CHM092 2 Matter Quan Config Periodicity

2-1

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Chapter 2

Electron Configuration

Chemical Periodicity

The Components of Matter

Quantum numbers


2-2

The Components of Matter, Quantum Numbers,

Electron Configuration,Chemical Periodicity

2.1 Elements and compounds

2.2 Atomic structure

2.3 Atomic number, mass number and isotopes

2.4 Bohrs atomic model

2.5 Quantum numbers

2.6 Electron configuration

2.7 The periodic table

2.8 Periodic Trend: atomic and ionic radius, ionization energy and electron affinity


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Molecular View of

Elements and Compounds


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2.1 Components of Matter

Element - the simplest type of substance with unique physical and

chemical properties. An element consists of only one type of atom. It

cannot be broken down into any simpler substances by physical or

chemical means. Examples Na, Mg , Zn , S, C , Cl2 , H2 . Can be

divided into atomic element and molecular element.


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Compound - a substance composed of two or

more elements which are chemically combined in

fixed or constant proportions by mass.

It has a definite chemical formula. Can be

classified into molecular compound (chemically link

by covalent bonds) and ionic compounds.

Examples: H2O, NH3, H2SO4 , K2O and NaCl.

Mixture - a group of two or more

elements and/or compounds that

are physically intermingled.

Constituents in the mixture can be

separated by physical means.


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Allowed to react chemically

therefore cannot be separated

by physical means.

Figure 2.1 The distinction between mixtures and compounds.

S

Fe

Physically mixed therefore can

be separated by physical

means; in this case by a

magnet.


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Molecule - a structure that consists of two or

more atoms of the same or different element that

are chemically bound together by covalent bonds

and thus behaves as an independent unit.

Elements like H2, O2, N2 and the halogens e.g.

F2 occur as diatomic molecules and are refer to

as molecular elements

Other elements occur as polyatomic molecules:

O3, P4, S8, Se8

HCl , CO2 and CH4 are molecular compounds.

Atom - the smallest particle of an element that gives the characteristic

properties of that element. It is also the smallest particle of an element that

can combine with itself or with other atoms in a chemical reaction.

Examples: Iron (Fe), sodium (Na), aluminium (Al), sulphur (S),

oxygen (O) and phosphorus (P).


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Elements that occur as molecules.

1A 2A 3A 4A 5A 6A 7A 8A

(1) (2) (13) (14) (15) (16) (17) (18)

H2

N2 O2 F2

P4 S8 Cl2

Se8 Br2

I2

diatomic molecules

tetratomic molecules

octatomic molecules

P4

S8


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Properties of some elements and an ionic compound

2.1


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2.2 The law of mass conservation:

Mass remains constant during a chemical reaction.


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The total mass of substances does not change during a

chemical reaction.

reactant 1 + reactant 2

total mass

product

total mass =

calcium oxide + carbon dioxide calcium carbonate

CaO + CO2 CaCO

3

56.08g + 44.00g 100.08g


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No matter the source, a particular compound is composed of

the same elements in the same parts (fractions) by mass.

Example : Calcium carbonate, CaCO3

Analysis by Mass

(grams/20.0g)

Mass Fraction

(parts/1.00 part)

Percent by Mass

(parts/100 parts)

8.0 g calcium

2.4 g carbon

9.6 g oxygen

20.0 g

40% calcium

12% carbon

48% oxygen

100% by mass

0.40 calcium

0.12 carbon

0.48 oxygen

1.00 part by mass

2.3 Law of Definite (or Constant) Composition:


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2.4 Structure of Atom ---- General features of the atom today.

The atom is an electrically neutral, spherical entity composes of a positively charged central nucleus surrounded by one or more

negatively charge electrons.

The atomic nucleus consists of protons and neutrons.


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Table 2.2 Properties of the Three Key Subatomic

Particles

Charge Mass

Relative

1+

0

1-

Absolute(C)*

+1.60218x10-19

0

-1.60218x10-19

Relative(amu)

1.00727

1.00866

0.00054858

Absolute(g)

1.67262x10-24

1.67493x10-24

9.10939x10-28

Location

in the Atom

Nucleus

Outside

Nucleus

Nucleus

Name(Symbol)

Electron (e-)

Neutron (n0)

Proton (p+)

* The coulomb (C) is the SI unit of charge.

The atomic mass unit (amu) equals 1.66054x10-24 g.


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Atomic number (Z) -- the total number of protons (p+) in the nucleus of

each atom of an element. Mass number (A) -- the sum of all the protons (p+) and neutrons (no)

present in the nucleus of an atom.

Mass number (A) = number of protons + number of neutrons

= atomic number (Z) + number of neutrons

Number of neutrons = A Z

Nuclear charge -- the total positive charge contributed by all the

protons in the nucleus of an atom. For example, a sodium atom has a nuclear charge of +11 because it

has 11 protons in its nucleus.

Atomic symbol a definite symbol for every element, sometimes known as element symbol. Example, carbon (C),

Magnesium (Mg)

Definations


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2.5 Symbolic representation of an

atom, isotopic or nuclear symbol

X = Atomic symbol of the element

A = mass number; A = Z + n

Isotope = atoms of an element with the same number

of protons, but different number of neutrons.

e.g 1H, 2H and 3H. 12C, 13C and 14C.

Since chemical properties of an element are primarily

determined by the number of electrons, so all isotopes of

an element have nearly identical chemical behavior, but

different physical properties which involve masses

Z = atomic number

n = number of neutrons in the nucleus

A

Z X The Symbol of the Atom or Isotope


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Sample Problem 2. 1 Determining the Number of Subatomic

Particles in the Isotopes of an Element

PROBLEM: Silicon(Si) is essential to the computer industry as a major component of semiconductor chips. It has three naturally

occurring isotopes: 28Si, 29Si, and 30Si. Determine the number

14

of protons, neutrons, and electrons in each silicon isotope.

