Atomic Structure:

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Chapter 3 Atomic Structure Images of the
Invisible

• Outline
• 1. Introduction
• 2. Discovery of electrons and Plum Pudding model
  of atoms
• 3. Rutherford model of atoms
• 4. Bohr model of atoms
• 5. Quantum model of atoms
• 6. Periodic table and electronic configuration

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1. Introduction

• The previous chapter gives evidences for the
  existence of atoms from
• Chemical reactions
• Scanning tunneling microscope
• Chemistry studies behaviors of atoms
• There is a need to know the structures of atoms.
• However, atoms are very small.
• How can we know the structures of atoms?
• What are the structures of atoms?

The purpose of this chapter is to answer these
questions.
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2. Discovery of electrons and Plum Pudding Model
of Atoms

• Volta invented electrochemical cells in 1800.

• Notations
• Electrodes
• Cathode negatively charged
• Anode positively charged
• Electrolyte
• Compound conducts electricity
• Made of cation and anion
• Ions (charged atom or group of atoms)
• Cation is positively charged moves to cathode
• Anion is negatively charged moves to anode

cation
anion
Humphry Davy (17781829) Michael Faraday
(17911867)
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Electrolysis experiments

• Electricity can split compounds. e.g.

Electrolysis of Water
Electrolysis of molten NaCl
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• Chemical reactions can produce electric current
  in battery, e.g.

Zn Cu2 -? Zn2 Cu
Chemical reaction ? current Current ? chemical
reaction Conclusion
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Cathode Ray Tube and discovery of electrons
Williams cooks (1832-1913) constructed a Cathode
Ray Tube

• When current is passed through a tube filled
  gas at low pressure (or vacuum)
• Cathode rays are generated
• Cathode rays travel in straight line unless some
  kind of external field is applied

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Thomsons Experiment

• When an external field is applied
• The ray beam was attracted by positive plate, and
  repelled by negative plate. ?negative charged
  particle.
• Same particle regardless of gas or electrode.
  ?electron
• Cathode ray beam of electrons
• He determined mass-to-charge ratio, but not mass
  or charge.

By Joseph John Thomson in 1897 Won Nobel Prize
in Physics in 1906
Conclusion
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Electron Charge

• 1909, Robert Millikan Using the Oil Drop
  Experiment, he determined the charge of an
  electron.
• The mass of an electron can also be calculated
  based on mass-to-charge ratio.

Charge of e- smallest possible deference in
charge between two droplets

• Mass of the electron
• 9.1 x 1028 g
• Would take 1 x 1027 to make 1 gram of electrons

By Robert A. Millikan (18681953) Won Nobel Prize
in Physics in 1923
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PLUM PUDDING MODEL OF ATOMS
In 1904,Thompson develops the idea that an atom
was made up of electrons scattered unevenly
within an elastic sphere surrounded by a soup of
positive charge to balance the electron's charge
like plums surrounded by pudding.
PLUM PUDDING MODEL
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3. Observation of positively charged particles
and Rutherford Model of Atoms
The plum pudding model was abandoned, to account
for other experimental observations.
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Goldsteins Experiment Observation of positive
particles
1886, Goldstein observed positive rays using a
perforated cathode (gas discharge tube).

• Positive particle of rays flows in the opposite
  direction toward the cathode.
• Mass of the particle is dependent on gas present
  in the tube.
• For H, the particle has a mass 1837 times that of
  electron.

Conclusion
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Discovery of Radioactivity

• In 1895, Becquerel found that uranium ores would
  spontaneous emit radiation.
• Marie and Pierre Curie isolated other elements
  that behaved like uranium.
• Radioactivity spontaneous emission of radiation
  from certain unstable elements

Antoine Henri Becquerel (18521908) Shared Nobel
Prize in Physics in 1903 with Marie Curie and
Pierre Curie
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Type and Behavior of Radioactivity
Three commonly found types of radioactivity
a Helium nuclei He2
b electron
g ray
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Types of Radioactivity

• Three commonly found types

Helium nuclei He2
Named by Ernest Rutherford (18711937)
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Rutherfords Experiment
Using an apparatus similar to that shown, Ernest
Rutherford discovered the atomic nucleus.
They fired a-particle (Helium nuclei) at a piece
of gold foil. While most of the helium nuclei
passed through the foil, a small number of them
were deflected and, to their surprise, some
helium nuclei bounced straight back.
Hans Geiger (18821945) Ernest Marsden (18891970)
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Model Explaining Rutherfords Experiment

