Chemistry Test Two Comprehensive Guide
Chemistry Test Two Comprehensive Guide 14054
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This 22 page Study Guide was uploaded by Kazendi Simon on Saturday March 5, 2016. The Study Guide belongs to 14054 at Georgia State University taught by Dr. Fernando in Spring 2016. Since its upload, it has received 77 views. For similar materials see Principles of Chemistry 1 in Chemistry at Georgia State University.
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Chemistry Test Two Comprehensive Guide Chapter 3: Molecules, Compounds, and Chemical Equations Key Vocabulary Terms: Compounds: Are composed of atoms held together by chemical bonds. Chemical bonds: Result from the attractions between the ‘charged’ particles (the electrons and protons) that compose atoms; there are two types: Ionic and Covalent Bonds. Ionic bonds: Occur between metals and nonmetals involve the transfer of electrons from one atom to another. Oppositely charged ions attract one another by electrostatic forces and form an ionic bond. i. When a metal interacts with a nonmetal, it can transfer one or more of its electrons to the nonmetal. ii. The metal atom then becomes a cation. iii. The nonmetal atom becomes an anion. Covalent bonds: Occur between two or more nonmetals involve the sharing of electrons between two atoms. i. When a nonmetal bonds with another nonmetal, neither atom transfers its electron to the other; the bonding atoms share some of their electrons. ii. The covalently bound atoms (molecular compounds) compose a molecule, a free unit which represents the actual number of atoms (molecular formula). Chemical formula: Indicates the elements present in the compound and the relative number of atoms or ions of each. i. Water is represented as H O. 2 Types of Chemical Formulas & Models The type of formula we use depends on how much we know about the compound and how much we want to communicate. Chemical formulas can generally be categorized into three different types: o Empirical formula: Gives the relative number of atoms of each element in a compound. For C 4 ,8the greatest common factor is 4. The empirical formula is therefore CH . 2 Empirical formula communicates the least amount of information. o Molecular formula: Gives the actual number of atoms of each element in a molecule of a compound. For CCl ,4the only common factor is 1, so the empirical formula and the molecular formula are identical. o Structural formula: Uses lines to represent covalent bonds and shows how atoms in a molecule are connected or bonded to each other. A structural formula communicates the most information Molecular Models: Is a more accurate and complete way to specify a compound. BallandStick Molecular Model: Represents atoms as balls and chemical bonds as sticks; how the two connect reflects a molecule’s shape. o The balls are typically color coded to specific elements. SpaceFilling Molecular Model: Atoms fill the space between each other to more closely represent our best estimates for how a molecule might appear if scaled to visible size. Ionic Bonds & Compounds Ionic Compounds in the solid phase is composed of a lattice (a regular threedimensional array of alternating cations and anions). o Ionic compounds do not exist as free units, due to nondirectional nature of the ionic bond, any one ion is surrounded by 4, 6 or 8 oppositely charged ions. o Ionic pattern repeats itself in all dimensions forming a crystal so the formula cannot represent the actual number of combining ions and we simply represent the ratio of ions as empirical formula. Ionic Compound Composition: Ionic compounds are composed of cations (usually a metal) and anions (usually one or more nonmetals) bound together by ionic bonds. o The basic unit of an ionic compound is the formula unit, the smallest, electrically neutral collection of ions. o The ionic compound table salt, with the formula unit NaCl, is composed of Na+ and Cl– ions in a onetoone ratio. Polyatomic Ions: Many common ionic compounds contain ions that are themselves composed of a group of covalently bonded atoms with an overall charge. o Polyatomic molecules: Phosphorus [P], Sulfur [S], and Selenium [Se]. o NaNO con3ains Na+ and NO . 