Chapter 5 – Periodic Law

Honors Chemistry, Chapter 5Page 1

Chapter 5 – Periodic Law


Mendeleev’s Periodic Table

• In 1869, a Russian chemist, Dmitri Mendeleev published the first periodic table

• Mendeleev arranged the elements by properties rather than by atomic mass

• His procedure left several empty spaces where elements were predicted to fill in when they were discovered


Mendeleev’s Periodic Table


Moseley and the Periodic Law

• Henry Moseley, working with Ernest Rutherford, examined the x-ray spectra of 38 elements and discovered that Mendeleev’s order was by charge in the nucleus rather than atomic mass.

• This confirmed Mendeleev’s principle of chemical periodicity.


Periodic Law

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

• This lead to the formation of our periodic table.


Periodic Table

• The periodic table is arranged with elements appearing in order of their atomic number so that elements with similar properties fall in the same column or group.


Families of Elements

• The noble gases, group 18.

• The halogens (F, Cl, Br, I, At), group 17.

• The lanthanides (Ce through Lu)

• The actinindes (Th through Lr)


Periodicity

Atomic # Difference Atomic #

He 2 3 Li

Ne 10 11 Na

Ar 18 19 K

Kr 36 37 Rb

Xe 54 55 Cs

Rn 86 87 Fr

8

8

18

18

32


Chapter 5, Section 1 Review1. How were Mendeleev and Moseley

involved in the development of the periodic table?

2. Describe the modern periodic table.3. Explain how the periodic law can be

used to explain the physical and chemical properties of elements.

4. Explain how the elements belonging to a group are interrelated in terms of atomic number.


s-Block Elements

• The group 1 and group 2 elements are the s-block.

• The group 1 elements are called the alkali metals (H, Li, Na, K, Rb, Cs, Fr).

• The group 2 elements are called the alkaline earth metals (Be, Mg, Ca, Sr, Ba, Ra) (note that He is NOT in group 2)

• Note that H is a special case and really doesn’t fit in either group 1 or 17.


Sample Problem 5-1

• What is the element with the electron configuration: [Xe]6s2 ?

The highest level with electrons is n=6. The second element is Ba.

• What is the electron configuration of the group 1 element in the 3rd period?

[Ne]3s1 is sodium.


d-Block Elements

• d-Block elements are located in group 3 to group 12.

• These have properties of metals are called the transition elements.

• Although the s level is not always filled, the sum of the s and d electrons equals the group number.

• Exceptions: Pd is [Kr] 4d10 5s0

Pt is [Xe] 4f14 5d9 6s1


Sample Problem 5-2

• An element has the electron configuration [Kr] 4d5 5s1. What is the period? Block? Group? What is the element?

period = 5

block = d

group = 6

element = Mo


p-Block Elements

• The p-block elements are in groups 13 through 18 where the p orbitals are filled.

• The p-block and the s-block are called the main-group elements.

• The p-block contains all non-metals, all metalloids (B, Si, Ge, As, Sb, and Te) and a few metals.

• Group 17 is called the halogens• Group 18 is called the Noble gases.


Group Number, Blocks, Electron Configurations

Group Number Group Configuration Block1,2 ns1,2 s

3-12 (n-1)d1-10ns0-2 d13-18 ns2np1-6

p


Sample Problem 5-3

• What is the outer electron configuration of the period 2 element in group 14?

[He]2s2 2p2

Carbon


Sample Problem 5-4

• For the electron configurations given:– Name the block and group– Identify the element as metal, non-metal or

metalloid– Describe it as likely to be high or low

reactivity

a. [Xe] 4f14 5d9 6s1 c. [Ne]3s2 3p6

b. [Ne] 3s2 3p5 d. [Xe] 4f6 6s2


Chapter 5, Section 2 Review

1. Describe the relationship between electrons in the sub-levels and the length of each period of the periodic table.

2. Locate and name the four blocks of the periodic table. Explain the reasons for these names.

3. Discuss the relationship between group configurations and group numbers.


Review continued

4. Describe the locations in the periodic table and the general properties of the alkali metals, the alkaline-earth metals, the halogens, and the noble gases.