PLAN: We have to use the atomic number (14) and atomic masses

(28,29,30)

SOLUTION: The atomic number of silicon is 14. Therefore

28Si has 14p+, 14e- and 14n0 (28-14)

29Si has 14p+, 14e- and 15n0 (29-14)

30Si has 14p+, 14e- and 16n0 (30-14)


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Sample Problem 2.2 Calculating the Atomic Mass of an Element

PLAN: We have to find the weighted average of the isotopic masses, so we multiply each isotopic mass by its fractional abundance

and then sum those isotopic portions.

PROBLEM: Silver(Ag: Z = 47) has 46 known isotopes, but only two occur naturally, 107Ag and 109Ag. Given the following mass

spectrometric data, calculate the atomic mass of Ag:

Isotope Mass(amu) Abundance(%) 107Ag

109Ag

106.90509

108.90476

51.84

48.16

SOLUTION:

mass portion from 107Ag =

106.90509amu x 0.5184 = 55.42amu

mass portion from 109Ag = 108.90476amu x 0.4816 = 52.45amu

atomic mass of Ag = 55.42amu + 52.45amu = 107.87amu

mass(g) of each

isotope

portion of atomic mass

from each isotope atomic mass

Ga-69 = 60.11%, 68.9256 amu

Ga-71 = 39.89%, 70.9247 amu

atomic mass, amu

isotope masses,

isotope fractions

avg. atomic mass

Sample Problem 2.3

Given the following mass spectrometric data for Ga,

calculate the atomic mass of Ga in amu


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Formation of Ions

an atom is electrically neutral because

number of protons number of electrons

( positive charges) ( negative charges)

Atoms do not lose or gain protons in any chemical reactions but

lose or gain electrons and become ions.

releases electron/electrons positive ion (cation) no. of electrons in an atom > no. of electrons in its ion

Charge = no. of electrons in atom no. of electrons in its ion

Eg: Al 3e Al3+ Charge on aluminium ion= 13 10 = +3

accepts electron/electrons negatively charged ion (anion) no. of electrons in an atom < no. of electrons in its ion

Charge = no. of electrons in an atom no. of electrons in its ion

Eg: S + 2e S2 Charge on sulphide ion = 16 18 = 2

Nuclear or isotopic symbol of an ion

=

Atom


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Fill in the blanks with the appropriate response.

19 (a) 9W + e ________

40

(b) _____ e 19 X+

24 (c) 12Y + 2n ______

Sample Problem 2.4

Aluminium atom and ion, atomic no = 13 , mass no = 27

Nuclear symbol of aluminium atom = 27 Al Aluminium ion = 27Al3+

13 13

Sulphur atom and ion, atomic no = 16 , mass no = 32

Nuclear symbol of sulphur atom = 32S sulphur ion = 32S2-

16 16


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(a) The element Bi has a nuclear charge of +83. How many

electrons are there in the ion Bi3+?

(b) The total number of neutrons in the nucleus of ion X3+ is 1 more

than its number of protons. The mass number of the element X

is 9 times the charge on the ion. Write the appropriate symbol

for the ion X.

(c) The mass number of element R is 60. The atom has the same

number of neutrons, protons and electrons. What is the

complete symbol of atom R?

Sample Problem 2.5


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2.6 The Modern Reassessment of the Atomic Theory

1. All matter is composed of atoms. The atom is the smallest body that

retains the unique identity of the element.

2. Atoms of one element cannot be converted into atoms of another

element in a chemical reaction. Elements can only be converted

into other elements in nuclear reactions.

3. All atoms of an element have the same number of protons and

electrons, which determines the chemical behavior of the element.

Isotopes of an element differ in the number of neutrons, and thus

in mass number. A sample of the element is treated as though its

atoms have an average mass.

4. Compounds are formed by the chemical combination of two or more

elements in specific ratios.


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2.7 The modern periodic table.


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Metals, metalloids, and nonmetals.

Chromiu

m

Copper Cadmium

Lead

Bismuth

Boron

Silicon Arsenic

Antimony

Tellurium

Carbon

(graphite)

Sulfur

Chlorine Bromine

Iodine

The Atomic Symbols

Some symbols are one capital letter, like C, S, and I. Others are two

letters, and the second is lowercase, like Br and Sr

Some symbols come from the element s name, like C for carbon. Others come from the Latin name of the element, like Au for gold (aurum), Cu for

copper (cuprium), Fe (Ferum), Sn (Stanum), Pb (Plumbum) etc

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2.8 Bohrs Atomic Model of Hydrogen

The single electron in H atom moves around the nucleus in orbits or energy levels

The orbits are pictured as concentric circles around the nucleus. Energy levels are arranged in parallel lines.

Bohrs major idea was that the energy of the atom was quantized (discrete), known as photons and that the

amount of energy in the atom was related to the

electrons position in the atom

quantized means that the atom could only have very specific amounts of energy

The electrons travel in orbits that are at a fixed distance from the nucleus

therefore the energy of the electron was proportional to the distance the orbit was from the nucleus

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Figure 2.2 Principal Energy Levels in Hydrogen

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Electron Transitions

For transition to a higher energy state, the electron must gain the correct amount of energy corresponding to the

difference in energy between the final and initial states

Electrons in high energy states are unstable. They are in the excited state, and tend to lose energy and fall back to

lower energy states

Electrons emit radiation when they jump from an orbit with higher energy down to an orbit with lower energy

the emitted radiation was a photon of light

the distance between the orbits determined the energy of the photon of light produced

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Figure 2.3 Quantum leap

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Many of the properties of atoms are related to the energies of the electrons

Schdingers Equation allows us to calculate the probability of finding an electron with a particular amount

of energy at a particular location in the atom which is

refer to as orbital.

