17: Reference
- Page ID
- 306813
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Metric Prefixes
| Prefix | Symbol | Multiple | Multiple |
|---|---|---|---|
| Exa | E | 1018 | 1,000,000,000,000,000,000 |
| Peta | P | 1015 | 1,000,000,000,000,000 |
| Tera | T | 1012 | 1,000,000,000,000 |
| Giga | G | 109 | 1,000,000,000 |
| Mega | M | 106 | 1,000,000 |
| kilo | k | 103 | 1,000 |
| hecto | h | 102 | 100 |
| deka | dk | 101 | 10 |
| <base unit> | 100 | 1 | |
| deci | d | 10-1 | 0.1 |
| centi | c | 10-2 | 0.01 |
| milli | m | 10-3 | 0.001 |
| micro | µ | 10-6 | 0.000 001 |
| nano | n | 10-9 | 0.000 000 001 |
| pico | p | 10-12 | 0.000 000 000 001 |
| femto | f | 10-15 | 0.000 000 000 000 001 |
| atto | a | 10-18 | 0.000 000 000 000 000 001 |
Polyatomic Ions
| Ion Name | Ion Formula |
|---|---|
| ammonium | NH4+ |
| cyanide | CN- |
| hydroxide | OH- |
| nitrate | NO3- |
| nitrite | NO2- |
| sulfate | SO42- |
| sulfite | SO32- |
| hydrogen sulfate (bisulfate) | HSO4- |
| carbonate | CO32- |
| hydrogen carbonate (bicarbonate) | HCO3- |
| phosphate | PO43- |
| hydrogen phosphate | HPO42- |
| dihydrogen phosphate | H2PO4- |
| permanganate | MnO4- |
| perchlorate | ClO4- |
| chlorate | ClO3- |
| chlorite | ClO2- |
| hypochlorite | ClO- |
Gas Laws Summary
Boyle’s Law – Pressure vs. Volume
- Pressure and Volume are inversely proportional – as one increases, the other decreases by the same factor. \(P \propto 1/V\)
- \( PV = k\) (a constant) (when T and n are held constant)
- \(P_1 V_1 = P_2 V_2\)
Charles’s Law – Volume vs. Temperature
- Volume and Temperature (in Kelvin) are directly proportional – as one increases, the other increases by the same factor. \( V \propto T\)
- \( \frac{V}{T} \) = constant (when P and n are held constant)
- \(\frac{V_1}{T_1} = \frac{V_2}{P_2}\)
Gay-Lussac’s Law – Pressure vs. Temperature
- Pressure and Temperature are directly proportional. \( P \propto T\)
- \( \frac{P}{T} \) = constant (when V and n are held constant)
- \(\frac{P_1}{T_1} = \frac{P_2}{P_2}\)
Combined Gas Law – Pressure, Temperature, and Volume
- This combines the other three gas laws into a single ratio.
- \( \frac{PV}{T} \) = constant (when n is held constant)
- \( \frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2} \)
Avogadro’s Law – Moles vs. Volumes
- The number of moles of any gas is directly proportional to the volume of the sample. \( n \propto V\)
- \( \frac{n}{V}\) = constant (when P and T are held constant)
- \( \frac{n_1}{V_1} = \frac{n_2}{V_2} \)
- At STP, 1.00 mole of a gas occupies 22.4 Liters.
Ideal Gas Law
- The ideal gas law relates all of these different properties of a gas – pressure (P), volume (V), temperature (T in Kelvin), and number of moles (n) into a single equation that can be used to describe any ideal gas sample under any conditions.
- \( PV = nRT\)
- R is the ideal gas constant. There are different values for R, depending on what units the various properties are expressed in. Common R values are given at the bottom of the other side.
