CHEM 1106
Giga (G)
109
Mega (M)
106
Kilo (k)
103
Deci (d)
10-1
Centi (c)
10-2
Milli (m)
10-3
Micro (μ)
10-6
Nano (n)
10-9
Pico (p)
10-12
Femto (f)
10-15
Density equation
Mass (m)
Volume (v)
Fahrenheit to Celsius equation
Celsius to Fahrenheit equation
Celsius to Kelvin equation
x + 273
Another name for a homogeneous mix of metals
Alloy or solution
Definition of homogeneous
Same appearance and texture throughout
Definition of heterogeneous
Variable in texture and appearance
Definition of physical change
Alters the appearance of a material only
Definition of chemical change
Creates a new product or substance
Definition of a pure element
One type of element
Definition of a pure compound
Two or more elements in a fixed ratio
Properties of gas particles
Particles in rapid motion experiencing many collisions.
Properties of liquid particles
Particles close together without a regular arrangement.
Properties of solid particles
Particles tight together with regular arrangement.
Separation techniques for gases and liquids
Distillation, chromatography, membrane filteration
Separation techniques for solids
Sublimation, paper/gel chromatography, solubility diferences
Dalton theorized
Each element s made of atoms
Lavoisier stated
atoms are neither created or destroyed (Law of conservation of mass)
Proust stated
New compounds can be formed and compounds consist of specific types and numbers of atoms (Law of constant composition)
Dalton stated
Manipulating reaction parameters can create different compounds from the same atoms (Law of multiple proportions)
Discovered by J.J. Thompson
Electrons (e-) found using magnets and cathode-ray tubes
Discovered by R. Millikan
Electron (e-) mass found by exposing oil to e- and observing the variable rate at which they fell (or rose).
Discovered by E. Rutherford
Protons found by exposing gold foil to alpha-decay and measuring the angles of deflection.
Charge of proton
(+) positive
Discovered by J. Chadwick
Neutrons found by exposing beryllium foil to alpha-decay which knocked out a neutron that collided with a wax layer and knocked out a detectable proton.
Definition of an isotoope
The same element with differing numbers of neutrons.
Periodic orientation of groups (families)
Vertical
Periodic orientation of periods
Horizontal
The metalloids
Boron, silicon, geranium, arsenic, antimony, tellurium, and astatine
Have No Fear Of Ice Cold Beer
Mnemonic used to remember diatomic elements H2, N2, F2, O2, I2, Cl2, Br2
Properties of metals
Ductility, malleability, conductivity, luster. Forms Ionic bonds with non-metals (electrolytes).
Properties of non-metals
Naturally occur as solid, liquid, and gas. Form molecular compounds with non-metals.
Example of isotopes
13C, 14C, 12C
Definition of beta (β)-minus decay
The decay of a neutron into a proton. Releases an electron (e-).
Definition of beta (β)-plus decay
The decay of a proton into a neutron. Releases a positron (e+)
Example of beta (β)-minus decay
Example of beta ( β)-plus decay
( β)-decay tissue penetration
5mm
Released in alpha (α)-decay
Example of alpha (α)-decay
Tissue penetration alpha (α)-decay
Approximately 3cm
Result from decay of meta-stable nuclei or when positrons (e+) collide with electrons e-
Gamma (γ)-radiation
Example decay of meta-stable nuclei
Tissue penetration of gamma (γ)-radiation
50cm
Definition of half-life
Amount of time it takes 1/2 of nuclei to decay
Symbol for half-life
t1/2
Characteristics of ionic (metal to non-metal) compounds
Brittle and conductive when aqueous
Characteristics of molecular (non-metal to non-metal) compounds
hard or soft, not conductive
How to find an elements charge, not including transition metals.
