Combo with "AP Chem Cations and Anions" and 13 others
Avogadro's Law
"Equal volumes of gases at the same temp and pressure have the same number of molecules!"
ammonium dichromate
(NH₄)₂Cr₂O₇
ammonium sulfide
(NH₄)₂S
formal charge=
(valence e-) - [(nonbonding e-) + (.5 bonding e-)]
Define chemical formula and give an example.
- A chemical formula is the representation of chemical substances using their chemical symbols and appropriate subscripts for the numbers of atoms. - A simmple formula is Ca(NO₃)₂ - This formula indicates a compound with one calcium atom, two nitrogen atoms, and six oxygen atoms.
Define a hydrate and give an example.
- A hydrate is a substance that contains a fixed number of water molecules. - The water molecules are written separately from the formula itself and connected to it with a dot in the center of the line between the chemical formula and the water molecules. - For example, cobalt(II) chloride hexahydrate is CoCl₂∙6H₂O.
Define metal.
- A metal is a substance with characteristic properties of high electrical conductivity, malleability, and a metallic-silver or yellow luster. - A metal can also be described as the nucleus and core electrons in a "sea of valence electrons."
Define metalloid.
- A metalloid is an element that has properties of both metals and nonmetals. - They are useful as semiconductors. - Examples are silicon and germanium, which are used for transistors and integrated circuits.
Define proton.
- A proton is one of the three particles that make up the atom along with an electron and a neutron. - The proton has a positive charge, equal in magnitude (but with the opposite sign) to the charge of the electron. - The number of protons is equal to the atomic number of an element. - Protons, along with neutrons, are located in the nucleus and make up the bulk of an atom's mass.
Describe the three states of matter.
- A solid retains both volume and shape and molecules are relatively rigid in a crystal lattice. - A liquid retains volume but not shape, molecules move freely in close contact, and strong attractive forces are present. - A gas assumes the volume and shape of a container, molecules are almost totally independent of each other, there is little in the way of attractive forces, and it is highly compressable.
What is an amorphous substance?
- Amorphous means without structure. - Amorphous substances have no long-range crystal structure. - Glass is an example of an amorphous substance.
Define electron, neutron and proton.
- An electron occupies space outside the nucleus. Mass = 0, charge = 1-, symbol is e- - A neutron is an uncharged particle in the nucleus. Symbol is ¹n, mass = 1 - A proton is positively charged article in the nucleus. Symbol is ¹₁p, mass =1, charge = 1+ Number of protons defines the identity of an atom
What concept includes allotropes?
- An element exhibits allotropy when it can exist as two or more distinct chemical structures. - For example: the three allotropes of carbon, graphite, diamond, and C₆₀(buckminsterfullerence). - Not all elements have allotropes.
Define ion.
- An ion is an element that has lost or gained one or more electrons. - A cation has lost one or more electrons. - An anion has gained one or more electrons. - A polyatomic ion is a group of elements bound together covalently that also carries a charge.
What is the concept behind a chemical compound?
- Chemists envision a chemical compound as a substance that has a fixed ratio (by either mass or atoms) of two or more different atoms. - Chemical compounds have definite physical properties that can be used to identify the compound. - Chemical compounds also have characteristic chemical properties that indicate what reactions they participate in and what reactions they do not.
Define electron configuration.
- Electron configuration is a listing of the electrons within an atom based upon the sublevels that are filled and relative energies of these sublevels. The electron configuration for silicon is 1s²2s²2p⁶3s²3p². - It is also the sequence in which the sublevel fills can be read, row by row, from the periodic table.
Define element.
- Element is the term given to any one of the 118 distinct particles of known atoms. - Each had distinct chemical and physical properties. - Elements are organized on the basis of atomic numbers in the periodic table.
Define malleable.
- Malleable is a property of metals. - Malleable means that the substance can be hammered into new shapes.
Define metallic crystal.
- Metallic crystals form from the metals in the periodic table. - These crystals are malleable, ductile, and conduct electricity. - Chemists view metallic crystals as a lattice of nuclei and core electrons in a "sea" of mobile valence electrons.
- Describe where on the periodic table metals, nonmetals, and metalloids are found. - Describe their properties.
- Metals are found on the lower left of the periodic table, nonmetals on the upper right of the table, and metalloids are on either side of a staircase line that starts between boron and aluminum, aluminum to silicon, and then to the bottom of the table. - Metals typically have a silvery luster (except gold and copper), conduct electricity easily, and are malleable and ductile. - Nonmetals have none of the properties that metals have. - Metalloids have properties in between metals and nonmetals. They are semiconductors of electricity.
Define noble gases.
- Noble gases are the elements in the last group in the periodic table (Group ViiiA or 18). - They are unusually stable elements and all, except He, have ns², np⁶ valence electrons. - The noble gases include: - helium, He - neon, Ne - argon, Ar - krypton, Kr - xenon, Xe - radon, Rn
Describe the nuclear charge and the concept of effective nuclear charge.
- Nuclear charge is the number of positive charges in the nucleus. This is the same as the number of protons in the nucleus (Z) and is also the atomic number. - The concept of effective nuclear charge is that inner (core) electrons shield outer electrons from the nuclear charge. The outer electrons are attracted by a nuclear that is approximately equal to the number of valence electrons. - The atomic radius decreases from left to right across the periodic table because the effective nuclear charge increases as the number of valence electrons increases. The increase in effective nuclear charge increases the attractive force on the valence electrons, resulting in a smaller atom.
Amphoteric and amphiprotic substances have the same underlying concept. Explain.
- Substances that can both donates and accept protons are amphiprotic. Example: H₂PO₄₋. - Substances that can react as both an acid and a base are amphoteric. Example: Al₂O₃.
Define Bohr atom.
- The Bohr atom is the model of the atom developed by Niels Bohr. - This model viewed electrons circling the nucleus like a miniature solar system. - Each orbit had a definite energy and electrons moving from one orbit to another and either absorbed or emitted the energy difference between the orbits. - This theory replicated the Rydberg Constant to less than +-0.5%.
What is the Pauli Exclusion Principle?
- The Pauli Exclusion Principle requires that no two electrons in an atom may have the same set of four quantum numbers: n, l, ml, and ms. - This is also means that no two electrons with the same wave equations can coexist. - Another interpretation is that no two electrons can occupy exactly the same space at the same time.
Define periodic table, period, and group.
- The arrangement of the elements in an orderly fashion that shows the relationships of their electronic, chemical, and physical characteristics is the modern periodic table. - Each row in the table is called a period. - Each column in the talbe is called a group (and sometimes family).
What is the concept behind the atomic number?
- The atomic number is the number that specifies the position of an element in the periodic table. - It is also a number representing the number of protons in the nucleus of an atom. - The atomic number defines the identity of an element. The basic ida is that each element must fall in some sort of order. Mendeleev originally based the order on atomic masses. It was later found that the order should be based on the number of protons in the nucleus of an element.
- What is the formula for aluminum fluoride? - What are the symbols for the ions that make up this compound?
- The formula for aluminum fluoride is AlF₃. - The ions that make up this compound are Al³⁺ and F⁻.
- What is the formula for aluminum hydrogen sulfate or aluminum bisulfate? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for aluminum hydrogen sulfate is Al(HSO₄)₃. - The ions that make up this compound are Al³⁺ and HSO₄⁻. - The oxidation number of S is +6. - The HSO₄⁻ ion is often called the bisulfate ion.
- What is the formula for aluminum nitrate nonahydrate? - What are the symbols for the ions that make up this compound?
- The formula for aluminum nitrate nonahydrate is Al(NO₃)₃∙9H₂O. - The ions that make up this compound are Al³⁺ and NO₃⁻.
- What is the formula for aluminum sulfate? - What are the symbols for the ions that make up this compound?
- The formula for aluminum sulfate is Al₂(SO₄)₃. - The ions that make up this compound are Al³⁺ and SO₄²⁻.
- What is the formula for ammonium nitrate (fertilizer)? - What are the symbols for the ions that make up this compound?
- The formula for ammonium nitrate is NH₄NO₃. - The ions that make up this compound are NH₄⁺ and NO₃⁻.
- What is the formula for ammonium sulfate? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nitrogen in the ammonium ion?
- The formula for ammonium sulfate is (NH₄)₂SO₄. - The ions that make up this compound are NH₄⁺ and SO₄²⁻. - The oxidation number of N is -3.
- What is the formula for ammonium sulfite? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for ammonium sulfite is (NH₄)₂SO₃. - The ions that make up this compound are NH₄⁺ and SO₃²⁻. - The oxidation number of S is +4.
- What is the formula for barium hydroxide octahydrate? - What are the symbols for the ions that make up this compound?
- The formula for barium hydroxide octahydrate is Ba(OH)₂∙8H₂O. - The ions that make up this compound are Ba²⁺ and OH⁻.
- What is the formula for calcium bromate? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for calcium bromate is Ca(BrO₃)₂. - The ions that make up this compound are Ca²⁺ and BrO₃⁻. - The oxidation number of Br is +5.
- What is the formula for calcium carbonate (limestone)? - What are the symbols for the ions that make up this compound?
- The formula for calcium carbonate is CaCO₃. - The ions that make up this compound are Ca²⁺ and CO₃²⁻.
