Chemistry 101 EXAM STUDY GUIDE

Strong Acids undergo complete ionization: HI HBr HCl HNO₃ HClO₃ HClO₄ H₂SO₄

Strong Bases

do not undergo complete ionization: KOH NaOH LiOH CsOH RbOH Ba(OH)₂ Ca(OH)₂ Sr(OH)₂

Strong Electrolytes

conduct electricity and undergo complete dissociation: HCl HNO₃ HClO₄ H₂SO₄ NaOH Ba(OH)₂

Weak Electrolytes

conduct some electricity and undergo some dissociation: CH₃COOH HF HNO₂ NH₃ H₂O

Non-Electrolytes

do not conduct electricity and may undergo complete/some dissociation: (NH₂)₂CO (urea) alcohols sugars

Soluble Compounds

...

Insoluble Compounds

...

Accuracy vs Precision

A- how close to a target value P- how close the measurements are to each other (not necessarily the target value)

Temperature Conversions

C->F: (C)(9/5)+32 F->C: (F-32)(5/9) C->K: C+273.15

Kinds of Energies

thermal- from random motion of molecules radiant- from sun chemical- from bonds of substances nuclear- from neutrons and protons potential- from object's position

Unit Ladder

10^18 E-exa 10^15 P-peta 10^12 T-tera 10^9 G-giga 10^6 M-mega 10^3 k-kilo 10^2 h-hecto 10^1 da-deka - 10^-1 d-deci 10^-2 c-centi 10^-3 m-milli 10^-6 -micro 10^-9 n-nano 10^-12 p-pico 10^-15 f-femto 10^-18 a-atto

Conversions of Units

671 kg -> ___mg (big unit) -> (small unit) (3) - (-3) = +6 always subtract 671,000,000 becomes 6.71x10^8 OR: 43 cm -> ___km (small unit) -> (big unit) (-2) - (3) = -5 always subtract .00043 becomes 4.3x10^-4

SigFig Rules

when adding or subtracting, use the the least # of decimal spaces. (3.081-2.0108=1.070) when dividing or multiplying, use the least digits overall. (3.90x2.1=8.2)

Extensive Properties

how much of the material (mass, length, volume)

Intensive Properties

temperature, color, density

JJ Thompson

experiment: cathode ray tube "electrons exist"

Milikan

experiment: oil drop "mass of electron"

Rutherford's Model of an Atom

experiment: gold foil "atoms are made up almost entirely out of empty space." atomic radii= 100pm or 1x10^-10m nuclear radii- 5x10^-3pm or 5x10^-15 m

Atomic Number

= # of protons (possibly # of electrons, if no charge on element)

Dalton's Atomic Theory

elements are composed of atoms. all atoms in one element are identical, but vary from other elements. compounds are composed of atoms of two or more elements. in chemical reactions atoms are not created/destroyed.

Mass Number

= # of protons + # of neutrons

Isotope

atoms of the same element (same number of protons) with different # of neutrons.

Average Atomic Mass

weighted average of all naturally occurring isotopes in one element. example: 12C is 98.9% 13C is 1.10% (.989)(12amu)+(.0110)(13amu)=12.01 amu

Diatomic Molecules

only two atoms (not coefficient, but rather the subscript) example: N₂ or F₂ or O₂ and so on

Polyatomic Molecules

more than two atoms (not coefficient, but rather the subscript) example: H₂O or O₃ or NH₃ and so on

Wavelength λ

the distance between identical points in waves. (λ) aka: lambda speed (c) of wave= λ x v (<- "nu")

Amplitude

vertical distance from midline to peak/trough

Frequency v

# of waves passing through a point within one second. (v) aka: nu SI Unit- Hz "Hertz" v = c / λ

Maxwell

"visible light consists of electromagnetic waves" (c) or speed of light= 299,792,458 m/s ΔE = h x c / λ

Electromagnetic Spectrum

energy and wavelength are inversely proportional. energy and frequency are directly proportional.

