Chemistry Topic 4 Matriculation

7/24/2019 Chemistry Topic 4 Matriculation

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Periodic Table

Topic 4

Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.


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The Periods in the Periodic Table

2

1. The modern periodic table is constructed based on the proton

numbers of the elements.2. The modern periodic law states that the properties of the

elements are a periodic function of their proton number.

3. The horizontal rows are called periods. There are seven

periods in the periodic table.


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3

When the Elements Were Discovered


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4

1. The short periods consist of period 1, 2 and 3

a.

Period 1 -1s orbital is filled. Only two elements, H2(1s1

) and He (1s2

).b.

Period 2 In the outermost shell, 2s orbital is filled first, followed by the

2p orbitals. There are eight elements, from Li (1s22s1) to Ne (1s22s2

2p6).

c.

Period 3 In the outer most shell, 3s orbital is filled first, followed by the

3p orbitals. There are eight elements, from Na (1s22s22p63s1) to Ar(1s22s22p63s23p6).

The Short Periods


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5

1. The short periods consist of period 4, 5, 6 and 7.

a.

Period 4 and 5 consists of 18 elements.

b.

Period 6 consists of 32 elements. Series of 15 elements known asLanthanides is removed from this period and placed at the bottom.

c.

Period 7 consists of 32 elements. Series of 15 elements known as

Actinides is removed from this period and placed at the bottom. The

number of elements in period 7 is slowly increasing due to discovery of

new elements. These new elements are given temporarily names suchas ununbium Uub by the IUPAC.

The Long Periods


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6

1. The vertical columns in the periodic table are called group.

2. In the past these groups are numbered from I to VIII except the

d-block elements. The new system by IUPAC number these

group from 1 to 18. With group 3 to 12 are called d-block

elements.

The Groups in the Periodic Table


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7

3. Elements in the same group have the same number ofelectrons in the valence shell and have similar chemical

properties even though their physical properties are different.

4. The first element in each group usually has special properties.

5. There are often significant differences in chemical properties

between the first and the second members of the group.

The Groups in the Periodic Table


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8

1. The periodic table is also classified into main blocks based onorbital types.

a.

S-block outermost electrons in s orbitals

b.

P-block outermost electrons in p orbitals

c.

D-block outermost electrons in d orbitals

d.

F-block outermost electrons in f orbitals

The s, p, d, f block in the Periodic Table


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9

Electron Configurations of Cations and Anions

Na [Ne]3s1 Na+ [Ne]

Ca [Ar]4s2 Ca2+ [Ar]

Al [Ne]3s

2

3p

1

Al

3+

[Ne]

Atoms lose electrons so thatcation has a noble-gas outer

electron configuration.

H 1s1 H- 1s2or [He]

F 1s22s22p5 F- 1s22s22p6 or [Ne]

O 1s22s22p4 O2- 1s22s22p6or [Ne]

N 1s22s22p3 N3- 1s22s22p6or [Ne]

Atoms gain electrons

so that anion has anoble-gas outer

electron configuration.

Of Representative Elements


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10

+1

+2

+3

-1

-2

-3

Cations and Anions Of Representative Elements


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11

Electron Configurations of Cations of Transition Metals

When a cation is formed from an atom of a transition metal,electrons are always removed first from the nsorbital and

then from the (n 1)dorbitals.

Fe: [Ar]4s23d6

Fe2+: [Ar]4s03d6or [Ar]3d6

Fe3+: [Ar]4s03d5or [Ar]3d5

Mn: [Ar]4s23d5

Mn2+: [Ar]4s03d5or [Ar]3d5


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Example 3.1

12


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An atom of a certain element has 15 electrons. Withoutconsulting a periodic table, answer the following questions:

(a)What is the ground-state electron configuration of theelement?

(b) How should the element be classified?

Exercise 3.1


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Strategy

(a) We start writing the electron configuration with principal

quantum number n = 1 and continuing upward until all theelectrons are accounted for.

(b) What are the electron configuration characteristics of

representative elements? transition elements? noble gases?

