chpt 4

Weak acid examples...

(a) fruits such as oranges, lemons, and grapefruit contain the weak acid citric acid (b) vinegars contain the weak acid acetic acid

The guidelines that are used to assign oxidation numbers to each element in a molecule or ion.

1. The oxidation number of an atom in an elemental substance is zero 2. The oxidation number of a monatomic ion is equal to the ion's charge 3. Oxidation numbers for common nonmetals are usually assigned as follows: -Hydrogen: +1 when combined with nonmetals, -1 when combined with metals -Oxygen: -2 in most compounds, sometimes -1 (so-called peroxides, O2 2-), very rarely -1/2 (so-called superoxides, O2-), positive values when combined with F (values vary) -Halogens: -1 for F always, -1 for other halogens except when combined with oxygen or other halogens (positive oxidation numbers in these cases, varying values) 4. The sum of oxidation numbers for all atoms in a molecule or polyatomic ion equals the charge on the molecule or ion Note: the proper convention for reporting charge is to write the number first, followed by the sign (e.g., 2+), while the oxidation number is written with the reversed sequence, sign followed by number (e.g., +2_. This convention aims to emphasize the distinction between these two related properties.

Fundamental aspects of chemical equation:

1. The substances undergoing reaction are called reactants, and their formulas are placed on the left side of the equation. 2. The substances generated by the reaction are called products, and their formulas are placed on the right side of the equation. 3. Plus signs (+) separate individual reactant and product formulas, and an arrow (--->) separates the reactant and product (left and right) sides of the equation 4. The relative numbers of reactant and product species are represented by coefficients (numbers placed immediately to the left of each formula). A coefficient of 1 is typically omitted.

balancing by inspection

A balanced chemical equation often may be derived from a qualitative description of some chemical reaction by a fairly simple approach known as balancing by inspection. Consider as an example the decomposition of water to yield molecular hydrogen and oxygen. This process is represented qualitatively by an unbalanced chemical equation: refer to pic

Precipitation reactions and solubility rules

A precipitation reaction is one in which dissolved substances react to form one (or more) solid products. Many reactions of this type involve the exchange of ions between ionic compounds in aqueous solutions and are sometimes referred to as double displacement, double replacement, or metathesis reactions. These reactions are common in nature and are responsible for the formation of coral reefs in ocean waters and kidney stones in animals. They are used widely in industry for production of a number of commodity and specialty chemicals. Precipitation reactions also play a central role in many chemical analysis techniques, including spot tests used to identify metal ions and gravimetric methods for determining the composition of matter. The extent to which a substance may be dissolved in water, or any solvent, is quantitatively expressed as its solubility, defined as the maximum concentration of a substance that can be achieved under special conditions. Substances with relatively large solubilities are said to be soluble. A substance will precipitate when solution conditions are such that its concentration exceeds its solubility. Substances with relatively low solubilities are said to be insoluble and these are the substances that readily precipitate from solution. More info on these important concepts is provided in a later chapter on solutions. For purposes of predicting the identities of solids formed by precipitation reactions, one may simply refer to patterns of solubility that have been observed for many ionic compounds.

formation of solid lead iodide

A vivid example of precipitation is observed when solutions of potassium iodide and lead nitrate are mixed, resulting in the formation of solid lead iodide: 2KI(aq) + Pb(NO3)2(aq) -> PbI2(s) + 2KNO3(aq) This observation is consistent with the solubility guidelines: the only insoluble compound among all those involved is lead iodide, one of the exceptions to the general solubility of iodide salts. The net ionic equation representing this reaction is: Pb2+(aq) + 2I-(aq) -> PbI2(s) Lead iodide is a bright yellow solid that was formerly used as an artist's pigment known as iodide yellow. The properties of pure PbI2 crystals make them useful for fabrication of X-ray and gamma ray detectors.

