AP Chemistry Unit 2

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AP ChemistryUnit 2

• Orbitals and Bonding
• Chapters 8, 9, 7

2
Electronegativity
3
Electronegativity

• Developed by Linus Pauling
• Measures the ability for an atom in a compound to
  attract electrons from another atom in the
  compound
• AKA Measures how badly they want another electron

4
Electronegativity

• High electronegativities mean they will more
  readily take an electron, lower means less likely
• Electronegativities near or below 2 means they
  will readily give up electrons

5
Electronegativity

• Tend to increase across a period and decrease
  down a group
• Upper right of the p block will very easily take
  electrons
• Lower left of s block easily lose electrons

6
Bonding
7
Bonding

• 2 bonding extremes
• Ionic
• Covalent
• Everything is in between these

8
Ionic Bonding

• One atom steals the electrons of another
• Attraction between anions and cations
• Large difference in electronegativies
• ?EN gt 1.7
• Usually s block and nonmetal

9
Ionic Bonding

• Na has EN of 0.9
• Cl has EN of 3.0
• ?EN will always be positive, this determines
  cation and anion
• ?EN ENanion ENcation
• Use table on page 353

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Covalent Bonding

• Sharing electrons
• Small difference in EN
• ?EN lt 1.2

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Covalent Bonding

• Cl2
• ?EN 0
• HCl
• H has EN of 2.1
• Cl has EN of 3.0
• ?EN of 0.9

12
Polar Covalent Bonding

• Unequal sharing
• Moderate difference between atoms
• 1.2 lt ?EN lt 1.7

13
Polar Covalent Bonding

• LiI
• Li has EN of 1.0
• I has EN of 2.5
• ?EN of 1.5

14
Metallic Bonding

• Lots of cations (metals that have lost electrons)
  all have electrons flying around between them
• These arrangements make metals malleable,
  ductile, and good conductors
• Malleable ability to be hammered into sheets
• Ductile ability to be drawn into wires

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Bonding Visuals
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Bonding Visuals
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Practice

• What type of bonds are these?
• NaAt
• FrBr
• MgCl2
• AgBr2
• HgO
• H2O
• YN
• C2H2

18
Bond Polarities
19
Bond Polarities

• All covalent bonds have some small degree of
  polarity
• Except homonuclear diatomics
• A molecule with distinct positive and negative
  regions are dipolar, and have a dipole moment

20
Bond Polarities

• Sum of bond dipoles gives the overall dipole of
  the compound
• Draw the dipole of each bond
• Equal dipoles in opposite directions cancel each
  other out

21
Showing Dipole Moments

• Plus sign goes on the d atom with an arrow
  extending out and pointing to the d- atom

22
Practice

• Are these polar molecules?
• Ammonia
• Carbon tetrachloride
• Perchloric acid

23
Intermolecular Forces
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Intermolecular Forces

• Forces between compounds, not within
• Forces within a compound are intramolecular

25
Dipole-Dipole

• Strongest force between polar molecules
• Uneven distribution of charges
• Polar covalent bonding

26
Hydrogen Bonding

• A specific type of dipole-dipole force
• Only occurs in H-O, H-F, and H-N bonds
• Hydrogens are attracted to the lone pairs of
  another molecule

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London Dispersion Force

• Attraction due to momentary dipole created by the
  natural movements of electrons
• Very weak, occurs in all atoms and compounds

28
Molecular Geometry
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VSEPR Theory

• Valence Shell Electron-Pair Repulsion Theory
• The repulsion of the electrons to each other
  makes the bonds between atoms desirable to be as
  far apart as possible.

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VSEPR Theory

• Lone pairs need to be accounted for as well
• Geometry can be determined based on the number of
  atoms and lone pairs bonded to the central atom
  (stearic ) and the number of lone pairs
• Larger molecules have multiple central atoms

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Example

• Draw the Lewis Structure for water (H2O)
• According to VSEPR Theory how do we get the
  electrons farthest apart?
• Hint Think 3-Dimensionally
• 4 bonds, 2 lone pairs
• Bent Structure

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Practice

• NH3
• CO2
• NaAt
• PCl3
• HCP
• SF6
• CH4

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Lewis Structures

• Remember what they are?
• Follow the Octet Rule - normally

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Example
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Practice

• Draw these Lewis Structures
• O2
• CH4
• C2H5OH
• HCl
• N2

37
Bond Hybridization
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Hybridization

• Geometry determined by hybridization of the bond
• A hybrid bond is a combination of orbitals to
  create a new orbital

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Example

• Draw methanes (CH4) orbital notation
• 2s and 2p merge to form a 2sp3
• An sp3 sublevel has 4 orbitals, as many as there
  were before, just at a single energy

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Hybridization
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Valence Bond Theory
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Remember Lewis Structures?