- What is the formula for calcium chlorate? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for calcium chlorate is Ca(ClO₃)₂. - The ions that make up this compound are Ca²⁺ and ClO₃⁻. - The oxidation number of Cl is +5.
- What is the formula for calcium chloride dihydrate? - What are the symbols for the ions that make up this compound?
- The formula for calcium chloride dihydrate is CaCl₂∙2H₂O. - The ions that make up this compound are Ca²⁺ and Cl⁻.
- What is the formula for calcium hydrogen sulfite? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for calcium hydrogen sulfite is Ca(HSO₃)₂. - The ions that make up this compound are Ca²⁺ and HSO₃⁻. - The oxidation number of S is +4.
- What is the formula for calcium nitride? - What are the symbols for the ions that make up this compound?
- The formula for calcium nitrde is Ca₃N₂. - The ions that make up this compound are Ca²⁺ and N³⁻.
- What is the formula for calcium oxide (lime)? - What are the symbols for the ions that make up this compound?
- The formula for calcium oxide is CaO. - The ions that make up this compound are Ca²⁺ and O²⁻.
- What is the formula for calcium sulfate (plaster of Paris)? - What are the symbols for the ions that make up this compound?
- The formula for calcium sulfate is CaSO₄. - The ions that make up this compound are Ca²⁺ and SO₄²⁻.
- What is the formula for chromium(III) phosphate? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for chromium(III) phosphate is CrPO₄. - The ions that make up this compound are Cr³⁺ and PO₄³⁻. - The oxidation number of P is +5.
- What is the formula for chromium(III) sulfite? - What are the symbols for the ions that make up this compound?
- The formula for chromium(III) sulfite is Cr₂(SO₃)₃. - The ions that make up this compound are Cr³⁺ and SO₃²⁻.
- What is the formula for cobalt(III) chloride? - What are the symbols for the ions that make up this compound?
- The formula for cobalt(III) chloride is CoCl₃. - The ions that make up this compound are Co³⁺ and Cl⁻.
- What is the formula for copper sulfate pentahydrate? - What are the symbols for the ions that make up this compound?
- The formula for copper sulfate pentahydrate is CuSO₄∙5H₂O. - The ions that make up this compound are Cu²⁺ and SO₄²⁻.
- What is the formula for gold(III) phosphate? - What are the symbols for the ions that make up this compound?
- The formula for gold(III) phosphate is AuPO₄. - The ions that make up this compound are Au³⁺ and PO₄³⁻.
- What is the formula for iron(III) nitrate? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for iron(III) nitrate is Fe(NO₃)₃. - The ions that make up this compound are Fe³⁺ and NO₃⁻. - The oxidation number of N is +5.
- What is the formula for lead(II) acetate? - What are the symbols for the ions that make up this compound?
- The formula for lead(II) acetate is Pb(C₂H₃O₂)₂. - The ions that make up this compound are Pb²⁺ and C₂H₃O₂⁻.
- What is the formula for lead(IV) phosphate? - What are the symbols for the ions that make up this compound?
- The formula for lead(IV) phosphate is Pb₃(PO₄)₄. - The ions that make up this compound are Pb⁴⁺ and PO₄³⁻.
- What is the formula for lithium bromide? - What are the symbols for the ions that make up this compound?
- The formula for lithium bromide is LiBr. - The ions that make up this compound are Li⁺ and Br⁻.
- What is the formula for lithium bromite? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for lithium bromite is LiBrO₂. - The ions that make up this compound are Li⁺ and BrO₂⁻. - The oxidation number of Br is +3.
- What is the formula for lithium carbonate? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for lithium carbonate is Li₂CO₃. - The ions that make up this compound are Li⁺ and CO₃²⁻. - The oxidation number of C is +4.
- What is the formula for magnesium chloride hexahydrate? - What are the symbols for the ions that make up this compound?
- The formula for magnesium chloride hexahydrate is MgCl₂∙6H₂O. - The ions that make up this compound are Mg²⁺ and Cl⁻.
- What is the formula for magnesium hydroxide? - What are the symbols for the ions that make up this compound?
- The formula for magnesium hydroxide is Mg(OH)₂. - The ions that make up this compound are Mg²⁺ and OH⁻.
- What is the formula for magnesium phosphide? - What are the symbols for the ions that make up this compound?
- The formula for magnesium phosphide is Mg₃P₂. - The ions that make up this compound are Mg²⁺ and P³⁻.
- What is the formula for mercury(I) iodide? - What are the symbols for the ions that make up this compound?
- The formula for mercury(I) iodide is Hg₂I₂. - The ions that make up this compound are Hg₂²⁺ and I⁻.
- What is the formula for nickel(II) hydrogen carbonate or nickel(II) bicarbonate? - What are the symbols for the ions that make up this compound?
- The formula for nickel(II) hydrogen carbonate is Ni(HCO₃)₂. - The ions that make up this compound are Ni²⁺ and HCO₃⁻. The HCO₃⁻ ion is also called bicarbonate.
- What is the formula for potassium dichromate? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for potassium dichromate is K₂Cr₂O₇. - The ions that make up this compound are K⁺ and Cr₂O₇²⁻. - The oxidation number of Cr is +6.
- What is the formula for potassium dihydrogen phosphate? - What are the symbols for the ions that make up this compound?
- The formula for potassium dihydrogen phosphate is KH₂PO₄. - The ions that make up this compound are K⁺ and H₂PO₄⁻.
- What is the formula for potassium permanganate? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for potassium permanganate is KMnO₄. - The ions that make up this compound are K⁺ and MnO₄⁻. - The oxidation number of Mn is +7.
- What is the formula for rubidium perbromate? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for rubidium perbromate is RbBrO₄. - The ions that make up this compound are Rb⁺ and BrO₄⁻. - The oxidation number of Br is +7.
- What is the formula for silver chloride? - What are the symbols for the ions that make up this compound?
- The formula for silver chloride is AgCl. - The ions that make up this compound are Ag⁺ and Cl⁻.
- What is the formula for silver chromate? - What are the symbols for the ions that make up this compound?
- The formula for silver chromate is AgCrO₄. - The ions that make up this compound are Ag⁺ and CrO₄²⁻.
- What is the formula for sodium chloride? - What are the symbols for the ions that make up this compound?
- The formula for sodium chloride is NaCl. - The ions that make up this compound are Na⁺ and Cl⁻.
- What is the formula for sodium chromate? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for sodium chromate is Na₂CrO₄. - The ions that make up this compound are Na⁺ and CrO₄²⁻. - The oxidation number of Cr is +6.
- What is the formula for sodium hydrogen phosphate? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for sodium hydrogen phosphate is Na₂HPO₄. - The ions that make up this compound are Na⁺ and HPO₄²⁻. - The oxidation number of P is +5.
- What is the formula for sodium hypobromite? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for sodium hypobromite is NaBrO. - The ions that make up this compound are Na⁺ and BrO⁻. - The oxidation number of Br is +1.
- What is the formula for sodium hypochlorite? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for sodium hypochlorite is NaClO. - The ions that make up this compound are Na⁺ and ClO⁻. - The oxidation number of Cl is +1.
- What is the formula for sodium nitrite (food preservative)? - What are the symbols for the ions that make up this compound?
- The formula for sodium nitrite is NaNO₂. - The ions that make up this compound are Na⁺ and NO₂⁻.
- What is the formula for sodium nitrite? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for sodium nitrite is NaNO₂. - The ions that make up this compound are Na⁺ and NO₂⁻. - The oxidation number of N is +3.
- What is the formula for sodium oxalate? - What are the symbols for the ions that make up this compound?
- The formula for sodium oxalate is Na₂C₂O₄. - The ions that make up this compound are Na⁺ and C₂O₄²⁻.
- What is the formula for sodium oxide? - What are the symbols for the ions that make up this compound?
- The formula for sodium oxide is Na₂O. - The ions that make up this compound are Na⁺ and O²⁻.
- What is the formula for strontium iodide? - What are the symbols for the ions that make up this compound?
- The formula for strontium iodide is SrI₂. - The ions that make up this compound are Sr²⁺ and I⁻.
- What is the formula for tin(II) carbonate? - What are the symbols for the ions that make up this compound?
- The formula for tin(II) carbonate is SnCO₃. - The ions that make up this compound are Sn²⁺ and CO₃²⁻.
- What is the formula for titanium(II) chlorite? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for titanium(II) chlorite is Ti(Clo₂)₂. - The ions that make up this compound are Ti²⁺ and ClO₂⁻. - The oxidation number of Cl is +3.
- What is the formula for tungsten(VI) oxide? - What are the symbols for the ions that make up this compound?
- The formula for tungsten(VI) oxide is WO₃. - The ions that make up this compound are W⁶⁺ and O²⁻.
- What is the formula for zinc sulfate? - What are the symbols for the ions that make up this compound? - What is the oxidation number of the nonoxygen element in the polyatomic ion?
- The formula for zinc sulfate is ZnSO₄. - The ions that make up this compound are Zn²⁺ and SO₄²⁻. - The oxidation number of S is +6.
Define nucleus.
- The nucleus is the center of the atom that contains the protons and neutrons. - It comprises an extremely small fraction of the atom's volume. - The nucleus is extremely dense while the rest of the atom is primarily empty space.
Define and describe the four quantum numbers.
