Electronic Structure and Periodicity

Electronic Structure

The properties of atoms arise from the interactions between their nuclei and electrons..  Atoms are made up of a  positively charged nucleus. In the nucleus are positively charged protons and  neutral neutrons . Negatively charged electrons  orbit around the nucleus. Electrons can be easily added to or removed from most atoms.

According to Coulomb's Law, like charges repel each other and unlike charges attract each other. The higher the charge, the greater the attraction and repulsion, and the greater the distance between the charges, the less the  attraction and repulsion. The properties of atoms can be explained by opposite charges attracting each other, and like charges repelling each other.
In an atom, electrons arrange themselves into shells, subshells,  and  orbitals.. Each orbital can contain up to two electrons.  s subshells contain one orbital (up to 2 electrons), p subshells contain three orbitals (up to 6 electrons), d subshells contain five orbitals (up to 10 electrons). Larger subshells (f, g...) are rarely used in introductory chemistry.

Electron Configuration and the Periodic Table

Electron configuration is the distribution of electrons of an atom or molecule in atomic or molecular orbitals.  In the distribution of electrons,  the following should be remembered :

• The order of increasing energy in  given subshells is:
                                    1s < 2s < 2p < 3s < 3p < 4s < 4d < 4p < 5s
• Lower energy shells and subshells are filled first before  filling up the next higher energy level and subshells
• Core electrons are electrons in the filled shells and they are tightly bound to the atom
• Valence electrons are electrons at the outermost shell. These  electrons are more easily removed than the core electrons.

In the electron configuration of sodium  (  1s22s22p63s1 ) :
In 1s22s22p6 – there are 10 ( 2+,2 + 6 ) core electrons and 1s, 2s, and 2p are 
                        filled shells
in 3s1 – the lone electron in 3s is the valence electron and 3s is the valence shell

Electron Configurations of Ions

Ions are formed when atoms gain or lose electrons. A cation forms when one or more electrons are removed from a parent atom. For main group elements, the electrons that were added last are the first electrons removed. For transition metals and inner transition metals, however, electrons in the s  orbital are easier to remove than the d  or f  electrons. An anion forms when one or more electrons are added to a parent atom. The added electrons fill in the order predicted by the Aufbau principle.

Example :  Write the electron configuration and orbital diagram of Na  and  Na + .

Answer:
1.   Na:    1s22s22p63s1       2.  Na+ :    1s22s22p6  ( sodium cation loses 1 electron )

   1s     2s         2p          3s                                           1s        2s             2p

Periodic Table

Early attempts were made to classify elements ever since the discovery of metals. Some grouped metals separately from non-metals and some  applied formulation of arithmetic to the known atomic weights of certain  elements.
J.W. Dobereiner in 1817 discovered that when closely related elements are grouped in a set of three, the atomic weight of the middle element had an atomic weight almost the arithmetic mean of the other two elements in that group. Ex.

Element Lithium Sodium Potassium
Atomic weight 6.94 22.99 39.10
Mean atomic weight ----- 23.02 ------

He called such a group of three elements a Triad. He could group only a few elements due to lack of knowledge of correct atomic weights of the elements at that time.

In 1863, J.A.R. Newlands developed a system of classification of elements and entitled it Law of Octaves. He arranged the elements in such a way that every eight element had similar properties, like the notes of music. The law could not apply to a large number of known elements. However, the law indicated very clearly the recurrence of similar properties among the arranged elements. Thus the periodicity was visualized for the first time in a meaningful way.

More significant results were obtained when Lother Meyer’s work reflecting the periodicity was found to be based on physical properties of elements. He clearly showed that certain properties showed a periodic function.

Mosely, in 1913, showed that the properties of the elements are periodic function of their atomic numbers. He arranged all the elements according to the increasing atomic number.

Atomic number itself was discovered in 1913 by a team lead by Mosely. Table based on atomic number is term as Modern Periodic Table.

The arrangement of elements in the periodic table is a perfect matching of electron configuration of the elements and their physical and chemical properties. Some important considerations of the modern atomic structure applied to the classification are;
• an atom loses electrons from or gains electrons in the outermost orbit of an atom during a chemical reaction
• the sharing of an electron or electrons by an atom with other atom or atoms is largely through the outermost orbit. Thus the electrons in the outermost orbit of an atom largely determine the chemical properties of the elements

It can be concluded that the elements possessing identical outer electronic configuration should possess similar physical and chemical properties and therefore they
should be placed together for an easy and systematic study.

Structural Features of the Periodic Table

• There are 18 vertical columns called groups. They are numbered from 1 to 18. Each  group has a unique configuration.
• There are seven rows of boxes filled with elements. These rows are called periods.
• First period                        -  consists of only 2 elements.
Second and third periods -  consist of 8 elements each
Fourth and fifth period  -  consist of 18 elements each.
Sixth period and seventh period -  consist of 32 elements each

The International Union of Pure and Applied Chemistry (IUPAC) has announced that recently discovered elements 113, 115, 117 and 118 will now be known as nihonium (Nh: Z = 113 ),  moscovium ( Mc: Z = 115 ) , tennessine ( Ts : Z = 117 ),  and oganesson ( Og : Z = 118 ). The four elements, which complete the seventh row of the periodic table.

