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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 IsotopesIsotopes Have the same number of protons and electrons

but have different amounts of neutrons. Radioisotopes – give off radioactivity when they

decay

Rutherford Model – Planetary Model of the Atom

Protons

Neutrons

Electrons

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

Representing Atoms

XZ

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 Spectrum of Radiated 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

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

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 Atomic radius

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