SCH4U - Unit 1
Redox Reactions
In redox reactions, the charge of elements change
All single replacement reactions are redox
Finding oxidation and reduction in an equation:
1- Assign oxidation numbers to all the atoms
An atom in its elemental form has an ON = 0
A monoatomic ion has an ON = ionic charge
The sum of ON values for atoms in a molecule = 0
The sum of ON values in a polyatomic ion = ionic charge
Follow the group-specific rules in this order, as applicable
Oxygen: -2 (except for -1 in peroxides)
Hydrogen: -1 with metals, +1 with nonmetals
Halogens: -1 with metals and nonmetals, but will be positive when with oxygen
Alkaline Earth Metals: +2 in all compounds
Alkali Metals: +1 in all compounds
Use the group-specific rules first, then apply all remaining oxidation numbers so as to balance the oxidation numbers in a given atom or molecule
Assign oxidation numbers as if it were for only one of that atom, not the total number of that atom
Only consider subscripts and coefficients when considering oxidation number balancing
2 - Determine which elements are being oxidized, and which ones are being reduced
Oxidation - The loss of electrons, and thus becoming more positive
Reducing Agent - The particle that causes another particle to be reduced by losing electrons, which will become more positive on the reactant side
Reduction - The gain of electrons, and thus becoming more negative
Oxidizing Agent - The particle that causes another particle to be oxidized by gaining electrons, which will become more negative on the product side
To balance redox reactions, write two separate half reactions with only the elements that have been oxidized and reduced, and balance separately
Balance all elements that are not hydrogen and oxygen using coefficients
Balance all oxygen atoms by adding H2O molecules, where each molecule accounts for one oxygen atom
Balance all hydrogen atoms by adding H+, where each ion accounts for one hydrogen atom
FOR BASES ONLY Based on how many H+ ions are present, add the same number of OH- ions to each side. Combine the all possible H+ and OH- ions to form water molecules
Count the total electric charge on each side from the ions. Add electrons (e-) to one side to balance the charges
Multiply both half reactions by coefficients so that they each have the same number of added electrons
Put the two half reactions together by putting both reactants on one side and both products on the other
Cancel the electrons and combine any repetitive molecules, such as H2O
Quantum Numbers
Quantum numbers are determined by the last orbital and electron placement of an element
Principle Quantum Number “n” - Whole numbers starting at 1 that represent increasing energy levels
As energy levels increase, the level is farther away from the nucleus
Energy levels go from level 1 to level 6, which indicates quantum number “n”
Second Quantum Number “l” - The angular momentum number that gives us the shape of an atomic orbital
s - spheres - 2 electrons, 1 orbit
Located in the alkali/alkaline earth metals
l = 0
p - dumbbell - 6 electrons with 3 orbits
Located on the right side of the table
l = 1
d - flower - 10 electrons with 5 orbits
Located in the transition metals
l = 2
f - complex - 14 electrons, 7 orbits
Located in the lanthanide and actinide series
l = 3
Third Quantum Number “ml” - The orientation of an atom
Depends on the last orbital and specific orbit the last electron is placed in
s - 0
p - -1 0 1
d - -2 -1 0 1 2
f - -3 -2 -1 0 1 2 3
Fourth Quantum Number “ms” - The two possible electrons spins, either +1/2 or -1/2
If the last electron only half-fills the orbital (↑_), it is +1/2
If the last electron fills the orbital (↑↓), it is -1/2
Spectroscopy
Photon - A particle of light that behaves both as a particle and a wave
Use the equations E = hv and c = λv to calculate properties of light
E = Energy of a Photon = Joules
h = Planck’s Constant = 6.626 × 10-34Js
c = Speed of Light = 2.998 × 108 m/s
v = Frequency = Hz = cycles/second
λ = Wavelength = nanometres
One metre is 109 nanometres
High energy light, such as blue light and ultraviolet, has shorter wavelengths
Low energy light, such as red light and infrared, has longer wavelengths
Photoelectric Effect - The observable spark that occurs when a certain frequency of light strikes an object and ejects electrons
The higher the frequency, the more electrons would be ejected
Electrons highly attracted to the nucleus require more energy to be ejected
The required energy helps identify the orbitals that the electrons are in
It also correlates to the visible colours that are shown when energy is absorbed or emitted
Photoelectron Spectroscopy - An analytical technique that shows how much energy is required to remove an electron by sending X-rays and UV light towards a compound
Ionization Energy - The binding energy of electrons; how much energy is required to remove an electron from the nucleus
Ionization Energy (Binding Energy) = Photon Energy - Kinetic Energy
I.E. = E (or hv) - K.E.
Elight = Ebinding + Ekinetic
On a PES chart, greater ionization energy indicates a lower energy level and orbital, because they are closer to the nucleus
When an electron is bound closely to the nucleus, more energy is required to remove an electron
The height of the peaks on a PES chart indicate how many electrons have been removed
Beer’s Law - How much light a sample absorbs is proportional to the concentration of the sample
A = εBC
A - absorbance (no units)
ε - molar absorption coefficient (a constant that is unique to each solution, measured in 1/(mol/L * cm)
B - length of path or cuvette in cm
C - concentration of the solution in mol/L
UV/Vis Spectroscopy - The process of shining wavelengths of light in the UV and Visible range to see which wavelengths can excite electrons, and thus which ones can be absorbed or emitted
It looks at electrons that may be promoted/excited (electronic transitions)
When a compound absorbs one wavelength of light in the visible light spectrum, it will correlate to a colour
The colour of light that the compound emits is not absorbed by the compound, but reflected, and thus the visible colour
The light absorbed and the light emitted, if on the visible light spectrum, will be complementary colours
Absorbs red = emits/appears green, and vice versa
Absorbs orange = emits/appears blue, and vice versa
Absorbs yellow = emits/appears violet, and vice versa
Infrared Spectroscopy - The use of energy to study the stretching and vibration frequencies within a molecule, that depends on different bonds and functional groups
Higher frequencies of energy generally correlate to a stronger bond
In an IR chart, higher peaks (troughs) indicate a greater reliability
Alcohols, Carboxylic Acids (O-H)
Amines, Amides (N-H)
Non-Aromatic Compounds (C-H)
Aldehydes (C-H)
Alkynes (C≡C)
Various (C=O)
Ketones (C=O, where C is single bonded to 2 R’s)
Aldehydes (C=O, where C is single bonded to 1 R and 1 Hydrogen)
Carboxylic Acids (O=C-O-H, where C is single bonded to 1 R)
Esters (O=C-O-R, where C is single bonded to 1 R)
Amides (O=C-N, where C is single bonded to 1 R, and N is single bonded to 2 R’s)
Acid Chlorides (O=C-Cl, where C is single bonded to 1 R)
Alkenes (C=C)
Various (C-O-R)
Various (C-Cl)
Microwave Spectroscopy - The use of energy to determine bond rotations, which can indicate if a bond is a single, double, or triple
Nuclear Magnetic Resonance - Detects signals from protons in the nucleus, which indicates how many atoms are present in a molecule
It is especially effective at determining organic compounds
It determines the number and location of hydrogen atoms in a molecule
Mass Spectrometry - Breaks an atom into fragments to determine the mass and abundance of isotopes
Electron beams break apart the molecule, and the cations are deflected off of a magnetic field, which indicates the isotope
Diatomic elements will show two very distinct peaks - one for the atom itself, and one for the diatomic molecule