Final Review Chemistry
Law of Conservation of Mass
"In any chemical reaction the total mass of substances present after the reaction is the same as the total mass of substances before the reaction."
How many protons, neutrons and electrons are there in 35/17 Cl-
# protons = Z = 17 # neutrons = A - Z = 35 -17 = 18 # electrons = # protons - charge = 17 - (-1) = 18
Decaborane, B10H14, was used as a fuel for rockets in the 1950s. It reacts violently with oxygen, O2, to produce B2O3 and water. Calculate the percentage by mass of B10H14 in a fuel mixture designed to ensure that B10H14 and O2 run out at exactly the same time. (Such a mixture minimizes the mass of fuel that a rocket must carry.)
% by mass B10H14 = 25.8
An Ammonuim Dichromate "Volcano"
(NH4)2Cr2O7 ---> N2 + 4H2O + Cr2O3
Bromine
(a diatomic liquid)
Xenon
(a monoatomic gas)
Gold
(a precious metal)
Germanium
(a semiconductor element)
What are the molarities of the following solutes when dissolved in water? (a) 2.25 × 10−4 mol CH3CH2OH in 125 mL of solution (b) 57.5 g (CH3)2CO in 525 mL of solution (c) 18.5 mL of C3H5(OH)3 (d = 1.26 g/mL) in 375 mL of solution
(a) 0.00180M (b) 1.886M (c) 0.675M
What are the molarities of the following solutes? (a) aspartic acid (H2C4H5NO4) if 0.405 g is dissolved in enough water to make 100.0 mL of solution (b) acetone, (CH3)2CO, (d = 0.790 g/mL) if 35.0 mL is dissolved in enough water to make 425 mL of solution (c) diethyl ether, (C2H5)2O, if 8.8 mg is dissolved in enough water to make 3.00 L of solution
(a) 0.0304M (b) 1.12M (c) 4.0x10-5 M
Excess NaHCO3 is added to 525 mL of 0.220 M Cu(NO3)2. These substances react as follows: Cu(NO3)2(aq)+2NaOHCO3(s)→CuCO3(s)+2NaNO3(s)+H2O(l)+CO2(g) (a) How many grams of the NaHCO3(s) will be consumed? (b) How many grams of CuCO3(s) will be produced?
(a) 19.4g NaHCO3 (s) consumed (b) 14.3g CuCO3 (s) produced
Determine the mass, in grams, of (a) 7.34 mol NO2; (b) 4.220 × 10^25 O2 molecules; (c) 15.5 mol CuSO4 · 5 H2O; (d) 2.25 × 10^24 molecules of C2H4(OH)2.
(a) 338 g NO2 (b) 2242 g O2 (c) 3.87x10^3 g CuSO4 ⋅ 5H2O (d) 232 g C2H4(OH)2
Write balanced equations based on the information given. (a) solid magnesium + nitrogen gas → solid magnesium nitride (b) solid potassium chlorate → solid potassium chloride + oxygen gas (c) solid sodium hydroxide + solid ammonium chloride → solid sodium chloride + gaseous ammonia + water vapor (d) solid sodium + liquid water → aqueous sodium hydroxide + hydrogen gas
(a) 3Mg (s) + N2 (g) --> Mg3N2 (s) (b) 2KClO3 (s) --> 2KCl (s) + 3O2 (g) (c) NaOH (s) + NH4Cl (s) --> NaCl (s) + NH3 (g) + H2O (g) (d) 2Na (s) + 2H2O (l) --> 2NaOH (aq) + H2 (g)
For the mineral torbernite, Cu(UO2)2(PO4)2 · 8 H2O, determine (a) the total number of atoms in one formula unit (b) the ratio, by number, of H atoms to O atoms (c) the ratio, by mass, of Cu to P (d) the element present in the greatest mass percent (e) the mass required to contain 1.00 g P
(a) 41 total atoms (b) Ratio H:O = 16:20 or 4:5 (c) Ratio by mass Cu:P = 1.026 (d) U has greatest mass % = 50.77% (e) 15.1 g of torbernite required
All of the following minerals are semiprecious or precious stones. Determine the mass percent of the indicated element. (a) Zr in zircon, ZrSiO4 (b) Be in beryl (emerald), Be3Al2Si6O18 (c) Fe in almandine (garnet), Fe3Al2Si3O12 (d) S in lazurite (lapis lazuli), Na4SSi3Al3O12
(a) 49.765% Zr (b) 5.03004% Be (c) 33.658% Fe (d) 6.6622% S
Determine the mass, in grams, of (a) 2.10 × 10^&2 mol S8 (b) 5.02 × 10^22 molecules of palmitic acid, C16H32O2 (c) a quantity of the amino acid histidine, C6H9N3O2, containing 2.95 mol N atoms
(a) 5.39 g S (b) 21.4 g palmitic acid (c) 152.6 g C6H9N3O2
In many communities, water is fluoridated to prevent tooth decay. In the United States, for example, more than half of the population served by public water systems has access to water that is fluoridated at approximately 1 mg F− per liter. (a)What is the molarity of F− in water if it contains 1.2 mg F− per liter? (b) How many grams of solid KF should be added to a 1.6 × 108 L water reservoir to give a fluoride concentration of 1.2 mg F− per liter?
(a) 6.3x10-5 M (b) 5.9x105 g KF
For the compound Ge[S(CH2)4CH3]4, determine (a) the total number of atoms in one formula unit (b) the ratio, by number, of C atoms to H atoms (c) the ratio, by mass, of Ge to S (d) the number of g S in 1 mol of the compound (e) the number of C atoms in 33.10 g of the compound
(a) 69 atoms (b) Ratio C:H = 5:11 (c) Ratio by mass Ge:S = 0.566 (d) 28.264 g S (e) 8.212x10^23 C atoms
Para-cresol (p-cresol) is used as a disinfectant and in the manufacture of herbicides. A 0.4039 g sample of this carbon-hydrogen-oxygen compound yields 1.1518 g CO2 and 0.2694 g H2O in combustion analysis. Its molecular mass is 108.1 u. For p-cresol, determine its (a) mass percent composition; (b) empirical formula; (c) molecular formula.
