Question of the Day

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Question of the Day

• Draw a sketch of what you think an atom looks
  like. Be prepared to describe your drawing to the
  class.

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Chapter 11 Modern Atomic Theory

• The History of Atomic Theory

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Atomic Models

• This model of the atom may look familiar to you.
  This is the Bohr model. In this model, the
  nucleus is orbited by electrons, which are in
  different energy levels.
• A model uses familiar ideas to explain unfamiliar
  facts observed in nature.
• A model can be changed as new information is
  collected.

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A Very Early Model

• The atomic model has changed throughout the
  centuries, starting in 400 BC, when it looked
  like a billiard ball ?

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Who are these men?
In this lecture, well learn about the men whose
quests for knowledge about the fundamental nature
of the universe helped define our views of the
atomic world.
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Democritus
400 BC

• This is the Greek philosopher Democritus who
  began the search for a description of matter more
  than 2400 years ago.
• He asked Could matter be divided into smaller
  and smaller pieces forever, or was there a limit
  to the number of times a piece of matter could be
  divided?

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Atomos

• His theory Matter could not be divided into
  smaller and smaller pieces forever, eventually
  the smallest possible piece would be obtained.
• This piece would be indivisible.
• He named the smallest piece of matter atomos,
  meaning not to be cut.

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Atomos

• To Democritus, atoms were small, hard particles
  that were all made of the same material but were
  different shapes and sizes.
• Atoms were infinite in number, always moving and
  capable of joining together.

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• This theory was ignored and forgotten for
  more than 2000 years!

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Why?

• The eminent philosophers of the time, Aristotle
  and Plato, had a more respected, (and ultimately
  wrong) theory.

Aristotle and Plato favored the earth, fire, air
and water approach to the nature of matter. Their
ideas held sway because of their eminence as
philosophers. The atomos idea was buried for
approximately 2000 years.
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Daltons Model

• In the early 1800s, the English Chemist John
  Dalton performed a number of experiments that
  eventually led to the acceptance of the idea of
  atoms.

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Daltons Theory

• He deduced that all elements are composed of
  atoms. Atoms are indivisible and indestructible
  particles.
• Atoms of the same element are exactly alike.
• Atoms of different elements are different.
• Compounds are formed by the joining of atoms of
  two or more elements.

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.

• This theory became one of the foundations of
  modern chemistry.

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Thomsons Plum Pudding Model

• In 1897, the English scientist J.J. Thomson
  provided the first hint that an atom is made of
  even smaller particles.

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Thomson Model

• He proposed a model of the atom that is sometimes
  called the Plum Pudding model.
• Atoms were made from a positively charged
  substance with negatively charged electrons
  scattered about, like raisins in a pudding.

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Thomson Model

• Thomson studied the passage of an electric
  current through a gas.
• As the current passed through the gas, it gave
  off rays of negatively charged particles.

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Thomson Model
Where did they come from?

• This surprised Thomson, because the atoms of the
  gas were uncharged. Where had the negative
  charges come from?

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Thomson concluded that the negative charges came
from within the atom. A particle smaller than
an atom had to exist. The atom was divisible!

• Thomson called the negatively charged
  corpuscles, today known as electrons.
• Since the gas was known to be neutral, having no
  charge, he reasoned that there must be positively
  charged particles in the atom.
• But he could never find them.

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Rutherfords Gold Foil Experiment

• In 1908, the English physicist Ernest Rutherford
  was hard at work on an experiment that seemed to
  have little to do with unraveling the mysteries
  of the atomic structure.

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• Rutherfords experiment Involved firing a stream
  of tiny positively charged particles at a thin
  sheet of gold foil (2000 atoms thick)

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• Most of the positively charged bullets passed
  right through the gold atoms in the sheet of gold
  foil without changing course at all.
• Some of the positively charged bullets,
  however, did bounce away from the gold sheet as
  if they had hit something solid. He knew that
  positive charges repel positive charges.

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• http//chemmovies.unl.edu/ChemAnime/RUTHERFD/RUTHE
  RFD.html

• http//chemmovies.unl.edu/ChemAnime/RUTHERFD/RUTHE
  RFD.html

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• This could only mean that the gold atoms in the
  sheet were mostly open space. Atoms were not a
  pudding filled with a positively charged
  material.
• Rutherford concluded that an atom had a small,
  dense, positively charged center that repelled
  his positively charged bullets.
• He called the center of the atom the nucleus
• The nucleus is tiny compared to the atom as a
  whole.

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Rutherford

• Rutherford reasoned that all of an atoms
  positively charged particles were contained in
  the nucleus. The negatively charged particles
  were scattered outside the nucleus around the
  atoms edge.

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Bohr Model

• In 1913, the Danish scientist Niels Bohr proposed
  an improvement. In his model, he placed each
  electron in a specific energy level.

