chapter 4-dmc 101-arrangement of electron.pdf

7/30/2019 Chapter 4-DMC 101-Arrangement of Electron.pdf

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CHAPTER 4: THE ARRANGEMENT OF

ELECTRONS IN ATOMS

Reaction happens WHEN electronin atoms

interact.

Electromagnetic Radiation

When electron interacts produces

electromagnetic radiation

Electromagnetic radiation carries energy Example:

visible light, x-rays, radio waves

Audi Majdan - DMC 101 - KLIUC

Chemical

Reaction

Interaction of

electrons in atoms=

1


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Properties of electromagnetic radiation:

(a)

Move through vacuum at speed oflight, c = 3.00 x 108 m/s

(b) Wave-like properties

Properties Symbol Unit Pronunciation

frequency (Hz) nuwavelength (m) lambda

Audi Majdan - DMC 101 - KLIUC 2


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Relationship:

= c

c (speed of light) = 3.00 x 108 m/s

Electron interaction depends on

arrangement of electron in atoms.

The arrangements of electronin atoms

provides information on:(a) Number of electron

(b) Position of electron

(c) Energypossess by atom



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Quantum and Photons

Electron orbiting the nucleus have it's

own energy Energy can be released orabsorbed by the

atoms during the interaction

Quantum = a unit of energy released or

absorbed by atoms

A photon = a quantum of energy

E = h or E = hc

E = Energy of a quantum

= frequencyh = Planck's constant

= 6.63 x 10-34 Joule seconds (Js)

Example:

Calculate the smallest amount of energy (one

quantum) that an object can absorb from yellow

light with a wavelength of 589 nm.



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Solution

1) Frequency:

= c

= c/

= (3.00 x 108 m/s)

(589 x 10-9 m)

= 5.09 x 1014 s-1

2) Energy:

E = h

E = (6.63 x 10-34 Js) x (5.09 x 1014 s-1)

E = 3.37 x 10-19 J

Exercise 1

A laser emits light with a frequency of 4.69 x

1014 s-1. What is the energy of one quantum ofthe laser?

Electrons Arrangement in Atom



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Several systems are created to know the

LOCATION ofELECTRONs in ATOM

Bohr's Model of Atom

Dalton's model of atomprovides the basisof atom structure

Bohr's model of atom is an advanced

model of atom structure

Characteristics of Bohr's Model:

(a) Electrons orbiting the nucleus(b) The electrons are restricted only to orbits

of certain radius and energies

Audi Majdan - DMC 101 - KLIUC

ELECTRONS

ARRANGEMENT

Single electron atom Multiple electrons atom

Bohrs Model Schrodingers Model

6


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(c) An electron orbiting in an orbit has a

fixed energy state and will not release or

absorb energy

The energy possess by the electron:

RH = Rydberg constant

= 2.18 x 10-18 J (for calculation of E)

or

= 1.097 x 107 m-1 (for calculation of )

n = principle quantum number

= orbits for the electron

Energies always in negative value

n=1 an electron in the first allowed orbit

(closest to the nucleus)

n=2 an electron in the next allowed orbit

further from the nuclei n=3,4,5



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The energies of allowed orbits for the

electrons:

The lower the energy (more negative), the

more stable the atom



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Ground state = electron is in the lowest

energy state (n=1) the most stable state

Electrons can jump from one orbit to

another

Excited state = electron is in other state

than n=1 (n 2)

When n = infinity

Electron completely separated from the

atom

E = (-2.18 x 10-18 J) x (1/)

= 0

Reference or zero energy state = the state

in which the electron is separated from the

nucleus



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Electron moves to higher state (higher n)

= Energyabsorbed

Electron moves to lower state (lower n)

= Energyemitted

The difference in energy levels between the

final and initial orbits:

E = Ef - Ei

or

ni = initial orbit

nf = final orbit

E=+ve when energy absorbed:nf> ni = electron moves to higher n



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E=-vewhen energy emitted:ni > nf = electron falls to lower n

Example:

Calculate the energy of light that corresponds to

the transition of the electron from the n=4 to the

n=2 state of the hydrogen atom. Is the lightabsorbed or emitted by the atom?

