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

Bonding and Hybridization

  • Localized/Valence Electron Theory - All bonds are localized bonds formed between electrons from different atoms that attracted to the other nuclei, thus causing orbitals to overlap

    • Includes Lewis structures, VSEPR (Valence Shell Electron Pair Repulsion), and hybridization

    • Does not explain properties and electronic transitions accurately

  • Lewis Structures - Diagrams that describe the arrangement of valence electrons, but not the shape or geometry

  • VSEPR Theory - Each atom in a molecule will minimize the repulsion of valence electrons by forming an ideal geometry

    • Valence Shell Electron Pair Repulsion

    • Electron Geometry - The shape of all electron groups around the central atom

    • Molecular Geometry - The shape of all electron pairs around the central atom

      • Symmetrical molecular geometries cancel out the charges, so the net permanent dipole moment is zero (nonpolar), and the entire molecule has no charge

        • The centre of negative charge lies on the centre of positive charge

        • Symmetry requires all bonds to have the same difference in electronegativity

        • If all individual bonds are non-polar, the entire molecule will be non-polar, regardless of the shape

          • There can still be a total charge, like in the case of an ion

      • Asymmetrical molecular geometries do not cancel out the charges, so there is a permanent dipole moment (polar), and the entire molecule has a slight charge

        • The centre of negative charge lies underneath the centre of positive charge

      • Ionic compounds always form the same shape, so their molecules do not need to be considered for electron and molecular geometries

      • Molecular shapes can determine dipoles and functions

      • Because all electrons have the same negative charge, they will repel each other to maintain distance in a 3D space

Electron Groups - Electron Geometry (Hybridization Orbitals)

Molecular Geometry - (Central Atom, Bonding Groups, Non-Bonding Groups)

Approximate Bond Angle

Diagram

2 - Linear (sp)

Linear (AX2)

180

3 - Trigonal Planar (sp2)

Trigonal Planar (AX3)

120

3 - Trigonal Planar (sp2)

Bent (AX2E)

120

4 - Tetrahedral (sp3)

Tetrahedral (AX4)

109.5

4 - Tetrahedral (sp3)

Trigonal Pyramidal (AX3E)

109.5

4 - Tetrahedral (sp3)

Bent (AX2E2)

109.5 (104.5 for water)

5 - Trigonal Bipyramidal (sp3d)

Trigonal Bipyramidal (AX5)

90 and 120

5 - Trigonal Bipyramidal (sp3d)

Seesaw (AX4E)

90 and 120

5 - Trigonal Bipyramidal (sp3d)

T-Shaped (AX3E2)

90 and 120

5 - Trigonal Bipyramidal (sp3d)

Linear (AX2E3)

180

6 - Octahedral (sp3d2)

Octahedral (AX6)

90

6 - Octahedral (sp3d2)

Square Pyramidal (AX5E)

90

6 - Octahedral (sp3d2)

Square Planar (AX4E2)

90

6 - Octahedral (sp3d2)

T-Shaped (AX3E3)

90

6 - Octahedral (sp3d2)

Linear (AX2E4)

180

Highlighted geometries will have NO permanent dipole moment

  • Bond types

    • ∆ EN > 2 = ionic

    • 2 > ∆EN > 1.7 = ionic if it includes metal, covalent if no metal is present

    • 1.7 > ∆EN > 0.5 = polar covalent

      • Not evenly shared

    • ∆EN < 0.5 = nonpolar covalent

      • Very evenly shared

  • Physical bonds (intermolecular forces)

    • London Dispersion Forces - The temporary dipole created by electron movement, found in all molecules

    • Dipole-Dipole Forces - The forces between molecules with permanent dipoles, where the negative poles attract to the positive poles

    • Hydrogen Bonding - A type of dipole-dipole force, which results from a hydrogen atom being bonded to a highly electronegative atom such as nitrogen, oxygen, or fluorine

    • Ion-Dipole Force - The attraction between an ion and a polar molecule that causes them to dissolve

  • Formal charge checks will indicate if a molecule can break the octet rule

    • FC = Total valence electrons - bonded pairs - individual non-bonding electrons

  • Hybridization Theory - The blending of orbitals to create new orbitals with the characteristic of the parent orbitals, that also have a better surface area and overlap for bonding with nearby atoms

    • The number of orbitals an atom needs is the number of orbitals it must hybridize

      • This number of orbitals needed is dependent on the amount of electron groups/electron geometry