chemistry in everyday life- airbags

Airbags are a safety feature provided in most automobiles since the 1990s. The effective operation of an airbag requires that it be rapidly inflated with an appropriate amount (volume) of gas when the vehicle is involved in a collision. This requirement is satisfied in many automotive airbag systems through use of explosive chemical reactions, one common choice being the decomposition of sodium azide, NaN3. When sensors in the vehicle detect a collision, an electrical current is passed through a carefully measured amount of NaN3 to initiate its decomposition. 2NaN3(s) --> 3N2(g) + 2Na(s) This reaction is very rapid, generating gaseous nitrogen that can deploy and fully inflate a typical airbag in a fraction of a second. Among many engineering considerations, the amount of sodium azide used must be appropriate for generating enough nitrogen gas to fully inflate the air bag and ensure its proper function. For example, a small mass (~100 g) of NaN3 will generate approximately 50 L of N2.

Limiting reactant

Consider another food analogy, making grilled cheese sandwiches. 1 slice of cheese + 2 slices of bread --> 1 sandwich Provided with: 28 slices of bread and 11 slices of cheese. We can make: 11 sandwiches with 6 slices of bread left over. Consider this concept now with regard to a chemical process, the reaction of hydrogen with chlorine to yield hydrogen chloride: H2(g) + Cl2(g) --> 2HCl(g) The balanced equation shows the hydrogen and chlorine react in a 1:1 stoichiometric ratio. If these reactants are provided in any other amounts, one of the reactants will nearly always be entirely consumed, thus limiting the amount of product that may be generated. This substance is the limiting reactant, and the other substance is the excess reactant. Identifying the limiting and excess reactants for a given situation requires computing the molar amounts of each reactant provided and comparing them to the stoichiometric amounts represented in the balanced chemical equation. For example, imagine combining 3 moles of H2 and 2 moles of Cl2. This represents a 3:2 ratio (or 1.5:1) ratio of hydrogen to chlorine present for the reaction, which is greater than the stoichiometric ratio of 1:1. Hydrogen, therefore, is present in excess, and chlorine is the limiting reactant. Reaction of all the provided chlorine (2 mol) will consume 2 mol of the 3 mole of hydrogen provided, leaving 1 mol of hydrogen unreacted. An alternative approach to identifying the limiting reactant involves comparing the amount of product expected for the complete reaction of each reactant. Each reactant amount is used separately to calculate the amount of product that would be formed per the reaction's stoichiometry. The reactant yielding the lesser amount of product is the limiting reactant. For example, in the previous paragraph, complete the reaction of the hydrogen would yield mol HCl produced = 3 mol H x (2 mol HCl/1 mol H2) = 6 mol HCl Complete reaction of the provided chlorine would produce mole of HCl produced = 2 mol Cl2 x (2 mol HCl/1 mol Cl2) = 4 mol HCl. The chlorine will be completely consumed once 4 moles of HCl have been produced. Since enough hydrogen was provided to yield 6 moles of HCl, there will be unreacted hydrogen remaining once this reaction is complete. Chlorine, therefore, is the limiting reactant and hydrogen is the excess reactant.

