Chemistry Form 6 Sem 2 03

PRE-U CHEMISTRY SEMESTER 2 CHAPTER 3 : PERIODIC TABLE : PERIODICITY

3.0





Introduction to Inorganic Chemistry Inorganic chemistry deals with the properties of all of the elements in the periodic table. These elements range from highly reactive metals, such as sodium, to noble metals, such as gold. The nonmetals include solids, liquids, and gases, and range from the aggressive oxidizing agent fluorine to unreactive gases such as helium. Although this variety and diversity are features of any study of inorganic chemistry, there are underlying patterns and trends which enrich and enhance our understanding of the discipline. These trends in reactivity, structure, and properties of the elements and their compounds provide an insight into the landscape of the periodic table and provide a foundation on which to build understanding. The periodic table provides an organizing principle that coordinates and rationalizes the diverse physical and chemical properties of the elements. Periodicity is the regular manner in which the physical and chemical properties of the elements vary with atomic number

PERIOD 2 Element

Li

Be

B

C

N

O

F

Ne

Proton number

3

4

5

6

7

8

9

10

Atomic radius (nm)

0.152

0.111

0.086 0.077 0.073 0.062 0.032 0.029

Melting point (oC)

180

1287

2076

3500

-220

-218

-210

-249

1st ionisation energy (kJ/mol)

519

900

799

1090

1400

1310

1680

2080

Electronegativ ity

0.98

1.57

2.04

2.55

3.04

3.44

3.90

--

Classification

Metal

Metal -loid

Non metal

PERIOD 3 Element

Na

Mg

Al

Si

P

S

Cl

Ar

Proton number

11

12

13

14

15

16

17

18

Atomic radius (nm)

0.186

0.160

0.143

0.118

0.108

0.106

0.099

0.088

Melting point (oC)

98

650

660

1423

44

120

-101

-189

1st ionisation energy (kJ/mol)

494

736

577

786

1060

1000

1260

1520

Electronegativ ity

0.9

1.2

1.5

1.8

2.1

2.5

3.0

-

Classification

Metal

Metal -loid

Non metal

3.1 Variation of physical properties of group and period
1. Atomic radius – half of the distance between the nuclei of the two closest@ identical atom (or half of the closest internuclear distance) Type Diagram Explanation • Covalent radius (for metalloid and non-metal)

•

Atomic nucleus • Metallic Radius (for metal)

•

In the case of covalent molecule, atomic radius is also known as covalent radius. Covalent radius is half the distance between the nuclei of 2 identical atoms covalently bonded. Or simply, covalent radius is half of the bond length between 2 covalently bonded identical atoms. Metallic radius is define as half the distance between the nuclei of neighbouring metal ion in a crystal lattice of a metal. Usually, the metallic radius is greater than covalent radius.







The atomic size of an element is determined by 2 factors.
The screening effect of the inner shell electrons which makes the atomic size larger. The screening effect is the result of the mutual repulsion between the electrons in the inner shell with those in the outer shell. Filled inner shells “shield” the outer electrons more effectively than electrons in the same sub-shell shield each other.
The nuclear charge which pulls all the electrons closer to nucleus. As a result of the increasing nuclear charge, atomic size becomes smaller. Hence when 2 factors combine  effective nuclear charge, Zeff
Zeff = Z (nuclear charge) – σ (screening constant)













The trend of atomic radius when gases down to group  atomic increase radius …………………… Explanation : When going down to group, nuclear charge increase as number of proton increase together with number of electrons. However, as more electrons filling the shells, the screening effect also increase. Consequently caused the effect nuclear charge decrease and outer most shell electrons are not hold tightly by the nucleus. For these reason, atomic radius increase The trend of atomic radius across the period. When across the decrease period 3, atomic radius ………………. Explanation : When going across period, nuclear charge increases as number of protons increase together with the number of electrons. However, the screening effect remain almost constant because electrons are filling in the same shell. This will caused the effective nuclear charge increases gradually resulting the atomic radius to decrease.

Atomic radius increase

Atomic radius decrease

2.





Ionic radius Ionic radius measures from the ion’s nucleus to the outermost shell. Diagram below shows the ionic radius for 2 cations from Period 4 and 3 anions from Period 3.

