HL Chemistry: Unit 2

Light

Believed to exist in tiny "packets" or photons, known to exhibit properties of both particles and waves, a form of electromagnetic radiation

Median

Matter needed for waves like sound to propagate

Why are wavelengths shorter than visible light dangerous?

More energy

Greater potential to ionize (break bonds in atoms, lifting it from a bond)

Wave propagation

The way waves travels

Wavelength (λ/lambda)

Distance over which the waves shape repeats, from crest to crest or trough to trough

Has periodicity

frequency

Number of waves that occur per second

Equations to calculate wave properties

c = λf

- Inverse relationship between wavelength and frequency

ΔE = hf

- ΔE: Change in energy

- h: Planck's constant (6.63 * 10-34 JS)

Diffraction grating

White light consists of all wavelengths of light, able to diffract light using a prism and able to see the wavelengths contained within

The diffracted rays from top to bottom go from longest to shortest wavelength

Continuous spectrum

Contains all wavelengths in a given region

Ex. White light, sunlight

Line spectrum

Lines of light or no light at specific wavelengths

Each line is ONE WAVELENGTH, ONE COLOUR, ONE FREQUENCY, ONE AMOUNT OF ENERGY

The emission and absorption spectrum of any given element have lines in identical locations

Emission spectrum

A gas is given energy until it glows, the emitted light from the hot gas is then passed through a prism, only specific wavelengths of light are emitted by the atoms in the gas

Absorption spectrum

A source of white light is passed through cold gas source then through a prism, atoms in the cold gas blocks/absorbs specific wavelengths of light

Ground State

State electrons in stable atoms and ions are in

Exciting Electrons

When energy is added, the electron is excited/promoted (heat, electricity, visible light) to a higher energy level, but the electron (the atom) does not want the extra energy. Promoted electrons will "drop" back down to lower energy levels until they reach the ground state, and the excess energy is "emitted" in the form of light.

- The fall/excitation of a specific distance will always emit/absorb the same specific amount of energy, we can only see some of this emission

- Greater the distance an electron jumps down, the greater the light energy emitted

Quantized/Discrete

Electrons are only found in specific/fixed energy levels in the atom

Specific/Limited number of amount of energy that the electron can lose or gain

Hydrogen spectrum

Ultraviolet region: n = 1

Visible region: n = 2

Infrared region: n = 3

Convergence

Lines get closer together as energy increases, frequency increases and wavelength decreases

Decreasing distance between lines of light on the spectra

Energy levels get closer together (converge) as the energy level number increases

Ionization energy

As n → ∞ reaches the limit to which convergence occurs, indicates the energy required to completely remove an electron from the atom

Graph and extrapolate to when change in energy lines converge to when difference is 0

Heisenberg's uncertainty principle

Within a given energy level, the exact position of an electron cannot be determined, electrons exist in a region of probability

Schrodinger's equation

Predicted that electrons are waves (differential calculus)

Electrons are also particles because they have mass and charge

Allows us to find the probability of finding an electron with a particular energy in a particular location

Answer to equation → electrons can be described by 4 distinct quantum numbers

Orbital

Areas of probability

Sub-energy level

Distance to nucleus

Principle quantum number

energy level (n)

Electron's average distance from the nucleus

Equal to number of different orbital types that exist at each energy level

Angular-momentum quantum number

shape/secondary distance of orbital (ℓ)

Maximum: ℓ = n - 1

The higher the ℓ value, the further it is from the nucleus (radius increases)

Each value of ℓ corresponds to the specific shape of the given orbital

"s" orbital

ℓ = 0

Spherical shaped

Angular-momentum

Can hold up to 2 electrons

As the value of n increases, the radius of the s-orbital increases

"p" orbital

ℓ = 1

Barbell shape

3 different orbitals/orientations (Px, Py, Pz) corresponding to 3-dimensions

Each orientation (orbital) can hold up to 2 electrons

Maximum of 6 electrons in the p (ℓ=1) sub-shell

All equally distanced from the nucleus

"d" orbital

ℓ = 2

Various shapes

5 different orbitals/orientations

Each orientation can hold up to 2 electrons

Maximum of 10 electrons in the d sub-shell

"f" orbital

ℓ = 3

Various shapes

7 different orbitals/orientations

Each orientation can hold up to 2 electrons

Maximum of 14 electrons in the sub-shell

Magnetic quantum number

orientation of orbital (mℓ)

