Stpm Chemistry Form 6 Notes w1l47

STPM Chemistry Form 6 Notes – Terminology and Concepts: The Periodic Table (Part 1) The History of Periodic Table Döbereiner: Döbereiner’s Triads In 1817, Johann Wolfgang Döbereiner a German chemist (1780 – 1849) discovered that trends in certain properties of selected groups of elements and densities for some of these triadsfollowed a similar pattern. Newlands: Octave Law In 1865, John Alexander Reina Newlands an English chemist (1837 – 1898) devised a Periodic Table of elements arranged in order of their relative atomic masses and name it law of octaves. Law of octaves states that “any given element will exhibit analogues behaviour to the eighth element following it in the table“. Lothar Meyer: Atomic Volume of the Elements In 1870, Julius Lothar Meyer a German chemist (1836-1907) described 28 elements and arranged the atomic volume (relative mass/density) of the elements against their relative atomic masses where similar chemical and physical properties are repeated at periodic intervals. Meyer presented a series of maxima and minima curve. At the peaks of the curve, Meyer discovered that the most electropositive elements appeared at the peak (Li, Na, K, Rb and Cs). Mendeleev’s Periodic Table In 1869, Dimitri Mendeleev a Russian chemist (1836 – 1907) published the periodic table of all known elements and predicted several new elements to complete the table that formed the basis of the modern Periodic Table. The elements were arranged into Periods (horizontal rows)and Groups (vertical columns). The arrangement of the elements in groups of elements is in the order of their atomic weights corresponds to their valences. Mendeleev predicted new elements, namely eka-silicon (germanium), eka-aluminium (gallium), and eka-boron (scandium). Moseley: Proton Number In 1914, Henry Gwyn Jeffrey Moseley an English physicist (1887 – 1915) discovered the relationship between an element‟s X-ray wavelength and its atomic number (Z), Moseley demonstrated the arrangement by nuclear charge rather than related atomic mass. In Moseley‟s experiment the fast moving electrons strike a solid anode. From it, an X-ray spectrum is produced. Through this, Moseley‟s measurement of atomic numbers had an experimentally measurable basis and enable scientists to arrange the elements in the modern Periodic Table in order of increasing proton (atomic number). That‟s it for now on the history. Next part, you‟ll learn about the long and short periods, the „spdf‟ blocks, groups in the periodic table, atomic radius, ionic radius, electronegativity and ionisation energy of an atom. (Rather challenging, so make sure not to miss the next part only on Berry Berry Easy)

STPM Chemistry Form 6 Notes – Terminology and Concepts: The Periodic Table (Part 2) The Modern Periodic Table The modern Periodic Table is constructed on the basis of the proton (atomic) numbers of the elements and their electronic configuration

A) The short periods

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Period 1 – the 1s orbital is being filled. Period 2 – the 2s orbital is filled first, followed by the 2p orbital in the outermost shell (8 elements). Period 3 – the 3s orbital is filled first, followed by the 3p orbital in the outermost shell (8 elements).

B) The long periods

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Period 4 – the 4s, 3d and 4p orbitals are involved in the outermost shell (18 elements). Period 5 – the 5s, 4d and 5p orbitals are involved in the outermost shell (18 elements). Period 6 – the 6s, 4f, 5d and 6p orbitals are involved in the outermost shell (32 elements) and one series known as the lanthanides. Period 7 – the 7s, 5f, 6d and 7p orbitals are involved in the outermost shell (increasing due to the discovery of new elements) and one series known as the actinides.

C) The groups in the periodic table The s-block elements:

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Group 1 – general electronic configuration ns1. Group 2 – ns2.

The p-block elements

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Group 13 – ns2np1. Group 14 – ns2 np2. Group 15 – ns2 np3. Group 16 – ns2 np4. Group 17 – ns2 np5. Group 18 – ns2 np6.

The d-block elements

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Between Group 2 and Group 13 that the d orbitals are partially occupied.

The f-block elements

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Lanthanides (15 elements) – 4f orbitals are partially filled and must have a 6s2. Actinides (15 elements) – 5f orbitals are partially filled and must have a 7s2. All actinide elements are radioactive.

Some definitions 1.

Atomic Radius – half the distance between the nuclei of the two closest atoms in an element.

2.

Atomic Radii – decrease across a period from left to right in the periodic table & increase down a group in the periodic table.

3.

Ionic Radius – measure of the size of an atoms ion in a crystal lattice. Cation is smaller than the corresponding metal atom and anion is larger than the corresponding nonmetal atom.

4.

Electronegativity – measures the ability of an atom to attract to itself the electron pair forming a covalent bond. The greater the electronegativity of an atom, the greater the atom attraction for electrons.

