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Principles of Chemistry: The Molecular Science

Chapter 5

  • p. 143-145 Intro to acids (contributes H+) and bases (contributes OH-)
  • p. 146 -148 Neutralization reactions. Acid + Base –> H20 + Salt
  • p. 150 Gas-forming exchange reactions, e.g.:
    • calcium carbonate wtih hydrochloric acid yields gaseous CO2 and aqueous calcium chloride aka
    • CaCO3 + 2 HCl (aq) –> CaCl2 (aq) + intermediate H2CO3 –> CaCl2 (aq) + H2O (liquid) + CO2 (gaseous)

Redox Reactions, p. 151-158

  • See JH handwritten notes 8/06/2020 for redox diagram which draws on material from: Alberts p.71, Sadava p.167, MSJ p. 151-155, especially the table on MSJ p. 155.
  • See also simplified redox rules specific to Organic Reactions per L.G. Wade, p. 465:
    • Oxidation (falling down rusting-red): addition of O or O2; Addition of X2 where X = halogen; loss of H2
    • Reduction (rising up blue-heaven): addition of H2, addition of H-; loss of O2, loss of O; loss of halogens aka loss of X2
    • Neither Oxidation or Reduction reaction
      • addition or loss of proton H+
      • addition or loss of hydroxide ion -OH
      • addition or loss of water molecule H2O
      • addition or loss of Hydrogen+Halogen aka HX
  • p. 152-154, common oxidizing and reducing agents
  • LEO = Oxidation causes Loss of Electrons aka combines with oxygen
  • GER = Reduction causes Gain of Electrons aka loss of oxygen from the molecule aka reduces the postive charge.
    • eg, see what happens to Hydrogen in magnesium reaction on p. 646. The “positive +1 charge” charge of aqueous HCl reduces its positivity to 0 in the resulting H2 gas.

The Four Quantum Numbers: n, ℓ, mℓ, ms – MSJ Chapter 7.5, p. 237-242

  • List of four quantum numbers is here and a general description of atomic orbitals is on this wiki page.

Principal quantum number “n”

  • Values = {1, 2, 3, 4,…} aka which electron shell, principal energy level, or “size” of the electron orbital. See also the wiki article.
  • n=1 has 1 subshell s with 2 electrons.
  • n=2 has 2 subshells. 2s2 + 2p6 = total of 8 electrons in this shell.
  • n=3 has 3 subshells. 3s2 + 3p6 + 3d10 = total of 18 electrons in this shell.
  • Etc.

Azimuthal quantum number “ℓ” aka subshell

  • Specifies the quantum number for an atomic orbital which determines its orbital angular momentum as well as the shape of the electron orbital: s, p, d, f, g, etc…
  • ℓ=0 = subshell s, maximum of 2 electrons. Spherical shape.
  • ℓ=1 = subshell p, maximum of 6 electrons. Dumbell shapes, one for each axis: px, py, pz
  • ℓ=2 = subshell d, maximum of 10 electrons. For shape, look at the third row of this Wiki image
  • ℓ=3 = subshell f, maximum of 14 electrons. For shape, look at the fourth row of this Wiki image

Magnetic quantum number “mℓ”

  • Values of mℓ; are dependent on the subshell / azimuthal quantum number l and describes the orientation of each atomic orbital “lobe”. See sample values in the table in this wiki article or in MSJ p. 240, Table 7.3.
  • For subshell s where ℓ=0, there is only one value for mℓ which is mℓ = 0. Holds 2 electrons
  • For subshell p where ℓ=1, there are 3 possible values for mℓ which is mℓ = {-1,0,+1}. Holds 6 electrons.
  • For subshell d where ℓ=2, there are 5 possible values for mℓ which is mℓ = {-2,-1,0,+1,+1,+2}. Holds 10 electrons.
  • For subshell f where ℓ=3, there are 7 possible values for mℓ which is mℓ = {-3,-2,-1,0,+1,+1,+2,+3}. Holds 14 electrons.