Due to the angular momentum (rotation around the nucleus) electrons produce a magnetic field

The actual value of mℓ is determined by the chosen value of ℓ

mℓ = (-ℓ) ... 0 ... (+ℓ)

- ℓ = 1 → mℓ = -1, mℓ = ∅, mℓ = +1 (Px, Py, Pz)

Degenerate orbital

Orbitals within a sub energy level

Spin quantum number

(ms)

Refers to the characteristic of the electrons not the characteristics of the orbit

Electrons behave as though they are spinning

ms = +½ (clockwise spin)

ms = -½ (counter clockwise spin)

The first electron in an orbital always started with a positive ms value

Aufbau principle

Each additional electron added will occupy the next lowest available orbital

Overlap in some sub-shell due to convergence

Pauli exclusion principle

No electrons in an atom can have the same 4 quantum numbers

Hund's rule

Electrons in p, d, f orbitals can't be paired until each of the orbitals at the same sub-energy level contain at least 1 electron

Abbreviated notation of electron configuration

Start with the closest noble gas

Finish the configuration from there

Ex. Carbon 1s2 2s2 2p2 → [He] 2s2 2p2

Isoelectronic

Different elements with the same electronic configuration

To gain stability, atoms become isoelectronic with noble gases

Ex. F- (ion) 1s2 2s2 2p6 → Ne 1s2 2s2 2p6

Ions

Full, empty, or ½ filled sub-energy levels are also stable electron configurations

Ex. Zn: [Ar] 4s2 3d10 → it takes too much energy to lose 12 valence electrons

- Most d-block elements loses 2 of its "s" valence electrons to form form a 2+ ion, emptying its highest energy level

Ions Exceptions

Both copper and chromium "promote" one 4s electron a short distance to the 3d to gain half-filled and full sub-shell stability

3d and 4s shells are so close to each other, it requires little energy to be promoted, and 3d is almost filled, so it takes an electron from the 4s (it gets promoted), to give 1 half orbital and 1 full orbital

Copper: [Au] 4s2 3d9 → 4s1 3d10

Patterns

It takes more energy to remove an electron from a lower sub-energy level than a higher one (ionization energy)

- Lower energy levels are closer to the nucleus

- Electrons have less energy in lower levels

Removing an electron to reveal a stable state (full sub-energy level or ½ filled sub-energy level) is easier than removing an electron from a stable state

A more positive ion has a stronger attraction to (negatively charged) electron than a less positive ion

Ex. Despite the fact that B has more protons (nuclear charge) than Be, its first removed electron is further from the nucleus (and more shielding) and reveals a full sub-energy level → it is easier to remove B's 1st electron

As you go up in energy, difference between energy decreases → convergence

Transition Metals

Has an atom or at least one stable ion with incompletely filled d-orbitals (new)

All transition metals are in the d-block but not all d-block elements are transition metals

Ex. Atoms → Zn: [Ar] 4s2 3d10, Cu: [Ar} 4s1 3d10 (full d-sub-energy level → not a transition metal), Sc: [Ar] 4s2 3d1 (incomplete d-sub-energy level)

Ex. Ions → Zn2+: [Ar] 3d10 (full d-sub-energy level → not a transition metal), Sc3+: [Ne] 3s2 3p6 (empty d-sub-energy level)

Most transition metals can form a stable 2+ ion

Transition metal patterns going across period 4 d-block

Sc → Mn

- Max charge possible increases

- Number of possible ions increases

- Because there are more d electrons to lose

Fe → Zn

- Larger nuclear charge, more difficult to lose electrons

- Stability of paired d electrons

Oxidation state

Redox term: A reversible chemical reaction where one reaction is an oxidation and the reverse is a reduction

Often the same as charge

Charge (sign, #) vs. Oxidation state (#, sign)

Magnetic Properties

Unpaired d-electrons allow transition metals to have magnetic properties