An orbital is a region (space) around the nucleus in which an electron of certain energy may be found. an orbital

characterizes the energy of the electrons reside in it.

2.9 quantum-mechanical model

The quantum-mechanical model explains the

manner in which electrons exist and behave in atoms

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The size, shape, and orientation in space of an orbital are determined by three integer terms called quantum

numbers n, and m.

In quantum mechanics, the distribution of electrons in an

atom or an atomic orbital is specified by quantum

numbers n, and m.

They are called the principal quantum number (n), the

angular momentum quantum number (), and the magnetic quantum number (m).

These quantum numbers will be used to describe atomic

orbitals and to label electrons that reside in them. A

fourth quantum number the spin quantum number (ms) describes the spin of a specific electron and completes the description of electrons in atoms.

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Figure 2.3

Quantum staircase.

Principal quantum number or principal energy levels (n)

As n gets larger, energy difference between orbitals

gets smaller

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A positive integer of 1,2,3 . infinity. It specifies the energy of an orbital.

Energy increases with the value of n. The higher is the energy of the orbital, the less stable are the electrons which reside in it.

It relates to the average distance of the electron in a particular orbital from the nucleus. The larger n is, the greater is the

average distance of an electron in the orbital from the nucleus

The maximum number of electrons that can reside in each principal shell is given the formula 2n2 in which n is the principal

quantum number of the orbital.

n Name of principal shell Distance from nucleus Energy

1 K Closest to nucleus Lowest energy

2 L

3 M

4 N Furthest from nucleus Highest energy

increases increases

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It designates the particular sub-energy level or subshell within a principal energy level. It describes the geometrical shape of the orbital.

has possible integral values from 0 to (n1). If n = 1, there is only one possible value of , that is, = 11 = 0. If n = 2, there are 2 values of (that is, = 0 and = 21 = 1) and so on. The value of is generally designated by the letters, s, p, d, ....

The energy of the subshells within the same principal quantum number, n, increases according to the order of s < p < d < f and so

on.

Angular momentum quantum number ()

0 1 2 3 4

Name of subshell s p d f g

Maximum number of electrons in

the subshell

2 6 10 14 18

Energy Low High increases

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designates a particular orbital within a given sublevel. It describes the orientation of the orbital in space relative to other orbitals.

the value of m depends on the value of the angular momentum quantum number, . For a certain value of , m can have any integral value ranging from to +. There are (2 + 1) integral values of m in each .

Examples: If = 0, m = 0; (2 + 1 = 1 value of m ) If = 1, m = 1, 0, +1; (2 + 1 = 3 values of m ) If = 2, m = 2, 1, 0, +1, +2; (2 + 1 = 5 values of m )

All the orbitals (different values of m ) within the same subshell (same value of () are degenerate orbitals which means they have the same energy but different orientation in space.

An atomic orbital can accommodate a maximum number of two electrons.

Magnetic quantum number (m)

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Table 2.3 The Hierarchy of Quantum Numbers for Atomic Orbitals

Name, Symbol

(Property) Allowed Values Quantum Numbers

Principal, n

(size, energy)

Angular

momentum, l

(shape)

Magnetic, ml

(orientation)

Positive integer

(1, 2, 3, ...)

0 to n-1

-l,,0,,+l

1

0

0

2

0 1

0

3

0 1 2

0

0 -1 +1 -1 0 +1

0 +1 +2 -1 -2

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Describing an Orbital

Each set of n, l, and ml describes one orbital

Orbitals with the same value of n are in the same principal energy level

Orbitals with the same values of n and l are said to be in the same sublevel

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Sample Problem 2.6

SOLUTION:

PLAN:

Determining Quantum Numbers for an Energy Level

PROBLEM: What values of the angular momentum (l) and magnetic (ml)

quantum numbers are allowed for a principal quantum number (n) of

3? How many orbitals are allowed for n = 3?

Follow the rules for allowable quantum numbers found in the text.

l values can be integers from 0 to n-1; ml can be integers from -l

through 0 to + l.

For n = 3, l = 0, 1, 2

For l = 0 ml = 0

For l = 1 ml = -1, 0, or +1

For l = 2 ml = -2, -1, 0, +1, or +2

There are 9 m values and therefore 9 orbitals with n = 3.

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Sample Problem 2.7

SOLUTION:

PLAN:

Determining Sublevel Names and Orbital Quantum

Numbers

PROBLEM: Give the name, magnetic quantum numbers, and number of orbitals

for each sublevel with the following quantum numbers:

(a) n = 3, l = 2 (b) n = 2, l = 0 (c) n = 5, l = 1 (d) n = 4, l = 3

Combine the n value and l designation to name the sublevel.

Knowing l, we can find ml and the number of orbitals.

n l sublevel name possible ml values # of orbitals

(a)

(b)

(c)

(d)

3

2

5

4

2

0

1

3

3d

2s

5p

4f

-2, -1, 0, 1, 2

0

-1, 0, 1

-3, -2, -1, 0, 1, 2, 3

5

1

3

7

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Figure 2.4

1s 2s 3s

Probability of finding an electron in the s-orbitals in different

energy levels = 0

Each principal energy level has one s orbital

Lowest energy orbital in a principal energy

state

Spherical

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Figure 2.5 The 2p orbitals, =1

Each principal energy state above n = 1 has three p orbitals, ml = 1, 0, +1

Each of the three orbitals points along a different axis : px, py, pz

2nd lowest energy orbitals in a principal energy state

Two-lobed

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Each principal energy state above n = 2 has five d orbitals : ml = 2, 1, 0, +1, +2

Four of the five orbitals are aligned in a different plane

the fifth is aligned with the z axis, dz squared

dxy, dyz, dxz, dx squared y squared

3rd lowest energy orbitals in a principal energy level

Mainly four-lobed

one is two-lobed with a toroid

d orbitals, =2

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Figure 2.6 = 2, d orbitals

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Electrons spin on their own axes and it is this motion that causes the electron to behave like a magnet.

