chapter 5 the periodic law. history of the periodic table u 1869 – dmitri mendeleev published his...

Chapter 5Chapter 5The Periodic LawThe Periodic Law

History of the Periodic Table

1869 – Dmitri Mendeleev published his periodic table.

He arranged it by grouping together the elements that had similar properties, and by increasing atomic masses.

His periodic table left empty spaces for new elements that would be discovered.

Mendeleev’s List of elements in Russian Circa 1869

Periodic Table in English (Circa 1891)

Periodic Table circa 1898


1911 – Henry Moseley (a student of Ernest Rutherford)

He rearranged a few elements on the periodic table so that elements were arranged by increasing atomic number rather than by atomic mass.


1944 – Glenn T. Seaborg rearranged the periodic table to make it look like it does today.

He moved the Actinide Series and the Lanthanide Series elements to the bottom of the periodic table.

Periodic Table Circa 1944

Modern Periodic Table

Parts of the Periodic Table

The periodic table can be divided and labeled using several methods.

Elements are arranged:

Vertically columns are called Groups or Families

Horizontal rows are called Periods or Series



Metals Non-metals Metalloids

Color the following groups on your periodic table and make a key to show what your colors mean

•Metals

•Non-metals

•Metalloids (have properties of both metals and non-metals)

Periodic Families Alkali Metals Alkaline Earth Metals Halogens Noble gases

There are many families on the periodic table but we will focus on the main families in this class.

Color the following families on your periodic table and make a key to show what your colors mean

•Alkali Metals (Group 1)

•Alkaline Earth Metals (Group 2)

•Halogens (Group 17)

•Noble Gases (Group 18)


Main Group or Representative Elements

Transition Metals Inner Transition Metals

Color your periodic table and make a key: Your colors do not have to match the ones above

The Periodic Law

The physical and chemical properties of the elements are periodic functions of their atomic numbers

Periodic Trends (across the row)If you understand the trends on the periodic

table, you can predict almost anything about any element on the periodic table.

We will study:• Atomic Radii• Ionic Radii• Valence Electrons• Reactivity• Electronegativity• Electron Configuration

What are the trends among the elements for atomic radius?

The atomic radius of an element is one half of the distance between the nuclei of two atoms of the same element when the atoms are joined.In short; how big an atom is

– Atomic radii are often measured in picometers (pm).

Atomic Radii Group Trends (up or down a column)

Atomic radius increase as you move down because each atom has another energy level,

so the atoms get bigger as you go down or smaller as you go up.

HLi

Na

K

Rb

Atomic Radii Periodic Trends (from side to side across period)

As you go from left to right across a period, the radius gets smaller.

As you go from right to left across a period, the radius gets bigger.

Electrons are in same energy level.

Na Mg Al Si P S Cl Ar

Atomic Radii Periodic Trends (from side to side across period)

Continued As you go from left to right across the period the

atomic number is increasing, so in neutral atoms the electron number increases as well.

The more electrons the stronger the negative charge on the outside, and the more attracted the outside of the atom is to the positive center.

So the atoms get smaller.

Summary: trends in atomic radius

•Use one color pencil to draw the arrows and write the trend for atomic radius

•Remember if it decreases from left to right then it increases from right to left

•Make sure to make a key for this periodic table so you know which arrows are for which trends

CATIONS

Cations are positively charged ions 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 Radius Trends

ANIONS

Anions are negatively charged ions Anions form by gaining electrons. Anions are bigger than the atom they

come from. Nonmetals form anions. Anions of ‘main’ groups elements have

noble gas configuration.

Ionic Radii Group trends (up and down a column)

Adding an energy level with every period

Ions get bigger as you go down.

Just like atomic radius gets bigger as you increase the number of energy levels

Li1+

Na1+

K1+

Rb1+

Cs1+

Ionic Radii Periodic Trends(from side to side through period)

Across the period, nuclear charge increases so they get smaller.

Energy level changes between anions and cations.

Li1+

Be2+

B3+

C4+

N3- O2- F1-

Period trends in ionic radius From left to right across a period, two trends are

visible: (opposite when going right to left)– A gradual decrease in the size of the positive ions

(cations)– A decrease in the size of the negative ions (anions)

What are the trends among the elements for ionization energy?

The energy required to remove an electron from an atom is called ionization energy.

– This energy is measured when an element is in its gaseous state.

