Chapter 3. Elements, Atoms, Ions, and the Periodic Table

 

The Periodic Law and the Periodic Table

In the early 1800's many elements had been discovered and found to have different properties. In 1817 Dφbreiner's triads –with regularly varying properties:  (Mg, Ca, Ba) (F, Cl, Br) and (S Se Te).1865:  Newlands – "law of octaves", about 55 elements: pattern of reactivity follows after 8 elements. However, no one had found a clear "order" in their properties until  Mendeleev, Dmitri (1834-1907) arranged 63  then known elements in the order of increasing atomic mass in a periodic table  and showed some chemical properties would reappear periodically. In certain cases, he placed a lighter slightly heavier element before a lighter element so that the chemical properties of the vertical columns would be preserved. Even though in a different  and much less clear form Meyer, Lothar (1830–1895) also  came up with a graph showing periodic properties similar to Medeleev. 

In Mendeleev's table, there was a gap. He purposely left blank position in his table so that the consistent vertical columns with the same chemical properties would be preserved. These missing elements were later discovered.

 

 The periodic law is an organized "map" of the elements that relates their structure to their chemical and physical properties. The periodic table is the result of the periodic law, and provides the basis for prediction of such properties as relative atomic and ionic size, ionization energy, and electron affinity, as well as metallic or non‑metallic character and reactivity.

The modern periodic table exists in several forms. The most important variation is in group numbering. The tables in the text use the two most commonly accepted numbering systems.

 

Numbering Groups in the Periodic Table

Periods and Groups

Periods are the horizontal rows of elements in the periodic table; the columns represent groups or families.

Elements in a vertical group have similar chemical properties. The vertical groups are currently named by numbers ranging from 1 to 18. An older way to identify the vertical groups is to use a Roman number and the capital letters A or B. Vertical groups of main group elements (or representative elements) were given a Roman numeral plus the letter A. Vertical groups of transition elements were given a Roman numeral plus the letter B.

 Representative elements are elements that always lose or gain the same number of electrons in chemical reactions.
Transition elements are elements that can lose or gain variable numbers of electrons in chemical reactions.

The lanthanide series and the actinide series are parts of periods 6 and 7, respectively, and groups that have been named include the alkali metals, the alkaline earth metals, the halogens, and the noble gases. Group A elements are called representative elements; Group B elements are transition elements. Metals, metalloids, and nonmetals can be identified by their location on the periodic table.

These groups are number from 1 - 18, left to right and groups have their Roman numbers and A or B classification..

 

Name

Elements

Common Valence

 

 

 

Electron Configuration

Group 1 (IA)  -

Alkali metal:

Li, Na, K Rb, Cs, Fr

ns1

Group 2 (IIA) -

Alkaline earth metals:

Be, Mg, Ca, Sr, Ba, Ra

ns2

Group 13 (IIIA) -

No specific name

B, Al, Ga, In, Tl

ns2 3p1

Group 14 (IVA) -

No specific name

C, Si, Ge, Sn, Pb

ns2 3p2

Group 15 (VA) -

No specific name

N, P, As, Sb, Bi

ns2 np3

Group  16 (VIA) -

No specific name

O, S, Se, Te,  Po

ns2 np4

Group 17 (VIIA) -

Halogens:

Cl, Br, I, At

ns2 np5

Group 18 (VIIIA) -

Noble gases:

He, Ne, Ar, Kr, Xe, Rn

ns2 np6

 

In addition to groups in the periodic table there are three blocks of elements called transition elements (which are labeled with B), Lanthanides and Actinides ( placed bottom of the table.

Metals, Nonmetals and Metalloids

Most of the elements in the periodic table are metals. Note the stair step line in the periodic table.  Elements to the left of the line are metals. Elements to the right of the line are nonmetals. In between metal and non-metals there are semi-metals or metalloids. Metals lose electrons and nonmetals gain electrons.

Ionic Compounds are formed when electrons are exchanged in this way between metals and nonmetals. 

Covalent or Molecular Compounds are formed between non metals and non metals react by sharing electrons.

Atomic Number and Atomic Mass

         The atomic number (Z) of an element represents the number of protons in the nucleus of atoms of that specific element. No two element has that same number of protons. Atomic number after it was discovered proved to be the best order without any discrepancies to arrange the elements in the periodic table and is shown on top of the space for each element. The atomic number will always be a whole number value without decimals.  At the bottom average atomic mass calculated based on isotopes of each elements is written.

 

 

Problem: Pick the a) representative elements, b) transition elements, c) inert gas elements, d) elements that from anions, e) semi- metals, and  f) elements that from cations from  the following list: Ca, Si, K, Ar, Cu, Fe Zn, Ge, Kr, Cl, O, F.