It describes the two possible spinning motions of an electron, one clockwise and the other counterclockwise.

ms takes the values of + and . These values correspond to the two possible spinning motions of the electron, that is,

clockwise () and counterclockwise () .

Sample Problem 2.8

Electron Spin Quantum Number (ms)

Choose the sets of quantum numbers which are possible.

n m ms

2 2 0 -

2 3 2 -

4 2 1 +

4 1 -2 -

2 1 0 -

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n m ms

e1 3 2 0

e2 2 0 1

e3 1 0 0

e4 2 1 0

e5 1 1 1

(b) Correct the unacceptable sets of quantum numbers which you have

chosen in (a).

(c) Choose from e1 to e5

(i) an electron which has the highest energy.

(ii) two electrons which occupy the same orbital.

(iii) an electron which resides in d subshell.

(iv) an electron which resides in p subshell and in energy level L.

(a) Refer to the five sets of quantum numbers given below for electrons

e1 to e5 of an atom. Choose the sets of quantum numbers which are

not possible.

Sample Problem 2.9

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Table 2.4 Summary of Quantum Numbers of Electrons in Atoms

Name Symbol Permitted Values Property

principal n positive integers(1,2,3,) orbital energy (size)

angular

momentum

l integers from 0 to n-1 orbital shape (The l values

0, 1, 2, and 3 correspond to

s, p, d, and f orbitals,

respectively.)

magnetic m integers from -l0+l orbital orientation

spin ms + or - direction of e- spin

2.10 Electron Configuration and orbital diagram

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Distribution of electrons among the atomic orbitals

(a) Electron configuration in sublevel notation uses

numbers to designate the principal energy levels and

the letters, s, p, d, and f to identify the sublevels. A

superscript number following the letter indicates the

number of electrons in the designated subshell. For

example, the electron configuration of nitrogen atom

7N, is 1s2 2s2 2p3.

(b) The orbital diagram uses boxes to indicate orbitals

within subshells and arrows to represent electrons in

these orbitals. The directions of the arrows represent

the directions of the electron spins. The orbital diagram

for some atoms in the Periodic Table is shown in Table

2.5.

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Figure 2.7 Electron Configuration of H atom

principal energy level

of orbital occupied by

the electron

sublevel of orbital

occupied by the

electron

number of electrons in

the orbital 1s1

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Table 2.5 Electron configurations in sublevel notation and

orbital diagrams for some elements

Elements Energy

levels

Full Electron

configuration

Orbital diagram

11Na 2.8.1 1s2 2s2 2p6 3s1

5B 2.3 1s2 2s2 2p1

7N 2.5 1s2 2s2 2p3

8O 2.6 1s2 2s2 2p4

12Mg 2.8.2 1s2 2s2 2p6 3s2

In order to write the electron configurations in the sublevel notation for

the atoms, we need to describe three basic principles that govern the

distribution of electrons among atomic orbitals. The principles are:

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(1) Electrons occupy orbitals of the lowest energy available the Aufbau Principle

The energies of principal quantum numbers, n, increases with the

values of n, Therefore, the energy of each energy level increases

according to the order K < L < M < N

The order of filling of subshells can be predicted by using the Aufbau Principle shown below:

The orbitals will be filled up following the direction of the arrows.

Aufbau Principle

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

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Start by drawing a diagram

putting each energy shell on

a row and listing the sublevels,

(s, p, d, f), for that shell in

order of energy (left-to-right)

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d

7s

Next, draw arrows through

the diagonals, looping back

to the next diagonal

each time

Figure 2.8

Order of Sublevel filling of electrons when writing the electron

configuration in Sublevel notations.

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Illustrating Orbital Occupancies

The electron configuration

n l # of electrons in the sublevel

as s,p,d,f

The orbital diagram (box or circle)

Order for filling energy sublevels with

electrons

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(3) Hunds Rule

- Electrons will occupy all orbitals of the same energy level singly

and with same or parallel spin before they start pairing up

- This rule can be rationalised as follows: Two electrons with identical

charges tend to repel each other. Therefore, electrons prefer to

occupy orbitals separately as long as empty orbitals of the

appropriate energy are available.

(2) Paulis Exclusion Principle

- No two electrons in an atom can have four identical quantum

numbers n, , ml and ms.

- An atomic orbital can accommodate only up to two electrons

and these electrons must have opposite spins.

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The symbols given below describe four elements W, X, Y and Z (not the

actual chemical symbols of the elements).

17W2- 15X 41Y2+ 31Z

8 7 20 15

a) Calculate the number of neutrons in ion W2-.

b) Write the electron configuration i) of atom Y in sublevel notation

ii) in orbital diagram of atom Z.

c) Give the set quantum number (n, , m, ms) for the unpaired valence electrons which occupy the highest sub-shell of atom W.

d) How many electrons are there in the ion Z3- which have the

quantum numbers of = 1 and m = 0?

How many electrons are there in

i) Calcium ion which has the quantum numbers of m = 0 and ms = -?

ii) Fluorine atom which has the quantum numbers of = 1, m = 1?

Sample Problem 2.10

Sample Problem 2.11

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Core electrons are electrons that reside in the inner energy levels or lowest energy level of an atom.

Valence electrons are electrons that reside in the outer shell (principal shell containing electrons with the highest

quantum number) of an atom. They are the electrons that

can be involved when atoms participate in chemical

reactions or in chemical bonding.

2.11 Core Electrons and Valence Electrons

Sample Problem 2.12 Determining Quantum Numbers from Orbital Diagrams

PROBLEM: Write a set of quantum numbers for the third electron and a set

for the eighth electron of the F atom.

PLAN: Use the orbital diagram to find the third and eighth electrons.

