SCH4U
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Transcript of SCH4U
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SCH4U
http://www.youtube.com/watch?v=-d23GS56HjQ
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Dalton’s Theory Matter is made up of indestructible atoms. Law of definite proportions:
Elements combine in a characteristic ratio Law of multiple proportions:
Some elements have more than one combining capacity
Law of conservation of mass: Atoms cannot be created nor destroyed
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Thomson’s Theory “The Raisin Bun” model:
+ and – charges are mixed together Gave us electrons Atoms can gain or lose electrons to form
ions
Said that the identity of an element was based on its number of electrons
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Rutherford’s Model Atoms have a tiny nucleus which contains
positive & neutral charges and makes up the majority of the mass of the atom
Electrons are negative and occupy most of the volume of the atom.
Protons tell us the identity of the element
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Atoms and IsotopesIsotopes Have the same number of protons and electrons
but have different amounts of neutrons. Radioisotopes – give off radioactivity when they
decay
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Rutherford Model – Planetary Model of the Atom
Protons
Neutrons
Electrons
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Particle Mass (kg)
Location Charge
Proton (p+)
1.673 x 10-27 Nucleus +1
Electron (e-)
9.109 x 10-31
Orbitals outside nucleus
-1
Neutron (n0)
1.675 x 10-27 Nucleus 0
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Representing Atoms
XZ
A
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Problems - Revisited SPIRAL DEATH!!!!
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To solve this problem… we need a little bit more of an insight into two phenomena:
LIGHT
ENERGY
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Light is a Wave!Huygens, Newton
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Light is a Particle! (The Photoelectric Effect)
• The ejection of electrons from a metal surface when light strikes it
• Certain types of light cause ejection, others don’t
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Max Planck Spectrum of Radiated energy and intensity
Quantum: unit or package of energy (plural quanta)
Energy is quantize – can only have allowed values
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Planck Equation Energy is equal to the frequency of the radiation
times Planck’s constant (h) h = 6.64×10-34 J∙s
Energy is QUANTIZED – it comes in packets and the smallest packet is equal to Planck’s constant Only multiples of this number are allowed –
nothing more
𝐸=h𝑓
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Photons By extension, light is also a quantize, since it is a
type of energy
Photon: unit of light energy Or particles of light energy
(Used to describe the photoelectric effect)
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Homework Page 142 #1-7
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Bohr’s Model of the Atom Limitations of the Rutherford Model
Electrons orbiting around a nucleus should lose energy and spiral into the nucleus
Electrons should be attracted to proton and collapse in to the nucleus
SPIRAL DEATH
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Atomic Spectra Continuous Spectrum: an emission spectrum
that contains all the wavelengths of light in a specific region of the electromagnetic spectrum
Line Spectrum: emission spectrum that contains only specific wavelengths characteristic of the element being studied
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Hydrogen Emission Spectrum
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Reason?
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Different for Each Element
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Bohr’s Postulates First Postulate:
e- do not radiate energy as they orbit the nucleus. Each orbit corresponds to a state of constant energy (called stationary state).
Basically energy states (or levels)
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Second Postulate: e- can change their energy only by
undergoing a transition from one stationary state to another
Basically, give the e- a quantum of energy and it’ll jump up to the next energy level, when it loses the quantum it falls back down, releasing a photon
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Bohr-Rutherford Model
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Successes and Failures of the Bohr Model Works well at predicting properties and
periodicity of the elements
Problem: everything was a little bit off after Hydrogen.