(a) 77.83% C, 7.4641% H, 14.71% O (b) C7H8O (c) C7H8O
What reagent solution might you use to separate the cations in each of the following mixtures? [Hint: Refer to Exercise 23.] (a) PbSO4(s) and Cu(NO3)2(s) (b) Mg(OH)2(s) and BaSO4(s) (c) PbCO3(s) and CaCO3(s)
(a) Add H2O. Cu(NO3)2 will dissolve, while PbSO4 will not dissolve (b) Add HCl (aq). Mg(OH)2 will dissolve, but BaSO4 will not dissolve (c) Add HCl (aq). Both carbonates dissolve, but PbCl2 will precipitate while CaCl2 remains dissolved
Predict in each case whether a reaction is likely to occur. If so, write a net ionic equation. (a) AgNO3 (aq) + CuCl2 (aq) → (b) Na2S (aq) + FeCl2 (aq) → (c) Na2CO3 (aq) + AgNO3 (aq) →
(a) Ag+ (aq) + Cl- (aq) --> AgCl (s) (b) S2- (aq) + Fe2+ (aq) --> FeS (s) (c) CO32- (aq) + 2Ag+ (aq) --> Ag2CO3 (s)
Complete each of the following as a net ionic equation. If no reaction occurs, so state. (a) Ca2+ + 2I− + 2Na+ + CO32− → (b) Ba2+ + S2− + 2Na+ + SO42− → (c) 2K+ + S2− + Ca2+ + 2Cl− →
(a) Ca2+ (aq) + CO32- (aq) --> CaCO3 (s) (c) Ba2+ (aq) + SO42- (aq) --> BaSO4 (s) (c) no reaction occurs (no precipitate forms)
Supply the formula for the acids: (a) hydrofluoric acid; (b) nitric acid; (c) phosphorous acid; (d) sulfuric acid.
(a) HF(aq) (b) HNO3 (c) H3PO3 (d) H2SO4
Hydrogen and oxygen usually have oxidation states of +1 and −2, respectively, in their compounds. The following cases serve to remind us that there are exceptions. What are the oxidation states of the atoms in each of the following compounds? (a) MgH2; (b) CsO3; (c) HOF; (d) NaAlH4.
(a) Mg = +2 ; H = -1 (b) Cs = +1 ; O = -1/3 (c) H = +1 ; F = -1 ; O = 0 (d) Na = +1 ; H = -1 ; Al = +3
Write formulas for the compounds: (a) magnesium perchlorate; (b) lead(II) acetate; (c) tin(IV) oxide; (d) hydroiodic acid; (e) chlorous acid; (f) sodium hydrogen sulfite; (g) calcium dihydrogen phosphate; (h) aluminum phosphate; (i) dinitrogen tetroxide; (j) disulfur dichloride.
(a) Mg(ClO4)2 (b) Pb(C2H3O2)2 (c) SnO2 (d) HI (aq) (e) HClO2 (f) NaHSO3 (g) Ca(H2PO4)2 (h) AlPO4 (i) N2O4 (j) S2Cl2
Every antacid contains one or more ingredients capable of reacting with excess stomach acid (HCl). The essential neutralization products are CO2 and/or H2O. Write net ionic equations to represent the neutralizing action of the following popular antacids. (a) Alka-Seltzer (sodium bicarbonate) (b) Tums (calcium carbonate) (c) milk of magnesia (magnesium hydroxide) (d) Maalox (magnesium hydroxide, aluminum hydroxide) (e) Rolaids [NaAl(OH)2CO3]
(a) NaHCO3 (s) + H+ (aq) --> Na+ (aq) + H2O (l) + CO2 (g) (b) CaCO3 (s) + 2H+ (aq) --> Ca2+ (aq) + H2O (l) + CO2 (g) (c) Mg(OH)2 (s) + 2H+ (aq) --> Mg2+ (aq) + 2H2O (l) (d) Mg(OH)2 (s) + 2H+ (aq) --> Mg2+ (aq) + 2H2O (l)Al(OH)3 (s) + 3H+ (aq) --> Al3+ (aq) + 3H2O (l) (e) NaAl(OH)2CO3 (s) + 4H+ (aq) --> Al3+ (aq) + Na+ (aq) + 3H2O (l) + CO2 (g)
Balance the following equations by inspection. (a) SO2Cl2 + HI → H2S + H2O + HCl + I2 (b) FeTiO3 + H2SO4 + H2O → FeSO4⋅7 H2O + TiOSO4 (c) Fe3O4 + HCl + Cl2 → FeCl3 + H2O + O2 (d) C6H5CH2SSCH2C6H5 + O2 → CO2 + SO2 + H2O
(a) SO2Cl2 + 8HI --> H2S + 2H2O + 2HCl + 4I2 (b) FeTiO3 + 2H2SO4 + 5H2O --> FeSO4⋅7H2O + TiOSO4 (c) 2Fe3O4 + 12HCl + 3Cl2 --> 6FeCl3 + 6H2O + O2 (d) 2 C6H5CH2SSCH2C6H5 + 39O2 --> 28 CO2 + 4 SO2 + 14 H2O
Name these compounds: (a) Ba(NO3)2; (b) HNO2; (c) CrO2; (d) KIO3; (e) LiCN; (f) KIO; (g) Fe(OH)2; (h) Ca(H2PO4)2; (i) H3PO4; (j) NaHSO4; (k) Na2Cr2O7; (l) NH4C2H3O2; (m) MgC2O4; (n) Na2C2O4.
(a) barium nitrate (b) nitrous acid (c) chromium(IV) oxide (d) potassium iodate (e) lithium cyanide (f) potassium hypoiodite (g) iron(II) hydroxide (h) calcium dihydrogen phosphate (i) phosphoric acid (j) sodium hydrogen sulfate (k) sodium dichromate (l) ammonium acetate (m) magnesium oxalate (n) sodium oxalate
Which solution has the greatest [SO42−]? (a) 0.075 M H2SO4; (b) 0.22 M MgSO4; (c) 0.15 M Na2SO4; (d) 0.080 M Al2 (SO4)3; (e) 0.20 M CuSO4.