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9F

1s 2s 2p

SOLUTION: The third electron is in the 2s orbital. Its quantum numbers

are

n = l = ml = ms= 2 0 0 + or -

The eighth electron is in a 2p orbital. Its quantum numbers are

2 n = l = ml = ms= -1, 0, or +1 1 + or -

2.12 Periodic Classification of Elements

The periods and the groups of elements in the periodic table correlate

closely with the electron configurations of the elements concerned. The

length of each block of elements in the periodic table is the maximum number of electrons the sublevel can hold

Choose only one for m and one for ms

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(1) The period number is the same as the principal quantum

number, n, of the electrons in the outermost principal shell.

Example:

All elements in the third period have one or more electrons with n = 3

and none with a higher value of n. The period begins with Na (1s2 2s2

2p6 3s1), and ends with Ar (1s2 2s2 2p6 3s2 3p6).

The next subshell to fill after 3p is 4s, so the next element after Ar is K.

K which is at the beginning of the fourth period has an electron

configuration of 1s2 2s2 2p6 3s2 3p6 4s1.

(2) The periodic table group number of an A group element (main

group or representative elements) is the same as the number

of outer shell electrons or valence electrons of the element.

Example:

All the elements in Group IA have a single electron in an s orbital of

the outermost principal shell and all the noble gases except He in

group 8A have 8 outershell electrons.

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Figure 2.9

A periodic table of partial

ground-state electron

configurations

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s1

s2

d1 d2 d3 d4 d5 d6 d7 d8 d9 d10

s2 p1 p2 p3 p4 p5

p6

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

1

2

3

4

5

6

7

Figure 2.10 Number of electrons in the sub-

shells of the different blocks

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It is a systematic classification and arrangement of the elements

according to increasing atomic numbers. According to the type of

subshell being filled, the elements can be divided into categories -

the representative elements, the noble gases, the transition

elements, the lanthanides, and the actinides. Elements are

arranged in vertical columns called groups and horizontal rows

called periods.

2.13 Periodic Table

Elements in the same group of the periodic table possess same number of valence electrons, therefore similar valence

shell electron configurations. As a result elements in the same

group show similar chemical properties.

Some groups are given special names as shown :

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Group Special name Group Special name

IA Alkali metal VIIA Halogen

IIA Alkali earth metals VIIIA Noble/rare gases

Generally, metal atoms have small numbers of electrons in their

valence shells. Except for hydrogen and helium, all s block elements

(groups I and II) are metals. All d and f block elements are metals.

A few of the p block elements like Al, Ga, Pb, Sn, In and Bi are also

metals.

Hydrogen is a group IA element but not an alkali metal because

it does not have any of the chemical characteristics of a metal. It

is a nonmetal.

Metalloids are elements that have the physical appearance of

metals but some nonmetallic properties. This group of elements

separates the metals from the nonmetals in the periodic table. Boron

(B), silicon, (Si), germanium (Ge) are some examples of metalloids.

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Nonmetal atoms generally have larger numbers of electrons in

their valence shell than do metals, except H. Non-metals are all p block elements (groups IVA to VIIA).

Examples: H2, N2, O2, F2 and Cl2 = diatomic gas

Noble gases = monatomic gas

C, P4, S8, I2 = solids with low melting points

Br2 = liquid.

The noble gases (the group VIIIA elements) all have a

completely filled p subshell except He. The electron

configurations are 1s2 for He and ns2 np6 for the other noble

gases, where n is the principal quantum number for the

outermost shell.

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Figure 2.11 Electron Configuration from the Periodic Table

P = [Ne]3s23p3

P has five valence electrons

3p3

P

Ne

1

2

3

4

5

6

7

1A

2A 3A 4A 5A 6A 7A

8A

3s2

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As = [Ar]4s23d104p3

As has five valence electrons

4s2

Ar 3d10

4p3

As

1

2

3

4

5

6

7

1A

2A 3A 4A 5A 6A 7A

8A

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Figure 2.12

Orbital occupancy for the first 10 elements, H through Ne.

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Table 2.6

3p

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Figure 2.13 Condensed electron configurations in subshell

notation in the first three periods.

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Transition Elements

For the d block metals, the principal energy level is one less than valence shell

one less than the Period number

sometimes an s electron is promoted to d sublevel

4s 3d

Zn

Z = 30, Period 4, Group 2B

[Ar]4s23d10

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Irregular Electron Configurations

Due to sublevel splitting, the 4s sublevel is lower in energy than the 3d; and therefore the 4s fills before

the 3d. But the difference in energy is not large

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Some of the transition metals have irregular electron configurations in which the ns only partially fills before the

(n1)d or doesnt fill at all

Their electron configuration has stability associated with half-filled or completely filled subshell.

Anomalous Electron Configurations

Expected

Cr = [Ar]4s23d4

Cu = [Ar]4s23d9

Mo = [Kr]5s24d4

Pd = [Kr]5s24d8

Half-filled or full subshells

Cr = [Ar]4s13d5

Cu = [Ar]4s13d10

Mo = [Kr]5s14d5

Pd = [Kr]5s04d10

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Electron Configurations of Transition Metal Cations

When transition metals form cations, the first electrons removed are the valence electrons (the ns subshell), even though other electrons were added after

Electrons may also be removed from the sublevel closest to the valence shell {the (n-1)d} sublevel after the valence electrons

The iron atom has two valence electrons

Fe atom = 1s22s22p63s23p64s23d6

When iron forms a cation, it first loses its valence electrons

Fe2+ cation = 1s22s22p63s23p63d6

It can then lose 3d electrons

Fe3+ cation = 1s22s22p63s23p63d5

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Table 2.7

3d

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Table 2.8

4p

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Figure 2.14 The relation between orbital filling and the

periodic table

Aufbau principle

Diamagnetic elements are elements in which all the valence electrons are

paired which means their subshells are complete. Paramagnetic elements

consists one or more unpaired electrons in their outermost sublevels.

