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176 Cards in this Set
- Front
- Back
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pico |
10^-12 |
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nano |
10^-9 |
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micro |
10^-6 |
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milli |
10^-3 |
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kilo |
10^3 |
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Mega |
10^6 |
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Accuracy |
The answer is correct (accurite) |
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Precision |
The answer is repeatable |
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Types of matter |
Pure substances and Mixtures |
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Types of Pure substances |
Elements -- Things on the periodic table Compounds -- Things created with elements |
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Types of Mixture |
Homogeneous -- Mixture of a single compound Heterogeneous -- Mixture of multiple compounds |
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Periodic Table: Metals |
Things in the middle |
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Periodic Table: Metalloids |
The things on the "Staircase" |
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Periodic Table: Nobel Gases |
The things in the last column (18), barely react with anything
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Periodic Table: Halogens |
2nd to last column. Like to connect with things |
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Periodic Table: Non-metals |
Things above the staircase, used in organic molecules |
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Periodic Table: Alkali Metals |
First column, very reactive |
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Periodic Table: Alkaline Earth Metals |
Second column, reactive |
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Law of Definite Proportions |
Molecules of the same type will always have the same mass, by this we can identify them.
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Law of Multiple Proportions |
Atomic ratios are unique to each molecule. I.E., a molecule of a certain type will always be made of the same atoms and the same # of each type. |
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Protons |
Reside inside the n.ucleus, Have a positive charge. |
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Neutrons |
Reside in the nucleus, Majority of the mass of an atom. NO CHARGE |
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Electrons |
Orbit around the central nucleus. Participate in chemical reactions and bonding. Have negative charge. |
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Isotope |
An atom that has the normal # of protons for an element, but has more or less neutrons than normal. |
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Ions |
Atoms that have more or less electrons than normal |
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Average Atomic Mass |
Atomic mass is a weighted average, meaning that it is calculated using values and percentages of how much they matter. In chemistry, it depends on how often that element is found in nature. |
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Light: Wavelength |
Length from edge to edge |
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Light: Frequency |
How often the waves cycle, determine energy |
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Light: Relationship of frequency and wavelength |
Speed of Light = frequency * wavelength |
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Light: Energy of light |
E = h(Constant) * frequency |
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Quantization |
The idea that things can be measures in very small steps. In light, the idea started after observing "Blackbody" radiation, or the light from hot things. It proved that light was quantizable, and thus proved the existence of the proton. |
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Bohr Model |
Complicated equation that predicts electron location. PSI = allowed electron energies PSI^2 = Probability density of where they will be located |
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Debrogli Wavelength |
Wavelength of matter: Wavelength = H/mV |
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Paramagnetism |
The atom contains at least one unpaired electron |
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Diamagnetism |
The atom contains only paired electrons |
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Electron Configurations |
Fill the 4 layers: s, p, d, f linearly, it won't progress to the next until the previous is full. Fill in the following order: S1, S2, P2 s -- 2 electrons p -- 6 electrons d -- 10 electrons f -- 14 electrons |
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Expanding the octet |
Only happens in the third period or below. Common in sulfar and phosphorus. ONLY DO IT TO MINIMIZE FORMAL CHARGE. |
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Ground state |
When an electron is in the lowest available electron orbital. |
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Excited electron state |
An electron has been energized to be in a higher than normal orbital. |
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Ionization Energy |
Energy needed to remove an electron. Higher the farther to the right you go on the periodic table as the atoms are bigger and pull more on their electrons. |
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Electron Affinity |
Energy needed to add an electron. Generally higher as you move to the right on the periodic table, except for the Noble gases, who don't want more electrons at all. |
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Electronegativity |
The combination of Electron Affinity and Ionization energy. |
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Lewis structures of Molecules |
A visual way to represent the way the atoms are bonded together to form the molecule. |
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Bond Order |
Equal to the sum of all bonds in an atom divided by the number of bonds. |
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Resonance |
The idea that two molecules can have the same atomic placement, but different bonds. |
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Formal charge |
Done on lewis structures to estimate the stability of the molecule. Each atom is assigned a formal charge, and then they are added together to get a total formal charge. Done by calculating how many electrons an atom has normally, and then subtracting how many electrons it has on the lewis structure.