(d) has the greatest concentration of sulfate ions at 0.24M
Extensive Property
(depends on the amount of a substance) Length Volume Mass Heat Capacity
Intensive Property
(does not depend on the amount of a substance) Density Temperature Color Melting and boiling points
Arrange the following species in order of increasing (a) number of electrons; (b) number of neutrons; (c) mass. 112/50 Sn 40/18Ar 122/52 Te 59/29 Cu 120/48 Cd 58/27 Co 39/19 K
***Canvas messed up the formatting of the superscripts and subscripts, so only the element symbol is listed here. Please refer to your book for details*** A. Ar < K < Co < Cu < Cd < Sn < Te B. K < Ar < Cu < Co < Sn < Te < Cd C. K < Ar < Co < Cu < Sn < Cd < Te
Isotopic Masses/Mass Spectrometry•
- A beam of gaseous ions pass through a magnetic/electric field which separates ions into components with different masses. - Each mass which is present in the sample will show up on a photographic plate or another detector as lines. - You can relate the path traveled in the instrument to the mass of the ion and the intensity of the line to the amount of that particular mass which was present. - The sample shown below was for the mercury ion. You can clearly see the different isotopes of mercury on the figure to the right.
Chemical reactions are often accompanied by any of the following:
- A color change - Formation of a solid - Evolution of a gas - Evolution or absorption of heat or light
Metathesis (double displacement)
- AB + CD --> AD + CB - Acid - BAse - Precipitation
Sodium Chloride
- An extended array of Na+ and Cl- ions - The simplest formula unit is NaCl
Isotopes
- An isotope of an element is an atom which contains a different number of neutrons than another atom of the same element - Consequently the mass of each atom of the same element does not need to be exactly the same
Thomson's experiments led him to propose that:
- Cathode rays are made of particles not waves - Since the cathode rays do not depend on the composition of the cathode they are the fundamental negatively charged particles found in all atoms. - Cathode rays were then termed electrons in 1874. - The mass of electrons are much less than protons
Solubility Rules
- Compounds that are soluble• - Containing group 1 metals and the ammonium ion - Nitrates, perchlorates, and acetates - Compounds that are mostly soluble - Chlorides, bromides, and iodides, except those of Pb(II), Ag(I), and Hg(I) which are insoluble - Sulfates, except those of Ca(II), Sr(II), Ba(II), Pb(II), and Hg(I) - Compounds that are mostly insoluble - Hydroxides and sulfides (except for those of group 1 and ammonium. Sulfides of group 2 metals are soluble and the hydroxides of calcium, strontium and barium are slightly soluble) - Carbonates and phosphates (except for group 1 and ammonium) - Salts containing Pb(II), Ag(I), and Hg(I)
Dalton's Atomic Theory
- Dalton used the two laws that we just described to formulate his atomic theory. His theory involved 3 assumptions: - Each chemical element is composed of minute, indestructible particles called atoms. Atoms can neither be destroyed nor created during a chemical reaction. - All atoms of an element are alike in mass (weight) and other properties but the atoms of one element are different from those of all the other elements. - In each of their compounds, different elements combine in a simple numerical ratio: for example, one atom of A combines with one atom of B to form AB. Or one atom of A combines with 2 atoms of B to form AB2
Charge and Magnetic Interactions
- Electric charge is assigned either a positive or a negative sign - Positive and negative charges attract each other - Like charges (+ and + or - and -) repel each other
Ionic Compounds
- Electron is transferred from the metal to the nonmetal and ions result - Positive ion is called a cation - Negative ion is called an anion• Made up of positive and negative ions joined together by Coulombic attraction - Usually from a metal and a nonmetal. - The compound is best described using the empirical formula.
Thomson's "Plum Pudding" Model
- Electrons floated in a nebulous cloud of positive charge- - Much like Jello® with fruit embedded in it or a popular English dessert known as Plum Pudding.
Solutions
- Homogenous mixtures of solvent and solute - Liquid, solid or gas phase
molecular compound.
- Made up of molecules which typically consist of a small number of non-metal atoms held together by covalent bonds - Hydrazine is an example of a molecular compound
Activity Series of Nonmetals
- Most active (most strongly oxidizing) nonmetals appear on top, and least active nonmetals appear on the bottom. - For example, the series predicts that Cl2 will displace Br- and I- from solution, because Cl2 appears above Br2 and I2
Joseph Proust: Constant Composition
- Proust's experiments with copper carbonate demonstrated the law of constant composition or the law of definite proportions - In effect, all molecules of water are 11.19% hydrogen by mass and 88.81% oxygen by mass.
Avogadro's Number (Na) and the Mole
- The mole is the amount of a substance that contains the same number of entities (often atoms of molecules) as there are carbon-12 atoms in exactly 12 g of carbon-12. It turns out there are 6.02214199 x 1023 atoms/mol of a substance and this number is called Avogadro's constant. - A mole then is just basically a counting unit for how many of an item is present.
The Atomic Model Begins to Take Shape
- The nucleus is at the center of the atom and only takes up a small fraction of the total volume. - The electrons are found surrounding the nucleus - If the nucleus was the size of a penny, the atom would be the size of two soccer fields.
Dilutions
- a small volume of the liquid chemical is diluted in a larger volume of solvent to achieve a certain ratio - M1V1 = M2V2
Molecular mass (weight)
- the mass of a molecular compound (weighted average of naturally occurring isotopes) (u) - Usually MW = n x FW : n = integer
Molarity
- the number of moles of solute per liter of solution - Thus, molarity is a measure of the concentration of the solution. - Units for molarity is moles/liter or often expressed (M)
Combustion
- the process of burning something - reaction in which a substance reacts with oxygen gas, releasing energy in the form of light and heat.
-ate
-ic acid
-ite
-ous acid
In the year 2000, the Guinness Book of World Records called ethyl mercaptan, C2H6S, the smelliest substance known. The average person can detect its presence in air at levels as low as 9 × 10−4 μmol/m3. Express the limit of detectability of ethyl mercaptan in parts per billion (ppb). (Note: 1 ppb C2H6S means there is 1 g C2H6S per billion grams of air.) The density of air is approximately 1.2 g/L at room temperature.
0.05 ppb
A 10.00 mL sample of 2.05 M KNO3 is diluted to a volume of 250.0 mL. What is the concentration of the diluted solution?
0.0820M
The concentration of Mn2+ (aq) can be determined by titration with MnO4− (aq) in basic solution. A 50.00 mL sample of Mn2+(aq) requires 78.42 mL of 0.04997 M KMnO4 for its titration. What is [Mn2+ in the sample? Mn2+ + MnO4- ---> MnO2 (s) (not balanced)
0.1176M Mn2+
If 18.2 mL H2O evaporates from 1.00 L of a solution containing 15.5 mg K2SO4/mL, what is K+] in the solution that remains?