Exercise 3.1


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Solution

(a)We know that for n = 1 we have a 1s orbital (2 electrons); for

n = 2 we have a 2s orbital (2 electrons) and three 2p orbitals (6electrons); for n = 3 we have a 3s orbital (2 electrons). Thenumber of electrons left is 15 !12 = 3 and these three

electrons are placed in the 3p orbitals. The electron

configuration is 1s22s22p63s23p3.

(b) Because the 3p subshell is not completely filled, this is a

representative element. Based on the information given, we

cannot say whether it is a metal, a nonmetal, or a metalloid.

Exercise 3.1


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Quick Check 3.1


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17

1. Elements are classified into three categories: metals,metalloids and non metals.

2.

Metalloids (semi-metals) are elements that have both the

properties of metals and non-metals. They are poor conductors

of electricity, however conductivity may increase as

temperature increases and is affected markedly by impurities.

Metals, Metalloids and Non-metals in the Periodic Table


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18

3. In general, the chemical properties of metallic elementsdepend on the ability of the atoms to lose one or more

electrons. And their physical properties vary considerably.

4. Acid-base property of the oxide can also be used to classified

elements. Metals form basic oxides, non-metals form acidic

oxides and metalloids form amphoteric oxides.



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5. Metals like Na, Mg and Al conduct electricity in both solid andmolten states because they have delocalized electrons when

electric potential or voltage is applied.

6. N to F in period 2 and P to Cl in period 3 are non-metals. All

their valence electrons are used to form covalent bond. There

are no free no mobile electrons to conduct electricity, therefore,

they are non-conductors.7. Ne and Ar are noble gases which have stable octet electronic

configuration. They are also non-conductor of electricity.



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20

1. The elements in each group have similar chemical propertiesbecause they have similar electronic configurations.

2.

All elements in group 1 are metals and the outermost shell of

each elements has one electron. Their general electronic

configuration is ns1

3. Similarly, all element in group 13 have three electrons in their

outermost shell and their general electronic configuration is ns2

np1.

Electronic Configurations and the Positions of Elements in the Periodic Table

6


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ns1

ns2

ns2np1

ns2np2

ns2np3

ns2np4

ns2np5

ns2np6

d1

d5

d10

4f

5f

Ground State Electron Configurations of the Elements

E l 3 3 P i i i P i di T bl


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Example 3.3 Position in Periodic Table

22

E l 3 4 P iti i P i di T bl


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23



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24



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25



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26

Q i k Ch k 3 2


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Quick Check 3.2

At i d I i R dii


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Atomic and Ionic Radii1. The atomic radius is defined as

half the distance between the

nuclei of the two closest andidentical atoms.

3. An atom that contains only a few

electrons is not necessarily smaller

than atom which contains moreelectrons.

3. The atomic radius of an element is determined two factorsa.

Screening effect of inner shell electronsi.

Effects make atomic radius largerii.

Due to mutual repulsion between the inner shell electrons and theouter shell electrons

iii.

Filled inner shells shield the outer electrons from the nucleus moreeffectively than do electrons in the same subshell

b.

Nuclear chargei.

Pulls all the electrons closer to the nucleus. As the charge

increases, the atomic radius decreases

Eff ti N l Ch


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Effective Nuclear Charge1. The effective nuclear charge, Zeff, for a given electron is given

Where Z=number of protons (actual nuclear charge) and

S=number of electrons in the inner orbitals (screening effects)

Zeff =Z!

S

Element Z Inner Zeff Radiuspm

Na 11 10 1 156Mg 12 10 2 136

Al 13 10 3 125

Si 14 10 4 117

V i ti i t i di i d


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Variation in atomic radius across periods

4.

At the same time, more protons are added to the nucleuswhich increases the effective nuclear charge for the electronsin the valence shell.

5. This will cause the valence shell electrons to be drawn closerto the nucleus and thus decreases the size of the atomicradius.

1. There is a gradual decrease inatomic radius across a periodfrom left to right.

2.

When moving from left to right,the number of proton andelectron increases by one.

3.

Electron is added to the sameshell at about the samedistance from the nucleus.

Thus the electrons arerelatively ineffective atshielding each other and thescreening effect remainsconstant.

