Chapter 5 - The Periodic Law

I: History of the Periodic Table

• by 1860, sixty elements had been discovered
  
  
• for each the properties and compounds had to be discovered
  
  
• no accurate way to determine the atomic mass or the number of atoms in a particular compound
  
  
• in 1860, a gathering of chemists in Germany, to settle problems dealing with atomic mass, etc.
• Cannizzaro - method for measuring the relative mass of an atom
  
  
• thus the relationship between relative mass and other properties could be pursued.

A: Mendeleev and Chemical Periodicity

Russian chemist, Dmitri Mendeleev, decided to use the new atomic masses in textbook he was writing.

planned to organize elements according to their properties

when the elements were arranged by increasing atomic mass, certain similarities in their chemical properties appeared at regular intervals

periodic - definition

Figure 5-1 page 123

His table resulted in elements with similar properties were grouped together and he published his finding in 1869.

Figure 5-2 page 124 - Mendeleev's original table

He used the atomic masses as a guide but he let the properties determine the exact grouping of elements.

e.g. Tellurium (Te) and Iodine (I)

Left empty spaces for elements he said were not yet discovered e.g. Scandium (Sc), Gallium (Ga), Germanium (Ge). When they were discovered, their properties were very similar to those predicted by Mendeleev.

His table was accepted by the scientific community as a valuable tool.

Two questions that remained were a) why could most of the elements be arranged in the order of increasing atomic mass, but a few could not? b) What was the reason for chemical periodicity?

B: Moseley and the Periodic Law

The first question above was answered by Henry Moseley, working with Ernest Rutherford. He made the connection that elements in the periodic table were arranged in increasing order according to nuclear charge, or the number of protons in the nucleus. (atomic number).

Moseley's ideas led to the modern definition of atomic number and to the idea that atomic number is the basis for the organization of the periodic table.

He explained the tellurium/iodine question using atomic numbers.

Periodic Law - statement

C: The Modern Periodic Table

Since Mendeleev, chemists have found new elements and synthesized new ones in the lab.

Periodic Table - definition

1. The Noble Gases

1894 - John Strutt and William Ramsay discovered argon

Escaped notice because of it was totally unreactive.

In 1868 helium had been discovered as a component of the sun - how?

1895 - Ramsay showed that helium existed on earth.

To place these two elements into the periodic table he proposed a new group, what is now group 18.

1898 - Ramsay discovered krypton and xenon.

1900 - Friedrich Dorn discovered radon

2. The Lanthanides

lanthanides - definition

All are very similar in chemical and physical properties, thus separating them and identifying them was tedious.

They are part of period 6.

3. The Actinides

actinides - definition

They are part of period 7.

Lanthanides and actinides are at the bottom to save space. They would normally fit between groups 3 and 4.

4. Periodicity

Figure 5-4 page 126

The differences in atomic numbers for the two groups follow the same pattern: between first and second element, between second and third element, between third and fourth element, etc.

The key is looking for the patterns.

Homework: Chapter 5: 5.1

II: Electron Configuration and the Periodic Table

Group 18 - octet - stable

Configuration of valence electrons determines the atom's chemical properties.

A: Period and Blocks on the Periodic Table

Vertical columns (group or family) have similar chemical properties.

Horizontal rows are series or periods.

Length of each period is determined by the number of electrons that can occupy the sublevels being filled in that period.

Table 5-1 page 128

1st period 2 elements filling s sublevel

2nd period: 8 elements filling s and p sublevel

3rd period: 8 elements filling s and p sublevel

4th period: 18 elements filling s and d and p sublevels

Figure 5-5 page 129

1. The s Block Elements: Groups 1 and 2

Both groups are active metals: group 1 being more active than group 2.

Group 1 = alkali metals

• Group 1: each element has one electron in an s sublevel.
• Tend to lose that electron to form a +1 ion and achieve an octet.
• In the elemental state (uncombined) they have a silvery appearance and can be cut with a knife.
• Not found in nature in the elemental (uncombined) state. Instead they react vigorously with nonmetals and form compounds - the way they are found in nature.
• React strongly with water to form hydrogen gas and an aqueous solution of alkalis (solutions of a base).
• Stored under kerosene to protect them from reacting with the moisture in the air.
• Moving down the column the melting points decrease.

Group 2 = alkaline earth metals

• each element has two electrons in an s sublevel
• tend to lose two electrons to form a +2 ion and achieve an octet.
• harder, denser, and stronger than alkali metals
• have higher melting points than alkali metals
• less reactive than alkali metals
• not found in nature in the elemental (uncombined) state.

2. Hydrogen and Helium

Hydrogen has one electron in the s sublevel of the first energy level.

Usually above but separated from Group 1 metals. Does not have same properties.

Properties of hydrogen do not resemble those of any other group.

Helium has two electrons in the s sublevel of the first energy level.

It is part of group 18 (noble gases). Its filled s sublevel gives it special stability like the other members of group 18.

Different than group 2 metals since they have an unfilled p sublevel of the energy level n -- have a filled s sublevel of energy level n.