- The principal quantum number, n, indicates the shell (or principal energy level) of the atom min which and electron resides and is a rough measure of the distance from the nucleus. n can have any value from 1 to infinity. - The azimuthal quantum number, l, specifies the sublevel (or orbital) that an electron is located in. This number may be any number from zero up to n-1. This number specifies the shape of an atomic orbital. - The magnetic quantum number,m, indicates the orientation of a sublevel in space. The value of the m, quantum number can be any number from -l to zero to +l. - The spin quantum number,ms, indicates the spin of the electron. An electron is not actually spinning but it is a useful way of visualizing this property. Acceptable values for spin are +¹/₂ and -¹/₂. - Each electron possesses four quantum numbers. Each electron in an atom must have a set of four quantum numbers different from all other electrons in that atom.
What are the subatomic constituents of the atom?
- The subatomic constituents of the atom are protons and neutrons in the nucleus and electrons around the nucleus. - The atomic number Z=# of protons. - The mass number A=# protons + # of neutrons.
Define transition elements and inner transition elements.
- Transition elements are sometimes called d-block elements because the last electrons added to form the element are electrons occupying d orbitals. -They occupy the 10 groups in the center of the periodic table. -They often form colored ionic compounds. -They often have multiple possible oxidation states. -They often form polyatomic anions. - Inner transition elements have electrons occupying f orbitals as the last electrons added to their electron configuration. -They occupy the 14 groups that are often placed outside the periodic table. -They contain many of the radioactive and manmade elements.
Define weight and mass.
- Weight is the force developed due to the gravitational attraction of two masses toward each other. The weight of an object changes as the gravitational acceleration changes. weight=force=mass X acceleration of gravity - Mass is the amount of material in a sample. Mass does not change with the acceleration of gravity.
colligative properties - examples
- vapor pressure lowering - freezing point depression - boiling point elevation - osmotic pressure
changes in temperature (le chat)
1. exothermic: shifts left (rcnts > prod + heat) 2. endothermic: shifts right (heat + rcnts > prod)
changes in pressure/volume (le chat)
1. increase volume, decrease pressure, reaction shifts to side with more moles. 2. decrease volumes, increase pressure, reaction shifts to side with lesser moles. 3. if moles are equal, no shift occurs.
Entropy increases when:
1. molecules increase 2. temp increases 3. l - g, s- g, s - l
left to right
1. radius decreases 2. ionization energy, electronegativity, and electron affinity increase
top to bottom
1. radius increases 2. ionization energy, electronegativity, and electron affinity decrease
D
2 K + 2 H2O ---> 2 K+ + 2 OH¯ + H2 When 0.400 mole of potassium reacts with excess water at standard temperature and pressure as shown in the equation above, the volume of hydrogen gas produced is (A) 1.12 liters (B) 2.24 liters (C) 3.36 liters (D) 4.48 liters (E) 6.72 liters
E
6 I¯ + 2 MnO4¯ + 4 H2O(l) ---> 3 I2(s) + 2 MnO2(s) + OH¯ Which of the following statements regarding the reaction represented by the equation above is correct? (A) Iodide ion is oxidized by hydroxide ion. (B) MnO4¯ is oxidized by iodide ion. (C) The oxidation number of manganese changes from +7 to +2. (D) The oxidation number of manganese remains the same. (E) The oxidation number of iodine changes from -1 to 0.
G = 0
@ equilibrium
C
A 1.0 L sample of an aqueous solution contains 0.10 mol of NaCl and 0.10 mol of CaCl2. What is the minimum number of moles of AgNO3 that must be added to the solution in order to precipitate all of the Cl¯ as AgCl(s) ? (Assume that AgCl is insoluble.) (A) 0.10 mol (B) 0.20 mol (C) 0.30 mol (D) 0.40 mol (E) 0.60 mol
C
A 2.00-liter sample of nitrogen gas at 27 °C and 600. millimeters of mercury is heated until it occupies a volume of 5.00 liters. If the pressure remains unchanged, the final temperature of the gas is (A) 68 °C (B) 120 °C (C) 477 °C (D) 677 °C (E) 950. °C
C
A gaseous mixture containing 7.0 moles of nitrogen, 2.5 moles of oxygen, and 0.50 mole of helium exerts a total pressure of 0.90 atmosphere. What is the partial pressure of the nitrogen? (A) 0.13 atm (B) 0.27 atm (C) 0.63 atm (D) 0.90 atm (E) 6.3 atm
E
A hot-air balloon rises. Which of the following is the best explanation for this observation? (A) The pressure on the walls of the balloon increases with increasing tempearature. (B) The difference in temperature between the air inside and outside the ballon produces convection currents. (C) The cooler air outside the balloon pushes in on the walls of the ballon. (D) The rate of diffusion of cooler air is less than that of warmer air. (E) The air density inside the ballon is less than that of the surrounding air.
C
A hydrocarbon gas with an empirical formula CH2 has a density of 1.88 grams per liter at 0 °C and 1.00 atmosphere. A possible formula for the hydrocarbon is (A) CH2 (B) C2H4 (C) C3H6 (D) C4H8 (E) C5H10
D
A measured mass of an unreactive metal was dropped into a small graduated cylinder half filled with water. The following measurements were made. Mass of metal = 19.611 grams Volume of water before addition of metal = 12.4 milliliters Volume of water after addition of metal = 14.9 milliliters The density of the metal should be reported as (A) 7.8444 grams per mL (B) 7.844 grams per mL (C) 7.84 grams per mL (D) 7.8 grams per mL (E) 8 grams per mL
anion
A negatively charged ion
cation
A positively charged ion
C
A sample of 0.0100 mole of oxygen gas is confined at 37° C and 0.216 atmosphere. What would be the pressure of this sample at 15° C and the same volume? (A) 0.0876 atm (B) 0.175 atm (C) 0.201 atm (D) 0.233 atm (E) 0.533 atm
D
A sample of 3.30 grams of an ideal gas at 150.0 °C and 1.25 atmospheres pressure has a volume of 2.00 liters. What is the molar mass of the gas? The gas constant, R, is 0.0821 L atm mol¯1 K¯1). (A) 0.0218 gram/mole (B) 16.2 grams/mole (C) 37.0 grams/mole (D) 45.8 grams/mole (E) 71.6 grams/mole
C
A sample of 61.8 g of H3BO3, a weak acid is dissolved in 1,000 g of water to make a 1.0-molal solution. Which of the following would be the best procedure to determine to molarity of the solution? (Assume no additional information is available.) (A) Titration of the solution with standard acid (B) Measurement of the pH with a pH meter (C) Determination of the boiling point of the solution (D) Measurement of the total volume of the solution (E) Measurement of the specific heat of the solution
C
A sample of an ideal gas is cooled from 50.0 °C to 25.0 °C in a sealed container of constant volume. Which of the following values for the gas will decrease? I. The average molecular mass of the gas II. The average distance between the molecules III. The average speed of the molecules (A) I only (B) II only (C) III only (D) I and III (E) II and III
E
A solution is known to contain an inorganic salt of one of the following elements. The solution is colorless. The solution contains a salt of (A) Cu (B) Mn (C) Fe (D) Ni (E) Zn
A
A white solid is observed to be insoluble in water, insoluble in excess ammonia solution, and soluble in dilute HCl. Which of the following compounds could the solid be? (A) CaCO3 (B) BaSO4 (C) Pb(NO3)2 (D) AgCl (E) Zn(OH)2
A
A yellow precipitate forms when 0.5 M NaI(aq) is added to a 0.5 M solution of which of the following ions? A) Pb2+(aq) B) Zn2+(aq) C) CrO42¯(aq) D) SO42¯(aq) E) OH¯(aq)
C₂H₃O₂⁻¹
Acetate
A
Adding water to some chemicals can be dangerous because large amounts of heat are liberated. Which of the following does NOT liberate heat when water is added to it? (A) KNO3 (B) NaOH (C) CaO (D) H2SO4 (E) Na
B
After completing an experiment to determine gravimetrically the percentage of water in a hydrate, a student reported a value of 38 percent. The correct value for the percentage of water in the hydrate is 51 percent. Which of the following is the most likely explanation for this difference? A) Strong initial heating caused some of the hydrate sample to spatter out of the crucible. B) The dehydrated sample absorbed moisture after heating. C) The amount of the hydrate sample used was too small. D) The crucible was not heated to constant mass before use. E) Excess heating caused the dehydrated sample to decompose.
Silver
Ag
Silver
Ag +1
silver nitrite
AgNO₂
Aluminum
Al
Aluminum
Al +3
aluminum hydroxide
Al(OH)₃
aluminum selenide
Al2Se3
aluminum fluoride
AlF₃
Define alkali metals
Alkali metals are the elements in the first group (column) of the periodic table (IA or 1). All: - have ns¹ electrons as valance electrons - are extremely reactive elements - have low ionization energies - have low electronegativities - are metals that form 1+ ions
Define alkaline earth metals.
Alkaline earth metals are the elements in the second group (column) of the periodic table. All: - have ns² electrons for valence electrons - are very reactive elements - are metals that form 2+ ions
Al⁺³
Aluminum
aluminum chromate
Al₂(CrO₄)₃
NH₂⁻¹
Amide
NH₄⁺¹
Ammonium
C
An impossible electronic configuration (A) 1s2 2s2 2p5 3s2 3p5 (B) 1s2 2s2 2p6 3s2 3p6 (C) 1s2 2s2 2p6 2d10 3s2 3p6 (D) 1s2 2s2 2p6 3s2 3p6 3d5 (E) 1s2 2s2 2p6 3s2 3p6 3d3 4s2
Define isotope.