Answer:

a)      representative elements: Ca, Cl, O, F

b)      transition elements: Cu, Fe

c)      inert gas elements:  Ar, Kr

d)      elements that from anions: O, F

e)      semi- metals: Si, Ge

f)        elements that from cations: Ca, K, Cu, Fe Zn

 

 

 

Look on a periodic chart at the elements listed below. Do you know how to find an elements atomic number?

Problem: Use your periodic table to find the symbol, atomic number and atomic mass rounded to two decimal place of each of the following elements:

a)      Magnesium  b) Neon  c) Selenium  d) Gold

Answer

            Mg, atomic number = 12, mass = 24.31 amu

            Ne, atomic number = 10, mass = 20.18 amu

            Se, atomic number = 34, mass 78.96 amu

            Au, atomic number 79, mass 197.0 amu

 

Electron Arrangement and the Periodic Table

Bohr concluded that the energy levels of an atom can handle only a certain number of electrons at a time.

The Quantum Mechanical Atom

J. J. Thomson had demonstrated the particle properties of the electron earlier.

Because electrons can exhibit diffraction patterns, they have a dual nature of both wave and particle.

In 1924, Louis de Broglie suggested that the electron should have wave properties.

Light waves exhibit "diffraction."     

Erwin Schrodinger developed equations to describe the regions around the nucleus where electrons had the probability of being 95% of the time.

 

These regions of high probability for finding an electron around the nucleus were called orbitals.  Three dimensional models of the probability regions or orbitals can be constructed. Electron cloud representations are used to show the space that can be occupied by electrons in different energy levels.

 

Building Atoms by Orbital Filling

Schrodinger's work showed that each orbital could have a maximum of two electrons. Energy levels could contain different numbers of orbitals. Energy levels further from the nucleus can accommodate more orbitals than energy levels nearer the nucleus.

Energy levels can have sublevels when multiple orbitals are present.

Orbital Shapes 

Sub-level

Shape

# of orbitals/energy level

Picture

s

spherical

1

p

dumbbell 

3

d

complex          

5

f

very complex 

7

 

Energy Levels

n=1

n=2

n=3

n=4

number of sublevels

one

two

three

four

Sublevel Names

s

 s and p

s, p and d

s. p, d and f

Sublevels and orbitals

1s (1)

2s(1) 2p(3)

3s(1) 3p(3) 3d(5)

4s(1) 4p(3) 4d(5)  4f(7)

Number orbitals

1

 4

9

16

maximum number of electrons per sublevel

2(2n2)

2 + 6= 9(2n2)

2 + 6 +  10= 18 (2n2) 

2 +  6 +  10 + 14 = 32 (2n2)



 

The maximum number of electrons that can be in an energy level is 2n2, where n is equal to the energy level being considered.
 

Energy Level

maximum number of electrons
in an Energy Level

# of
Sublevels

sublevels names

maximum number of electrons per sublevel

n = 1

2

1

s

2

n = 2

8

2

s, p

2, 6 =8

n = 3

18

3

s, p, d

2, 6, 10 =18

n = 4

32

4

s, p, d, f

2, 6, 10, 14 =32

 

Problem:

How many electrons are found:

Within principle shells? a) n = 1 b) n = 2 c) n = 3 d) n =4  e) n = 5

Answer:

a.       n = 1; 2n2 = 2(1)2 = 2

b.      n = 2; 2n2 = 2(2)2 = 8

c.       n = 3; 2n2 = 2(3)2 = 18

d.      n = 4; 2n2 = 2(4)2 = 32

e.       n = 5; 2n2 = 2(5)2 = 50

 

Problem: With in a sub-shells: a) s,   b) p    c) d,    d) f

Answer:  a) s = 2, b) p=6 c) d=10, d) f= 14 

 

Problem: With in a Orbital?  

Answer:  Two electrons.

Energy Levels and Sublevels

A sublevel is a part of a principal energy level and is designated s, p, d, and f.  Each sublevel may contain one or more orbitals, regions of space containing a maximum of two electrons with their spins paired.

Schrodinger's work showed that

 

 

Building Atoms by Orbital Filling

Amazingly, the "electron configurations" of the elements are "embedded" in the Periodic Table.
Honk, if you can see this "embedded" information in the Periodic Table?
Analogy:  The periodic table is actually a packing slip that tells how the electrons are packed around the nucleus.

 

Electronic Configuration - the arrangement of electrons, in orbits or orbitals, around a nucleus of an atom.

 

Electron Configuration and the Aufbau ( Building Up)  Principle

A scheme used by chemist to obtain electronic configuration of a multi-electron atom in the ground state by filling atomic orbital starting with lowest energy.

 

1s 2s 2p3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d… (building up principle)

 

 If two or more orbitals exist at the same energy level, they are degenerate.  Do not pair the electrons until you have to.

