Solubility Rules:
-NO3-1 all soluble
-OH- all insoluble except Group 1, Ba(OH)2, Sr(OH)2, Ca(OH)2
-Cl- all soluble except AgCl, Hg2Cl2 and PbCl2
-CO3-2 all insoluble except NH4+ and Group 1
-SO4-2 all soluble except BaSO4, PbSO4 and SrSO4
-S-2 all insoluble except Group 1 and Group 2 and NH4+
Polyatomics:
Acids:
-Acetate (C2H3O2-1)
-Carbonate (CO3-2)
-Arsenate (AsO4-3)
-HCl = hydrochloric
-Chlorate (ClO3-1)
-Chromate (CrO4-2)
-Arsenite (AsO3-3)
-H2SO4 = sulfuric
-Chlorite (ClO2-1)
-Dichromate (Cr2O7-2)
-Phosphate (PO4-3)
-HNO3 = nitric
-Cyanide (CN-1)
-Oxalate (C2O4-2)
-Phosphite (PO3-3)
-CH3COOH = acetic
-Hydroxide (OH-1)
-Sulfate (SO4-2)
-Ammonium (NH4+1)
-H3PO4 = phosphoric
-Nitrate (NO3-1)
-Sulfite (SO3-2)
-Bicarbonate (HCO3-1)
-H2CO3 = carbonic
-Nitrite (NO2-1)
-Nitrite (NO2-1)
Matter is the physical material of the universe and exists in 3 states: solid, liquid, gas
Substance is matter that has a uniform and definite composition
Mixture is a physical blend of substances
Homogenous Mixture has uniform composition throughout (solution) Heterogeneous Mixture does not
Physical Properties: observed or measured
Chemical Properties: ability to undergo chemical change
Law of Conservation of Mass: matter/mass can be neither created nor destroyed
Extensive Properties: mass and volume
Intensive Properties: composition
Solid: definite shape, volume, not easily compressed Liquid: indefinite shape, definite volume, not easily compressed
Gas: indefinite shape, volume, easily compressed
*Particles are always moving in S, L, G
Indications of Chemical Reaction: dissolving, releasing/absorption of heat, production of gas, color change, precipitate
Qualitative give descriptions, no numbers
Quantitative number-based
Accuracy is how close to a single measurement
Precision is how close several measurements are together
Error = |expt – accepted| % Error = |expt-accepted|/accepted * 100 Percent Yield = expt/theoretical * 100
micro=106
1 kilo(unit) = 1000 (unit) 1 (unit) = 1000 milli
1 (unit) = 100 centi
K = C + 273
BP=373KDensity = mass/volume
C = K – 273
H2O FP=273K
Sig Figs: For addition & subtraction use least # of decimal places
nano=109
Density of H2O = 1 g/mL
/ For Multiplication & Division use least # of sig figs
Dalton’s Atomic Theory: 1) all elements are composed of atoms
2) atoms of same element are identical
3) atoms can physically mix/chemically combine
4) chemical rxns happen; atoms don’t change
Proton = pos., change = radioactivity
in whole #d ratios
Neutron = neutral, change = isotope
Electron = neg., change = ion
AMU = atomic mass unit = 1/12th Carbon 12 atom (average atomic mass)
Electrons found in electron cloud, move around in orbital, movement creates electricity
Quantum Mechanical Model: atom has no definite shape, increase in # of electrons = increase in size & energy
Sublevels = s (max 2e-), p (max 6e-), d (max 10e-), f (max 14e-)
Energy Levels (n) = 1,2,3,4,5,6,7 (rows on PT)
Aufbau Principle: electrons enter @ lowest E level
Pauli Exclusion: atomic orbital may have at most 2 electrons
Hund’s Rule: when electrons occupy orbitals, one electron enters each orbital until all orbitals contain one electron
**Exceptions based on Aufbau principle; due to stable electron-electron interactions; now unstable
__________________________________________________________________________________________________
Frequency = Speed of Light/Wave Length
Speed of Light = 3.0 x 108 m/s or 3.0 x 1010 cm/s
Atomic Spectra Principles:
1) Light is a form of energy called Electromagnetic Radiation
2) Waves have wavelength and frequency
3) Electromagnetic Radiation includes broad spectrum of radiant E
4) Speed of wave = wavelength x frequency
5) Energy is directly proportional to frequency; freq. & w.l. are inversely
*When you excite electrons you increase energy which is given off as heat and light = atomic spectra
