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Chemistry Chapter 4
Reactions in Solution: Aqueous Chemistry in Nature
| Question | Answer |
|---|---|
| Solvent (4.1) | The substance present in the greatest number of moles in a solution. |
| Solute (4.1) | All substances dissolved in a solvent. |
| Aqueous Solution (4.1) | A solution in which the solvent is water, H₂O. |
| Solid Solutions (4.1) | Uniform mixtures of solid substances. Often found in earth's crust. Conversely, gaseous solutions make up our atmosphere. |
| Acid-Base Reactions (4.1) | Reactions involving the transfer of H⁺ ions. |
| Precipitation Reactions (4.1) | Involves pairs of dissolved cations and anions that collide to form only slightly soluble ionic compounds that precipitate from (settle out of) solution. |
| Redox Reactions (4.1) | Elements are reduced, which means that their atoms gain electrons, while others are oxidized, which means that their atoms lose electrons. |
| Concentration of a Solution (4.2) | The amount of solute in a given amount of solvent or solution. Qualitatively, solutions are described as concentrated or dilute if they contain relatively high or low ratios of solute to solvent. |
| Mass-to-Volume Rations (4.2) | A way to describe the concentration of a solution (ex. milligrams of solute per 1 Liter of solution). |
| Molarity (4.2) | Expressed as (M) meaning moles of solute/liter of solution. |
| Millimolarity (4.2) | 1 mM = (M)10⁻³ |
| Micromolarity (4.2) | 1 μM = (M)10⁻⁶ |
| Nanomolarity (4.2) | 1 nM = (M)10⁻⁹ |
| Picomolarity (4.2) | 1 pM = (M)10⁻¹² |
| Mass-to-Mass Ratios (4.2) | Two commonly used mass-to-mass ratios are parts per million (ppm) and parts per billion (ppb). For example, there could be 1 part solute for every million parts of solution. A concentration of 1 ppb is 1/1000 as concentrated as 1 ppm. |
| Dilution (4.3) | The process of lowering the concentration of a solution by adding solvent to a known volume of the initial solution. This can be expressed like this: [Vᶠⁱⁿᵃˡ x Mᶠⁱⁿᵃˡ = Vⁱⁿⁱᵗⁱᵃˡ x Mⁱⁿⁱᵗⁱᵃˡ]. |
| Absorbance (4.3) | A measure of the intensity of the color. Expressed using the variable (A). |
| Beer's Law (4.3) | Relates absorbance to three quantities: concentration (c), the path length that the light travels through the solution (b), and a constant (called the molar absorptivity, or ϵ) characteristic of the dissolved solute. |
| Electrolytes (4.4) | Any solute that imparts electrical conductivity to an aqueous solution. Molecular compounds tend to not be conductive. Solutes that do not form ions when dissolved in water are called nonelectrolytes. |
| Strong Electrolytes (4.4) | Dissociates completely in water, which means it completely breaks up into its component ions. |
| Weak Electrolytes (4.4) | Only a small fraction of a weak electrolyte’s molecules dissociate. |
| (Brønsted-Lowry) Acid (4.5) | A proton donor. Usually this proton is represented as H⁺ (aq) which combines with a water molecule to create a H₃O⁺ hydronium ion. |
| (Brønsted-Lowry) Base (4.5) | A proton acceptor. Strong bases include the hydroxides of group 1 and group 2 metals, which dissociate completely when dissolved in water. (OH⁻) Hydroxide ions are bases because they can accept protons. This definition permits other substances that do not consist of hydroxide ions to be categorized as bases as well. |
| Neutralization Reaction (4.5) | The acid and the base no longer exist as a result of this reaction. Produces water and a salt, "using up" the remaining ions. |
| Salt (4.5) | A substance formed along with water as a product of a neutralization reaction. |
| Molecular Equation (4.5) | Each reactant and product is written as a neutral compound (no actual species). |
| Overall Ionic Equation (4.5) | Shows the reactants and products as the particles that exist in solution, which means that any substance that dissociates completely is written as individual ions. Distinguishes soluble ionic substances from molecular substances within the solution. Any nonelectrolytes or weak electrolytes are written as neutral molecules. |
