Elements: The most simple, cannot be broken down, the basis of matter.

Mixtures: 2+ substances physically bonded NOT in fixed ratio. They retain individual properties and can be physically separated.

Compounds: 2+ elements chemically bonded, fixed ratio. They have new properties, and can't be physically broken.

Kinetic Molecular Theory: Explains physical properties of matter & change of state. Uses symbols & good pics.

Temperature (T): Average kinetic energy of particles.

Kelvin (K): SI unit of temperature with same increments as Celsius (C).

• K = C + 273

Atom: Negatively charged, dense nucleus; occupy space outside nucleus.

Subatomic particles: Protons (+), neutrons (neutral), electrons (-).

• Relative masses/charges of subatomic particles should be known

Atomic Mass (A): Mass of electron negligible.

Atomic number (Z): Number of protons.

Isotopes: Atoms of same element with different number of neutrons -> same chemical properties, different physical properties.

Emission spectra: Produced when photons are emitted from atoms as electrons in excited states return to lower energy levels.

Line emission spectrum of hydrogen: Provides evidence for the existence of electrons in discrete energy levels.

Relationship between energy, wavelength, frequency: Which converge at higher energies.

Continuous spectrum vs. line spectrum: Energy across electromagnetic spectrum (EMS).

Emission spectrum of hydrogen atom: Lines & energy transitions to 1st, 2nd, 3rd energy levels.

Main energy level: Given integer number, n. Can hold maximum of 2n2 electrons.

Division of main energy level into s, p, d, f sublevels: Of successively higher energy.

Orbitals: Have defined energy, shape for e-'s, and orientation and chemical environment; can hold 2 e-'s of opposite spin.

Aufbau principle: e-'s occupy each energy level to reduce max e-, Hund's rule, Pauli exclusion principle to reduce e- configuration.

e- configuration: (long) for atoms and ions up to Z=36; (full e- configuration), condensed ([Ne]3s2).

Ionization Energy: Energy required to remove one mol e-'s from one mol of gaseous atoms (kJ mol-1).

• Can be calculated from emission spectra.

• X(g) -> X+(g) + e-

Exceptions: To the Aufbau principle, there are two exceptions.

• Exception #1

  • Be -> B: new e- is easier to remove

  • Mg -> Al: must be at a higher energy

  • Evidence for sublevels -> p is higher energy than s

• Exception #2: Hund's rule

2nd IE: X+(g) -> e- + X2+(g)

The mole (mol): An SI unit for an amount of something.

• 6.02 * 1023 "entities" (cations/molecules/ions/e's...).

Ar: Relative atomic mass = weighted average of isotope abundances.

Mr: Relative molecular mass = add Ar for each atom in compound.

• Mr = relative molecular mass = the mass of "C"

Mr: Molar mass = # HRS units of 1 mol.

Empirical formula: Simplest whole # ratio of elements in a compound.

• Steps for converting to empirical formula:

  1. Treat percentages like grams

  2. Convert g -> mol of atoms for each element

  3. Divide all by smallest mol amount

  4. Multiply all to get whole numbers

Molecular formula: Actual formula.

Molar concentration of solutions:

• Concentration mol dm-3 = mol/ dm3 solvent

Ideal gases: Particles moving with elastic collisions

• Breaks down @ low temps: intermolecular forces.

• Breaks down @ low volume: negligible particle volume

For the same gas sample (constant n):

• P1V1/T1 = P2V2/T2

Ionic Bonding: Typically a metal & nonmetal.

• Charged ions arranged in a lattice

• Lattice enthalpy: strength of ionic bond, and becomes stronger with 1) higher charges and 2) smaller ions.

Physical properties of ionic compounds:

• Volatility: Low b/c strong bonding (can't move)

• Electrical conductivity: NO as solid (ions can't move) yes molten/liquid (ions can move)

• Solubility: Based on lattice enthalpy; more soluble in polar solvent

• Brittle = solids fracture when force applied

Metallic Bonding: An electrostatic attraction between a metal lattice & a "sea" of delocalized e's.

Physical Properties of metallic bonding:

• High melting & boiling points (smaller ions & higher charges = stronger)

• Electrical conductivity: yes solid (delocalized e's can move) yes liquid

• Thermal conductivity (good at it) -> delocalized e's can carry/transfer heat quickly

• Malleability & ductility

  • Can be beaten into a wire -> b/c cations can slide past each other without breaking a bond

  • Can be hammered -> into a drawn

Covalent Bonding: (includes metalloids)

• Typically 2 nonmetals

• An electrostatic attraction between shared pair(s) of e-'s & positive nuclei on either side of the e-'s.

Coordinate bond: When one atom donates both shared e-'s.

Electronegativity: Relative attraction an atom has for a shared pair of e-'s; polarity of a bond is based on EN difference

VSEPR Shapes: Linear, trigonal planar, bent, tetrahedral, trigonal pyramidal, bent.

Molecular Polarity: Based on bond polarities & shape; larger molecules = stronger dispersion forces

Size & polarity determine volatility: -> polar = dipole-dipole = hydrogen bonding -> high melting/boiling points (solids); non-polar = dispersion forces -> generally low melting/boiling -> volatile.

Giant covalent structures:

• No conductivity (except graphite)

• Si, SiO2 -> sand

Allotropes of carbon:

• Diamond, graphite, fullerenes

Sigma bonds (σ): Head-on, linear combination of orbitals.

