periodicity 1

8/13/2019 Periodicity 1

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Syllabus

Periodic variation in physical properties of the

elements H to Ar

Variations in first ionization enthalpies, atomic radii,electronegativitiesand melting points.

Interpretation of these variations in terms of structure andbonding.

Periodic relationship among the oxides of theelements Li to Cl

Bonding and stoichiometric composition of the oxidesofthese elements, and their behaviour with water, diluteacids and alkalis.
http://localhost/var/www/apps/conversion/Clevel/ppt-ch7/PerTable.exe


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The Periodic Table

The elements are arranged in the order of atomic number

Do. Q. 1b


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Elements were first arranged in order of increasing atomic

massesby Dimitri Mendeleev(1834 - 1907)

The elements were observed to repeat their properties periodically

(a) (b)

Periodic Table (early forms)


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Modern periodic table: (p.1)

Rows periods Columns groups

Classified into 4 areas:

p-blocks-block

d-block

transition elements

f-block inner transition elements


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s-block elements (p.2):

Group IA: alkali metals

1 ein outermost shell (ns1)

(e.g. Li, Na, K)

Group IIA: alkaline earth metals

2 ein the outermost shell (ns2)

(e.g. Be, Mg, Ca)

p-block elements:

Groups IIIA, IVA, VA, VIA, VIIA, 0

Group VIIA : halogens (ns2np5)

Group 0 : noble gases (ns2 np6)

d-block elements:

Electronic configuration : (n1)d1ns2 to (n1)d10ns2

(Group IIIB) (Group IIB)

Transition elements


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f-block elements:

Lanthanide series and actinide series :

4fand 5forbitals are filled up with 1 to 14 e

-

inner-transition elements

Aims of Periodic Table: (p. 4):

1. Similar elements to be grouped together as families;

2. Gradual changes in properties such as electronegativity,

ionization enthalpies.


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Covalent radius is defined as half the internuclear distance

between two covalently bonded atoms in a molecule of the

element.

Atomic Radius

(p.9) How can scientists measure the sizes of atoms?

(1) For non-metals, atomic radius refers to the covalent radius:

(2) For metals, atomic radius refers to the metallic radius:

Metallic radius is defined as half the internuclear distancebetween atoms in a metall ic crystal.


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Across the period, the atomic radii decreaseprogressively

Variation in atomic radius of the first 20 elements

Down the group, the atomic radii increaseprogressively
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(1) Screening/Shielding of electrons (repulsion between

electrons)

(2) Attraction of the nucleus (protons) for the electrons

The atomic radius is governed by two factors: (p.6, notes p.9)


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Down a group:

Increase in number of electron shells.

Increase shielding effect from inner shells electrons.

Across a period:

Electrons add to the same outermost shellnot much

increase in shielding effect

More protonsgreater attraction to eoutweighs increase

in shielding effectsmaller size

Decrease along the period of transition series is small:

Electrons are added to inner d-orbitalsscreen the

outermost electron shell.


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Melting point (notes p.11)

Melting Temperature depends on the magnitude offorcesbetween particles

Metals: Metallic Bond

Giant Covalent Crystals: Covalent Bond Molecular Crystals:

Van der Waals forces

permanent dipolepermanent dipole


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Variation in melting point of the first 20 elements


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Structure and Bonding (p.4)


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1. Steady increasein melting point from Li to Beand Na to Al

no. of outermost/delocalised electrons increases

no. of protons increases

metallic bond strength increases

- similar forces exist in liquidmelting point not very high.

2. Carbon and silicon correspond to the maximain Periods

2 and 3

both have giant covalent structures. Atoms are held

together by strong covalent bonds.

Across a Period (p.11)


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1. Metallic structures

Very strong metallic bond because of the availability of d

electrons / orbitals for metallic bonding.

Maximum reaches in the middle except manganese as it has

stable half-filled structure.

2. Carbon and silicon correspond to the maximain Periods

2 and 3

both have giant covalent structures. Atoms are held

together by strong covalent bonds.

Across the Period of transition elements (p.13)


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More on carbon: (p.7, 11)

Carbon has two allotropes: graphite and diamond.

Which of them is more stable?