Plank

"radiant energy emitted by an object at a certain temperature is dependent on its wavelength." (h) aka: planks constant h= 6.63x10^-34 E = h x v

Einstein

"light has the nature of particles and waves" photon- particle of light KE = h x v - w (<- work) v = w /h

Bhor

"light is emitted as one electron from an energy level moves to a lower energy level" E = -R x ( 1/n² - 1/n² ) initial final (R) aka: Rydberg's Constant R= 2.18x10^-18 (n) aka: principal quantum number

De Broglie

"electron energy is quantized because it is both a wave and a particle" so, λ = h / m x v (<- velocity here) when given mass (m), use this equation to calculate the wavelength. the more mass, the smaller the wavelength

Heisenberg's Uncertainty Principle

a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time. Δx · Δ(mv) ≥ h/4π solve for the values directed

Quantum Numbers

n, l, ml, ms

n

principal quantum number. tells the distance of an electron from the nucleus may be 1,2,3,4

l

angular momentum quantum number. tells the shape of the orbital the electron occupies at least n - 1 = l ( if n=0, l cannot be +1) *if going by block elements: S-Block= 0 P-Block= 1 D-Block= 2 F-Block= 3

ml

magnetic quantum number. tells the orientation of the orbital range: -l -> +l

ms

spin quantum number. tells the spin of the electron only +1/2 or -1/2

Pauli Exclusion Principle

"the set must have at least one difference." set of n,l,ml,ms can be 4,1,-1,+1/2

Hunds Rule

"the most stable arrangement of electrons in a sub shell is the one with the greatest number of parallel spins." so, pair up last if "_ _ _" fill the shell with one electron each then write in the counter electron. if set of n,l,ml,ms is 4,1,-1,+1/2 the magnetic spin (ms) must be -1/2 on the other in order to pair.

Afbau

"fill up electrons in the last energy orbitals first, then move up."

Orbitals

energy of orbitals depend on the principal quantum number (n). in a multielectron atom, energy depends on n AND l. strength: filled > empty > half filled > partial filled

Magnetic Properties of Orbitals

unpaired electron shells are paramagnetic paired electron shells are diamagnetic

Electron Configuration

how the electrons are distributed among the various atomic orbitals in an atom. 2s² 2= principal quantum # (n) s= angular momentum (l) ²= number of electrons in that subshell 2s² is the EC for Be.

Ionization Energy

minimum amount of energy required to remove an electron. the trend on the periodic table is increasing as right and up. the larger radii, the easier to remove. radii increases left and down

Electron Affinity

negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion. if F (g) + e- -> F-(g) the ΔH = -328 so EA = +328 kJ/mol the trend on the periodic table is increasing as right and up. (forming a bond)

Electron Negativity

the availability to attract electrons toward itself in a chemical bond. (already in a bond)

Lattice Energy

energy required to separate one mol of solid ionic compound into gaseous ions. E = K x ( (Q++Q-) / r ) E = potential energy Q+ = cation charge Q- = anion charge r = radius K = LE increases as Qtotal increases or r decreases

Ionic Bonding

electrostatic force holding together an ionic compound. metal + non metal

Covalent Bonding

electrostatic force holding together a covalent compound. non metal + non metal

Polar Bonds

are covalent bonds. have greater electron density around one of the two or more atoms within the bond.

Classification of Bonds

electronegativity diference: x < .5 - covalent (share electrons) .5 < x < 1.5 - polar covalent (partial electron transfer) x > 1.5 - ionic (electron transfer)

Bond Lengths

single > double > triple

Bond Strengths

triple > double > single

Bond Enthalpy

enthalpy change needed to break a specific bond. triple > double > single

Enthalpies of a Reaction

ΔH of a reaction = sum of products - sum of reactants aka: ΔH(rxn) = Σ(prod) - Σ(reac)

Naming Ionic Compounds

if monatomic, add the ending -ide to the anion. -BaCl₂ is barium chloride if polyatomic, use the name of the whole anion. -NaNO₃ is sodium nitrate with transition metals, identify charge of metal with roman numerals: FeCl₂ is iron(ii) chloride

Naming Covalent Compounds

use prefixes to identify the number of atoms of each element. element furthest left & down in a group is placed first in the name.