Dmitri Mendeleev arranged elements by ATOMIC MASS
Periods are rows / energy levels (1-7)
Henry Moseley arranged by ATOMIC NUMBER
Groups are vertical columns (18); A = representative elements, B = transitions
Periodic Law states when elements are arranged in order of incr. atomic #, there is a periodic pattern in properties
Noble Gases – inert gases (stable) VIIIA
Halogens – salt formers VIIA
Alkali Metals – most reactive IA
Transitions – group B
Lanthanides (inner transition) 4f
Alkaline Earth Metals – IIA
& Actinides (inner transition) 5f
Atomic & Ionic Radii – distance bet. 1 nuclear to another
Ionic Size – size of a charged atom
Ionization Energy – amt of energy to remove 1 electron
Electronegativity – amt of energy to attract 1 electron
*Flourine (F) has highest Ionization E & Electronegativity
*Francium (Fr) has greatest atomic/ionic radii & size
Cations give up electrons (metals) Anions take electrons (nonmetals)
Molecular Compounds: composed of molecules, only nonmetals, named using prefixes
Ionic Compounds: composed of ions, have metal as cation, name without prefixes
Monary Cmpd. – 1 element Binary Cmpd. – 2 elements Ternary Cmpd. – 3 elements Diatomics: H2, O2, F2, Br2, Cl2, I2, N2
Prefixes: Mono, Di, Tri, Tetra, Penta, Hexa, Hepta, Octa, Nona, Deca Endings: -ide = periodic table -ite/ate: polyatomic
Ionic bonding – transfer of electrons, makes ionic compounds; cations must be metal
Ions bond to complete their Octet Rule (have full outer electron shell)
Seen through Lewis Dot Diagram
Metallic bonding – a more specific type of ionic bonding
Salts – metal cation + halogen; most are solid at STP, in 3D pattern, conduct electricity as liquid, ductile & malleable
Law of Definite Properties – masses of elements are in SAME PROPORTIONS within same compound
Law of Multiple Proportions – can compare masses of same element between similar compounds via small #d ratios
Covalent bonding – the sharing of electrons; 1 pair = 2 e-, 2 pair = 4 e-, 3 pair = 6 eNonbonding Sites: IA (1; 1 valence e), IVA (4; 4 valence e), VA (3; 5 valence e), VIA (2; 6 valence e), VIIA (1; 7 valence e)
Coordinate Covalent –sharing of unshared pairs (NH3 -> NH4+1)
Resonance – 2+ orientations of same compound (CO 3-2)
Paramagnetism – strong attraction forces to ext. magnetic field (NO2) Dimagnetism – “balanced sharing” (F2)
VSEPR Theory – valence shell electron pair repulsion theory; e- repel each other so bond angles adjust
Trigonal Planar =120 deg. Tetrahedral =109.5 deg. Trigonal Bipyramidal =90 & 120 deg. Linear =180 deg. Bent =120 deg.
Bond Polarity – polar = unequal pulling, nonpolar = equal pulling
Bond Dissociation Energy – energy to break bond (kJ)
Intermolecular Attractions – Ionic > Hydrogen > Dipole > Dispersion
Van der Waals – weakest attractions
Dispersion Forces – weakest, caused by e- motion
Dipole Interactions – polar molecules attracted to e/o
Hydrogen Bonds – H covalently bonded to very EN atom and bonded weakly to unshared e- pair of an EN atom
Avogadro’s Number = 6.02 x 1023 representative particles
THE MOLE ROAD:
Moles  Grams = x Molar Mass
Moles  Rep. Particles = x Av. #
Moles  Volume @ STP = x 22.4
Grams  Moles = / Molar Mass
Rep. Particles  Moles = / Av. #
Volume @ STP  Moles = / 22.4
Percent Composition:
1. Determine MM of each compound
2. Determine % of each element in compound
Percentage  Formula:
1. Assume % = grams
2. Convert grams to moles of each element
3. Divide all mole values by smallest mole value to get “multiplier” for each element
Empirical  Molecular:
1. Find Empirical Formula MM
2. MF mm/EF mm = multiplier
**Know how to balance equations
Types of Reactions: 4 base and 2 special = 6 total
Reactants  Products
1. Synthesis – combination rxn [A+BAB]
2. Decomposition – breakdown rxn [ABA+B]
3. Single Replacement – [A+BYAY+B]
4. Double Replacement – [AX+BYAY+BX]
5. Acid-Base – similar to DR [ACID+BASEH2O+SALT] 6. Combustion – 2 types; only thing that changes is hydrocarbon