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SAMPLE PROBLEM 2.13 Determining Electron Configuration

PLAN:

SOLUTION:

PROBLEM: Using the periodic table on the inside cover of the text, give the full

and condensed electrons configurations, partial orbital diagrams

showing valence electrons, and number of inner electrons for the

following elements:

(a) potassium (K: Z =19) (b) molybdenum (Mo: Z = 42) (c) Stannum (Sn: Z = 50)

Use the atomic number for the number of electrons and the periodic

table for the order of filling for electron orbitals. Condensed

configurations consist of the preceding noble gas and outer electrons.

(a) for K (Z = 19)

1s22s22p63s23p64s1

[Ar] 4s1

4s1

condensed configuration

partial orbital diagram

full configuration

There are 18 inner electrons.

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SAMPLE PROBLEM 2.13 continued

(b) for Mo (Z = 42)

1s22s22p63s23p64s23d104p65s14d5

[Kr] 5s14d5

(c) for Sn (Z = 50)

[Kr] 5s24d105p2



full configuration

5s1



full configuration 1s22s22p63s23p64s23d104p65s24d105p2

There are 36 inner electrons

and 6 valence electrons.

There are 46 inner electrons

and 4 valence electrons.

5s2 5p2 4d10

4d5

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SAMPLE PROBLEM 2.14

Writing Electron Configurations of Main-Group Ions

PLAN:

PROBLEM: Using condensed electron configurations, write reactions for the

formation of the common ions of the following elements:

(a) Iodine (Z = 53) (b) Potassium (Z = 19) (c) Indium (Z = 49)

Ions of elements in Groups 1A(1), 2A(2), 6A(16), and 7A(17) are usually

isoelectronic with the nearest noble gas.

Metals in Groups 3A(13) to 5A(15) can lose their np or ns and np

electrons.

SOLUTION:

(a) Iodine (Z = 53) is in Group 7A(17) and will gain one electron to be isoelectronic

with Xe: I ([Kr]5s24d105p5) + e- I- ([Kr]5s24d105p6)

(b) Potassium (Z = 19) is in Group 1A(1) and will lose one electron to be isoelectronic

with Ar: K ([Ar]4s1) K+ ([Ar]) + e-

(c) Indium (Z = 49) is in Group 3A(13) and can lose either one electron or three

electrons: In ([Kr]5s24d105p1) In+ ([Kr]5s24d10) + e+

In ([Kr]5s24d105p1) In3+([Kr] 4d10) + 3e-

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From left to right across a period, there is a transition from metals to

metalloids to nonmetals. There is also a gradual periodic variation in

the characteristic physical properties of metallic and nonmetallic

elements across a period.

Factors Affecting Atomic Orbital Energies

The Effect of Electron Repulsions (Shielding) (a)

Additional electron in the same orbital/energy sublevel (i)

An additional electron raises the orbital energy through electron

-electron repulsions.

Additional electrons in inner orbitals/energy sublevels (ii)

Inner electrons shield outer electrons more effectively than do

electrons in the same sublevel. This is due to the fact that repulsive

forces between electrons in different sublevels are stronger than

that between electrons in the same sublevel.

2.14 Periodic Trends

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The presence of shielding electrons reduces the electrostatic attraction between the positively charged protons in the nucleus and the outer

electrons.

The shielding effect increases down a group as each succeeding member has one inner shell more than the preceding member of the

group, OR the quantum number of the valence electron increases.

The Effect of Nuclear Charge (Zeffective) (b)

The repulsions between electrons in different energy shells cause the

valence electron to have a net reduced attraction to the nucleus it is shielded from the nucleus

The total amount of attraction that an electron especially the valence

electron feels for the nucleus is called the effective nuclear charge of the

electron. The shielding causes a reduction on the attractive forces

between the nuclear charge and the valence electron.

All the elements across a period have the same shielding effect because the number of inner shells remains the same across a period.

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Figure 2.15 The effect of nuclear charge on orbital energy.

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Figure 2.16 Shielding

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The effective nuclear charge, Zeff, acting on an electron is the actual nuclear charge, Z, less the screening effect of inner

electrons in the atom.

Zeff = Z number of inner shell electrons

Example: Sodium atom with 1 valence electron in 3s sublevel which is shielded from the positively charged nucleus by the 1s

and 2s electrons. Therefore, experience an effective nuclear

charge (Zeff) of +11 10 = +1.

Across a period, the value of Z increases while the number of inner electrons remains the same. Effective nuclear charge

increases from left to right of a period of representative elements

in the periodic table.

The effective nuclear charge remains the same from top to bottom within a vertical group A of the periodic table. The

effective nuclear charge calculated for elements Be, Mg and Ca

in group II are shown to be equal to +2.

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The atomic radius of an element is one-half the distance

between the two nuclei in two adjacent metal atoms or two like

atoms joined into a particular diatomic molecule.

Element Nuclear

charge

Total core electron

(inner-shell electrons)

Effective nuclear

charge

4Be + 4 -2 +2

12Mg +12 -10 +2

20Ca +20 -18 +2

13Al +13 -10 +3

15P +15 -10 +5

17Cl +17 -10 +7

Table 2.9 The effective nuclear charge of elements

2.15 Atomic Radius

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Figure 2.17 Defining metallic and covalent radii

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Atomic radii of the representative elements decrease from left to right in a horizontal row of the periodic table (refers

to as a period of the periodic table) because of a steady

increase in the effective nuclear charge across the

period which causes the valence electrons to be held more

strongly by the nucleus.

Atomic radii increase from top to bottom within a group of the periodic table because the principal

quantum number, n, of the valence electrons increases

down a group which causes the outer electrons to be

farther from the nucleus.

