Chapter 7_Jan14 New Version (Shortened)

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CHAPTER 7CHAPTER 7

Periodic Properties of the Elements

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CONTENTCONTENT

7.1 Development of the Periodic Table7.2 Effective nuclear charge and the

Sizes of Atoms7.3 Ionization Energy7.4 Electron Affinities7.5 Metals, Nonmetals, and Metalloids7.6 Group Trends for the Active Metals7.7 Group Trends for Selected Nonmetals

7.6 and 7.7 sub-chapters : SELF-LEARNING !!!

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Learning outcomes

Able to relate effective nuclear charge to size, ionization energy and electron affinity of elements in periodic table

To differentiate metal, non-metal and metalloid

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7.1 Development of The 7.1 Development of The Periodic TablePeriodic Table

The periodic table was first developed by Mendeleev and Meyer on the basis of the similarity in properties and reactivities exhibited by certain elements.

Elements in the same column of the periodic table have the same number of electrons in their valence orbitals.

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Con’t: 7.1 Development of Con’t: 7.1 Development of The Periodic TableThe Periodic Table

Group 1A: Group 6A:

3Li - [He] 2s 1 8O - [He] 2s 2 2p 4

11Na - [Ne] 3s 1 16S - [Ne] 3s 2 3p 4

19K - [Ar] 4s 1

37Rb - [Kr] 5s 1

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7.2 Effective Nuclear 7.2 Effective Nuclear Charge and Sizes of AtomsCharge and Sizes of Atoms

7.2.1 Effective Nuclear Charge

The net positive charge experienced by an electron on a many-electron atom.

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.

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Con’t: 7.2.1 Effective Nuclear Charge

• The shielding effect is called the screening effect.

• The nuclear charge experienced by an electron depends on its distance from the nucleus and also the screening effect.

Zeff = Z – S

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The principal quantum number, n, of the valence orbitals of the atoms changes from top to bottom of the Periodic Table.

All orbitals with the same value of n are referred to as a shell.

7.2.2 Electron Shells in 7.2.2 Electron Shells in AtomsAtoms

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Con’t: 7.2.2 Electron Shells in Con’t: 7.2.2 Electron Shells in AtomsAtoms

Consider the noble gases:Nuclear charge

He 1s 2 2+Ne 1s 2 2s 2 2p 6 10+Ar 1s 2 2s 2 2p 6 3s 2 3p 6 18+

Plot radial electron-density graph for electron distribution in these atoms.

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Con’t:

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The maxima (peaks) appear based on electrons having the same quantum number, n.

n = 1 - 1st peak (1s)n = 2 - 2nd peak (2s 2p)n = 3 - 3rd peak (3s 3p 3d)

Con’t: 7.2.2 Electron Shells in Con’t: 7.2.2 Electron Shells in AtomsAtoms

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From the graph:

1s electrons for Helium show a maximum in radial electron density - 0.3 Å

1s electrons for Argon - 0.05 Å only

Con’t: 7.2.2 Electron Shells in AtomsCon’t: 7.2.2 Electron Shells in Atoms

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Question

• Why is the 1s shell in Argon so much

closer to the nucleus than the 1s shell

in Helium?

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Answer

Nuclear charge of He is 2+ and Ar 18+.

1s electrons (innermost electrons of the atom), are not shielded from the nucleus.

Thus, as the nuclear charge increases (2+ 18+), the 1s electrons are “pulled” closer and closer to the nucleus.

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7.2.3 Atomic Sizes7.2.3 Atomic Sizes

Atoms do not have boundaries that fix their sizes.

Atomic radius is used as a mean to estimate the radius of an atom.

To determine atomic radii, assume atoms are spheres that touch each other when they are bonded together.

Example: C-C bond is 1.54 Å therefore radius of C atom is 0.77 Å

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Con’t: 7.2.3 Atomic SizesCon’t: 7.2.3 Atomic Sizes

1. Within each column (group) the atomic radius (and also the size of orbitals) tends to increase when n increases.

(a) n , orbital size (b) n , Zeff remains relatively

constant

n , Atomic radius

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Example:

n = 2 Li : 1.52 Ån = 3 Na : 1.86 Ån = 4 K : 2.27 Ån = 5 Rb : 2.47 Å

Radius increases

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2. Within each row (period), the atomic radius decrease as we move from left to right.