• Showing covalent bonds using electron dot
  structures simplified to show bonds
• 1 flaw Assumes shared electrons are between the
  atoms
• Localized electron model
• Are they?

44
Valence Bond Theory

• Allows bond angles to be calculated
• Takes in to account the fact that electrons are
  not localized
• Thanks to Heisenberg

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Sigma and Pi Bonds
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Sigma and Pi Bonds

• When orbitals overlap they make 2 special types
  of bonds
• Sigma bonds (s bonds) are formed along a line
  between the nuclei
• Pi bonds (p bonds) are formed by overlapping
  side-by-side p orbitals

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s Bonding

• Can be formed by overlapping orbitals along the
  axis between the 2 nuclei overlapping
• s, p, or a hybrid

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p Bonding

• An overlap of either of the p orbitals
  perpendicular to the axis between the two atoms
• ONLY p orbitals

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Sigma and Pi Bonding
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Applying s and p Bonding

• A single bond is a s bond
• A double bond is a s bond and a p bond
• A triple bond is a s bond and 2 p bonds
• Only unhybridized bonds form s or p bonds
• This portion of Valence Bond Theory explains why
  we cant have more than a triple bond, or 6
  shared electrons

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Connecting it All
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Determining Bonding

• The formula can be translated in to a molecular
  structure, showing single, double, and triple
  bonds
• The type of bond is translated s and p bonds for
  each atom
• In addition, the orbital notation can tell the
  hybridization and s and p bonds

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Examples

• Draw the structure, then write the orbital
  notations, determine hybridization, and bonding
• CH4
• CO2
• C2H2

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Practice

• Draw the structure, then write the orbital
  notations, determine hybridization, and bonding
• N2
• H2O
• HCl
• HCN

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Double Bonds
56
Double Bonds

• Double bonds have some special properties
• Double bonds usually only form between period 2
  nonmetals
• Double bonds crucial to all living organisms
• Silicon-based life cannot be as complex as
  carbon-based, even though silicon acts much the
  same as carbon

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Double Bonds

• They have a bond strength greater than a single
  bond, but less than 2 single bonds
• This is in part because p bonds are weaker than s
  bonds since they are only a side-by-side overlap
  on parallel axes rather than on the same axis

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Double Bonds

• Looking at molecules as 3D structures how can
  atoms move?
• In single bonds
• In double bonds

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Double Bonds

• Atoms cannot rotate around double bonds
• This is because p bonds dont allow rotations
• If groups larger than hydrogen are on the same
  side of a double bond it is referred to as cis-,
  if they are on opposite sides of a double bond
  they are called trans-

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Cis- and Trans-
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Practice

• Draw the cis- and trans- versions of each
• CH3CHCHOH
• C3H4Br2

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Bond Angles
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Bond Angles

• The geometry determines bond angles
• Steric determines geometry
• Atoms and lone pairs want to be as far apart as
  possible

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Bond Angles
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Bond Angles

• The angles given are maximums, lone pairs
  decrease the angles
• Water has a bond angle of 104.5 caused by 2 lone
  pairs pushing the hydrogens closer together

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Bond Length
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Bond Length

• The shorter a bond is, the stronger it is
• The more shared electrons there are the shorter
  the bond
• In general, the larger the difference in
  electronegativities, the longer the bond

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Molecular Orbital Theory
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Molecular Orbital Theory

• Based on quantum mechanics and Valence Bond
  Theory
• Lewis structures show lone pairs
• Not quite true

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Example

• Draw the Lewis structure of oxygen
• Are there lone pairs?

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Oxygen

• Oxygen is paramagnetic, which only happens with
  unpaired electrons
• Paramagnetism is the ability to be affected by a
  magnetic field, but do not hold their magnetic
  presence in the absence of an applied field
• Molecular Orbital Theory explains this

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Molecular Orbital Theory

• Developed in the 1920s by Robert Woodward and
  Roald Hoffman
• Also called MO Theory
• Describes the orbitals throughout a molecule and
  how they bond, draws on Valence Bond Theory

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MO Theory

• Electrons belong to the molecule, and shared not
  between any 2 atoms, but all of them attached in
  the molecule
• There arent really atom-specific orbitals, but
  more of molecular orbitals
• Electrons are delocalized, spread out over the
  whole molecule, not in bonds

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MO Theory

• MOs with a lower energy than the separate atomic
  orbitals are called bonding, MOs with higher
  energy are called antibonding
• The molecular orbital energy level diagram
  combines atomic energy levels and shows the
  created bonding and antibonding energy levels

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MO Diagram
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MO Diagram
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What does it mean?

