CHM11-3 Lecture 2

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LectureNo. 2(The Electronic Structure of Atoms)

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Content

Atomic Models

Quantum Mechanics

Electronic Configuration

Periodic Relations of Elements

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Atomic Models (Cont.)

John Dalton (1800)

He proposed a modern model based on experimentation.

Daltons Atomic Theory:

1. Elements are composed of extremely small particles called

atoms. All atoms of the same element are alike, and atoms

of different elements are different.

2. The separation of atoms and the union of atoms occur in

chemical reactions. In these reactions, no atom is created or

destroyed, and no atoms of one element are converted intoan atom of another element.

3. A chemical compound is the result of the combination of

atoms of two or more elements. A given compound always

contains the same kinds of atoms combined in the same

proportions.

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J. J. Thomson (1900)

He discovered electron and proposed a model of the

atom called Plum Pudding Model.

Atoms were made

from a positively

charged substancewith negatively

charged electrons

scattered about.

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Ernest Rutherford (1910)

He proposed that atoms are mostly empty space and

negative electrons orbit a positive nucleus using the

Gold Foil Experiment.

The Gold Foil Experiment.

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Rutherfords Model:

A nucleus exists in

the center of the

atom.The nucleus contains

protons and neutrons

which together

account for the mass.

Electrons, which occupy most of the total volume

of the atom, are outside the nucleus and move

rapidly around it.

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Neils Bohr (1913)

He proposed an improved atomic model.

The Bohrs Model.

Electrons move in

definite orbits aroundthe nucleus.

These orbits, or

energy levels are

located at certaindistances from the

nucleus.

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Nuclear Bohr Model of Hydrogen Atom:

Electrons normally

exist in the lowerenergy state (ground

state).

When an electron

jumps into higherenergy state it is

said to be in an

exited state.

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Werner Heisenberg (1925)Theorized that it is impossible to know simultaneously

both the velocity and position of a particles

(Heisenberg Uncertainty Principle).The probable location of an electron is based on how

much energy the electron has.

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Electron Cloud Model:

Electron cloud is a

space in which electron

are likely to be found.Electrons whirl about

the nucleus billion of

times in 1 second.

They are not movingaround in random

pattern.

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Quantum Mechanics

Electromagnetic Radiation

Light travels through space in a wave motion.

THEWAVENATURE OF LIGHT

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Quantum Mechanics (Cont.)

Definitions

Wavelength () the distance between two

similar points on two successive waves. It

should be expressed in meters. (1nm = 10-9 m).

Amplitude (a) height of a crest or depth of

a trough.

Intensity (brightness) of Radiation

is proportional to the square of amplitude (a2).

Speed of light (c)

equivalent to 2.998 x 108

m/sec.

c =

Frequency () is the

number of waves that pass a

given spot in a second, this is

the reciprocal of second (1/s =hertz).

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Colors of Visible Spectrum:

Color Wavelength, nm

Violet 400 450

Blue 450 500

Green 500 570

Yellow 570 590

Orange 590 620

Red 620 750


Sample Problem:

What is the frequency of red light with a wavelength of

700nm and violet light with a wavelength of 400nm?

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Heinrich Hertz

Generated electromagnetic waves with long wavelengths

larger than those of visible light and who demonstrated that

long wavelength radiation exhibits the same phenomena as

light does.

Electromagnetic

Radiation

(Spectra).

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MaxPlank

Proposed the quantum theory of radiant energy.

Suggested that radiant energy could be absorbed orgiven off only in definite quantities called quanta.

E= hv = hc /

hPlanks constant (6.626 x 10-34 J-s)

THEPARTICLENATURE OF LIGHT


Albert Einstein

Proposed that Planks quanta were discontinuous bits of

energy called photons.

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Sample Problem:Consider a violet light with a wavelength of 400nm.

Calculate the energy, in joules, of one photon of this

light.

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Isaac Newton

Showed that visible (white light) from the sun can be

broken down into its various component by a prism.

ATOMIC SPECTRA


Definitions

Continuous Spectrum spreading out

into a wide range band of the white light.

Emission Spectrum When an elementabsorbs sufficient energy (from a flame or

electric arc), it emits radiant energy in the

form of light. When this radiation is passed

through a prism, it separates into a

component wave length.

Absorption Spectrum when

continuous radiation (white light)

passes through a substance, certain

wavelengths of radiation may beabsorbed. These wavelengths are

characteristics of a substance that

absorbs the radiation and pattern of

these lines are referred to as an

absorption spectrum.

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Every element has its own unique line spectrum,

therefore these spectra are characteristics of an

atoms electronic structure .

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Bohrs Theory:

1. The electron of the hydrogen atom can exist only in certain

circular orbits (energy levels or shell).

2. The electron has a definite energy characteristic of the orbit

in which it is moving.

3. When an electron of an atom is as close to the nucleus, it isin the condition of the lowest energy called the ground state.

When an atoms are heated in an electric arc or Bunsen

burner, electron absorbs energy and jump to outer levels,

which are higher energy states (excited state).4. When an electron falls back to a lower level, it emits a

definite amount of energy. The energy difference between

the high-energy state and low-energy state is emitted in the

form of a quantum of light.

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Louis De Broglie

He reasoned that if light could show the behavior of

particles (photons) as well as waves, then perhaps an

electron, which Bohr had treated as particle, could behavelike a wave. (Dualistic Nature of Light).

QUANTUM MECHANICAL MODEL


Erwin Schrdinger

Formulated wave equation that relates the energy of the

electron to its position in the atom.