oxidation-reduction reactions background

Earth's atmosphere contains about 20% molecular oxygen, O2, a chemically reactive gas that plays an essential role in the metabolism of aerobic organisms and in many environmental processe that shape the world. The term oxidation was originally used to describe chemical reactions involving O2, but its meaning has evolved to refer to a broad and important reaction class known as oxidation-reduction (redox) reactions. A few examples of such reactions will be used to develop a clear picture of this classification. Some redox reactions involve the transfer of electrons between reactant species to yield ionic products, such as the reaction between sodium and chlorine to yield sodium chloride: 2Na(s) + Cl2(g) --> 2NaCl(s) It is helpful to view the process with regard to each individual reactant, that is, to represent the fate of each reactant in the form of an equation called a half-reaction: 2Na(s) --> 2Na+ (s) + 2e- Cl2(g) + 2e- --> 2Cl-(s) These equations show that Na atoms lose electrons while Cl atoms (in the Cl2 molecule) gain electrons, the "s" subscripts for the resulting ions signifying they are present in the form of a solid ionic compound. For redox reactions of this sort, the loss and gain of electrons define the complementary processes that occur: oxidation = loss of electrons reduction = gain of electrons In this reaction, then, sodium is oxidized and chlorine undergoes reduction. Viewed from a more active perspective, sodium functions as a reducing agent (reductant), since it provides electrons to (or reduces) chlorine. Likewise, chlorine functions as an oxidizing agent (oxidant), as it effectively removes electrons from (oxidizes) sodium. reducing agent = species that is oxidized oxidizing agent = species that is reduced Some redox processes, however, do not involve the transfer of electrons. Consider, for example, a reaction similar to the one yielding NaCl: H2(g) + Cl2(g) --> 2HCl(g) The product of this reaction is a covalent compound, so transfer of electrons in the explicit sense is not involved. To clarify the similarity of this reactioin the previous oen and permit an unambiguous definition of redox reactions, a property called oxidation number has been defined. The oxidation number (or oxidation state) of an element in a compound is the charge its atoms would possess if the compound was ionic.

Culinary aspects of chemistry

Examples of acid-base chemistry are abundant in the culinary world. One example is the use of baking soda, or sodium bicarbonate in baking. NaHCO3 is a base. When it reacts with an acid such as lemon juice, buttermilk, or sour cream in batter, bubbles of carbon dioxide gas are formed from decomposition of the resulting carbonic acid, and the batter "rises." Baking powder is a combo of sodium bicarbonate, and one or more acid salts that react when the two chemicals come in contact with water in the batter. Many people like to put lemon juice or vinegar, both of which are acids, on cooked fish. It turns out that fish have volatile amines (bases) in their systems, which are neutralized by the acids to yield involatile ammoium salts. This reduces the odor of the fish, and also adds a "sour" taste that we seem to enjoy. Pickling is a method used to preserve veggies using a naturally produced acidic environment. The vegetable, such as a cucumber, is placed in a sealed jar submerged in a brine solution. The brine solution favors the growth of beneficial bacteria and suppresses the growth of harmful bacteria. The beneficial bacteria feed on starches in the cucumber and produce lactic acid as a waste product in a process called fermentation. The lactic acid eventually increases the acidity of the brine to a level that kills any harmful bacteria, which require a basic environment. Without the harmful bacteria consuming the cucumbers they are able to last much longer than if they were unprotected. A byproduct of the pickling process changes the flavor of the veggies with the acid making them taste sour.

smallest whole-number coefficients

Finally, with regard to balanced equations, recall that convention dictates use of the smallest whole-number coefficients. Although the equation for the reaction between molecular nitrogen and molecular hydrogen to produce ammonia is, indeed, balanced, 3N2 + 9H2 -> 6NH3 the coefficients are not the smallest possible integers representing the relative numbers of reactant and product molecules. Dividing each coefficient by the greatest common factor, 3, gives the preferred equation: N2 + 3H2 -> 2NH3