• There are six names given to the groups or a cluster of groups on the basis of the similarity of their properties (  e.g., transition metals ).
Group 1:    Alkali Metals ( except hydrogen )
Group 2:    Alkaline Earth Metals
Group 3 - 12:  Transition Metals
Group  16: Chalcogens
Group 17: Halogens
Group 18: Noble Gases ( Inert gases )
• Lanthanides ( Inner Transition Elements – First series ) – elements  with atomic numbers 58 to 71
• Actinides ( Inner Transition Elements - Second series ) -  elements from atomic numbers 90 to 103

All elements except transition or inner transition elements are also collectively called Main Group Elements.

Valence Electrons and the Periodic Table

Valence  electrons are the outermost electrons in the valence shell (highest value of n) of a ground-state atom. They determine how an element reacts.

The number of valence electrons in an atom is reflected by its position in the periodic table of the elements (see Figure below). Across each row, or period, of the periodic table, the number of valence electrons in groups 1–2 and 13–18 increases by one from one element to the next. Within each column, or group, of the table, all the elements have the same number of valence electrons. This explains why all the elements in the same group have very similar chemical properties.
The group number of the representative elements  is  the number of valence electrons.

The representative elements are in columns A1 - A7. In an 18 column table the map is as follows:
A1 -  column 1 ( hydrogen + alkali metals ) have 1 valence
A2 = column 2 (  alkaline earth metals  ) have 2 valence
A3 = column 13 (  Boron family ) has 3 valence
A4 = column 14 ( Carbon group ) has 4 valence
A5 = column 15 ( Nitrogen family ) has 5 valence
A6 = column 16 ( Oxygen column : chalcogens ) have 6 valence
A7 = column 17 ( Fluorine column:  halogens ) have 7 valence


Based on the figure above, it shows the number of valence electrons of each element in relation to its group number. Any element in group 1 has just one valence electron. Any element in group 8 ( column 18 ) has eight valence electrons with the exception of He which has a total of 2 electrons.

Properties of Families of Elements

 Metalloids
  These are elements with  a mix  properties of nonmetal and metal properties. The metalloids are hybrid elements. The most important characteristic of metalloids is that they partially conduct electricity for semi-conductors.

Nonmetals
Nonmetallic elements are brittle. They  gain electrons to become anions.

Metals
Metallic elements are good electrical and heat conductors. They are   malleable and ductile,  possess metallic luster , opaque as thin sheet, and solid at room temperature ( except Hg ).  They normally lose electrons to become cations. 

Halogens
Elements under this group have seven valence electrons, which make them highly reactive for covalent bonds. Halogens form ionic bonds with all kinds of elements. Example is NaCl ( salt ).

Noble Gases
These are un-reactive elements because their outer shells are fully filled with eight electrons (except helium which has two).

Chalcogens
All chalcogens are very reactive to alkali earth metals. Lighter chalcogens such as sulfur and oxygen are non-toxic, and essential to all types of life. Heavier chalcogens like tellurium, selenium, and polonium are toxic and quite harmful. Each  of these elements have 6 valence electrons.
Alkali Metals
 Elements in the first group which have one valence electron. Alkali metals are very reactive because they only need to lose one electron to have a full shell.

Alkaline Earth Metals
These are elements under Group 2. They are also highly reactive, but not as reactive as the alkali metals, because they have two valence electrons. They are called alkaline metals because when they are mixed in solutions they form basic  solutions.

Transition Metals
 This make up the largest family in the periodic table. They are located between and including the following elements horizontally: scandium through copper, yttrium through silver, lanthanum through gold, actinium through all higher atomic numbers in that period. They have a lot of electrons (normally) and distribute them in many advanced/ complicated ways.

Lanthanides
These are groups of metals located on the second row from the bottom of the periodic table. They are fairly rare, their atomic numbers range from 57 (lanthanum) to 71 (lutetium). Some of these elements can be found in superconductors, glass production, or lasers.
Actinides
These are groups of metals in the bottom row of the periodic table. The actinide family contains fifteen elements starting with actinium through the entire row to lawrencium. All actinides are radioactive and some are not found in nature.
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Periodicity

Periodicity refers to trends or recurring variations in element properties with increasing atomic number. Periodicity is caused by regular and predictable variations in element atomic structure.
Mendeleev organized elements according to recurring properties to make a periodic table of elements. Elements within a group (column) display similar characteristics.

The rows in the periodic table (the periods) reflect the filling of electrons shells around the nucleus, so when a new row begins, the elements stack on top of each other with similar properties. For example, helium and neon are both fairly unreactive gases that glow when an electric current is passed through them. Lithium and sodium both have a +1 oxidation state and are reactive, shiny metals.

