Chem 1A Final

gas characteristics

• assumes volume/shape of container

• most compressible state of manner

• mixes evenly when two are confined in the same container

pressure

F/A; results of collisions between gas molecules and the surfaces around them

manometer

• measures pressure in a stationary or flowing fluid/gas

• when calculating h in a book problem, convert from cm to mmHg to atm to Pascals

• h+ P_gas= P_atm

• P_flask > P_atm then add

• P_flask < P_atm then subtract

Boyle’s Law

• relates pressure and volume

• P₁V₁=P₂V₂

• volume and pressure are inversely proportional

Charles’ Law

• relates volume and temperature

• V₁/T₁ = V₂/T₂

• directly proportional

• independent of pressure

Avagadro’s Law

as number of moles increase, volume also increase

ideal gas law

PV=nRT; temperature must be in kelvin, can be manipulated in order to fit the problem

gas density

d=PM/RT; highest molar mass = highest density

partial pressures

• P_total = P₁ + P₂

mole fraction

X_a = moles of a/ sum of moles of all components

• can be used for book problems and times by the total pressure in order to find the partial pressure of a gas

ideal gas/ kinetic molecular theory

does not actually exist but in theory we assume that gas particles have a negligible volume and no IMFS

• as well as KE_avg and temp being directly proportional

• these particles are in continuous movement 

• speed increases with temp

maxwell-boltzmann distribution

by making changes to temp, you change how slow/fast the particles are moving

graham’s law

rate₁/rate₂ = square root of M₁/M₂

collision rate

• number of collisions per second

• can find the relative frequency when multiplying relative n and relative u_avg ( these are found by dividing pressure by first flask pressure)

endothermic process

where system absorbs heat and heat flows from surrounding into system

exothermic process

system evolves heat, heat flows out of system into surroundings

wavelength

measured in meters, symbol is lambda

frequency

the number of waves that pass a fixed place in a given amount of time

• symbol = v

• measured in s^-1 or 1/s

photoelectric effect

the emission of electrons from surface, typically a metal, when light or other electromagnetic radiation if a specific frequency strikes it

duality of light

dual nature of light that is both a wave/particle

electron diffraction

pattern that shows electron is behaving like a wave

• particles have wavelike properties

absorption 

when an atom or molecule takes in a photon of light, causing an electron to jump from a lower energy level to a higher one.

excitation

the result of absorption: an electron has gained energy and is now in a higher (excited) energy state instead of its ground state.

observation effect

the act of measuring/observing subatomic particles changes its state or behavior

uncertainty principle

looks at uncertainty in the position/momentum of an electron 

Schrodinger’s equation

incorporates wavelike and particle behavior: used to calculate probabilities if where the electron is located

quantum number

numbers that make Schrodinger’s equation true

principle quantum number

first quantum number, abbreviated by n

• restriction - can have integer value of whole positive number

• n tells you energy of electron (aka what shell it is in)

• n = number in front of orbital (ex. in 1s n = 1)

angular momentum quantum number

symbol is l

• restriction - dependent on principle quantum number

• 0 = s orbital, 1 = p orbital, 2 = d orbital, 3 = f orbital

magnetic quantum number

symbol is m_l

• restriction - can have values of -l, -l + 1, -l + 2 …0…+1,+2,+l

• describes the orientation of the orbital in space

• s orbital is 0

• p orbital is -1, 0 ,1

• d orbtial is -2, -1, 0, 1, 2

• f orbtial is -3, -2, -1, 0, 1, 2, 3

magnetic spin quantum number

symbol is m_s

• describes the angular momentum of an electron’s spin around its own axis

• electron can either be +1/2 (spin up) or -1/2 (spin down)

pauli exclusion principle

no 2 electrons can have the same set of 4 quantum numbers

• orbital can only hold 2 electrons and they must have opposite signs

effective nuclear charge

Z_eff = Z - o (shielding constant)

atomic radius

measurement of an atom’s size

• decreases across a period from left to right, increases going down a group

ionization energy

the energy to remove an electron during a specific reaction

• decreases down a group and increases from left to right

• increases in a chemical reaction with more positive charges to an atom

electron affinity

ease in which an electron can be added during a specific reaction

• E_a = E(a)-E(a^-)

• more negative E_a, the easier it is to put an electron on to the atom

• increases from left to right across a period and decrease from top to bottom down a group

ionic radius

more protons means that the electrons are pulled in tighter and therefore results in a smaller ionic radius

• increases as you move down a group b/c new shells are added

• decreases across a period b/c positive nuclear charge pulls electrons more tightly

covalent bonding

electrons are shared equally (non metal bound to non metal)

• sharing of electrons to lower the energy of the system

• polar covalent bonding is unequal sharing of electrons due to electronegativity difference

dipole moment

separation of charges within a molecule creating positive end and negative end

electronegativity

the power of an atom to attract electrons when it is part of a compound

• F is the most electronegative

• down a group and left from fluorine electronegativity decreases

lewis dot structure

• determine number of valence

• arrange chemical symbols to show which are bonded (least electronegative in the middle and hydrogens at end of chain)

• distribute electron pairs so each bond has one

• distribute lone pairs/multiple bonds to satisfy octet

heisenberg uncertainty principle

you can only know the position or the velocity of a particle not both

d-block configuration: Cr and Cu exceptions

Chromium

• [Ar] 4s¹3d⁵ a half filled d⁵ subshell\

Copper

• [Ar]4s¹3d¹⁰ a fully-filled d¹⁰ subshell is more stable

formal charge

f= V - LPE - ½ BE

• formal charges on the atoms must add up to the charge on the ion

resonance structure

used to describe equivalent structures, average bond characteristics, energy of resonance hybrid is lower, minimize formal charges

electron deficient

have less than 8 electrons around one single atom and won’t complete octet

enthalpy

the amount of heat evolved or absorbed in a reaction

• measured in joules or kilojoules 

• equation is delta H = (bonds broken) - (bonds formed)