Equations for ionic reactions

Given the abundance of water on earth, it stands to reason that a great many chemical reactions take place in aqueous media. When ions are involved in these reactions, the chemical equations may be written with various level of detail appropriate to their intended use. To illustrate this, consider a reaction between ionic compounds taking place in an aqueous solution. When aqueous solutions of CaCl2 and AgNO3 are mixed, a reaction takes place producing aqueous Ca(NO3)2 and solid AgCl. CaCl2(aq) + 2AgNO3(aq) -> Ca(NO3)2(aq) + 2AgCl(s) This balanced equation, derived in the usual fashion, is called a molecular equation because it doesn't explicitly represent the ionic species that are present in the solution. When ionic compounds dissolve in water, they may dissociate into their constituent ions, which are subsequently dispersed homogeneously throughout the resulting solution. Ionic compounds dissolved in water are, therefore, more realistically represented as dissociated ions (a thorough discussion of this important process is provided in the chapter on solutions). Ionic compounds dissolved in water are, therefore, more realistically represented as dissociated ions, in this case: CaCl2(aq) -> Ca2+(aq) + 2Cl-(aq) 2AgNO3(aq) -2Ag+(aq) + 2NO3-(aq) Ca(NO3)2(aq)->Ca2+(aq) + 2NO3-(aq) Unlike these three ionic compounds, AgCl doesn ot dissolve in water to a significant extent, as signified by its physical state notation, s. Explicitly representing all dissolved ions results in a complete ionic equation. In this particular case, the formulas for the dissolved ionic compounds are replaced by formulas for their dissociated ion: Ca2+(aq) + 2Cl-(aq) + 2Ag+(aq) + 2NO3- (aq) -> Ca2+ (aq) + 2NO3-(aq) + 2AgCl(s) examining this equation shows that two chemical species are present in identical form on both sides of the arrow, Ca2+(aq) and NO3-(aq). These spectator ions- ions whose presence is required to maintain charge neutrality- are neither chemically nor physically changed by the process, and so they may be eliminated from the equation to yield a more succinct representation called a net ionic equation. refer to photo... Following the convention of using the smallest possible integers as coefficients, this equation is then written: Cl-(aq) + Ag+(aq) -> AgCl(s) This net ionic equation indicates that solid silver chloride may be produced from dissolved chloride and silver (I) ions, regardless of the source of these ions. These molecular and complete ionic equations provide additional info, namely, the ionic compounds used as sources of Cl- and Ag+.

Coefficients as ratios...

It is common practice to use the smallest possible whole-number coefficients in a chemical equation, as is done in this example. Realize, however, that these coefficients represent the relative numbers of reactants and products, and, therefore, may be correctly interpreted as ratios. Methane and oxygen react to yield carbon dioxide and water in a 1:2:1:2 ratio. This ratio is satisfied if the numbers of these molecules are, respectively, 1-2-1-2, or 2-4-2-4, or 3-6-3-6, and so on. Likewise, these coefficients may be interpreted with regard to any amount (number) unit, and so this equation may be correctly read in many ways, including: -one methane molecule and two oxygen molecules react to yield one carbon dioxide molecule and two water molecules. -one dozen methane molecules and two dozen oxygen molecules react to yield one dozen carbon dioxide molecules and two dozen water molecules. - one mole of methane molecules and 2 moles of oxygen molecules react to yield 1 mole of carbon dioxide molecules and 2 moles of water molecules.

Using fractions instead of integers as intermediate coefficients:

It is sometimes convenient to use fractions instead of integers as intermediate coefficients in the process of balancing a chemical equation. When balance is achieved, all the equation's coefficients may then be multiplied by a whole number to convert the fractional coefficients to integers without upsetting the atom balance. For example, consider the reaction of ethane (C2H6) with oxygen to yield H2O and CO2., represented by the unbalanced quation: C2H6 + O2 -> H2O + CO2. (unbalanced) Following the usual inspection approach, one might first balance C and H atoms by changing the coefficients for the two product species as shown. C2H6 + O2 -> 3H2O + 2CO2. (unbalanced) This results in seven O atoms o the product side of the equation, and odd number- no integer coefficient can be used with the O2 reactant to yield an odd number, so a fractional coefficient, 7/2, is used instead to yield a provisional balanced equation: C2H6 + 7/2O2 -> 3H2O + 2CO2 A conventional balanced equation with integer-only coefficients is derived by multiplying each coefficient by 2: 2C2H6 + 7O2 -> 6H2O + 4CO2