Variation in atomic radius across d block elements


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Variation in atomic radius across d-block elements

3. For d-block elements in the same row, additional electrons goto the inner subshell (3d orbital), the number of electrons in theoutermost subshell (4s orbital) remains constant. The inner 3d

electrons shield the 4s electrons from the nuclear charge moreeffectively than the outer electrons shield each other.4. Thus, increased in nuclear charge (proton numbers) is roughly

cancelled by the screening effect of the 3d electrons.5. Basically, the outermost electrons experience roughly similar

force of attraction going across the period, thus atomic radiusdoes not change very much for the first row.

1. Except Sc and Ti,the atomic radius ford-block elements are

approximatelyconstant across theperiod.

2. Decreasing atomicradius withincreasing proton

number does notapply for d-blockelements.

Variation in atomic radius descending a group


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Variation in atomic radius descending a group

1. On descending group, there is an increase in the atomicradius as the proton number increases. The higher theprincipal quantum number n of the valence electron, the largerthe atomic radius.

2.

The increase in atomic radius is due toa.

Increase in the number of electron shellsb.

Increase in the screening effect of the outer electrons by the innerelectrons as more completed shells are formed

3. The nuclear charge and the screening effect also increasesa.

Outer electrons enter new energy level

b.

Outer electrons are now screened by more electrons, thus decreasesthe effective nuclear charge, therefore atomic radius increases.

Comparison between the atomic radius of an elements and its ionic radius


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Comparison between the atomic radius of an elements and its ionic radius

1. For a given nuclear charge, the smaller number of electrons inan atom or ion, the smaller the repulsion between theelectrons, and the smaller the atomic or ionic size.

2.

The size of cations (+ve) are smaller than their neutral atoms

because they have smaller number of electrons.3. The size of anions (-ve) are larger than theirs neutral atomsbecause they have more electrons.

Example 3 8 Atomic Radius


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Example 3.8 Atomic Radius

34

Variation in the radius of isoelectronic species


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Variation in the radius of isoelectronic species1. Isoelectronic species are ions or molecules that have the same

number of electrons.2. For a given number of electrons, the higher the nuclear

charge, the higher the forces of attraction and the smaller theatomic or ionic radius.3. Table below shows the atomic or ionic radius of isoelectronic

elements with the same electronic configuration 1s22s22p6.

Example 3 9 Atomic Radius


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Example 3.9 Atomic Radius

36

Variation in ionic radius across Periods 2 and 3


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Variation in ionic radius across Periods 2 and 3

1. For Period 2 elementsa. The radii of +ve ions

decrease from Li

+

toBe3+

b. The radii increasefrom Be3+to N3-

c. The radii of ve ionsdecrease from N3-to

F-

2. For Period 3 elementsa. The radii of +ve ions decrease from Na+to Al3+

b. The radii increase from Al3+to P3-

c. The radii of ve ions decrease from P3-to Cl-

3.

The higher nuclear charge pulls the electron cloud closer to thenucleus, causing the ionic radius to decrease.4. The presence of additional electron shell increases the ionic

radii due to increase in screening effect.5. For a given nuclear charge, the larger the number of electrons

in an atom or ion, the greater the repulsion between electrons,and the larger the atomic or ionic radius.

Variation in ionic radius Descending a group


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Variation in ionic radius Descending a group

1. For ions of the same charge, the ionic size increases as wedescend a group because

a.

the number of energy levels increasesb.

the valence electrons are farther from the nucleusc. the valence electrons become more effectively shielded

from the nucleus by the increasing number of electrons

2. Both factors, increasing distance from the nucleus and theshielding effect, outweigh the effect of increasing nuclearcharge. As a result, the effective nuclear charge for theelectrons in the valence shell decrease.

The first and Second Ionization Energies


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The first and Second Ionization Energies1. The first ionization energy is the minimum energy required to

remove one mole of electrons from 1 mole of atoms in thegaseous state. The "H1is

2. The second first ionization energy is the minimum energyrequired to remove one mole of electrons from 1 mole of uni-positive ions in the gaseous state. The "H2is

3. The first and second ionization energies of iron are +762 kJ/mol and 1560 kJ/mol.

4. The third and subsequent ionization energies can be definedusing the same method.

5.

