SPRING FINAL REVIEW

Periodic table

• Alkali metals, alkaline earth metals, transition metals, halogens, noble gases

• Mass number = P + N

• Isotopes - atoms of an element with different numbers of neutrons

• Average atomic mass from weighted average of isotope mass and relative abundance (frequency)

Moles

• PV = nRT

• Avogadro’s number 6.022*10^23

• AT STP (1 atm, 273K), 22.4 L/mol

• Molarity M = moles/L

• Percent composition - divide the mass of each element/compound by the total molar mass of the substance

• Empirical formula is simplest ratio, molecular formula is actual formula for substance

Energy

• Electron potential energy increases with distance from nucleus

• Electron energy is quantized - can only exist at specific energy levels at specific intervals, not in between

• Coulomb’s law: F = kq1q2/(r^2) where F is electrostatic force

• Atoms absorb energy in the form of electromagnetic radiation as electrons jump to higher energy levels; when electrons drop levels (closer), atoms give off energy

Photoelectron spectroscopy

• energy measured in electronvolts (eV)

• Incoming radiation energy = binding energy + kinetic energy of the ejected electron

• Electrons that are further away from nucleus require less energy to eject, thus will move faster

• Photoelectron spectrum

  • Each section of peaks represents a different energy level (1, 2, 3, etc.)

  • Subshells within each energy level (shape of space electron can be found in orbiting nucleus) are represented by the peaks (1s, 2s, 2p, etc.)

  • s(2) - first subshell, p(6) - second subshell

  • Height of peaks determines number of electrons in subshell (ex. Peak of p subshell in energy level 2 will be 3x that of s subshell)

Electron configuration

• Electron configuration - spdf - shorthand with noble gas first

• Configuration rules

  • Aufbau principle - electrons fill lowest energy subshells available first

  • Pauli exclusion principle - 2 electrons in same orbital cannot have same spin

  • Hund’s rule - Electrons occupy empty subshells first

• Zn +2, Ag +1, Al +3, Cd +2, most other transition metal charges vary

Periodic trends

• Electrons are more attracted if they are closer to the nucleus, or if there are more protons

• Electrons are repelled by other electrons - if there are electrons b/w the valence electrons and nucleus, the e- will be less attracted (shielding)

• Completed shells are very stable, completed subshells are also stable; atoms will add/subtract valence electrons to complete their shell

• INCREASING: atomic radius down left; ionization energy up right; electronegativity up right

  • Ionization energy - energy required to remove an electron from an atom

  • Electronegativity - how strongly the nucleus of an atom attracts electrons of other atoms in a bond

  • Electron affinity - energy change that occurs when an electron is added to an atom in the gas state (usually exothermic - energy is released)

Bonds

• Atoms are more stable with full valence shells

• Ionic bonds

  • Cation gives up electrons completely

  • Electrostatic attractions in a lattice structure

  • Metals and nonmetals (salts)

  • Coulomb’s law - greater charge leads to a greater bond/lattice energy (higher melting point)

  • If both have equal charges, smaller radius will have greater coulombic attraction

  • Ionic solid - electrons do not move around lattice; ionic solids are poor conductors of electricity; ionic liquids conduct electricity because ions are free to move around, though e- are still localized around particular atoms

• Metallic bonds

  • Sea of electrons model - positively charged core is stationary while valence electrons are very mobile

  • Metals bond to form alloys - interstitial alloy w/ metals of different radii; substitutional alloy w/ metals of similar radii

• Molecular covalent bonds

  • 2 atoms share electrons - both atoms achieve complete outer shells

  • 2 nonmetals

  • Creates molecules - 2+ atoms covalently bonded together

  • Single has 1 sigma bond - order 1, longest length, least energy; double has 1 sigma and 1 pi bond - order 2, int. length, int. energy; triple has 1 sigma and 2 pi bonds - order 3, shortest length, greatest bond energy

  • Bond forms when potential energy is at minimal level

    • Too close - potential energy is too high due to repulsive forces

    • Too far - potential energy is near 0 because attractive forces are very weak

    • Minimul PE occurs when repulsive and attractive forces are balanced

  • Network covalent bonds - lattice of covalent bonds - poor conductors, high melting and boiling points

Conductivity

• Conductivity of different substances in different phases

Lewis dot structures

• Resonance - for bond order calculations, average together all possible orders of a specific bond

• BORON (B) is stable with 6 electrons - only one that does not need a full octet