VESPR Theory

• bond angles and molecular shape

• electron pair repulsions: bonds are made of electrons and you want them as far apart as possible

• steric number: count all lone pairs and atoms that are connected to central atom

steric number 2

• (2 regions of electron density) becomes linear

• linear - molecular/electronic geometry

• 180 degrees

steric number 3

• 3 regions of electron density

• trigonal planar (electronic/molecular geometry)

• molecular geometry can also be bent if there is one lone pair

• 120 degrees

steric number 4

• 4 regions of electron density

• becomes tetrahedral (electronic/molecular)

• molecular geo is bent if there is 2 lone pairs

• trigonal pyramidal with one lone pair

steric number 5

• 5 regions of electron density

• becomes trigonal bipyramidal (electronic geometry)

• molecular geometry is seesaw with one lone pair

• square planar with 2 lone pairs

exceptions to the octet

• incomplete/deficient octet: some atoms are stable with fewer than 8 electrons (H=2, He=2, Li=2, Be=4, B=6)

• hypervalent/expanded octet: atoms in period 3 or higher can have more than 8 electrons (P, S, Cl, Br, I, Xe)

dipole moment (2)

• dipole arrows towards the more electronegative atom

  • no net dipole = non polar

  • facing away from each other = polar

• just because there are lone pairs does not mean it is polar

bond energies v bond length v bond multiplicity

increased multi, increased bond energy, decreased bond length

paramagnetic

atoms or molecules with unpaired electron spins, attracted to magnetic field

diamagnetic

all electron spins are paired, repelled by magnetic field

strong electrolytes

• strong acids: HCl, HBr, HNO3, H2SO4, HClO4

• strong bases: LiOH, NaOH, KOH, Ba(OH)2, Sr(OH)2, Ca(OH)2

• soluble salts: NaCL, KBr, LiNO3

weak electrolytes

• HF, CH2COOH, HCN, H2CO3, H3PO4

• NH3, CH3NH2

• PbCl2, Ag2SO4

nonelectrolytes

• most covalent compounds like sugar (C6H12O6), ethanol (C2H5OH), urea (NH2CONH2)

steric number 2

• zero lone pairs

• molecular/electronic geometry = linear

• bond angle = 180

• sp hybridization

steric number 3 (zero lone pairs)

• molecular/electronic geometry = trigonal planar

• bond angles = 120

• sp² hybridization

steric number 3 (one lone pair)

• molecular geometry = trigonal planar

• electronic geometry = bent

• bond angles: 120

• sp² hybridization

steric number 4 (no lone pairs)

• molecular/electronic geometry = tetrahedral

• bond angles: 109.5

• sp³ hybridization

steric number 4 (one lone pair)

• molecular geometry = tetrahedral

• electronic geometry = trigonal pyramidal

• bond angles: <109.5

• sp³ hybridization

steric number 4 (2 lone pairs)

• molecular geometry = tetrahedral

• electronic geometry = bent

• bond angles: «109.5

• sp³ hybridization

steric number 5 (no lone pairs)

• molecular/electronic geometry = trigonal bipyramidal

• bond angles: 90, 120

• sp³d hybridization

steric number 5 (one lone pair)

• molecular geometry = trigonal bipyramidal

• electronic geometry = see-saw

• bond angles: <90, <120, >180

• sp³d hybridization

steric number 5 (2 lone pairs)

• molecular geometry = trigonal bipyramidal

• electronic geometry = t-shaped

• bond angles: <90

• sp³d hybridization

steric number 5 (3 lone pairs)

• molecular geometry = trigonal bipyramidal

• electronic geometry = linear

• bond angles: 180

• sp³d hybridization

steric number 6 (no lone pair)

• molecular/electronic geometry = octahedral

• bond angles: 90

• sp³d² hybridization

steric number 6 (one lone pair)

• molecular geometry = octahedral

• electronic geometry = square pyramidal

• bond angles: 90, <90

• sp³d² hybridization

steric number 6 (2 lone pairs)

• molecular geometry = octahedral

• electronic geometry = square planar

• bond angles: 90,

• sp³d² hybridization

steric number 6 (3 lone pairs)

• molecular geometry = octahedral

• electronic geometry = t-shaped

• bond angles: 90, <180

• sp³d² hybridization

steric number 6 (4 lone pairs)

• molecular geometry = octahedral

• electronic geometry = linear

• bond angles: 180

• sp³d² hybridization

ammonium

NH₄⁺

nitrite

NO₂^-

nitrate

NO₃^-

sulfite

SO₃^2-

sulfate

SO₄^2-

hydrogen sulfate (or bisulfate)

HSO₄^-

hydroxide

OH^-

cyanide

CN^-

phosphate

PO₄^3-

hydrogen phosphate

HPO₄^2-

dihydrogen phosphate

H₂PO₄^-

carbonate

CO₃^2-

hydrogen carbonate (or bicarbonate)

HCO₃^-

acetate

C₂H₃O₂^-

permanganate

MnO₄^-

dichromate

Cr₂O₇^2-

chromate

CrO₄^2-

peroxide

O₂^2-

hypochlorite

ClO^-

chlorite

ClO₂^-

chlorate

ClO₃^-

perchlorate

ClO₄^-

oxidation states

• atoms in elemental form = 0

• atoms of Group 1 = +1

• atoms of Group 2 = +2

• hydrogen is either +1/-1

• fluorine is always -1

• oxygen is always -2 unless with F

law of conservation of mass

cannot create or destroy mass

law of constant composition

all samples of a given compound have the same proportions of their constituent elements