The (C-C) bond distance in graphite is 1.415 A while that in

diamond is 1.54 A.

Diamond is hard while graphite can be used as lubricant?

Why?

3. The melting points of elements from N to Neand P to

Arare relatively low

they exist as discrete moleculeswhich held by weak

van der Waals forces


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S has a highermelting point than P although the atomic

size of P is larger than S. (Why?)

S exists as S8moleculesin its molecular crystal whereasP exists as P4moleculesin its molecular crystal


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1. For Group I Metals, melting point decreases down the group.

Size increases

Shielding effect of inner shell e-increases

Metallic bond strength decreases

2. For halogens and noble gases, melting point increases down

the group.

Size increases

Van der Waals forces increases

Down a group (p.13)

Do Q. 2, Q.6 on p. 32


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a. Giant covalent structure

Large amount of energy used to

break the strong covalent bonds.

b. Metallic bond.

Valence electrons & Size Metallic bond

c. Metals: metallic bonds persist in

liquid.

Non-metals: weak van der Waals

forces.

d. Van der Waals forces determine

the m.p. in non-metals.

S8largest sizestrongest van

der Waals forces.


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F irst I onization Enthalpy (p.7, notes p.16)

X(g) X+(g) + e

The first ionization enthalpies of the first 20 elements


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Electronic configuration (s, p, half-filled?)

Nuclear charge

Screening/Shielding effect

Atomic radius

Fourfactors affecting the magnitude of ionization enthalpy:


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Variation in the first ionization enthalpy of thefirst 20 elements


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1. Noble Gases have the highest I.E. (p.17)

The electronic configuration of noble gases is very

stable(completely filled octet)

2. Alkali Metals have the lowest I.E. in a period (p.18)

It has the lowest effective nuclear charge in the period


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1. General increaseacross periods 2 and 3

increase in nuclear charge outweighs the increase in

shielding effect of additional electron of the same shell.

stronger attraction to outermost electrons

2. Irregularitieswith the general increase

Peaks in the general increase due to the extra stability

provided by full-filled ssub-shell(Be) and half-filled p

sub-shell (N)

3. Across a Period: (notes p.18)


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Explain why Group III elements have a lower

first I.E. than Group II elements? (p. 19)

i. Extra stability is gained for completely

filled s orbital in Group II elements.ii. For Group III elements, electron is

removed from p-orbital which is further

away (at higher energy level) from thenucleus and shielded by the s electrons.


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Explain why Group VI elements have a lower

first I.E. than Group V elements? (p.20)

i. Extra stability is gained for half filled p

orbital in Group V elements.ii. For Group VI elements, repulsion exists

between the first paired p-electrons


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Decreasedown a group

i. outermost electrons are further away from the

nucleus

ii. Shielding effect of inner shells electrons.

weaken attraction to outermost electrons

4. Down a group: (notes p.18)

Arrange the following in increasing first I.E.:

i. N, C, B

ii. B, Be, Li

iii. S+, S, S-


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Electronegativity (notes p.22)

Electronegativityis the measure of the relative tendency

of an atom to attract bond pair(s) electrons towards

itself in a covalent bond

Electronegativity values on an arbitrary scale from 0 to 4


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Electronegativity values of the first 20 elements


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Across the period, electronegativity increasesfromleft to right

Down the group, electronegativity decreases


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2. Decrease down groups

i. increase in size

ii. the increasing number of electron shells creates agreater shielding effect.

smaller attraction to bonding electrons

Explanations:

1. General increaseacross periods 2 and 3

increase in nuclear charge outweighs the increase in

shielding effect of additional electron of the same shell.

stronger attraction to outermost electrons


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Li and Mg; Be and Al; B and Si show similarproperties: like 1stI.E. and electronegativity

Why?

Shielding effect increases down a group andnuclear charge increases across a period.

Ionization enthalpy/ electronegativity of elementsdiagonally below one another are similar

They form bond with similar strength / show similarchemical properties (will be discussed later)

However, C and P, N and S showNOdiagonalrelationship. Why?

C and N has no low lying empty d-orbital

Diagonal Relationship (notes p. 23)


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The END

periodicity 1

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