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Molecular Geometry |
The shape that the molecule has based on a central atom. |
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Linear |
2 atoms -- 180 bond angle 2 atoms, 3 lone pairs -- 90, 120 bond angle |
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Trigonal Planar |
3 atoms, 120 bond angle |
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Tetrahedral |
4 atoms, 109.5 bond angle |
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Trigonal Pyramidal |
3 atoms, 1 electron, <109.5 bond angle |
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Trigonal bipyramidal |
5 atoms, 90, 120 bond angle |
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Octahedral |
6 atoms, 90 bond angle |
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Bent Shape |
2 atoms, 2 lone pairs, <109.5 bond angle |
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See Saw |
4 bonded atoms, 2 lone pairs |
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Steric Number |
# of atoms + # of lone pairs |
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Sigma bonds |
Formed with single bonds |
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Pi bonds |
Formed with double bonds, but they also form a sigma bond. |
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Polar Molecules |
Molecules that their shape and the difference in their electronegativities give them an overall dipole moment. |
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Molecular Orbital Theory |
Electrons fill the orbitals sequentially. |
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S Orbitals |
Shape -- O 1 Orbital Energy Level 1 0 angular nodes |
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P Orbitals |
Shape --> 8 3 orbitals energy level 2 1 angular node |
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D Orbitals |
shape --> 2 8's intertwined 5 orbitals energy level 3 2 angular nodes |
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F orbitals |
7 orbitals energy level 4 3 angular nodes |
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Bond Length |
Depends on the strength of the bond. Stronger bonds mean less distance between atoms. |
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Lattice Energy of Ionic compounds |
The smaller the ions are the stronger the Lattice Energy. This is because of Couloumb's law |
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Inter-Molecular Forces: From weakest to strongest |
1. London Dispersion (Van der Waals) 2. Dipole-induced Dipole 3. Hydrogen bonds 4. Dipole to Dipole 5. Ion to Dipole 6. Ionic |
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Inter-Molecular Forces: London Dispersion |
Non-polar Ex: propane |
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Inter-Molecular Forces: Dipole-Induced Dipole |
Only occurs between two different substances, non-polar |
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Inter-Molecular Forces: Hydrogen |
Occurs between Hydrogen and Nitrogen, Oxygen, or Florine |
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Inter-Molecular Forces: Dipole-dipole |
EX: Acetane |
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Inter-Molecular Forces: Ion to dipole |
Ion to polar Ion dissolved in another substance EX: Na+ in Water |
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Inter-Molecular Forces: Ionic |
Only occurs between two different substances. Anion to Cations |
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Boiling point and inter-molecular forces |
The stronger the bond, the higher the boiling point |
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Organic Functional Groups: Alkane |
R--H Molecule created with only Hydrogens and Carbons that form only single bonds |
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Organic Functional Groups: Alkene |
C2R4 Molecules created with one or more double bonds between carbons |
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Organic Functional Groups: Alkyne |
R--C==--C--R Molecules created with one or more triple carbon to carbon bonds |
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Organic Functional Groups: Aromatic |
C6R6 Arranged in a circle with alternating single and double bonds |
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Organic Functional Groups: Amine |
R--NH2 R--NHR R--NR2 |
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Organic Functional Groups: Alcohol |
R--OH |
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Organic Functional Groups: Ether |
R--O--R |
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Organic Functional Groups: Aldehyde |
R--C(=O)--H RC(O)H |
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Organic Functional Groups: Ketone |
R--C(=O)--R RC(O)R |
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Organic Functional Groups: Carboxylic Acid |
R--C(=O)--OH RC(O)OH |
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Organic Functional Groups: Ester |
R--C(=O)--OR RC(O)OR |
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Organic Functional Groups: Amide |
R--C(=O)--NH2 RC(O)NH2 |
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Empirical Formula |
The simplest formula for a compound. The subscripts are not divisible by a number |
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Molecular Formula |
A formula that gives the composition of one molecule. The subscripts are usually divisible by the same number |
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Percent to mass, mass to mole, divide by small, multiply to a whole |
Get Molecular Formula from masses |
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Heat |
Heat is a representation of how much energy each individual atom has |
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Internal energy |
How much energy a substance has |
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Work |
Force * Distance |
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Energy |
