CHEM12 UNIT 1

Photon

Unit of light energy

Quanta

Small amount of energy

Electron

A negatively charged subatomic particle

Radioactivity

Spontaneous decay or disintegration of the nucleus of an atom

Nucleus

the dense center of an atom with a positive charge

proton

a positively charged subatomic particle

neutron

an electrically neutral subatomic particle

isotopes

atoms with the same number of protons but different number of neutrons

atomic number (Z)

number of protons in a nucleus

mass number (A)

protons + neutrons in a radioisotope (an isotope that emits radioactive gamma rays and/or subatomic particles

photoelectric effect

electrons are emitted by matter that absorbs energy from shortwave electromagnetic radiation (visible/UV light)

ground state

lowest energy state of an atom

quantized

energy moves in distinct steps (levels) rather than moving continuously E

excited state

state where potential energy is higher than ground state

Dalton’s model of an atom

• “Billiard ball”

• Matter is made of small, invisible spheres

• Matter cannot be created, destroyed or divided

• Atoms of the same element are identical in all other properties

Thomson’s model of an atom

• Atoms are positively charged spheres with negatively charged particles embedded in them

Rutherford’s model of the atom

• Positively charged, dense, nucleus with electrons around it

• Shining alpha particles at golden foil, most went through but some had wild deflections

Successes of the Bohr model

• Lower energy levels filled first

• First energy level has max 2 electrons, second has max 8, third has max 18

• Atoms are arranged according to photons, neutrons, and electrons and elemental properties

• Electrons exist in discrete energy levels

Failures of the Bohr Model

• Confusing past the first 20 elements

• Energies are inconsistent for atoms with more than 1 electron

• Electrons don’t orbit the nucleus

Bohr’s model of the atom

• circular orbits with distinct energy levels

• only exists in allowed orbits

• can jump orbits by gaining or losing energy

Einstein’s contributions to the atom

• Electromagnetic radiation is a stream of photons, each photons has it's own quantum energy

• Energy changes at the atomic level only occur at a certain amount of energy

• Energy of a photon is transferred to the electron, breaking the electron free from the atom

  • A minimum amount is required to be met

• Photon strength depends on energy frequency

Limitations of Rutherford’s model

• if an atom is constantly accelerating, it should be emitting electromagnetic energy

• If it is constantly emitting photons, it should be losing energy and decaying into a nucleus

Spectroscopy

Analyzing spectrums to determine properties of their source

line/emission spectrum

unique sets of colour produced when light from an excited substance is passed through a prism

Quantized

It can only exist in discrete, specifics values rather than a continuous range

Quantum mechanics

Application of quantum theory to explain properties of matter

Orbital

Region around nucleus where electron has a high probability of being found

Electron probability density/electron probability distribution

Indicates regions with the greatest probability of finding an electron ,determined with wave equations

Quantum mechanica model

Model for the atom based on quantum theory and calculation of probabilities for the location of electrons.

Heisenberg’s uncertainty principle

The idea that it is impossible to determine the position and speed of an object at the same time due to it’s nature of being both a. wave and a particle.

Pauli exclusion principle

Electrons placed in lowest orbitals first, arrows point in opposite directions to indicate spin.

No two electrons in the same atom can be in the same quantum state, therefore no two electrons have the same four quantum numbers and orientation is a unique property.

Aufbau principle

Orbitals must be filled before moving to the next highest orbital

Hund’s rule

Place one electron in each orbital (letter) of the same before a second electron is added

Order of orbital placement

Analogous electron configurations

• Typically found in d subshell

• Half filled and filled subshells are more stable than unfilled shells

• Atoms that end in d^4 or d^9 will promote an electron from an s orbital

• Cr, Mo, W, Cu, Ag, Au 

Paramagnetism

• weak attraction/magnetic field due to unparied electrons in an atom

• Unpaired electrons all have the same spin (Hund’s) spinning generates magnetic field

Quantum numbers

Describes the quantum mechanical properties of orbitals

Principal quantum number (n)

Size and energy level (shell) of an orbital

• eg: (n)s

• Ranges from n=1 to n=∞

Secondary quantum number (l)

• Describes shape of orbital

• Ranges from 0 to (n-1)

  • l=0 orbital s

  • l=1 orbital p

  • l=2 orbital d

  • l=3 orbital f

magnetic quantum number (ml)

• Orientation of orbital

• (2l+1) possibilities ranging from -l to l

spin quantum number (ms)

• Orientation of the axis of electron spin

• Can only be +½ or -½

• Draw out diagram, if spin is facing downwards it has a negative axis