• Originally assumed that all particles in an atom
  were evenly spread out
• All particles should pass.
• Old model cannot explain the results of his
  experiment
• Needed a new model

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Rutherford Model of Atoms

• 1. The atom contains a tiny dense center called
  the nucleus.
• The amount of space taken by the nucleus is only
  about 1/10 trillionth the volume of the
  atom.
• 2. The nucleus has essentially the entire mass of
  the atom.
• The electrons weigh so little they contribute
  practically no mass to the atom.
• 3. The nucleus is positively charged.
• The amount of positive charge balances the
  negative charge of the electrons.
• 4. The electrons are dispersed in the empty space
  of the atom surrounding the nucleus.
• Like water droplets in a cloud.

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Rutherford model of atoms, pictorial presentation
-Most of mass and all the positive charge are
concentrated in a tiny core called the nucleus. -
The outer space contains electrons
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Interpretation of the results
Why most particles pass straight through? Why
some Particles are deflected and bounced back?
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Charge of nuclei

• Rutherford also suggested
• -The smallest positive particle found in
  Goldsteins experiment is a proton.
• Proton has a charge equal in magnitude to that of
  the electron and
• Proton has nearly the same mass of a hydrogen
  atom (ca. 1837 times that of electron).
• -Proton constitute the positive charge of
  nucleus.
• Heavier atoms contains more than one protons.

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But what about the Neutron?

• Except H, atomic nuclei are heavier than
  indicated by the number of positive charges
  (protons). For example,
• Helium has 2 protons in its nucleus.
• But, helium must have a mass of around double
  that.
• Where is the missing mass???????

Rutherford inferred that there are other
particles that are neutral and have approximately
the same mass as a proton. He called these
particles neutrons. The existence of neutrons
was confirmed in the 1930s by James Chadwick .
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Rutherford model of Atom, Summary

• Atoms are composed of three main piecesprotons,
  neutrons, and electrons.
• The nucleus contains protons and neutrons.
• The nucleus is only about 10-13 cm in diameter.
• The electrons move outside the nucleus with an
  average distance of about 10-8 cm.
• Therefore, the radius of the atom is about 105
  times larger than the radius of the nucleus.

• Number of protons in element number of
  electrons atomic number
• For a give element all atoms having the same
  atomic number

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A summary on subatomic particles
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Isotopes
For atoms of a given element, number of protons
number of electrons number of neutrons may not
be equal to number of protons

• Isotopes have the same atomic number but
  different mass number (same number of protons but
  different number of neutrons). e.g. for hydrogen,

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Nuclear Symbol
A mass number number of protons number
of neutrons number of nucleons Z atomic
number X symbol of the element e.g. Isotopes
of hydrogen
Hydrogen-1
Hydrogen-3
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4. The Bohr Model of atoms

• Each atom has a tiny positively charged nucleus,
  negatively charged particles called electrons.
• How electrons are arranged in an atom?

The Bohr Model of atoms a model proposed to
explain the emission spectrum of elements
(especially hydrogen).
Quantum mechanical model modern view on
structures of atoms
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Emission spectrum of elements
Flame Tests Various elements placed in a flame
will give different color of the flame.
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The colors of fireworks are attributable to
different elements
Why different elements give different color in
flame test and firework?
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Atomic Spectra and Bohr Model
Before discussing atomic spectra, Let's consider
the properties of light.
Light is Electromagnetic Radiation
Light has wavelength, frequency, speed

• frequency wavelength
• symbol n (Greek letter nu) l (Greek
  lambda)
• units cycles per sec Hertz distance (nm)

Visible light" corresponds to a wavelength range
of 400 - 700 nanometers (nm) and a color range of
violet through red.
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Continuous Spectra

• When white light or sunlight is passed through a
  prism, it produces a continuous spectrum of
  colors.

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Emission spectrum of atoms

• An emission spectrum of atoms usually shows a
  series of lines, i.e. only certain wavelengths
  are possible

Why lined spectrum?
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Bohr model
In 1913, Niels Bohr proposed a model for the
hydrogen atom which explained the spectrum of a
hydrogen atom based upon the following
assumptions

• The electron in a hydrogen atom travels around
  the nucleus is a circular orbit.
• More like shells than planet orbits.
• Each orbit corresponding to a shell labeled as
  1,2,3.
• Maximum number of electrons in an given shell
  2n2.
• The energy of the electron is directly
  proportional to its distance from the nucleus.
• The farther the electron is from the nucleus the
  more energy it has.