• Expanded octets - any atom of an element from n=3 or greater (those with a d subshell) can have [8,12] valence electrons on center atom

  • Noble gases form bonds by filling empty d orbital with electrons

• Formal charge - number of valence electrons minus assigned electrons (1 e- for each line “shared” bond) - 0 for neutral molecules

Molecular geometry (VSEPR)

Double and triple bonds have more repulsive strength than single bonds - occupy more space

Lone electron pairs have more repulsive strength than bonding pairs, so molecules with lone pairs will have slightly reduced angles between terminal atoms 

Hybridization - how many atoms are attached (sp, sp2, sp3, sp3d, etc.)

Polarity

• Covalent bond where electrons are unequally shared - polar covalent

• Dipoles are caused by polar covalent bonds - pair of opposite electric charges separated by some distance, like partial charges on atoms in a polar covalent bond

• If 2 identical atoms bond (ex. Cl-Cl) the electrons are equally shared, creating a nonpolar covalent bond with no dipole

• Bonds can be polar; so can molecules depending on the molecular geometry (and polarity of bonds - secondary)

• In polar molecules, more electronegative atoms will gain negative partial charge

  • Usually central atom will be positive - exception is hydrogen (terminal), which is usually positive since it has less electronegativity

Intermolecular forces

• Forces b/w molecules in a covalently bonded substance - need to be broken apart for covalent substances to change phases

• Changing phase: ionic substances break bonds b/w individual ions; covalent substances keep bonds inside a molecule in place but break bonds b/w molecules

• Dipole-dipole forces

  • Polar molecules - positive end of one molecule is attracted to negative end of another molecule

  • Greater polarity -> greater dipole dipole attraction -> larger dipole moment -> higher melting/boiling points

  • Relatively weak overall - melt and boil at low temps

• Hydrogen bonds

  • Special type of dipole-dipole attraction where positively charged hydrogen end of a molecule is attracted to negatively charged end of another molecule containing an extremely electronegative element (F, O, N)

  • Much stronger than normal dipole-dipole forces since a hydrogen atom “sharing”/giving up its lone e- to a bond is left w/ no shielding

  • Higher melting/boiling points than substances held together only by other types of IMF

• London dispersion forces

  • All molecules - very weak attractions due to random motion of electrons on atoms within molecules (instantaneous polarity)

  • Molecules w/ more e- experience greater LDF (more random motion)

  • Higher molar mass usually means greater LDF (as mass increases, e- increases for the molecule to remain electrically neutral)

• IMF strength

  • Ionic substances are generally solids at room temp - melting them requires lattice bonds to be broken - necessary energy determined by Coulombic attraction

  • Covalent substances (liquids at room temp) boil when IMF are broken; for molecules of similar size, from strongest to weakest: hydrogen bonds, permanent dipoles, LDF (temporary dipoles - greater for larger molecules)

  • Melting/boiling points of covalent substances are LOWER than for ionic substances

• Bonding/Phases

  • Substances w/ weak IMF (LDF) tend to be gases at room temp (N2); substances w/ strong IMF (hydrogen bonds) tend to be liquids at room temp (H2O)

  • Ionic substances do not experience IMF - since ionic bonds are stronger than IMF, ionic substances are usually solids at room temp

Vapor pressure

• Molecules in a liquid are in constant motion - if they hit the surface of the liquid with enough kinetic energy, they can escape the IMF holding them to other molecules and transition into the gas phase

• Vaporization (NOT boiling) - no outside energy is added

• Temperature and vapor pressure are directly proportional

• At the same temp, vapor pressure is dependent on strength of IMF (stronger IMF, lower vapor pressure)

Solution separation

• Solutes and solvents - like dissolves like

• Paper chromatography

  • Piece of filter paper with substance on the bottom is dipped in water

  • More polar components of substance travel further up the filter paper with the polar water

  • Distance substance travels up the paper measured by retention/retardation factor Rf = (distance traveled by solute - substance being separated)/(distance traveled by solvent front - water)

  • Stronger attraction - larger Rf

• Column chromatography

  • Column is packed with a stationary substance

  • separable solution (analyte) is injected, adhering to stationary phase 

  • another solution (eluent - liquid/gas) is injected into column

  • more attracted analyte molecules will move through faster and leave column first

• Distillation

  • Takes advantage of different boiling points of substances by boiling a mixture at an intermediate point

