#### Transcript Unit 1 Notes - Structure and Properties of Matter

```SCH4U
G Ra de 12
C He mistr Y
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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
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
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
Atoms and Isotopes
Isotopes
 Have the same number of protons and electrons
but have different amounts of neutrons.
decay
Rutherford Model – Planetary
Model of the Atom
Electrons
Protons
Neutrons
Particle
Proton
(p+)
Mass
(kg)
1.673 x
10-27
Location Charge
Nucleus
+1
-1
0
Electron
(e-)
9.109 x
10-31
Orbitals
outside
nucleus
Neutron
(n0)
1.675 x
10-27
Nucleus
Representing Atoms
Z
X
A
Problems - Revisited
 SPIRAL DEATH!!!!
 To solve this problem… we need a little bit more
of an insight into two phenomena:
 LIGHT
 ENERGY
Light is a Wave!
Huygens, Newton
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
Max Planck
energy and intensity
Quantum: unit or
package of energy
(plural quanta)
Energy is quantize – can
only have allowed values
Planck Equation
 Energy is equal to the frequency of the radiation
times Planck’s constant (h)
 h = 6.64×10-34 J∙s
𝐸 = ℎ𝑓
 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
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)
Homework
 Page 142 #1-7
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
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
Hydrogen Emission Spectrum
Reason?
Different for Each Element
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)
 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
Bohr-Rutherford Model
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.
Trends in the Periodic Table
 Ionization Energy
 Electron Affinity
 Electronegativity
Homework
THE QUANTUM
MECHANICAL MODEL OF
THE ATOM
And now for something
completely different…
Quantum Mechanics
 The application of quantum theory to explain
the properties of matter, particularly electrons
in atoms
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
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
Heisenberg Uncertainty
Principle
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
 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
Homework
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
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
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)
Values of l
Value
0
1
2
3
4
Letter
Used
s
p
d
f
g
Name
sharp
principal
diffuse
fundamental
 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
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
 What are possible values for ml if l is:
0
1
2
3
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
Homework
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
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
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
Rules
 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
Practice
 H, B, C, Ne
 Mg, P, Ar
 Ca, Mn, Zn, Ge, Kr
Electron Configurations
 Condensed versions of orbital diagrams and not
in
 Write the electron configuration for each of the
atoms above
Exceptions to the Rules
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