COMPLETE: Hydrocarbon + O2  CO2 + H2O
INCOMPLETE: Hydrocarbon + O2  CO + H2O
Solubility Constants and Rules
1-H2O is always a liquid
2-Gases are anything that says “gas”, diatomics, noble gases
3-Solids are all metals except Hg
4-Acids & Bases are always aqueous
Net Ionic Equations – 1. Break apart aq compounds 2. Eliminate spectator ions
3. Rewrite and balance
Stoichiometric Calculation Steps
1. Change GIVEN/HAVE into moles (reactants)
2. Convert HAVE of first reactant into NEED of second reactant via mole to mole ratio from balanced equation
HAVE > NEED = excess
b. HAVE < NEED = limiting reagent
3. Take the moles of LR and convert to moles of the specified product then change moles of product into grams
This is the THEORETICAL YIELD
4. Determine Percent Yield
(Actual Yield/Theoretical Yield) * 100
-------------------------------------------------------------------------------------------------------------------------------------------------------------------------------Molarity – tells concentration of ions in an aqueous solution
M = moles/liters
Kinetic Energy – tiny particles are in constant motion; particle movement affected by volume, pressure, temperature
Evaporation = LG
Vaporization = LG
Condensation = GL
Sublimation = go directly from solid to gas [ex: dry ice = solid CO2]
Glass = amorphous solid (super cooled liquid)
Plasma = gaseous mix of e- and positive ions; very HIGH temp. [ex: Aurora Borealis – the Northern Lights]
To speed up kinetic energy = get faster movement = heat it up
To slow down = become stationary = solidification
Measures entropy (measure of randomness)
For gases, collisions are elastic
Pressure: atm, kPa, mmHg, Tprr [1 atm = 101.3 kPa = 760 mmHg] result of collisions, no collisions in a vacuum
Temperature: Celsius, Fahrenheit, Kelvin; incr. temp = incr KE
Volume: mL or L
Overall Relationship: only 2 change at a time; P&V are inversely proportional; T changes in same direction but ½ rate
Avogadro’s Hypothesis: equal volumes of gases @ same T&P contain equal # of particles [1 mole = 22.4L @ STP]
Real vs. Ideal Gases: ideal is hypothetical
Boyle’s Law: P1V1=P2V2 [constant = T]
Charles’ Law: V1/T1 = V2/T2 [constant =P] OR P1/T1 = P2/T2 [constant = V]
Combined Gas Law: P1V1/T1 = P2V2/T2
Dalton’s Partial Pressure: Ptotal = P1 + P2 + P3 etc…
Ideal Gas Law: PV = nRT
P = pressure [atm (R=0.0821) or kPa (R=8.31)]
V= volume
n = moles
T = Kelvin
**Know Exothermic and Endothermic Reaction Graphs
Energy transferred through the form of HEAT – determined by changes in temp
Exothermic = ∆H is negative, heat LEAVES the system
Endothermic = ∆H is positive, heat ENTERS the system
Calorimeter – device that contains the heat within a unit, includes a lid and insulated holding device (Styrofoam cup)
Q = C x M x ∆T
q = heat in joules OR kilojoules
M = mass in grams
∆T = change in temp [final-initial]
C = specific heat constant – 4.18 joules for water, 1 calorie for water
∆Hfus = molar heat of fusion
∆Hcond = molar heat of condensation
∆Hsolid = molar heat of solidification
∆Hsoln = molar heat of solution
**Know heat to mole ratio
**Know Hess’s Law of Summation
Arrhenius (looks): Acid = H+ or H3O+; Base = OH-
∆Hvap = molar heat of vaporization
Bronsted-Lowry (behaves): Acid = Donates H+; Base = Accepts H+
Acids - # of Hs indicate strength of acid; more Hs means it can ionize more than once
HCl = monoprotic  ionize once H2SO4 = diprotic  ionize twice
H3PO4 = triprotic  ionize thrice
Bases - # of OHs tell solubility of base in water; more OHs means less soluble
pH = -log [H+]
pOH = -log [OH-]
pH + pOH = 14
Neutralization Reactions – Acid + Base  H2O + Salt
Saturated – cannot dissolve more solute
**Solubility Graph!!!
[H+][OH-] = 1 x 10-14
Titration – process of neutralizing an acid or base
Unsaturated – can dissolve more solute
Supersaturated – too much dissolved