  • Vapor is collected, cooled, and condensed back to a liquid separate of leftover liquid

Kinetic molecular theory

• Kinetic energy of a single gas molecule: KE = ½ mv^2

• Average kinetic energy of a gas depends on the temperature (directly proportional), not the identity of the gas (different gases will have same KE at same temp)

• Ideal gases have insignificant volume of molecules, no forces of attraction b/w molecules, and are in constant motion without losing KE

  • Deviations occur at low temperatures or high pressures (gas molecules are packed too tightly together)

    • Volume of gas molecules becomes significant (less free space for molecules to move around than predicted)

    • Gas molecules attract one another and stick together (real pressure is smaller than predicted pressure)

• Maxwell-boltzman diagrams

  • Higher temp -> greater KE -> greater range of velocity

  • Smaller masses, greater velocities to have same KE

• Effusion

  • Rate at which a gas escapes from a container through microscopic holes

  • High to low pressure

  • Greater speed, greater temp, greater rate of effusion

  • If at same temp, gas w/ lower molar mass will effuse first

Equations

• Ideal gas equation: PV = nRT

  • R=0.0821

• Combined gas law: P1V1/T1 = P2V2/T2

• Dalton’s law: P(total) = Pa + Pb + Pc + …

• Partial pressure: Pa = P(total)*(moles of gas A)/(total moles of gas)

• Density: D = m/V

  • From ideal gas law: Molar mass = DRT/P

• Electromagnetic spectrum

  • E=hv 

    • E = energy change; h = Planck’s constant 6.626*10^-34; v = frequency

  • C = lambda * v

    • C = speed of light 2.998*10^8; v = frequency; lambda = wavelength

• Beer’s law: A = abc

  • A = absorbance; a = molar absorptivity (constant depending on solution); b = path length of light through solution (constant); c = concentration of solution

  • Colorimetry - direct relationship b/w concentration and absorbance

Types of reactions

• Synthesis: everything combines to form one compound

• Decomposition: one compound + heat is split into multiple elements/compounds

• Acid-base rxn: Acid + base -> water + salt

• Oxidation-reduction (redox) rxn: changes the oxidation state of some species

• Combustion: substance w/ H and C + O2 -> CO2 + H2O

• Precipitation: aqueous solutions -> insoluble salt (+ more aq sometimes)

  • Can be written as net ionic - Those free ions not in net ionic are spectator ions

Solubility rules

• Alkali metal cations or ammonium (NH4+) cations are ALWAYS soluble

• Compounds with a nitrate (NO3-) anion are ALWAYS soluble

Common polyatomic ions

Calculations

• Percent error: 100 * abs(experimental - expected)/(expected)

• Combustion analysis - use law of conservation of mass (if x g of CO2 is produced, find g of C which will be starting amt)

• Gravimetric analysis - when asked to determine the identity of a certain compound, find g of component produced, then use mass percent (g found / total sample mass) and compare to mass percent of options (molar mass of component / molar mass of entire compound)

Oxidation states

• Neutral atoms not bonded to other atoms have an oxidation state of 0

• Monoatomic ions have an oxidation state equal to the charge on that ion (ex. Zn2+ will be +2)

• Oxygen is -2 (EXCEPTION: in hydrogen peroxide, H2O2, O is -1)

• Hydrogen is +1 w/ nonmetals, -1 with metals

• In absence of oxygen, most electronegative element in a compound will take an oxidation state equal to its usual charge (ex. F is -1 in CF4)

• IF none of the above rules apply, determine the oxidation state by adding up all elements’ oxidation states to 0/charge on ion

• C, N, S, P frequently vary oxidation states (low electronegativity)

Redox reactions

• Write full rxn as 2 half reactions (oxidation and reduction; OIL RIG)

• Add H2O to compensate for oxygen on one side

• Add H+ to compensate for H from H2O on other side

• Balance 2 half rxns to have the same number of electrons and add them together to produce one complete reaction

• ACIDIC: stop here

• BASIC: Add OH- to both sides - enough for all H+ on one side to be converted to H2O; then cancel out H2O so it only remains on one side

Acids and bases (briefly)

• Color change signals the end of a titration (can be redox or acid/base)

• Acids are capable of donating protons (H+); bases are capable of donating electrons

  • Species with the H+ ion are acids, same species but without H+ is a base - conjugate acid/base pairs

• Water can act as an acid or base - amphoteric