Atomic Structure

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Atomic Structure

• Unit II

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I. Historical Information
a) Democritus Greek philosopher conceptualized
the idea of the atom atomos means indivisible.
b) Dalton (1808) SOLID (HARD) SPHERE MODEL -
developed an early model of the atom he
described the atom as being a solid, indivisible
particle.
Dalton's Atomic Theory

• All elements are composed of atoms, which are
  indivisible and indestructible particles
  (Lavoisier).

2) All atoms of the same element are exactly
alike in particular, they all have the same mass
(Proust).
3) All atoms of different elements are different
in particular, they have different masses
(Proust).
4) Compounds are formed by the joining of atoms
of two or more elements.  In any compound , the
atoms of the different elements in the compound
are joined in a definite whole-number ratio, such
as 1 to 1, 2 to 1, 3 to 2, etc.
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c) Joseph John Thomson (1897)- Plumb Pudding
Model -

• proposed that the atom was a sphere of positive
  electricity (which was diffuse) with negative
  particles imbedded throughout after discovering
  the electron, a discovery for which he was
  awarded the Nobel Prize in physics in 1906.

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• Thomson used cathode ray tubes to characterize
  electrons.

• The rays are made up of electrons very small,
  negatively charged particles that are indeed
  fundamental parts of every atom.

http//www.chem.uiuc.edu/clcwebsite/video/Cath.mo
v
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d) Rutherford (1909) Solar (Planetary) System
Model - performed a set of experiments with
alpha particles and gold foil. He discovered
that atoms mostly consist of empty space and have
a small, dense, positively charged center
(nucleus).
Figure 1 Rutherfords Gold Foil Experiment

• Radioactive material releases alpha particles
  (a-particles) which have a net positive charge.

• The alpha particles were directed at a thin
  sheet of gold foil.

• Most of the alpha particles passed directly
  through the foil.

• Some of the alpha particles were deflected by
  the gold foil.

http//www.mhhe.com/physsci/chemistry/essentialche
mistry/flash/ruther14.swf
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Results of Rutherfords Gold Foil Experiment

• Rutherford concluded that

1. Since most of the positively charged alpha
particles passed directly through the foil, most
of the atom consists of empty space.
2. Since some of the positively charged alpha
particles were deflected by the foil, the gold
atom must consist of a small, dense, positively
charged center that is called the nucleus.
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FIGURE 2 Enlarged cross-section of the gold
foil in the apparatus, showing the deflection of
alpha particles by the nuclei of the gold atoms.
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e) Niels Bohr (1913) Electron-shell Model -
described the position of negatively charged
electrons in relation to the nucleus as a
function of their energy.

• Bohr stated that electrons are found in areas
  surrounding the nucleus called principal energy
  levels.

• The electrons in energy levels that are close to
  the nucleus of the atom have little free energy
  while those further from the nucleus have greater
  free energy.

• Lower level electrons can jump to higher
  levels by gaining energy. These are called
  excited electrons.

• Excited electrons will return to their
  original level (ground state) while giving off
  energy.

Figure 2 The Bohr Model of the atom
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f) The Electron Cloud Model (wave-mechanical or
quantum-mechanical model)

• 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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II. Subatomic Particles
a) Nucleons subatomic particles that are
located in an atoms nucleus includes protons
and neutrons.

• Proton positively charged has a mass of 1 amu.

• Neutron no charge has a mass of 1 amu.

b) Electrons small, negatively charged
subatomic particles that revolve around the
atoms nucleus at extremely high velocities.

• Electrons are located in energy levels around
  the nucleus.

• Electrons have a mass of 1/1836 amu.

Note Atomic Mass Unit amu or m
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KEY P proton N neutron e- electron
Figure 3 Diagram of Structure of a Carbon Atom
From http//www.spacesciencegroup.org/lessons/g
raphics/partsoftheatom.jpg
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III. Atomic Mass Unit (amu or m)
a) The amu is a unit used to describe the mass of
atoms.
b) Typical units of measurement for mass (grams)
are too large to describe that of the atom.

• For example, 1 amu has a mass equal to 1.66 x
  10-27 kg.

c) The basis for an amu is an atom of carbon.
d) The functional definition of an amu is 1/12
the mass of carbon-12.

• The 12 in C-12 represents the atomic mass of
  the element.