• The lowest energy is s, which is located between
  the nuclei
• There is also a higher antibonding orbital
  formed, s
• The total number of orbitals is conserved

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Bond Order

• Bond order is just the amount of bonds that link
  a pair of atoms
• Bond order ½ (number of bonding electrons
  number of antibonding electrons)
• When determining bond order must show s and p MOs

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Example

• H2

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Practice

• N2
• O2
• F2

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Bonds

• For s orbital interactions only 1 bonding and 1
  antibonding orbital are created
• s and s

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Bonds

• Up to 3 bonding with p orbitals
• 1 s and 2 p
• Also 3 antibonding possible
• 1 s and 2 p
• If 2 atoms with p orbitals interact there are 6
  orbitals being combined

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Practice

• H2
• N2
• CH bond
• OH bond

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Exceptions to the Octet Rule
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Exceptions to the Octet Rule

• Hydrogen and helium only need 2 valence electrons
• Boron and beryllium like 6
• This is called electron deficient
• Electron deficient elements are extremely
  reactive
• The more lone pairs another compound has, the
  more violent the reaction

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Practice

• Draw boron trichloride
• What is the geometry, steric ? What type of bond
  is it?
• Write the reaction showing the formation of boron
  trichloride. Was it energetic?

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Bond Energies
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Bond Energies

• Also called bond enthalpy
• The stronger the bond the more energy it has and
  the more is needed to break it

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Electromagnetic Radiation
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Electromagnetic Radiation

• Energy that acts as a wave
• Energy linked to wavelength and frequency of the
  light

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Electromagnetic Radiation

• All waves move at 3.00 x 108 m/s in a vacuum
• Wavelength (?) is the distance between
  corresponding points on a wave
• Frequency (?) is the number of waves per second
  that pass by a point

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Energy of Waves

• High frequency and low wavelength are high energy
• Low frequency and high wavelength are low energy

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Electromagnetic Spectrum
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Visible Light

• Only certain range detectible to human eyes
• 400-700 nm or 5 x 1014 1 x 1015 Hz
• Violet is the highest energy, red the lowest

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Ultraviolet and Infrared

• Ultraviolet (UV) means more than violet, higher
  energy and frequency, shorter wavelength
• This is why UV rays are damaging
• Infrared (IR) is less than red, lower energy and
  frequency, longer wavelength
• This is why IR remotes dont do damage

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Electromagnetic Spectrum

• Be able to rank from high to low by wavelength,
  frequency, or energy
• Pg 291

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Relations of Wavelength and Frequency

• The speed of light (m/s) is equal to the
  wavelength (m) times the frequency (1/s or Hz)
• c ??
• Which is where a physics joke comes from.

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Practice

• A wave has a frequency of 9.1 x 10-8 Hz, what is
  its wavelength and what type of wave in the
  electromagnetic spectrum is it?

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What is Matter?
100
The Modern Atom

• Bohr Model
• Atoms made of 3 particles

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Fundamental Particles

• These are the most basic particles
• Quarks make protons and neutrons

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Quarks

• 6 types
• Up, down, top, bottom, charm, and strange
• Combinations of 3 make up protons and neutrons

103
Light

• Originally thought to be a wave
• Discovered to also behave like a particle
• A packet of energy called a quantum is wavelike
• A photon contains a quantum and behaves like a
  particle

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The Energy of Light

• Quantum also happen to be multiples of the energy
  gained or lost by electrons as they move closer
  or further from the nucleus, emitted as photons
• Energy (J) equals Plancks Constant (Js) times
  the frequency (1/s or Hz)
• E h?
• ?E nh?

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Plancks Constant

• Developed by Max Planck
• Relates the energy to the frequency
• 6.626 x 10-34 Js

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Special Relativity
107
Special Relativity

• Developed by Einstein in 1905
• Mass defined by the energy and the speed of light
• m E/c2
• E mc2
• m h/?c

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Photons Have Mass?