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Trends in the Periodic Table Atomic radius
Ionization Energy
Electron Affinity
Electronegativity
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Homework
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THE QUANTUM MECHANICAL MODEL OF THE ATOM
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And now for something completely different…
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Quantum Mechanics The application of quantum theory to explain
the properties of matter, particularly electrons in atoms
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Schrodinger’s Standing Waves Louis De Broglie developed a theory that matter
can have wave-like properties
Schrodinger extended this theory to electrons bound to a nucleus Postulated that electrons resembled a
standing wave Certain orbitals exist at whole wavelengths of
electron vibrations
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Orbitals - Redefined
Orbital: region around the nucleus where there is a high probability of finding an electron
As per wave model of Schrodinger – because things are vibrating
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Heisenberg Uncertainty Principle
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Heisenberg Uncertainty Principle Heisenberg studied statistics and developed matrix
algebra
Developed a statistical approach to explaining how electrons works and realized…
IT IS IMPOSSIBLE TO KNOW THE EXACT POSITION AND SPEED OF ELECTRON AT A GIVEN TIME At best, we can describe the probability of finding
it at a specific place
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Wave functions: the mathematical probability of finding an electron in a certain region of space
Wave functions give us:
Electron probability densities: the probability of finding an electron at a given location, derived from wave equations
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Homework
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Quantum Numbers Quantum Numbers: numbers that describe the
quantum mechanical properties (energies) of orbitals
From the solutions to Schrodinger’s equation
The most stable energy states is called the ground state
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Principal Quantum Number (n) Integer number (n)
used to level the main shell or energy level of the electron
Describes size and energy of the atomic orbital
Increase number = increase energy, bigger
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Secondary Quantum Number, l Describes the shape of the orbital within each
shell
Each energy level contains several sublevels
Relates to the shape of the orbital
Can be any integer from 0 to (n-1)
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Values of l
Value 0 1 2 3 4
Letter Used s p d f g
Name sharp principal diffuse fundamental
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Each orbital is given a code:
Example If n = 1, l = 0 then we call it a 1s orbital
If n = 3, l = 2 then we call it a 3d orbital
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Magnetic Quantum Number, ml
Describes the orientation of the orbital in 3-space
Can be whole number integers from – l to + l
Example: if l = 1, then ml can be -1, 0, +1 There are 3 possible p orbitals
px, py, and pz
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What are possible values for ml if l is: 0 1 2 3
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Spin Quantum Number Electrons are basically little magnetics spin
around when placed in magnetic fields, they can have spin ‘up’ or spin ‘down’
ms can be either +1/2 or – 1/2
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Homework
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Electron Configurations and Energy Level Diagrams The four quantum numbers tell us about the
energies of electrons in each atom
Unless otherwise stated were are talking about ground state energies
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Energy Diagrams Describe how electrons fill orbitals using
quantum numbers
Electrons fill the lowest energy level orbitals first
Each shell is (for the most part) filled before moving to higher shells
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Rules Use circles (or boxes) to represent each orbital
in any given energy level and arrows for electrons
Unoccupied circles imply that there are no electrons in it
A circle can have at most two electrons in it; only if the arrows are pointing in opposite directions
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Rules to Remember Pauli exclusion Principle: no two electrons can
have the same 4 quantum numbers. Electrons in the same orbital can’t have the same spin
Hund’s Rule: One electron occupies each of several orbitals in the same energy level before a second can occupy the same orbital
Aufbau Principle: each electron is added to the lowest energy orbital avaible
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Building Orbital Diagrams
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Practice H, B, C, Ne Mg, P, Ar Ca, Mn, Zn, Ge, Kr
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Electron Configurations Condensed versions of orbital diagrams and not
in
Write the electron configuration for each of the atoms above
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Exceptions to the Rules Examine the allowed charges for Chromium and
Copper
Write the electron configuration for chromium and copper
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What actually happens? Why? Evidence suggests that half-filled and
filled orbitals are more stable than other orbitals, so electrons rearrange to give these configurations
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Explaining Ion Charges In order to get particular charges, entire energy
levels or sublevels get cleared first.
Use electron configuration theory to explain why: Zn Zn2+
Pb Pb2+ or Pb4+
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Explaining Trends in the Periodic Table Atomic Radius: size of the atom
Ionization Energy: energy needed to remove an electron from the outermost energy level from an electron in the gaseous state
Electron Affinity: change in energy that occurs when an electron is added to a gaseous atom