Balancing redox reactions via the half-reaction method

Redox reactions that take place in aqueous media often involve water, hydronium ions, and hydroxide ions as reactants or products. Although these species are not oxidized or reduced, they do participate in chemical change in other ways (e.g., by providing the elements required to form oxyanions). Equations representing these reactions are sometimes very difficult to balance by inspection, so systematic approaches have been developed to assist in the process. One very useful approach is to use the method of half-reactions, which involve the following steps: 1. Write the two half reactions representing the redox process 2. Balance all elements except oxygen and hydrogen 3. Balance oxygen atoms by adding H2O molecules 4. Balance charge by adding electrons 6. If necessary, multiply each half-reaction's coefficients by the smallest possible integers to yield equal numbers of electrons in each 7. Add the balanced half-reactions together and simplify by removing species that appear on both sides of the equation 8. For reactions occurring in basic media (excess hydroxide ions(, carry out these additional steps: a. Add OH- ions to both sides of the equation in numbers equal to the number of H+ ions. b. On the side of the equation containing both H+ and OH- ions, combine these ions to yield water molecules. c. Simplify the equation by removing any reductant water molecules 9. Finally, check that both the number of atoms and total charges are balanced

percent yield

The amount of product that may be produced by a reaction under specific conditions, as calculated per the stoichiometry of an appropriately balanced chemical equation, is called the theoretical yield of the reaction. In practice, the amount of product obtained is called the actual yield, and it is often less than the theoretical yield for a number of reasons. Some reactions are inherently inefficient, being accompanied by side reactions that generate other products. Others are, by nature, incomplete (consider the partial reactions of weak acids and bases discussed earlier in this chapter). Some products are difficult to collect without some loss, and so less than perfect recovery will reduce the actual yield. The extend to which a reaction's theoretical yield is achieved is commonly expressed as percent yield. Percent yield = (actual yield/theoretical yield) x 100% Actual and theoretical yields may be expressed as masses or molar amounts (or any other appropriate property; e.g., volume, if the product is a gas). As long as both yields are expressed using the same units, these units will cancel when percent yield is calculated.

balancing equations

The chemical equation described before is balanced, meaning that equal numbers of atoms for each element involved in the reaction are represented on the reactant and product sides. This is a requirement the equation must satisfy to be consistent with the law of conservation of matter. It may be confirmed by simply summing the numbers of atoms on either side of the arrow and comparing these sums to ensure they are equal. Note that the number of atoms for a given element can be calculated by multiplying the coefficient of any formula containing the element by the element's subscript in the formula. If an element appears in more than one formula on a given side of the equation, the number of atoms represented in each must be computed and then added together. For example, both product species in the example reaction, CO2 and H2O, contain the element oxygen, and so the number of oxygen atoms on the product side of the equation is (1 CO2 molecule x (2 O atom/CO2 molecule)) + (2 H2O molecules x (1 O atom/H2O molecules)) = 4.0 atoms This equation for the reaction between methane and oxygen to yield carbon dioxide and water is confirmed to be balanced per this approach, as shown here: refer to photo

Additional info in chemical equations

The physical states of reactants and products in chemical equations very often are indicated with a parenthetical abbreviation following the formulas. Common abbreviations include s for solids, l for liquids, g for gases, and aq for substances dissolved in water (aqueous solutions, as introduced in the preceding chapter). These notations are illustrated in the example equation here: 2Na(s) + 2H2O (l) -> 2NaOH(aq) + H2(g) This equation represents the reaction that takes place when sodium metal is placed in water. The solid sodium reacts with liquid water to produce molecular hydrogen gas and the ionic compound sodium hydroxide (a solid in pure form, but readily dissolved in water). Special conditions necessary for a reaction are sometimes designated by writing a word or symbol above or below the equation's arrow. For example, a reaction carried out by heating may be indicated by the uppercase Greek letter delta (triangle) over the arrow. Refer to pic

reaction yields

The relative amount of reactants and products represented in a balanced chemical equation are often referred to as stoichiometric amounts. All the exercises of the preceding module involved stoichiometric amounts of reactants. For example, when calculating the amount of product generated from a given amount of reactant, it was assumed that any other reactants required were available in stoichiometric amounts (or greater). In this module, more realistic situations are considered, in which reactants are not present in stoichiometric amounts.