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The Bohr Model of the Atomorbits and energy
levels

• There are a limited number of specific allowed
  energy levels, i.e., the orbits are quantized,
  the energy of electron in an atom is quantized.
• Each orbit has a specific amount of energy.
• The energy of each orbit (shell) is characterized
  by an integer nthe larger the integer, the more
  energy an electron in that orbit has.
• The integer, n, is called a quantum number.
• Each orbit (shell) has a corresponding energy
  (energy level)

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The Bohr Model of the Atomorbits and energy
levels
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Filling orbits with electrons

• Electrons prefer to be in the orbit (shell) of
  lowest energy
• The Maximum number of electrons in an given shell
  2n2.

• Ground state atoms with electrons in the lowest
  possible energy levels
• Excited state atoms with electrons not in the
  lowest possible energy levels

Ground state
Excited state
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Electronic Configuration
Electronic Configuration the arrangement of the
electrons in an atom.
Cl 2, 8, 7

• To give electron configuration of atoms in the
  ground state, add electrons to the lowest shell
  until filled, then go to the next shell.
• Electron configuration can be represented by a
  set of numbers.
• Electron configuration can also be represented
  by a diagram using Dots or Crosses to show
  electrons, and circles to show the shells. For
  example, Cl has 17 e-.

X
X
X
X
X
X
Cl
X
X
X
X
X
X
X
X
X
X
X
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Exercise
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Emission of light by atoms

• Normally electrons are in the lowest energy
  levels, (ground state).
• When a flame or other source of energy is
  absorbed by the electrons, they are promoted to a
  higher energy level (excited state).
• When an electron in an excited state returns to a
  lower energy state, it emits a photon of energy
  which may be observed as light.
• The energy of light absorbed or emitted is equal
  to the difference in the energy levels of the
  orbits between which the electron jumps of falls.
  

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Emission of light by atoms, illustrated in
another way

• Normally electrons are in the lowest energy
  levels, (ground state).
• When a flame or other source of energy is
  absorbed by the electrons, they are promoted to a
  higher energy level (excited state).
• When an electron in an excited state returns to a
  lower energy state, it emits a photon of energy
  which may be observed as light.
• The energy of light absorbed or emitted is equal
  to the difference in the energy levels of the
  orbits between which the electron jumps or falls.

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The Bohr Model of the Atom Hydrogen Spectrum

• Every hydrogen atom has identical orbits, so
  every hydrogen atom can undergo the same energy
  transitions.
• However, since the distances between the orbits
  in an atom are not all the same, no two leaps in
  an atom will have the same energy.
• The closer the orbits are in energy, the lower
  the energy of the photon emitted.
• Lower energy photon longer wavelength.
• Therefore, we get an emission spectrum that has a
  lot of lines that are unique to hydrogen.

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General cases
Why different elements placed in a flame will
give different color of flame?
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The Bohr Model of the AtomSuccess and Failure

• The Bohr model very accurately predicts the
  spectrum of hydrogen.
• However, it had its limitations
• Works only to single-electron atoms.
• Could not account for the intensities or the fine
  structure of the spectral lines.
• Could not explain the binding of atoms into
  molecules.
• A better theory was needed.

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5. A modern model Quantum Model
Also called Electron Cloud Model or Quantum
mechanical Model (1920's)- an atom consists of a
dense nucleus composed of protons and neutrons
surrounded by electrons that exist in different
clouds at the various energy levels.   Erwin
Schrodinger and Werner Heisenburg developed
probability functions to determine the regions or
clouds in which electrons would most likely be
found.
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The Quantum Model

• The quantum model of the atom is a
  probability-based model. It is composed of
  principle energy levels, sublevels, and atomic
  orbitals.
• Electrons are in atomic orbitals

Atomic orbital a region in space representing a
high probability of locating an
electron. Orbitals can have very different shapes
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The Quantum ModelQuantum Numbers

• In Schrödingers wave equation, there are 3
  integers, called quantum numbers, that quantize
  the energy.
• The principal quantum number, n, specifies the
  main energy level for the orbital.