Solutions of these equations give rise to quantum

numbers.

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ATOMIC ORBITALS

Bohrs Model OUTDATED!!!Schrodingers Equation Quantum Mechanical Model

Heisenberg Uncertainty Principle impossible to know

precisely both the velocity and locations of an electron at

the same time

n = 1 s 1

n = 2 s,p 1,3

n = 3 s,p,d 1,3,5n = 4 s,p,d,f 1,3,5,7

Principal

energy level

Sublevel No. of orbitals

in each sublevel

nucleus

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Atomic Orbital Shapes:

s and p Orbitals

d Orbitals

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QUANTUM NUMBERS set of numbers that describes an

electron orbital


1. First Quantum Number (Principal) (n) indicates the

main energy level by the electron. It defines the total energy

of the electrons and has values from (1 to 7).

2. Second Quantum Number (Azimuthal) (l) it describes

the way the electron moves around the nucleus or the shape

of the probability distribution. The values range from 0 to (n-1).

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n l Spectral Lines1 0 s type (sharp)

2 0

1

s

p type (principal)

3 0

1

2

s

p

d type (diffuse)

4 01

2

3

sp

d

ftype (fundamental)

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3. Third Quantum Number (Magnetic) (ml) it defines the

possible orientation of the electrons in space. The values arefroml through 0 to +l.

4. Fourth Quantum Number (Spin)

(ms)

it takes intoaccount the spinning of the electron around its own axis as it

moves about the nucleus. The spin is either clockwise or

counterclockwise. The values are +1/2 (clockwise) and -1/2

(counterclockwise).

l ml

Number of orbitals (2l + 1)

0 0 11 -1, 0, +1 3

2 -2,-1, 0, +1, +2 5

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Aufbaus Principle

Electrons willsuccessively occupy

the available orbitals

on order of

increasing energy.

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Quantum Number

ElectronsAddress

4 #s and no two electrons can have the same quantumnumbers

(n, l, ml, ms)

n = principal energy level

(cannot be zero) n = 1, 2, 3, 4, 5, 6, 7l = sublevel (s, p, d or f) l = n-1

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

ml = orbital ml = -l to +ls _ p _ _ _ d _ _ _ _ _ f _ _ _ _ _ _ _

0 -1 0 +1 -2 -1 0 +1 +2 -3 -2 -1 0 +1 +2 +3

ms

= spin + or -

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Example:Write the

possible set of

quantum

numbers for theelectrons in:

(a) 3s

(b) 3d

(c) 4f

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Aufbaus Principle

Electrons willsuccessively occupy

the available orbitals

on order of

increasing energy.

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Hunds Rule

When filling a set of degenerate energy levels, the

electron enter the orbitals singly, with spins in the same

direction (same as s number), until the set is half filled,

before they pair up with opposite spins.


Concepts:

Paulis Exclusion Principle

Each electron within a given atom must have a unique

set of the four quantum numbers.

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Refers to the arrangement of electrons in energy

levels.Methods ofWriting

1. Orbital method

2. Shell method3.Arrow Rectangular method

4. Core Method

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Electronic Configuration (Cont.)

Example (neutral atom):

Write the electronic configuration using the four

methods of the following elements:

(a) Br

(b) Ca

Example (monoatomic ions):

Give the electronic configuration of:

(a) V4+(b) Cl

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Magnetism

Paramagnetic substances that contain net unpaired

electrons and are attracted by a magnet.Ex. Li

Diamagnetic substances that do not contain net

unpaired electrons and are slightly repelled by amagnet.

Ex. Mg


RelatedConcepts:

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Periodic Relations ofElements

Periodic Classifications ofElements Elements may be classified according to their

electronic configurations.

1. The Noble Gases

They are also known as InertG

ases orGroup O elements. They are colorless monoatomic gases,

which are chemically unreactive and diamagnetic. They have

outer configurations of ns2np6 (except forHelium).

2. The Representative Elements These elements are foundin the A families of the periodic table. They exhibit a wide

range of chemical behavior and physical characteristics. The

chemistry of these elements depends upon the valence

electrons.

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3. The Transition Elements They are found in the B families

of the periodic table. All of these elements are metals and

most of them are paramagnetic and form highly colored,

paramagnetic compounds.

4. The Inner-Transition Elements These elements are foundat the bottom of the periodic table, but they belong to the 6th

and 7th periods after the elements of group IIIB. All inner-

transition elements are metal and are paramagnetic. Their

compounds are also paramagnetic and colored.

Periodic Relations ofElements (Cont.)

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Periodic Variation in Physical Properties1. Atomic Radius one-half the distance between the nuclei of

the two atoms in an elemental substance.


TREND:

LEFT RIGHT decreases

TOP

BOTTOM increases

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2. Ionic Radius one-half the distance between the nuclei of ametal and a non-metal.


TREND:

LEFT RIGHT decreases*

TOP BOTTOM increases

positive ions are smaller than

the metal atoms from which they

are formed negative ions are larger than

the nonmetal atoms from which

they are formed

* this happens when comparing both

metals and non-metals but by

comparing the metals and non-metals, nonmetals

have larger ionic radius than metals.

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3. Ionization Energy the minimum energy required to remove

an electron from a gaseous atom in its ground state.


TREND:LEFT RIGHT increases

TOP BOTTOM decreases

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4. ElectronAffinity measure of the energy change when

electron is added to a neutral atom to form a negative ion.

TREND:LEFT RIGHT increases

TOP BOTTOM decreases

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END of Lecture No. 2

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CHM11-3 Lecture 2