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IV. Atomic Mass and Atomic Number
a) Atomic Mass indicates the number of protons
plus neutrons in an atoms nucleus.

• The atomic mass for a given atom can vary.

• Atoms of the same element that have varying
  atomic masses are known as isotopes.

• Isotopes vary in the number of neutrons found
  within their nucleus.
• Ex 12C 13C 14C
• 6 6 6

Naturally Occurring Isotopes of Carbon

• Larger Number is the Atomic Mass / Smaller
  Number is the Atomic Number

p6 p6 p6
n6 n7
n8 e6 e6
e6
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b) Atomic Number
1. The atomic number of an elements indicates
the number of protons found in an atoms nucleus.
2. The atomic number for a given atom never
changes.
3. The atomic number is the identification tag of
a given element.
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V. Electron Configuration The Arrangement of
Electrons Around the Nucleus of the Atom
a) Electrons are arranged around an atoms
nucleus in energy levels.
b) The energy level that an atom belongs
determines the amount of free energy it has.
c) Electrons located in energy levels close to
the nucleus have little free energy. Electrons
that exist in energy levels that are far from the
nucleus have great amounts of free energy.
d) Electrons found in their ground state
(original energy level) can move to the excited
state (higher energy levels) by gaining energy.
Electrons in the excited state will return to
their ground state while losing a very specific
quantity (quanta) of light energy (photons).
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• e) Energy Level and Chemical Stability
• Each energy level can hold a certain amount of
  electrons.
• For example
• n energy level
• n Max. of electrons
• 1 2
• 2 8
• 3 18

2n2 maximum of electrons found in energy
level.

• Note For atoms to be stable they must have a
  complete outer energy level. Meaning they must
  contain the maximum number of possible electrons
  in their outermost energy level.
• The maximum number of electrons that the third
  energy level can hold is 18. However, for an
  atom to be stable it only needs 8 electrons in
  the third energy level.

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Electrons are located in energy levels, shells or
principal quantum numbers (n). There are 7
energy levels and each level can accommodate only
a limited number of electrons according to the
formula 2n2 where n energy level. Energy
level 1 can have a maximum of 2 e- 2(12)
2 Energy level 2 can have a maximum of  8 e-
2(22) 8 Energy level 3 can have a maximum of
18 e- 2(32) 18 Energy level 4 can have  a
maximum of 32 e- 2(42) 32 etc.
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f) Valence Electrons and Chemical Stability

• Valence electrons are those found in an atoms
  outer most energy level.

• Electrons can be gained, lost, or shared between
  atoms to complete their respective valence
  shells.

• When atoms have incomplete valence shells they
  are unstable.

• When an atom is unstable, it implies that it has
  a high energy content.

• When an atom completes its valence shell, energy
  is lost and the stability of the atom increases.

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g) Lewis Electron-Dot Structure

• Indicates the number of valence electrons that
  an atom has.
• For example
• 23Na p11
• 11 n12
• e11

2-8-1
of e in 3rd energy level
of e- in 1st energy level
of e- in 2nd energy level
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Lewis Electron-Dot Structure of Sodium
Na .
The one dot next to the chemical symbol of sodium
indicates that it contains only one valence
electron.
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35Cl p 17 17 n 18 e 17
2 - 8 - 7
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VI. Ions
a) Definition Ions are charged particles that
form when atoms either gain or lose electrons.

• When an atom gains an electron(s), it becomes a
  negatively charged ion (anion).

• When an atom loses an electron(s), it becomes a
  positively charged ion (cation).

