chapter 8 periodic properties & electron configurations bushra javed

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Chapter 8Periodic Properties & Electron Configurations

Bushra Javed

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Contents

• Electron Spin• Electron Configurations of elements• The development of Periodic Table• Periodic Trends

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Electron Configurations

• Quantum-mechanical theory describes the behavior of electrons in atoms

• The electrons in atoms exist in orbitals• A description of the orbitals occupied by

electrons is called an electron configuration

1s1principal energy level of orbital occupied by the

electron sublevel of orbital occupied by the electron

number of electrons in the orbital

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How Electrons Occupy Orbitals

• Calculations with Schrödinger’s equation show hydrogen’s one electron occupies the lowest energy orbital in the atom

• Schrödinger’s equation calculations for multi-electron atoms cannot be exactly solved due to electron-electron interactions in multi-electron atoms

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How Electrons Occupy Orbitals

• The interactions that occur in multi-electron atoms are due to :

• electron spin • & energy splitting of sublevels

The Property of Electron Spin

• Spin is a fundamental property of all electrons• All electrons have the same amount of spin• The orientation of the electron spin is quantized,

it can only be in one direction or its opposite– spin up or spin down

• The electron’s spin adds a fourth quantum number to the description of electrons in an atom, called the Spin Quantum Number, ms

– not in the Schrödinger equation

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The Property of Electron Spin• Experiments by Stern and Gerlach showed a

beam of silver atoms is split in two by a magnetic field.

• The experiment reveals that the electrons spin on their axis

• spinning charged particles generate a magnetic field

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Electron Spin

• If there is an even number of electrons, about half the atoms will have a net magnetic field pointing “north” and the other half will have a net magnetic field pointing “south”

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Electron Spin

The two possible spin orientations of an electron and the conventions for ms are illustrated here.

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ms, and Orbital Diagrams• ms can have values of +½ or −½• Orbital Diagrams use a square to represent each

orbital and a half-arrow to represent each electron in the orbital

• By convention, a half-arrow pointing up is used to represent an electron in an orbital with spin up

• Spins must cancel in an orbital– paired

Orbital Diagrams• We often represent an orbital as a square or a

circle and the electrons in that orbital as arrows– the direction of the arrow represents the spin of the

electron

orbital withone electron

unoccupiedorbital

orbital withtwo electrons

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Pauli Exclusion Principle• No two electrons in an atom may have the same set

of four quantum numbers

• Therefore no orbital may have more than two electrons, and they must have the opposite spins

• Knowing the number orbitals in a sublevel allows us to determine the maximum number of electrons in the sublevel

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Pauli Exclusion Principles sublevel has 1 orbital, therefore it can hold 2 electrons

p sublevel has 3 orbitals, therefore it can hold 6 electrons

d sublevel has 5 orbitals, therefore it can hold 10 electrons

f sublevel has 7 orbitals, therefore it can hold 14 electrons

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Sublevel Splitting in Multi-electron Atoms• The sublevels in each principal energy shell of

Hydrogen all have the same energy • We call orbitals with the same energy degenerate• For multi-electron atoms, the energies of the sublevels

are split– caused by charge interaction, shielding and penetration

• The lower the value of the l quantum number, the less energy the sublevel has

– s (l = 0) < p (l = 1) < d (l = 2) < f (l = 3)

Shielding & Effective Nuclear Charge• Each electron in a multi-electron atom experiences

both the attraction to the nucleus and repulsion by other electrons in the atom

• These repulsions cause the electron to have a net reduced attraction to the nucleus – it is shielded from the nucleus

• The total amount of attraction that an electron feels for the nucleus is called the effective nuclear charge of the electron

Shielding & Penetration

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penetration & Effective Nuclear Charge• The closer an electron is to the nucleus, the

more attraction it experiences

• The better an outer electron is at penetrating through the electron cloud of inner electrons, the more attraction it will have for the nucleus

