Atomic Structure

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Atomic Structure

• Mrs. Daniels Chemistry .2
• September 2002 Revised August 2006

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IN YOUR JOURNALS

• Describe what you think atoms are made of.
• Can you see an atom with your naked eye or under
  a microscope?
• Draw a picture of a simple atom and its
  components.

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Daltons Atomic Theory

• All elements are composed of tiny indivisible
  particles called atoms
• Atoms of the same element are identical to each
  other are different from those of other
  elements
• Atoms can physically mix together or chemically
  combine in whole number ratios to form compounds.
• Chemical reactions occur when atoms are
  separated, joined, or rearranged.

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How BIG or how small is an atom?

• Is a penny big or small?
• Imagine grinding up the penny into copper dust.
  Is each grain of copper big or small?
• A pure copper penny contains approximately 2.4 x
  1022 individual copper atoms
• Can you imagine dividing the penny into
  24,000,000,000,000,000,000,000 different parts?

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Cutting it down to size activity

• Pair up with the person sitting next to you
• You and your partner will be given two items a
  pair of scissors and a piece of paper
• Your task predict how many times you and your
  partner can cut this piece of paper in half with
  the scissors
• Each time, discard (set aside and well discard
  at the end) one half of the paper youve cut and
  continue on with one piece

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Cutting it down to size activity

• How many times were you able to cut the paper in
  half?
• Which pair was able to make the most cuts?
• How many times would you have to divide this
  original 8 1/2 x 11 piece of paper in order to
  get it to be the width of one atom?
• 31

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Atomic Structure

• First of all, an atom has no overall electric
  charge.
• Secondly, we know that an equal number of
  negative and positive particles combine to form a
  neutral particle
• Keeping this in mind, lets look at three
  subatomic particles.

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Atomic Structure

• Proton (p) A positively charged particle found
  in the central core of an atom (called the
  nucleus)
• Neutron (n0) A neutral particle found in the
  nucleus of an atom
• Electron (e-) A tiny negatively charged
  particle found outside of the atomic nucleus

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Atomic Structure

• The mass of a proton and a neutron is relatively
  equal
• However, the electron has a mass equal to 1/1840
  of a proton
• Which subatomic particles do you think take up
  the most space in the atom?

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Atomic Structure

• If we cant see these subatomic particles, how do
  we know they exist?

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Before we answer that

• The atoms were discussing each have a
  representative symbol as you should recall.
• See how many of these elements you can discover
  in the symbol recognition worksheet.
• Homework tonight
• Read pages 55-61

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J.J. Thomsons Cathode Ray Tube

• 1897 was a big year for J.J. He discovered the
  electron.
• Thomson passed electric current through gases
  under low pressure in a sealed glass tube
• At one end of the glass tube was an electrode
  with a positive charge (anode) and at the other
  end was a negative electrode (cathode)
• When electricity was passed through, a glowing
  beam formed between the cathode and the anode.
  This beam was then called a cathode ray.

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J.J. Thomsons Cathode Ray Tube

• Thomson found that the cathode ray was attracted
  to a positively charged metal plate
• Knowing that opposites attract, he concluded that
  the beam (cathode ray) is made up of tiny
  negatively charged particles moving at high speed
• These particles were called electrons

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Other Scientists

• Millikan measured the charge of an electron
• Moseley used an X-ray to determine the number of
  protons in an atom
• Rutherford used a gold foil experiment to
  determine that most of the atom is empty space
  and the tiny center of the atom is positively
  charged
• Chadwick demonstrated the existence of neutrons

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Organization of the Atom

• The protons and neutrons are tightly packed in a
  central core called the nucleus
• If an atom were the size of a football stadium,
  the nucleus would be a tiny marble sitting in the
  center of it
• The electrons are found in different layers
  (energy levels) of a cloud around the nucleus

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Atomic Mass

• Elements differ because of the number of protons
  they have
• The Atomic Number is the number of protons
• The number of electrons in an atom must equal the
  number of protons in order for the atom to be
  neutral
• The Mass Number is the whole number of protons
  plus neutrons in an atom

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Atomic Mass
O
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Mass Number
Chemical Symbol
8
Atomic Number
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Mass vs. Atomic Mass
Ne
20.18
Atomic mass
Chemical Symbol
10
Atomic Number
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Isotopes

• Isotopes of an atom occur when the number of
  neutrons changes
• Isotopes have the same chemical properties as the
  original atom because the charged particles
  remain the same

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Atomic Mass

• A weighted average mass of the atoms in a
  naturally occurring sample of an element is
  called the Atomic Mass
• This number represents the mass as well as the
  relative abundance of each isotope
• Since atoms are so small, grams are not typically
  used as units of mass
• Instead, an Atomic Mass Unit is used
  (mathematically defined as 1/12th of the mass of
  Carbon-12.)