(a) n constant, orbital size constant

(b) number of core electrons stay the same , nuclear charge (Z) acting on the electron valence increase

Zeff , Atomic radius

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same nZeff increases

B C N O F

0.88 Å 0.77 Å 0.75 Å 0.73 Å 0.71 Å

radius decreases

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Con’t:

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(Z = 3) Li 1s 2 2s 1 (1.52 Å)(Z = 4) Be 1s 2 2s 2 (1.13 Å)

For Li, 1s2 electrons shield the outer 2s1 electron from the 3+ (Z value) charge nucleus. The outer electron 2s1 electron experiences Zeff of slightly more than 1+.

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Beryllium, outer 2s2 electrons experience Zeff larger. The inner 1s2 electrons are shielding a 4+ nucleus. The Zeff experienced by the 2s electron is closer to 2+.

As the effective nuclear charge increases, the electrons are drawn closer to the nucleus.

The radius decreases as we move from left to right.

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7.3 Ionization Energy7.3 Ionization Energy

Ionization is a process of removing an electron from an atom or ion.

Ionization energy of an atom or ion - the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion.

The ease in removing electrons from an atom is an important indicator of the atom’s chemical behaviour.

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Con’t: 7.3 Ionization EnergyCon’t: 7.3 Ionization Energy

First ionization energy, I1 (or IE1) - energy needed to remove the first electron from a neutral gaseous atom.

Na (g) Na+ (g) + e-

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Second ionization energy, I2 (or IE2) - energy required to remove the second electron from a gaseous ion.

Na+ (g) Na2+(g) + e-

Ionization energies I1, I2 are always positive (endothermic) where energy is absorbed from the surrounding.

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Con’t:

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The greater the ionization energy, the more difficult to remove an electron.

Magnitude I1 < I2 < I3Reason: The positive nuclear charge remains the same, the number of electrons (produce repulsive interactions) decreases.

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Con’t:

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Sharp increase in ionization energy when an inner-shell electron (noble gas core) is removed.

Example:

Silicon: 1s 2 2s 2 2p 6 3s 2 3p 2 or [Ne] 3s 2 3p 2

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Ionization energy increases from: I1 I2 I3 I4 I5 (kJ/mol)

786 - 3230 4360 161003s 2 3p 2 3s 2 3p 1 3s 2 3s 1 1s 2 2s 2 2p

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I1 I4 : 786 kJ/mol 4360 kJ/mol

Loss of the four electrons in the 3s and 3p subshells.

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I4 I5 : 4360 kJ/mol 16100 kJ/mol

The inner shell 2p electron (core electron) is much closer to the nucleus - greater Zeff.

Large increase in ionization energy when electrons are removed from its noble gas core.

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The outermost electrons take part in:a) chemical bondingb) reaction

The core (noble-gas core) tightly bound to the nucleus.

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7.3.1 Periodic Trends in 7.3.1 Periodic Trends in Ionization EnergiesIonization Energies

Generally:

1. Within each row, the alkali metals show the lowest ionization energy and the noble gases the highest.

2. Each column (group), the ionization energy decreases with increasing atomic number.Ionization Energy He > Ne > Ar > Kr > Xe

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Con’t: 7.3.1 Periodic Trends Con’t: 7.3.1 Periodic Trends in Ionization Energiesin Ionization Energies

3. Ionization energy of the transition elements increase slowly from left to right.

4. The f -block elements show only a small variation in the values of I1.

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Attraction of electrons to the nucleus depends on:

The effective nuclear charge (Zeff)The average distance of the electron

from the nucleus (atomic radius).

Attraction , Ionization energy

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Move across a row: Increase in Zeff and decrease in

atomic radius.