+1 for each group moving right until carbon, then -1 until helium
Hydronium
H₃O⁺
Ammonium
NH₄⁺
Hydroxide
OH⁻
Cyanide
CN⁻
Nitrate
NO₃⁻
Nitrite
NO₂⁻
Carbonate
CO₃²⁻
Hydrogen Carbonate
HCO₃⁻ (g)
Sulfite
SO₃²⁻
Sulfate
SO₄²⁻
Phosphate
PO₄³⁻
Hydrogen Phosphate
HPO₄²⁻ (g)
Dihydrogen Phosphate
H₂PO₄⁻ (g)
Acetate
C₂H₃O₂⁻
Permanganate
MnO₄⁻
Chromate
CrO₄²⁻
Dichromate
Cr₂O₇²⁻
Perchlorate
ClO₄⁻
Chlorate
ClO₃⁻
Chlorite
ClO₂⁻
Hydrosulfuric Acid
H₂S (aq)
Hypochlorite
ClO⁻
Hydrofluoric Acid
HF (aq)
Hydrobromic Acid
HBr (aq)
Hydroiodic Acid
HI (aq)
Nitric Acid
HNO₃ (aq)
Hypochlorous Acid
HClO (aq)
Nitrous Acid
HNO₂ (aq)
Chlorous Acid
HClO₂ (aq)
Chloric Acid
HClO₃ (aq)
Sulfuric Acid
H₂SO₄ (aq)
Perchloric Acid
HClO₄ (aq)
Sulfurous Acid
H₂SO₃ (aq)
Phosphoric Acid
H₃PO₄ (aq)
Phosphorous Acid
H₃PO₃ (aq)
hydrochloric acid
HCl (aq)
hydrogen chloride
HCl (g)
hydrogen fluoride
HF (g)
hydrogen iodide
HI (g)
hydrogen bromide
HBr (g)
hydrogen sulfide
H2S (g)
hydroselenic acid
H2Se (aq)
hydrogen selenide
H2Se (g)
bromic acid
HBrO3 (aq)
iodic acid
HIO3 (aq)
carbonic acid
H2CO3 (aq)
bromous acid
HBrO2 (aq)
iodous acid
HIO2 (aq)
hypobromous acid
HBrO (aq)
hypoiodous acid
HIO (aq)
hypophosphorous acid
H3PO2 (aq)
perbromic acid
HBrO4 (aq)
periodic acid
HIO4 (aq)
bromate
BrO3 -
iodate
IO3 -
bromite
BrO2 -
iodite
IO2 -
phosphite
PO3 -3
hypobromite
BrO-
hypoiodite
IO-
hypophosphite
PO2 -3
perbromate
BrO4 -
periodate
IO4 -
Mnemonic to remember most oxyanions
Nick the Camel ate Clam for Supper in Phoenix
Equal to # of consenants
# of oxygen (nick the camel...)
Equal to # of vowels
# of charge (nick the camel...)
Occurs when oxyanion gains an oxygen
Gains prefix "per"
Occurs when a gaseous oxyanion loses an oxygen
Gains suffix "ite" (g)
Occurs when an oxyanion loses 2 oxygen
Gains prefix "hypo"
Occurs when an aqueous oxyanion loses an oxygen
Gains suffix "ous"
Occurs when an aqeuous diatomic element ("have no fear") has all charge replaced by hydrogen
Gains prefix "hydro"
Occurs when some charge of a gaseous oxyanion is replaced by hydrogen
gains prefix "hydrogen"
The ratio moles in H2O
2 mol Hydrogen
1
mol Oxygen
The ratio of moles in H2O2
2 mol Hydrogen
2 mol Oxygen
Equals molar mass of CH4 (MMCH4)
MMC + 4 x MMH (12.011 g/mol + 4 x 1.0079g/mol)
The unit for molar mass (MM)
grams (g)
moles (mol)
The symbol for molar mass
MM
What you do when given grams
CONVERT TO MOLES
Find the empirical formula given grams or moles of the elements in a compound
Divide all mol by lowest mol to get the number of each element in the compound. Multiply to as necessary to get whole number.
Find molecular formula
Multiply each empirical element by quotient of MMmolec. / MMemp.
The limiting reagent
The reactant completely consumed, leaving other reactants partially consumed.
Water in a solution.
The solvent
Sugar in a solution
The solute
The molarity at which a solute is considered insoluble in a solvent.
<0.01 M (mol/L)
Symbol for molarity
M
Equation for molarity (M)
moles (mol)
liter (L)
Example of alternate concentration concerning occupation of space
% by volume (vol/vol)
Example of alternate concentration concerning mass
% by mass (m/m or w/w)
Given moles of CO, equation to find grams of C
Given grams of CO, find moles of O
Given grams of Mg(OH)2, find grams of Mg
Find percent composition of Mg in Mg(OH)2
Find empirical formula given 13.65% carbon and 86.35% fluorine
Di
2
Tri
3
Tetra
4
Penta
5
Hexa
6
Definition of percent yield
The amount of theoretical yield that was actually obtained.