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Bell Work

• Calculate this students grade if the class is
  weighted as follows
• Tests 75 Homework 5
• Lab 10 Final exam 10
• Test scores 89, 84, 72, 90
• Lab 99, 100, 98, 99, 94, 97
• Homework 92, 93, 96, 98, 105, 94
• Final exam 90

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Bell Work

• Calculate this students grade if the class is
  weighted as follows
• Tests 75 Homework 5
• Lab 10 Final exam 10
• Test scores 89, 84, 72, 90 335/4 83.75
• Lab 99, 100, 98, 99, 94, 97 587/6 97.8
• Homework 92, 93, 96, 98, 105, 94 578/6
  96.3
• Final exam 90
• 83.75(.75) 97.8(.10) 96.3(.05) 90(.10)
• 86.4 B

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Bell Work

• Now if this teacher did NOT weight grades, what
  would this students grade be?
• Test scores 89, 84, 72, 90
• Lab 99, 100, 98, 99, 94, 97
• Homework 92, 93, 96, 98, 105, 94
• Final exam 90
• 1590/1700 93.5 Avery different

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Periodic Table of Elements

• What do you think of when you hear the word
  periodic?
• Periodic actually means occurring on a regular
  basis
• There are certain trends that exist on the
  periodic table that are consistent

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Periodic Table of Elements

• A horizontal row across the periodic table is
  called a period.
• When you read across the page, you eventually
  come the end of a sentence. At the end of a
  sentence is a period.
• A vertical column on the periodic table is called
  a family or group.

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Alkali Metals

• Group/Family I is called the Alkali Metals
• Hydrogen is not included in this group
• It is in a group of its own
• This family shares certain characteristics
  react vigorously with water, are metals, and
  have 1 e- in their outermost shell

Li
Na
K
Rb
Cs
Fr
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Alkaline Earth Metals

• The alkaline earth metals are all metals and all
  have 2 e- in their outermost shell
• The second family from the left of the periodic
  table

Be
Mg
Ca
Sr
Ba
Ra
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Transition Metals

• The transition metals are located in the center
  of the periodic table.
• They vary in their number of electrons, however,
  they all share in the common properties of
  metals.
• 80 of all of the elements are metals
• The inner transition metals are referred to as
  the rare earth elements. These are the two rows
  found at the bottom of the periodic table

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Metalloids

• Along the zigzag borders are the metalloids
• These share some properties of metals (some of
  the time)
• Aluminum is an exception it is a metal

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Non-metals

• In the upper right hand corner of the periodic
  table are the non-metals
• Typically non-lustrous and are poor conductors of
  electricity
• Halogens (group 7) include chlorine and bromine
• Noble Gases (group 8) Undergo few or no chemical
  reactions

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Puzzle Activity Instructions

• As a group, you are resonsible for
• Drawing each missing piece of your puzzle (be
  sure to number it) on the white paper
• Guess what design or picture is on the piece and
  then draw and color it.
• When youre finished, give the puzzle to your
  instructor. She will give you the puzzle pieces.
• Compare the real pieces to the ones youve drawn.
  Write down ANY differences.
• In your Journal, what did this activity have to
  do with Mendeleev and the first periodic table?

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Valence Electrons

• The shell or energy level (n) containing the
  outermost electrons for an element is called the
  valence shell
• The electrons in that shell are called valence
  electrons
• These electrons are the farthest from the atoms
  nucleus and are therefore the easiest to remove
• How many valence electrons do each of the alkali
  metals have?

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Valence Electrons

• The similarity in the of valence electrons
  causes members of the same family to share
  chemical behaviors
• Hydrogen is so tiny, however that it reacts very
  differently than other members of its family

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How many valence electrons ?

• How many valence electrons do each of the
  following have?
• Na
• O
• C
• Cl
• B

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Ionization Energy

• Some energy is required to remove an electron
  from that valence shell
• This energy is referred to as the ionization
  energy
• This energy is measured in Volts
• Valence electrons are much easier to remove than
  electrons closer to the nucleus and are therefore
  usually the only ones capable of being removed

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Octet Rule

• We will soon be talking about chemical bonding
• One important rule to remember is that atoms tend
  to want 8 electrons in their outermost shell
• This could mean that they give electrons up, take
  on electrons, or share electrons in order to
  achieve this goal
• Hydrogen Helium are exceptionsthey only want 2.

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Energy Levels or Orbits

• Each orbit around the nucleus has a very specific
  energy associated with it
• When an element was treated with heat or an
  electric current, where did the energy go?
• The electrons will absorb this energy
• If each energy level is assigned a specific
  amount of energy, what does the electron have to
  do in order to absorb the extra outside energy?