0.181M
What volume of 0.0665 M KMnO4 is necessary to convert 12.5 g KI to I2 in the reaction below? Assume that H2SO4 is present in excess. 2KMnO4+10KI+8H2SO4→6K2SO4+2MnSO4+5I2+8H2O
0.226L of 0.0665 M KMnO4
Assuming the volumes are additive, what is the [NO3−] in a solution obtained by mixing 275 mL of 0.283 M KNO3, 328 mL of 0.421 M Mg(NO3)2, and 784 mL of H2O?
0.255M
A 0.3126 g sample of oxalic acid, H2C2O4, requires 26.21 mL of a particular concentration of NaOH(aq) to complete the following reaction. What is the molarity of the NaOH(aq)? H2C2O4(s)+2NaOH(aq)→Na2C2O4(aq)+2H2O(l)
0.2649 M
A 0.406 g sample of magnesium reacts with oxygen, producing 0.674 g of magnesium oxide as the only product. What mass of oxygen was consumed in the reaction?
0.268 g oxygen
In normal blood, there are about 5.4 × 10^9 red blood cells per milliliter. The volume of a red blood cell is about 90.0 × 10^−12 cm^3, and its density is 1.096 g/mL. How many liters of whole blood would be needed to collect 0.5 kg of red blood cells?
0.9 L of blood or 9x102 mL of blood
Meth
1 carbon
The Greek Concept of Atomos: The Indivisible Atom
1) All matter is composed of atoms, which are bits of matter too small to be seen. These atoms CANNOT be further split into smaller portions. 2) There is a void, which is empty space between atoms. "Unless there is a void with a separate being of its own, 'what is' cannot be moved-nor again can it be 'many', since there is nothing to keep things apart." 3) Atoms are completely solid. "there can be no void inside an atom itself. Otherwise an atom would be subject to changes from outside and could disintegrate ." 4) Atoms are homogeneous, with no internal structure. 5) Atoms are different in their sizes, shapes, and weight.
Rules for Oxidation States
1. Any element/diatomic by itself = 0 2. Group 1 elements = +1 3. Group 2 elements = +2 4. Halogens = usually -1, +1 when combined with oxygen 5. Hydrogen with nonmetals = +1 6. Hydrogen with metals = -1 7. Oxygen = usually -2, -1 in H2O2 8. Fluorine = -1 9. Any ion with a charge = that charge 10. Any neutral compound with no charge = all atoms present's oxidation states must add up to 0 11. Any polyatomic ion with a charge = all atoms present's oxidation states must add up to that charge
Balancing Redox Reactions
1. Separate the two half reactions 2. Balance the atoms of each half reaction(may have to add H2O or H) 3.Balance the electrons of each half reaction 4.Combine the half reactions 5.Confirm mass and charge are equal
What is chemistry?
1. the science that systematically studies the composition, properties, and activity of organic and inorganic substances and various elementary forms of matter. 2. chemical properties, reactions, phenomena, etc.: the chemistry of carbon.
Calculate the mass of a block of iron (d = 7.86 g/cm^3) with dimensions of 52.8 cm × 6.74 cm × 3.73 cm.
1.04x104 g iron
A KMnO4(aq) solution is to be standardized by titration against As2O3(s). A 0.1078 g sample of As2O3 requires 22.15 mL of the KMnO4(aq) for its titration. What is the molarity of the KMnO4(aq)? 5As2O3 + 4MnO4- + 9H2O + 12H+ ---> 10H3AsO4 + 4Mn2+
1.968 x 10-2 M
Dec
10 carbons
The two naturally occurring isotopes of silver have the following abundances: 107Ag, 51.84%; 109Ag, 48.16%. The mass of 107Ag is 106.905092 u. What is the mass of 109Ag?
109Ag = 108.9 u
What volume of 0.0962 M NaOH is required to exactly neutralize 10.00 mL of 0.128 M HCl?
13.3 mL NaOH (aq) solution
How many kilograms of HNO3 are consumed to produce 125 kg Ca(H2PO4)2 in this reaction? Ca3(PO4)2+HNO3→;Ca(H2PO4)2+Ca(NO3)2 (not balanced)
135 kg HNO3
Eth
2 carbons
Alpha particles
2 fundamental units of positive charge and have essentially the same mass as helium atoms
How many Cu atoms are present in a piece of sterling-silver jewelry weighing 33.24 g? (Sterling silver is a silver-copper alloy containing 92.5% Ag by mass.)
2.4x1022 atoms Cu
Use the conventional atomic mass of boron to estimate the fractional isotopic abundances of the two naturally occurring isotopes, 10B and 11B. These isotopes have masses of 10.012937 u and 11.009305 u, respectively.
20% 10B and 80% 11B
Anhydrous CuSO4 can be used to dry liquids in which it is insoluble. The CuSO4 is converted to CuSO4 · 5 H2O, which can be filtered off from the liquid. What is the minimum mass of anhydrous CuSO4 needed to remove 12.6 g H2O from a tankful of gasoline?
22.3 g CuSO4
Lithopone is a brilliant white pigment used in water-based interior paints. It is a mixture of BaSO4 and ZnS produced by the reaction BaS(aq)+ZnSO4(aq)→ZnS(s)+BaSO4(s)lithopone How many grams of lithopone are produced in the reaction of 315 mL of 0.275 M ZnSO4 and 285 mL of 0.315 M BaS?
28.7g lithopone
Prop
3 carbons
A 1.562 g sample of the alcohol CH3CHOHCH2CH3 is burned in an excess of oxygen. What masses of CO2 and H2O should be obtained?
3.710 g CO2, 1.898 g H2O
But
4 carbons
Determine the only possible +2 ion for which the following two conditions are both satisfied: - The net ionic charge is one-tenth the nuclear charge. - The number of neutrons is four more than the number of electrons.
42 20 Ca2+ (Canvas can't format it correctly, but the 42 should be directly above the 20)
Pent
5 carbons
Household ammonia, used as a window cleaner and for other cleaning purposes, is NH3(aq). The NH3 present in a 5.00 mL sample is neutralized by 28.72 mL of 1.021 M HCl. The net ionic equation for the neutralization is NH3 (aq) + H+ (aq) ---> NH4+ (aq)
5.86M
Iodine-131 is a radioactive isotope that has important medical uses. Small doses of iodine-131 are used for treating hyperthyroidism (overactive thyroid) and larger doses are used for treating thyroid cancer. Iodine-131 is administered to patients in the form of sodium iodide capsules that contain 131 I− ions. Determine the number of neutrons, protons, and electrons in a single 131 I− ion.