      • One group can be: a single bond, double bond, triple bond, or a lone pair

        • In a double or triple bond, only one of these bonds (sigma bonds) will exist in a hybridized orbital, and the rest (pi bonds) will exist in non-hybridized p orbitals

        • If all of the bonds were in hybridized orbitals, they would repel one another

      • Lone pairs need an orbital, even if they are not bonding

    • The wave functions of each individual orbital overlap to create constructive interference

    • If the amount of hybridized orbitals needed are filled unequally (ex. the first is fully filled, while the last is empty), atoms can become “excited” and promote an electron from the original s orbital to the original, empty p orbital

      • Promotions of electrons can also help create an additional hybridized orbital by promoting an electron in an s or p orbital to a d orbital (used for trigonal bipyramidal or octahedral geometries for elements that are in the third period or below)

Molecular Orbital Theory

  • Molecular Orbital Theory - A method for determining the molecular structure of a compound in which electrons are not assigned to individual bonds between atoms, but are treated as moving under the influence of nuclei in the whole molecule

    • The electrons are free to travel around any atoms in the structure (they are not localized)

    • Molecular Orbital - An orbital formed from the interactions between two atomic orbitals that are merged when bonding, though they may or may not be occupied by electrons

      • Formed when the symmetries of atomic orbitals are compatible, there is a significant region of atomic orbital overlap, and the orbitals are close in energy levels

  • Bonding Orbitals - The area between two positive nuclei that have a high electron probability, represented by sigma (σ) in s and p levels, and pi (π) in p levels

  • Anti-Bonding Orbitals - The area away from the two positive nuclei that have a low electron probability, represented by sigma star (σ*) in s and p levels, and pi star (π*) in p levels

  • Bond Order - The calculation that tells us how stable a molecule is, and how likely it is to exist

    • Bond Order = (# of bonding e- - # of antibonding e-)/2

    • A bond order can never be negative

    • A bond order of 0 means the molecule is not observed

    • A bond order less than 1 is usually very unstable, because the bonds are comparatively weaker

    • A fractional bond order greater than one means the molecule is likely has a resonance structure

    • A whole number bond order is stable, and usually correlates with the number of bonds

    • A higher bond order correlates with a greater strength and a shorter bond

  • Molecular Orbital Diagram - A diagram which represents the bonding and anti-bonding orbitals of a molecule as a relationship between the non-hybridized orbitals of two atoms, compared with the potential energy of each

    • The higher the position, the greater the energy

    • It begins with 1s at the bottom, and moves upwards

    • Lower energy MO are more stable, and higher energy MO are more unstable

    • Bonding orbitals reduce energy, because they shield the positive nuclei, thus reducing repulsion and increasing the bond strength

    • Anti-bonding orbitals increase energy because they are constantly repelling and weakening the bond strength

      • Therefore, anti-bonding, star orbitals tend to be higher on the diagram than the bonding, non-star orbitals

  • P-orbital Orientation

    • The p orbital can lie along the x, y, or z axis, as denoted by npx, npy, and npz

    • When two atoms with p orbitals are in contact, the x and y p orbitals will overlap twice and create pi bonding orbitals and anti-pi bonding orbitals

    • pz orbitals, when overlapping with either an s or pz orbital, will only cross over in one spot and create stronger sigma bonding orbitals and sigma anti-bonding orbitals

    • In elements from Li2 to N2, the 2s and 2p orbitals are very close and influence each other, so through sp mixing, the 2pz orbital is weaker, and will have a higher energy than 2px and 2pz

  • Magnetism

    • Paramagnetic - A substance with at least one unpaired electron in its atomic or molecular orbitals that has a slight attraction to magnetic fields

    • Diamagnetic - A substance with all electrons paired in its atomic or molecular orbitals that has a slight repulsion to magnetic fields

  • Non-Bonding Orbitals

    • In heteronuclear diatomic molecules, the orbitals may not overlap symmetrically

    • The orbitals that are not involved with the bonding are placed in a non-bonding molecular orbital that is on the same level as the atomic orbital energy level

    • The orbitals that do bond are connected as normal to create a molecular orbital and labelled appropriately

Energy In and Energy Out

  • Energy In - The amount of energy required to break the chemical bonds of the reactant molecules to create a chemical reaction

  • Energy Out - The amount of energy released when the products are made through the forming chemical bonds

  • The energy required to break bonds is also the energy that is released when those bonds are formed

    • Weaker bonds require less energy to be broken and release less energy when formed