| Net ionic Equation (4.5) | Removes the terms that are equal on both sides, simplifying the equation to display only the net change. |
| Spectator Ions (4.5) | The ions that are removed when an overall ionic equation is simplified into a net ionic equation. These ions do not react in any significant way. |
| Hydrolysis (4.5) | The reaction of water with another material. The chemical breakdown of a compound due to reaction with water. |
| Monoprotic & Diprotic & Polyprotic Acids (4.5) | One molecule of a monoprotic acid donates one proton, while a diprotic molecule donates two. Polyprotic acids are usually much weaker. |
| Amines (4.5) | Characterized by having one or more of the hydrogen atoms in NH₃ replaced by an organic group. These are typically weak bases despite containing no hydroxide groups. |
| Amphiprotic Substances (4.5) | Can function as a proton acceptor (Brønsted–Lowry base) in solutions containing acid solutes or as a proton donor (Brønsted–Lowry acid) in solutions containing basic solutes. Water is an amphiprotic substance. |
| Titrations (4.6) | A way to use the reactions of acids and bases to analyze solutions and determine the concentrations of dissolved substances. A standard solution, meaning its concentration is known accurately, is called a titrant. |
| Equivalence Point (4.6) | The point in the titration when just enough standard solution has been added to completely react with all the solute in the sample. |
| End Point (4.6) | The point at which the indicator changes color. |
| Precipitate (4.7) | A solid that forms due to an ionic reaction. The liquid remaining is referred to as the "supernatant" liquid. |
| How to Determine When a Precipitate Forms (4.7) | 1. Recognize if they are both salts, which means that they are both soluble in water and therefore have dissociated into their respective cations and anions. 2. Determine whether any of the possible combinations of ions produces a slightly soluble product. When these ions collide with one another, do any of them form a slightly soluble ionic compound, that is, a precipitate? |
| Saturated Solution (4.7) | A solution containing the maximum amount of solute that can dissolve. Solubility depends on temperature; often, the higher the temperature, the greater the solubility of solids in water. |
| Supersaturated Solution (4.7) | Occasionally, more solute dissolves in a volume of liquid than the amount predicted by the solute’s solubility in that liquid. |
| Oxidation (4.8) | A reaction that increased the oxygen content of a substance. |
| Reduction (4.8) | Reactions in which the oxygen content of a substance is reduced. |
| Oxidation Number (O.N.) or Oxidation State (4.8) | In monatomic ions, the oxidation number is the charge on the ion. The oxidation numbers of the atoms in a neutral molecule sum to zero; those of the atoms in an ion sum to the charge on the ion. In compounds containing fluorine and one or more other elements, the oxidation number of the fluorine is always −1. In most compounds, the oxidation number of hydrogen is +1 and that of oxygen is −2. Unless combined with oxygen or fluorine, chlorine and bromine have an oxidation number of −1. |
| Using the Terms Oxidated and Reduced (4.8) | Because the oxidation number of sodium increases, sodium is said to be oxidized in this reaction. Likewise, because the oxidation number of chlorine decreases, chlorine is said to be reduced. |
| Oxidizing Agent (4.8) | A substance that contains the element being reduced. The oxidizingpage188 agent in any redox reaction is always reduced in the process. Any species that is reduced experiences a reduction in oxidation number, as is the case for the silver ion, which goes from [O.N. = +1] to [O.N. = 0]. |
| Reducing Agent (4.8) | The reducing agent is always oxidized and experiences an increase in oxidation number; in this reaction, the copper atom goes from O.N. = 0 to O.N. = +2. Every redox reaction has both an oxidizing agent and a reducing agent. |
| Missing Sections (4.8) | Half-Reactions, The Activity Series of Metals, Redox in Nature. |
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