Atomic Structure and the periodic table

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Atomic Structure and the periodic table

• Chapter 7

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Characterestics Of Light

• Spectroscopy is the study of the interaction
  between radiation (electromagnetic radiation, or
  light, as well as particle radiation) and
  matter.
• Spectrometry is the measurement of these
  interactions and an instrument which performs
  such measurements is a spectrometer or
  spectrograph.
• A plot of the interaction is referred to as a
  spectrum.

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• The spectrum of colors emitted by atoms that have
  been heated to high temperatures is called an
  atomic emission spectrum.
• Light is electromagnetic radiation, a wave of
  electric and magnetic field that travels at the
  speed of light(3.00?108m/s).

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Dispersion of light as it passes through a prism
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• Light waves travel at a constant speed. Because
  of this there is a one to one relationship
  between light's wavelength and its frequency. If
  waves are short, there must be more of them in a
  set amount of time to travel the same distance in
  that time (the same speed).
• The unit of frequency is Hertz(Hz) 1Hz1cycle/sec

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• The intensity of the light is proportional to the
  square of the amplitude of the wave.
• The distance between adjacent peaks gives its
  wavelength,?.
• ?c/?

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• Our eyes detect electromagnetic radiation with
  wavelengths from 700nm(red light) to 400nm(violet
  light). The electromagnetic radiation in this
  range is called as visible light.
• Ultraviolet radiation(UV) has a frequency higher
  than that of violet light its wavelength is less
  than about 400nm.
• Infrared radiation(IR) has a frequency lower than
  that of red light its wavelength is greater than
  about 800nm.
• Microwaves, used in radars and microwave ovens,
  have wavelengths in the millimeter to centimeter
  range.

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Class Practice

• What is the wave length of orange light of
  frequency 4.8?1014Hz?

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Quanta and photon

• Electromagnetic radiation is a stream of many
  packets of electromagnetic energy. These packets
  are called as photons.
• The photon is the elementary particle responsible
  for electromagnetic phenomena.
• The photon travels (in vacuum) at the speed of
  light, c.
• Energy is transferred in discrete amounts called
  as quantum(plural quanta)
• Like all quanta, the photon has both wave and
  particle properties (wave particle duality ).
• The more intense the light the greater the number
  of photons passing a given point each second.
• The relation between the energy , E of a photon
  and the frequency ?, of the radiation as Eh?
  where h is Planks constant, with a value of
  6.63?10-34 J.s

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Class Practice

• What is the energy in kilojoules per mole of
  photons of amber light of frequency 5.2?1014 Hz?
• A students favorite radio station broadcasts a
  signal at a frequency of 91.5MHz. What is the
  energy of the photons broadcast?

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Home work

• Page 314
• 7.12,7.14 and 7.18

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Atomic Spectra and energy levels

• The Lyman series is the wavelengths in the
  ultraviolet (UV) spectrum of the hydrogen atom,
  resulting from electrons dropping from higher
  energy levels into the n 1 orbit.
• The Balmer series is the wavelengths inthe
  visible light spectrum of the hydrogenatom,
  resulting from electrons falling fromhigher
  energy levels into the n 2 orbit.
• The Paschen series is the wavelengths in the
  infrared spectrum of the hydrogenatom, resulting
  from electrons falling fromhigher energy levels
  into the n 3 orbit.

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• Rydberg formula for hydrogen
• Where
• ?vac is the wavelength of the light emitted in
  vacuum, RH is the Rydberg constant for hydrogen.
  

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Bohrs atom

• The Bohr model of the hydrogen atom where
  negatively charged electrons confined to atomic
  shells encircle a small positively charged
  atomic nucleus, and that an electron jump between
  orbits must be accompanied by an emitted or
  absorbed amount of electromagnetic energy h?.
  The orbits that the electrons travel in are shown
  as grey circles their radius increases n2, where
  n is the principal quantum number. The 3?2
  transition depicted here produces the first line
  of the Balmer series, and for hydrogen (Z 1)
  results in a photon of wavelength 656 nm (red).