An isotope is an atom with a specific number of neutrons in addition to the protons and electrons defined by the atomic number. A given element may have several isotopes, each of which has a different number of neutrons while having he same number of electrons and protons. - For example: C-12, C-13, and C-14. Each has 6 electrons and 6 protons; however, they have 6,7, and 8 neutrons, respectively.
Sb⁺³
Antimony (III) or Antiminous
Sb⁺⁵
Antimony (V) or Antimonic
Argon
Ar
AsO₄⁻²
Arsenate
As⁻³
Arsenide
Arsenic
As
C
As the temperature is raised from 20 ° C to 40 ° C, the average kinetic energy of neon atoms changes by a factor of (A) 1/2 (B) [square root of](313/293) (C) 313/293 (D) 2 (E) 4
Arsenate
AsO4 -3
Astatine
At
D
At 25 °C, a sample of NH3 (molar mass 17 grams) effuses at the rate of 0.050 mole per minute. Under the same conditions, which of the following gases effuses at approximately one-half that rate? (A) O2 (molar mass 32 grams) (B) He2 (molar mass 4.0 grams) (C) CO2 (molar mass 44 grams) (D) Cl2 (molar mass 71 grams) (E) CH4 (molar mass 16 grams)
Gold
Au
Gold
Au +1 and +3
Boron
B
Borate
BO3 -3
Barium
Ba
barium hypochlorite
Ba(ClO)₂
barium hydroxide
Ba(OH)₂
barium oxide
BaO
Ba⁺²
Barium
Beryllium
Be
Be⁺²
Beryllium
Bismuth
Bi
BO₃⁻³
Borate
B⁺³
Boron
Bromine
Br
HypoBromite
BrO -1
Bromate
BrO3 -1
bromate
BrO₃⁻
BrO₃⁻¹
Bromate
Br⁻¹
Bromide
BrO₂⁻¹
Bromite
Carbon
C
Acetate
C2H3O2 -1
Oxalate
C2O4 -2
Tartrate
C4H4O6 -2
C
CCl4, CO2, PCl3, PCl5, SF6 Which of the following does not describe any of the molecules above? (A) Linear (B) Octahedral (C) Square planar (D) Tetrahedral (E) Trigonal pyramidal
D
CH3CH2OH boils at 78 °C and CH3OCH3 boils at - 24 °C, although both compounds have the same composition. This difference in boiling points may be attributed to a difference in (A) molecular mass (B) density (C) specific heat (D) hydrogen bonding (E) heat of combustion
methane
CH₄
Cyanide
CN -1
cyanide
CN⁻
Carbonate
CO3 -2
carbonate
CO₃²⁻
Calcium
Ca
calcium chlorite
Ca(ClO₂)₂
calcium oxide
CaO
Cd⁺²
Cadmium
Ca⁺²
Calcium
C
Can be used to predict that a gaseous carbon atom in its ground state is paramagnetic (A) Heisenberg uncertainty principle (B) Pauli exclusion principle (C) Hund's rule (principle of maximum multiplicity) (D) Shielding effect (E) Wave nature of matter
C⁻⁴
Carbide
D
Carbon dioxide, CO2(s) (A) Lattice of positive and negative ions held together by electrostatic forces. (B) Closely packed lattice with delocalized electrons throughout (C) Strong single covalent bonds with weak intermolecular forces. (D) Strong multiple covalent bonds (including bonds.) with weak intermolecular forces (E) Macromolecules held tgether with strong polar bonds.
CO₃⁻²
Carbonate
calcium phosphate
Ca₃(PO₄)₂
Cadmium
Cd
cadmium selenide
CdSe
Cs⁺¹
Cesium
A
Cesium chloride, CsCl (s) (A) Lattice of positive and negative ions held together by electrostatic forces. (B) Closely packed lattice with delocalized electrons throughout (C) Strong single covalent bonds with weak intermolecular forces. (D) Strong multiple covalent bonds (including bonds.) with weak intermolecular forces (E) Macromolecules held tgether with strong polar bonds.
ClO₃⁻¹
Chlorate
Cl⁻¹
Chloride
ClO₂⁻¹
Chlorite
CrO₄⁻²
Chromate
Cr⁺²
Chromium (II) or Chromous
Cr⁺³
Chromium (III) or Chromic
Chlorine
Cl
Hypochlorite
ClO -1
Chlorite
ClO2 -1
Chlorate
ClO3 -1
Perchlorate
ClO4 -1
hypochlorite
ClO⁻
chlorite
ClO₂⁻
chlorate
ClO₃⁻
perchlorate
ClO₄⁻
dichlorine heptoxide
Cl₂O₇
Cobalt
Co
Cobalt
Co +2 and +3
cobalt (III) nitrate
Co(NO3)3
Co⁺²
Cobalt(II) or Cobaltous
Co⁺³
Cobalt(III) or Cobaltic
C
Contains 1 sigma (s) and 2 pi (p) bonds (A) Li2 (B) B2 (C) N2 (D) O2 (E) F2
Cu⁺¹
Copper(I) or Cuprous
Cu⁺²
Copper(II) or Cupric
Chromium
Cr
Chromium
Cr +2 and +3
chromium(III) nitrate
Cr(NO₃)₃
chromium (II) carbonate
Cr2(CO3)3
DiChromate
Cr2O7 -2
Chromate
CrO4 -2
chromate
CrO₄²⁻
dichromate
Cr₂O₇²⁻
Cesium
Cs
Copper
Cu
Copper
Cu +1 and +2
copper (II) bicarbonate
Cu(HCO₃)₂
copper (II) nitrate
Cu(NO3)2
copper(II) bromide
CuBr₂
copper (I) chloride
CuCl
copper(I) selenide
Cu₂Se
CN⁻¹
Cyanide
acetate
C₂H₃O₂⁻
ethene
C₂H₄
ethane
C₂H₆
oxalate
C₂O₄²⁻
pentane
C₅H₁₂
Colligative Properties
Depend on the amount of solute particles
Cr₂O₇⁻²
Dichromate
H₂PO₄⁻¹
Dihydrogen phosphate
Positive deviation of Raoult's Law
Endothermic reaction. Weak interactions between solute and solvent. Higher tendency to escape into gas phase.
A
Equal numbers of moles of He(g), Ar(g), and Ne(g) are placed in a glass vessel at room temperature. If the vessel has a pinhole-sized leak, which of the following will be true regarding the relative values of the partial pressures of the gases remaining in the vessel after some of the gas mixture has effused? A) PHe < PNe < PAr B) PHe < PAr < PNe C) PNe < PAr < PHe D) PAr < PHe < PNe E) PHe = PAr = PNe
CH3-CH2-CH2-CH3 and CH (with 3 CH3 attached w/ a single bond)
Examples of isomers
Negative deviation of Raoult's Law
Exothermic reactions. Strong interaction between solute and solvent. Lower tendency to escape into gas phase
E
Explains the experimental phenomenon of electron diffraction (A) Heisenberg uncertainty principle (B) Pauli exclusion principle (C) Hund's rule (principle of maximum multiplicity) (D) Shielding effect (E) Wave nature of matter
Fluorine
F
Iron
Fe
Iron
Fe +2 and +3
iron(II) cyanide
Fe(CN)₂
iron(II) sulfide
FeS
FO₃⁻¹
Fluorate
F⁻¹
Fluoride
FO₂⁻¹
Fluorite
D
For the substance represented in the diagram, which of the phases is most dense and which is least dense at - 15 °C. Most Dense Least Dense (A) Solid Gas (B) Solid Liquid (C) Liquid Solid (D) Liquid Gas (E) The diagram gives no information about densities.
Francium
Fr
The frequency of a photon is____________ to its wavelength. a. directly proportional b. inversely proportional c. not related If a or b is chose, write the appropriate equation.
Frequency is inversely proportional to wavelength, as shown in the equation: v(frequency)=c
Gallium
Ga
gallium arsenide
GaAs
Germanium
Ge
Au⁺¹
Gold
B
Gold, Au(s) (A) Lattice of positive and negative ions held together by electrostatic forces. (B) Closely packed lattice with delocalized electrons throughout (C) Strong single covalent bonds with weak intermolecular forces. (D) Strong multiple covalent bonds (including bonds.) with weak intermolecular forces (E) Macromolecules held tgether with strong polar bonds.
Hydrogen
H
D
H2(g) + (1/2) O2(g) ---> H2O(l) DH° = - 286 kJ 2 Na(s) + (1/2) O2(g) ---> Na2O(s) DH° = - 414 kJ Na(s) + (1/2) O2(g) + (1/2) H2(g) ---> NaOH(s) DH° = - 425 kJ Based on the information above, what is the standard enthalpy change for the following reaction? Na2O(s) + H2O(l) ---> 2 NaOH(s) (A) -1,125 kJ (B) -978 kJ (C) -722 kJ (D) -150 kJ (E) +275 kJ
Dihydrogen phosphate
H2PO4 -1
hydrobromic acid
HBr
Hydrogen carbonate
HCO3 -1
hydrogen carbonate (bicarbonate)
HCO₃⁻
hydrochloric acid
HCl
perchloric acid
HClO₄
acetic acid
HC₂H₃O₂
iodous acid
HIO2
nitrous acid
HNO2
Hydrogen Phospate
HPO4 -2
Hydrogen Sulfide
HS -1
Hydrogen Sulfite
HSO3 -1
hydrogen sulfite (bisulfite)
HSO₃⁻
hydrogen sulfate (bisulfate)
HSO₄⁻
Define halogen.