3 Molecular Elements Elements that exist as diatomic molecules: Hydrogen [H], Nitrogen [N], Oxygen [O], Fluoride [F], Chloride [Cl], Bromide [Br], and Iodine [I]. o I Never Bring Classwork Home On Fridays. Ionic Compounds: Formulas and Names: Ionic compounds always contain positive and negative ions. In a chemical formula, the sum of the charges of the positive ions (cations) must equal the sum of the charges of the negative ions (anions). The formula of an ionic compound reflects the smallest wholenumber ratio of ions. The charges of the representative elements can be predicted from their group numbers. o Lithium [Li] is in group one (Akali Metals) has a charge of positive 1. o The representative elements forms only one type of charge. Transition metals tend to form multiple types of charges and can’t be predicted like the representative elements. Naming Binary Ionic Compounds of Type I Cations: Binary Compounds: Contain only two different elements. The names of binary ionic compounds take the following form: o First: Name of cation (metal) Second: Base name of anion (nonmetal) + ide o For example, the name for KCl consists of the name of the cation, potassium, followed by the base name of the anion, chlor, with the ending ide. o KCl is potassium chloride. Table 3.2 Metals Whose Charge is Invariant from One Compound to Another Metal Ion Name Group Number Li Li+ Lithium 1A + Na Na Sodium 1A K K+ Potassium 1A + Rb Rb Rudidum 1A Cs Cs + Cesium 1A +2 Be Be Beryllium 2A Mg Mg +2 Magnesium 2A +2 Ca Ca Calcium 2A Sr Sr +3 Strontium 2A +3 Ba Ba Barium 2A Al Al +3 Aluminum 3A +2 Zn Zn Zinc * Sc Sc+ 3 Scandium * + Ag ** Ag Silver * * The charge of these metals cannot be inferred from their group number ** Silver sometimes forms compounds with other charges, but these are rare. Table 3.3 Some Common Monoatomic Anions Nonmetal Symbol for Ion Base Name Anion Name Fluorine F Fluor Fluoride Chlorine Cl Chlor Chloride Bromine Br - Brom Bromide Iodine I Iod Iodide Oxygen O 2 Ox Oxide 2 Sulfur S Sulf Sulfide Nitrogen N 3 Nitr Nitrogen 3 Phosphorus P Phosph Phosphide Table 3.5 Some Common Polyatomic Ions Name Formula Name Formula Acetate C 2 3 2 Hypochlorite ClO 2 Carbonate CO 3 Chlorite ClO 2 Hydrogen Carbonate HCO 3 Chlorate ClO 3 (Bicarbonate) Hydroxide HO Perchlorate ClO 4 Nitrite NO 2 Permanganate MnO 4 Nitrate NO Sulfite SO 2 3 3 Hydrogen Sulfite 2 Chromate CrO 4 (Bisulfite) HSO 3 2 2 Dichromate Cr2O 7 Sulfate SO 4 Hydrogen Sulfate Phosphate PO 43 (Bisulfate) HSO 4 2 Hydrogen Phosphate HPO 4 Cyanide CN Dihydrogen Phosphate H 2O 32 Peroxide O 22 + + Ammonium NH 4 Hydronium H 3 Oxyanions: Most polyatomic ions are oxyanions, anions containing oxygen and another element. o Notice that when a series of oxyanions contains different numbers of oxygen atoms, they are named according to the number of oxygen atoms in the ion. If there are two ions in the series; the one with more oxygen atoms has the ending ate, and the one with fewer has the ending ite. o NO : i3 nitrate o NO : i2 nitrite Oxyanions that have more than two ions in the series then the prefixes hypo, meaning less than, and per, meaning more than, are used. ClO: is hypochlorite BrO: is hypobromite ClO 2 s chlorite BrO 2 is bromite ClO 3 chlorate BrO : is bromate 3 ClO 4 is perchlorate BrO 4 is perbromate Naming Ionic Compounds Containing Polyatomic Ions: We name ionic compounds that contain a polyatomic ion in the same way as other ionic compounds, except that we use the name of the polyatomic ion whenever it occurs. o For example: NaNO is na2ed according to its cation, Na sodium, and its polyatomic anion, NO , n2trite. Creating NaNO is2sodium nitrite. Type II Ionic Compounds This is where it gets a bit complex, some metals mainly the transition metals can form more than one type of cation. Some main group metals, such as Lead [Pb], Thallium [Tl], and Tin [Sn], form more than one type of cation. Naming Type II Binary Ionic Compounds: The full names for compounds containing metals that form more than one kind of cation have the following form: o Name of cation (Metal) > (charge of cation (metal) in roman numerals in parenthesis) > Base name of anion (nonmetal) + ide o The charge of the metal cation can be determined by inference from the sum of the charges of the nonmetal. Hydrated Ionic Compouds: Hydrates are ionic compounds containing a specific number of water molecules associated with each formula unit. o For example, the formula for epsom salts is MgSO • 7H 4 2 Its systematic name is magnesium sulfate heptahydrate o Basically the number in front of the Hydrogen take the prefix of the Ionic Compound. Common hydrate prefixes o Hemi = ½ o Mono = 1 o Di = 2 o Tri = 3 o Tetra = 4 o Penta = 5 o Hexa = 6 o Hepta = 7 o Octa = 8 o Nona = 9 o Deca = 10 Molecular Compounds: Formulas and Names The formula