53 protons, 54 electrons, 78 neutrons
Hex
6 carbons
How many atoms are present in a 1.50 m length of 20-gauge copper wire? A 20-gauge wire has a diameter of 0.03196 in., and the density of copper is 8.92 g/cm^3.
6.56x1022 atoms
The two naturally occurring isotopes of lithium , lithium 6 and lithium 7, have masses of 6.01513 and 7.01601 u respectively. Use the periodic table atomic mass of 6.941 u for lithium to determine the percent natural abundances for each of these isotopes.
6Li = 7.49% 7Li = 92.51%
Hept
7 carbons
What is the mass of a cube of osmium that is 1.25 inches on each side? The density of osmium is 22.48g/cm3.
719 g osmium
When a solid mixture of MgCO3 and CaCO3 is heated strongly, carbon dioxide gas is given off and a solid mixture of MgO and CaO is obtained. If a 24.00 g sample of a mixture of MgCO3 and CaCO3 produces 12.00 g CO2, then what is the percentage by mass of MgCO3 in the original mixture?
73.33%
Dichlorodifluoromethane, once widely used as a refrigerant, can be prepared by the reactions shown. How many moles of Cl2 must be consumed in the first reaction to produce 2.25 kg CCl2F2 in the second? Assume that all the CCl4produced in the first reaction is consumed in the second. CH4+Cl2→CCl4+HCl (not balanced)CCl4+HF→CCl2F2+HCl (not balanced)
74.4 mol Cl2
How many grams of commercial acetic acid (97% CH3COOH by mass) must be allowed to react with an excess of PCl3 to produce 75 g of acetyl chloride (CH3COCl), if the reaction has a 78.2% yield? CH3COOH+PCl3→ CH3COCl+H3PO3(not balanced)
76g commercial acetic acid
Iron metal reacts with chlorine gas. How many grams of FeCl3 are obtained when 515 g Cl2 reacts with excess Fe? 2Fe(s)+3Cl2(g)→2FeCl3(s)
785 g FeCl3
Oct
8 carbons
Solid silver oxide, Ag2O(s), decomposes at temperatures in excess of 300 °C, yielding metallic silver and oxygen gas. A 3.13 g sample of impure silver oxide yields 0.187 g O2g. What is the mass percent Ag2O in the sample? Assume that Ag2O(s) is the only source of O2g. [Hint: Write a balanced equation for the reaction.]
86.6% Ag2O
What volume of 2.00 M AgNO3 must be diluted with water to prepare 500.0 mL of 0.350 M AgNO3?
87.5mL of 2.00 M AgNO3
Cryolite, Na3AlF6, is an important industrial reagent. It is made by the reaction below. Al2O3(s)+6NaOH(aq)+12HF(g)→2Na3AlF6(s)+9H2O(l) In an experiment, 7.81 g Al2O3 and excess HF(g) were dissolved in 3.50 L of 0.141 M NaOH. If 28.2 g Na3AlF6 was obtained, then what is the percent yield for this experiment?
87.6%
Non
9 carbons
100. ml of a HCl solution is titrated with 0.125 M NaOH solution. At the equivilance point, 7.50 ml of the NaOH solution is used. What is the concentration (molarity) of the HCl solution?
9.44 x 10^-3 M
What is the total number of atoms in (a) 15.8 mol Fe; (b) 0.000467 mol Ag; (c) 8.5 × 10^−11 mol Na?
A 9.51x1024 atoms Fe B. 2.81x1020 atoms Ag C. 5.1x1013 atoms Na
Chemical Reactions
A chemical reaction is a process in which one or more substances is converted into one or more new substances.
Silver bromide (AgBr)
A photosensitive solid
Filtration
A process that separates materials based on the size of their particles.
Distillation
A process that separates the substances in a solution based on their boiling points
Neutralization
A reaction of an acid with a base, yielding a solution that is not as acidic or basic as the starting solutions were.
The scientific method
A series of steps followed to solve problems including collecting data, formulating a hypothesis, testing the hypothesis, and stating conclusions.
Base
A substance that provides hydroxide ions, OH-(aq), in aqueous solutions
Chromatography
A technique that is used to separate the components of a mixture based on the tendency of each component to travel or be drawn across the surface of another material.
Combination/Synthesis
A+B-->AB
Perform the following calculations; express each number and the answer in exponential form and with the appropriate number of significant figures. (a) (320×24.9)/0.080 = (b) (432.7×6.5 × 0.002300)/(62×0.103) = (c) (32.44+4.9−0.304)/82.94 = (d) (8.002+0.3040)/(13.4−0.066+1.02) =
A. 1.0x105 B. 1.0 C. 4.47x10-1 D. 5.79x10-1
In one experiment, the reaction of 1.00 g mercury and an excess of sulfur yielded 1.16 g of a sulfide of mercury as the sole product. In a second experiment, the same sulfide was produced in the reaction of 1.50 g mercury and 1.00 g sulfur. (a) What mass of the sulfide of mercury was produced in the second experiment? (b) What mass of which element (mercury or sulfur) remained unreacted in the second experiment?
A. 1.74 g of the sulfide of mercury B. 0.76 g sulfur unreacted
Perform the following conversions from non-SI to SI units. (Use information from the inside back cover, as needed.) (a) 68.4 in = cm (b) 94 ft = m (c) 1.42 lb = g (d) 248 lb = kg (e) 1.85 gal = dm^3 (f) 3.72 qt = mL
A. 174 cm B. 29 m C. 644 g D. 112 kg E. 7.00 dm3 F. 3.52x103 mL
Perform the following conversions. (a) 2.35 kg = g (b) 792 g = kg (c) 3869 mm = cm (d) 0.043 cm = mm
A. 2.35x103 g B. 0.792 kg C. 386.9 cm D. 0.43 mm
A non-SI unit of mass used in pharmaceutical work is the grain (gr) (15 gr = 1.0 g). An aspirin tablet contains 5.0 gr of aspirin. A 161 lb arthritic individual takes two aspirin tablets per day. (a) What is the quantity of aspirin in two tablets, expressed in milligrams? (b) What is the dosage rate of aspirin, expressed in milligrams of aspirin per kilogram of body mass? (c) At the given rate of consumption of aspirin tablets, how many days would it take to consume 1.0 kg of aspirin?