    • Strong bonds require more energy to be broken and release more energy when formed

    • Sigma bonds are strong, and have a lot of energy

    • Pi bonds will still add more energy to the bond, but not as much as the sigma bond

  • The equation ΔH = (Energy In - Energy Out) gives us the change in heat energy for a reaction, which will indicate id a reaction is spontaneous or not

    • Spontaneous Reaction - The reaction releases more energy than it requires to start (ΔH is negative)

      • Weakly bonded reactants, such as those in gasoline, require little energy to be broken, and then form strongly bonded products, thus releasing lots of energy

    • Not Spontaneous Reaction - The reaction releases less energy than it requires to start (ΔH is positive)

  • Calculating Bond Energies and Heat

    • Find all the reactants and products, and draw the appropriate Lewis Structures for each

    • Given the bond energies in a chart, solve the bond energy for each molecule

      • Multiply each type of bond in the molecule by the total number of times it is present

        • The number of times present = (the amount of times in the molecule) x (the amount of molecules in the reaction)

          • Ex. H2O has 2 O—H bonds, but 4 H2O molecules would have 8 O—H bonds

      • Add up each of these products, though if there is only one type of bond, there should only be one product

    • Add up the total energy for each molecule on the reactant side, and do the same on the product side

    • Substitute into the equation to find ΔH, and if it is spontaneous

Solid States, Bonding, and Coulomb’s Law

  • Physical Bond - An intermolecular bond that holds molecules together, such as London dispersion forces, dipole-dipole, hydrogen bonding, and ion-dipole

  • Chemical Bond - An intramolecular bond that holds atoms within a molecule together, such as covalent/molecular, ionic/electrostatic, and metallic bonds

  • Amorphous Solid - A solid without a well-defined 3-D unit structure

    • The strength depends on the intermolecular and/or intramolecular bonds that hold it together

  • Crystalline Solid - A solid made of repetitive units that may be visible on the macroscopic scale

    • In order from weakest to strongest, they are:

      • Molecular solids

      • Metallic solids

      • Ionic solids

      • Covalent network solids

  • Molecular Solid - A solid that consists of atoms or molecules held together by intermolecular forces with weak bonds

    • Held together by weak physical bonds

    • Low melting point

    • Brittle

    • Non-conducting

  • Metallic Solid - A solid that consists of a positive core of atoms held together by a surrounding sea of electrons with metallic bonds

    • The transfer of electrons creates very strong bonds

    • High melting point

    • Not brittle and very malleable, because the electrons are delocalized and can just “slide” across the whole solid without breaking bonds

    • Conducts electricity due to moving electrons

    • The metal atoms behave like positive ions

    • Substitutional Alloy - A combination of metals where the atoms of one metal are substituted by the atoms of another

    • Interstitial Alloy - A combination of metals where the atoms of one metal occupies the interstitial spaces in the lattice structure of the other metal

  • Ionic Solid - A solid that consists of cations and anions held together by the electrical attraction of opposite charges with intramolecular forces

    • High melting point

    • Hard and brittle

    • Non-conductive as a solid, but conductive as a liquid because the ions become dissociated

    • Ions with opposing charges neutralize the solid, so when they are broken along a face, the like charges line up and repel to create “flakes”

    • Coulomb’s Law - The closer two charges are, the stronger the force is between them

      • Ionic compounds have maximized attractive forces between cations and anions, where:

        • F=\frac{kQ_1Q_2}{r^2}

        • E=\frac{kQ_1Q_2}{r^{}}

        • A negative force means the ions are attracted to one another at an ideal distance for bond lengths, and it is stable

        • When the ions are too close, the positive nuclei repel, so the total energy is very high, and it is unstable

        • When the ions are too far, there is very little attraction between them, so energy is low

  • Covalent Network Solid - A solid that consists of atoms held together chemically in large networks or chains by intramolecular chemical covalent bonds

    • Very high melting point

    • Hard

    • Usually non-conductive

    • There are large numbers of chemical covalent bonds in multiple different directions

  • Lattice Structure Shapes

    • Lattice structures are measured in unit cells, which is a repeating pattern

    • Cubic unit cells are the most common and stable

      • Simple Cubic Unit Cell - The lattice points are at the corners

      • Body-Centred Cubic Unit Cell - The lattice point is at the centre of the cell, as well as the corners

      • Face-Centred Cubic Unit Cell - The lattice points at the centre of each face and in the corners