Solubility guidelines for ionic compounds

The solubility guidelines may be used to predict whether a precipitation reaction will occur when solutions of soluble ionic compounds are mixed together. One merely needs to identify all the ions present in the solution, and then consider if possible cation/anion pairing could result in an insoluble compound. For example, mixing solutions of silver nitrate and sodium fluoride will yield a solution containing Ag+, NO3-, Na+ and F- ions. Aside from the two ionic compounds originally present in the solutions, AgNO3 and NaF, two additional ionic compounds may be derived from this collection of ions: NaNO3 and AgF. The solubility guidelines indicate all nitrate salts are soluble but that AgF is one of the exceptions to the general solubility of fluoride salts. A precipitation reaction, therefore, is predicted to occur, as described in the following equations: NaF(aq) + AgNO3(aq) -> AgF(s) + NaNO3(aq). (molecular) Ag+(aq) + F-(aq) -> AgF(s)

Oxidation reduction (redox) reactions

Using the oxidation number concept, an all-inclusive definition of redox reaction has been established. Oxidation-reduction (redox) reactions are those in which one or more elements involved undergo a change in oxidation number. (While the vast majority of redox reactions involve changes in oxidation number for two or more elements, a few interesting exceptions to this rule do exist) Definitions for the complementary processes of this reaction class are correspondingly revised as shown here: oxidation = increase in oxidation number reduction = decrease in oxidation number Returning to the reactions used to introduce the topic, they may now be identified as redox processes. In the reaction between sodium and chlorine to yield sodium chloride, sodium is oxidized (its oxidation number increases from 0 in Na to +1 in NaCl) and chlorine is reduced (its oxidation number decreases from 0 in Cl2 to -1 in NaCl). In the reaction between molecular hydrogen and chlorine, hydrogen is oxidized (its oxidation number increases from 0 in H2 to +1 in HCl) and chlorine is reduced (its oxidation number decreases from 0 in Cl2 to -1 in HCl). Several subclasses of redox reactions are recognized, including combustion reactions in which the reductant (also called a fuel) and oxidant (often, but not necessarily, molecular oxygen) react vigorously and produce significant amounts of heat, and often light, in the form of a flame. Solid rocket-fuel reactions such as the one depicted in figure 4.1 are combustion processes. A typical propellant reaction in which solid aluminum is oxidized by ammonium perchlorate is represented by this equation: 10Al(s) + 6NH4ClO4(s) --> 4Al2O3(s) + 2AlCl3(s) + 12H2O(g) + 3N2(g) Single-displacement (replacement) reactions are redox reactions in which an ion in solution is displaced (or replaced) via the oxidation of a metallic element. One common example of this type of reaction is the acid oxidation of certain metals: Zn(s) + 2HCl(aq) --> ZnCl2(aq) + H2(g) Metallic elements may also be oxidized by solutions of other metal salts; for example: Cu(s) + 2AgNO3(aq) --> Cu(NO3)2 (aq) + 2Ag(s) This reaction may be observed by placing copper wire in a solution containing a dissolved silver salt. Silver ions in solution are reduced to elemental silver at the surface of the copper wire, and the resulting Cu2+ ions dissolve in the solution to yield a characteristic blue color.