Spin quantum number “ms”

  • Describes the spin of an electron, values of +1/2 or -1/2, and at most two electrons can occupy an atomic orbital. If they are paired, then the electrons must have opposite spin. Wikipedia uses the variable name s for spin quantum number but textbooks and Khan Academy use the ms notation instead.
  • Chapter 7, sections 7.9 - 7.12, p. 254 - 269
  • Examples of trends across periodic table: Isoelectricity / ion electron configs, size of atoms, size of associated ions, ionization energies, and electron affinities.

Isoelectronic Ions and Atoms

  • MSJ Chapter 7.8 Ion Electron Configurations, p. 251-252
  • Isoelectronic atoms and ions have the same number of electrons and in the same orbital configurations.
  • Group 1-3 metals do this by becoming positive cations (aka lose electrons).
  • Group 14-17 (aka 4A - 7A) nonmetals do this by become negative anions (aka gain electrons). They ultimately have the “desired” octet of valence electrons equivalent to the noble gas (Group 18) in their associated period. For Group 14-17 ions, they stay in the same period.
  • For Group 1-3 metals, these atoms become cations by losing electrons to get “demoted” back to the noble gas in the previous period. See illustration on p. 251; all the species in the same row have the same color and are isoelectronic with each other.

Electron Affinity

  • MSJ Chapter 7.12, p. 260 - 261.
  • Not to be confused with electronegativity, even though they seem similar in concept and share the same trends on the periodic table (aka highest in upper right near Fluorine, lowest near lower left near Francium).
  • Electronegativity is a dimensionless measure whose range was arbitrarily set by Linus Pauling and only relevant within a molecule.
  • In contrast, electron affinity is a directly measured unit of energy released when an electron is added to a neutral atom in the gaseous state to form a negative ion. aka, the amt of energy you need to *add to a anion to get it to give up its extra electron(s) to become a neutral atom.
  • Thus, main group metals in Groups 1 and 2 have low electron affinities because it’s super easy for them to give up electrons b/c they prefer to be positive cations.
  • High electron affinity atoms “want” electrons more, e.g., non metals like Flurorine in Group 17 (7A) with low periods like period 2 or 3. These elements strongly prefer to be negative anions and you need to add a lot of energy to it to have them give up their extra electrons to become neutral atoms.
  • Some good answers on this message board page.

Chapter 8: Covalent Bonding, Intro to Lewis Structures, Hydrocarbons, Molecular Orbital Theory

  • p. 297 - contrast between atomic orbital theory and molecular orbital theory in Section 8.10 to explain paramagnetism in O2 implying unpaired electrons, even though naive Lewis structure predicts diamagnetism
  • p. 277 saturated hydrocarbons like butane VS. 2methylpropane branched chain
  • p. 277 cyclopropane, cyclobutane, cyclopentane, cyclohexane
  • p. 278 C=O is a carbonyl group.
  • p. 278-279 Aldehydes are hydrocarbons that have a -CHO functional group aka an “aldehyde functional group” which necessarily includes a carbonyl.
  • Whereas alkanes ( p.277) are necessarily saturated (aka have maximum # of hydrogens possible aka no double bonds between carbons), alkenes (p. 280) are hydrocarbons that have at least 1 C=C double bond.
  • Alkynes are unsaturated hydrocarbons as well, but unlike alkenes, alkynes have triple bonds between carbon atoms.