Halogens are the elements in the next to last group of the periodic table (Group VIIA or 17). They are reactive elements with an ns², np⁵ valence electron structure. - Halogens commonly form salts (halide is another name for salt) with metals. - Halogens generally have high electronegativities. Fluorine has the highest electronegativity in the periodic table. - In simple compounds, halogens tend to form only one bond. However, halogens form multiple bonds with oxygen and other halogens.
D
Has a bond order of 2 (A) Li2 (B) B2 (C) N2 (D) O2 (E) F2
C
Has the largest bond-dissociation energy (A) Li2 (B) B2 (C) N2 (D) O2 (E) F2
Helium
He
Mercury
Hg
Mercury
Hg +2
mercury (I) chloride
Hg2Cl2 (mercury exists as Hg2 2+ ions)
mercury (I) chloride
Hg₂Cl₂
E
How many grams of calcium nitrate, Ca(NO3)2, contains 24 grams of oxygen atoms? (A) 164 grams (B) 96 grams (C) 62 grams (D) 50. grams (E) 41 grams
C
How many milliliters of 11.6-molar HCl must be diluted to obtain 1.0 liter of 3.0-molar HCl? (A) 3.9 mL (B) 35 mL (C) 260 mL (D) 1,000 mL (E) 3,900 mL
A
How many moles of solid Ba(NO3)2 should be added to 300. milliliters of 0.20-molar Fe(NO3)3 to increase the concentration of the NO3¯ ion to 1.0-molar? (Assume that the volume of the solution remains constant.) (A) 0.060 mole (B) 0.12 mole (C) 0.24 mole (D) 0.30 mole (E) 0.40 mole
Define Hund's Rule.
Hund's Rule states that all orbitals in a sublevel must fill with one electron before a second electron of opposite spin can be added to any orbital in that sublevel.
H⁻¹
Hydride
H⁺¹
Hydrogen
HSO₃⁻¹
Hydrogen Sulfite or Bisulfite
HCO₃⁻¹
Hydrogen carbonate or Bicarbonate
B
Hydrogen gas is collected over water at 24 °C. The total pressure of the sample is 755 millimeters of mercury. At 24 °C, the vapor pressure of water is 22 millimeters of mercury. What is the partial pressure of the hydrogen gas? (A) 22 mm Hg (B) 733 mm Hg (C) 755 mm Hg (D) 760 mm Hg (E) 777 mm Hg
HPO₄⁻²
Hydrogen phosphate
HSO₄⁻¹
Hydrogen sulfate or bisulfate
OH⁻¹
Hydroxide
BrO⁻¹
Hypobromite
ClO⁻¹
Hypochlorite
FO⁻¹
Hypofluorite
IO⁻¹
Hypoiodite
carbonic acid
H₂CO₃
chromic acid
H₂CrO₄
dihydrogen monoxide
H₂O
dihydrogen phosphate
H₂PO₄⁻
sulfurous acid
H₂SO₃
phosphoric acid
H₃PO₄
Iodine
I
B
I2(g) + 3 Cl2(g) ---> 2 ICl3(g) According to the data in the table below, what is the value of DH° for the reaction represented above? Bond Average Bond Energy (KJ/mole) I---I 149 Cl---Cl 239 I---Cl 208 (A) - 860 kJ (B) - 382 kJ (C) + 180 kJ (D) + 450 kJ (E) + 1,248 kJ
HypoIodite
IO -1
Iodate
IO3 -1
Periodate
IO4 -1
iodine dioxide
IO₂
iodate
IO₃⁻
E
If 87 grams of K2 SO4 (molar mass 174 grams) is dissolved in enough water to make 250 milliliters of solution, what are the concentrations of the potassium and the sulfate ions? [K+] [SO42¯] (A) 0.020 M 0.020 M (B) 1.0 M 2.0 M (C) 2.0 M 1.0 M (D) 2.0 M 2.0 M (E) 4.0 M 2.0 M
Tyndall Effect
If it is a colloid, the light will be reflected; if it is a homogeneous mixture the light will be invisible
Indium
In
E
In a molecule in which the central atom exhibits sp3d2 hybrid orbitals, the electron pairs are directed toward the corners of (A) a tetrahedron (B) a square-based pyramid (C) a trigonal bipyramid (D) a square (E) an octahedron
B
In a qualitative ananlysis for the presence of Pb2+, Fe2+, and Cu2+ ions in a aqueous solution, which of the following will allow the separation of Pb2+ from the other ions at room temperature? A) Adding dilute Na2S(aq) solution B) Adding dilute HCl(aq) solution C) Adding dilute NaOH(aq) solution D) Adding dilute NH3(aq) solution E) Adding dilute HNO3(aq) solution
D
In the periodic table, as the atomic number increases from 11 to 17, what happens to the atomic radius? (A) It remains constant. (B) It increases only. (C) It increases, then decreases. (D) It decreases only. (E) It decreases, then increases.
B
In which of the following compounds is the mass ratio of chromium to oxygen closest to 1.62 to 1.00 ? (A) CrO3 (B) CrO2 (C) CrO (D) Cr2O (E) Cr2O3
B
Indicates that an atomic orbital can hold no more than two electrons (A) Heisenberg uncertainty principle (B) Pauli exclusion principle (C) Hund's rule (principle of maximum multiplicity) (D) Shielding effect (E) Wave nature of matter
IO₃⁻¹
Iodate
I⁻¹
Iodide
IO₂⁻¹
Iodite
Iridium
Ir
Fe⁺²
Iron(II) or Ferrous
Fe⁺³
Iron(III) or Ferric
D
Is added to silicon to enhance its properties as a semiconductor (A) Pb (B) Ca (C) Zn (D) As (E) Na
D
Is used to explain the fact that the carbon-to-carbon bonds in benzene, C6H6, are identical (A) hydrogen bonding (B) hybridization (C) ionic bonding (D) resonance (E) van der Waals forces (London dispersion forces)
B
Is used to explain the fact that the four bonds in methane are equivalent (A) hydrogen bonding (B) hybridization (C) ionic bonding (D) resonance (E) van der Waals forces (London dispersion forces)
E
Is used to explain why iodine molecules are held together in the solid state (A) hydrogen bonding (B) hybridization (C) ionic bonding (D) resonance (E) van der Waals forces (London dispersion forces)
A
Is used to explain why the boiling point of HF is greater than the boiling point of HBr (A) hydrogen bonding (B) hybridization (C) ionic bonding (D) resonance (E) van der Waals forces (London dispersion forces)
Potassium
K
potassium nitride
K3N
potassium cyanide
KCN
potassium chloride
KCl
potassium chlorate
KClO3
potassium permanganate
KMnO₄
Krypton
Kr
Lanthanum
La
Pb⁺²
Lead(II) or Plumbous
Pb⁺⁴
Lead(IV) or Plumbic
Lithium
Li
lithium nitride
Li3N
lithium hydrogen carbonate
LiHCO₃
Li⁺¹
Lithium
Molar Mass of a Gas
MM=dRT/P (Molar Mass = "dirty pee")
Mg⁺²
Magnesium
Mn⁺²
Manganese (II) or Manganous
Mn⁺³
Manganese (III) or Manganic
D
Mass of an empty container = 3.0 grams Mass of the container plus the solid sample = 25.0 grams Volume of the solid sample = 11.0 cubic centimeters The data above were gathered in order to determinethe density of an unknown solid. The density of the sample should be reported as (A) 0.5 g/cm3 (B) 0.50 g/cm3 (C) 2.0 g/cm3 (D) 2.00 g/cm3 (E) 2.27 g/cm3
Mass Percentage
Mass solute/Total Mass Solution
Hg₂⁺²
Mercury(I) or Mercurous
Hg⁺²
Mercury(II) or Mercuric
C
Methane, CH4(s) (A) Lattice of positive and negative ions held together by electrostatic forces. (B) Closely packed lattice with delocalized electrons throughout (C) Strong single covalent bonds with weak intermolecular forces. (D) Strong multiple covalent bonds (including bonds.) with weak intermolecular forces (E) Macromolecules held tgether with strong polar bonds.
Magnesium
Mg
magnesium phosphide
Mg3P2
Manganese
Mn
Manganese
Mn +2 and +3
Permanganate
MnO4 -1
permanganate
MnO₄⁻
Molybdenum
Mo
A
Molecules that have planar configurations include which of the following? I. BCl3 II. CHCl3 III. NCl3 (A) I only (B) III only (C) I and II only (D) II and III only (E) I, II, and III
Molarity
Moles solute/ Liter solution
molality
Moles solute/Kg solvent
Nitrogen
N
nitrogen trichloride
NCl₃
Ammonium
NH4 +1
ammonium acetate
NH4C2H3O2
ammonium hydrogen sulfate
NH4HSO4
ammonia
NH₃
ammonium
NH₄⁺
Nitrite
NO2 -1
Nitrate
NO3 -1
nitrite
NO₂⁻
nitrate
NO₃⁻
Sodium
Na
sodium oxide
Na2O
sodium peroxide
Na2O2
sodium hydride
NaH
sodium dihydrogen phosphate
NaH2PO4
sodium peroxide
Na₂O₂
Neon
Ne
Nickel
Ni
Nickle
Ni +2
Ni⁺²
Nickel (II) or Nickelous
Ni⁺³
Nickel (III) or Nickelic
NO₃⁻¹
Nitrate
N⁻³
Nitride
NO₂⁻¹
Nitrite
Define noble gas.