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Wavelike properties of electrons

• Louis-Victor formulated the de Broglie
  hypothesis, claiming that all matter, not just
  light, has a wave-like nature he related
  wavelength (denoted as ? ),
• and momentum (denoted as p)
• This is a generalization of Einstein's equation
  above, since the momentum of a photon is given by
  pE/c and the wavelength by ?c/f, where c is the
  speed of light in vacuum.

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Class Practice

• Calculate the wavelength of a marble of mass 5.00
  g traveling at 1.00m/s.

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Models Of Atoms

• Our current model of a hydrogen atom was proposed
  by the Austrian scientist Erwin Schrodinger. He
  devised an equation (Schrodingers Equation)that
  enabled him to calculate the shape of the wave
  associated with any particle.
• Quantum theory is based on Schrodinger's
  equation H?E? in which electrons are considered
  as wave-like particles whose "waviness" is
  mathematically represented by a set of wave
  functions ?, obtained by solving Schrodinger's
  equation.

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Atomic orbitals

• The wave function for an electron is so important
  that it is given a special name an atomic
  orbital. Atomic orbitals have characterestic
  energy and shapes. The different shapes are
  identified by different letters.
• An s orbital is a spherical cloud that becomes
  less dense as the distance from the nucleus
  increases. p-orbital is a cloud with two lobes
  on opposite sides of the nucleus. The lobes are
  separated by a nodal plane. The d and f orbital
  have complicated shapes.

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The shapes of different orbitals

• As for p-orbital the opposite shades of color
  denote opposite signs of wave.
• The location of an electron is best described as
  a cloud of probable locations.

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Home work

• Page 314
• 7.30,7.34,7.36

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Energy levels

• Each orbital has a characteristic energy. When
  Schrodinger solved his equation, he found that
  the allowed energies of the electron in a
  hydrogen atom are given by the expression E-hRH
  /n² n1,2
• RH, is called as the Rydbergs constant. The
  experimental value of Rydbergs
  constant3.29?1015 Hz
• Whenever the atom emits a photon of radiation,
  the energy emitted is equal to the difference in
  two of the allowed energy levels. Each transition
  corresponds to a line in the spectrum of atomic
  hydrogen.

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Class Practice

• Calculate the wavelength of a photon emitted by a
  hydrogen atom when an electron makes a transition
  from an orbital with n3 to one with n2.

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Quantum Numbers

• Schrodinger found that each atomic orbital is
  identified by three numbers called quantum
  numbers. The quantum number n is called the
  principal quantum number. The orbital angular
  momentum quantum number, l, specifies the shape
  of the orbital. The magnetic quantum number,ml
  specifies the individual orbital of a particular
  shape.

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• An orbital is specified by three quantum numbers
  orbitals are organized into shells and subshells.
  
• An electron has the property of spin the spin
  is described by the quantum number ms which may
  have one of the two values.

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Class Practice

• Calculate the total number of orbitals in a shell
  with n6.

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Orbital energies

• In a many-electron atom, because of the effects
  of penetration and shielding, s-electrons have a
  lower energy than p-electrons of the same shell
  the order of increasing orbital energies within a
  given shell is sltpltdltf.

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Paulis Exclusion Principle

• No more than two electrons may occupy any given
  orbital. When two electrons do occupy one
  orbital, their spins must be paired.

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Hunds Rule

• If more than one orbital in a subshell is
  available, electrons will fill empty orbitals
  before pairing in one of them.

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• The ground state electron configuration of an
  atom of an element with atomic number Z is
  predicted by adding Z electrons to available
  orbitals so as to obtain the lowest total energy.

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Class Practice

• Predict the ground state electron configuration
  of a sulfur atom and draw the box diagram.
• Predict the ground state electron configuration
  of a magnesium atom.

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Electron configuration of an ion.

• To predict the electron configuration of a cation
  , remove outermost electrons in the order np,
  ns,and(n-1)d
• For an anion, add electrons until the next noble
  gas configuration has been reached.

