Upcoming SlideShare
×

Like this document? Why not share!

# Periodic table - periodic properties

## on Jan 07, 2014

• 470 views

NCERT unit 3 periodic classification of elements.

NCERT unit 3 periodic classification of elements.

### Views

Total Views
470
Views on SlideShare
459
Embed Views
11

Likes
1
18
0

### 2 Embeds11

 http://chemncert.blogspot.in 8 http://www.weebly.com 3

### Report content

• Comment goes here.
Are you sure you want to

## Periodic table - periodic propertiesDocument Transcript

• Periodic table and periodic properties Unit III - Classification of Elements and Periodicity in Properties          Johann Dobereiner classified elements in group of three elements called triads. In Dobereiner’s triad the atomic weight of the middle element is very close to the arithmetic mean of the other two elements. Dobereiner’s relationship is referred as Law of triads. Since Dobereiner’s Law of triads worked only for few elements, it was dismissed. Chancourtois arranged elements in order of increasing atomic weights and made a cylindrical table of elements. John Newland arranged the elements in the increasing order of atomic weight and noted that the properties of the every eighth element are similar to the first one. This relationship is known as “Law of octaves”. Lothar Meyer proposed that on arranging the elements in order of increasing atomic weights similarities appear at a regular interval in physical and chemical properties. According to Mendeleev’s periodic law the physical and chemical properties of elements are periodic functions of their atomic weights. Merits of Mendeleev’s periodic table:     Mendeleev’s periodic table was very helpful in remembering and studying the properties of large number of elements Mendeleev’s periodic table helped in correcting the atomic masses of some of the elements like gold, beryllium and platinum based on their positions in the periodic table. Mendeleev could predict the properties of some undiscovered elements like scandium, gallium and germanium. By this intuition, he had left gaps for the undiscovered elements while arranging elements in his periodic table. Demerits of Mendeleev’s periodic table:       Position of hydrogen is not correctly defined in periodic table. It is placed in group I though it resembles both group 1 and 17. In certain pairs of elements increasing order of atomic masses was not obeyed. For example argon (Ar, atomic mass 39.9) is placed before potassium (K, atomic mass 39.1). Isotopes were not given separate places in the periodic table although Mendeleev's classification is based on the atomic masses. Some similar elements are separated and dissimilar elements are grouped together. For example copper and mercury resembled in their properties but had been placed in different groups. On the other hand lithium and copper were placed together although their properties are quite different. Mendeleev did not explain the cause of periodicity among the elements. Lanthanides and actinides were not given a separated position in the table. Modern Periodic Table:             According to Modern periodic law the physical and chemical properties of the elements are periodic functions of their atomic numbers. Modern periodic table is also referred to as long form of periodic table. Horizontal rows in the periodic table are called periods. Vertical columns in the periodic table are called groups. In the modern periodic table there are 7 periods and 18 groups. The period number corresponds to highest principal quantum number of elements. First period contains 2 elements. – also called as shortest period Second and third period contains 8 elements. – called as shorter periods Fourth and fifth period contains 18 elements. – called as longer periods Sixth period contains 32 elements. – called as longest period In the modern periodic table, 14 elements of both sixth and seventh periods i.e. lanthanides and actinides respectively are placed separately at the bottom of the periodic table. Elements with atomic number greater than 92 are called transuranic elements. According to IUPAC, until a new element’s discovery is proved and its name is officially recognized it is given a temporary name. This nomenclature is based Latin words for their numbers.
• Periodic table and periodic properties Classification of elements into blocks:                     The modern periodic table is divided into four main blocks – s -block, p-block, d-block and f-block depending on the type of orbital that are being filled with exception of hydrogen and helium. The elements in which last electron enter the s-orbital of their outermost energy level are called s-block elements. The s-block consists of two groups, Group-1 and Group-2. The elements of Group-1 are called alkali metals and have ns1 as the general outer electronic configuration. The elements of Group-2 are called alkaline earth metals and have ns2 as the general outer electronic configuration. The elements in which last electron enter the p-orbital of their outermost energy level are called p-block elements. The p-block elements constitute elements belonging to group 13 to 18. Elements of s-block and p-block are collectively called representative element The outermost electronic configuration of p-block elements varies from ns2np1 to ns2np6. Elements of group 18 having ns2np6 configuration are called noble gases. Elements of group 17 are called halogens. Elements of group 16 are called chalcogens. Number of valence electrons in group =Group number -10 for elements belonging to group 13 to 18. 36. Elements in which the last electron enters d-orbitals of penultimate energy level constitute d-block elements. Elements of group 3 to 12 in the centre of periodic table constitute the d-block elements. General outer electronic configuration of d-block elements is (n-1) d 1-10 ns 1-2. d-block elements constitute transition series elements. The name “transition series” is derived from the fact the d-block elements represent transition in character from reactive metals (belonging to group1 and 2 constituting s-block) on one side of the periodic table to non-metals (belonging to group 13 to 18 constituting p-block) on other side of the periodic table. Elements in which last electron enters f-orbitals are called f-block elements. Elements of Lanthanide series have general outer electronic configuration of 4f 1-14 5d 0-1 6s2. Elements of Actinide series have general outer electronic configuration of 5f 1-14 6d 0-1 7s2. Elements in lanthanide and actinide series are called inner transition series.