for a molecular compound cannot readily be determined from its constituent elements because the same combination of elements may form many different molecular compounds, each with a different formula. o Nitrogen and oxygen form all of the following unique molecular compounds: NO, NO2, N2O, N2O3, N2O4, and N2O5. o Molecular compounds are composed of two or more nonmetals. o Generally, write the name of the element with he smallest group number first. If the two elements lie in the same group, then write the element with the greatest row number first. o The prefixes given to each element indicate the number of atoms present. These prefixes are the same as those used in naming hydrates. Acids: Acids are molecular compounds that release hydrogen ions (H+) when dissolved in water. Acids are composed of hydrogen, usually written first in their formula, and one or more nonmetals, written second. – HCl is a molecular compound that, when dissolved in water, forms H+ (aq) and Cl– (aq) ions, where aqueous (aq) means dissolved in water. Acids • Acids are molecular compounds that form H+ when dissolved in water. – To indicate the compound is dissolved in water (aq) is written after the formula. » A compound is not considered an acid if it does not dissolve in water. • Sour taste • Dissolve many metals – such as Zn, Fe, Mg; but not Au, Ag, Pt • Formula generally starts with H – e.g., HCl, H2SO4 Chemistry Test Two Comprehensive Guide Chapter 4: Chemical Quantities and Aqueous Reactions Key Vocabulary Terms: Law of Conservation of Mass: states that matter can be changed from one form into another, mixtures can be separated or made, and pure substances can be decomposed, but the total amount of mass remains constant. o Balancing equations by balancing atoms The study of the numerical relationship between chemical quantities in a chemical reaction is called stoichiometry. The reactant that limits the amount of product is called the limiting reactant. Reactants not completely consumed are called excess reactants. The theoretical yield is the amount of product that can be made in a chemical reaction based on the amount of limiting reactant. The actual yield is the amount of product actually produced by a chemical reaction. The percent yield is calculated as follows: actual yield o theoreticalyield x100 Reaction Stoichiometry: The coefficients in a chemical reaction specify the relative amounts in moles of each of the substances involved in the reaction. 2 molecules of C H8 r18ct with 25 molecules of O to 2orm 16 molecules of CO and 18 2 molecules of H O2 o 2C H 8l)18 25O (g) 2 16CO (g) + 28H O(g) 2 The ratio of the coefficients acts as a conversion factor between the amount in of the reactants and products. Example: Suppose That We Burn 22.0 Moles of C H ; how8Ma18 Moles of CO Form? 2 o 2C H (l) + 25O (g) 16CO (g) + 18H O(g) 8 18 2 2 2 16molCO 2 o 22.0molC H 8 18 =176molCO 2 2molC H 8 18 o The combustion of 22 moles of C H ad8s 186 moles of CO to the at2osphere. Solutions: When table salt is mixed with water, it seems to disappear or become a liquid, the mixture is homogeneous. o Homogeneous mixtures are called solutions. The component of the solution that changes state is called the solute. The component that keeps its state is called the solvent. If both components start in the same state, the major component is the solvent. Solution Concentration: Solutions are often described quantitatively, as dilute or concentrated. o Dilute solutions have a small amount of solute compared to solvent. o Concentrated solutions have a large amount of solute compared to solvent. Solution Concentration: Molarity: A common way to express solution concentration is molarity (M). o Molarity is the amount of solute (in moles) divided by the volume of solution (in liters). amountof solute(¿mol) o Molarity (M = volumeof solution(¿L) Remembershe prefers¿usemLsobeready¿convert o Solution Dilution: Often, solutions are stored as concentrated stock solutions. To make solutions of lower concentrations from these stock solutions, more solvent is added. o The amount of solute doesn’t change, just the volume of solution: The concentrations and volumes of the stock and new solutions are inversely proportional: o Moles solute in solution I = moles solute in solution II o M x1V =1 xV 2 2 Electrolyte and Nonelectrolyte Solutions: Materials that dissolve in water to form a solution that will conduct electricity are called electrolytes. Materials that dissolve in water to form a solution that will not conduct