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Move down a column:The atomic radius increases, while

Zeff remains constant


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Irregularities:a) Decrease in ionization energy from:

Beryllium [He] 2s 2 Boron [He] 2s 2 2p1

o The electrons in the filled 2s orbital shielding the electrons in 2p.

o Zeff decreases from 2+ 1+.

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b) Decrease in ionization energy from:

Nitrogen [He] 2s 2 2p 3 Oxygen [He] 2s 2 2p 4

o Due to repulsion of paired electron in the p4 configuration.

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Example 1

Arrange the following atoms in order of increasing first ionization energy:

Ne, Na, P, Ar, K

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Example 1 (Answer)

Use trends to predict.1. Na, P and Ar are in the same row.

exhibit the order of Na<P<Ar(I1 increases from left to right)

2. Ne is a noble gas and above Ar. Ne exhibits the greater ionization energy.(I1 decreases from top to bottom)

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Example 1 (Answer)

3. K below Na.

I1 K is less than Na.

K < Na < P < Ar < Ne

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7.4 Electron Affinities7.4 Electron Affinities

Measure of attraction towards electron or the ease of an atom to gain electron.

The energy change (E) that occurs when an electron is added to a gaseous atom is called the electron affinity.

Cl(g) + e- Cl-(g)

[Ne] 3s2 3p5 [Ne] 3s2 3p6

E(energy) = EA1 = -349 kJ/mol

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Con’t: 7.4 Electron AffinitiesCon’t: 7.4 Electron Affinities

Energy is released when an electron is added.

The electron affinity of Cl is -349 kJ/mol. The greater the attraction, the more negative

the electron affinity will be.EA1 = -x

the greater the affinity the more negative the value easier to gain electron

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Halogen (F, Cl, Br, I) - by gaining an electron, a halogen atom forms a stable negative ion that has a noble gas configuration.Noble gas configuration: Ne : 1s2 2s2 2p6

Ar : 1s2 2s2 2p6 3s2 3p6

Cl(g) + e- Cl-(g)

[Ne]3s23p5 [Ne] 3s2 3p6 or [Ar]

p subshell (orbital 3p) is filled to form a stable negative ion similar to Ar.

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Positive Value of the Electron Affinity

The noble gases posses positive values of electron affinity.

The anion is higher in energy than the atom.Ar(g) + e- Ar-

(g) EA1 > 0

[Ne] 3s 2 3p 6 [Ne] 3s 2 3p 6 4s 1

As EA1 > 0, the Ar- is not stable and will not form.

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The addition of an electron to a noble gas requires the electron to reside in a new, higher energy subshell or in 4s orbital.

Occupying a higher-energy subshell is unfavourable.

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Generally, electron affinity becomes increasingly negative from left to right i.e from alkali metals to halogens.

Electron affinities of Be : [He] 2s2 and Mg: [Ne] 3s2 are positive. The added electron would reside in an empty p subshell that is higher in energy (Hund’s first rule).

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Group 5A

N, P, As, Sb

N [He] 2s 2 2p 3 0P [Ne] 3s 2 3p 3 -72 kJ/molAs [Ar] 4s 2 3d 10 4p 3 -78 kJ/molSb [Kr] 5s 2 4d 10 5p 3 -103 kJ/mol

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These elements have half-filled p subshells. The added electron must be placed in an orbital that is already occupied - resulting larger electron-electron repulsion (Hund’s first rule).

As we proceed from top to bottom, the average distance of the added electron from the nucleus increases. The electron-nucleus attraction decreases.

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N: [He] 2s22p3 Sb : [Kr] 5s 2 4d 10 5p 3

strong electron-electron repulsion

lower electron-electron

repulsion

will not accept electron

EA1 > 0 EA1 = -103 kJ/mol

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Con’t:

g

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You can explain:F [He] 2s 22p 5 - 328 kJ/mol

Br [Ar] 4s 2 3d1 0 4p 5 - 325 kJ/mol

I [Kr] 5s 2 4d 10 5p 5 - 295 kJ/mol

5p 5 - away from nucleus, therefore less nucleus

attraction.

