U1L1: The Atom
Dalton’s Billard Ball Model (1805) Thomson’s “Plum Pudding” Model (1897)
- All matter is made of atoms - Cathode rays were deflected by charged
- Atoms of an element are identical objects
- Compounds are formed through different - Made of negatively charged particles
elements combined in constant ratios - The discovery that the negative subatomic
- Atoms are not created nor destroyed, but are particles (electrons) were scattered around a
rearranged in reactions positive cloud within atoms
- Theory: atoms are INDIVISIBLE and - Recognized electrons as components of atoms
INDESTRUCTIBLE - No nucleus; did not explain experimental
- Atoms aren’t indivisible, they’re observations
composed from subatomic particles
Hantaro Nagaoka’s “Saturnian” Model (1903) Rutherford’s Nuclear Model (1911)
- Refined Thomson’s model to resemble a - Gold Foil Experiment
miniature solar system - Alpha particles were fired at thin metal foils
- Centre has a large positive charge with - Small number of particles were deflected at a
negative charges orbiting it large angle while most went through
- Dense nucleus surrounded by a cloud of
Gilbert Lewis Newton’s Dot Diagrams (1916) negative charges
- Chemical bonding only involves valence
electrons (outermost electrons in the last - Positive charge was localised in the nucleus of
energy level) an atom
- Did not explain why electrons remain in orbit
around the nucleus
Bohr’s Planetary Model (1915) Schrodinger’s Quantum Model (1926)
- Energy orbit the nucleus in energy “shells” - Electrons move in waves, not set paths around
- Electrons travel within energy level without the nucleus
losing or gaining energy - Electrons can behave like atom and an electron
- Greater distance from the nucleus = greater simultaneously
energy - “Cloud of probability” illustrates the
possibilities of where the electron may lie as
- Proposed stable electron orbits; explained the their distinct position is unknown
emission spectra of some elements
- Move of electrons within orbit shells should - Electrons do not move around the nucleus in
emit energy and collapse into the nucleus orbits but in clouds where their positions are
(model did not work well for heavier atoms) uncertain
U1L2: The Quantum Mechanical Model
,Bohr Diagram show energy levels that electrons occupy
- Higher level = higher energy
Absorption: when elements are “excited” electrons can
absorb wavelengths and move to a higher energy level
Emission: after “excitation” electrons emit specific
wavelengths of light and go back to lower energy
positions
- As hydrogen electrons absorb energy, they
jump into an outer ring
- As they lose energy, photons of light are
emitted with a specific amount of
energy/specific wavelength which causes them
to fall back into an inner ring (energy loss).
- Sum of visible photon appears as a single color
until they are separated through a prism
- Each element has a unique emission spectrum
Absorption and Emission Spectra
- Every element is unique because of the unique
electrons in its orbitals
Behavior of Electrons
- Particle Wave Duality: electrons behave like
both particle and a wave
- If a particle passes through a small slit, it
would be detected in a small region on the
other side of the slit
- If a wave passes through a small slit, it would
spread out across a larger region on the other
side of the slit
If there are two slits, waves will experience an
interference pattern
Heisenberg’s Uncertainty Principle:
- Due to its wave and particle nature, we can
never know both the position and speed of an
electron
Orbits Orbitals
- Electrons orbit in a rigid structure (line/circle) - Regions of space around the nucleus
- Simple planar representation of an electron (electrons)
- Lines in an orbit are distinct - 3D model (involves a cloud)
, - 2D model - Orbitals do not have discrete/clear lines
- Electrons behave like a particle (has a distinct - Probability cloud (electrons can be anywhere
position) within the cloud)
- Electrons behave like particles but has a wave
function
U1L3: Quantum Numbers
Heisenberg’s Uncertainty Principle
1. The velocity of an electron is related to its wave “The more accurately we know the position of an
nature electron, the less accurately we can know its velocity”
2. The position of an electron is related to its
particle nature
Probability Distribution Map:
- How likely it is for an electron to exist in a
certain place
- Darker regions = higher probability
- Lighter regions = lower probability
Orbitals: probability distribution map for electrons
Nodes: Points where electrons cannot be located
- More nodes, higher energy
1. s subshell: 1 orientation
2. p subshell: 3 orientations
3. d subshell: 5 orientations
4. f subshell: 7 orientations
Quantum Numbers:
- A set of four numbers used to describe a
specific electron within orbitals (probability
clouds)
- Each electron has its own specific set of
quantum numbers
1. Principal Quantum Number (n)
2. Angular Momentum Quantum Number (l)
3. Magnetic Quantum Number (ml)
4. Spin Quantum Number (ms)
The Principal Quantum Number (n) The Angular Momentum Quantum Number (l)
- Overall size and energy of an orbital - Shape of the orbital
- Distance of the electron from nucleus - Subshells within the main energy level
- Closer to nucleus = lower energy - Values (n - 1): l = 0, 1, 2, …
- Values: n = 1, 2, 3, 4, …