• Penetration causes the energies of sublevels in the same principal level to not be degenerate

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Effect of Penetration and Shielding• In the fourth and fifth principal levels, the effects of

penetration become so important that the s orbital lies lower in energy than the d orbitals of the previous principal level

• The energy separations between one set of orbitals and the next become smaller beyond the 4s

– the ordering can therefore vary among elements– causing variations in the electron configurations of

the transition metals and their ions

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Filling the Orbitals with Electrons• Energy levels and sublevels fill from lowest energy

to highs → p → d → fAufbau Principle

• Orbitals that are in the same sublevel have the same energy

• No more than two electrons per orbitalPauli Exclusion Principle

• When filling orbitals that have the same energy, place one electron in each before completing pairsHund’s Rule

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Filling the Orbitals with ElectronsThe lowest-energy configuration of an atom is called its ground state.

Any other allowed configuration represents an excited state.

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Electron Configuration of Atoms • The electron configuration is a listing of the sublevels

in order of filling with the number of electrons in that sublevel written as a superscript.Kr = 36 electrons = 1s22s22p63s23p64s23d104p6

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‘Building up’ order of Filling in sublevels (Ground State Electron Configurations)

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d

7s

Start by drawing a diagramputting each energy shell ona row and listing the sublevels, (s, p, d, f), for that shell in order of energy (left-to-right)

Next, draw arrows throughthe diagonals, looping back to the next diagonaleach time

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Electron Configuration & the Periodic Table

• The Group number corresponds to the number of valence electrons

• The length of each “block” is the maximum number of electrons the sublevel can hold

• The Period number corresponds to the principal energy level of the valence electrons

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s1

s2

d1 d2 d3 d4 d5 d6 d7 d8 d9 d10

s2p1 p2 p3 p4 p5

p6

f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 f14d1

1234567

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Blocks of Elements

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Blocks of Elements

For main-group (representative) elements, an s or a p subshell is being filled.

For d-block transition elements, a d subshell is being filled.

For f-block transition elements, an f subshell is being filled.

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Electron ConfigurationsExample 1Write the ground state electron configuration of the chlorine atom, Cl,

1s2 2s2 2p6 3s2 3p5

For chlorine, Cl, Z = 17.

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Example 2Write the ground state electron configuration of the manganese atom, Mn

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Example 3: What is the ground-state electron configuration of tantalum (Ta)?

a. 1s22s22p63s23p64s23d104p65s24d105p6 6s25d104f7

b. 1s22s22p63s23p64s23d104p65s24d105p6 6s24f145d3

c. 1s22s22p63s23p64s23d104p65s24d105p6 6s2 5d3

d. 1s22s22p63s23p64s23d104p65s24d105p6 6s24f14

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Electron ConfigurationsExample 4Which of the following electron configurations corresponds to the ground state electron configuration of an atom of a transition element?

a. 1s22s22p2

b. 1s22s22p63s23p5

c. 1s22s22p63s23p64s2

d. 1s22s22p63s23p63d54s2

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Exceptions to the ‘Building up’ order of Filling

• About 21 of the predicted configurations are inconsistent with the actual configurations observed.

One Possible Reason:• Half filled and fully filled subshells are highly Stable

• Cr, Cu ,Ag and U are some of the exceptions.


Chromium (Z=24) and copper (Z=29) have been found by experiment to have the following ground-state electron configurations:

Cr: 1s2 2s2 2p6 3s2 3p6 3d5 4s1

Cu: 1s2 2s2 2p6 3s2 3p6 3d10 4s1

In each case, the difference is in the 3d and 4s subshells.

