Chem 1 Exam 2

Formal charge

valence electrons - # electrons around atom after bonds are split

Criteria for choosing more important resonant form

smaller formal charges preferable, same nonzero FC on adjacent atoms are not preferable, more negative FC should reside on a more electronegative atom

Exceptions to octet rule - less than 8 electrons

hydrogen, lithium, helium (2), beryllium and boron (4 or 6)

Exceptions to octet rule - more than 8 electrons

third row and lower because they have d orbitals - ex. Phosphorous, sulfur, iodine

Exceptions to octet rule - odd number of electrons

NO

2 bonded atoms 0 lone pairs

linear molecular and electron geometry, 180°, sp

3 bonded atoms 0 lone pairs

trigonal planar molecular and electron geometry, 120°, sp2

2 bonded atoms 1 lone pair

bent molecular geometry, trigonal planar electron geometry, <120°, sp2

4 bonded atoms 0 lone pairs

tetrahedral molecular and electron geometry, 109.5°, sp3

3 bonded atoms 1 lone pair

trigonal pyramidal molecular geometry, tetrahedral electron geometry, ~107°, sp3

2 bonded atoms 2 lone pairs

bent molecular geometry, tetrahedral electron geometry, ~105°, sp3

5 bonded atoms 0 lone pairs

trigonal bipyramidal molecular and electron geometry, 90°, 120°, 180°, sp3d

4 bonded atoms 1 lone pair

seesaw molecular geometry, trigonal bipyramidal electron geometry, <120°, <90°, sp3d

3 bonded atoms 2 lone pairs

t-shaped molecular geometry, trigonal bipyramidal electron geometry, ~90°, 180°, sp3d

2 bonded atoms 3 lone pairs

linear molecular geometry, trigonal bipyramidal electron geometry, 180°, sp3d

6 bonded atoms 0 lone pairs

octahedral molecular geometry, octahedral electron geometry, 90°, 180°, sp3d2

5 bonded atoms 1 lone pair

square pyramidal molecular geometry, octahedral electron geometry, <90°, <180°, sp3d2

4 bonded atoms 2 lone pairs

square planar molecular geometry, octahedral electron geometry, 90°, 180°, sp3d2

Sigma bond

end-to-end overlap, all single bonds

Pi bonds

orbitals parallel to each other (sideways overlap)

Double bond

1 sigma 1 pi

Triple bond

1 sigma 2 pi

Dipole-dipole forces

positive pole of one polar molecule attracts negative pole of another (only in polar molecules)

Hydrogen bonding

H atom covalently bonded to N, O, or F, increase in H bonds = increase in boiling point

London dispersion forces

instantaneous dipole in one particle induces a dipole in another → attraction between them (in all particles but stronger for larger, more polarizable particles)

Order of strength of forces (highest to lowest)

H bonds > dipole-dipole > london dispersion

Stronger intermolecular forces

higher BP, viscosity, surface tension, lower vapor pressure

Strong electrolytes

conduct electricity very well, dissociate (break apart) completely in aqueous solution, strong acids and strong bases

Weak electrolytes

do not conduct electricity very well because they produce relatively fewer ions when dissolved in water, weak acids and weak bases

Non-electrolytes

do not conduct electricity because they do not produce ions when dissolved in water

Atom in its elemental form or pure form

ON = 0

Monatomic ion

ON = charge on ion

Group 1A

ON = +1 in all compounds

Group 2A

ON = +2 in all compounds

Hydrogen

ON = +1 combined with nonmetals, -1 combined with metals

Fluorine

ON = -1 in all compounds

Oxygen

ON = -2 in all compounds except with F, -1 in peroxides (ex. H2O2, Na2O2, BaO2)

Group 7A

ON = -1 in combination with metals, nonmetals (except O), and other halogens lower in the group

Sum of all ON values for atoms in a compound

= 0, = charge on ion in polyatomic ion

Arrhenius acid

produces H+ ions in water

Arrhenius base

produces OH- ions in water

Bronsted lowry acid

proton (H+) donor

Bronsted lowry base

proton (H+) acceptor

Strong acids

HCl, H2SO4, HI, HNO3, HBr, HClO3, HClO4

Strong bases

LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

Neutralization reaction

acid reacts with base to form salt and water

Nonmetal oxides form __ in water

acids

Metal oxides form __ in water

bases

Oxidation

increase in ON, loss of electrons

Reduction

decrease in ON, gain of electrons

Oxidizing agent

undergoes reduction (usually non-metal)

Reducing agent

undergoes oxidation (usually metal)

Quick way to tell if reaction is redox

something is a pure element and then becomes a compound or vice versa

Precipitation reaction

2 soluble ionic compounds react to form an insoluble product (precipitate)

Bond order

½(# bonding electrons - # antibonding electrons)

Positive bond order

molecule should be stable

0 or negative bond order

molecule should not be stable

High bond order

stronger and shorter bond

Dimagnetic

all paired electrons

Paramagnetic

at least one unpaired electron

Ideal gas

particles have no intermolecular forces and no volume (high temp, low pressure)

NO₃⁻

Nitrate

NO₂⁻

Nitrite

SO₃²⁻

Sulfite

SO₄²⁻

Sulfate

HSO₄⁻

Hydrogen sulfate (or bisulfate)

OH⁻

Hydroxide

CN⁻

Cyanide

PO₄³⁻

Phosphate

HPO₄²⁻

Hydrogen phosphate

H₂PO₄⁻

Dihydrogen phosphate

NH₄⁺

Ammonium

CO₃²⁻

Carbonate

HCO₃⁻

Hydrogen carbonate (or bicarbonate)

CH₃COO⁻ (or C₂H₃O₂⁻)

Acetate

MnO₄⁻

Permanganate

Cr₂O₇²⁻

Dichromate

CrO₄²⁻

Chromate

O₂²⁻

Peroxide

ClO⁻

Hypochlorite

ClO₂⁻

Chlorite

ClO₃⁻

Chlorate

ClO₄⁻

Perchlorate