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Home work

• Page 315
• 7.42,7.44,7.48,

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Periodicity of Atomic Properties

• Atomic radius The atomic radius is defined as
  half the distance between the centers of
  neighboring atoms.
• Atomic radii generally increase down a group and
  decrease from left to right.

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Ionic radius

• The ionic radius of an element is its share of
  the distance between neighboring ions in an ionic
  solid.
• Ionic radii generally increase down a group and
  decrease from left to right across a period.
  Cations are smaller than their parent atoms and
  anions are larger.

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Class Practice

• Arrange each pair of ions in order of increasing
  ionic radius(a)Mg2 and Ca2
• (b)O2- and F-

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Ionization energies

• The ionization energy of an element is the energy
  needed to remove an electron from an atom of the
  element in the gas phase.
• For the first ionization energy we start with a
  neutral atom.

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First and second ionization energies

• For the first ionization energy,I1, we start with
  the neutral atom, for example, for copper,
  Cu(g)?Cu(g) e-(g)
• I1 (energy of Cu e-)-(energy of Cu)
• The experimental value for copper is
  785kJ/mol.The second ionization energy I2 is the
  energy required to remove an electron from a
  singly charged gas phase cation. For copper
• Cu(g)? Cu2(g) e-(g)
• I2(energy of Cu2 e-)-(energy of Cu)For copper
  I2 1958kJ/mol

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• The second ionization energy is always greater
  than the first because once the first electron is
  pulled off, you are now trying to remove the
  second electron from a positively charge ion.
  Because of the electrostatic attraction between
  and -, it is more difficult to pull an electron
  away from a positively charge ion than a neutral
  atom.

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• The first ionization energy is highest for
  elements close to helium and is lowest for
  elements close to cesium. Second ionization
  energies are higher than first ionization
  energies (of the same element). An ionization
  energy is very high if the electron is to be
  expelled from a closed shell.

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Electron affinity

• The electron affinity is a measure of the energy
  change when an electron is added to a neutral
  atom to form a negative ion. For example, when a
  neutral chlorine atom in the gaseous form picks
  up an electron to form a Cl- ion, it releases an
  energy of 349 kJ/mol or 3.6 eV/atom.
• An electron affinity of -349 kJ/mol and this
  large number indicates that it forms a stable
  negative ion.
• Small numbers indicate that a less stable
  negative ion is formed. Groups VIA and VIIA in
  the periodic table have the largest electron
  affinities.

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• Elements with the highest electron affinities are
  those close to oxygen, fluorine, and chlorine.
  Group 17 atoms can acquire one electron with a
  release of energy and group 16 atoms can accept
  two electrons with chemically attainable
  energies. The halogens typically form X- ions and
  group 16 elements form X2- ions.

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Diagonal relationship

• On moving diagonally across the periodic table,
  the elements show certain similarities in their
  properties which are quite prominent in some
  cases as shown below. This is called a diagonal
  relationship.

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• A Diagonal Relationship is said to exist between
  certain pairs of diagonally adjacent elements in
  the second and third periods of the periodic
  table. These pairs (Li Mg, Be Al, B Si
  etc.) exhibit similar properties for example,
  Boron and Silicon are both semiconductors form
  halides that are hydrolysed in water and have
  acidic oxides.
• Such a relationship occurs because crossing and
  descending the periodic table have opposing
  effects.
• On crossing a period of the periodic table, the
  size of the atoms decreases, and on descending a
  group the size of the atoms increases. Similarly,
  on moving along the period the elements become
  progressively more covalent, less reducing and
  more electronegative whereas on descending the
  group the elements become more ionic, more basic
  and less electronegative.

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• Thus, on both descending a group and crossing by
  one element the changes cancel each other out,
  and elements with similar properties which have
  similar chemistry are often found - the atomic
  size, electronegativity, properties of compounds
  (and so forth) of the diagonal members are
  similar.

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Homework

• Page 315
• 7.62,7.64,
• Page 316
• 7.70,7.78,7.80,7.90