrons and 5 energy levels• Xe


• The heaviest halogen…• At (astatine)• The triplet with the average atomic weight of

35.5…• Cl• The least reactive Chalcogen• Po (polonium)• The group that fills the s2 valence orbital• Alkaline Earth Metals• A third period metalloid• Si

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Bonding

• See interactive

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Types of Bonding

• Ionic– Electrons transfer from one atom to

another creating + and – ions.

• Covalent– Atoms share electrons to create a

molecule.

• Metallic– Many atoms share electrons

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Types of Bonding• Ionic

– Electrons transfer from one atom to another – creating + and – ions.

+ -

e-

+ energy

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Ionization Energy

• The energy required to remove the outermost e- in an atom.

Ionization Energy

0

500

1000

1500

2000

2500

0 10 20 30 40

Atomic Number

Ioni

zati

on E

nerg

y (k

J/m

oL)

Series1Hydrogen

Helium

Lithium

Neon

Sodium

Argon

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• Why are some e- removed more easily?– Electrons that are farther away from the nucleus and

that have more E levels between them and the nucleus

• Low ionization energy – characteristic of METALS.

• High ionization energy – characteristic of NONMETALS.

• Removing successive electrons is more difficult, but follows the same overall pattern.

• Na + Energy Na+ + e- 119 Kcal / mol

• Na+ + Energy Na++ + e- 1090 Kcal/ mol

• Na+++ Energy Na++++ e- 1652 Kcal/ mol

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The energy released / absorbed when an electron is accepted by a neutral atom

+ -

e-

+ energy

Ionization E removes e- and forms + ion

Electron affinity is the E released when the neutral atom accepts the freed e-

and becomes -

linked

Electron Affinity

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Electron Affinity

Electron Affinity

-400

-300

-200

-100

0

100

0 10 20 30 40

Atomic Number

Elec

tron

Affi

nity

(kJ/

moL

)

Series1

Lithium

Fluorine

Sodium

Chlorine

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Electron Affinity

Increases across a periodDecreases

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• For atoms that have - valences:• Atom + e- A- + E• exothermic - energy released• (electron affinity)• stable product• Atom + e- + E A-• endothermic - energy required• unstable product

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• Covalent– Atoms share electrons to create a

molecule.Shared e-’s

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Electronegativity

• the attraction of an atom for a shared pair of electrons

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ElectronegativityElectronegativity

00.5

11.5

22.5

33.5

44.5

0 10 20 30 40

Atomic Number

Elec

tron

egati

vity

Series1

Lithium

Sodium

FluorineChlorine

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Table of Electronegativities

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Electronegativity

• Types of Covalent Bonds:• pure covalent - relatively even sharing

of e-• polar covalent - uneven sharing of e-•

• 0 - .5 ....... pure covalent• .5 - 1.7.....polar covalent• > 1.7 ....... ionic bond

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Atomic Radius [size]Atomic Radius

0

50

100

150

200

250

0 10 20 30 40

Atomic Number

Ato

mic

Rad

ius

(pm

)

Series1

LithiumSodium

Fluorine

Chlorine

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Atomic Radius [size]

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• Down a group – E levels are added.

• Across a period – Increased attraction between the E levels

and the nucleus causes the size to decreasedecrease.

• Pauli Repulsion Theory• As the number of electrons

increases so does the repulsion between the electrons; this may help account for the irregular increase in the radii.

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Ions [size]• Increases down a Group• Decreases across a Period

• Metal atoms lose electrons– become positive (cation) – Cations are SMALLER than the atoms

from which they come.• Nonmetal atoms gain electrons

– become negative (anion) – Anions are LARGER than the atoms

from which they come.

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DensityDensity

-10123456789

10

0 10 20 30 40

Atomic Number

Den

sity

(g/c

m^3

)

Series1

Sodium

Boron

FluorineLithium

Aluminum

Chlorine

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Density

• Here the density of each period is graphed individually

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Density

• Generally, density– increases down a group– Increases across the metals in a

period, and then decreases across the nonmetals

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M.P and B.P.Melting Point/Boiling Point

-500

0

500

1000

1500

2000

2500

3000

3500

0 5 10 15 20 25 30 35 40

Atomic Number

Tem

pera

ture

°C

Melting Point

Boiling Point

Lithium

Boron

Neon

Aluminum

Sodium Argon

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M.P. and B.P.

• Generally, like density, M.P. and B.P.– increases down a group– Increases across the metals in a

period, and then decreases across the nonmetals

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Configurations and the Periodic Table

• Electrons that reside in the outermost shell of an atom are called valence electrons.

– These electrons are primarily involved in chemical reactions.

– Elements within a given group have the same “valence shell configuration.”

– This accounts for the similarity of the chemical properties among groups of elements.

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• The following slide illustrates how the periodic table provides a sound way to remember the Aufbau sequence.– In many cases you need only the configuration of

the outer electrons.– You can determine this from their position on the

periodic table.– The total number of valence electrons for an

atom equals its group number.

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s1

s2 p1 p2 p3 p4 p5

p6

s2

d1 d2 d3 d4 d5 d6 d7 d8 d10d9 SS

fillinfillingg

P

filling

5d

f filling

** 6d

4d

1

3d

2

3

4

5

6

7

*

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The Elements

• The Elements in Song

Periodic Table

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