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Bohr Model

• According to Bohrs atomic model, electrons move
  in definite orbits around the nucleus, much like
  planets circle the sun. These orbits, or energy
  levels, are located at certain distances from the
  nucleus.

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Wave Model
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The Wave Model

• Todays atomic model is based on the principles
  of wave mechanics.
• According to the theory of wave mechanics,
  electrons do not move about an atom in a definite
  path, like the planets around the sun.

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The Wave Model

• In fact, it is impossible to determine the exact
  location of an electron. The probable location of
  an electron is based on how much energy the
  electron has.
• According to the modern atomic model, at atom has
  a small positively charged nucleus surrounded by
  a large region in which there are enough
  electrons to make an atom neutral.

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Electron Cloud

• A space in which electrons are likely to be
  found.
• Electrons whirl about the nucleus billions of
  times in one second
• They are not moving around in random patterns.
• Location of electrons depends upon how much
  energy the electron has.

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Electron Cloud

• Depending on their energy they are locked into a
  certain area in the cloud.
• Electrons with the lowest energy are found in the
  energy level closest to the nucleus
• Electrons with the highest energy are found in
  the outermost energy levels, farther from the
  nucleus.

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Indivisible Electron Nucleus Orbit Electron Cloud
Greek X
Dalton X
Thomson X
Rutherford X X
Bohr X X X
Wave X X X
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Question of the Day

• Name one individual from yesterdays lecture and
  describe his version of the atom.

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Sec. 11.2 Energy Light

• Electromagnetic Radiation
• Energy being transmitted from one place to
  another by light
• Characterized by wavelength (? lambda) the
  distance between two consecutive peaks
• Frequency of a wave (v nu) how many waves pass
  a given point during a time period
• Speed how fast a given peak travels

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Electromagnetic Radiation

• Waves have a frequency
• Use the Greek letter nu, ?, for frequency, and
  units are cycles per sec
• All radiation ? ? c where c velocity
  of light 3.00 x 108 m/sec

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Electromagnetic Radiation

• Most subatomic particles behave as PARTICLES and
  obey the physics of waves.

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Electromagnetic Spectrum

• Long wavelength --gt small frequency
• Short wavelength --gt high frequency

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Electromagnetic Spectrum
In increasing energy, ROY G BIV
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Electromagnetic Radiation
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Quantum mechanics

• Black body radiation was explained by Planck ,
  chuncks of energy
• quanta
• E hn h 6.63 x 10 -34
• lh/p p is the light momentum
• c ln c is the speed of light

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Electromagnetic radiation.
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Sec. 11.3 Emission of Energy by Atoms

• When atoms receive energy from a source they
  become excited.
• They release that energy in the form of light
  it is carried away by a photon
• The photons of red light carry less energy than
  photons of blue light because red light has a
  longer wavelength than blue does.

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11.4 Energy Levels and Line Spectra

• Lowest possible energy state for an atom is
  ground state
• When atoms absorb energy they enter an excited
  state
• As they return to a lower energy state
• Atoms emit light in quantized amounts.

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An excited lithium atom emitting a photon of red
light to drop to a lower energy state.
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An excited H atom returns to a lower energy level.
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Line Emission Spectra of Excited Atoms

• Excited atoms emit light of only certain
  wavelengths
• The wavelengths of emitted light depend on the
  element.

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Spectrum of Excited Hydrogen Gas
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Light Spectrum Lab!
Slit that allows light inside
Line up the slit so that it is parallel with the
spectrum tube (light bulb)
Scale
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Light Spectrum Lab!
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Line Spectra of Other Elements
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Sec. 11.7 Orbitals

• Orbitals area of probability that an electron
  can be found
• An electron can be found in this area 90 of the
  time, 10 elsewhere
• Many times graphics of orbitals are shown with
  fuzzy edges

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Arrangement of Electrons in Atoms in Energy Levels

• Electrons in atoms are arranged as
• LEVELS (n)
• SUBLEVELS (l)
• ORBITALS (ml)

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Energy Levels

• Each energy level has a number called the
  PRINCIPAL QUANTUM NUMBER, n
• Currently n can be 1 thru 7, because there are 7
  periods on the periodic table

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Energy Levels
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QUANTUM NUMBERS

• The shape, size, and energy of each orbital is a
  function of 3 quantum numbers which describe the
  location of an electron within an atom or ion
• n (principal) ---gt energy level
• l (orbital) ---gt shape of orbital
• ml (magnetic) ---gt designates a particular
  suborbital
• (The fourth quantum number is not derived from
  the wave function)
• s (spin) ---gt spin of the electron (clockwise
  or counterclockwise ½ or ½)