electricity are called nonelectrolytes. Ionic substances such as sodium chloride that completely dissociate into ions when they dissolve in water are strong electrolytes. o In contrast to sodium chloride, sugar is a molecular compound. o Most molecular compounds (except for acids), dissolve in water as intact molecules. Binary Acids: Acids are molecular compounds that ionize when they dissolve in water. o The molecules are pulled apart by their attraction for the water. o When acids ionize, they form H+ cations and also anions. o The percentage of molecules that ionize varies from one acid to another. Acids that ionize virtually 100% are called strong acids. + − HCl (aq) H (aq) + Cl (aq) Acids that only ionize a small percentage are called weak acids. HF (aq) ↔ H (aq) + F (aq) Strong and Weak Electrolytes: Strong electrolytes are materials that dissolve completely as ions. o Ionic compounds and strong acids o Solutions conduct electricity well Weak electrolytes are materials that dissolve mostly as molecules, but partially as ions. o Weak acids o Solutions conduct electricity, but not well o When compounds containing a polyatomic ion dissolve, the polyatomic ion stays together. + − HC H2O 3(a2) H (aq) + C H O 2(a3)2 Dissociation and Ionization When ionic compounds dissolve in water, the anions and cations are separated from each other this is called dissociation. o Na S(aq) 2Na (aq) + S (aq)2– 2 When compounds containing polyatomic ions dissociate, the polyatomic group stays together as one ion. + 2− o Na SO2 (aq4 2Na (aq) + SO 4 (aq) When strong acids dissolve in water, the molecule ionizes into H+ and anions. + 2− H 2O 4aq) Na (aq) + SO 4 (aq) The Solubility of Ionic Compounds: When an ionic compound dissolves in water, the resulting solution contains: o Not the intact ionic compound itself, but its component ions dissolved in water. However, not all ionic compounds dissolve in water. o If we add AgCl to water, for example, it remains solid and appears as a white powder at the bottom of the water. In general, a compound is termed soluble if it dissolves in water and insoluble if it does not. Precipitation Reactions: Precipitation reactions are reactions in which a solid forms when we mix two solutions. o Reactions between aqueous solutions of ionic compounds produce an ionic compound that is insoluble in water. The insoluble product is called a precipitate. No Precipitation Means No Reaction: Precipitation reactions when two aqueous solutions are mixed. Predicting Precipitation Reactions 1. Determine what ions each aqueous reactant has. 2. Determine formulas of possible products. o Exchange ions. o (+) ion from one reactant with (–) ion from other. o Balance charges of combined ions to get the formula of each product. 3. Determine solubility of each product in water. Use the solubility rules. If product is insoluble or slightly soluble, it will precipitate. 4. If neither product will precipitate, write no reaction after the arrow. 5. If any of the possible products are insoluble, write their formulas as the products of the reaction using (s) after the formula to indicate solid. Write any soluble products with (aq) after the formula to indicate aqueous. 6. Balance the equation. 7. Remember to only change coefficients, not subscripts. Representing Aqueous Reactions: An equation showing the complete neutral formulas for each compound in the aqueous reaction as if they existed as molecules is called a molecular equation. o 2KOH(aq) + Mg(NO ) (aq) 2KNO (aq) + Mg(OH) (s) 3 2 3 2 In actual solutions of soluble ionic compounds, dissolved substances are present as ions. Equations that describe the material’s structure when dissolved are called complete ionic equations. Ionic Equation: Rules of writing the complete ionic equation: o Aqueous strong electrolytes are written as ions. Soluble salts, strong acids, strong bases o Insoluble substances, weak electrolytes, and nonelectrolytes are written in molecule form. Solids, liquids, and gases are not dissolved, hence molecule form Ionic Equation: Notice that in the complete ionic equation, some of the ions in solution appear unchanged on both sides of the equation. These ions are called spectator ions because they do not participate in the reaction. Basically spectator ions are the ions seen on both the reactant and yield side that is crossed out when finding the Net Ionic Equation. Acid–Base and GasEvolution Reactions Two other important classes of reactions that occur in aqueous solution are 1. Acid–base reactions 