5.

Ionisation Energy of an Atom – measures of its tendency to lose electrons. The larger the ionisation energy, the more difficult it is to remove an electron.

Before we end, ask yourself too, “Why the the atomic radii decrease across a period and increase down a group in the periodic table?”. If you cannot answer it, the answers will be revealed in the Part 3.

STPM Chemistry Form 6 Notes – Terminology and Concepts: The Periodic Table (Part 3) Periodicity of Atomic Radius Atomic radii for elements in Periods 2 and 3

Elements Li Be B C N O F Na Mg Al P S Cl

Atomic radius (pm) 152 112 80 77 74 74 72 156 136 125 110 104 99

Atomic radii can be classified into three categories:

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Covalent radius: Metallic radius Van der Waals radius

Effecting factors of the atomic radius:



Screening effect of the inner shell electrons: negatively-charged shells repel one another and are being pushed further away from the nucleus; screening effect increase; and size of the atoms increase.



Nuclear charge (number of protons in the nucleus) that pulls all the electrons closer to the nucleus: The higher the nuclear charge; the stronger the attraction between nucleus and the electron cloud; and the size of the atom decrease.

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Effective nuclear charge = No. of protons – No. of inner electrons

A) Atomic radius across a period Example: Period 2 (Li, Be, B, C, N, O, F, Ne) and Period 3 (Na, Mg, Al, Si, P, S, Cl, Ar)

Across the period:

    

Number of protons increase by one. Number of electrons increase by one. Screening effect does not affect much (same quantum shell). Nuclear charge increase (stronger attraction between nucleus and electron cloud). Size of the atoms decrease.

B) Atomic radius down a group Example: Group 2 (Be, Mg, Ca, Sr, Ba)

Down the group:

   

Screening effect increase. Nuclear charge increase. Effective nuclear charge decrease. Size of the atoms increase (the increase in the screening effect is larger than the increase in the nuclear charge).

C) Ionic radius (radius of a cation or or an anion) across Period 3

Ion Na+ Mg2+ Al3+ P3S2Cl-

Ionic radius 0.095 0.065 0.050 0.212 0.184 0.181

No. of electrons 10 10 10 18 18 18

No. of protons 11 12 13 15 16 17

Isoelectronic – species have the same number of electrons and the same electronic configuration. When given number of electrons (Na+, Mg2+, Al3+) or (P3-, S2-, Cl-)

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higher the nuclear charge, higher the force of attraction smaller the atomic size or ionic size.

When given nuclear charge,

  

larger the number of electrons in an atom or an ion, greater the repulsion between electrons larger the atomic or ionic size.

Conclusion:

 

Cationic size decreases (increasing proton number). Anionic size decreases (increasing proton number).

D) Ionic radius down a group Example: Group 2 (Be2+, Mg2+, Ca2+, Sr2+, Ba2+) & Group 17 (F-, Cl-, Br-, I-) Going down the Group 2 and Group 17:

  

Each successive ion has one additional shell filled with electrons. Screening effect increase Ionic size increase.

STPM Chemistry Form 6 Notes – Terminology and Concepts: The Periodic Table (Part 4) A) Boiling Point, Melting Point and Enthalpy of Vaporisation Enthalpy of vaporisation – the heat energy required to convert 1 mol of a liquid to its vapour at the boiling point of the liquid. Example: Period 2 (Li, Be, B, C, N, O, F, Ne) & Period 3 (Na, Mg, Al, Si, P, S, Cl, Ar) Across the period: the element become less metallic.

Li, Be, Na, Mg and Al are metals (metal lattice):

 

B.P, M.P and enthalpy of vaporisation increase and the atoms are held together by strong metallic bond. Increasing the number of valence electrons cause the strength of the metallic bond increase.

B, C (graphite) and Si are metalloids & C (diamond) is non-metal (giant covalent molecule):

  

B.P, M.P and enthalpy of vaporisation are very high. The atoms are held together by strong covalent bonds which form giant covalent structure(crystal lattice structure) in a 3-D structure. All the covalent bonds are needed to be broken before the solid melts.

N, O, F, P, S, Cl are non-metallic elements (simple molecular structure):

  

B.P, M.P and enthalpy of vaporisation are relatively low that involves only the breaking of weak Van der Waals forces. N2, O2, F2, P4, S8, Cl2 consist of small and discrete molecules. The covalent bonds within the molecules are very strong, but the Van der Waals forces of attraction between the molecules are very weak.

Ne and Ar are non-metallic (monoatomic structure):

 

B.P, M.P and enthalpy of vaporisation are very low. Noble gases are uncombined atoms and have very weak Van der Waals forces of attraction between the atoms.