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7.5 Metals, Nonmetals and 7.5 Metals, Nonmetals and MetalloidsMetalloids

Properties of individual atoms:

1. Atomic radii2. Ionization energies3. Electron affinities

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7.5.1 Metals7.5.1 Metals

Properties: Metals conduct heat and electricity. They are malleable - can be pounded into

thin sheets. Ductile - can be drawn into wire. Shiny luster All are solids (except Hg) High melting point (except Cs, Ga)

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Con’t: 7.5.1 MetalsCon’t: 7.5.1 Metals

Properties (con’t): Metals have low ionization energies -

oxidised (lose electrons) when undergo chemical reaction.

Metal oxides are ionic solids (basic)

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Charge:

Alkali metals 1+.Alkali earth metal 2+.Transition metals ions 2+, 1+, 3+.

Able to form more than one positive ion.

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Basic oxide:

Most metal oxides are basic oxides. They dissolve in water, react to form metal hydroxides.Metal Oxide + Water Metal HydroxideNa2O(s) + H2O(l) 2NaOH(aq)

CaO(s) + H2O(l) Ca(OH)2(aq)

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React with acids to form salts and water:

Metal Oxide + Acid Salt + Water

MgO(s) + 2HCl(aq) MgCl2(aq) + H2O(l)

NiO(s) + H2SO4(aq) NiSO4(aq) + H2O(l)

Na2O(s) + H2SO4(aq) Na2SO4(aq) + H2O(l)

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Example 2Example 2

a) Write the chemical formula of aluminium oxide.

b) Write the balanced chemical equation for the reaction of aluminium oxide with nitric acid.

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Example 2 (Answer)

a) Aluminium has a 3+ charge, Al3+.The oxide ion is O2-

Al2O3

b) Metal oxides react with acids to form salts and water.

Al2O3(s) + 6HNO3(aq) 2Al(NO3)3(aq) + 3H2O(l)

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7.5.2 Nonmetals7.5.2 Nonmetals

Not lustrous (not shiny). Poor conductors of heat and electricity. Melting points - generally lower than those

of metals (except diamond : 3570 C). Seven nonmetals exist as diatomic

molecules.oH2,N2, O2, F2, Cl2 - gases.

oBr2 - liquid

oI2 - volatile solid

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Con’t: 7.5.2 NonmetalsCon’t: 7.5.2 Nonmetals

o Nonmetals, reacting with metals, gain electrons and become anions.

Metal + Nonmetal Salt

2Al(s) + 3Br2(l) 2AlBr3(s)

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Nonmetals gain electrons to fill their outer p subshell giving a noble-gas electron configuration.

Most nonmetal oxides are acidic oxides (molecular substance).Those dissolve in water react to form acids.

Example:Selenium dioxide - SeO2

Tetraphosphorus hexoxide - P4O6(s)

Tetraphosphorus decoxide - P4O10(s)

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Reaction with water:SeO2(s) + H2O(l) H2SeO3 (aq)

P4O6(s) + 6H2O(l) 4H3PO3(aq)

P4O10(s) + 6H2O(l) 4H3PO4(aq)

Nonmetal oxide + base salt + waterCO2(g) + 2NaOH(aq) Na2CO3(aq) + H2O(l)

SO3(g) + 2KOH(aq) K2SO4(aq) + H2O(l)

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7.5.3 Metalloids7.5.3 Metalloids

Metalloids have properties intermediate between those of metals and nonmetals.

They may have some characteristic metallic properties but lack others.

Example: silicon looks like a metal but it is brittle rather than malleable and a much poorer conductor of heat and electricity than metals.

Several metalloids are semiconductors ( most notably silicon).

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7.6 Group Trends for The 7.6 Group Trends for The Active MetalsActive Metals

7.6.1 Group 1A: The Alkali Metals

The alkali metals are soft metallic solids. Possess characteristics such as silvery, metallic

luster and high thermal and electrical conductivities.

As moving down the group, we observe expected trends such as increasing atomic radius and decreasing first ionization energy.

