Review

• Test Monday April 13th

• MC (10 marks)• Short answer (37 marks)

Review

• History– Dalton• All matter is composed of indivisible particles called

ATOMS• Atoms of a given element are identical in mass and

properties• Compounds are formed by a combination of 2 or more

atoms• Atoms cannot be created or destroyed

History

• J.J Thompson– Using a cathode ray, Thompson discovered the

subatomic particle called the ELECTRON were present in atoms• Small• Negatively charged

– PLUM PUDDING MODEL

History

• Rutherford – Conducted gold foil experiment – Sent a positive stream of particles through a piece of

gold foil– Predicted that all particles should pass through

however, every now and then a particle would deflect back

– Concluded:• Atoms are mostly empty space• Center of the atom is a dense positive area called the

NUCLEUS

History

•Bohr• Electrons orbit around the nucleus in shells (solar

system)• 7 different shells or orbits

• Electrons in different orbits have different energy levels• electrons would only occupy the lowest possible energy level • Electrons can move up a level if the lower levels were full

History

• Schrödinger– WAVE-MECHANICAL MODEL• Electrons move in areas called ORBITALS which is an

area of high probability of finding an electron • Energy can behave as both WAVES and PARTICLES

Light and Energy • WAVES – Amplitude- height of the waves from the origin to the

crest– Wavelength (λ lambda) – distance between crests

– Frequency – (f) is the number of wave cycles to pass in a given period of time

Light and Energy

• Higher Frequency Higher Energy

Light and Energy

• Bohr– Electrons orbiting in shells around the nucleus– Energy levels differ from one another– When an atom absorbs energy, electrons are

promoted to higher energy levels.– When the atom releases the absorbed energy, the

electron falls back down to lower energy levels and EM radiation is given off (sometimes light)

– These amounts are unique to every element

n=1

n=2

n=3

n=4

Spectrum

UV

IR

Vi s ible

Ground State

Excited State

Excited State

Excited State unstable and drops back down

•Energy released as a photon

•Frequency proportional to energy drop

Excited State

But only as far as n = 2 this time

Light and Energy

• Emission spectrum of each element is unique to that element

Electron Configurations

• Energy Levels– N= 1,2,3,4

• Sub-level– Shapes– How many orbitals – How many electrons in each sub-shell

Electron Configurations

Electron Configurations

• How electrons fill the orbitals is based on 3 rules:– The Aufbau principle– The Pauli Exclusion principle– Hund’s rule

Orbital Diagram

Electron Configurations

• The Pauli exclusion principle– Atomic orbital may describe at most 2 electrons

– Electrons must have opposite spins– ↑↓

Electron Configurations

• Hund’s rule• In a set of orbitals, the electrons will fill the

orbitals in a way that would give the maximum number of parallel spins (maximum number of unpaired electrons)

• ↑ ↑ ↑

Writing Electron Configurations

• Full electron configuration Atoms and Ions– Al: 1s2 2s2 2p6 3s2 3p1

– Al3+: 1s22s22p6

• Orbital Configuration (arrows)• Noble Gas Simplification for Atoms and Ion– [He] 2s22p6

Electron Configurations

• Identify– valence electrons– What element (based on electron configuration)• 1s22s1

– Invalid electron configurations

Periodic Trends

• Groups• Vertical column of the periodic table• Indicates valence electrons

• Periods • Horizontal row of the periodic table • Period will indicate number of energy levels around the

nucleus

Periodic Trends

• Atomic radius – increase as we go down a group– decrease we go across a period

• Ionic Radius– decreases when being converted to a positive ion– increases when being converted to a negative ion

Periodic Trends

• Electronegativity– decreases as you down a group– increases as you move across a period

• Ionization Energy – decreases as you go down a group– Ionization energy increases as you move across a

period