The lower the ionization energy of an element, the more easilyits electron can be removed. As the ionization energies

decrease, the metals become more metallic.

M(g)!M+

(g)+ e"

M+

(g)!M2+

(g)+ e"

Fe(g)!

Fe+

(g)+ e"

Fe+

(g)!

Fe2+

(g)+ e"

Factors that affect Ionization Energies


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Factors that affect Ionization Energies1. The level of difficulty in removing electron from the atom or

from the ion depends on the strength of the attractive forcesbetween the electron and the nucleus.

2.

Hence, the ionization energy is influenced bya.

Distance of the outer electron from the nucleusb.

Size of the nuclear chargec.

Screening effect of the electrons in the inner shells

3. Atomic radiusa.

The attraction of the positive nucleus for the negatively charge electron

decreases as the distance increases and this causes ionization energydecreases. Therefore, the ionization energy decreases as the atomicradius increases.

4.

Nuclear chargea.

When nuclear charge becomes more positive, its attraction on the outershell electrons increases. This causes the ionization energy increases.

b.

Mg (1s2

2s2

2p6

3s2

) has a nuclear charge of +12, and Na (1s2

2s2

2p6

3s1) has a nuclear charge of +11. Both have the same number ofelectron in the inner shell, Mg has higher effective nuclear charge.Then the ionization energy of Mg is higher than Na.

5. Screening effect

Factors that affect Ionization Energies


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Factors that affect Ionization Energies5. Screening effect

a.

The valence electrons are shielded from the attraction of the nucleus bythe screening effect (effect of repulsion) of the electrons in the innershells.

b.

The screening effect by the inner electrons is more effective if theseelectrons are closer to the nucleus. Because of thisi.

Electron in the lower principle quantum number have a strongershielding effect compared with electrons in shell of higher principlequantum number.

ii.

Electrons in the same shell exert a very small screening effect on

each other.6.

Ionization energies always increase because7. In the case of Al (1s22s22p63s23p1), there is a large increase

between the first and the second ionization energies becausethe electron in removed not only from a positive ions Al+, butalso from a filled 3s orbital which is more stable.

8.

The fourth ionization energy of Al is very high because toremove the fourth electron requires breaking into inner shellwhich has stable noble gas electronic configuration (1s22s22p6).

Variation in first ionization energy across a Period


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Variation in first ionization energy across a Period1.

With minor exceptions, whenmoving across a period fromleft to right, the ionizationenergy increases and theelements become lessmetallic and more nonmetallic.

2. When moving across aperiod from left to right, thenuclear charge increase, theatomic radius decreases, butthe screening effect remainsalmost constant.

3. When the effective nuclear charge increases, and theoutermost shell electrons are more tightly held by the nucleus,the ionization energy increases.

Variation in first ionization energy involving anomalous behavior between group


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gy g g p

2 and 13

1. Between Be and B, B hassmaller atomic radius,therefore we would expectthat the ionization energy ofB would be higher than Be.

2. The ionization energy of B islower than Be because Be(1s22s2) loses a 2s electronwhere as B (1s22s2 2p1)loses a 2p electron.

3. More energy is required toremove an electron from thelower energy 2s orbital thanfrom the 2p orbital.

4.

The electrons in the filled 2s orbital are more effective atshielding the electron in the 2p orbital than they are atshielding each other. Thus, the single electron in the 2psubshell are better shielded than the 2s2electrons. Therefore,less energy is needed to remove a single 2p electron than apaired of 2s electron.

5.

The same apply between Mg and Al.

Variation in first ionization energy involving anomalous behavior between group


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15 and 16

1. The first ionization energy ofO (1s22s2 2px2 2py1 2pz1) islower than N (1s22s2 2px1

2py1 2pz1) , even though bothhas same energy level of 2porbital electrons, because Nhas a half filled electronicconfiguration which is morestable.

2.

Moreover, O has twoelectrons in the same 2pxorbital which result in greaterrepulsion effects, and thismake it easier to remove theelectron.