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The Quantum ModelQuantum Numbers, Continued

• Each principal energy shell has one or more
  subshells.
• The number of subshells the principal quantum
  number.
• The quantum number that designates the subshell
  is often given a letter.
• s, p, d, f.
• Each kind of sublevel has orbitals with a
  particular shape.
• The shape represents the probability map.
• 90 probability of finding electron in that
  region.

shell a group of orbitals having the same
n subshell a group of orbitals having the same
n and similar shape
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• Each subshell contains one or more orbitals.
• s subshells have 1 orbital.
• p subshells have 3 orbitals.
• d subshells have 5 orbitals.
• f subshells have 7 orbitals.

s orbital
p orbital
d orbital
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Energy levels of atomic orbitals of atoms
Orbitals with same n have the same energy.
Orbitals in different subshell of same n have
different energy.
The subshells in a shell of H all have the same
energy, but for many electron atoms the subshells
have different energies. s lt p lt d lt f.
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The Bohr Model vs.the Quantum-Mechanical Model

• Both the Bohr and quantum-mechanical models
  predict the spectrum of hydrogen very accurately.
• Only the quantum-mechanical model predicts the
  spectra of multi-electron atoms.

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Fill the orbitals with electrons

• Electrons in atoms are in atomic orbitals. Each
  orbital can have at most two electrons.
• The electron in an atom will fill in orbitals of
  lower energy first.

General sequence 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d
5p
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Each energy shell and subshell has a maximum
number of electrons it can hold. s , p , d
, f . Based on the number of orbitals in
the subshell.

• s subshells have 1 orbital.
• p subshells have 3 orbitals.
• d subshells have 5 orbitals.
• f subshells have 7 orbitals.

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Electron Configurations

• Electron Configurations the arrangement of the
  electrons in an atom, usually given by a listing
  of the subshells in order of filling with the
  number of electrons in that subshell written as a
  superscript.
• Normally refers to ground state

e.g. He, Z 2, 2 electrons
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Electronic configuration
Another example electron configuration of
nitrogen,
N Z 7, 7 electrons
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Exercise 1. What is the electronic configuration
of a Na atom?
Z 11, Na has ___ e?
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• Exercise 2. Write the electronic configurations
  for
• (a) B, Z 5 (b) Ar, Z 18 (c) Fe, Z 26.

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Abbreviated electronic 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.

• He, Z 2, 1s2
• B, Z 5 1s22s22p1 He2s22p1
• Ar, Z 18 1s22s22p63s23p6
• Fe, Z 26 1s22s22p63s23p64s23d6

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6. Periodic Table and electron configuration
groups
Notice the labels of groups
It is composed vertical columns called groups or
families and horizontal rows called periods.
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Periodic Table and electron configuration
Convention used in this course
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Periodic Table and electron configuration
The location of an elements in the Periodic Table
is related to its electron configuration.
Characteristic outmost electronic configurations
(n-1)dxns2
ns2npx
nsx
or ns2(n-1)dx
(n-2)fxns2
or ns2(n-2)fx

• Elements in the same group usually have similar
  electron configuration.

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Classification of elements
s and p block main group elements d block
transition metals f block inner transition
metals (lanthanides and actinides)
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Some groups in the periodic table have special
names

• Alkali Metals Group 1A Alkaline Earth Metals
  Group 2A
• Halogens Group 7A Noble Gases Group 8A

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Write abbreviated Electron Configurations with
the help of periodic Table.
Characteristic electronic configurations
xns2npx (periods 2-3) x(n-1)d10ns2npx
(period 4-5) x(n-2)f10(n-1)d10ns2npx (period
6,7)
x(n-1)dxns2 (periods 4-5) or x(n-2)f10(n-1)dxn
s2 (period 6,7)
xnsx

• For main group elements, Group of e- in ns,
  np orbitals

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Write abbreviated Electron Configurations with
the help of periodic Table. e.g. Zinc
Zn Ar4s23d10
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Exercise. Write abbreviated Electron
Configurations for the following elements (a) Si,
(b) As, (c) Re
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Exceptions in electronic configurations
in Period 4
Not Ar4s23d4
Not Ar4s23d9
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Exceptions in electronic configurations
In Period 5
Not Kr5s24d3
Not Kr5s24d4
Not Kr5s24d6
Not Kr5s24d8
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Electron Configurations and valence electrons

• Valence Electrons
• Valence electrons are the electrons in the
  outmost shell of an atom. e.g.
• These are the electrons that are gained, lost, or
  shared in a chemical reaction.
• Elements in a group or family have the same
  number of valence electrons.

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Valence electrons
For s and p blocks (main group) electrons in ns
np orbitals groups
e.g. N, group 5A ,V.E. 5 I, group 7A, V.E.
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