Alkali metals have low I1 values, the outer s electrons can be removed relatively easy.

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Con’t: 7.6.1 Group 1A: The Con’t: 7.6.1 Group 1A: The Alkali MetalsAlkali Metals

As a result, the alkali metals are all very reactive, readily losing one electron to form ions with a 1+ charge:

M M+ + e-

*M represents any one of the alkali metals

The alkali metals are the most active metals and thus exist in nature only as compounds.

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Electrolysis is the process used to obtain metals from compounds.

The chemistry of the alkali metals is dominated by the formation of 1+ cations.

The metals combine directly with most nonmetals.

Examples:2M(s) + H2(g) 2MH(s)

2M(s) + S(s) M2S(s)

2M (s) + Cl2(g) 2MCl(s)

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The hydrides of the alkali metals (LiH, NaH, and so forth): hydrogen is present as H-, called the hydride ion.

Note the difference between hydride ion H- and hydrogen ion H+.

The alkali metals react vigorously with water to produce hydrogen gas and solutions of alkali metal hydroxides.2M(s) + 2H2O(l) 2MOH(aq) + H2(g)

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This reaction is very exothermic (heat is released).

This reaction is most violent for the heavier members of the group - weaker hold on the single outer-shell electron.

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Reaction with oxygen:1. Li reacts with oxygen to form lithium oxide,

which contains O2- ion. 4Li(s) + O2(g) 2Li2O(s) lithium oxide

2. Other alkali metals react with oxygen to

form metal peroxides, which contain O22- ion.

2Na(s) + O2(g) Na2O2(s) sodium peroxide

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Potassium, rubidium, and cesium also form compounds that contain the O2

- ion, called superoxides.K(s) + O2(g) KO2(s) potassium superoxide

As the alkali metals are extremely reactive toward water and oxygen, they are usually stored in hydrocarbon, such as kerosene or mineral oil.

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Alkali metals salts and their aqueous solutions are colourless unless they have a coloured anion eg yellow CrO4

2-.

When alkali metal compounds are placed in a flame, they emit characteristic colours.

(Li: red, Na: yellow, K: blue)

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7.6.2 Group 2A: The 7.6.2 Group 2A: The Alkaline Earth MetalsAlkaline Earth Metals

The elements are all solids with typical metallic properties.

Compare to elements in Group 1A (alkaline metals), alkaline earth metals are harder, more dense and melt at higher temperatures.

Their I1 are low, but not as low as those of alkali metals.

Less reactive than alkali metals.

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Con’t: 7.6.2 Group 2A: The Con’t: 7.6.2 Group 2A: The Alkaline Earth MetalsAlkaline Earth Metals

Be and Mg are the least reactive. Be does not react with water or steam. Mg does not react with water but does react

with steam to form Magnesium oxide and hydrogen. Mg(s) + H2O (g) MgO (s) + H2 (g)

The other elements react readily with water (less reactive than alkali metals).Ca (s) + 2H2O (s) Ca(OH) 2 (aq) + H2 (g)

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Con’t: 7.6.2 Group 2A: The Con’t: 7.6.2 Group 2A: The Alkaline Earth MetalsAlkaline Earth Metals

Pattern in the reactivity of the alkaline earth metals - the tendency to lose their two outer s electrons and form 2+ ions.

Example:Mg(s) + Cl2(g) MgCl2(s)

2Mg(s) + O2(g) 2MgO(s)

Like the 1+ ions of the alkali metals, the 2+ ions of the alkaline earth elements have a noble gas configuration.

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cont

r

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cont

r

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7.7 Group Trends for 7.7 Group Trends for Selected NonmetalsSelected Nonmetals

7.7.1 Hydrogen The first element in the periodic table. 1s1 electron configuration. Placed above alkali

metals. Unique element, nonmetal, exists as diatomic

gas, H2(g), under most conditions. The ionization energy of hydrogen, 1312

kJ/mol, is markedly higher than that of the active metals.

Reason: the complete absence of nuclear shielding of its sole electron.