3.

Therefore, the decrease in the first ionization energy on goingfrom N to O is due to the repulsion of paired electrons in the2p4 configuration of the O atoms.

4. The same apply between P and S.

Variation in first ionization energy on descending a group


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Variation in first ionization energy on descending a group

1.

On descending a group, the atomic radius increases as moreelectrons are added to successive energy levels and thiscauses the screening effect to increase.

2.

The further the outermost electron is from the nucleus, thesmaller is the attraction force between the nucleus and theelectron, and the more easily the electron can be removed.

Example 3 10


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Example 3.10

46

Successive ionization energy, electronic configuration and the position of an


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element in the Periodic Table

1. Successive ionizationincrease because thepositive charge on the ionproduced increases as theeach electron is removed,making successive electronsincreasingly difficult toremove.

2. The study of successiveionization energies ofelements proves theexistence of energy levels inan atom.

3.

Consider the log plot of successive ionization energy of Na.

There is relatively large increase in ionization energy when thesecond and the tenth electron is removed. This suggests thatthe nucleus of the Na atom is surrounded by electrons whichare group into three energy levels, called shells.

4.

The plot also suggests that the 2ndelectron and the 10thisnearer to the nucleus than the 1stand the 9thelectron.

1s22s22p63s1

Example 3 11


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Example 3.11

48

Electronegativity


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Electronegativity

49

1. The electronegativity of an atom is the ability or power of theatom in a covalent bond to attract pairs of shared electrons to

itself.2. The electronegativity is measured using Paulings scale andvalue 4 is the highest electronegativity.

3. Consider the covalent molecule HCla.

Cl is more electronegative than H

b.

Cl attract shared electrons away from H to itselfc.

The H atom in the HCL acquires a partial +ve charge and the Cl atomacquires a partial ve charge The HCL molecule is represented as H#+-Cl#-

4. The greater the electronegativity of an atom, the greater its

ability to attract electron to itself.5. The electronegativity of an atom is different from its electronaffinity. Electronegativity measures the ability of an atom in amolecule to attract pairs of shared electrons, whereas electronaffinity measures the ability of a single gaseous atom to gain

an electron to form a negative ion.

Electronegativity


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Electronegativity

50

6. The electronegativity of an element is related to its electronaffinity and its ionization energy. A highly electronegative atom

will have aa.

Very negative first electron affinity

b.

High ionization energy

Because it will attract electrons from other atoms but resist

having its electrons pulled away.

Variation in electronegativity across a period and descending a group


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g y p g g p

51

1. Electronegativity increases across a period from left to right,that is from metallic to non metallic.

2. Electronegativity decreases down a group.3. However, there are some exception

4.

The electronegativity of H to Ca above shows that the electronegativities and ionization energy variations has the sametrend because both values depend on nuclear charge and theatomic radius.

The periodicity of the physical properties of elements across


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The periodicity of the physical properties of elements acrossperiod 3 and down group 1 and 17

52

1.

The more an element exhibits physical properties of metals,the greater is its metallic character. In general, the metallic

charactera.

Decrease across a periodb.

Increase down a group

2. The metallic character of elements can be compared in termsof first ionization energiesa.

The first ionization energy increases across a period and decreasesdown a group

b.

Metals (group 1) have low ionization energiesc.

Non metals (group 17) have high ionization energy

The periodicity of the physical properties of elements across


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p y p y p pperiod 3 and down group 1 and 17

53

3. On descending agroup, the metallic

character increasesand the first ionizationenergy decreases. Ondescending, group 1reactivity with O2andH2O increases. Themore reactive a metal,the greater the metalliccharacter.

4. For group 17, ondescending a group,the non metallic

character decreasesand the metalliccharacter increases.

5. The trend of increasingmetallic character isclearer for group 14

and 15.

Variation in melting point and boiling point across period 3


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1. For period 3, the melting and boiling pointa.

Increase from Na metal to the giant molecular Si, and thenb.

Decrease sharply from Si to the simple non metallic molecule P

is caused by the changes in the structure and bonding of theelements across the period.