• e- density is on the bond axis (between nuclei)

• Single bond is always sigma bond.

Pi bonds (π): Lateral combination of p-orbitals.

• e- density on opposite sides of bond axis.

• Always form as 2nd or 3rd bond.

Hybridization: Mixing different orbital types.

Resonance structures: When there is more than one position for a double bond.

Formal Charge: A man-made concept to "determine" which Lewis structure is "preferred".

• FC = V - N - 1/2B

• Calculate for each atom

• Preferred structure is closer to all zeroes

Polymers: Made of many repeating units (monomers).

• Join together chemically

• Form macromolecules

Condensation Polymers: Form by removing H2O.

Addition Polymers: Form by breaking a double bond.

Transition Elements: Shows variable oxidation states with at least one having an incomplete d sublevel.

All can be +2, b/c they lose their 4s e's first (easiest to take); there is not a big difference in energy between 4s & 3d.

Have magnetic properties (ferro/para/diamagnetic).

Form colored compounds

• Form complex ions with ligands

Colored Complexes: In a free ion, the five d-orbitals are degenerate (equal in energy).

• A ligand approaches to form a coordination bond with the ion

• As the ligand approaches, the d-orbitals split

• Light passing through the compound is absorbed & promotes an electron to the higher d-orbital

• The non-absorbed colors are transmitted; we see the color complementary to the light/was absorbed (color wheel).

  • Factors affecting color seen:

    1. Geometry of the complex ion formed (we know octahedral)

    2. Which transition metal

    3. Oxidation state of the metal

    4. The ligand (strengths given below color wheel)

Stereoisomers: Same connectivities & bond multiplicities -> different 3-D arrangement.

Cis/trans isomers: Alkenes & cycloalkanes

• Have free rotation

Optical isomers: Can bend polarized light

• Contain a chiral carbon (attached to 4 different groups)

• Enantiomers = non-superimposable "mirrored" images

Mass spectroscopy: Used to find molecular mass and fragmentation patterns.

• M+ = molecular ion

IR Spectroscopy: Makes their bonds vibrate

• Greenhouse gases absorb IR wavelengths of light which stops the light/ heat from escaping back out into space

Proton nuclear magnetic resonance (1H NMR):

• What H environments exist

• How many H environments (number of peaks)

• How many H's in each environment (integral trace)

• How many H's are in the adjacent environment (splitting pattern)

Atomic Radius: 1/2 of distance between 2 nuclei.

First Ionization Energy: The energy required to remove 1 mol of e-'s from 1 mol gaseous atoms (kJ mol-1).

• X -> X+ + e-

Electron affinity: Energy change to add 1 mole e- to 1 mol gaseous atoms.

• X(g) + e- -> X-(g)

Oxidation States: Number to show e-'s "transferred" when forming a bond

• *Elements always +1 (except in hydrides, H-)

• *Oxygen always -2 (except -1 in peroxides, H2O2)

• *Hydrogen always +1 (except -1 in hydrides e.g., LiH)

Chemical reactions: Involve the transfer of energy between the system and surroundings; no energy is destroyed.

Temperature: Average kinetic energy.

Heat: Total energy (kJ).

Exothermic: If energy is transferred to the surroundings, the surroundings will get warmer, the surroundings get colder.

Endothermic: If the surroundings transfer energy to the reaction, the surroundings get colder.

Activation energy: Energy needed to break ionic lattice.

Reaction coordinate:

Bond Enthalpies: ΔH = energy within a bond in a gaseous molecule.

• It's an average value

Covalent bond enthalpy: Energy within a bond of a gaseous molecules.

• Form/from gaseous atoms.

Atomization energy: Energy needed to form 1 mol gaseous atoms from a standard state element; always endothermic.

Ionization energy: Energy needed to remove 1 mole e- from 1 mol gaseous atoms; always endothermic.

Lattice enthalpy: Forming/breaking an ionic lattice from/into gaseous ions.

• Forming bonds = very exothermic

• Breaking bonds = endothermic

• Magnitude is based on ionic bond strengths: stronger when smaller & has higher charge

Combustion: Reactive metals, nonmetals, hydrocarbons in excess O2 to form oxides.

• ΔHcΘ is the enthalpy change when 1 mol of substance completely combusts at standard conditions.

  • -kJ mol-1

Incomplete combustion: In limited O2, form some amounts of CO/C.

• And/or soot & less energy is released

Fossil fuels: Variety compounds formed through decomposition of biological materials.

• Reduction of biological complexes.

  • Contain C, H, O, N, S & metals

Petroleum: Mostly hydrocarbons

Natural gas: Mostly CH4, small sulfur compounds.

Coal: Mostly C, small hydrocarbons, sulfur compounds & metals.

Biofuels: Made from plant/animal matter; derived directly/ indirectly from sunlight.

Fuel Cells: Electrochemical energy comes from chemical reaction.

• Combustion where oxidation is occurring - the anode

Greenhouse gases: The greenhouse effect occurs when CO2, CH4 or nitrogen oxides absorb IR radiation emitted from Earth's surface after they've been heated by the sun's rays.

• Greenhouse gases emit IR radiation in all directions so net effect is higher solar energy trapped in atmosphere; global temp becomes higher

• Why greenhouse gases absorb infrared radiation: Molecules can vibrate as the bonds in them stretch & bend. The energy used here is infrared energy. If the stretching & bending involves a change in dipole moment, the vibrations are IR active.