q + work |
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Specific Heat Capacit |
Specific heat is a measurement of the amount of energy needed to heat a substance 1 degree celcius |
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Heat flow in heating and cooling |
Temperature increases (or decreases) linearly with the slope of the specific heat until it reaches a phase change. It then sits there until the necessary energy is expended (or added) to change the mass into the new phase. |
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Enthalpy |
Energy needed for a reaction |
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Calculating bond enthalpies |
Bond's broken - bond's formed |
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Dihydrogen Phosphate |
H2PO4^- |
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Acetate |
C2H3O2^- CH3COO^- |
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Hydrogen Sulfite |
HSO3^- |
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Hydrogen sulfate |
HSO4^- |
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Hydrogen Carbonate |
HCO3^- |
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Nitrite |
NO2^- |
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Nitrate |
NO3^- |
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Cyanide |
CN^- |
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Hydroxide |
OH^- |
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Permanganate |
MnO4^- |
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Hypchlorite |
ClO^- |
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Chlorite |
ClO2^- |
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Chlorate |
ClO3&- |
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Perchlorate |
ClO4^- |
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Azide |
N3^- |
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Thicyanate |
SCN^- |
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Hydrogen Phosphate |
HPO4&2- |
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Oxalate |
C2O4^2- |
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Sulfite |
SO3^2- |
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Sulfate |
SO4^2- |
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Carbonate |
CO3^2- |
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Chromate |
CrO4^2- |
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Dichromate |
Cr2O7^2- |
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Silicate |
SiO3^2- |
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Disulfide |
S2^2- |
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Phosphate |
PO4^3- |
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Phosphite |
PO3^3- |
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Ammonium |
NH4^+ |
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Percent Yield |
Percent Yield means the percent of how much was produced compared to how much it should have produced. |
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Limiting Reactants |
The limiting reactant of a molecular formula is the part that will get used up first in a reaction. To figure out the limiting reactant, you must calculate how much could be produced with each reactant if the others were infinite. The one that produces the least is the limiting reactant |
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Enthalpies of Formation |
This is the energy required to combine elements into a molecule. Specifically if it says formation, it means it was measured at standard pressure and standard temperature. 1 atm and 20 degrees celsius. |
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Heat of vaporization |
The temperature and energy required to change a liquid to a gas |
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Condensation |
When a gas becomes a liquid again |
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Sublimation |
Going from a solid to a liquid |
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Fusion |
Going from a liquid to a solid |
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Entropy |
A measurement of disorder in a substance |
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Factors that affect entropy |
Temperature, volume, Number of independantly moving molecules |
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Entropy changes in reversable processes |
Entropy doesn't increase. Perfectly reversable is impossible as it would mean no energy is lost |
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Spontaneity and gibb's free energy |
A negative G means the reaction is spontaneous dG = dH - Tds G -- Gibbs free energy H -- dQ, change in heat energy T -- Temperature S -- Entropy |
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Temperature dependance of Gibb's free energy |
At a phase change, Gibbs free energy is 0. At high temperatures, a reaction is more likely to be spontaneous |
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Density of ideal gases |
Use the equation density = mass/volume with the combined gas law |
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Stoichiometry in gas phase reactions |
You know that you can calculate how much product you will get from a certain reaction |
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Combined Gas law |
PV/nT = R P1V1/N1T1 = P2V2/N2T2 |
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Dalton's law of partial pressures |
Pressures add linearly. I.E. the total pressure is a combination of each smaller pressure as calculated by the combined gas law |
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Real vs Ideal gases |
Ideal gases make the assumption that all collisions between atoms are elastic and lose no energy. They also assume no attraction between atoms. This breaks down at high pressures due to a lot of collisions |
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Velocity of gases |
As a general rule, smaller atomic mass means that the gas moves quicker.