Metallic Bond


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55

Metallic Bond

1.

In the metallic structure, each

positively charged metal ion is

attracted to a cloud of negative

electron which is responsible for

the bonding in metals.

2.

The melting point of Al (1s22s2 2p6

3s23p1) is only slightly higher than

Mg (1s2

2s2

2p6

3s2

), implying,that Al atom does not use all the

three valence electrons for metallic

bonding.

3. The boiling points of metals are very much higher than theirmelting pointa. This implies that most of the metallic bonds still exist in the

liquid state

b. However, when the liquid changes into vapor, the atomsmust be separated to a considerable distance and thisinvolves breaking all the metallic bonds.

Giant Covalent and Simple Molecular Structure


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56

Giant Covalent and Simple Molecular Structure

1. Si has very high melting and boiling pointbecause it has strong covalent bond in the

giant crystal lattice.2.

Nearly all the bonds must be broken before

the solid melts. As a result, the boiling points

are not very much higher than their melting

points.3.

All the non metallic elements (P4, S8and Cl2) in period 3 formsimple molecular structure and consists of small and discretemolecules.

4. The covalent bonds within the molecule are very strong but theVan der Waals forces of attraction between the molecules are

very weak, thus the relatively low melting point.5.

The Van der Waals forces of attraction increases as themolecular size increases, that is as the relative molecularmass increases.

6.

The boiling of elements with simple molecular structures onlyinvolves overcoming the weak Van der Waals forces. Thisexplain why the boiling takes place at low temperature and notthat much higher than the melting point.

Variation in melting and boiling point descending a group


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57

Variation in melting and boiling point descending a group

1. Group 1 (alkali metals)have low melting andboiling points compared to

other metals such as Fe.2.

Descending a group, themelting and boiling pointdecrease because theattractive forces betweenthe atoms becomes

weaker as the atomicradius becomes larger.

3. Group 17 (halogen) consists of small molecules. The forces ofattractions are weak and so halogens have low melting andboiling points.

4. Descending a group, the melting and boiling points increases.

This is because as the molecules get larger, the Van der Waalsforces between the neighboring molecules increase.5. Hence, Fl2and Cl2are gases, Br2is a liquid and I2is a solid at

room temperature.

The acid-base character of the oxides of period 3 elements


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p

58

1.

Metals form basic oxides whereas non metals form acidicoxides.

2. Metal oxides are ionic compound, whereas non metal oxidesare molecular covalent compound.

3. Si is metalloid. Its oxide SiO2 is a giant covalent molecule.4. Across period 3, the metallic character decreases as the

ionization energy increases. As a result, the metal oxidesbecome more acidic.

Basic Oxides and Acidic Oxides


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59

1. Metal oxides react with acids to form salt and water only andare called basic oxides and they are ionic solids with highmelting and boiling points.

2. Na2O in an alkali because it dissolves readily in water to formalkaline solution.

3. Non metals burn in O2 to form acidic oxides. Acidic oxides aresimple covalent molecules which exist as gases or solids(P4O10) with low melting points.

4. Acidic oxides dissolve in water to form acids.

5. SiO2is an acidic oxide. It is insoluble in water but dissolves inhot concentrated NaOH to form salt and water.

Na2O s( )+2HCl aq( )! 2NaCl aq( )+H2O l( )

Na2

O s( )+H2O l( )! 2NaOH aq( )

SiO2 s( )+ 2NaOH aq( )! Na2SiO3 aq( )+H2O l( )

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

Amphoteric Oxides


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60

1. Amphoteric oxides acts as both acidic and basic oxidesdepending on the conditions. They behave as acids when theyreact with bases and behave as bases when they react with

acids.2.

For example, Al2O3reacts with HCl to form AlCl3

3. Al2O3dissolve in NaOH solution to form NaAlO2

Al2O

3 s( )+ 6HCl aq( )! 2AlCl3 aq( )+3H2O l( )

3Al2O

3 s( )+ 6NaOH aq( )! 2NaAlO2 aq( )+3H2O l( )

Quick Check 3.4


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Chemistry Topic 4 Matriculation

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