 

 

 

 Problem:  What is the electron configuration of a) K and  b) P?

Answer:

         Using Aufbau principle or periodic table

 

            a.   Potassium: 1s2, 2s2, 2p6, 3s2, 3p6, 4s1

            b.   Phosphorus:  1s2, 2s2, 2p6, 3s2, 3p3

Problem:  Examine the electron configurations below, and name the element.
1s2 2s2           1s2 2s2 2p3            1s2 2s2 2p6 3s2 3p1

Answer:  Going through the periodic table.

1s2 2s2  (He)         1s2 2s2 2p3  (N)           1s2 2s2 2p6 3s2 3p1 (Al)

Problem: If a neutral atom in its ground state contains only 5 electrons in its outermost p sublevel, it is an atom in what "vertical group" of elements?

Answer:  group 17 or VIIB or Halogen family.

Problem: If a neutral atom in its ground state contains 2 electrons in its outermost s sublevel, it is an atom in what "vertical group" of elements?

Answer:  group 2 or IIA or Alkaline Earth family.

Problem: State what is similar and what is different about the electron configuration of fluorine and chlorine.

Answer:  F: 1s2 2s2 2p5   Cl: 1s2 2s2 2p6 3s2 3p5 valance shell electron configuration is similar but electron configurations are different.

Problem: Fluorine and chlorine have similar chemical properties. Oxygen and sulfur have similar chemical properties. However, oxygen and sulfur have chemical properties different from fluorine and chlorine. What does electron configuration have to do with this observation?

Answer:          F:         1s2 2s2 2p5       Cl:        1s2 2s2 2p6 3s2 3p5  both have same vala.nce electron configurations and similar chemical properties.

                        O:        1s2 2s2 2p4       S:         1s2 2s2 2p6 3s2 3p4 both have same valance electron configurations and similar chemical properties..

However, two groups F, Cl: ns2 np5  and O, S: ns2 np4   have different valance electron configurations  creating different chemical properties.

 

Valence Electrons

            The outermost electrons in an atom are valence electrons. For representative elements, the number of valence electrons in an atom corresponds to the group or family number. If atoms of different elements have the same electron arrangement in their valence shell electrons, then they can have similar chemical properties even if their atomic numbers or atomic masses are quite different.  Metals tend to have fewer valence electrons than nonmetals. Valence electrons are involved in chemical interactions and bonding (valence comes from the Latin valere, "to be strong"). Valence shell electrons are available to be lost, gained, or shared in chemical reactions.

Problem: How many total electrons and valance electrons are in the following atoms: 

a)  K, b) F,  c) P, d) O and e) Ca

Answer

For counting valance electrons go to the period the element is found and count ( excluding transition element blocks) from left to right until element is found.

            a.   Total electrons = 19 (same as atomic number), valence electrons = 1

            b.   Total electrons =   9 (same as atomic number),  valence electrons = 7

            c.   Total electrons = 15 (same as atomic number),  valence electrons = 5

            d.   Total electrons =  8 (same as atomic number),   valence electrons = 6

e.   Total electrons = 20 (same as atomic number),  valence electrons = 2

 

Abbreviated Electron Configurations

Abbreviated electronic configuration is separating valance electrons from core electrons and designating core electrons as a noble gas.

 

E.g.      What is the abbreviated electron configurations of a) K, b) P and Sn?

Answer

First, obtain the electron configuration then find the valence electrons. 

                  a.   Potassium (K):             1s2, 2s2, 2p6, 3s2, 3p6, 4s1

                  b.    Phosphorus (P):           1s2, 2s2, 2p6, 3s2, 3p3

                        c.    Tin (Sn):                      1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p2

Second, lump all non valance electrons as “core” abbreviated as a noble gas con figuration.

            a.   Ar:              1s2, 2s2, 2p6, 3s2, 3p6                           = [Ar]

            b.   Ne:             1s2, 2s2, 2p6, 3s2                                              =  [Ne]

c.    Kr:            1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6      =  [Kr]

Final answer

            a.   Potassium (K):        [Ar] 4s1

            b.   Phosphorous (P):    [Ne] 3s2, 3p3

c.    Tin (Sn):                [Kr] 5s2, 4d10, 5p2

 

      Electron configuration of the elements is predictable, using the Aufbau Principle. Knowing the electron configuration, we can identify valence electrons and begin to predict the kinds of reactions that the elements will undergo.

            Elements in the last family, the noble gases, have either two or eight valence electrons. Their most important properties are their extreme stability and lack of reactivity. A full energy level is responsible for this unique stability.