Atomic Radii

• The atomic radius is defined as one-half the distance between the nuclei of identical atoms that are bonded together.

• The trend to smaller atoms across a period is caused by the increasing positive charge.

• In general, the atomic radii of main group elements increase down a group.


Sample Problem 5-5

• Of the elements Mg, Cl, Na, and P which has the largest atomic radius? Why?

Na Sodium

• Of the elements Ca, Be, Ba, and Sr which has the largest atomic radius? Why?

Ba Barium


Ionization Energy

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

• Ionization is any process which results in the formation of an ion.

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

A + energy A+ + e-


Trends in Ionization Energy• In general, in main-group elements,

ionization energy increases across a period due to the increase in nuclear charge.

• Among main-group elements, ionization generally decreases down the group. Moving down a group the electrons are in higher orbitals are partially shielded from the nucleus making them easier to remove.


Removing Electrons from Positive Ions

• The energy to remove the first electron from a neutral atom is the first ionization potential, IE1

• The energy to remove the second electron is the second ionization potential, IE2

• The energy to remove the third electron is the third ionization potential, IE3

• Etc.


Trends In Ionization Energy

• In general, it is harder to remove the second electron than the first because of the stronger effective nuclear charge (the nuclear charge minus the electron shielding)

• Once an ion reaches a noble gas configuration, removal of the next electron requires a much higher input of energy.


Ionization Energies for Period 3

IE in KJ/mol Na Mg Al Si

IE1 496 738 578 787

IE2 4562 1451 1817 1577

IE3 6912 7733 2745 3232

IE4 9544 10540 11578 4356

IE5 13353 13628 14831 16091


Sample Problem 5-6

• Element A has an ionization energy of 419 KJ/mol. Element B has an ionization energy of 1000 KJ/mol. For each element, is it likely to be in the s-block or the p-block? Which element is more likely to form a positive ion?

Element A: s-block, likely to form a + ion

Element B: p-block


Electron Affinity• Electron affinity is energy change that

occurs when an electron is acquired by a neutral atom.

A + e- A- + energy• In general, electron affinities increase

(become more negative) from left to right across the periodic table.

• With less regularity, electron affinity generally decreases down the periodic table.


Ionic Radii

• A positive ion is called a cation.

• A negative ion is called an anion.

• Cationic and anionic radii decrease from left to right across the periodic table due to the increase in nuclear charge.

• In general ionic radii increase down a group since the outer electrons are further from the nucleus.


Valence Electrons

• The electrons in an atom which are available to be lost, gained, or shared are called valence electrons.

• Valence electrons are often located in incompletely filled main-energy levels (s and p electrons)


Electronegativity

• Electronegativity is an arbitrary scale invented by Linus Pauling

• Electronegativity is the measure of how strongly an atom attracts the electrons in a molecule.

• Electronegativity increases from left to right on the periodic table

• Electronegativity decreases from top to bottom on the periodic table


Sample Problem 5-7

• Among Ga, Br, and Ca, which has the highest elecronegativity? Explain in terms of periodic trends.

All are in the 4th period.

Br because Br is farthest right. Electronegativity increases from left to right.


Properties of d-Block Elements

• The properties of d-block elements (all metals) vary less and with less regularity than those of the main-group elements.

• Atomic radii generally decrease from left to right across the periods of the d-block because of the increase in nuclear charge.

• d-Block Ionization energies and electronegativities generally increase from left to right across the periods.


Chapter 5, Section 3 Review

1. Define atomic and ionic radii, ionization energy, electron affinity, and electronegativity.

2. Compare the periodic trends of atomic radii, ionization energy, and electronegativity, and state the reasons for these variations.


Review Continued

3. Define valence electrons, and state how many are present in atoms of each main group element.

4. Compare the atomic radii, ionization energies and electronegativities of the d-block elements with those of the main group elements.

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Chapter 5 – Periodic Law