• Yes and no
• While in motion they have a mass relative to
  other particles and other interactions, but if it
  stops it has no mass

109
Wave-Particle Duality

• Developed by Einstein
• Light behaves both as a wave and a particle
• Quanta act like waves and are a unit of energy
• Photons act as particles and are made of quanta
• Photons have 0 mass and are pure energy

110
Wave-Particle Duality

• 1924 Louis de Broglie realized electrons also
  acted like a wave as well as a particle
• Waves confined to a set space have only certain
  frequencies, much like electrons do
• de Broglie pointed out that electrons acted like
  waves confined to the space around an atom

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de Broglie Waves

• All matter can behave as a wave
• Wavelength of a particle determined by the mass
  and speed of the particle
• ? h/mv

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Diffraction

• Splitting of light due to its wave-like properties

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Interference

• Constructive waves combine and increase
  intensity, waves have similar displacement
• Destructive waves cancel each other out,
  opposite displacement

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Line Emission Spectra
115
Line-Emission Spectra

• Energy added to gas atoms and electrons emit
  photons
• Electrons promoted from the ground state to an
  excited state
• Ground state is the lowest energy level
• Excited state is any higher energy state

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Line-Emission Spectra

• Each element has a unique set of lines

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Line-Emission Spectra

• Lyman Series
• UV
• Balmer Series
• Visible, most commonly used
• Paschen Series
• IR

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Energy Levels
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Energy Levels
120
Niels Bohr

• Proposed set orbitals for electrons
• These were based on the line-emission spectrum of
  different elements
• Photons emitted have specific energies based on
  how much energy is lost in falling to a lower
  energy level

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Quantum Theory

• When electrons lose energy they emit it as
  photons
• Studies the behavior of atoms and their energies
  in relation to photons and quanta of energy

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Quantum Numbers
123
Schrödingers Equation

• Erwin Schrödinger developed an equation to
  mathematically show electrons act as wave
• Solutions prove that only certain energies are
  allowed
• Gives probability regions, called orbitals, where
  electrons can be
• Probability regions are where electrons are 95
  likely to be

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Solutions to Schrödingers Equation

• Gives 4 quantum numbers
• Principal quantum number n
• Angular momentum quantum number l
• Magnetic quantum number m
• Spin quantum number

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Principal Quantum Number

• n gt 0
• Gives the energy level of the electron
• As n increases so does energy and distance from
  the nucleus
• Electrons allowed per level is 2n2 up to 32

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Principle Quantum Number
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Angular Momentum Quantum Number

• 0 l n-1 3
• Gives the sublevel
• All sublevels but n1 have multiple orbitals

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Angular Momentum Quantum Number
129
s Sublevel Orbitals

• Spherical
• Holds 2 electrons

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p Sublevel Orbitals

• 3 orbitals
• Named by which axis they lay along
• Holds 6 electrons total

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d Sublevel Orbitals

• 5 orbitals
• Named for orientation
• Holds 10 electrons total

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f Sublevel Orbitals

• 7 orbitals
• Holds 14 electrons total

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f Sublevel Orbitals
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Orbitals

• s is lowest energy, then p, d, and f

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Magnetic Quantum Number

• -l m l
• Gives orientation of the orbitals
• Gives number of allowed orbitals
• 2l 1

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Spin Quantum Number

• Only 2 possible values, relating to the two
  magnetic poles
• 1/2 or -1/2
• Each orbital can have 2 electrons and they must
  have opposite spins

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Periodicity
138
Atom Radii

• Defined by ½ the distance between the nuclei of
  identical atoms bonded together
• Radii get smaller across a period and larger down
  a group
• Smallest radii are at the upper right, largest at
  the lower left

139
Atomic Radii
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Ionic Radii

• Cations smaller than their neutral version
• More protons than electrons, held tighter
• Anions larger than their neutral version
• More electrons than protons, held less tightly
• Increase down a group

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Ionic Radii
142
Ionization Energy

• The energy required to remove an electron from an
  atom
• Increases for each successive electron removed
• The stronger the electron is held to the atom the
  higher the ionization energy
• Only deals with loss of an electron
• Increases across the period and decrease down the
  group

143
Heisenberg Uncertainty Principle

• It is impossible to determine both the position
  and velocity of a particle.
• Also stated Observing a system disturbs it in
  such a way that any data collected cannot be
  certain.

144
2 Visualizations

• Explained by Schrödinger

145
Double Slit Experiment

• A beam of light is shined on a paper with 2
  parallel slits
• What should show on the screen behind the paper
  with slits?

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Schrödingers Cat

• A radioactive isotope is placed in a container
  with a Geiger Counter connected to a poison. If
  the counter detects radioactive decay the poison
  will be released and kill the cat in the
  container. If not the cat will be fine. The box
  is sealed with the cat inside. Is the cat alive
  or dead?

148
Schrödingers Cat

• The cat is both alive and dead until observed

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End of Unit

• Questions?