Chapter 9: Molecular Structures, VSEPR, Orbital Hybridization, Noncovalent Interactions, and Intermolecular Forces

  • MSJ p. 310, Figure 9.4 table listing names and shapes of common molecular structures around a central atom: linear, trigonal planar, trigonal bipyramidal, tetrahedral, octahedral.
  • p. 316, Figure 9.6 table illustrating more exotic shapes with 5 or 6 shared electron bonds and/or lone pairs surrounding a central atom: trigonal bipyramidal, seesaw, T-shaped, linear, octahedral, square pyramidal, square planar

    Hybridization of orbitals, sigma bonds, pi bonds, steric number

  • MSJ Chapter 9
  • Lone pair = pair of valence electrons that are not bonded/shared covalently with another atom. Eg., ammonia NH3 has steric number of 4 = 3 sigma bonds to hydrogen + 1 lone pair.
  • Valence Shell Electron-Pair Repulsion (VSEPR) theory. See Wiki article and MSJ Chapter 9.2 (p. 309-318)
  • Interesting summary of sp linear, sp2 trigonal, sp3 tetrahedral at this page
  • See also these videos at Khan Academy:

Section 9.5: Molecular Polarity

  • p. 325 caused by both difference in electronegativity between different atoms in a molecule and the organization of charges in the molecular structure
  • For example, compare CF4 is nonpolar because of molecular symmetry while CHF3 is polar because the hydrogen atom throws off tetrahedral symmetry - p. 327

Section 9.6: Noncovalent Interactions

  • Noncovalent interactions refer to all forces of attraction apart from covalent bonds, ionic bonds, and metallic bonds. p. 329
  • Noncovalent bonds can refer to both intramolecular and intermolecular forces
  • Primary categories: London Forces, Dipole-Dipole Forces, and Hydrogen Bonds

Chapter 12: Chemical Kinetics, Rate Law, Mechanisms, Catalysts, Enzymes

  • Example of methane combustion in bunsen burner with and without fire. CH4 + 2 O2 –> CO2 + 2 H2O - p. 415
  • Four factors that impact speed of reaction when all reactants and products are in the same phase (eg all are gases or all are in solution) aka in a homogeneous reaction:
    1. molecular structure and bonds
    2. concentration
    3. temperature
    4. presence and concentration of catalyst (if any)
  • In addition, if we are talking about input/output that are in different phases aka a heterogeneous reaction:
    1. area of reaction surface
    2. nature of reaction surface

Example of crystal violet Cv+ reaction rate (p. 417)

  • Colorful crystal violet dye (abbreviated Cv+) combining with hydroxide ion (OH-) to produce colorless product CvOH\ on p. 417-419. The full formula is crystal violet chloride combining with sodium hydroxide to form the CvOH and sodium chloride:
    • CvCl (ions in solution) + NaOH (ions in solution) –> CvOH (colorless product) + NaCl (ions in solution)
  • The notation for “concentration of” uses brackets. e.g., [Cv+] is pronounced “concentration of crystal violet dye”.
  • Average reainction rate is delta [Cv+] / delta time. And since the concentration of Cv+ decreases over time but reaction rate is usually measured as a positive rate, the final equation has a negative in it. Average reaction rate from time t1 to t2 = - Δ[Cv+] / Δt

Example of NO2 –> NO + O2, (p. 419)

  • The crystal violet Cv+ example from p. 417, one mole of each reactant (aqueous Cv+ and aqueous OH-) produces one mole of the colorless product CvOH. A nice one-to-one ratio so the reaction rate is the same regardless of whether we are measuring [Cv+ cation in solution], [OH- anion in solution], or [CvOH molecule]
  • In contrast, consider the reaction 2 NO2 (gas) –> 2 NO (g) + O2 (g). There is a 2:1 ratio between the reactant NO2 and the product O2. This means the reaction rate seems different depending on if we measure [NO2] or [O2].
  • To resolve this, we define reaction rate using stoichiometry with Equation 12.3 (p. 419):
    • Let lowercase a,b,c,d be coefficients and upper case A,B,C,D be chemicals in the following reaction formula: aA +bB –> cC + dD
    • Then the uniform reaction rate is equivalent for the following:
      • Rate = - (1/a)(Δ[A]/Δt) reactant A
      • Rate = - (1/b)(Δ[B]/Δt) reactant B
      • Rate = + (1/c)(Δ[C]/Δt) product C
      • Rate = + (1/d)(Δ[D]/Δt) product D