Ion

Anion

Cation

Diagram

Going across Period

Going down Group

When going across P3- , S2- , Cl- , K+ , Ca2+ the ionic radius decrease When going across these ions, the nuclear charge increase, since the number of protons increase. However, all these ions are isoelectronic (have the same number of electrons), hence the screening effect of these ions remain constant. This will caused the effective nuclear charge to increase, which result the ionic radius decrease when going across these ions. When going down to any group, ionic radius decrease (e.g. :Group 1 Li+, Na+, K+, Rb+, Cs+ ) This is due to, as nuclear charge increase, more electrons filling in the shell, which caused the screening effect to increase gradually. As a result, the effective nuclear charge decrease, hence caused the ionic radius to decrease




When the atomic and ionic radius of an element were to compare, student must know why does the atomic radius of an element is greater/smaller than its ionic radius, by using the screening effect and effective nuclear charge Cation

Trend

Atomic radius is larger than cation radius

Using Mg and Mg2+ as example ; Electronic configuration of Mg is 1s22s22p63s2. When 2 electrons were Explanation donated and form Mg2+ (1s22s22p63s2), the effective nuclear charge increase as the number of shell decrease, which will decrease the screening effect.

Anion

Atomic radius is smaller than anionic radius Using P and P3- as example ; Electronic configuration of P is 1s22s22p63s23p3. As P accept 3 electrons and form P3- (1s22s22p63s23p6), the effective nuclear charge decrease as the number of shell increase, which will decrease the effective nuclear charge.

3.


Melting point Table below shows the melting point of elements across Period 3

Bonding and Forces

Period

Explanation Elements : Lithium (Li) and Beryllium (Be) Valence electrons of Li and Be are 2s1 and 2s2 respectively. Since Be delocalise more electrons than Li, so melting point of Be is higher than Li

Metallic bonding - formed 2 when electrostatic forces is formed between the delocalised electrons and the positive ion. When Elements : Sodium (Na) , Magnesium (Mg) and electrons were delocalised Aluminum (Al) from a metal, it formed an Valence electrons of Na, Mg and Al are 3s1 , 3s2 electron sea thus interacting 3 and 3s23p1 respectively. Since Na, Mg and Al with the positive ion formed delocalised 1, 2 and 3 electrons respectively, so as a result of donating melting point increase Na < Mg < Al electrons. Thus, the more the electrons delocalised Example : Between Be and Mg by the metal, stronger the Valence electrons of Be and Mg are 2s2 and 3s2 electrostatic forces, Between respectively, indicate they are from the same stronger the metallic bond Period Group. Since Be has smaller metallic radius than Mg, hence greater electrostatic forces, so higher the melting point.

Bonding and Forces

Gigantic structure - each atom are strongly held by using covalent bond (depending on the number of valence electrons that are able to form covalent bond) hence forming a gigantic network which are very stable and required high temperature to break the covalent bond within the gigantic network.

Period

Explanation

2

Elements : Boron and Carbon Valence electrons of B and C are 2s22p1 & 2s22p2 respectively, hence B form strong covalent with 3 other boron atoms(via sp2 hybridisation), while C form strong covalent bond with 4 other carbon atoms (via sp3 hybridisation). More energies required to break more covalent bonds form between C, hence C has a higher melting point than B

3

Element : Silicon Valence electron of Si is 3s23p2. So, each Si can form strong covalent bond with 4 other Si atom (via sp3) to form a gigantic covalent network, hence required high temperature to break it.

Example : C and Si Valence electrons of C and Si are 2s22p2 & 3s23p2 Betwee respectively. Both of them form sp3 hybridisation between each atom. Since bond length between C-C is n Period shorter than Si-Si, hence stronger covalent bond is form between C-C. That is why carbon has a higher melting point than silicon

Bonding and Forces

Simple molecules - Nonmetal (except for C) tend to form simple molecule between them by using covalent bond. These molecules are hold weakly by using weak Van Der Waals forces between them, hence it required relatively low temperature to break the weak intermolecular forces between them

Period

Explanation

2

Elements : Nitrogen, Oxygen, Fluorine, Neon. Nitrogen, Oxygen, Fluorine exist as diatomic molecule, as N2 O2 and F2, while Neon exist as monoatomic Ne. Boiling point increase from Ne