– The energy required to remove the first electron from an atom is called the first ionization energy.

Ionization Energy Group Trend(up and down the column)

As you go down a group the atomic radius increases so the valence electrons (-) are further away from the nucleus (+).

This makes the attraction between the nucleus and the valence electrons weaker.

Ionization Energy Group Trend(up and down the column)

It is easier to pull a valence electron out of its energy level the further away it is from the nucleus. This means it requires LESS energy.

Therefore as you move down a group ionization energy decreases.

As you move up a group ionization energy increases.

Ionization Energy Periodic Trend(from side to side through period)

In general, the first ionization energy of representative elements tends to increase from left to right across a period. This trend can be explained by the nuclear charge/positive charge in the nucleus.

– The nuclear charge increases across the period.– As a result, there is an increase in the attraction

of the nucleus for an electron.– Thus, it takes more energy to remove an

electron from an atom.

Summary: trends in ionization energy

•Use a new colored pencil to draw the ionization energy arrows on the same periodic table you labeled your atomic radii trend above



Electronegativity The tendency for an atom to attract

electrons to itself when it is chemically combined with another element.

High electronegativity means it pulls the electron toward it.

Electronegativity Group Trend(up and down the column)

The further down a group, the farther the electron is away from the nucleus, and the more electrons an atom has.

More willing to share. Low electronegativity.

Electronegativity Periodic Trend (from side to side across row)

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.

ElectronegativityElectronegativity Values for Selected Elements

H2.1

Li1.0

Be1.5

B2.0

C2.5

N3.0

O3.5

F4.0

Na0.9

Mg1.2

Al1.5

Si1.8

P2.1

S2.5

Cl3.0

K0.8

Ca1.0

Ga1.6

Ge1.8

As2.0

Se2.4

Br2.8

Rb0.8

Sr1.0

In1.7

Sn1.8

Sb1.9

Te2.1

I2.5

Cs0.7

Ba0.9

Tl1.8

Pb1.9

Bi1.9

Summary: Trends for Electronegativity

•Use a new colored pencil to draw the electronegativity arrows on the same periodic table you labeled your atomic radii and ionization energy trend above



Summary: periodic trends

Valence Electrons

Define: The electrons available to be lost, gained, or shared in the formation of compounds.

The electrons in the highest energy level

Valence Electrons

Periodic Trends:

Group 1 = 1 valence electron = 1+ Oxidation Number

Group 2 = 2 valence electrons = 2+ Oxidation Number

Group 13 = 3 valence electrons = 3+ Oxidation Number

Group 14 = 4 valence electrons = 4+/4- Oxidation Number

Group 15 = 5 valence electrons = 3- Oxidation Number



Group 18 = 8 valence electrons = 0 Oxidation Number

1+

H

Li

Na

Be

2+

Mg

B

3+

Al

C

4-4+

Si

N

3-

O

2-

F

1-

Ne

0

He

P S Cl

K

Rb

Cs

Fr

Ca

Sr

Ba

Ra

Ga

In

Tl

Ge

Sn

Pb

As Se

Ar

Br Kr

Sb Te I Xe

Bi Po At Rn

Reactivity

Reactivity increases as you go down the columns of metallic elements.

Reactivity decreases as you go down the columns of non-metallic elements.

Watch the video to see what that means.

Electron Configuration and the Periodic Table

Electron Configurationof Main Group Elements

Group Period #(+) Example 1 s1 Na = 3s1

2 s2 Ba = 6s2

13 s2p1 Ga = 4s24p1

15 s2p3 Sb = 5s25p3

17 18

s2p5 Br = 4s24p5

s2p6 Rn = 6s26p6

Electron ConfigurationTransition Elements

Period # s2 + Period # (-1) d1 – 10

Examples: Sc = 4s23d1

Zn = 4s23d10

Mo = Ir =

5s24d46s25d7

Metals-Physical/Chemical Nonmetals-Physical/Chemical

Good conductors of heat and electricity

Poor conductors of heat and electricity

Luster (shininess) No luster (dull) often colored

Easily lose electrons (form cations (+) ions)

Tends to gain or share electrons

Corrodes easily Brittle

Ductile and Malleable Non Ductile and Non Malleable

Solid at STP (except Hg)

Solids, liquids, and gases

chapter 5 the periodic law. history of the periodic table u 1869 – dmitri mendeleev published his...

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