Sections 12.2 - 12.4: Rate Laws for 0th-, 1st-, and 2nd-order Reactions; Unimolecular, Bimolecular, and Termolecular Elementary Reactions; Ea, exothermic and endothermic rxns(p. 422-436)

  • Constant k is the specific rate a reacion will run at a given temperature, multiplied by a relevant concentration of a reactant or product.
    • Zeroth Order Reaction = k * [A]0 = k * 1 = k. means that reaction proceeds independently of concentration of A.
    • First Order Reaction = k * [A]1 reaction directly proportional to concentration of A
    • Second Order Reaction = k * [A]2 means reaction proportional to square of [A]
    • For more info see this page
  • All reactions can be decomposed to elementary reactions (aka subreactions). which are themselves unimolecular, bimolecular, or termolucular (requiring 3 reactants).
    • Per Wikipedia, “Reactions of higher molecularity [than 3] are not observed due to very small probability of simultaneous interaction between 4 or more molecule.”
    • See this video for a good explanation of how to express rate laws for elementary reactions involving 1, 2, or 3 reactants.
  • Activation Energy (abbreviated Ea) needed to get over the hump plus if the end state is lower energy, the reaction is exothermic. Conversely if final product is higher energy state, then the reaction is endothermic. (p, 434)
  • Intermediate form in the process of moving from reactant to product is called the transition state aka activated complex. (p. 433)
  • See example of transition between cis-2-butene to trans-2-butene on p. 433-434.

Section 12.5: Arrhenius Equation to determine k

  • k = Ae -Ea/RT where:
    • k = rate constant
    • A = Arrhenius frequency factor, which depends on (a) how often molecules collide when all concentratrions are at 1 mol per liter and on (b) whether molecules are properly oriented when they collide
    • Ea = activation energy
    • R = gas law constant, 8.314 joules/ (mol * temperature in degrees Kelvin)
    • T = temperature in degrees Kelvin

Chapter 13: Chemical Equilibrium

Chapters 14-15: Acids, Bases, Buffers

Chemistry notes from Sadava

Chapter 2: Review of General Chemistry

Section 2.1: How does Atomic Structure Explain the Properties of Matter?

  • Review of atomic orbitals p. 24-25
    • An atomic orbital can hold at most two electrons.
    • Remember:
      • First Shell: 1s2 = total of 2 electrons in this shell. Only an s-orbital shape and 1 orbital.
      • Second Shell: 2s2 + 2p6 = total of 8 electrons in this shell. Two subshells. The s subshell has one spherical orbital in it. The p subshell has 3 dumbbell-shaped orbitals. Thus there are a total of 4 orbitals.
      • Third Shell: 3s2 + 3p6 + 3d10 = total of 18 electrons in this shell. Three subshells: s, p, d. As expected, the s subshell contains one s orbital. The p subshell contains 3 p orbitals. The d subshell has 5 d orbitals. 1+3+5= 9 orbitals which collectively hold a maximum of 18 electrons.
      • Fourth Shell: 4s2 + 4p6 + 4d10 + 4f14 = total of 32 electrons in this shell. 1+3+5+7 = 16 orbitals which can hold a maximum of 32 electrons
      • Fifth Shell: 5s2 + 5p6 + 5d10 + 5f14 + 5g18 = total of 50 electrons in this shell.
    • Filling order is based on the Aufbau principle aka the Madelung energy ordering rule:
      • 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, 8s, etc.