V = sqrt (3RT/M) |
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Standard Temperature and Pressure (STP) |
Standard Temperature: 273.15 K Standard Pressure: 1 atm |
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Effusion of gases (When a gas is ejected through a small hole) |
Graham's Law of Effusion(Rate of Effusion)1/(Rate of Effusion)2 = sqrt((Molar Mass)2/(Molar Mass)1) (3/2)kT = (1/2)mv^2 k -- boltzmann constant T -- temperaturem -- mass v -- velocity |
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Saturated solution |
A solution that contains the maximum concentration of a solute possible at a given temperature |
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Unsaturated solution |
A solution that contains less than the maximum quantity of solute predicted to be soluble in a given volume of solution at a given temperature |
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Supersaturated Solution |
A solution that contains more than the maximum quantity of solute predicted to be soluble in a given volume of solution at a given temperature. The only way to create a supersaturated solution is to heat the liquid and add more solute until it is saturated, then the temperature is reduced, resulting in a higher than normal saturation |
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Molarity |
Moles of solute / Volume of solution in liters Results in Moles per liter |
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Acid |
Substance that releases H+ into water |
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Base |
Substance that releases OH- into water |
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pH |
# < 7 -- Acid # > 7 -- Base |
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Electrolytes and conductivity |
Things that conduct electricity in a solution Include: Soluable Ionic compounds Strong acids and bases |
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Soluble Compounds |
Anything with an alkali metal is soluable. Nitrates, bicarbonates, and Clorates. All compounds containing the following ions are soluble in water: Cations: group 1 ions (alkali metals) and NH4+ Anions: NO3- and CH3COO- (acetate) |
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Generally soluble compounds |
Halides: Unless paired with Ag+, Hg^2+, or Pb^2+ Sulfates (SO4^2-): Unless paired with Ag+, Ca^2+, Sr^2+, Ba^2+, Hg2^2+, and Pb^2+ Compounds containing the following anions are soluble except as noted: Group 17 ions (halides), except the halides of Ag+, Cu+, Hg2^2+, and Pb^2+SO4^2-, except the sulfates of Ba^2+, Ca^2+, Hg2^2+, Pb^2+, and Sr^2+ |
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Net Ionic equations |
The idea is that you only write the ions you care about from the compounds |
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Spectator ions |
Ions that don't participate in the acid/base balancing |
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Precipitation reactions |
Anything that isn't soluble will precipitate, or separate and fall to the bottom of the solution. Base it on the solubility rules explained above. |
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Stoichiometry in acid base reactions |
Basically your goal would be to create water. Sometimes called a double replacement because 2 atoms switch partners |
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Get the conjugate Acid |
Add a hydrogen to the base |
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Get the conjugate base |
Remove a single proton (hydrogen) from the acid |
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Dilutions of solutions |
ppm -- parts per million -- mg/L cc -- mL -- cm^3 |
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pOH |
Opposite of pH. Basically the pOH of something is 14 -pH |
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Neturaliation reactions |
When an acid and a base combine to form water. There must be an H to balance every OH |
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Acid strength and k values |
K values tell how likely it is to break apart. Rouphly it tells the proportion of atoms to break apart (forming ions) to those that don't. |
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Factors that affect acid strength |
Bond strength to they hydrogen. The weaker the bond, the stronger the acid. A highly polar bond between the hydrogen and another molecule will make a stronger acid as the hydrogen will more likely separate. |
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Oxidation Reactikon |
Reaction where the focused atom loses O.N. electrons. Or when a substance combines with Oxygen to have more than it did before |
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Reduction Reaction |
When the focused atom gains O.N. electrons. Or when a substance loses oxygen and leaves a substance and oxygen. |
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Oxidizing agent |
Atom in reaction that loses O.N. electrons |
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Reducing agent |
Atom in reaction that gains O.N. electrons |
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Nomenclature of Alkanes |
1. Find the longest chain in the molecule 2. Number the chain from the end nearest the first substituent encountered 3. List the substituents as a prefix along with the number of the carbons to which they are attached |
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Structural isomers |
Same atom placement but different bonds |
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Geometric isomers |
Same connectivity, but different arrangement in space |
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Accounting for real gases: Van der Waal's Equation |
P = (nRT)/(V-nb) - (n^2a)/V^2 |
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Endothermic reaction |
Heat enters the system and work is done on the system |
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Exothermic reaction |
Heat leaves the system and work is done by the system |