A. 6.7x102 mg B. 9.2 mg aspirin/kg body weight C. 1.5x103 days
Express the result of each of the following calculations in exponential form and with the appropriate number of significant figures. (a) (4.65 × 10^4) × (2.95 × 10^-2) × (6.663 × 10^-3) × 8.2 = (b) 1912 × (0.0077 × 10^4) × (3.12 × 10^−3) ------------------------------------------- = (4.18 × 10^−4)^3 (c) (3.46 × 10^3) × 0.087 × 15.26 × 1.0023 = (d) (4.505 × 10^−2)^2 × 1.080 × 1545.9 ------------------------------------------ = 0.03203 × 10^3 (e) (−3.61 × 10^−4) + square root of (3.61 × 10^−4)^2 + 4 (1.00) (1.9 × 10^−5) --------------------------------------------------------------------------- = 2 × (1.00)
A. 7.5x101 B. 6.3x1012 C. 4.6x103 D. 1.058x10-1 E. 4.2x10-3
Refer to the periodic table inside the front cover and identify (a) the element that is in group 14 and the fourth period (b) one element similar to and one unlike sulfur (c) the alkali metal in the fifth period (d) the halogen element in the sixth period
A. Ge B. Similar to S: O, Se, Te. Unlike S: most other elements, especially metals such as Na, K, Rb C. Rb D. At
How many significant figures are shown in each of the following? If this is indeterminate, explain why. (a) 450; (b) 98.6; (c) 0.0033; (d) 902.10; (e)0.02173; (f) 7000; (g) 7.02; (h) 67,000,000
A. indeterminate, 2 or 3 sig figs B. 3 sig figs C. 2 sig figs D. 5 sig figs E. 4 sig figs F. indeterminate, 1-4 sig figs G. 3 sig figs H. indeterminate, 2-8 sig figs
Hydrogen and chlorine atoms react to form simple diatomic molecules in a 1:1 ratio, that is, HCl. The percent isotopic abundances of the chlorine isotopes are 35Cl and 37Cl are estimated to be 75.77% and 24.23%, respectively. The percent isotopic abundances of 2H and 3H are estimated to be 0.015% and less than 0.001%, respectively. (a) How many different HCl molecules are possible, and what are their mass numbers (that is, the sum of the mass numbers of the H and Cl atoms)? (b) Which is the most abundant of the possible HCl molecules? Which is the second most abundant?
A. six different HCl molecules are possible: 1H35Cl mass = 36, 2H35Cl mass = 37, 3H35Cl mass = 38, 1H37Cl mass = 38, 2H37Cl mass = 39, 3H37Cl mass = 40 B. most abundant is 1H35Cl. second most abundant is 1H37Cl.
Decomposition
AB->A+B
Oxidizing agent
Accepts electrons and becomes reduced.
C2H3O2-
Acetate
NH4+
Ammonium
Ions
An atom of an element can also gain or lose electrons from the electron cloud causing the ion to have a net negative or positive charge. This is called an ion. The number of protons, and consequently the composition of the element does not change when an ion is formed.
Xenon difluoride (XeF2)
An unstable solid
Matter
Anything that has mass and takes up space
Binary Acids
Are acids that consist of two elements, usually hydrogen, and one of the halogens.
Solid
Atoms or molecules in close contact, often arranged in a complex structure known as a crystal. Solids occupy a definite shape.
A piece of gold (Au) foil measuring 0.25 mm × 15 mm × 15 mm is treated with fluorine gas. The treatment converts all the gold in the foil to 1.400 g of a gold fluoride. What is the formula and name of the fluoride? The density of gold is 19.3 g/cm3.
AuF3, gold(III) fluoride
Barium oxide
BaO
CsCN
Cesium Cyanide
Why study chemistry?
Chemistry can be useful in explaining the natural world, preparing people for career opportunities, and producing informed citizens.
ClO3-
Chlorate
ClO2-
Chlorite
Inorganic compounds
Compounds that do not contain carbon
Isomers
Compounds with the same formula but different structures.
CN-
Cyanide
Some Allotropes of Carbon
Diamond, graphite, buckminsterfullerene
Reducing agent
Donates electrons and becomes oxidized.
The following observations were made for a series of five oil drops in an experiment similar to Millikan's (see Figure 2-8). Drop 1 carried a charge of 1.28 × 10−18 C drops 2 and 3 each carried 1/2 the charge of drop 1; drop 4 carried 1/8 the charge of drop 1; drop 5 had a charge four times that of drop 1. Are these data consistent with the value of the electronic charge given in the text? Could Millikan have inferred the charge on the electron from this particular series of data? Explain.
Drop 1: 1.28x10-18 C = 8e Drops 2&3: 6.40x10-19 C = 4e Drop 4: 1.60x10-19 C = 1e Drop 5: 5.12x10-18 C = 32e These values are consistent with the charge that Millikan found for that of the electron, and he could have inferred the correct charge from this data since they are all multiples of e.
Electrical Conductivity in Water
Electrical current is supported by ions in solution - the more ions, the greater the current
Organic compounds abound in nature
Fats, carbohydrates and proteins are foods. Propane, gasoline, kerosene, oil. Drugs and plastics
Iron (III) sulfide
Fe2S3
Liquid
Flow and assume the shape of the container.
Molar mass
Formula or molecular mass (weighted average of naturally occurring isotopes) in grams/mole
molecular compounds
H2O, CH4, NH3
Common strong acids
HCl, HBr, HI, H₂SO₄, HNO₃, HClO₃, HClO₄
A 0.696 mol sample of Cu is added to 136 mL of 6.0 M HNO3. Assuming the following reaction is the only one that occurs, will the Cu react completely? 3Cu(s)+8HNO3(aq)→3Cu(NO3)2(aq)+4H2(l)+2NO(g)
HNO3 (aq) is the limiting reactant, it will be completely consumed, leaving some Cu unreacted
OH-
Hydroxide
Millikan's Experiment
In his experiment, ions are produced by ionizing radiation such as X-rays. The ions which are produced can "attach" to oil droplets which Millikan produced through a device much like a perfume sprayer. If ions attach to a droplet, they will become electrically charged. When the droplets are then placed in the electric field between the plates, the rate of fall of the droplets will change depending on the magnitude and sign of the charges on the droplets. Millikan analyzed data for a very large number of droplets and concluded that the magnitude of the charge on a droplet is -1.6022 x 10-19 C.