• Periodic table and periodic properties Periodicity:   The recurrence of similar properties of elements after certain regular intervals when they are arranged in order of increasing atomic number is called periodicity. The cause of periodicity of properties of elements is due to the repetition of similar electronic configuration of their atoms in the outermost energy shell after certain regular interval. Effective Nuclear Charge: • In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. • The nuclear charge that an electron experiences depends on both factors. • Effective nuclear charge is the net positive charge experienced by an electron. • The effective nuclear charge is not the same as the charge on the nucleus because of the effect of the inner electrons. • The electron is attracted to the nucleus, but repelled by the inner-shell electrons that shield or screen it from the full nuclear charge. • The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of electrons in the • • • spherical volume out to the electron in question. As the average number of screening electrons (S) increases, the effective nuclear charge (Zeff) decreases. Zeff = Z – S As the distance from the nucleus increases, S increases and Zeff decreases. S is called the screening constant which represents the portion of the nuclear charge that is screened from the valence electron by other electrons in the atom. The value of S is usually close to the number of core electrons in an atom. • Metal, Non-metals and Metalloids: • Metallic character refers to the extent to which the element exhibits the physical and chemical properties of metals. • Metallic character increases down a group. • Metallic character decreases from left to right across a period. Metals: • Metals are shiny and lustrous, malleable and ductile. • Metals are solids at room temperature (exception: mercury is liquid at room temperature; gallium and caesium melt just above room temperature) and have very high melting temperatures. • Metals tend to have low ionization energies and tend to form cations easily. • Metals tend to be oxidized when they react. • Compounds of metals with non-metals tend to be ionic substances. • Metal oxides form basic ionic solids. • Most metal oxides are basic: Metal oxide + water ------> metal hydroxide : Na2O(s) + H2O (l) ------> 2NaOH (aq) • Metal oxides are able to react with acids to form salts and water: Metal oxide + acid ------> salt + water: MgO(s) + 2HCl (aq) ------> MgCl2 (aq) + H2O (l) Non-metals: • Non-metals are more diverse in their behaviour than metals. • In general, non-metals are non-lustrous, are poor conductors of heat and electricity, and exhibit lower melting points than metals. • Seven non-metallic elements exist as diatomic molecules under ordinary conditions: • H2(g), N2(g), O2(g), F2(g), Cl2(g), Br2(l), I2(s) • When non-metals react with metals, non-metals tend to gain electrons: • • Metal + non-metal ------> salt 2Al(s) + 3Br2 (l) ------> 2AlBr3 (s) Compounds composed entirely of non-metals are molecular substances. Most non-metal oxides are acidic: Non-metal oxide + water ------> acid P4O10(s) + 6H2O (l) ------> 4H3PO4 (aq)
• Periodic table and periodic properties CO2 (g) + H2O (l) ------> H2CO3 (aq) • Non-metal oxides react with bases to form salts and water: Non-metal oxide + base ------> salt + water CO2 (g) + 2NaOH (aq) ------> Na2CO3 (aq) + H2O (l) Metalloids: • Metalloids have properties that are intermediate between those of metals and non-metals. • Example: Si has a metallic lustre but it is brittle. • Metalloids have found fame in the semiconductor industry. • In general metallic character increases down the group and decreases along period. • In general non-metallic increases along a period and increases along group. Periodic Trends in Atomic Radii: • Atomic size varies consistently through the periodic table. • As we move down a group the atoms become larger. • As we move across a period atoms become smaller. • There are two factors at work: • The principal quantum number, n, and • The effective nuclear charge, Zeff. • As the principal quantum number increases (i.e., we move down a group), the distance of the outermost electron from the nucleus becomes larger. Hence the atomic radius increases. • As we move across the periodic table, the number of core electrons remains constant, however, the nuclear charge increases. Therefore, there is an increased attraction between the nucleus and the outermost electrons. This attraction causes the atomic radius to decrease. Periodic Trends in Ionic Radii: • Ionic size is important in predicting lattice energy. • In determining the way in which ions pack in a solid. • Just as atomic size is periodic, ionic size is also periodic. • In general, Cations are smaller than their parent atoms. • Electrons have been removed from the most spatially extended orbital. • The effective nuclear charge has increased. • Therefore, the cation is smaller than the parent atom. • Anions are larger than their parent atoms. • Electrons have been added to the most spatially extended orbital. • This means total electron-electron repulsion has increased. • Therefore, anions are larger than their parent atoms. • For ions with the same charge, ionic size increases down a group. • All the members of an isoelectronic series have the same number of electrons. • As nuclear charge increases in an isoelectronic series the ions become smaller: • Covalent radius for a homo nuclear molecule is defined as one half of the distance between the centres of nuclei of two similar atoms bonded by single covalent bond.