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QUANTUM NUMBERS

• So if two electrons are in the same place at the
  same time, they must be repelling, so at least
  the spin quantum number is different!
• The Pauli Exclusion Principle says that no two
  electrons within an atom (or ion) can have the
  same four quantum numbers.
• If two electrons are in the same energy level,
  the same sublevel, and the same orbital, they
  must repel.
• Think of the 4 quantum numbers as the address of
  an electron Country gt State gt City gt Street

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Types of Orbitals

• The most probable area to find these electrons
  takes on a shape
• So far, we have 4 shapes. They are named s, p,
  d, and f.
• No more than 2 e- assigned to an orbital one
  spins clockwise, one spins counterclockwise

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Orbitals and the Periodic Table

• Orbitals grouped in s, p, d, and f orbitals
  (sharp, proximal, diffuse, and fundamental)

s orbitals
d orbitals
p orbitals
f orbitals
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Types of Orbitals (l)
s orbital
p orbital
d orbital
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Relative sizes of the spherical 1s, 2s, and 3s
orbitals
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p Orbitals

• this is a p sublevel with 3 orbitals
• These are called x, y, and z

There is a PLANAR NODE thru the nucleus, which is
an area of zero probability of finding an electron
3py orbital
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p Orbitals

• The three p orbitals lie 90o apart in space

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d Orbitals

• d sublevel has 5 orbitals

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The shapes and labels of the five 3d orbitals.
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f Orbitals

• For l 3, ---gt f sublevel with 7 orbitals

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Electron Configurations

• A list of all the electrons in an atom (or
  ion)..
• Must go in order (Aufbau principle)
• 2 electrons per orbital, maximum
• We need electron configurations so that we can
  determine the number of electrons in the
  outermost energy level. These are called valence
  electrons.
• The number of valence electrons determines how
  many and what this atom (or ion) can bond to in
  order to make a molecule

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
4f14 etc.
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Diagonal Rule

• Steps
• Write the energy levels top to bottom.
• Write the orbitals in s, p, d, f order. Write
  the same number of orbitals as the energy level.
• Draw diagonal lines from the top right to the
  bottom left.
• To get the correct order, follow the arrows!

1 2 3 4 5 6 7
s
s 2p
s 3p 3d
s 4p 4d 4f
By this point, we are past the current periodic
table so we can stop.
s 5p 5d 5f 5g?
s 6p 6d 6f 6g? 6h?
s 7p 7d 7f 7g? 7h? 7i?
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Why are d and f orbitals always in lower energy
levels?

• d and f orbitals require LARGE amounts of energy
• Its better (lower in energy) to skip a sublevel
  that requires a large amount of energy (d and f
  orbtials) for one in a higher level but lower
  energy
• This is the reason for the diagonal rule! BE SURE
  TO FOLLOW THE ARROWS IN ORDER!

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Electron Configurations

• 2p4

Number of electrons in the sublevel
Energy Level
Sublevel
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
4f14 etc.
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Lets Try It!

• Write the electron configuration for the
  following elements
• H
• Li
• N
• Ne
• K
• Zn
• Pb

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Shorthand Notation

• A way of abbreviating long electron
  configurations
• Since we are only concerned about the outermost
  electrons, we can skip to places we know are
  completely full (noble gases), and then finish
  the configuration

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Shorthand Notation

• Step 1 Find the closest noble gas to the atom
  (or ion), WITHOUT GOING OVER the number of
  electrons in the atom (or ion). Write the noble
  gas in brackets .
• Step 2 Find where to resume by finding the next
  energy level.
• Step 3 Resume the configuration until its
  finished.

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Shorthand Notation

• Chlorine
• Longhand is 1s2 2s2 2p6 3s2 3p5
• You can abbreviate the first 10 electrons with a
  noble gas, Neon. Ne replaces 1s2 2s2 2p6
• The next energy level after Neon is 3
• So you start at level 3 on the diagonal rule (all
  levels start with s) and finish the configuration
  by adding 7 more electrons to bring the total to
  17
• Ne 3s2 3p5

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Practice Shorthand Notation

• Write the shorthand notation for each of the
  following atoms
• Cl
• K
• Ca
• I
• Bi

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Question of the Day

• Which of the following orbital designations is
  (are) not correct? Why?
• A. 2p
• B. 1d
• C. 3f
• D.4s

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Electrons Valence

• Electrons are divided between core and valence
  electrons
• B 1s2 2s2 2p1
• Core He , valence 2s2 2p1

Br Ar 3d10 4s2 4p5 Core Ar 3d10 ,
valence 4s2 4p5
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Rules of the Game

• No. of valence electrons of a main group atom
  Group number (for A groups)

Atoms like to either empty or fill their
outermost level. Since the outer level contains
two s electrons and six p electrons (d f are
always in lower levels), the optimum number of
electrons is eight. This is called the octet
rule.
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Keep an Eye On Those Ions!