Salts
Ionic compounds that can be formed by replacing one or more of the hydrogen ions of an acid with another positive ion
potassium carbonate
K2CO3
How do the Pieces Fit?•
Matter cannot be composed of just negative particles! There must be positive charges to offset the electrons negative charges.
CH3OH
Methanol
Molecules
Molecules are discrete forms of matter consisting of two or more atoms bound by covalent bonds. Molecules are new forms of matter with unique structural and physical properties.
Elemental Allotropes
Molecules of the same elemental composition but with different masses and connectivity. - O2 (oxygen) and O3 (ozone) - P4 (white phosphorus) and P∞ (red phosphorus) - Sulfur rings: S5, S6, S7, S8, ... - C60, C70, diamond, graphite
Reactions in Solution: Displacement
More reactive metals will displace less reactive metals from compounds
There are many diatomic elements
N2, O2, F2, Cl2, Br2, I2, .......
NH3(aq) conducts electric current only weakly. The same is true for CH3COOH(aq). When these solutions are mixed, however, the resulting solution is a good conductor. How do you explain this?
NH3 (aq) is a weak base, CH3COOH (aq) is a weak acid. The reaction produces a solution of NH4CH3COO (aq) which is a salt and a strong electrolyte
Most soluble salts
NaCl, KF, Cu(NO3)2, etc
Common strong bases
NaOH, KOH, Ca(OH)2, Ba(OH)2
There are only a few monoatomic elements:
Ne, Ar, Kr, Xe, Ar, Hg
The electrolyte in a lead storage battery must have a concentration between 4.8 and 5.3 M H2SO4 if the battery is to be most effective. A 5.00 mL sample of a battery acid requires 49.74 mL of 0.935 M NaOH for its complete reaction (neutralization). Does the concentration of the battery acid fall within the desired range? [Hint: Keep in mind that the H2SO4 produces two H+ ions per formula unit.]
No, the concentration is 4.65M
Molar Mass
Numerically the same as molecular mass but expressed in g/mol
Elemental molecules
O2, N2, P4, S8
Many non-metallic elements are polyatomic:
P4, S8, C60, .......
ClO4-
Perchlorate
Allotropes of Phosphorus
Red, white, violet, and black
Dibutyl succinate is an insect repellent used against household ants and roaches. Its composition is 62.58% C, 9.63% H and 27.79% O. Its experimentally determined molecular mass is 230 g/mol. What are the empirical and molecular formulas of dibutyl succinate?
Step 1: Determine the mass of each element in a 100g sample. C 62.58 g H 9.63 g O 27.79 g Step 2: Convert masses to amounts in moles. Step 3: Write a tentative formula. Step 4: Divide formula subscripts by smallest number Step 5: Convert to a small whole number ratio. Step 6: Determine the molecular formula. Empirical formula mass is 115 g/mol. Experimental molecular mass is 230 g/mol.
Cathode Rays
Streams of electrons that are produced when a high voltage is applied to electrodes in an evacuated tube
Atomic Mass
The average mass of all the isotopes of an element
Temperature Scales
The temperature scale is determined by the melting (0o C) and boiling point of water (100o C). The interval between these points is divided up into 100 equal increments each called one Celsius degree
S2O32-
Thiosulfate
Rutherford's Experiment
This led Rutherford to propose that atoms: - Have most of the mass and all of their positive charge centered in a very small region called the nucleus. - The atom is mostly empty space - The magnitude of the positive charge is different for different atoms and the number of positive charges is often approximately one-half the atomic weight. - There are as many electrons outside the nucleus as there are units of positive charge in the nucleus to maintain electrical neutrality. - The positively charged particles were called protons.
Exact Mass
Use the most abundant isotopes (12C61H1216O6)
Molecular Mass
Use the naturally occurring mixture of isotopes
Ionic Solutes
When many ionic compounds are exposed to water they break up into ions in a process called hydration
Mixture
When pure substances (i.e. elements and compounds) are mixed (but not reacted)
Observing Atoms with STM
With scanning tunneling microscopy (STM), single atoms can be observed and even manipulated!!
The element X forms the compound XOCl2 containing 59.6% Cl. What is element X?
X is the element Sulfur
physical change
a change in which the physical properties of a substance may change, but the composition of the substance remains the same.
Analytical Balance
a device for determining mass that gives highly accurate measurements
Substance
a form of matter in which the atoms are combined in constant fixed proportions and have distinct properties. Pure by definition.
ductility
a measure of the substances mass per unit volume, usually expressed in grams per milliliter of volume.
Combustion Analysis
a method of obtaining empirical formulas for unknown compounds, especially those containing carbon and hydrogen, by burning a sample of the compound in pure oxygen and analyzing the products of the combustion reaction
Maldonite (AuBi2)
a mineral containing gold and bismuth
heterogeneous mixture:
a mixture in which the components are separated into distinct regions with differing physical properties. (Examples: a salad, bowl of soup, smog, mud)
homogeneous mixture
a mixture whose properties and composition is uniform throughout. (Examples: salt in water, air, alloys)
chemical property
a property describing a particular substance's reactivity. In effect, the substance's ability or inability to react under a certain condition.
Physical properties
a property that a sample of matter displays without a change in its composition. (examples include: malleability, ductility, density, state of matter)
Precipitation reaction
a reaction in which an insoluble substance forms and separates from the solution
Compound
a substance in which two or more elements are combined with one another in a specific ratio.
Element
a substance made up of only one type of atom. Each square on the periodic table represents one element
Acid
a substance which provides hydrogen ions, H+(aq), in aqueous solutions
Oxidation-Reduction (single displacement)
a type of chemical reaction that involves a transfer of electrons between two species.
Assigning Oxidation States What is the oxidation state of the underlined element in each of the following? a) P4; b) Al2O3; c) MnO4-; d) NaH.
a) P4 is an element. P OS = 0 b) Al2O3: O is -2. O3 is -6. Since (+6)/2=(+3), Al OS = +3. c) MnO4-: net OS = -1, O4 is -8. Mn OS = +7. d) NaH: net OS = 0, rule 3 beats rule 5, Na OS = +1 and H OS = -1.
