Announcement

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Announcement

• Homework is due at the beginning of class.
• Put it on your desk during the first two minutes.
• HW detentions emailed at the beginning of class.
• You will need a homework detention rescinding
  form signed (after class!) if you didnt
  complete your HW by the first two minutes of
  class.

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Review

• Review subatomic particles, mass number, atomic
  number, average atomic mass

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The Periodic Tableof the Elements
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Dmitri Mendeleev (1869)

• Organized the elements by atomic mass and other
  properties
• Looked for repeating patterns
• When arranged by increasing atomic mass,
  similarities in chemical properties appeared at
  regular intervals.
• Periodic things that repeat at regular intervals

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The Periodic Table is Born

• Mendeleev then created a table where elements
  with similar properties were grouped together
  the first periodic table of the elements (1869)

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The First Periodic Table
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• Atomic mass
• Always increasing?
• Empty spaces
• Predicted they would be filled by elements with
  certain properties
• By 1886 these spaces were filled, by elements
  with those properties!

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Another Contribution

• Henry Moseley arranged elements by their atomic
  number, not by atomic mass
• Periodic Law the physical and chemical
  properties of the elements are periodic functions
  of their atomic numbers
• Periodic functions repeating patterns
• Related to the number and location of electrons

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Increasing Atomic Number (and, more or less,
atomic mass)
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Periods (across rows)
Groups (down columns)
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Group Numbers
1
18
2
13
14
15
16
17
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• Elements in the same group have similar chemical
  and physical properties.

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Areas of the PTE
Non-metals
Metals
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A quick tour of the periodic table
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Metals

• Solid at room temperature
• Conduct electricity
• Shiny
• Ductile and malleable can be made into wires and
  beaten into sheets

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Group 1 Alkali Metals

• Extremely reactive
• Why?
• Dont exist alone naturally

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sodium
potassium
cesium (l)
rubidium
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• http//video.google.com/videoplay?docid-213426665
  4801392897ei59rqSt_wCNSclAfixvCdCQqalkalimeta
  lshlen

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Group 2 Alkaline Earth Metals

• Very reactive

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Non-Metals

• Liquids or gases at room temperature
• Dont conduct electricity

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Group 17 Halogens

• Very reactive non-metals
• Need only one electron to fill valence shell
• All exist as diatomic molecules
• Cl2, Br2, I2 (all colored)

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Halogens colored diatomic molecules
fluorine pale yellow gas
chlorine yellow-green gas
bromine red-brown liquid
iodine black-purple liquid purple gas
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Group 18 Noble Gases

• Not reactive
• Gases at room temperature
• Stable electron arrangement

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The Giant Periodic Table Project
6 C 12.01
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The Giant Periodic Table Project

• Be neat
• Be bold! We need to be able to read this from
  anywhere in the room!
• Be creative, but not cluttered.
• Be accurate.
• Atomic mass to TWO digits (Ex 12.01) is ideal.

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Electron Configuration

• The arrangement of electrons in an atom
• Each elements atoms are different
• How do we figure out what the electron
  configuration looks like?

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So where are the electrons?

• Electrons do NOT orbit like planets.

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Energy Levels

• Major shells (layers) around the nucleus
• lower levels filled before higher levels are
  filled
• 1st 2 electrons
• 2nd 8 electrons
• 3rd 8 electrons

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Practice with Simple Electronic Structure

• H
• He
• B
• Mg
• Li
• Be

• O
• Ne
• Cl

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• Electron Apartment Building

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Electronic Structure

• Energy levels
• Sub-levels
• Orbital
• Two electrons per orbital

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Inside energy levels

• Each energy level has one or more sub-level made
  from different shaped orbitals.

Energy Level Number of Sub-levels
1 1 (s)
2 2 (s, p)
3 3 (s, p, d)
4 4 (s, p, d, f)
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• Sub-levels are made up of one or more orbitals

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What is an orbital?

• Orbital An area where you expect to find the
  electron.

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Sub-levels

• Different shapes
• s sphere
• one orbital
• p figure eight
• three orbitals
• d
• five orbitals
• f
• seven orbitals

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p Sub-level

• p sub-level has three orbitals
• px, py, pz

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Energy Level Types of sub-levels Total orbitals Electron capacity (orbitals x2)
1 s 1 2
2 s, p 13 4 8
3 s, p, d 135 9 18
4 s, p, d, f 135716 32
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How do we figure out where the electrons are?

• 1. Figure out the energy levels of the orbitals
• 2. Add electrons to the orbitals according to
  three rules
• Three Rules for Electron Configuration

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1 Aufbau Principle

• An electron goes to the lowest-energy orbital
  that can take it.

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2 Pauli Exclusion Principle

• No two electrons can have exactly the same
  configuration description
• Can have the same orbital, but must have opposite
  spins.

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3. Hunds Rule

• Orbitals of equal energy are each occupied by one
  electron before any orbital is occupied by a
  second electron
• All electrons that are by themselves in an orbit
  must have the same spin.