• Electrons are lost or gained like they always are
  with ions negative ions have gained electrons,
  positive ions have lost electrons
• The electrons that are lost or gained should be
  added/removed from the highest energy level (not
  the highest orbital in energy!)

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Keep an Eye On Those Ions!

• Tin
• Atom Kr 5s2 4d10 5p2
• Sn2 ion Kr 5s2 4d10
• Note that the electrons came out of the highest
  energy level, not the highest energy orbital!

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Keep an Eye On Those Ions!

• Bromine
• Atom Ar 4s2 3d10 4p5
• Br- ion Ar 4s2 3d10 4p6
• Note that the electrons went into the highest
  energy level, not the highest energy orbital!

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Try Some Ions!

• Write the longhand notation for these
• F-
• Li
• Mg2
• Write the shorthand notation for these
• Br-
• Ba2
• Al3

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Orbital Box Diagrams

• Graphical representation of an electron
  configuration
• One arrow represents one electron
• Shows spin and which orbital within a sublevel
• Same rules as before (Aufbau principle, d4 and d9
  exceptions, two electrons in each orbital, etc.
  etc.)

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Orbital Diagrams

• One additional rule Hunds Rule
• In orbitals of EQUAL ENERGY (p, d, and f), place
  one electron in each orbital before making any
  pairs
• All single electrons must spin the same way
• I nickname this rule the Monopoly Rule
• In Monopoly, you have to build houses EVENLY.
  You can not put 2 houses on a property until all
  the properties has at least 1 house.

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Lithium

• Group 1A
• Atomic number 3
• 1s22s1 ---gt 3 total electrons

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Carbon

• Group 4A
• Atomic number 6
• 1s2 2s2 2p2 ---gt
• 6 total electrons

Here we see for the first time HUNDS RULE. When
placing electrons in a set of orbitals having the
same energy, we place them singly as long as
possible.
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Draw these orbital diagrams!

• Oxygen (O)
• Chromium (Cr)
• Mercury (Hg)

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Question of the Day

• Draw the orbital box notation for Ga

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Sec. 11.11 Periodic Trends

• Atomic and ion size
• Ionization energy
• Electronegativity

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Atomic Size

• Size goes UP - going down a group.
• Because electrons are added further from the
  nucleus, there is less attraction. This is due to
  additional energy levels and the shielding
  effect. Each additional energy level shields
  the electrons from being pulled in toward the
  nucleus.
• Size goes DOWN - going across a period.

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Atomic Size

• Size decreases across a period owing to increase
  in the positive charge from the protons. Each
  added electron feels a greater and greater
  charge because the protons are pulling in the
  same direction, where the electrons are scattered.

Large
Small
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Which is Bigger?

• Na or K ?
• Na or Mg ?
• Al or I ?

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Ion Sizes
Does the size go up or down when losing an
electron to form a cation?

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Ion Sizes
Forming a cation.
Li,152 pm
3e and 3p

• CATIONS are SMALLER than the atoms from which
  they come.
• The electron/proton attraction has gone UP and so
  size DECREASES.

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Ion Sizes

• Does the size go up or down when gaining an
  electron to form an anion?

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Ion Sizes
Forming an anion.

• ANIONS are LARGER than the atoms from which they
  come.
• The electron/proton attraction has gone DOWN and
  so size INCREASES.
• Trends in ion sizes are the same as atom sizes.

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Trends in Ion Sizes
Figure 8.13
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Which is Bigger?

• Cl or Cl- ?
• K or K ?
• Ca or Ca2 ?
• I- or Br- ?

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Ionization Energy
IE energy required to remove an electron from
an atom (in the gas phase).

• Mg (g) 738 kJ ---gt Mg (g) e-
• This is called the FIRST ionization energy
  because we removed only the OUTERMOST electron

Mg (g) 1451 kJ ---gt Mg2 (g) e- This is
the SECOND IE.
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Trends in Ionization Energy

• IE increases across a period because the positive
  charge increases.
• Metals lose electrons more easily than nonmetals.
• Nonmetals lose electrons with difficulty (they
  like to GAIN electrons).

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Which has a higher 1st ionization energy?

• Mg or Ca ?
• Al or S ?
• Cs or Ba ?

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Electronegativity, ?

• ? is a measure of the ability of an atom in a
  molecule to attract electrons to itself.

Concept proposed by Linus Pauling 1901-1994
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Periodic Trends Electronegativity

• In a group Atoms with fewer energy levels can
  attract electrons better (less shielding). So,
  electronegativity increases UP a group of
  elements.
• In a period More protons, while the energy
  levels are the same, means atoms can better
  attract electrons. So, electronegativity
  increases RIGHT in a period of elements.

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Electronegativity
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Which is more electronegative?

• F or Cl ?
• Na or K ?
• Sn or I ?

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The End !!!!!!!!!!!!!!!!!!!