Solution

ni = 4

nf = 2

RH = 2.18 x 10-18 J (for calculation of E)

E = (2.18 x 10-18 J) x ((1/16)-(1/4))= -4.09 x 10-19 J

(light is emitted)



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Exercise 2

Determine the energy of light that corresponds

to the transition of electron from n=3 to n=5

state for hydrogen atom. Identify whether the

energy is absorbed or released.

Bohr's Model for Hydrogen atom

Line Spectra

Line spectra = spectrum of coloured line

with different and separated wavelengths

when electron fall (move to lower orbit)

Example:

A rainbow represents the spectrum of

wavelengths of light contained in the light

emitted by the sun

Forhydrogen atom: 5 series ofline spectra

observed depending on the type of electrontransition



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Series nf ni Spectrum

Region

Lymann 1 2,3,4.. Ultra violet

Balmer 2 3,4,5.. Visible light

Paschen 3 4,5,6.. Infrared

Brackett 4 5,6,7.. Infrared

Pfund 5 6,7,8.. Infrared

To determine wavelengths of the spectrum:

=

22

111

fi

H

nnR

Example:

Calculate the wavelength of energy released

during a transition from the n=2 to n=4.

=

22

111

fi

H

nnR

RH = 1.097 x 107 m-1 (for calculation of )

=

22

7

4

1

2

110097.1

1

( )1875.010097.11

7

=



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16

1006.21

= m

m710862.4 =

Exercise 3

Calculate the wavelength of the hydrogen

emission line that corresponds to the transition

of the electron from the n = 3 to the n =1 state.

Schrodinger's Model of Atom

Difference between Bohrs Model and

Schrodingers Model

Properties Bohr Schrodinger

Type of atom Single electron

atom

Multiple

electrons atom



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Electrons

Location

In a defined

orbit

Probability

electron

distribution

only

Notation

system

One quantum

number:

n

Four quantum

numbers:

n, l, ml, ms

Probability distribution of an electron in the

ground state in a hydrogen atom:



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Uncertainty Principle: it is impossible to

simultaneously know both the position

and momentum of an object as small asan electron

Schrodingers Notation System

The Bohrs model: one quantum number (n)to describe an orbit

The Schrdinger model: four quantum

numbers:


More probability

to find an

electron

Less probability to

find an electron


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(a) Principal quantum number (n),

(b) Orbital quantum number (l),

(c) Magnetic quantum number (ml)

(d) Spin quantum number (ms)

to describe an orbital.

(a) The principle quantum number (n)

n = shell (the energy level/orbits) in the

atom

n = 1, 2, 3, ...

Example: n = 1 (electron is at 1st shell)

As n increases, the electron density is

further away from the nucleus

As n increases, the electron has a higher

energy and is less tightly bound to the

nucleus

(b) The azimuthal quantum number (l)



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l= subshell in a shell

The value ofldepends on n

For every n, there are n value of l

l= (n-1)

lis referred to by a letter:

s = 0, p =1, d =2, f =3

l 0 1 2 3 4

notation s p d f g

Example:

n = 1 l= (n-1)

= 1-1= 0

(n = 1: one subshelllwith value of 0)

Notation: 1s

n = 2 l= 2-1


n = 1 l= 0


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= 1

(n = 2: two subshelllwith value of 0 and 1)

Notation: 2s (n = 2 and l= 0 )

Notation: 2p (n = 2 and l= 1)

n lSubshell

notation

1 0 1s2

0 2s

1 2p

3

0 3s

1 3p

2 3d

40 4s1 4p

2 4d

3 4f

(c) The magnetic quantum number (ml)

ml = orbitals in a subshell Orbital is the place where electrons located

ml = all values between 'l', 0, -'l'



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nValues

ofl

Subshell

Designat

ion

Values

ofml

Number of

Orbitals in

Subshell

Total

number

of orbitals

in shell

1 0 1s 0 1 1

2 0 2s 0 14

1 2p 1,0,-1 3

3 0 3s 0 1

91 3p -1,0,1 3

2 3d-2,-

1,0,1,25

4 0 4s 0 1

16

1 4p -1,0,1 3

2 4d-2,-

1,0,1,25

3 4f

-3,-2,-

1,0,1,2,

3

7



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(d) The electron spin quantum number (ms)

ms = direction of spin of the electron

ms = + or -

Explanation on The Quantum Number

Electron shell = consisting ofsubshell

Subshell = consisting oforbitals

Orbitals = consisting ofelectrons


+

-


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Example:

All orbitals that have principle quantum

number of 3 (n=3) is said to be in third shell.