Homework: Chapter 5: 5.2

3. The d-block Elements: Groups 3-12

After the ns sublevel is filled, electrons go into the (n-1)d sublevel until the d sublevel is filled. After the (n-1)d sublevel is filled the np sublevel is filled.

• d-block elements are metals are are called transition elements or transition metals.
• They represent a transition from metals on the left of the periodic table to nonmetals on the right.
• good conductors of electricity
• high luster
• less reactive than alkali metals and alkaline earth metals
• some are unreactive and exist in nature in the elemental (uncombined) form e.g. palladium, platinum and gold

Homework: Chapter 5: 5.3

4. The p-Block Elements: Groups 13-18

electrons enter the p sublevel only after the s sublevel of the same energy level is filled

p-block and s-block elements make up what is called the main-group elements

• properties vary
• on right side - nonmetals except hydrogen and helium
• the six metalloids (boron, silicon, germanium, arsenic, antimony, and tellurium) are p-block elements
• metals on on the left, bottom

Group 17 - Halogens

• most reactive nonmetals
• react vigorously with most metals to form a group of compounds called salts
• fluorine and chlorine are gases at room temperature
• bromine is a reddish liquid
• iodine is a dark purple solid that sublimes
• astatine is a synthetic element which is a solid

Metalloids or Semiconducting Elements

• on line separating nonmetals and metals in p-block -- stepped red line
• mostly brittle solids
• some properties of metals and some properties of nonmetals
• electrical conductivity between metals and nonmetals

metals of p block are

• generally harder and denser than the s-block alkaline earth metals, but softer and less dense than the d-block metals
• except for bismuth, found in nature in combined state
• when purified as pure element, they are stable in the presence of air

Homework: Chapter 5: 5.4

The f-Block elements: Lanthanides and Actinides

• between groups 3 and 4 in the sixth and seventh periods
• fourteen f-block elements in each series
• lanthanides are shiny metals similar in reactivity to the group 2 alkaline earth metals
• the actinides are all radioactive
• the first four of the actinides are found naturally on earth; the remaining actinides are made in the laboratory

Homework: Chapter 5: 5.5

III: Electron Configuration and Periodic Properties

This section deals with the relationship between the periodic law and the electron configurations of the elements.

A: Atomic Radii

atomic radius - definition

1. Periodic Trends

figure 5-13 page 141

figure 5-14 page 142

across second period - gradual decrease in atomic radii

reason - increasing positive charge of the nucleus and added electrons are entering same energy level

2. Group Trends

group 1 - atomic radii of the main group elements increase down a group

reason - added electrons are entering higher energy levels farther from nucleus - electrons don't feel added pull of nucleus

group 13 - gallium smaller than aluminum not larger

reason - gallium has 10 d electrons and therefore ten extra protons pulling the electron cloud closer to the nucleus

summary - interplay between the distance of the added electrons from the nucleus and the increased number of protons in the nucleus

Homework: Chapter 5: 5.6

B: Ionization Energy

If you supply enough energy to an atom (A) it will lose an electron and form a positive ion:

A + I.E. ---> A⁺ + e^-

Ion - definition

Ionization - definition

Want to compare ease with which atoms give up electrons so compare ionization energies.

ionization energy - definition

figure 5-15 page 143

figure 5-16 page 144

1. Period Trends

First element in series (alkali metal) has lowest I.E. and last element has highest I.E. (noble gases).

Low I.E. indicates element wants to lose the electron. - more stable after electron is lost

High I.E. Indicates element does not want to lose electron. - less stable if electron is lost.

Trend - increase in I.E. as you move across period.

Reason - increased nuclear charge and electrons entering same energy level (same distance from nucleus) means nucleus is holding electron more tightly and therefore harder to take electron from atom.

Metals tend to have lower I.E. than nonmetals.

2. Group Trends

Trend - for main group elements the I.E. decreases down the family.

Reason - as you move down the family the electron you are removing is in a different energy level and, therefore, farther from the nucleus, therefore, easier to remove. The farther from the nucleus the valence electrons are the more they are shielded by inner shell electrons and the less the valence electrons feel the pull of the nucleus.

3. Removing Electrons from Positive Ions

Can remove electrons from neutral atoms, +1 ions, +2 ions, +3 ions.

Referred to as first ionization energy (neutral atom), second ionization energy (+1 ion), third ionization energy (+2 ion), etc.

Second ionization energy always larger than first because you are removing an electron from a positive ion rather than a neutral atom. In the ion the protons are holding the fewer number of electron tighter than in the neutral atom where you have a one to one, proton to electron ratio.

Each successive ionization energy increases for the same reason.

table 5-3 page 145

First ionization energy of group 18 higher than the first ionization energy of any other group in the series - energy needed to break an octet.

Difference between first and second I.E. of Li.

Difference between second and third I.E. Of Be.

Difference between I.E.. Of berrylium and boron.

Difference between I.E.. Of nitrogen and that of oxygen.

Homework: Chapter 5: 5.7

C: Electron Affinity

Electron Affinity - definition

Most atoms release energy when they acquire an electron.

A + e^- ---> A^- + electron affinity (energy)

Some atoms need to absorb energy to accept the electron.