Naming Acids

oxoacid- addition of H+ oxoanion- removal of H+

Oxoacid

HClO₄: per- -ic HClO₃: -ic HClO₂: -ous HClO: hypo- -ous

Oxoanion

ClO₄-: per- -ate ClO₃-: -ate ClO₂-: -ite ClO: hypo- -ite

Dipole Moments

compounds with electron rich regions creating a force to the more electronegative side. (μ) aka: dipole μ = Q x r dipole increases in electron rich regions.

Lone Pair + Lone Pair

109.5

Lone Pair + Bonding Pair

107.3

Bonding Pair + Bonding Pair

104.5

Molecular Geometry

geometry: bond angles: linear 180 trigonal planar 120 bent <120 tetrahedral 109.5 trigonal pyramidal <109.5 bent <109.5 trigonal bipyramidal 90, 120, 180 seesaw 90, 120, 180

Hybridization

process of orbital mixing. new atomic orbitals are formed from 2 or 3 different kinds of orbitals. sp , sp² , sp³ , sp³d , sp³d²

sp

80

sp²

120

sp³

109.5

Sigma Bonds

end to end overlap. form a single bond. (longer=weaker)

Pi Bonds

overlap of two regions. form double or triple bonds. (one sigma bond + as many pi bonds) (shorter=stronger)

VB Theory

"covalent bonds form when orbitals of two atoms overlap" 1. must consist of opposing spins of bonding electrons. 2. increase in overlap = increase in stability/strength of the bond. 3. different atomic orbitals account for different bond angles.

AMU

1 amu = 1.66x10^-24 g or 1 g = 6.022x10^23 amu

Avogadros Number

6.022x10^-23

Mass to Moles (n)

mass = m molar mass = M n = m/M

Moles to Mass (m)

moles = n molar mass = M m = n x M

Moles to Atoms (N)

moles = n avogadros # = Na N = n x Na

Atoms to Moles (n)

atoms = N avogardos # = Na n = N / Na

Percent Composition of an Element in a Compound

( moles x mm of element / mm of compound ) x100

Empirical Formulas from Mass Percent

mass % -> mass in g (m/M)-> mol of element -> mol ratio of element (mol/total mol) -> empirical formula

Grams to Mol / Mol to Grams

in stoich: mass of compound A (m/M) -> mol of compound A -> mol of compound B (using ratio from balanced equation B/A) -> mass of compound B (ratio x M)

Limting Reagent

reactant used up first in the reaction.

Theoretical Yield

amount of product that would result is all of the limited reactant reacted. aka: max obtainable result

Actual Yield

amount of product actually obtained from the reaction.

Percent Yield

( actual / theoretical ) x 100

Solute

substance in small amount

Solvent

substance in large amount

Solution

mixture of two or more substances

Concentration

amount of solute present in a given quantity of solution. aka: molarity

Molarity

M = mol of solute / volume solution

Dilution

procedure for preparing a less concentrated solution from a more concentrated solution m1 v1 = m2 v2 m = mol v = volume

Solubility

max amount of solute that will dissolve in a given quantity of solvent at a specific temperature. like dissolves like

Arrhenius Acids/Bases

A- produces H+ (H₃O+ in H₂O) B- produces OH- in H₂O

Bronsted Acids/Bases

A- proton donor B- proton acceptor example: NH₃+H₂O -> NH₄+OH NH₃ - acceptor (B) H₂O - donor (A) NH₄ - conjugate (A) OH - conjugate (B)

Reduction

gain of electron (more - charge)

Oxidation

loss of electrons (more + charge)

Reduction/Oxidation Reagents

the compound causing reduction / oxidation aka: what is O is the R agent and what is R is the O agent

Tritrations

a solution of known concentration is added ti another unknown solution concentration until the chemical reaction between the two is complete. to solve: g / mm = mol then mol / M = v or M x v = mol then g / mol = mm

Equivalence Point

point at which the reaction is complete

Indicator

substance changing color at or near the equivalence point.