B) Electrical Conductivity

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All metals (Li, Be, Na, Mg and Al) are good conductors either in the solid or molten state. Metals have delocalised electrons which will move freely across the metal in the solid lattice structure when an electrical potential or voltage is applied.

 

Non-metals (C diamond, N, O, F, Ne, P, S, Cl, Ar) are non-conductors. All the valence electrons in non-metals are used to form covalent bonds between atoms and there are no mobile electrons in the structure. Ne and Ar are noble gases and have the stable octet electronic configuration and do not have any mobile electrons.

 

Metalloids (C graphite, Si) are semi-conductor. Conductivity of metalloid increases with the increasing of temperature.

C) Ionisation Energy Ionisation energy of an element – the amount of energy required to pull one electron off an atom. Down the group in the periodic table, ionisation energy decreases because of the screening effect / shielding effect (electrons in low-energy levels repel electrons in higher-energy levels away from the nucleus) First ionisation energy of an element – the minimum energy required to remove 1 mol ofelectrons from 1 mol of atoms in the gaseous state. M(g) –> M+(g) + e, ΔH = first ionisation energy Second ionisation energy of an element – the minimum energy required to remove 1 mol of electrons from 1 mol of unipositive ion in the gaseous state. M+(g) –> M2+(g) + e, ΔH = second ionisation energy Third ionisation energy of an element M2+(g) –> M3+(g) + e, ΔH = third ionisation energy Fourth ionisation energy of an element M3+(g) –> M4+(g) + e, ΔH = fourth ionisation energy i) Factors affecting ionisation energy

  

Distance of the outer electrons from the nucleus (atomic size) Size of the nuclear charge (nuclear charge) Screening effect of the electrons in the inner shells (screening effect)

ii) Ionisation energy across a period 2 and period 3

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The first ionisation energy increase with increasing proton numbers for the elements (atomic size decreases, nuclear charge increases and the screening effect remains the same).



First ionisation of Be (period 2) and Mg (period 3) is higher than expected because the first electron to be removed is from a fully filled s orbital.

  

Be: 1s2 2s2 Mg: 1s2 2s2 2p6 3s2 First ionisation of N (period 2) and P (period 3) is is higher than expected because the first electron to be removed is from a half filled p orbital.

 

N: 1s2 2s2 2p3 P: 1s2 2s2 2p6 3s2 3p3

So there you go. Not so difficult as you though, isn‟t it? Try and revise all the way from the first post on this chapter about the periodic table for STPM Chemistry Form 6.

STPM Chemistry Form 6 Notes – Terminology and Concepts: The Periodic Table (Part 5 – Final) A) Electronegativity Electronegativity – measure how easy it is for an atom to gain electrons and how much an atom will pull electrons away from other atoms it has bonded to / covalent bond (similar to electron affinity but the difference is electron affinity deals with isolated atoms in the gas phase). Across the periodic table (left to right) 1.

Electronegativity increases.

2.

Left side: prefer to lose electrons.

3.

Right side: prefer to gain electrons.

4.

Noble gases: no electron affinity.

Down the periodic table 1.

Electronegativity decreases.

2.

It is because the shielding effect (nuclear charge increases but screening effect increase and the atomic size increases and as a result, the effective charge decrease).

B) Electron Affinity Electron affinity – the energy change that occurs when a gaseous atom picks up an extra electron. First electron affinity is exothermic: Example: O (g) + e –> O- (g) First electron is pulled/attracted by the positively charged oxygen atom nucleus. Second electron affinity is endothermic: Example: O- (g) + e –> O2- (g) Second electron is repelled by the existing negative charge on the oxygen ion. Across the periodic table (left to right) 1.

Left: elements want to lose electrons to be the nearest noble gas. As result, not much energy is released when these elements gain an extra electron. Electron affinity to be slightly negative.

2.

Right: elements want to gain electrons to be the nearest noble gas. As result, a very high energy to be released. Electron affinity to be more negative.

Down the periodic table 1.

Elements want to gain electron less (shielding effect)

2.

Bottom: elements have less negative electron affinities.

C) Variation of the Period of d-block Element Across the periodic table – First series (left to right) 1.

Atomic size is approximately the same (except Sc and Ti).

2.

Effective nuclear charge remains almost.

3.

High melting points and boiling points (except Zn).

4.

Density increases (but decreases for zinc).

5.

1st and 2nd ionisation energies of the elements increase slightly (as the proton numbers increase)

6.

3rd and 4th ionisation energies of the elements increase drastically.

Please revise previous parts in this series of notes if you want to understand the complete idea (for STPM level) on the periodic table.