Standard Temperature and Pressure (STP)
- Standard Temperature = 0 °C = 273.15 K
- Standard Pressure = 1.000 atm = 760.0 mmHg = 101300 Pa = 101.3 kPa
Values of the Ideal Gas Constant (R)
- R = 0.08206 \( \frac{\text{L atm}}{\text{mol K}} \)
- R = 8314 \( \frac{\text{L Pa}}{\text{mol K}} \)
- R = 62.37 \( \frac{\text{L mmHg}}{\text{mol K}} \)
- R = 8.314 \( \frac{\text{L kPa}}{\text{mol K}} \) = 8.314 \( \frac{\text{J}}{\text{mol K}} \)
| Property | Conversion Factor |
|---|---|
| Mass |
2.205 lb = 1 kg 1 lb = 16 oz |
| Volume |
1.06 qt = 1 L 1 qt = 4 cups 1 fl oz = 29.6 mL 1 tsp = 5 mL |
| Length |
1 mi = 1.6 km 1 mi = 5280 ft 1 in = 2.54 cm 1 ft = 12 in. |
| Energy |
1 cal = 4.187 J 1 Calorie = 1000 calories |
Radiation Activity Units
- 1 becquerel (Bq) = 1 disintegration per second
- 1 Curie (Ci) = 3.7 x 1010 disintegration per second
- 1 Ci = 1000 mCi
- 1 Ci = 1,000,000 µCi
Solutions
- Concentration = \( \frac{\text{amount of solute}}{\text{amount of solution}} \)
- Molarity, M = \( \frac{\text{mole solute}}{\text{L solution}} \)
- % concentration = \( \frac{\text{parts of solute}}{\text{100 parts of solution}} \)
- Dilution Factor = \( \frac{V_f}{V_i} \)
- % (m/v) = \( \frac{\text{g solute}}{\text{mL solution}} \times 100 \)
- % (m/m) = \( \frac{\text{g solute}}{\text{g solution}} \times 100\)
- % (v/v) = \( \frac{\text{mL solute}}{\text{mL solution}} \times 100 \)
- \( C_{initial} \times V_{initial} = C_{final} \times V_{final}\)
- ppm = \( \frac{\text{g solute}}{\text{mL solution}} \times 1,000,000\)
- ppb = \( \frac{\text{g solute}}{\text{mL solution}} \times 1,000,000,000\)
Calculating pH and pKa
- pH = -log[H3O+]
- [H3O+] = 10-pH
- pKa = -log Ka
Supplemental Notes: Naming
In this exercise, you will practice naming and writing chemical formulas for many inorganic compounds, both ionic and molecular. Before beginning the exercise, you should carefully read all the sections of your text (or notes) on the names and formulas of ionic compounds, simple covalent compounds, and acids. The following is a brief summary of the Nomenclature rules for each of these types of compounds.
Ionic Compounds
- Composed of metal cations and non-metal anions, or, of polyatomic ions.
- Names and formulas always start with the positively charged cation.
- Ions are combined in ratios so that the final ionic compound is neutral.
- Never use prefixes in the names of ionic compounds. The cation name is simply combined with the anion name only.
- If the cation is capable of having more than one possible charge, the cation charge is included in the name as a Roman numeral in brackets (Stock system).
- Several ion names, charges and formulas are provided in the following tables. They must be memorized as soon as possible.
| Al3+ | Aluminum | Pb2+ | Lead(II); Plumbous |
| NH4+ | Ammonium | Pb4+ | Lead(IV); Plumbic |
| As3+ | Arsenic(III) | Li+ | Lithium |
| Ba2+ | Barium | Mg2+ | Magnesium |
| Cd2+ | Cadmium | Mn2+ | Manganese(II) |
| Ca2+ | Calcium | Mn4+ | Manganese(IV) |
| Cr2+ | Chromium(II) | Hg22+ | Mercury(I); Mercurous |
| Cr3+ | Chromium(III) | Hg2+ | Mercury(II); Mercuric |
| Cr6+ | Chromium(VI) | Ni2+ | Nickel(II); Nickelous |
| Co2+ | Cobalt(II); Cobaltous | K+ | Potassium |
| Co3+ | Cobalt(III); Cobaltic | Ag+ | Silver |
| Cu+ | Copper(I); Cuprous | Na+ | Sodium |
| Cu2+ | Copper(II); Cupric | Rb+ | Rubidium |