Noble gases are in the last group of the periodic table (Group VIIIA or 18). - They are unusually unreactive with an octet of valence electrons ns²np⁶. - Helium was first found in the sun's spectrum. - Xenon was the first noble gas that was made into a compound.
The van't Hoff factor is 1 for?
Nonelectrolytes, Not ionic
Neptunium
Np
dinitrogen tetrahydride
N₂H₄
dinitrogen trioxide
N₂O₃
Oxygen
O
E
O || CH3-C-CH2-CH3 The organic compound represented above is an example of (A) an organic acid (B) an alcohol (C) an ether (D) an aldehyde (E) a ketone
Peroxide
O2 -2
Hydroxide
OH -1
hydroxide
OH⁻
D
Of the following molecules, which has the largest dipole moment? (A) CO (B) CO2 (C) O2 (D) HF (E) F2
C
On a mountaintop, it is observed that water boils at 90°C, not at 100°C as at sea level. This phenomenon occurs because on the mountaintop the A) equilibrium water vapor pressure is higher due to the higher atmospheric pressure B) equilibrium water vapor pressure is lower due to the higher atmospheric pressure C) equilibrium water vapor pressure equals the atmospheric pressure at a lower temperature D) water molecules have a higher average kinetic energy due to the lower atmospheric pressure E) water contains a greater concentration of dissolved gases
C₂O₄⁻²
Oxalate
O⁻²
Oxide
Phosphorus
P
Gay-Lussac's Law
P/T = k P₁/T₁ = P₂/T₂
diphosphorus pentoxide
P2O5
PyroPhosphate
P2O7
phosphorus pentachloride
PCl₅
Phosphite
PO3 -3
Phosphate
PO4 -3
phosphite
PO₃³⁻
phosphate
PO₄³⁻
Boyle's Law
PV = k P₁V₁ = P₂V₂
Ideal Gas Law
PV=nRT
Lead
Pb
Lead
Pb +2 and +4
lead (IV) sulfide
PbS2
Palladium
Pd
BrO₄⁻¹
Perbromate
ClO₄⁻¹
Perchlorate
FO₄⁻¹
Perfluorate
IO₄⁻¹
Periodate
MnO₄⁻¹
Permanganate
O₂⁻²
Peroxide
PO₄⁻³
Phosphate
P⁻³
Phosphide
Polonium
Po
K⁺¹
Potassium
A
Predicts that it is impossible to determine simultaneously the exact position and the exact velocity of an electron (A) Heisenberg uncertainty principle (B) Pauli exclusion principle (C) Hund's rule (principle of maximum multiplicity) (D) Shielding effect (E) Wave nature of matter
Factors that affect the formation of a solution
Pressure Temperature Nature of Solute Nature of Solvent
Platinum
Pt
Plutonium
Pu
Combined Gas Law
P₁V₁/T₁ = P₂V₂/T₂
Universal Gas Constant
R=0.0821 (L*atm)/(mol*K)
Radium
Ra
Ra⁺²
Radium
Rubidium
Rb
Radon
Rn
Rb⁺¹
Rubidium
Sulfur
S
spontaneous if:
S > 0, H < 0, G < 0
Thiocyanate
SCN -1
thiocyanate
SCN⁻
sulfur diflouride
SF2
sulfur hexaflouride
SF6
Sulfate
SO4 -2
sulfite
SO₃²⁻
sulfate
SO₄²⁻
Antimony
Sb
Scandium
Sc
Sc⁺³
Scandium
Selenium
Se
selenium tetrabromide
SeBr4
Selenate
SeO4 -2
selenate
SeO₄²⁻
Se⁻²
Selenide
Silicon
Si
Silicate
SiO3 -2
silicate
SiO₃²⁻
SiO₃⁻²
Silicate
Si⁻⁴
Silicide
Ag⁺¹
Silver
Tin
Sn
Tin
Sn +2 and +4
tin (II) flouride
SnF2
tin(II) phosphate
Sn₃(PO₄)₂
Na⁺¹
Sodium
Strontium
Sr
strontium flouride
SrF2
Sr⁺²
Strontium
SO₄⁻²
Sulfate
S⁻²
Sulfide
SO₃⁻²
Sulfite
2nd Law of Thermodynamics
Suni=Ssurr+Ssys = 0 for nonspontaneous Suni=Ssurr+Ssys > 0 for spontaneous
Tellurium
Te
Thorium
Th
C
The Lewis dot structure of which of the following molecules shows only one unshared pair of valence electron? (A) Cl2 (B) N2 (C) NH3 (D) CCl4 (E) H2O2
D
The SbCl5 molecule has trigonal bipyramid structure. Therefore, the hybridization of Sb orbitals should be (A) sp2 (B) sp3 (C) dsp2 (D) dsp3 (E) d2sp3
Boiling-Point Elevation
The boiling point is always higher for the solution than the pure solvent
C
The cooling curve for a pure substance as it changes from a liquid to a solid is shown right. The solid and the liquid coexist at (A) point Q only (B) point R only (C) all points on the curve between Q and S (D) all points on the curve between R and T (E) no point on the curve P\ R______S \Q/ \T
A
The electron-dot structure (Lewis structure) for which of the following molecules would have two unshared pairs of electrons on the central atom? (A) H2S (B) NH3 (C) CH4 (D) HCN (E) CO2
E
The energy change that occurs in the conversion of an ionic solid to widely separated gaseous ions (A) Activation energy (B) Free energy (C) Ionization energy (D) Kinetic energy (E) Lattice energy
C
The energy required to convert a ground-state atom in the gas phase to a gaseous positive ion (A) Activation energy (B) Free energy (C) Ionization energy (D) Kinetic energy (E) Lattice energy
Freezing-Point Depression
The freezing point is always lower for a solution of a nonvolatile solute than the pure solvent
E
The ground-state configuration for the atoms of a transition element (A) 1s2 2s2 2p5 3s2 3p5 (B) 1s2 2s2 2p6 3s2 3p6 (C) 1s2 2s2 2p6 2d10 3s2 3p6 (D) 1s2 2s2 2p6 3s2 3p6 3d5 (E) 1s2 2s2 2p6 3s2 3p6 3d3 4s2
B
The ground-state configuration of a common ion of an alkaline earth element (A) 1s2 2s2 2p5 3s2 3p5 (B) 1s2 2s2 2p6 3s2 3p6 (C) 1s2 2s2 2p6 2d10 3s2 3p6 (D) 1s2 2s2 2p6 3s2 3p6 3d5 (E) 1s2 2s2 2p6 3s2 3p6 3d3 4s2
B
The ground-state configuration of a negative ion of a halogen (A) 1s2 2s2 2p5 3s2 3p5 (B) 1s2 2s2 2p6 3s2 3p6 (C) 1s2 2s2 2p6 2d10 3s2 3p6 (D) 1s2 2s2 2p6 3s2 3p6 3d5 (E) 1s2 2s2 2p6 3s2 3p6 3d3 4s2
Raoult's Law
The higher the concentration of solute particles, the less the solvent is at the interface and the lower the vapor pressure
C
The ionization energies for element X are listed in the table above. On the basis of the data, element X is most likely to be (A) Na (B) Mg (C) Al (D) Si (E) P
B
The mass of element Q found in 1.00 mole of each of four different compounds is 38.0 grams, 57.0 grams, 76.0 grams, and 114 grams, respectively. A possible atomic weight of Q is (A) 12.7 (B) 19.0 (C) 27.5 (D) 38.0 (E) 57.0
B
The melting point of MgO is higher than that of NaF. Explanations for this observation include which of the following? I. Mg2+ is more positively charged than Na+ II. O2¯ is more negatively charged than F¯ III. The O2¯ ion is smaller than the F¯ ion (A) II only (B) I and II only (C) I and III only (D) II and III only (E) I, II, and III
E
The metal calcium reacts with molecular hydrogen to form a compound. All of the following statements concerning this compound are true EXCEPT: (A) Its formula is CaH2. (B) It is ionic. (C) It is solid at room temperatur(E) (D) When added to water, it reacts to produce H2 gas. (E) When added to water, it forms an acidic solution.
B
The phase diagram above provides sufficient information for determining the (A) entropy change on vaporization (B) conditions necessary for sublimation (C) deviations from ideal gas behavior of the gas phase (D) latent heat of vaporization (E) latent heat of fusion
1st Law of Thermodynamics
The principle of the conservation of energy. Energy can be transferred and transformed, but it cannot be created or destroyed.