The third electron shell ('n'=3) consists of the

3s, 3p and 3dsubshells

The 3s subshellcontains 1 orbital,

the 3p subshellcontains 3 orbitals

the 3d subshellcontains 5 orbitals

The third electron shellis comprised of

nine different orbitals, but all orbital with

the same energy (true for hydrogen only)

Each subshell is divided into orbitals:

Subshell Number of orbitals

s 1

p 3

d 5

f 7



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Energies of all hydrogen electron orbitals

until n=3:

At room temperatures all hydrogen atoms

are in their ground states

The electron may be promoted to an excitedstate by the absorption of a photon

Example:

a) Predict the number of subshells in the

fourth shell, n=4.

b) Give the label for each of the subshells

in (a)

c) How many orbitals are there in each of

the subshells in (c)?



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Answer:

a) n = 4 l= n-1

= 3

Therefore: 0, 1, 2, 3

4 subshells

b) s = 0, p = 1, d = 2, f = 3

Therefore: 4s, 4p, 4d, 4f

c) 4s l= 0, value forml ( from -l to l)

= 0 one 4s orbital

4p three 4p orbitals;

4d five 4d orbitals;

4f seven 4f orbitals

Exercise 4

a) What is the designation for the subshell

with n =5 and l=1?

b)How many orbitals are in this subshell?

c) Indicate the values ofml for each ofthese orbitals.

Orbitals in Multiple-Electron Atoms



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The hydrogen atom only 1 electron.

All subshells ofhydrogen atoms with the

same principle quantum number (n) have thesame energetic level

In multiple electrons atom: electron-

electron repulsion influences the energy

levels of the orbitals. Example:

Energy state 2s orbital < 2p orbital in a

many-electron atom

Electron spin and the Pauli Exclusion

Principle



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Pauli exclusion principle = No two

electrons in an atom can have the same

set of four quantum numbers (n, l, mland

ms)

For any given orbital has fixed values of:

n

l

ml

ms quantum number is use to differentiate

between 2 electron in an orbital (+ and

- )

Only two electrons at most can occupy the

same orbital, and they have opposite valuesfor magnetic spin

Example:

2 electron in one orbital can have quantum

number of

(1, 0, 0, +) & (1, 0, 0, -)

Electron Configurations



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Electron configuration = the way electrons

are distributed among the orbitals

Orbitals are occupied in order of

increasing energy, with no more than two

electrons per orbital

Example

Lithium

1. Lithium has 3 electrons.2. Start : Place two electrons in the 1s,

lowest energy orbital. These two electrons

have opposite ms (+ & -).

3. Place the third electron in the next

highest energy level orbital = the 2s orbital:



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The arrows indicate the value of the

magnetic spin (ms) quantum number

(up for + and down for - )

Diagram representation:



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1s2 2s1


Li

1s 2s

Orbital

Diagram

Electronic

Configuration


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Writing electronic configurations: Period 1, 2,

3

All 2p orbitals are equal energy any 2p

orbital can be filled

Carbon has 6 electrons. 3 possible

arrangements can be made in 2p orbital.

Which one is true?


(a) (b) (c)


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Hund's rule:

For degenerate orbitals, the lowest energy is

attained when the number of electrons with thesame spin is maximized

The second 2p electron in Carbon is placed

in another 2p orbital, but with the same

spinas the first 2p electron

Core electrons: the electrons in the stable

(Noble gas) configuration

Valence electrons: the electrons in the outer

shell (beyond the stable core)



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Exercise 5

1. Draw the orbital diagram

representation for the electron

configuration of oxygen, atomic

number 8.

2. Write the electron configuration of

phosphorus, element 15.

chapter 4-dmc 101-arrangement of electron.pdf

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