2. Gasevolution reactions. Acid–base Reaction: An acid reacts with a base and the two neutralize each other, producing water (or in some cases a weak electrolyte). o An acid–base reaction is also called a neutralization reaction. Acid–Base and GasEvolution Reactions: In a gasevolution reaction, a gas forms resulting in bubbling. In both acid–base and gasevolution reactions, as in precipitation reactions, the reactions occur when the anion from one reactant combines with the cation of the other. o Many gasevolution reactions are also acid–base reactions. Acids and Bases in Solution + Acids ionize in water to form H ions. o More precisely, the H from the acid molecule is donated to a water molecule to + form hydronium ion, H O 3 Bases dissociate in water to form OH ions. o Bases, such as NH that do not contain OH ions, produce OH by pulling H off 3 water molecules. In the reaction of an acid with a base, the H from the acid combines with the OH from the base to make water. The cation from the base combines with the anion from the acid to make the salt. GasEvolving Reactions Some reactions form a gas directly from the ion exchange. Other reactions form a gas by the decomposition of one of the ion exchange products into a gas and water. Oxidation–Reduction Reactions: The reactions in which electrons are transferred from one reactant to the other are called oxidationreduction reactions. o These are also called redox reactions. o Many redox reactions involve the reaction of a substance with oxygen. 4Fe(s) + 3O (2) 2Fe O 2s)3(rusting) 2C H (l) + 25 O2 (g) 16CO (g) + 18H O (g) (combustion) 8 18 2 2 2H 2g) + O 2g) 2H O(g2 Combustion as Redox Redox without Combustion: Reactions of Metals with Nonmetals: Consider the following reactions: o 4Na(s) + O (g2 → 2Na O(s) 2 o 2Na(s) + Cl (g2 → 2NaCl(s) The reactions involve a metal reacting with a nonmetal. In addition, both reactions involve the conversion of free elements into ions. Redox Reaction: The transfer of electrons does not need to be a complete transfer (as occurs in the formation of an ionic compound) for the reaction to qualify as oxidation–reduction. o For example, consider the reaction between hydrogen gas and chlorine gas: H 2(g) + Cl 2 (g) 2HCl (g) When hydrogen bonds to chlorine, the electrons are unevenly shared, resulting in an increase of electron density (reduction) for chlorine and a decrease in electron density (oxidation) for hydrogen. Oxidation and Reduction: To convert a free element into an ion, the atoms must gain or lose electrons. o Of course, if one atom loses electrons, another must accept them. Reactions where electrons are transferred from one atom to another are redox reactions. Atoms that lose electrons are being oxidized, while atoms that gain electrons are being reduced. o 2Na(s) + Cl (2) → 2Na Cl (s) – + – Na → Na + 1e (oxidation) Cl2 + 2e → 2 Cl (reduction) Oxidation States: For reactions that are not metal + nonmetal, or do not involve O , we need a method for 2 determining how the electrons are transferred (oxidation state the electron flow in the reaction). o Even though they look like them, oxidation states are not ion charges! Identifying Redox Reactions Oxidation: An increase in oxidation state Reduction: A decrease in oxidation state Carbon changes from an oxidation state of 0 to an oxidation state of +4. o Carbon loses electrons and is oxidized. Sulfur changes from an oxidation state of 0 to an oxidation state of –2. o Sulfur gains electrons and is reduced. Redox Reactions: Oxidation and reduction must occur simultaneously. o If an atom loses electrons another atom must take them. The reactant that reduces an element in another reactant is called the reducing agent. o The reducing agent contains the element that is oxidized. The reactant that oxidizes an element in another reactant is called the oxidizing agent. o The oxidizing agent contains the element that is reduced. + – o 2Na(s) + Cl (2) → 2 Na Cl (s) Na is oxidized, while Cl is reduced. Na is the reducing agent, and Cl 2 the oxidizing agent. Combustion Reactions: Combustion reactions are characterized by the reaction of a substance with to form one or more oxygencontaining compounds, often including water. o Combustion reactions also emit heat. For example, natural gas (CH )4reacts with oxygen to form carbon dioxide and water: Combustion: Ethanol, the alcohol in alcoholic beverages, also reacts with oxygen in a combustion reaction to form carbon dioxide and water.
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