D
The simplest formula for an oxide of nitrogen that is 36.8 percent nitrogen by weight is (A) N2O (B) NO (C) NO2 (D) N2O3 (E) N2O5
A
The structural isomers C2H5OH and CH3OCH3 would be expected to have the same values for which of the following? (Assume ideal behavior.) (A) Gaseous densities at the same temperature and pressure (B) Vapor pressures at the same temperature (C) Boiling points (D) Melting points (E) Heats of vaporization
C
The system shown above is at equilibrium at 28 °C. At this temperature, the vapor pressure of water is 28 millimeters of mercury. The partial pressure of O2(g) in the system is (A) 28 mm Hg (B) 56 mm Hg (C) 133 mm Hg (D) 161 mm Hg (E) 189 mm Hg
D
The volume of distilled water that should be added to 10.0 mL of 6.00 M HCl(aq) in order to prepare a 0.500 M HCl(aq) solution is approximately A) 50.0 mL B) 60.0 mL C) 100. mL D) 110. mL E) 120. mL
C
The weight of H2SO4 (molecular weight 98.1) in 50.0 milliliters of a 6.00-molar solution is (A) 3.10 grams (B) 12.0 grams (C) 29.4 grams (D) 294 grams (E) 300. grams
NCS⁻¹
Thiocyanate
S₂O₃⁻²
Thiosulfate
Titanium
Ti
Sn⁺²
Tin(II) or Stannous
Sn⁺⁴
Tin(IV) or Stannic
Thallium
Tl
D
Types of hybridization exhibited by the C atoms in propene, CH3CHCH2, include which of the following? I. sp II. sp2 III. sp3 (A) I only (B) III only (C) I and II only (D) II and III only (E) I, II, and III
Uranium
U
C
Utilized as a coating to protect Fe from corrosion (A) Pb (B) Ca (C) Zn (D) As (E) Na
A
Utilized as a shield from sources of radiation (A) Pb (B) Ca (C) Zn (D) As (E) Na
Vanadium
V
Charles's Law
V/T = k V₁/T₁ = V₂/T₂
Volume/Volume Percentage
Volume Solute/Volume Solvent
Tungsten
W
A
What is the most electronegative element? (A) O (B) La (C) Rb (D) Mg (E) N
B
When a hydrate of Na2CO3 is heated until all the water is removed, it loses 54.3 percent of its mass. The formula of the hydrate is (A) Na2CO3 . 10 H2O (B) Na2CO3 . 7 H2O (C) Na2CO3 . 5 H2O (D) Na2CO3 . 3 H2O (E) Na2CO3 . H2O
B
When a solution of sodium chloride is vaporized in a flame, the color of the flame is (A) blue (B) yellow (C) green (D) violet (E) White
E
Which element exhibits the greatest number of different oxidation states? (A) O (B) La (C) Rb (D) Mg (E) N
D
Which of the elements above has the smallest ionic radius for its most commonly found ion? (A) O (B) La (C) Rb (D) Mg (E) N
E
Which of the following acids can be oxidized to form a stronger acid? (A) H3PO4 (B) HNO3 (C) H2CO3 (D) H3BO3 (E) H2SO3
C
Which of the following actions would be likely to change the boiling point of a sample of a pure liquid in an open container? I. Placing it in a smaller container II. Increasing the number of moles of the liquid in the container III. Moving the container and liquid to a higher altitude (A) I only (B) II only (C) III only (D) II and III only (E) I, II, and III
E
Which of the following compounds is ionic and contains both sigma and pi covalent bonds? (A) Fe(OH)3 (B) HClO (C) H2S (D) NO2 (E) NaCN
A
Which of the following conclusions can be drawn from J. J. Thomson's cathode ray experiments? (A) Atoms contain electrons. (B) Practically all the mass of an atom is contained in its nucleus. (C) Atoms contain protons, neutrons, and electrons. (D) Atoms have a positively charged nucleus surrounded by an electron cloud. (E) No two electrons in one atom can have the same four quantum numbers.
A
Which of the following does NOT behave as an electrolyte when it is dissolved in water? (A) CH3OH (B) K2CO3 (C) NH4Br (D) HI (E) Sodium acetate, CH3COONa
A
Which of the following gases deviates most from ideal behavior? A) SO2 B) Ne C) CH4 D) N2 E) H2
D
Which of the following has the lowest conductivity? (A) 0.1 M CuS04 (B) 0.1 M KOH (C) 0.1 M BaCl2 (D) 0.1 M HF (E) 0.1 M HNO3
E
Which of the following is a correct interpretation of the results of Rutherford's experiments in which gold atoms were bombarded with alpha particles? (A) Atoms have equal numbers of positive and negative charges. (B) Electrons in atoms are agganged in shells. (C) Neutrons are at the center of an atom. (D) Neutrons and protrons in atoms have nearly equal mass. (E) The positive charge of an atom is concentrated in a small region.
C
Which of the following is lower for a 1.0-molar aqueous solution of any solute than it is for pure water? (A) pH (B) Vapor pressure (C) Freezing point (D) Electrical conductivity (E) Absorption of visible light
A
Which of the following is true at the triple point of a pure substance? (A) The vapor pressure of the solid phase always equal the vapor pressure of the liquid phase. (B) The temperature is always 0.01 K lower that the normal melting point. (C) The liquid and gas phases of the substance always have the same density and are therefore indistinguishable. (D) the solid phase always melts if the pressure increases at constant temperature. (E) The liquid phase always vaporizes if the pressure increases at constant temperature.
A
Which of the following molecules has a dipole moment of zero? (A) C6H6 (benzene) (B) NO (C) SO2 (D) NH3 (E) H2S
D
Which of the following pairs of liquids forms the solution that is most ideal (most closely follows Raoult's law)? A) C8H18(l) and H2O(l) B) CH3CH2CH2OH(l) and H2O(l) C) CH3CH2CH2OH(l) and C8H18(l) D) C6H14(l) and C8H18(l) E) H2SO4(l) and H2O(l)
C
Which of the following sets of quantum numbers (n, l, ml, ms) best describes the valence electron of highest energy in a ground-state gallium atom (atomic number 31) ? (A) 4, 0, 0, 1/2 (B) 4, 0, 1, 1/2 (C) 4, 1, 1, 1/2 (D) 4, 1, 2, 1/2 (E) 4, 2, 0, 1/2
E
Which of the following solids dissolves in water to form a colorless solution? (A) CrCl3 (B) FeCl3 (C) CoCl2 (D) CuCl2 (E) ZnCl2
Write the name for the formula: P₄O₁₀ Write the formula for the name: dinitrogen pentoxide
Write the formula for the name: tetraphosphorous decaoxide Write the name for the formula: N₂O₅
Write the name for the formula: AsF₃ Write the formula for the name: phosphorous pentafluoride
Write the formula for the name: arsenic trifluoride Write the name for the formula: PF₅
Write the name for the formula: BeCl₂ Write the formula for the name: arsenic trioxide
Write the formula for the name: beryllium dichloride Write the name for the formula: AsO₃
Write the name for the formula: CS₂ Write the formula for the name: boron trifluoride
Write the formula for the name: carbon disulfide Write the name for the formula: BF₃
Write the name for the formula: N₂O₄ Write the formula for the name: sulfur hexachlride
Write the formula for the name: dinitrogen tetroxide Write the name for the formula: SCl₆
Write the name for the formula: H₂O₂ Write the formula for the name: nitrogen monoxide
Write the formula for the name: hydrogen peroxide Write the name for the formula: NO
Write the name for the formula: H₂S(g) Write the formula for the name: antimony trichloride
Write the formula for the name: hydrogen sulfide Write the name for the formula: SbCl₃
Write the name for the formula: Nl₃ Write the formula for the name: carbon monoxide
Write the formula for the name: nitrogen triiodide Write the name for the formula: CO
Write the name for the formula: O₃ Write the formula for the name: nitrogen dioxide
Write the formula for the name: ozone Write the name for the formula: NO₂
Write the name for the formula: SiC Write the formula for the name: carbon tetrachloride
Write the formula for the name: silicon carbide Write the name for the formula: CCl₄
Write the name for the formula: SiO₂ Write the formula for the name: sulfur dioxide
Write the formula for the name: silicon dioxide Write the name for the formula: SO₂
Write the name for the formula: SO₃²⁻ Write the formula for the name: carbon dioxide
Write the formula for the name: sulfite ion Write the name for the formula: CO₂
Write the name for the formula: XeF₄ Write the formula for the name: sulfur trioxide
Write the formula for the name: xenon tetrafluoride Write the name for the formula: SO₃
E
X = CH3-CH2-CH2-CH2-CH3 Y = CH3-CH2-CH2-CH2-OH Z = HO-CH2-CH2-CH2-OH Based on concepts of polarity and hydrogen bonding, which of the following sequences correctly lists the compounds above in the order of their increasing solubility in water? (A) Z < Y < X (B) Y < Z < X (C) Y < X < Z (D) X < Z < Y (E) X < Y < Z
Xenon
Xe
Zn⁺²
Zinc
Zinc
Zn
Zinc
Zn +2
zinc sulfide
ZnS
Zirconium
Zr
D
__ Cr2O72¯ + __ e¯ + __ H+ ---> __ Cr3+ + __ H2O(l) When the equation for the half reaction above is balanced with the lowest whole-number coefficients, the coefficient for H2O is (A) 2 (B) 4 (C) 6 (D) 7 (E) 14
weak base
a base that reacts with water to produce hydroxide ions to only a slight extent in aqueous solution
indicator
a chemical that changes color and is used to mark the end point of a titration
oxidation state
a concept that provides a way to keep track of electrons in oxidation--reduction reactions according to certain rules
reduction
a decrease in oxidation state ( a gain of electrons)
colloid
a homogeneous mixture in which the dispersed particles do not settle out.