Exothermic Reaction

process of system giving off heat to the surroundings. (-)

Endothermic Reaction

process of system gaining heat from the surroundings. (+)

Energy Level Diagrams

more stable compounds are located lower on the diagram. exo- products are lower on the diagram than reactants (-) endo- products are higher on the diagram than reactants (+)

Internal Energy

ΔU

q

heat exchange between the system and surroundings gives off- (-q) gains- (+q)

w

work done by or on the system done by- (-w) done on- (+w)

Piston Movement

outwards- (-w) inward- (+w)

Thermodynamics Equations

ΔU = q + w w = -P x ΔV

Converting L x atm -> J

multiply x 101.3 J

State Functions

properties determined by the state of the system regardless of the how the conditioned was achieved. examples: ΔU , ΔV , ΔP , ΔT WORK is not a state function

Specific Heat Capacity

the amount of heat needed to raise the temperature of one mol of the substance by one degree C.

Equation for Specific Heat

q = mcΔt

ΔHfusion

(melting) on diagram B to C to solve: solve for q = mcΔt and add kJ to the kJ of (g/mm = mol x heat of fusion)

ΔHvaporization

(boiling) on diagram C to D to solve: solve for q = mcΔt and add kJ to the kJ of (g/mm = mol x heat of vaporization)

Melting

solid -> liquid

Freezing

liquid -> solid

Vaporization

liquid -> gas

Condensation

gas -> liquid

Sublimation

solid -> gas

Deposition

gas -> solid

Gas Unit Converions

1 atm = 760 mmHg = 760 torr = 101,325 Pa = 14.7 psi

R

R = .08206 L x atm / mol x K R = 8.314 J / mol x K

Ideal Gas Law Equation

PV/nRT = PV/nRT cancel units not needed and solve for those asked

V and P Relationship Boyle's Law

inversely proportional

V and T Relationship Charles' / Gay-Lussac's Law

directly proportional

V and n Relationship Avogadros Law

directly proportional

STP Conditions

t = 0 C p = 1 atm

Pressure

increase in distance from sea level = decrease in pressure

Partial Pressure

PP = mol fraction x P total

Mole Fraction

mole faction = mole of gas / mol total

Manometers

P of gas = Ph (closed tube) P of gas = Ph + P atm (open tube) as Ph increases, the volume of the sample decreases

Density Equations

d = m / V d = P x mm / R x T

Velocity

KE = 1/2mv² V = √(3RT/M)

Gas Diffusion

gradual mixing of molecules of one gas with molecules of another by virtue of their kinetic energy. equation: r1/r2 = √mm2/mm1

Gas Effusion

process by which a gas under pressure escapes from one compartment to another through a small opening. equation: t2/t1 = √mm2/mm1

Van der Waals

non ideal gas equation: (P + an²/v²)(v-nb) = nRT

Intermolecular Forces

attracted forces between molecules .

Ion-Ion

ionic molecule + ionic molecule

Hydrogen Bond

H - O H - N H - F

Dipole-Dipole

polar molecule + polar molecule - lone pairs on central atom

Induced Dipole / London Dispersion

non polar molecule + non polar molecule (covalent molecule + covalent molecule) -hydrocarbon -diatomic molecule -no lone pairs on central atom ALL have dispersion forces

Determining Strength of I.M. Forces

strength: ion-ion > hydrogen bond > dipole-dipole > dispersion if the same molecular force: higher molecular weight- the stronger the force EXCEPT with ion-ion with ion-ion: find charges within elements of compound and multiply them with each other- larger product is stronger.

Surface Tension

amount of energy required t increase the surface of a liquid by unit area. strong IM forces = high surface tension

Cohesion

IM force between like molecules

Adhesion

IM force between unlike molecules

Viscosity

resistance to flow strong IM forces = high viscosity

Simple Cubic

unit cell. s = 2r (s) aka: side (r) aka: radius

Body Centered Cubic

unit cell. s = 4r / √3 (s) aka: side (r) aka: radius

Face Centered Cubic

unit cell. s = √8r (s) aka: side (r) aka: radius

Substitutional Alloy

takes up space normally occupied by host metal. hence: sub

Interstitial Alloy

occupies spaces between atoms of host metal. hence: in

Semiconductor

has a wide band gap allowing few electrons to jump onto conduction band

n-type

has extra electrons that allows for conductivity

p-type

has a "hole" or is missing electron that allows conductivity

Vapor Pressure

ln(p1/p2) = (ΔHvap/R) x (1/T₁ - 1/T₂) "the higher the everything, the lower the VP"

R

R = .08206 L x atm / mol x K R = 8.314