8 | 33

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• Expected• Cr = [Ar]4s23d4

• Cu = [Ar]4s23d9

• Mo = [Kr]5s24d4

• Ru = [Kr]5s24d6

• Pd = [Kr]5s24d8

• Found Experimentally• Cr = [Ar]4s13d5

• Cu = [Ar]4s13d10

• Mo = [Kr]5s14d5

• Ru = [Kr]5s14d7

• Pd = [Kr]5s04d10

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Example 5Which of the following electron configurations represents an allowed excited state of the indicated atom?

a. He: 1s2

b. Ne: 1s2 2s2 2p6

c. Na: 1s2 2s2 2p6 3s2 3p2 4s1

d. P: 1s2 2s2 2p6 3s2 3p2 4s1

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The Noble Gas coreElectron Configuration

• The noble gases have eight valence electrons.– except for He, which has only two

electrons• We know the noble gases are

especially non-reactive– He and Ne are practically inert

• The reason the noble gases are so non-reactive is that the electron configuration of the noble gases is especially stable

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Noble Gas Core Electron Configuration

• A short-hand way of writing an electron configuration is to use the symbol of the previous noble gas in [] to represent all the inner electrons, then just write the last set.

Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1 = [Kr]5s1

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Noble Gas Core Electron ConfigurationExample 6Which ground-state electron configuration is incorrect?

a. Fe: [Ar] 3d5

b. Ca: [Ar] 4s2

c. Mg: [Ne] 3s2

d. Zn: [Ar] 3d10 4s2

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Writing Electron Configurations

The pseudo-noble-gas core includes the noble-gas subshells and the filled inner, (n – 1), d subshell.

For bromine, the pseudo-noble-gas core is[Ar]3d10

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Hund’s rule

Hund’s rule states:

that the lowest-energy arrangement of electrons in a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin before pairing electrons

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Example 7Draw an orbital diagram for nitrogen.

1s 2s 2p

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Electron ConfigurationsExample 8Which of the following electron configurations or orbital diagrams are allowed and which are not allowed by the Pauli exclusion principle? If they are not allowed, explain why?

a.1s22s12p3

b.1s22s12p8

c. 1s22s22p63s23p63d8 d.1s22s22p63s23p63d11

e. 1s 2s

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Electron ConfigurationsExample 9write the complete ground state orbital diagram and electron configuration of potassium

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Valence Electrons• The electrons in all the sublevels with the

highest principal energy shell are called the valence electrons

• Electrons in lower energy shells are called core electrons

• Chemists have observed that one of the most important factors in the way an atom behaves, both chemically and physically, is the number of valence electrons

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Valence- shell configurationFor main-group elements, the valence configuration is in the form

nsAnpB

The sum of A and B is equal to the group number.

So, for an element in Group VA of the third period, the valence configuration is

3s23p3

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Valence-shell configurationExample 10What are the valence-shell configuration for arsenic and zinc?

Arsenic is in period 4, Group VA.Its valence configuration is 4s24p3.

Zinc, Z = 30, is a transition metal in the first transition series. Its noble-gas core is Ar, Z = 18.Its valence configuration is 4s23d10.

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Valence-shell configurationExample: 11What is the valence -shell configuration of Tc, Z = 43

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Mendeleev’s Periodic Table

• Order elements by atomic mass• Saw a repeating pattern of properties • Periodic Law – when the elements are

arranged in order of increasing atomic mass, certain sets of properties recur periodically

• Put elements with similar properties in the same column

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Mendeleev’s Periodic Table

• Used pattern to predict properties of undiscovered elements

• Mendeleev’s Periodic Law allows us to predict what the properties of an element will be based on its position on the table

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Mendeleev's Predictions

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Everyone Wants to Be Like a Noble Gas! The Alkali Metals

• The alkali metals have one more electron than the previous noble gas

• In their reactions, the alkali metals tend to lose one electron, resulting in the same electron configuration as a noble gas– forming a cation with a 1+ charge

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Periodic Trends

Periodic law states that when the elements are arranged by atomic number, their physical and chemical properties vary periodically.

We will look in more detail at three periodic properties:

1. atomic radius, 2. ionization energy, and 3. electron affinity.