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Energy Levels or Orbits

• It has to jump to the next energy level located
  farther from the nucleus
• This is now an excited electron
• Excited electrons are very unstable and cannot
  remain in the excited state
• They must return to their original orbit or
  ground state

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Energy Levels or Orbits

• In order for it to return to its ground state, it
  must give off the exact amount of energy it
  picked up from the outside source
• When it returns to its ground state, it emits or
  gives off the energy in the form of light and
  heat
• The light emitted by excited electrons in atoms
  is not a continuous spectrum (all the colors) but
  a line spectrum (only certain wavelengths)
• No two elements have the same line spectrum

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Visible Light Spectrum

• Reminder Light is a form of energy
• Review from gradeschool
• ROY G BIV
• Violet light has higher energy than red light
• There is an inverse relationship between light
  wavelengths and energy
• So as the wavelength of light gets larger, the
  energy of light gets smaller

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Other atomic models As you may recall

• Before Bohr, there was Thompson and Rutherford
• Thompson proposed that an atom was a ball of
  positive charges which contained several
  electrons
• Rutherford, with his gold foil experiment, showed
  that the bulk of the atoms mass was concentrated
  in a small, positively charged region called the
  nucleus

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Quantum Mechanical Model

• Bohrs model gave rise to the quantum mechanical
  model
• When Bohr proposed that the energy required to
  excite an electron (which was then later emitted)
  was quantized
• There is a specific amount of energy required in
  order for an electron to become excited and move
  to the next energy level
• BUT, each orbit or energy level has its own
  requirements. They are not all the same.

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Quantum Mechanical Model

• Differing from Bohrs model, the quantum
  mechanical model suggests that the electrons
  dont just follow an exact path around the
  nucleus like our planets do around the sun
• Instead, the true location of the electron is
  uncertain and only a probability of its location
  is mapped
• This idea lends to the analogy of a cloud (the
  more dense the cloud, the higher the probability
  of finding the electron there)

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Sublevels

• Each energy level (n) is made up of one or more
  subshells or energy sublevels
• The number of energy sublevels is the same as the
  number of the energy level (n)
• So, the 3rd energy level has 3 sublevels the 5th
  energy level has 5 sublevels and so on.
• The sublevels are designated s, p, d, and f.

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Orbitals

• As you proceed beyond the 3rd energy level,
  overlapping of sublevels occurs and becomes more
  complex as you increase the energy level number.
• The s sublevels have only one orbital
• The p sublevels have 3 orbitals
• The d sublevels have 5 orbitals
• The f sublevels have 7 orbitals

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Orbitals

• s orbitals are spherical
• p orbitals are dumbbell-shaped with three
  different spatial orientations
• d orbitals are interesting 4 of the 5 kinds of
  d orbitals are clover-leafed and the fifth has
  two opposite nodes with a ring in between
• f orbitals are too difficult to visualize

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RULES FOR FILLING ATOMIC ENERGY LEVELS

• Electrons fill up the energy sublevels
• The lowest energy sublevel must be completely
  filled before the next higher sublevel can begin
  to be filled. (Aufbau principle)
• Each orbital can hold a maximum number of 2
  electrons of opposite spin (Pauli exclusion
  principle)
• Due to their negative charge, electrons repel one
  another. They will not pair up in an orbital of
  any given sublevel until all orbitals in that
  sublevel have been half-filled. (Hunds rule)

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Electron Configuration

• There is a pattern that can be used to help you
  remember which energy sublevel is next in line
• 7s 7p 7d 7f 7g
• 6s 6p 6d 6f 6g
• 5s 5p 5d 5f 5g
• 4s 4p 4d 4f
• 3s 3p 3d
• 2s 2p
• 1s

g is theoretical and is not used in current
electron configurations
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Electron Configuration

• Remember, the maximum number of electrons an s
  sublevel can hold is 2.
• The p 6 The d 10 The f 14
• There is an easier way to indicate which
  sublevels are filled compared with drawing out
  the line diagrams each time
• This is called Electron Configuration

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Electron Configuration

• Electron configuration is a shorthand way of
  showing which orbitals of each sublevel are
  filled
• When done correctly, the sum of the superscripts
  of all orbitals equals the number of electrons in
  the atom
• For example the electron configuration for
  phosphorus is
• P 1s22s22p63s23p3
• Add the superscripts 22623 15 e- in P

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Exceptions to the rule

• Chromium and Copper have unusual electron
  configurations
• They do not follow Aufbaus energy diagram
• Write down the electron configuration for Cr
• What does it end with?
• The true electron configuration for Cr is
  1s22s22p63s23p64s13d5
• Likewise, the electron configuration for Cu ends
  with 3d10 with only 1 electron in the 4s level

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