LC

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 (CC)

      • 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

Bonding and Hybridization