CERAMIC SCIENCE - 1

8/7/2019 CERAMIC SCIENCE - 1

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CRYSTAL CHEMISTRY

Electronic Configuration of the Elements

Hydrogen through Krypton

Here's a useful table for your chemistry homework or

general use! This is a compilation of the electron configurations of the

elements up through number 104, broken into three pages (the table was too

large for anything less). To arrive at the electron configurations of atoms, you

must know the order in which the different sublevels are filled. Electrons enter

available sublevels in order of their increasing energy. A sublevel is filled or

half-filled before the next sublevel is entered. For example, the s sublevel can

only hold two electrons, so the 1s is filled at helium (1s2). Thep sublevel can

hold six electrons, the dsublevel can hold 10 electrons, and the fsublevel can

hold 14 electrons. Common shorthand notation is to refer to the noble gas

core, rather than write out the entire configuration. For example, the

configuration of magnesium could be written [Ne] 3s2, rather than writing out

1s22s

22p

63s

2.

No. Element K L M N O P Q

1 2 3 4 5 6 7

s s p s p d s p d f s p d f s p d f s

1 H 1

2 He 2

3 Li 2 1

4 Be 2 2

5 B 2 2 1

6 C 2 2 2

7 N 2 2 3

8 O 2 2 4

9 F 2 2 5

10 Ne 2 2 6

11 Na 2 2 6 1

12 Mg 2 2 6 2

13 Al 2 2 6 2 1

14 Si 2 2 6 2 2

15 P 2 2 6 2 3


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16 S 2 2 6 2 4

17 Cl 2 2 6 2 5

18 Ar 2 2 6 2 6

19 K 2 2 6 2 6 - 1

20 Ca 2 2 6 2 6 - 2

21 Sc 2 2 6 2 6 1 2

22 Ti 2 2 6 2 6 2 2

23 V 2 2 6 2 6 3 2

24 Cr 2 2 6 2 6 5* 1

25 Mn 2 2 6 2 6 5 2

26 Fe 2 2 6 2 6 6 2

27 Co 2 2 6 2 6 7 2

28 Ni 2 2 6 2 6 8 229 Cu 2 2 6 2 6 10 1*

30 Zn 2 2 6 2 6 10 2

31 Ga 2 2 6 2 6 10 2 1

32 Ge 2 2 6 2 6 10 2 2

33 As 2 2 6 2 6 10 2 3

34 Se 2 2 6 2 6 10 2 4

35 Br 2 2 6 2 6 10 2 5

36 Kr 2 2 6 2 6 10 2 6

*Note Irregularity

The electron configuration of an atom is the particular distribution of electrons

among available shells. It is described by a notation that lists the subshell

symbols, one after another. Each symbol has a subscript on the right giving the

number of electrons in that subshell. For example, a configuration of the

lithium atom (atomic number 3) with two electrons in the 1s subshell and one

electron in the 2s subshell is written 1s22s

1.


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sublevel orbital maximum # of electrons

s 1 2

p 3 6

d 5 10

f 7 14

The notation for electron configuration gives the number of electrons in each

subshell. The number of electrons in an atom of an element is given by the

atomic number of that element.

On the left we have a diagram to show how the orbitals of a subshell are

occupied by electrons. On the right there is a diagram for the filling order of

electrons in a subshell.


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Here are some examples that show how to use the filling order diagram to

complete the electron configuration for a certain substance.

Element # of Electrons in Element Electron Configuration

He 2 1s2

Li 3 1s22s

1

Be 4 1s22s

2

O 8 1s22s

22p

4

Cl 17 1s22s22p63s23p5

K 19 1s22s

22p

63s

23p

64s

1


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Often times you will be asked to find the electron configuration for something

that looks like this:

53I

The 53 denotes the number of electrons in an atom of iodine. You would now

proceed to do the electron configuration by looking at the filling order chart.