Effective nuclear charge (Zeff) is the positive charge felt by an electron.

Na

Mg

Al

Si

11

12

13

14

10

10

10

10

1

2

3

4

186

160

143

132

Zeff Core Z Radius (pm)

Zeff = Z - s 0 < s < Z (s = shielding constant)

Zeff Z number of inner or core electrons

Periodic Trends in the Size of Atoms and

Effective Nuclear Charge

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Figure 2.18

Atomic radii of the main-

group and transition

elements.

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Figure 2.19 Periodicity of atomic radius

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SAMPLE PROBLEM 2.15 Ranking Elements by Atomic Size

PLAN:

SOLUTION:

PROBLEM: Using only the periodic table, rank each set of main group

elements in order of decreasing atomic size:

(a) Ca, Mg, Sr (b) K, Ga, Ca (c) Br, Rb, Kr (d) Sr, Ca, Rb

Elements in the same group increase in size and you go down;

elements decrease in size as you go across a period.

(a) Sr > Ca > Mg These elements are in Group 2A(2).

(b) K > Ca > Ga These elements are in Period 4.

(c) Rb > Br > Kr Rb has a higher energy level and is far to the left.

Br is to the left of Kr.

(d) Rb > Sr > Ca Ca is one energy level smaller than Rb and Sr.

Rb is to the left of Sr.

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Figure 2.20 Depicting ionic radius.

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2.16 Trends in Ionic Radius

Ions in same group have same charge

Ion size increases down the column higher valence shell, larger

Cations smaller than neutral atoms; anions larger than neutral atoms

Cations smaller than anions except Rb+ & Cs+ bigger or same size as F and O2

Larger positive charge = smaller size of cation for isoelectronic species

isoelectronic = same electron configuration

Larger negative charge = larger size of anion for isoelectronic species

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INCREASE

INCREASE

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INCREASE

INCREASE

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Ionic Radius ()

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2.17 Trends in Cation Radius

When atoms form cations, the valence electrons are removed. The rest of the electrons will be attracted more strongly and closer to the nucleus.

Result: cations are smaller than their parent atoms

These new valence electrons also experience a larger effective nuclear charge than the old valence electrons, shrinking the ion even more

Traversing down a group increases the (n 1) level or inner shells causing the cations to get larger.

Traversing to the right across a period increases the effective nuclear charge for isoelectronic cations, causing the cations to get smaller. Example: Na+ > Mg2+ > Al3+

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2.18 Trends in Anion Radius

When atoms form anions, electrons are added to the valence shell, Zeff remains unchanged. Repulsion between

electrons in the valence shell causes the valence electron

to be further from the nucleus.

These new valence electrons experience a smaller effective nuclear charge than the old valence electrons, increasing the size

Results: Anions are larger than their parent atoms

Traversing down a group increases the n level, causing the anions to get larger

Traversing to the right across a period increases the effective nuclear charge for isoelectronic anions, causing

the anions to get smaller. The more negative charge the

anion the larger the size. Example: N3 > O2 > F

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Figure 2.21 Ionic vs. atomic radius.

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SAMPLE PROBLEM 2.16 Ranking Ions by Size

PLAN:

SOLUTION:

PROBLEM: Rank each set of ions in order of decreasing size, and explain your

ranking:

(a) Ca2+, Sr2+, Mg2+ (b) K+, S2-, Cl - (c) Au+, Au3+

Compare positions in the periodic table, formation of positive and

negative ions and changes in size due to gain or loss of electrons.

(a) Sr2+ > Ca2+ > Mg2+

(b) S2- > Cl - > K+

These are members of the same Group (2A/2) and

therefore decrease in size going up the group.

The ions are isoelectronic; S2- has the smallest Zeff and

therefore is the largest while K+ is a cation with a large Zeff

and is the smallest.

(c) Au+ > Au3+ The higher the + charge, the smaller the ion.

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Zr4+, Ti4+, Hf4+

Na+, Mg2+, F, Ne

I, Br, Ga3+

same column & charge,

therefore Ti4+ < Zr4+ < Hf4+

isoelectronic,

therefore Mg2+ < Na+ < Ne < F

Ga3+ < Br < I

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SAMPLE PROBLEM 2.17

PROBLEM: Rank each set of ions in order of increasing size, and explain your

ranking:

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The first ionization energy is the minimum energy required to

remove the first valence electron from the gaseous atom in its

ground state.

M(g) 1e M+

(g) H = First ionization energy

The second ionization energy is the energy required to remove

the second electron from the gaseous positive ion in its ground

state.

M+(g) 1e M 2+

(g) H = Second ionization energy

2.19 Ionization Energy

The ionization energy is the minimum energy required to

remove an electron from a ground state atom (ion) in the

gaseous state. The greater the ionization energy of an atom, the

more inclined the atom is to retain its electrons.

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The first ionization energies generally increase going from

left to right through a period in the periodic table.

As the atomic radii decrease from left to right of a period, the outer

electrons are more tightly held to the nucleus and higher ionization

energies have to be supplied to remove the first electron.

The first ionization energies decrease with elements moving down a group.

As the atomic radii of the elements within a group increase from

top to bottom, the average distance between the valence electrons

and the nucleus increase resulting in weaker nucleus-valence

electron attraction. Therefore, lower ionization energy down a

group. Consequently, the metallic character of the elements as well

as the reactivity of the metals increase down a group.

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Figure 2.22 Periodicity of first ionization energy (IE1)

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Anomalies in the Variation of the First Ionization

Energy Across a Period

The increase in the 1st ionization energy (I.E.) across a period is not uniform. Two anomalies occur in Periods 2 and 3:

From groups 2A to 3A (13) and groups 5A (15) to 6A (16)

B and Al (gp 3A) with smaller atomic radius are expected to have higher I.E. than Be and Mg (gp 2A), respectively, but the

reverse occurs.