| Au3+ | Gold(III); Auric | Sr2+ | Strontium |
| H+ | Hydrogen | Sn2+ | Tin(II); Stannous |
| Fe2+ | Iron(II); Ferrous | Sn4+ | Tin(IV); Stannic |
| Fe3+ | Iron(III); Ferric | Zn2+ | Zinc |
| C2H3O2- | Acetate | I- | Iodide |
| AsO43- | Arsenate | IO4- | Periodate |
| BO33- | Borate | MoO42- | Molybdate |
| B4O72- | Tetraborate | N3- | Nitride |
| Br- | Bromide | NO2- | Nitrite |
| BrO- | Hypobromite | NO3- | Nitrate |
| BrO3- | Bromate | C2O42- | Oxalate |
| CO32- | Carbonate | O2- | Oxide |
| HCO3- | Bicarbonate; Hydrogen carbonate | O22- | Peroxide |
| Cl- | Chloride | MnO4- | Permanganate |
| ClO- | Hypochlorite | P3- | Phosphide |
| ClO2- | Chlorite | PO43- | Phosphate |
| ClO3- | Chlorate | HPO42- | Hydrogen phosphate |
| ClO4- | Perchlorate | H2PO42- | Dihydrogen phosphate |
| CrO42- | Chromate | Se2- | Selenide |
| Cr2O72- | Dichromate | S2- | Sulfide |
| C6H5O72- | Citrate | SO32- | Sulfite |
| CN- | Cyanide | SO42- | Sulfate |
| F- | Fluoride | HSO3- | Bisulfite; Hydrogen sulfite |
| H- | Hydride | HSO4- | Bisulfate; Hydrogen sulfate |
| OH- | Hydroxide | S2O32- | Thiosulfate |
| SCN- | Thiocyanate |
Examples
- K2S
- 2 K+ cations and 1 S2- anion
- potassium sulfide
- FeCl3
- 1 Fe3+ cation and 3 Clanions
- iron(III) chloride, or ferric chloride
- Mg3(PO4)2
- 3 Mg+2 cations and 2 PO4 3- anions
- magnesium phosphate
Simple Covalent (Molecular) Compounds
- Composed of non-metal atoms only.
- The more metallic non-metal is written first.
- Prefixes are used in the name to indicate the number of each atom present. A list of prefixes 1-10 (and 12) is provided below, which must be memorized.
- The prefix “mono” is dropped if there is only one of the first element.
- The name of the second element always ends in __ide.
| 1 | Mono |
| 2 | Di |
| 3 | Tri |
| 4 | Tetra |
| 5 | Penta |
| 6 | Hexa |
| 7 | Hepta |
| 8 | Octa |
| 9 | Nona |
| 10 | Deca |
| 12 | Dodeca |
Examples
- P4S3
- 4 P atoms and 3 S atoms
- tetraphosphorus trisulfide
- N2O
- 2 N atoms and 1 O atom
- dinitrogen monoxide
- BrCl5
- 1 Br atom and 5 Cl atoms
- bromine pentachloride
Acids
- Composed of hydrogen cations and non-metal anions or polyatomic anions.
- H always leads the formula.
- Acids are in the aqueous state.
- Ions are combined in ratios so that the final acid is neutral.
- The acid name depends on the name of the anion involved:
Examples
- HBr (aq)
- 1 H+1 cation and 1 Br-1 anion (bromide)
- hydrobromic acid
- HNO3 (aq)
- 1 H+1 cation and 1 NO3 -1 anion (nitrate)
- nitric acid
- H2SO3 (aq)
- 2 H+1 cations and 1 SO3 -2 anion (sulfite)
- sulfurous acid
Hydrates
- Hydrates are solid substances that also contain water molecules in their crystal structure.
- The number of water molecules for each formula unit is fixed.
- The formula is written the same as any other compound, but then we add a dot (•), (note that this is NOT a multiplication sign!) followed by the number of water molecules per ionic formula unit and the symbol H2O.
- The name is the same as for any other compound, but we add the word “hydrate” at the end with a Latin prefix to indicate the number of water molecules per ionic formula unit.
Examples
- CuSO4 • 5 H2O
- 1 Cu2+ ion, 1 SO4 -2 ion, and 5 water molecules
- Copper(II) sulfate pentahydrate
- Sodium carbonate decahydrate
- Na+ ion(s), CO3 -2 ions(s), and 10 water molecules
- Na2CO3 • 10 H2O