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Con’t: 7.7.1 HydrogenCon’t: 7.7.1 Hydrogen

React with other nonmetals to form molecular compounds.

2H2(g) + O2(g) 2H2O(l)

Hydrogen reacts with other active metals to form solid metal hydrides, contain the hydride ion, H-.

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7.7.2 Group 6A: The Oxygen 7.7.2 Group 6A: The Oxygen GroupGroup

Increase in metallic character as moving down the group.

Oxygen is colourless gas, the rest are solids.

Oxygen, sulfur, and selenium are typical nonmetals.

Tellurium is a metalloid. Polonium is a metal.

OSSeTePo

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Con’t: 7.7.2 Group 6A: The Con’t: 7.7.2 Group 6A: The Oxygen GroupOxygen Group

Oxygen occurs in two forms, O2 and O3.

O2 - dioxygen (normally called as oxygen)

O3 – ozone

Ozone - toxic and pungent gas. It is also formed from O2 in electrical discharge, eg lightning storm.3O2 (g) 2O3(g)

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Oxygen has a great tendency to attract electrons from other elements ( to oxidise them).

Form oxide, O2- ion. This ion has a noble gas configuration and thus stable.

Peroxide, O22- and superoxide, O2

- ions often react with themselves to produce O2- and O2.

Eg.: 2H2O2(aq) 2H2O(l) + O2(g)

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After oxygen, the most important element is sulfur.

The most common and stable is the yellow solid, S8.

Sulfur is written simply as S(s). Sulfur has the tendency to gain electron

forming sulfides, S2-.Eg.: 2Na(s) + S(s) Na2S(s)

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Most sulfur in nature is present as metal sulfides.

The chemistry of sulfur is more complex than that of oxygen.

Sulfur can be burned in oxygen to produce sulfur dioxide, the main pollutant.

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7.7.3 Group 7A: The 7.7.3 Group 7A: The HalogensHalogens

As we move from group 6A to 7A, the nonmetallic behaviour of the elements increases.

All the halogens are typical nonmetals (except At: metalloid).

Melting points and boiling points increase with increasing atomic number.

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Con’t: 7.7.3 Group 7A: The Con’t: 7.7.3 Group 7A: The HalogensHalogens

Fluorine and chlorine - gas at room temperature.

Bromine - liquid Iodine – solid (sublime easily) Each element consists of diatomic

molecules: F2, Cl2, Br2, and I2.

F2 - pale yellow gas; Cl2(g) - yellow-green colour; Br2(l) - reddish brown; I2(s) - greyish black.

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The halogens have highly negative electron affinities.

The chemistry of halogens are dominated by their tendency to gain electrons from other elements to form halide ions:

X2 + 2e- 2X-

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Fluorine and chlorine are more reactive than bromine and iodine.

2Na(s) + F2(g) 2NaF(s)

2H2O(l) + 2F2(g) 4HF(aq) + O2(g)

Fluorine gas is difficult and dangerous to be used in laboratory because it is very reactive.

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Chlorine is the most industrially useful of the halogens.

Chlorine is usually produced by electrolysis of molten NaCl:

electricity

2NaCl(aq) + 2H2O(l) 2NaOH(aq) + H2(g) + Cl2(g)

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Chlorine react slowly with water to form relatively stable aqueous solutions of HCl and HOCl (hypochlorous acid):

Cl2(g) + H2O(l) HCl(aq) + HOCl(aq)

The halogens react directly with most metals to form ionic halides. Also react with hydrogen to form gaseous hydrogen halide compounds.

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cont

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7.7.4 Group 8A: The Noble 7.7.4 Group 8A: The Noble GasesGases

The elements are all nonmetals that are gases at room temperature.

They are all monoatomic (consist of single atoms rather than molecules).

The noble gases have completely filled s and p subshells.

All elements have high first ionization energies, the values decrease as moving down the group.

The noble gases are exceptionally unreactive. Also called inert gases.

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cont

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END OF CHAPTER 7

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Chapter 7_Jan14 New Version (Shortened)