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Atomic RadiusWhile an atom does not have a definite size, we can define it in terms of covalent radii (the radius in covalent compounds).

Trend in Atomic Radius – Main Group

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• There are several methods for measuring the radius of an atom, and they give slightly different numbers

• van der Waals radius = nonbonding• covalent radius = bonding radius

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• Atomic Radius Increases down groupvalence shell farther from nucleuseffective nuclear charge fairly close

• Atomic Radius Decreases across period (left to right)adding electrons to same valence shelleffective nuclear charge increasesvalence shell held closer

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TrendsWithin each group (vertical column), the atomic radius increases with the period number.

This trend is explained by the fact that each successive shell is larger than the previous shell.

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Within each period (horizontal row), the atomic radius tends to decrease with increasing atomic number (nuclear charge).

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A representation of atomic radii is shown below.

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Example 12:Refer to a periodic table and arrange the following elements in order of increasing atomic radius: Br, Se, Te.

35Br

34Se

52Te

Te is larger than Se.Se is larger than Br.

Br < Se < Te

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Example 13An atom of which of the following elements has the largest atomic radius?

a) Rbb)Clc) Mgd) P

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Example 14An atom of which of the following elements has the smallest atomic radius?

a) Mgb) Clc) Asd) Rb

Trend in Ionization Energy

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• Minimum energy needed to remove an electron from an atom or ion in the gaseous state• endothermic process• valence electron easiest to remove, lowest IE

• M(g) + IE1 M1+(g) + 1 e- • M+1(g) + IE2 M2+(g) + 1 e-

first ionization energy = energy to remove electron from neutral atom; 2nd IE = energy to remove from 1+ ion; etc.

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• Going down a group, first ionization energy decreases.

• This trend is explained by understanding that the smaller an atom, the harder it is to remove an electron, so the larger the ionization energy.

• In general : E1< IE1< IE3

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Trend in Ionization EnergyGenerally, ionization energy increases with atomic number.

Ionization energy is proportional to the effective nuclear charge divided by the average distance between the electron and the nucleus.

Because the distance between the electron and the nucleus is inversely proportional to the effective nuclear charge, ionization energy is inversely proportional to the square of the effective nuclear charge.


Example 15Refer to a periodic table and arrange the following elements in order of increasing ionization energy: As, Br, Sb.

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Sb is larger than As.As is larger than Br.

Ionization energies: Sb < As < Br

35Br

33As51Sb

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Small deviations occur between Groups IIA and IIIA and between Groups VA and VIA.

Examining the valence configurations for these groups helps us to understand these deviations:

IIA ns2 IIIA ns2np1

VA ns2np3

VIA ns2np4

It takes less energy to remove the np1 electron than the ns2 electron.

It takes less energy to remove the np4 electron than the np3 electron.

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Example16An atom of which of the following elements has the largest second ionization energy?

a) Lib) Cc) Fd) Be

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Example 17Which of the following elements has larger firstIE and WHY ?

a) Alb) Nc) Od) K

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Electron affinity (E.A.)

The energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion.

A negative energy change (exothermic) indicates a stable anion is formed. The larger the negative number, the more stable the anion. Small negative energies indicate a less stable anion.

A positive energy change (endothermic) indicates the anion is unstable.

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Trends in Electron affinity (E.A.)

Electron affinities in the main-group elements show a periodic variation when plotted against atomic number, although this variation is somewhat more complicated than that displayed by ionization energies.

In a given period, the electron affinity rises from the Group IA element to the Group VIIA element but with sharp drops in the Group IIA and Group VA elements.

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Trends in Electron affinity (E.A.)

Broadly speaking, the trend is toward more negative electron affinities going from left to right in a period.

Let’s explore the periodic table by group.