Periodicity helped scientists find new elements because they could be expected to display certain characteristics based on the location they would take in the periodic table. Now that the elements have been discovered, scientists  used periodicity to make predictions about how elements will behave in chemical reactions and their physical properties. Periodicity helps chemists predict how the new, super heavy elements might look and behave.

Periodic Variation in Atomic Properties

The valence electrons of an element is determined by the element’s position in the periodic table.  Because the valence electrons determine the reactivity and chemical properties of an element, the periodic table is used as a map to determine the general properties of any element.
Elements are generally classified as either metals, nonmetals, or metalloids. Elements in each classification behave slightly in different ways. The number of protons in the nucleus, number of shells , and shielding effect  are the three factors that help in the prediction of the trends in the periodic table. Below are some atomic properties and their periodic trend.

Atomic Radius
Atomic size is the distance from the nucleus to the valence shell where the valence electrons are located. Atomic radius is a more definite and measurable way of defining atomic size. It is the distance from the center of one atom to the center of another atom.
From top to bottom of the periodic table, the atomic radii increase. This is because the valence electron shell is getting  larger and there is a larger principal quantum number, so the valence shell lies physically farther away from the nucleus.
Going across a row on the periodic table, left to right,  although the valence shell maintains the same principal quantum number, the number of protons( the nuclear charge ) is increasing .  The increasing positive charge casts a tighter grip on the valence electrons, thus, the atomic radii decrease

Ionization Energy
Ionization energy is the energy required to completely remove an electron from a gaseous atom or ion. The Ionization Energy is always positive.
From top to bottom of the periodic table, it becomes easier to remove an electron from an atom (i.e., IE decreases) because the valence electron is farther away from the nucleus, so ionization energy decreases.

From left to right,  the electrons get drawn closer in. It takes more energy to remove an electron so ionization energy increases.

Electron Affinity
Electron affinity is the energy change that occurs when an electron is added to a gaseous atom. It is the opposite of ionization energy. Electron affinity can further be defined as the enthalpy change that results from the addition of an electron to a gaseous atom. It can be either positive or negative value. The greater the negative value, the more stable the anion is.

Electronegativity
Electronegativity is the measurement of an atom to compete for electrons in a bond. The higher the electronegativity, the greater its ability to gain electrons in a bond. Electronegativity determines  polar and nonpolar molecules. Electronegativity is related with ionization energy and electron affinity. Electrons with low ionization energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted by the positive nucleus on the negative electrons.

Metallic Character
The metallic character is used to define the chemical properties that metallic elements present. Generally, metals tend to lose electrons to form cations. Nonmetals tend to gain electrons to form anions. They also have a high oxidation potential therefore they are easily oxidized and are strong reducing agents. Metals also form basic oxides; the more basic the oxide, the higher the metallic character.
From left to right of the periodic table,  the metallic character decreases, because the elements easily accept electrons to fill their valance shells. These elements take on the nonmetallic character of forming anions. From bottom to top, the metallic character decreases, due to the greater pull that the nucleus has on the outer electrons. This greater pull makes it harder for the atoms to lose electrons and form cations.

Summary of periodic trend of atomic properties;

Atomic Property         From Left to Right   From Bottom to Top
1.  Atomic Radius         Decreases           Decreases
2.  Ionization Energy Increases                   Increases
3.  Electron Affinity Increases                   Increases
4.  Electronegativity Increases                   Increases
5.  Metallic Character Decreases           Decreases


Predicting the Properties of Elements Based on Its Position in the Periodic Table

Elements are organized by period and group, with the period corresponding to the principal energy level, and the group relating to the extent the subshells are filled.
The properties of an atom relate directly to the number of electrons in various orbitals, and the periodic table is much like a road map among those orbitals such that chemical properties can be deduced by the position of an element on the table. The electrons in the outermost or valence shell are especially important because they can engage in the sharing and exchange that is responsible for chemical reactions.
Using periodic trends, the periodic table can help predict the properties of various elements and the relations between properties. It therefore provides a useful framework for analyzing chemical behavior and is widely used in chemistry and other sciences.
Atomic Orbitals
The electrons in the partially filled outermost shell (or shells) determine the chemical properties of the atom. It is called the valence shell. Each shell consists of one or more subshells, and each subshell consists of one or more atomic orbitals. The properties of an atom depend ultimately on the number of electrons in the various orbitals, and on the nuclear charge which determines the compactness of the orbitals.
Periods 1, 2, & 3
Shown below are the first three rows of the periodic table ( Period 1 , 2 , and  3 ) which  corresponds to the principal energy levels ( n =1, n=2, and n=3 ).


Shown above are the electron configuration of the 18 elements. In each electron configuration , the principal energy level ( n )  and the valence electrons are shown. As mentioned, the properties of an element are dependent on the valence electrons.
The chemical properties are influenced by the valence (outer shell) electrons. Elements with the same number of valence electrons fall in the same column and exhibit similar chemical properties up to a point. For example, lithium, sodium, potassium, rubidium and cesium in column 1 are valence +1 and have similar alkaline chemistry tending to form ionic compounds that ionize and disperse in water solution.