A + e^- + electron affinity (energy) ---> A^-

The ions formed are unstable and emit the electron spontaneously.

figure 5-17 page 147

In tables not equations: energy will have a negative value for exothermic reactions; energy will have a positive value for endothermic reactions

1. Periodic Trends

Halogens gain electron most readily - they have a large negative electron affinity.

Why does giving off a large amount of energy indicate the atom wants the electron?

Across the p block of any series, adding an electron produces greater negative values.

Not the case between group 14 and group 15. Why?

2. Group Trends

Trends are not as regular as for I.E.

Generally, down the group, adding the electron becomes more difficult.

Reason: increase in nuclear charge going down the group; increase in atomic radius down the group.

Size effect is the predominant factor.

Many exceptions in the transition metals.

3. Adding Electrons to Negative Ions

Hard to add an electron to a negative ion - repulsion. Thus, all second electron affinities are positive - the atom will not accept the electron without the addition of energy.

Homework: Chapter 5: 5.8

D: Ionic Radii

Cation - definition

A cation is formed when a neutral atom loses one or more electrons.

Radius of the ion is always smaller than the radius of the atom. Why?

Anion - definition

Formed when one or more electrons are added to a neutral atom.

Negative ion is always larger than the radius of the atom. Why?

1. Periodic Trends

Metals tend to form cations; nonmetals tend to form anions.

Cationic radii decrease across a period because the electron cloud shrinks due to the increasing nuclear charge acting on the electrons in the same energy level.

Beginning with group 15 - atoms assume stable noble gas configurations by gaining electrons.

Anionic radii decrease across the period from there. Same reason as with cations.

2. Group Trends

Gradual increase in ionic radius down the group - outer shell electrons are in higher energy levels.

Homework: Chapter 5: 5.9

E: Valence Electrons

Chemical compounds form when electrons are lost, gained, or shared between atoms.

Valence Electrons - definition

Usually found in incompletely filled main energy levels.

Main group elements - those electrons in s and p sublevel of outermost energy level.

Table 5-4 page 150

Elements in groups 13-18 have a number of valence electrons equal to the group number minus 10.

Homework: Chapter 5: 5.10

F: Electronegativity

The valence electrons hold atoms together in chemical compounds.

Valence electrons are not always mid-way between the two atoms. They are sometimes closer to one atom than another. This affects the chemical properties of the compound.

Electronegativity - definition

Element with highest electronegativity is fluorine. (note spelling)

Scale of electronegativity is arbitrary. Highest value is 4.0 given to fluorine. It is a relative scale.

1. Periodic Trends

figure 5-20 page 151

Electronegativity tends to increase going across the periods. Some exceptions.

The active metals are the least electronegative.

The active nonmetals are the most electronegative.

Down a group the electronegativities tend to decrease down a group or remain about the same. Why?

Some noble gases do not form compounds and so do not have electronegativities assigned.

Highest electronegativities are in the upper right of the periodic table and lowest are in the lower left.

Figure 5-21 page 152

Homework: Chapter 5: 5.11

G: Periodic Properties of the d and f Block Elements

d block elements vary less and with less regularity than the main group elements. e.g. flat curves in figure 5-14 and 5-16.

For d block elements, electrons in both the Ns and (n-1)d sublevel are available to interact with the surroundings. Electrons in the incompletely filled d sublevels are responsible for many characteristic properties of the d block elements.

1. Atomic Radii

generally decrease across a period

the decrease is less than that for main group elements

in figure 5-14 the radii decrease then increase slightly across each of the four periods that contain d block elements -- as the number of electrons in the d block sublevel increases, the radii increase because of repulsion among the electrons.

2. Ionization Energy

I.E. generally increase across the periods

I.E. increases going down group

s electrons are less shielded from nuclear charge by d electrons

3. Ion Formation and Ionic Radii

Order of removing electrons.

Highest energy level then highest energy sublevel where s < p < d < f

e.g. Fe

Most d block elements form 2+ ions because they lose the two electrons in the s sublevel.

Cations have smaller radii than atoms.

4. Electronegativity

d block elements have electronegativities between 1.1 and 2.54.

Values increase as radii decreases and vice versa.

Homework: Section Review page 154 #1-3

end of notes

Periodic means recurring at regular intervals. back

Periodic Law state the physical and chemical properties of the elements are periodic functions of their atomic numbers. Back

The periodic table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group. Back

The lanthanides are the fourteen elements with atomic numbers from 58 to 71.back

The actinides are the fourteen elements with atomic numbers from 90 to 103.back

Ionization energy is the energy required to remove on electron from a neutral atom of an element. Back

An ion is an atom or group of bonded atoms that has a positive or negative charge.back

Any process that results in the formation of an ion is referred to as ionization.back

Atomic Radius is one-half the distance between the nuclei of identical atoms that are bonded together. Back

Electron affinity is the energy change that occurs when an electron is acquired by a neutral atom.back

A cation is any positive ion. Back

An anion is any negative ion. Back

Valence electrons are those electrons available to be lost, gained, or shared in the formation of chemical compounds. Back

Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons. Back