The I.E. of B and Al are lower than that in Be and Mg due to

the valence electrons in the filled 2s and 3s orbitals of Be and Mg, respectively are more stable than the single valence

electron in the partially filled 2p and 3p orbitals of B and Al,

respectively.

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The single electron in the 2p and 3p orbitals are better shielded by the inner electrons than the electrons in 2s and

3s orbitals.

The I.E. of the gp 6A elements are expected to have higher energy than the gp 5A elements but the reverse occur because:

Half-filled p orbitals have special stability.

In addition,the valence electrons in the 2p and 3p orbitals of N and P, respectively are in 3 separate orbitals with minimum

repulsion, whereas, the paired electrons in one of the p

orbitals of O and S experience strong repulsion.

Therefore , it is easier to remove the valence electron from O and S than the valence electron of N and P which are more

stable with the least repulsion among the valence electrons.

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Figure 2.23 First ionization energies of the main-group elements.

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Figure 2.24 The first three ionization energies of beryllium (in

MJ/mol).

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Figure 2.25 Similar reactivity within a group

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SAMPLE PROBLEM 2.18 Ranking Elements by First Ionization Energy

PLAN:

SOLUTION:

PROBLEM: Using the periodic table only, rank the elements in each of the

following sets in order of decreasing IE1:

(a) Kr, He, Ar (b) Sb, Te, Sn (c) K, Ca, Rb (d) I, Xe, Cs

IE decreases as you proceed down in a group; IE increases as

you go across a period.

(a) He > Ar > Kr

(b) Te > Sb > Sn

(c) Ca > K > Rb

(d) Xe > I > Cs

Group 8A(18) - IE decreases down a group.

Period 5 elements - IE increases across a period.

Ca is to the right of K; Rb is below K.

I is to the left of Xe; Cs is furtther to the left and

down one period.

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Table 2.10

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SAMPLE PROBLEM 2.19 Identifying an Element from Successive

Ionization Energies

PLAN:

SOLUTION:

PROBLEM: Name the Period 3 element with the following ionization energies

(in kJ/mol) and write its electron configuration:

IE1 IE2 IE3 IE4 IE5 IE6

1012 1903 2910 4956 6278 22,230

Look for a large increase in energy which indicates that all of the

valence electrons have been removed.

The largest increase occurs after IE5, that is, after the 5th valence

electron has been removed. Five electrons would mean that the

valence configuration is 3s23p3 and the element must be

phosphorous, P (Z = 15).

The complete electron configuration is 1s22s22p63s23p3.

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For multiple charge anions, the electrons are added stepwise with a

different electron affinity for each step. Consider the formation of an

oxide ion from an oxygen atom.

O (g) + e O

(g) H = 142 kJ O(g) + e O

2(g) H = +745 kJ

2.20 Electron Affinity

Electron affinity is the energy change that occurs when a free

electron is accepted by an atom in its gaseous state. The more

energy that is released, the larger the electron affinity.

The more negative the electron affinity, the greater the tendency of

the atom to accept an electron.

Example : F(g) + e F(g) H = 320 kJ

Li(g) + e Li(g) H = 61 kJ

Na(g) + e Na(g) H = 54 kJ

Ne(g) + e Ne(g) H = +29 kJ

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The electron affinity values become more negative from left to right across a period.

As the atomic radius of the element decreases, the attractive forces

of the nucleus increases, hence the tendency to accept electrons

increases. The electron affinities of metals (left) are generally more

positive (or less negative) than those of nonmetals.

Group 5A generally lower EA than expected because extra electron

must pair.

Group 2A and 8A generally very low EA because added

electron goes into higher energy level or sublevel

The second electron affinity is a positive quantity because an

electron approaches an ion with a net charge of 1. It is strongly repelled, and work must be done to force the extra electron onto

the O (g) ion.

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All noble gases exist as monatomic species because they are very unreactive and have little or no tendency to combine among

themselves or with other elements.

The electron configurations of the noble gases show that their

atoms have completely filled outer s and p subshells, indicating

great stability. Thus, the group VIIIA ionization energies are among

the highest of all elements, and they have no tendency to accept

extra electrons (smallest electron affinity).

Within a group, electron affinity values of elements become less negative from top to bottom.

The increase in atomic radii causes the attractive forces of the

nucleus to decrease and so there is less tendency to accept an

electron.

the atom with the Highest Electron Affinity in any period = halogen

Ionisation energy and electron affinity apply only to isolated gaseous atoms, and not directly to atoms in molecules.

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Figure 2.26 Electron affinities of the main-group elements.

-2

-5

-10

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Figure 2.27

Trends in three atomic properties.

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Figure 2.28 Trends in metallic behavior.

Metallic Character

Metallic character is how closely an elements 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

Measuring the magnetic behavior of a sample.

The apparent mass of a

diamagnetic substance is

unaffected by the magnetic

field.

The apparent mass of a

paramagnetic substance

increases as it is attracted by the

magnetic field.

Magnetic Properties of Transition Metal

ions

Magnetic behavior can provide evidence for the

electron configuration of a given ion.

Ti (Z = 22)

4s 4p 3d

4s 4p 3d

Ti2+

Ti2+ has 2 unpaired electrons and is paramagnetic,

providing evidence that the 4s electrons are lost before

the 3d electrons.

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Alkali Metals Table 2.11

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Figure 2.29 The trend in acid-base

behavior of element

oxides.

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Figure 2.30 Main-group ions and the noble gas configurations.

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Choose the more metallic element in each pair

i) Sn or Te ii) P or Sb iii) Ge or In iv) S or Br v) Mg or Al

vi) Si or Sn vii) Br or Te viii) Se or I

SAMPLE PROBLEM 2.20

CHM092 2 Matter Quan Config Periodicity

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Transcript of CHM092 2 Matter Quan Config Periodicity