stoichiometry

a balanced chemical equation provides a great deal of info in a very succinct format. Chemical formulas provide the identities of the reactants and products involved in the chemical change, allowing classification of the reaction. Coefficients provide the relative numbers of these chemical species, allowing a quantitative assessment of the relationships between the amounts of substances consumed and produced by the reaction. These quantitative relationships are known as the reaction's stoichiometry, a term derived from the Greek words stoicheion (meaning "element") and metron (meaning "measure". The general approach to using stoichiometric relationships is similar in concept to the way people go about many common activities. Food prep, for example, offers an appropriate comparison. A recipe for making eight pancakes calls for 1 cup pancake mix, 3/4 cup milk, and one egg. The "equation" representing the preparation of the pancakes per this recipe is 1 cup mix + 3/4 cup milk + 1 egg --> 8 pancakes If two dozen pancakes are needed for a big family breakfast, the ingredient amounts must be increased proportionally according to the amounts given in the recipe. For example, the number of eggs required to make 24 pancakes is (24 pancakes) x (1 egg/8pancakes) = 3 eggs Balanced chemical equations are used in much the same fashion to determine the amount of one reactant required to react with a given amount of another reactant, or to yield a given amount of product, and so forth. The coefficients in the balanced equation are used to derive stoichiometric factors that permit computation of the desired quantity. To illustrate this idea, consider the production of ammonia by reaction of hydrogen and nitrogen. N2(g) + 3H2(g) --> 2NH3(g) This equation shows that ammonia molecules are produced from hydrogen molecules in a 2:3 ratio, and stiochiometric factors may be derived using any amount (number) unit: 2NH3 molecules/3H2 molecules or 2 dozen NH3 molecules/3 dozen H2 molecules or 2 mol NH3 molecules/3 mol H2 molecules These stoichiometric factors can be used to compute the number of ammonia molecules produced from a given number of hydrogen molecules, or the number of hydrogen molecules required to produce a given number of ammonia molecules. Similar factors may be derived for any pair of substances in a chemical equation.

base

a base is a substance that will dissolve in water to yield hydroxide ions, OH-. The most common bases are ionic compounds composed of alkali or alkaline earth metal cations (groups 1 and 2) combined with the hydroxide ion- for example, NaOH and Ca(OH)2. Unlike the acid compounds discussed previously, these compounds do not react chemically with water; instead they dissolve and dissociate, releasing hydroxide ions directly into the solution. For example, KOH and Ba(OH)2 dissolve in water and dissociate completely to produce cations (K+ and Ba2+, respectively) and hydroxide ions, OH-. These bases, along with the other hydroxides that completely dissociate in water, are considered strong bases. Consider as an example the dissolution of lye (sodium hydroxide) in water: NaOH(s) -> Na+ (aq) + OH- (aq) This equation confirms that sodium hydroxide is a base. When dissolved in water, NaOH dissociates to yield Na+ and OH- ions. This is also true for any other ionic compound containing hydroxide ions. Since the dissociation process is essentially complete when ionic compounds dissolve in water under typical conditions, NaOH and other ionic hydroxides are all classified as strong bases. Unlike ionic hydroxides, some compounds produce hydroxide ions when dissolved by chemically reacting with water molecules. In all cases, these compounds react only partially and so are classified as weak bases. These types of compounds are also abundant in nature and important commodities in various technologies. For example, global production of the weak base ammonia is typically well over 100 metric tons annually, being widely used as an agricultural fertilizer., a raw material for chemical synthesis of other compounds, and an active ingredient in household cleaners. When dissolved in water, ammonia reacts partially to yield hydroxide ions, as shown here: NH3(aq) + H2O(l) <--> NH4+(aq) + OH- (aq) This is, by definition, an acid-base reaction, in this case involving the transfer of H+ ions from water molecules to ammonia molecules. Under typical conditions, only about 1% of the dissolved ammonia is present as NH4+ ions.

neutralization reaction

a neutralization reaction is a specific type of acid-base reaction in which the reactants are an acid and a base (but not water), and the products are often a salt and water acid + base --> salt + water to illustrate a neutralization reaction, consider what happens when a typical antacid such as milk of magnesia (an aqueous suspension of solid Mg(OH)2) is ingested to ease symptoms associated with excess stomach acid (HCl): Mg(OH)2(s) + 2HCl(aq) -> MgCl2(aq) + 2H2O(l) Note that, in addition to water, this reaction produces a salt, magnesium chloride.