1s22s

22p

63s

23p

64s

23d

104p

65s

24d

105p

5

Periodicity

With increasing atomic number, the electron configuration of the atoms

display a periodic variation. Because of this the elements show periodic

variations of both physical and chemical behavior. The periodic law is a lawstating that when the elements are arranged by atomic number, their physical

and chemical properties vary periodically. We are going to be looking at three

physical properties of an atom: atomic radius, ionization energy, and electron

affinity.

Atomic Radius

The size of the electron cloud increases as the principal quantum numberincreases. Therefore, as you look down the periodic table, the size of atoms in

each group is going to increase. When you look across the periodic table, you

see that all the atoms in each group have the same principal quantum number.

However, for each element, the positive charge on the nucleus increases by

one proton. This means that the outer electron cloud is pulled in a little tighter.

One periodic property of atoms is that they tend to decrease in size from left


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to right across a period of the table. So finally we have a good definition for

how the atomic radii increases: the atomic radii increases top to bottom and

right to left in the periodic table.

Ionization Energy

The energy needed to remove the most loosely held electron from an atom is

known as ionization energy. Ionization energies are periodic. The ionization

energy tends to increase as atomic number increases in any horizontal row or

period. In any column or group, there is a gradual decrease in ionization energy

as the atomic number increases. Metals typically have a low ionization energy.

Nonmetals typically have a high ionization energy.

Electron Affinity

The attraction of an atom for an electron is called electron affinity. Metals

have low electron affinities while nonmetals have high electron affinities. The

general trend as you go down a column is a decreasing tendancy to gain

electrons. As you go across a row there is also a trend for a greater attraction

for electrons.

IONIC BOND

The Ionic Bond: Ionic bonds are formed when there is a complete transfer of

electrons from one atom to another, resulting in two ions, one positively

charged and the other negatively charged. For example, when a sodium atom

(Na) donates the one electron in its outer valence shell to a chlorine (Cl) atom,

which needs one electron to fill its outer valence shell, NaCl (table salt)

results. Ionic bonds are often 4-7 kcal/mol in strength.


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Ionic Bonding

Key Concepts

y An ionic solid is made up of positive ions (cations) andnegative ions (anions) held together by electrostatic forces in

a rigid array or lattice.y Ionic bonding refers to the electrostatic attraction between

cations and anions.y The physical properties of ionic compounds are:

o High melting and boiling pointso Ionic solids do not conduct electricity (they are insulators).o When molten (liquid) ionic compounds conduct electricity.o When dissolved in water to form an aqueous solution ionic compounds

conduct electricity.o Hardo Brittle

Physical Properties of Ionic Compounds

Melting Point

Ionic compounds have high melting points.

The electrostatic attraction (ionic bond) between cations and anions is strong. It takes alot of energy to overcome this attraction in order to allow the ions to move more freely

and form a liquid.The factors which affect the melting point of an ionic compound are:

y The charge on the ions.In general, the greater the charge, the greater the electrostatic attraction, the

stronger the ionic bond, the higher the melting point.

The table below compares the melting point and ion charges for sodium chlorideand magnesium oxide.

Ionic Compound Melting Point (oC) Cation Charge Anion Charge

NaCl 801 +1 -1

MgO 2800 +2 -2

yMgO has a higher melting point than NaCl because 2 electrons are transferredfrom magnesium to oxygen to form MgO while only 1 electron is transferred fromsodium to chlorine to form NaCl.

y The size of the ions.Smaller ions can pack closer together than larger ions so the electrostatic

attraction is greater, the ionic bond is stronger, the melting point is higher.The melting point of Group IA (alkali) metal fluorides is compared to the ionicradius of the cation in the table below.


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When a stress is applied to the ionic lattice, the layers shift slightly.The layers are arranged so that each cation is surrounded by anions in the lattice. If the

layers shift then ions of the same charge will be brought closer together.

Ions of the same charge will repel each other, so the lattice structure breaks down intosmaller pieces.

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