Percent Yield
actual yield/theoretical yield x 100
ROH
alcohol
RCH=O
aldehyde
R2C=CR2
alkene
RC≡CR
alkyne
AlF3
aluminum fluoride
RNH2
amine
NH4Cl
ammonium chloride
Net ionic equation
an equation for a reaction in solution showing only those particles that are directly involved in the chemical change
Ionic Equation
an equation in which ions are explicitly shown
Oxyanions
anions derived of a non-metal bonded to one or more oxygen atoms
Limiting Reagent
any reactant that is used up first in a chemical reaction; it determines the amount of product that can be formed in the reaction
Ba(OH)2
barium hydroxide
CO2
carbon dioxide
CO
carbon monoxide
Organic compounds
carbon-based molecules
CO32-
carbonate ion
A 15.2 L sample of chloroform at 20 °C has a mass of 22.54 kg. What is the density of chloroform at 20 °C, in grams per milliliter?
chloroform density = 1.48 g/mL
HClO2
chlorous acid
HClO3
choric acid
Acids
compounds that form hydrogen ions when dissolved in water
Hydrates
compounds that have a specific number of water molecules attached to them
Bases
compounds that reduce the concentration of hydrogen ions in a solution.
Cu2O
copper (I) oxide
Cu(C2H3O2)2
copper (II) acetate
CuSO4•5 H2O
copper(II) sulfate pentahydrate
The molar mass
defined as the mass of one mole of atoms is numerically identical to the atomic mass and is expressed in g/mol.
A pycnometer (see Exercise 78) weighs 25.60 g empty and 35.55 g when filled with water at 20 °C. The density of water at 20 °C is 0.9982 g/mL. When 10.20 g of lead is placed in the pycnometer and the pycnometer is again filled with water at 20 °C, the total mass is 44.83 g. What is the density of the lead in grams per cubic centimeter?
density = 11 g/mL or 11 g/cm3 (because 1 mL = 1 cm3)
Mass
describes the quantity of matter in an object.
Titration Reactions
determine the concentration of an unknown species by chemical reaction
The volume of a red blood cell is about 90.0 × 10^−12 cm^3. Assuming that red blood cells are spherical, what is the diameter of a red blood cell in millimeters?
diameter = 5.56x10-3 mm
Cr2O72-
dichromate
H2PO4-
dihydrogen phosphate ion
N2O
dinitrogen monoxide
N2O5
dinitrogen pentoxide
N2O4
dinitrogen tetroxide
N2O3
dinitrogen trioxide
octa
eight
Covalent Bonding
electrons are shared between atoms. Usually formed between two nonmetals.
Ionic Bonding
electrons are transferred from one atom to another atom to form ions. Formed between a metal and a nonmetal.
H2C=CH2
ethylene
Gas
expands to fill container, will change volume readily with external force (pressure) -compressibility, large intermolecular distance compared to a liquid and solid.
penta
five
tetra
four
RX
haloalkaline
Beta particles
have the same properties as an electron
ate
higher oxidation state
Accuracy
how close the measured number is to the actual value
HCl
hydrochloric acid
HF
hydrofluoric acid
HCO3-
hydrogen carbonate (bicarbonate) ion
HPO42-
hydrogen phosphate ion
HSO4-
hydrogen sulfate (bisulfate)
HSO3-
hydrogen sulfite (bisulfite)
ClO-
hypochlorite
HClO
hypochlorous acid
Anhydrous sodium sulfate, Na2SO4, absorbs water vapor and is converted to the decahydrate, Na2SO4 · 10 H2O. How much would the mass of 36.15 g of anhydrous Na2SO4 increase if converted completely to the decahydrate?
increase in mass = 45.85 g
Triple Beam Balance
instrument used to measure mass
CH3I
iodemethane
Spectator ions
ions that do not participate in a reaction
R2C=O
ketone
ite
lower oxidation state
Balance
measure weight, but are precisely calibrated to account for earth's gravitational field and their readout is mass
Digital Analytical Balance
measures mass
HCHO
methanal
CH3NH2
methylamine
Electrolytes
minerals that help maintain the body's fluid balance
NO3-
nitrate ion
NO2-
nitrite ion
NO2
nitrogen dioxide
NO
nitrogen monoxide
Gamma rays
not particles but rather is electromagnetic radiation of extremely high penetrating power. Gamma rays are not affected by electric fields
mono
one
chemical change
one or more types of atoms or molecules are converted to a new substance with a different composition.
Composition
parts or components which make up a sample of matter
HClO4
perchloric acid
PO43-
phosphate ion
PCl5
phosphorus pentachloride
PCl3
phosphorus trichloride
K2Cr2O7
potassium dichromate
(CH3)2CO
propanone
CH3C≡CH
propyne
Properties of Subatomic Particles
protons- +1, 1 neutrons- 0,1 electrons- -1, 1/1836
Properties
qualities or attributes which we can use to distinguish one sample of matter from another.
Germanium Tetrachloride (GeCl4)
reactive liquid
hepta
seven
Structural formula
shows the order in which atoms are bonded together in a molecule and what types of bonds are present
hexa
six
NaClO3
sodium chlorate
NaClO2
sodium chlorite
NaClO
sodium hypochlorite
NaNO3
sodium nitrate
NaNO2
sodium nitrite
NaClO4
sodium perchlorate
Na2S2O3
sodium thiosulfate
Strong Electrolytes
substances that completely dissociate into ions when they dissolve in water
Indicators
substances that have distinctly different colors in acidic and basic media
SO42-
sulfate ion
SO32-
sulfite ion
nomenclature.
systematic method of naming compounds
Functional Groups
the components of organic molecules that are most commonly involved in chemical reactions
Oxidation State
the condition of an atom expressed by the number of electrons that the atom needs to reach its elemental form
Precision
the degree of reproducibility of the measurement
Molecular formula
the formula of an actual molecule of that compound
Formula mass(weight)
the mass of a formula unit (weighted average of naturally occurring isotopes) in atomic mass units (u)
Equivalence point
the point at which the two solutions used in a titration are present in chemically equivalent amounts
Empirical formula
the simplest formula for a compound, it shows the atoms present and their relative number with the subscripts being reduced to the lowest whole number ratios
molecule
the smallest dicrete unit of an element or a compound which contains two or more atoms. A molecule is the smallest unit of a substance which retains the chemical properties of that substance.
Malleability
the substances ability to be shaped by external forces.(example: hammering copper metal into a thin sheet)
Half Reactions
the two parts of an oxidation-reduction reaction, one representing oxidation, the other reduction
tri
three
di
two
Density
typically expressed in terms of g/cm3 or g/mL
state of matter
whether or not the substance is a solid, liquid or gas.
ZnSO4•7 H2O
zinc sulfate heptahydrate