acid-base reaction

an acid-base reaction is one in which a hydrogen ion, H+, is transferred from one chemical species to another. Such reactions are of central importance to numerous natural and technological processes, ranging from the chemical transformations that take place within cells and the lakes and oceans, to the industrial-scale production of fertilizers, pharmaceuticals, and other substances essential to society. The subject of acid-base chemistry, therefore, is worthy of thorough discussion, and a full chapter is devoted to this topic later in text. For purposes of brief introduction, we'll consider only the more common types of acid-base reactions that take place in aqueous solutions. In this context, an acid is a substance that will dissolve in water to yield hydronium ions, H3O+. As an example, consider this equation: HCl(aq) + H2(aq) -> Cl-(aq) + H3O+(aq) The process represented by this equation confirms that hydrogen chloride is an acid. When dissolved in water, H3O+ ions are produced by a chemical reaction in which H+ ions are transferred from HCl molecules to H2O molecules. The nature of HCl is such that its reaction with as just described is essentially 100% efficient: virtually every HCl molecule that dissolves in water will undergo this reaction. Acids that completely react in this fashion are called strong acids, and HCl is one among just a handful of common acid compounds that are classified as strong. A far greater number of compounds behave as weak acids and only partially react with water, leaving a large majority of dissolved molecules in their original form and generating a relatively small amount of hydronium ions. Weak acids are commonly encountered in nature, being the substances partly responsible for the tangy taste of citrus fruits, the stinging sensation of insect bites, and the unpleasant smells associated with body odor. A familiar example of a weak acid is acetic acid, the main ingredient in food vinegars: CH3CO2H(aq) + H2O(l) <--> CH3CO2-(aq) + H3O+(aq) When dissolved in water under typical conditions, only about 1% of acetic acid molecules are present in the ionized form, CH3CO2-. The use of a double arrow above denotes the partial reaction aspect of this process, a concept addressed fully in the chapters on chemical equilibrium.

chemical equation

an earlier chpt introduced the use of element symbols to represent individual atoms. When atoms gain or lose electrons to yield ions, or combine with other atoms to form molecules, their symbols are modified or combined to generate chemical formulas that appropriately represent these species. Extending this symbolism to represent both identities and the relative qualities of substances undergoing a chemical (or physical) change involves writing and balancing a chemical equation. Consider as an example the reaction between one methane molecule (CH4) and two diatomic oxygen molecules (O2) to produce one carbon dioxide molecule (CO2) and two water molecules (H2O). Photo: upper half represents chemical equation and lower half represents space-filling molecular model

stomach antacids

our stomachs contain a solution of roughly .03 M HCl, which helps us digest the food we eat. The burning sensation associated with heartburn is a result of the acid of the stomach leaking through the muscular valve at the top of the stomach into the lower reaches of the esophagus. The lining of the esophagus is not protected from the corrosive effects of stomach acid the way the lining of the stomach is, and the results can be very painful. When we have heartburn, it feels better if we reduce the excess acid in the esophagus by taking an antacid. As you may have guessed, antacids are bases. One of the most common antacids is calcium carbonate, CaCO3. The reaction, CaCO3(s) + 2HCl(aq) <--> CaCl2(aq) + H2O(l) + CO2(g) not only neutralizes stomach acid, it also produces CO2(g), which may result in a satisfying belch. Milk of Magnesia is a suspension of the sparingly soluble base magnesium hydroxide, Mg(OH)2. It works according to the reaction: Mg(OH)2(s) <--> Mg2+(aq) + 2OH-(aq) The hydroxide ions generated in this equilibrium then go to react with the hydronium ions from the stomach acid, so that: H3O+ + OH- <-->2H2O(l) This reaction does not produce carbon dioxide, but magnesium-containing anatacids can have a laxative effect. Several antacids have aluminum hydroxide, Al(OH)3, as an active ingredient. The aluminum hydroxide tends to cause constipation, and some antacids use aluminum hydroxide in concert with magnesium hydroxide to balance the side effects of the two substances.

common strong acids

refer to pic