Dalton’s Billard Ball Model (1805) Thomson’s “Plum Pudding” Model (1897)
- All matter is made of atoms - Cathode rays were deflected by charged
- Atoms of an element are identical objects
- Compounds are formed through different - Made of negatively charged particles
elements combined in constant ratios - The discovery that the negative subatomic
- Atoms are not created nor destroyed, but are particles (electrons) were scattered around a
rearranged in reactions positive cloud within atoms
- Theory: atoms are INDIVISIBLE and - Recognized electrons as components of atoms
INDESTRUCTIBLE - No nucleus; did not explain experimental
- Atoms aren’t indivisible, they’re observations
composed from subatomic particles
Hantaro Nagaoka’s “Saturnian” Model (1903) Rutherford’s Nuclear Model (1911)
- Refined Thomson’s model to resemble a - Gold Foil Experiment
miniature solar system - Alpha particles were fired at thin metal foils
- Centre has a large positive charge with - Small number of particles were deflected at a
negative charges orbiting it large angle while most went through
- Dense nucleus surrounded by a cloud of
Gilbert Lewis Newton’s Dot Diagrams (1916) negative charges
- Chemical bonding only involves valence
electrons (outermost electrons in the last - Positive charge was localised in the nucleus of
energy level) an atom
- Did not explain why electrons remain in orbit
around the nucleus
Bohr’s Planetary Model (1915) Schrodinger’s Quantum Model (1926)
- Energy orbit the nucleus in energy “shells” - Electrons move in waves, not set paths around
- Electrons travel within energy level without the nucleus
losing or gaining energy - Electrons can behave like atom and an electron
- Greater distance from the nucleus = greater simultaneously
energy - “Cloud of probability” illustrates the
possibilities of where the electron may lie as
- Proposed stable electron orbits; explained the their distinct position is unknown
emission spectra of some elements
- Move of electrons within orbit shells should - Electrons do not move around the nucleus in
emit energy and collapse into the nucleus orbits but in clouds where their positions are
(model did not work well for heavier atoms) uncertain
U1L2: The Quantum Mechanical Model
,Bohr Diagram show energy levels that electrons occupy
- Higher level = higher energy
Absorption: when elements are “excited” electrons can
absorb wavelengths and move to a higher energy level
Emission: after “excitation” electrons emit specific
wavelengths of light and go back to lower energy
positions
- As hydrogen electrons absorb energy, they
jump into an outer ring
- As they lose energy, photons of light are
emitted with a specific amount of
energy/specific wavelength which causes them
to fall back into an inner ring (energy loss).
- Sum of visible photon appears as a single color
until they are separated through a prism
- Each element has a unique emission spectrum
Absorption and Emission Spectra
- Every element is unique because of the unique
electrons in its orbitals
Behavior of Electrons
- Particle Wave Duality: electrons behave like
both particle and a wave
- If a particle passes through a small slit, it
would be detected in a small region on the
other side of the slit
- If a wave passes through a small slit, it would
spread out across a larger region on the other
side of the slit
If there are two slits, waves will experience an
interference pattern
Heisenberg’s Uncertainty Principle:
- Due to its wave and particle nature, we can
never know both the position and speed of an
electron
Orbits Orbitals
- Electrons orbit in a rigid structure (line/circle) - Regions of space around the nucleus
- Simple planar representation of an electron (electrons)
- Lines in an orbit are distinct - 3D model (involves a cloud)
, - 2D model - Orbitals do not have discrete/clear lines
- Electrons behave like a particle (has a distinct - Probability cloud (electrons can be anywhere
position) within the cloud)
- Electrons behave like particles but has a wave
function
U1L3: Quantum Numbers
Heisenberg’s Uncertainty Principle
1. The velocity of an electron is related to its wave “The more accurately we know the position of an
nature electron, the less accurately we can know its velocity”
2. The position of an electron is related to its
particle nature
Probability Distribution Map:
- How likely it is for an electron to exist in a
certain place
- Darker regions = higher probability
- Lighter regions = lower probability
Orbitals: probability distribution map for electrons
Nodes: Points where electrons cannot be located
- More nodes, higher energy
1. s subshell: 1 orientation
2. p subshell: 3 orientations
3. d subshell: 5 orientations
4. f subshell: 7 orientations
Quantum Numbers:
- A set of four numbers used to describe a
specific electron within orbitals (probability
clouds)
- Each electron has its own specific set of
quantum numbers
1. Principal Quantum Number (n)
2. Angular Momentum Quantum Number (l)
3. Magnetic Quantum Number (ml)
4. Spin Quantum Number (ms)
The Principal Quantum Number (n) The Angular Momentum Quantum Number (l)
- Overall size and energy of an orbital - Shape of the orbital
- Distance of the electron from nucleus - Subshells within the main energy level
- Closer to nucleus = lower energy - Values (n - 1): l = 0, 1, 2, …
- Values: n = 1, 2, 3, 4, …
