Periodic table
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Periodic Table. The most useful tool in the Lab. Early Organization. J.W. Dobereiner (1829) organized elements in triads Triad – three elements with similar properties (ex: Cl, Br, I) J.R. Newlands (1864) organized elements in octaves Octave – repeating group of 8 elements.

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Periodic Table

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Periodic Table

The most useful tool in the Lab


Early Organization

  • J.W. Dobereiner (1829) organized elements in triads

    • Triad – three elements with similar properties (ex: Cl, Br, I)

  • J.R. Newlands (1864) organized elements in octaves

    • Octave – repeating group of 8 elements


Development of the PeriodiceTable

  • Dmitri Mendeleev taught chemistry in terms of properties.

  • Mid 1800’s - molar masses of elements were known.

  • Wrote down the elements in order of increasing mass.

  • Found a pattern of repeating properties.


Mendeleev’s Table

  • Grouped elements in columns by similar properties in order of increasing atomic mass.

  • Found some inconsistencies - felt that the properties were more important than the mass, so switched order.

  • Also found gaps.

  • Must be undiscovered elements.

  • Predicted their properties before they were found.


The Modern Periodic Table

  • Henry Moseley – British physicist

  • Arranged elements according to increasing atomic number

  • The arrangement today

  • Symbol, atomic number & mass


The New Way

  • Elements are still grouped by properties.

  • Similar properties are in the same column.

  • Order is by increasing atomic number.

  • Added a column of elements (noble gases)

  • Weren’t found because they are unreactive.


Organization

  • Horizontal rows = periods

    • There are 7 periods

    • Each period represents an energy level

    • Every element in the same period has

      • the same # of energy levels and

      • the same core electron configuration


Organization

  • Vertical column = group or family

    • Similar physical & chemical prop.

    • Same # of valence electrons

    • Same common oxidation state

    • Identified by number & letter


  • Horizontal rows are called periods

  • There are 7 periods


  • Group 1A are the alkali metals

  • Group 2A are the alkaline earth metals


  • Group 7A is called the Halogens

  • Group 8A are the noble gases


  • These are called the inner transition elements, and they belong here

The group B are called the transition elements


1A

8A

2A

3A

4A

5A

6A

7A

  • The elements in the A groups are called the representative elements

    outer s or p filling


Lanthanides – the 4f orbital fills for these elements


Actinide series – the 5f orbitals are being filled for these elements.


Types of elements

  • Metals

  • Non-metals

  • Metalloids or semi-metals


Metals

  • Good conductor of heat and electricity

  • Malleable

  • Ductile

  • High tensile strength

  • High luster

  • Solid at room temperature

  • React by losing electrons


Nonmetals

  • Poor conductors of heat and electricity

  • React by gaining electrons

  • Some gases (O, N, Cl); some are brittle solids (S); one is a fuming dark red liquid (Br)


Semi-Metals

  • Heavy, stair-step line

  • Metalloids border the line

    • Properties intermediate between metals and nonmetals

    • Learn the general behavior and trends of the elements, instead of memorizing each element property

  • B, Si, Ge, As, Sb, Te


Families

  • Group IA – alkali metals

    • most reactive metals

    • Silvery in appearance

    • Soft

    • Combine easily with non-metals

    • Melting point is higher than the boiling point of water

    • Have 1 valence electron


Families

  • Group 2 – Alkaline Earth Metal Family

    • Harder, stronger, denser, higher melting point, and less reactive than alkali

    • Usually not found as free elements, but as compounds

    • Have 2 valence electrons


Families

  • Group 7 – Halogens

    • Most reactive family

    • Non-metals

    • Have seven valence electrons

  • Group 8 – Noble Gas

    • Inert, unreactive

    • Have full set of valence electrons


H

1

Li

3

Na

11

K

19

Rb

37

Cs

55

Fr

87

1s1

1s22s1

1s22s22p63s1

1s22s22p63s23p64s1

1s22s22p63s23p64s23d104p65s1

1s22s22p63s23p64s23d104p65s24d10 5p66s1

1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1


S- block

Alkali metals all end in s1

Alkaline earth metals all end in s2

really should include He, but it fits better later.

He has the properties of the noble gases.

s1

s2


He

2

Ne

10

1s2

1s22s22p6

1s22s22p63s23p6

1s22s22p63s23p64s23d104p6

1s22s22p63s23p64s23d104p65s24d105p6

1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6

Ar

18

Kr

36

Xe

54

Rn

86


p1

p2

p6

p3

p4

p5

The P-block


1

2

3

4

5

6

7

  • Each row (or period) is the energy level for s and p orbitals.


Areas of the periodic table

  • Group A elements = s & p blocks

  • representative elements

    • Wide range of phys & chem prop.


Transition Metals -d block

s1 d5

s1 d10

d1

d2

d3

d5

d6

d7

d8

d10


  • d orbitals fill up after previous energy level, so first d is 3d even though it’s in row 4.

1

2

3

4

5

6

7

3d


f6

f13

f1

f2

f3

f4

f5

f7

f8

f10

f12

f14

f11

f9

F - block

  • inner transition elements


1

2

3

4

5

6

7

  • f orbitals start filling at 4f

4f

5f


Atomic Size

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Radius

  • First problem: Where do you start measuring from?

  • The electron cloud doesn’t have a definite edge.

  • Atomic Radius = half the distance between two nuclei of a diatomic molecule.


Trends in Atomic Size

  • Influenced by three factors:

    1. Energy Level

    • Higher energy level is further away.

      2. Charge on nucleus

    • More charge pulls electrons in closer.

      3. Shielding effect

(blocking effect)


WHAT HAPPENS TO ATOMIC RADII?

  • Does a negative ion (anion) get larger or smaller?

  • Does a positive ion (cation) get larger or smaller?


Trends in Ionic Size

  • Cations form by losing electrons.

  • Cations are smaller than the atom they come from.

  • Metals form cations.

  • Cations of representative elements have noble gas configuration.


Ionic size

  • Anions form by gaining electrons.

  • Anions are bigger than the atom they come from.

  • Nonmetals form anions.

  • Anions of representative elements have noble gas configuration.


WHAT IS IONIZATION ENERGY?

  • The energy required to remove an electron

  • Which element has the highest ionization energy? Why?


What determines Ionization Energy?

  • The greater the nuclear charge, the greater IE.

  • Greater distance from nucleus decreases IE

  • All the atoms in the same period have the same energy level.

  • But, increasing nuclear charge

  • So IE generally increases from left to right.


Ionization Energy

  • The energy required to remove the first electron is called the first ionization energy

  • Thesecond ionization energy is the energy required to remove the second electron.

  • Always greater than first IE.

  • The third IE is the energy required to remove a third electron.

  • Greater than 1st or 2nd IE.


Driving Force

  • Full Energy Levels require lots of energy to remove their electrons.

  • Noble Gases have full orbitals.

  • Atoms behave in ways to achieve noble gas configuration.


WHAT IS ELECTRONEGATIVITY?

  • The ability of an atom to pull off an electron.

  • Which element has the highest electronegativity? Why?


Periodic Trend

  • Metals are at the left of the table.

  • They let their electrons go easily

  • Low electronegativity

  • At the right end are the nonmetals.

  • They want more electrons.

  • Try to take them away from others

  • High electronegativity.


Trends in Electron Affinity

  • The energy change associated with adding an electron to a gaseous atom.

  • Easiest to add to group 7A.

  • Gets them to full energy level.

  • Increase from left to right: atoms become smaller, with greater nuclear charge.

  • Decrease as we go down a group.


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