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Energies of sub-levels
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Multiple choice practice
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1. Who arranged the periodic table by atomic number?

1. Dmitri Mendeleev
2. Henry Mosley
3. Ernest Rutherford
4. Democritus

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1. Which property is not increasing from left to
   right on the periodic table?

• Atomic mass
• Atomic number
• Number of electrons
• Reactivity

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1. Which is not a property of the metals?

• Shiny
• Gas at room temperature
• Conduct electricity
• Solid at room temperature

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1. Which is a property of non-metals?

• Conduct electricity
• Gas or liquid at room temperature
• To the left of the staircase line
• Malleable and Ductile

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1. Which is not a property of the alkali metals?

• Solid at room temperature
• Not reactive
• Extremely reactive
• Never exist alone in nature

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1. Which accurately describes the maximum number of
   electrons in each energy level?

• 2, 8, 8
• 8, 8, 8
• 2, 2, 8
• 8, 2, 2

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1. Which lists the sublevels in the correct order?

• d, p, s, f
• s, p, f, d
• p, s, f, d
• s, p, d, f

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1. What is the name of the following orbital shape?

• s
• p
• d
• f

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1. How many electrons can fit in the s
   sub-level/orbital?

• 1
• 2
• 3
• Eight million

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1. How many electrons can fit in any single orbital?

• 1
• 2
• 3
• Eight million

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1. How many electrons total can fit in the p
   sub-level?

• 2
• 4
• 6
• 14

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1. How many d orbitals make up the d sub-level?

• 2
• 3
• 5
• 7

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1. How many f orbitals make up the entire f
   sub-level?

• 2
• 3
• 5
• 7

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1. How many electrons total can fit in the entire d
   sub-level?

• 2
• 6
• 10
• 14

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1. Which rule states that the lower energy levels
   are filled before the higher ones?

• Aufbau Principle
• Pauli Exclusion Principle
• Hunds Rule

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1. Which rule states that electrons are alone unless
   there is no space left, and then they can be with
   another electron?

• Aufbau Principle
• Pauli Exclusion Principle
• Hunds Rule

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1. Which rule states that no two electrons can have
   the same configuration (that is, be in the same
   space at the same time)?

• Aufbau Principle
• Pauli Exclusion Principle
• Hunds Rule

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Orbital Notation

• Take these notes on Orbital Notation paper

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• What do the orbitals/sub-levels "look like" when
  you put them all together?

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1 Aufbau Principle

• An electron goes to the lowest-energy orbital
  that can take it.

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2 Pauli Exclusion Principle

• No two electrons can have exactly the same
  configuration description
• Can have the same orbital, but must have opposite
  spins.

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3. Hunds Rule

• Orbitals of equal energy are each occupied by one
  electron before any orbital is occupied by a
  second electron
• All electrons that are by themselves in an orbit
  must have the same spin.

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How do we write electron configurations?

• Orbital Notation
• Electron Configuration Notation
• Noble Gas Notation (shorthand)
• All ways to communicate where the electrons are
  in any atom.

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• s sphere
• one orbital 2 electrons
• p figure eight
• three orbitals 6 electrons
• d
• five orbitals 10 electrons
• f
• seven orbitals 14 electrons

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Orbital Notation

• Electron configuration goes below a line or box
• Arrows representing electrons go on the lines or
  in the boxes

1s 2s 2p 3s 3p
4s
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Be, Si, Ne, Cl
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Write the orbital notationHomework notebook!

1. He
2. N
3. Ne
4. Na
5. S
6. Zn
7. I

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Electron Configuration Notation

• Main energy level Sub-levelElectrons
• Carbon (6 electrons) 1s22s22p2
• Aluminum (13 electrons)
• 1s22s22p63s23p1
• Oxygen
• Argon
• Copper
• Zirconium

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Noble Gas Notation (Shorthand)

• Noble gasses have totally filled outer orbitals.
• If Ne (a noble gas) is 1s22s22p6, we can
  abbreviate Na (Sodium) as Ne3s1.
• Sodium has one more electron than Neon, so its
  Noble Gas Notation is Neon plus one electron in
  the s sublevel of the third energy level.

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Practice

• Fe (Iron)
• Electron Configuration Notation
• Noble Gas Notation
• K (Potassium)
• Electron Configuration Notation
• Noble Gas Notation
• Li (Lithium)
• Be (Beryllium)

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Orbital Notation

• Boron
• 1s22s22p1
• Atomic Number?
• How many electrons?
• Orbital notation

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• Aufbau Principle start with 1s and work up in
  energy level
• Pauli Exclusion Principle No two elements can
  have the same arrangement of electrons
• Hunds Rule Fill in one electron per orbital
  first, then go back.

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s, p, d and f blocks of the periodic table
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Long form periodic table
groups
periods
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• Electronic structures are related to the position
  of the elements on the periodic table
• s-block s orbitals are filled
• p-block p orbitals are filled
• etc.

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