Equation for percent yield
Actual Yield
Theoretical Yield
Commonly referred to as Redox
Reduction - oxidation reactions
Reduction Involves Gain (RIG)
Refers to a reactant gaining an electron
Oxidation Involves Loss (OIL)
Refers to a reactant losing an electron
True about the reactant that oxidizes another reactant
Is the reactant being reduced
True about the reactant that reduces another reactant
Is the reactant being oxidized
The charge of a reducing agent
Increases (+) because it loses an electron
The charge of an oxidizing agent
Decreases (-) because it gains an electron
Standard pressure (atmospheres)
1 atm
Standard pressure supports how much water
10.33m
Standard pressure (mm Hg)
760 mm Hg
Standard pressure (kPa)
101.3 kPa
Standard pressure (torr)
760 torr
Standard pressure (psi)
14.7 psi
Equation for Boyles Law
P1V1 = P2V2
Equation for Charles law
V1
/ T1 = V2 / T2
Equation for Avogadros law
V1 / n1 = V2 / n2
Avogadros number
6.022 x 10^23
Standard temperature and pressure (STP)
0oC and 1 atm
Ideal gas constant (R)
0.08206 (Lxatm)/(molxK)
Amontons Law
P1 / T1 = P2 / T2
Ideal gas law
PV = nRT
The units for pressure and temperature in the combined gas law
atm, oK
Combined gas law
P1V1 / T1 = P2V2 / T2
Equation given pressure (atm), moles (mol), and temperature (K)
P1 / N1T1 = P2 / N2T2
Equation for Daltons law of partial pressure
Ptotal = P1 + P2
Related to partial pressure ratios
Mole (mol) ratio of gases
Equation for gas density with 2 variables
Mass (g) / volume (L)
Equation for gas density (g/L) with 4 variables
Pressure (atm) x Molar mass
(g/mol)
constant R (Latm/molK) x temperature (K)
P x MM
R x T
All gas density relationships
d = m / v = PMM / RT
Diameter of nucleus
10-14m
Volume occupied by electrons
10-10m to 5 x 10-10m
Acronym for wave energies (low to high)
ROY G BIV
Lowest wave energy
Radio waves
Highest wave energy
Gamma waves
Plancks constant
h = 6.626 x 10-34J-s
Speed of light
3.0 x 108
Relationship between wavelength and energy
Inversely proportional v = 1/λ
Equation for energy
E = h x c / λ
Equation for frequency (v) without energy
v = c / λ
Equation for wavelength (λ) without energy
λ = c / v
Units for frequency (v)
1 / seconds(s) or s-
What electrons do when emitting light
Jump down an energy level or orbital
Line spectra Bohr couldn't explain
>He
Found electrons act like waves
Louis De Broglie
De Broglie found electrons act like waves through
Passing β-radiation through a crystal causes diffraction.
Heisenburgs uncertainty principal
Cannot know both electron speed and location.
Idea Heisenburgs uncertainty principal was founded on
The energy required to view an electrons location changes its speed.
Developed by Schroedinger
Wave mechanics
Wave mechanics concerns
Electron energy, mass, and space occupied in probabilities
Definition of an orbital
Area an electron has 90% probability of being found
Orbital with 2 electrons
s orbital
Orbital with 6 electrons
p orbital
Orbital with 10 electrons
d orbital
Orbital with 14 electrons
f orbital
Energy level the s orbital starts at
1
Energy level the p orbital starts at
2, after 2s2
Energy level the d orbital starts at
3, after 4s2
Energy level the f orbital starts at
4, after 6s2
Hund's rule
Each available orbital in an energy level is filled with 1 electron before doubling up
Pauli exclusion principal
No electron in an atom has the same 4 quantum numbers (energy level [n], orbital [l], spatial orientation [m1], spin [ms])
Shorthand for core electrons configurations
Written as last noble gas ([Ar]4s1 for K)
Shared by columns in periodic table
# of valence electrons
Core electrons shield valence electrons from
Nuclear force of the nucleus
Definition of ionization energy
Energy to remove an electron
Rubidium I.E., in relation to hydrogen
Lower ionization energy
I.E increases
From francium to helium
Definition of electron affinity (EA)
Ability of an atom to add electrons