Ex 1 Na Na 1e-
Sodium atom
Sodium Ion
Ex 2 1e- Cl Cl-
Chlorine atom
Chlorine ion
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Ex 1 23Na Element Sodium
11 p11 2-8-1
n12
e11
First Energy Level 2 electrons
11p 12n
Second Energy Level 8 electrons
Third Energy Level 1 electron
THE THIRD ENERGY LEVEL ONLY HAS ONE ELECTRON!!
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Ex 2 35Cl Element Chlorine
17 p17
2-8-7 n18
e17
First Energy Level 2 electrons
17p 18n
Second Energy Level 8 electrons
THE THIRD ENERGY LEVEL HAS SEVEN ELECTRONS!!
Third Energy Level 7 electrons
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Question Why do atoms of sodium and chlorine
form bonds?
Sodium Chlorine
Sodium gives an electron to chlorine. Sodium
becomes a cation while chlorine becomes an anion.
A bond is formed that connects the two atoms
called an ionic bond.
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Na Cl-
Ionic Bond
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Do Now 1. Which ion has the same electron
configuration as an atom of He? a) H- b) O-2 c)
Na2 d) Ca2
2. Compared to an element of calcium-40, an atom
of potassium-39 contains fewer a) protons b)
neutrons c) occupied sublevels d) occupied
principal energy levels
3. Which electron configuration is correct for a
sodium ion? a) 2-7 b) 2-8 c) 2-8-1 d) 2-8-2
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4. If 75.0 of the isotopes of an element have a
mass of 35.0 amu and 25.0 of the isotopes have a
mass of 37.0 amu, what is the atomic mass of the
element? a) 35.0 amu b) 35.5 amu c) 36.0 amu
d) 37.0 amu
5. Which Lewis electron-dot structure is drawn
correctly for the atom it represents?
1 2 3 4
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VII. Ground versus Excited State

• Electrons have the ability to move between
  energy levels.

• The ground state refers the energy level that
  a given electron usually belongs to.

• The excited state refers to a higher energy
  level that a given electron moves to as it gains
  energy.

• Electrons that have gained energy and jumped
  to the excited state will eventually return to
  the ground state and release its energy in the
  form of a photon (light energy).

• As a result of this process, different atoms
  will give off (emit) unique (characteristic)
  colors (bands) of light. This is called the
  Bright Line Spectrum.

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• Base your answers to the following questions on
  the diagram below, which shows bright-line
  spectra of selected elements.
•                                                  
                            
• Identify the two elements in the unknown
  spectrum.
• Explain how a bright-line spectrum is produced,
  in terms of excited state and ground state.

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Example When subjected to a flame, solutions
containing certain metals have characteristic
colors corresponding to the energy released when
excited electrons return to lower energy levels.
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From http//www.chemcool.com/regents/atomicconce
pts/aim3.htm
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VIII. Advanced Electron Configuration
a) In addition to describing the location of
electrons around the atoms nucleus in terms of
their principal energy levels, it can be defined
in terms of sublevels and orbitals.
b) Sublevels

• Each principal energy level has one or more
  sublevels.

• The sublevels are designated by the following
  symbols
• s, p, d, and f.

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c) Each sublevel has a characteristic shape and a
specific number of orbitals.

• The maximum number of electrons per orbital is
  TWO.

1. s- sublevel

• shape circular
• of orbitals 1
• maximum number of electrons 2

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2. p-sublevel

• shape dumbbells
• of orbitals 3
• maximum number of electrons 6

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3. d-sublevel

• shape complex
• of orbitals 5
• maximum number of electrons 10

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4. f-sublevel

• shape complex
• of orbitals 7
• maximum number of electrons 14

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n sublevels (l) of orbitals of electrons
1 s 1 2 2 s,
p 4 8 3 s, p, d 9 18 4
s, p, d, f 16 32
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e) Rules of Sublevel and Orbital Filling

• electrons will always be placed into energy
  levels of lowest energy.

1. Electron filling for elements with atomic
numbers of 18 and below

• The lower energy levels (ones closer to the
  atoms nucleus) must be completely filled before
  electrons can enter higher energy levels.
• Ex The first energy level (n 1) must be
  filled before electrons are placed into the
  second energy level (n2).

• Sublevels of lower energy (ones closer to the
  atoms nucleus) must be filled before electrons
  can enter those of higher energy.
• Ex The s-sublevel must be filled before the
  p-sublevel.

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Practice Problem 1
Electron Configuration
Sodium
23 Na
P N e
11 12 11
Basic
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2 -
8 -
1
First Energy Level
Third Energy Level
Second Energy Level
Orbital Notation
1s2 2s2 2p6 3s1

• The superscripts represent the number of
  electrons in the respective sublevels.

s-sublevel
p-sublevel
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Practice Problem 2
Electron Configuration
35Cl
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Basic
2 8 - 7
Orbital Notation
1s2 2s2 2p6 3s2 3p5
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2. Rules for Electron filling for elements with
atomic numbers of 19 and above
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http//www.chemcool.com/regents/atomicconcepts/ato
micconcepts.htm