if Kc is large:
a large amount of products are produced, reactions lies far to the RIGHT
weak electrolyte
a material which, when dissolved in water, gives a solution that conducts only a small electric current
strong electrolyte
a material, when dissolved in water, gives a solution that conducts an electric current very efficiently
strong base
a metal hydroxide salt that completely dissociates into its ions in water
polar molecule
a molecule that has a permanent dipole moment
volumetric analysis
a process involving titration of one solution with another
oxidizing agent (electron acceptor)
a reactant that accepts electrons from another reactant
reducing agent (electron donor)
a reactant that donates electrons to another substance to reduce the oxidation state of one of its atoms
precipitation reaction
a reaction in which an insoluble substance forms and separates from the solution
oxidation-reduction (redox) reaction
a reaction in which one or more electrons are transferred
aqueous solution
a solution in which water is the dissolving medium or solvent
supersaturated solution
a solution that contains MORE dissolved solute than a saturated solution under the same conditions
unsaturated solution
a solution that contains less solute than a saturated solution under the same conditions
saturated solution
a solution that contains the maximum amount of dissolved solute
Ideal solution
a solution that obeys Raoult's law. The solute and solvent infractions are very similiar
standard solution
a solution whose concentration is accurately known
solute
a substance dissolved in liquid to form a solution
acid
a substance that marks the end point of an acid-base titration by changing color
base
a substance that produces hydroxide ions in aqueous solution, a proton acceptor
nonelectrolyte
a substance that, when dissolved in water, gives a nonconducting solution
titration
a technique in which one solution is used to analyze another
concentration (le chat)
add more products, reaction shifts left. add more reactants, reaction shifts right.
boiling point elevation
add solute, higher boiling point elevation
freezing point depression
add solute, lower freezing point depression
vapor pressure
add solute, lower vapor pressure
solubility
amount of substance needed to form a saturated solution (with a specific amt of solvent at a specific temp)
strong acid
an acid that completely dissociates to produce an H+ ion and the conjugate base
weak acid
an acid that dissociates only slightly in aqueous solution
neutralization reaction
an acid--base reaction
net ionic equation
an equation for a reaction in solution, where strong electrolytes are written as ions, showing only those components that are directly involved in the chemical change
formula equation
an equation representing a reaction in solution showing the reactants and products in undissociated form, whether they are strong or weak electrolytes
complete ionic equation
an equation that shows all substances that are strong electrolytes as ions
oxidation
an increase in oxidation state (a loss of electrons)
Cr²⁺
blue
Cu²⁺
blue
V³⁺
blue
Mn⁴⁺
brown
Cu⁺
colorless
Sc³⁺
colorless
V⁴⁺
colorless
Zn²⁺
colorless
vapor pressure _________ as imf _________
decreases, increases
viscosity _________ as temperature _________
decreases, increases
colligative properties
def properties of a solution that depend on the number of solute particles present and not on the type
does not
dilution with water does or doesn't alter the numbers of moles of solute present
H > 0
endothermic (+)
H < 0
exothermic (-)
m=moles of a solute / liters of solution
formula for molarity
moles of solute after dilution=moles of solute before dilution
formula of moles of solute after dilution
Cr³⁺
green
Fe²⁺
green
Mn⁶⁺
green
Ni²⁺
green
V³⁺
green
strength of IMFs
h-bonds > ion-dipole > dipole-dipole > LD
how to make a supersaturated solution
heat the solvent add maximum amount of solute cool down the solution
______ the concentration, _______ the chance of collisions, _______ the reaction rate
higher, greater, greater
______ temp, ______ the number of collisions, _______ the chance of collisions at the reaction site
higher, higher, greater
heats of transition _______ as imf _________
increase, increases
viscosity _________ as complexity of molecules _______
increases, increase
surface tension _________ as imf _________
increases, increases
viscosity _________ as imf _________
increases, increases
AgCl₂
insoluble
BaSO₄
insoluble
CaSO₄
insoluble
Hg₂Cl₂
insoluble
PbCl₂
insoluble
PbSO₄
insoluble
SrSO₄
insoluble
carbonates (except group 1 and ammonium)
insoluble
hydroxides (except group 1 and ammonium)
insoluble
oxides(except group 1 and ammonium)
insoluble
phosphates (except group 1 and ammonium)
insoluble
sulfides (except group 1,2 and ammonium)
insoluble
oxides (O-2)
insoluble; except group 1A and group IIA which reacts with H2O
hydroxides (OH-)
insoluble; except group 1A, Ca+2, Ba+2
carbonates (CO3-2)
insoluble; except group 1A, NH4+
phosphates (PO4-3)
insoluble; except group 1A, NH4+
sulfides (S-)
insoluble; except group 1A, group IIA, and NH4+
spectator ions
ions present in solution that do not participate directly in a reaction
nature of reactions
large, complex molecules react slower due to greater chance for collisions
immiscible
liquids that ARE NOT soluble in each other
miscible
liquids that dissolve freely in one another in any proportion
Molarity (M)
moles of solute per liter of solution M = mol / L
molarity
moles of solute per volume of solution in liters
soluble
most nitrate( NP3) salts are_______
Avogadro's Law
n/V = k n₁/V₁ = n₂/V₂
G > 0
non spontaneous
if Kc is small:
not much product is produced, reaction lies far to the LEFT
dissolve many substance
one of the most important properties of water is that it
Cr₂O₇²⁻
orange
solution equilibrium
physical state in which the opposing processes of dissolution and crystallization of a solute occur at equal rates
how to mix a proper molar concentration solution
pick the size of volumetric flask to fit the solution sample size fill about half full of distilled water add the calculated amount of solute and mix until dissolved add distilled water to the fill line
Co²⁺
pink
Mn²⁺
pink
Henry's law
solubility of a gas is DIRECTLY PROPORTIONAL to the partial pressure of the gas above the liquid
Group 1 hydroxides (stong bases)
soluble
Group 1A
soluble
HBr (strong acid)
soluble
HCl (stong acid)
soluble
HClO₃(strong acid)
soluble
HClO₄(strong acid)
soluble
HI (strong acid)
soluble
HNO₃(strong acid)
soluble
H₂SO₄(strong acid)
soluble
NH4+
soluble
acetates (C₂H₃0₂)
soluble
anythiung compound with ammonium (NH₄⁺)
soluble
group 1 compouds
soluble
group 2 hydroxides below Mg
soluble
most chlorides
soluble
most sulfates
soluble
nitrates (NO₃⁻)
soluble
NO3-
soluble, except when paired with Cu+, Ag+, Pb+, Hg2+2
sulfate (SO4-)
soluble; except Pb+2, Ca+2, Sr+2, Ba+2
acetate (C2H3O2-)
soluble; except when paired with Cu+, Ag+, Pb+, Hg2+2
perchlorate (ClO4-)
soluble; except when paired with Cu+, Ag+, Pb+, Hg2+2
hydration
solution process with water as teh solvent
catalyst
speeds up a reaction by lowering the activation energy
G < 0
spontaneous
HClO3
strong acid; chloric
HBr
strong acid; hydrobromic
HCl
strong acid; hydrochloric
HI
strong acid; hydroiodic
HNO3
strong acid; nitric
HClO4
strong acid; perchloric
H2SO4
strong acid; sulfuric
Ba(OH)2
strong base
Ca(OH)2
strong base
Sr(OH)2
strong base
group IA hydroxides
strong bases (Li - Cs)
electrical conductivity
the ability to conduct an electric current
solubility
the amount of a substance that dissolves in a given volume of solvent at a given temperature
Arrhenius
the basis for the conductivity properties of a solution was first correctly identified by
solvent
the dissolving medium in a solution
osmotic pressure
the greater the difference in dissolved particles (across a membrane) the greater the osmotic pressure
boiling point elevation
the higher the number of dissolved particles the higher the boiling point
freezing point depression
the higher the number of dissolved particles the lower the freezing point
vapor pressure lowering
the higher the number of dissolved particles the lower the vapor pressure
hydration
the interaction between solute particles and water molecules
ammonia
the most common weak base is
water
the most important substances on earth is
endpoint
the point in a titration at which the indicator changes color
stoichiometric (equivalence) point
the point in a titration when enough titrant has been added to react exactly with the substance in solution being titrated
dilution
the process of adding solvent to lower the concentration of solute in a solution
precipitate
the solid formed in the precipitation reaction
half-reactions
the two parts of an oxidation--reduction reaction, one representing oxidation, the other reduction
Mn⁷⁺
violet
CH3COOH
weak acid; acetic
NH4+
weak acid; ammonium
HCN
weak acid; hydrocyanic
HF
weak acid; hydrofluoric
H2S
weak acid; hydrogen sulfide
HS-
weak base; HS- ion
NH3
weak base; ammonia
cations, anions
when ionic substances dissolve in water they break up into individual _______ and _______
Co³⁺
yellow
CrO₄²⁻
yellow
Fe³⁺
yellow
V⁵⁺
yellow