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Trend in Electron Affinity (EA)

Example 18An atom of which of the following elements has themost negative electron affinity?

a) Kb) Clc) Brd) See) N

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Groups IIA and VIIIA do not form stable anions; their electron affinities are positive.

Group Valence Anion ValenceIA ns1 ns2 stableIIIA ns2np1 ns2np2 stableIVA ns2np2 ns2np3 stableVA ns2np3 ns2np4 not so stable

VIA ns2np4 ns2np5 very stableVIIA ns2np5 ns2np6 very stable

Except for the members of Group VA, these values become increasingly negative with group number.

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Metallic CharacterElements with low ionization energies tend to be metals. Those with high ionization energies tend to be nonmetals. This can vary within a group as well as within a period.

Metallic Character• Metallic character is how closely an element’s

properties match the ideal properties of a metal– more malleable and ductile, better conductors, and

easier to ionize• Metallic character decreases left-to-right across

a period– metals are found at the left of the period and

nonmetals are to the right• Metallic character increases down the column

– nonmetals are found at the top of the middle Main Group elements and metals are found at the bottom

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Oxides of Main Group ElementsOxidesA basic oxide reacts with acids. Most metal oxides are basic. If soluble, their water solutions are basic.

An acidic oxide reacts with bases. Most nonmetal oxides are acidic. If soluble, their water solutions are acidic.

An amphoteric oxide reacts with both acids and bases.

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Trends in Metallic Character

Group IA, Alkali Metals (ns1)These elements are metals; their reactivity increases down the group.

The oxides have the formula M2O.

Hydrogen is a special case. It usually behaves as a nonmetal, but at very high pressures it can exhibit metallic properties.

Reactions of Alkali Metals with Water• Alkali metals are oxidized to the 1+ ion• H2O is split into H2(g) and OH− ion• The Li, Na, and K are less dense than the water

– so float on top• The ions then attach together by ionic bonds• The reaction is exothermic, and often the

heat released ignites the H2(g)

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Group IIA, Alkaline Earth Metals (ns2)These elements are metals; their reactivity increases down the group.

The oxides have the formula MO.

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Group IIIA (ns2np1)Boron is a metalloid; all other members of Group IIIA are metals.

The oxide formula is R2O3.

B2O3 is acidic.Al2O3 and Ga2O3 are amphotericThe Group IIIA oxides are basic.

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Group IVA (ns2np2)Carbon is a nonmetal; silicon and germanium are metalloids; tin and lead are metals.

The oxide formula is RO2 and, for carbon and lead, RO.

CO2, SiO2, and GeO2 are acidic (decreasingly so).SnO2 and PbO2 are amphoteric.

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PbO (yellow) PbO2

(dark brown)

SiO2 (crystalline solid quartz)

SnO2 (white)

Some oxides of Group IVA

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Group VA (ns2np3)

Nitrogen and phosphorus are nonmetals; arsenic and antimony are metalloids; bismuth is a metal.

The oxide formulas are R2O3 and R2O5, with some molecular formulas being double these.

Nitrogen, phosphorus, and arsenic oxides are acidic.Antimony oxides are amphoteric.Bismuth oxide is basic.

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Group VIA, Chalcogens (ns2np4)

Oxygen, sulfur, and selenium are nonmetals; tellurium is a metalloid; polonium is a metal.

The oxide formulas are RO2 and RO3.

Sulfur, selenium, and tellurium oxides are acidic except for TeO2, which is amphoteric. PoO2 is also amphoteric.

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Group VIIA, Halogens (ns2np5)

These elements are reactive nonmetals, with the general molecular formula being X2. All isotopes of astatine are radioactive with short half-lives. This element might be expected to be a metalloid.

Each halogen forms several acidic oxides that are generally unstable.

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Group VIIIA, Noble Gases (ns2np6)

These elements are generally unreactive, with only the heavier elements forming unstable compounds. They exist as gaseous atoms.

chapter 8 periodic properties & electron configurations bushra javed

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