Table 2.1 Chemical Bonds - p. 26

Type of Bond Description Bond Energy (kcal/mol)
Covalent bond Sharing of electron pairs 50-110
Ionic bond Example NaCl — electrons stolen to complete octets in valence shell to produce +cations and -anions 3-7
Hydrogen bond Attraction between bonded H+ atom and electronegative atom. Important for secondary structure within a protein molecule but also occurs between molecules 3-7
Hydrophobic interaction Nonpolar substances in presence of water (or other polar substrate) 1-2
van der Waals interaction Interaction of electrons of nonpolar substances 1

Concept of Electronegativity - p. 28

  • When two atoms of different elements are covalently bonded in a molecule, their electrons are not equally shared. So one of the positively charged nuclei may exert a greater attractive force on the shared electron pair than the other nucleus. This attractive force is called electronegativity.
  • p. 28: atoms with similar electronegativity = nonpolar covalent bond. Atoms with different electronegativity produce a polar covalent bond.
  • Extreme electronegativity = powerful positive proton nuclues which leads to ability to attract greater than “fair share” of shared electrons in a molecule
  • Maps to having almost having enough electrons in valencce shell to complete s2p6 octet
  • In general, most electronegative columns in the periodic table are Groups 16 & 17 because they have 6 & 7 electrons, respectively in their outer valence shell.
  • Conversely, the least electronegative groups in the periodic table are Groups 1 & 2 which respectively have 1 & 2 electrons in their outermost shell and these are the easiest electrons to lose.
  • At the limit, when superstrong-electronegative and superweakly-electronegative elements interact, they don’t just form a hyper-polar covalent bond; instead actually turn into an Ionic Bond.
  • good illustration of this in MSJ p. 288, Figure 8.5.

    Ionic Bonds - p. 28-29

  • When 2 interacting atoms have a large enough difference in electronegativity, ions form as the more powerfully electronegative nucleus strips electrons away from the weakly electronegative (aka positive atom).
  • Example of sodium Na and chlorine Cl
    • Na has its 1st two shells filled (aka 1s2 + 2s2 + 2p6). Also, as a Group I element, it also has an extra electron in the next subshell: 3s2 aka the first electron to appear in Shell M. It is easy to lose this electron.
    • Meanwhile, Cl only has 7 electrons in its 3s2 + 3p6 subshells. But ideally it would “like” 8 electrons to complete subshell 3p6.
    • So in solution, Cl’s more electronegative (aka more positive) nucleus “steals” the sole electron in Na’s outermost subshell 3s2.
    • This produces a positive cation Na+ and a negative anion Cl- .
  • This also explains the octet rule for molecules formed by atoms where s2 and p6 subshells are concerned.

Hydrogen bonds

Section 2.4: H20 has special properties - p. 32-34

  • Electronegativity causes strongly polar covalent bonds. Water has a tetrahedral shape with oxygen in the center, two of the tetrahedral vertices occupied by hydrogen atoms, and the two remaining vertices occupied by an unbonded electron pair. Because oxygen has a high electronegativity relative to hydrogen, there is a partial delta-positive charge on each hydrogen. Conversely, there is a partial delta-negative charge on each unbonded electron pair on the other two vertices.
  • Highly polar bonds then cause multiple and powerful hydrogen bonds between water molecules. Can also refer to these hydrogen bonds as intermolecular forces.
    • In a liquid state, water molecules are connected to each other by chains of powerful but intermittent hydrogen bonds. Special properties of liquid water thanks to multiple strings of hydrogen bonds:
      • High surface tension
      • High cohesion (water to water) and high adhesion (water to other substances) which are critical to upward water flow in plant xylem
      • High surface tension (water striders) and allows beads of water droplets (ants carrying a ball of water)
      • Lower vapor pressure (aka evaporation)
      • High specific heat
      • High boiling point aka heat of vaporization
    • When water freezes and turns into a solid, the daisy chain of randomly appearing/disappearing hydrogen bonds form into more permanent hydrogen bonds. These new hydrogen bonds follow from the partial delta-charges; the result is a “hollow” crystal structure for ice which is why is less dense than liquid water. Each water molecule is connected to 4 other water molecules in a 3-d tetrahedral structure.