 

The Octet Rule

         Noble gases are non-reactive because they all have a complete outer shell. An atom chemically reacts to fill its valance shell.  A full valance shell contains eight electrons there fore the name octet. The octet rule tells us that in chemical reactions atoms of elements will gain, lose or share the minimum number of electrons necessary to achieve the electron configuration of the nearest noble gas. 

octet rule - the rule which predicts that atoms form the most stable molecules or ions when they are surrounded by eight electrons in their highest occupied energy (valance) level.

 

Electronic configuration of ions

Series of negative ions, noble gas atom, and positive ions with the same number electrons and electronic configuration. Electron configuration of ions is obtained by adding more electrons (anions) or removing electrons (cations) from a neutral atom. In the process atoms achieves a noble gas electron configuration.

 Group 1 (or IA), Alkali Metals  have one valence electron.

They all form +1 cations when the single valence electron is lost.

Metals lose electrons and achieve electron configuration of preceding noble gas.

E.g. Potassium (K):     

                  K ΰ  K+ (cation) +  e-

Oxygen (O):     

                  O + 2e- ΰ  O2- (anion) 

Metallic elements tend to form cations and nonmetals form anions that are isoelectronic with their nearest noble gas neighbor.

 

Isoelectronic electronic configurations

 If atom and a cation or anion have same number of electrons they are called isoelectronic.

E.g.  K+ and Ar          

         O2-  and Ne

Problem: Which of the following are isoelectronic: F–,Cl–, K+, Ar

Answer:

a.       F–, 10e–; Cl–, 18 e–; Not isoelectronic

b.      K+, 18 e–; Ar, 18e–; Isoelectronic         

 

Ion Formation and the Octet Rule

Metals lose electrons and achieve a an octet of valance electrons similar to electron configuration of preceding noble gas.

 

E.g. Potassium (K):      [Ar] 4s1

                  K ([Ar] 4s1) ΰ  K+  ([Ar])  +  e-

Oxygen (O):      [He] 2s2 2p4

                  O ([He] 2s2 2p4)  + 2e- ΰ  O2-  ([Ne]) 

 

a.      I–  (54 e–)   = Xe=  1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6

            b.   Ba2+ (54 e–)= Xe=  1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6

            c.   Se2– (36 e–) =Kr =  1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6

            d.   Al3+ (10 e–)=Ne =   1s2, 2s2, 2p6

 

 

Trends in the Periodic Table

Atomic Size

Atomic size increases from top to bottom but decreases from left to right in the periodic table. Cations are smaller than the parent atom. Anions are larger than the parent atom. Ions with multiple positive charge are even smaller than their corresponding monopositive ion; ions with multiple negative charge are larger than their corresponding less negative ion.

Problem: Arrange the following list of elements in order of increasing atomic size.

a)      Al, Si, P, Cl, S

b)      In, Ga, Al, B, Tl

c)      Sr, Ca, Ba, Mg, Be

d)      O,N, Sb, Bi, As

Answer:

a.       (Smallest) Cl, S, P, Si, Al (Largest)

b.      (Smallest) B, Al, Ga, In, Tl ( Largest)

c.       (Smallest) Be, Mg, Ca, Sr, Ba (Largest)

d.      (Smallest) N, P, As, Sb, Bi (Largest)

Ionization Energy

 The energy required to remove an electron from an atom in the gas phase.

      The energy required to remove an electron from the atom is the ionization energy. Descending a group, the ionization energy decreases. Proceeding across a period, the ionization energy increases.

Problem:  Arrange the following list of elements in order of increasing ionization energy.

a)      N,F, O

b)      Li, K, Cs

c)      Br, I, Cl

Answer:

a)      (Smallest) N, O, F(Largest)

b)      (Smallest) Cs, K, Li (largest)

d)      (Smallest) Cl, Br, I  (Largest)

Electron Affinity

      The energy released when a single electron is added to neutral atom in the gaseous state is known as the electron affinity. Electron affinities generally decrease proceeding down a group and increase proceeding across a period.

      Exceptions exist for periodic trends. They are generally small anomalies, and do not detract from the predictive power of the periodic table.

Problem:  Arrange the following list of elements in order of increasing ionization energy.

a.       Na, Li, K

b.      Br, F, Cl

c.       S, O, Se

Answer:

a.             (Smallest) Li, Na, K (Largest)

b.             (Smallest) F, Br, Cl (Largest)

c.             (Smallest) Se, S, O (Largest)

 

Ion Size

Ions follows same trends as for atomic radius in a group, fro example taking oxide and sulfide ion: radius of  O2- < S2-.

Cation or positive ions have fewer electrons than neutral atom and nuclear charge being same attract remaining electrons strongly making cation smaller than the neutral atom

Anions or negative ions larger than neutral atom.  Anions are larger than the atoms from which there are formed. Adding electrons  to an atom increases the repulsion between electrons.   Anion has a harder time holding on to the electrons.