Chapter 5 Chemistry: Atoms First Julia Burdge & Jason Overby Ionic and Covalent Compounds Kent L. McCorkle Cosumnes River College Sacramento, CA
5 Ionic and Covalent Compounds 5.1 Compounds 5.2 Lewis Dot Symbols 5.3 Ionic Compounds and Bonding 5.4 Naming Ions and Ionic Compounds Formulas of Ionic Compounds Naming Ionic Compounds 5.5 Covalent Bonding and Molecules Molecules Molecular Formulas Empirical Formulas 5.6 Naming Molecular Compounds Specifying Numbers of Atoms Compounds Containing Hydrogen Organic Compounds
5 Ionic and Covalent Compounds 5.7 Covalent Bonding in Ionic Species Polyatomic Ions Oxoacids Hydrates Familiar Inorganic Compounds 5.8 Molecular and Formula Masses 5.9 Percent Composition of Compounds 5.10 Molar Mass Interconverting Mass, Moles, and Number of Particles Determination of Empirical Formula and Molecular Formula from Percent Composition
Compounds 5.1 A compound is a substance composed of two or more elements combined in a specific ratio and held together by chemical bonds. Familiar examples of compounds are water and salt (sodium chloride).
Lewis Dot Symbols 5.2 • • • B • • • • • • B B B • • • When atoms form compounds, it is their valence electrons that actually interact. ALewis dot symbolconsists of the element’s symbol surround by dots. Each dot represents a valence electron. Boron 1s22s22p1 3 valence electrons Lewis dot symbol for boron other reasonable Lewis dot symbols for boron
Lewis Dot Symbols Na: 1s22s22p63s1 Na+: 1s22s22p6 10 electrons total, isoelectronic with Ne Cl: 1s22s22p63s23p5 Cl‒: 1s22s22p63s23p6 18 electrons total, isoelectronic with Ar Atoms combine in order to achieve a more stable electron configuration. Maximum stability results when a chemical species is isoelectronic with a noble gas.
Lewis Dot Symbols Lewis dot symbols of the main group elements.
• • •• • • • • • • B C N • • 1s22s22p1 1s22s22p2 1s22s22p3 • Na •• • • O •• Lewis Dot Symbols Dots are not paired until absolutely necessary. 5 valence electrons; first pair formed in the Lewis dot symbol For main group metals such as Na, the number of dots is the number of electrons that are lost. For nonmetals in the second period, the number of unpaired dots is the number of bonds the atom can form.
Na • Na 1s22s22p63s1 Na+ 1s22s22p6 •• 2‒ •• • • O • • • • O •• •• O2‒ 1s22s22p6 Lewis Dot Symbols Ions may also be represented by Lewis dot symbols. Remember the charge Na+ Core electrons not represented in the Lewis dot symbol Valence electron lost in the formation of the Na+ ion. O 1s22s22p4
Worked Example 5.1 Write Lewis dot symbols for (a) fluoride ion (F-), (b) potassium ion (K+), and (c) sulfide ion (S2-). Strategy Starting with the Lewis dot symbols for each element, add dots (for anions) or remove dots (for cations) as needed to achieve the correct charge on each ion. Don’t forget to include the appropriate charge on the Lewis dot symbol. Solution (a) (b) K+ (c) Think About It For ions that are isoelectronic with noble gases, cations should have no dots remaining around the element symbol, whereas anions should have eight dots around the element symbol. Note, too, that for anions, we put square brackets around the Lewis dot symbol and place the negative charge outside the brackets. Because the symbol for a common cation such as the potassium ion has no remaining dots, square brackets are not necessary.
Ionic Compounds and Bonding 5.3 Na • Na+ + e− •• •• − • • • • + e− Cl Cl • • • •• •• •• •• − • • • • Na • Na+ + + Cl Cl • • • •• •• Ionic bondingrefers to the electrostatic attraction that holds oppositely charged ions together in an ionic compound. The attraction between the cation and anion draws them together to form NaCl
Ionic Compounds and Bonding The resulting electrically neutral compound, sodium chloride, is represented with the chemical formula NaCl. The chemical formula, or simply formula, of an ionic compound denotes the constituent elements and the ratio in which they combine.
− − − + + + + + − − − − + + Ionic Compounds and Bonding A three-dimensional array of oppositely-charged ions is called a lattice. Lattice energy is the amount of energy required to convert a mole of ionic solid to its constituent ions in the gas phase. NaCl(s) Na+(g) + Cl−(g) Hlattice = +788 kJ/mol
Q1 Q2 d Ionic Compounds and Bonding The magnitude of lattice energy is a measure of an ionic compound’s stability. Lattice energy depends on the magnitudes of the charge and on the distance between them. Q = amount of charge d = distance of separation
Ionic Compounds and Bonding The magnitude of lattice energy is a measure of an ionic compound’s stability. Lattice energy depends on the magnitudes of the charge and on the distance between them.
Worked Example 5.2 Arrange MgO, CaO, and SrO in order of increasing lattice energy. Strategy Consider the charges on the ions and the distances between them. Apply Coulomb’s law to determine the relative lattice energies. All three compounds contain O2- and all three cations are +2. Recalling that lattice energy increases as the distance between ions decreases, we need only consider the radii of the cations as all three contain the same anion. From Figure 4.13, the ionic radii are 0.72 Å (Mg2+), 1.00 Å (Ca2+), and 1.18 Å (Sr2+). Solution MgO has the smallest distance between ions, whereas SrO has the largest distance between ions. Therefore, in order of increase lattice energy: SrO < CaO < MgO. Think About It Mg, Ca, and Sr are all Group 2A metals, so we could have predicted this result without knowing their radii. Recall that ionic radii increase as we move down a column in the periodic table, and charges that are farther apart are more easily separated (meaning the lattice energy will be smaller.) The lattice energies of SrO, CaO, and MgO are 3217, 3414, and 3890 kJ/mol, respectively.
Naming Ions and Ionic Compounds 5.4 A monatomic ion is named by changing the ending of the element’s name to –ide. Cl– is chloride O2– is oxide Some metals can form cations of more than one possible charge. Fe2+ : ferrous ion [Fe(II)] Fe3+ : ferric ion [Fe(III)] Mn2+ : manganese(II) ion Mn3+ : manganese(III) ion Mn4+ : manganese(IV) ion
Naming Ions and Ionic Compounds Formulas for ionic compounds are generally empirical formulas. Ionic compounds are electronically neutral.
Al3+ O2– Al2O3 Formulas of Ionic Compounds In order for ionic compounds to be electronically neutral, the sum of the charges on the cation and anion in each formula must be zero. Aluminum oxide: Sum of charges: 2(+3) + 3(–2) = 0
Naming Ions and Ionic Compounds To name ionic compounds: 1) Name the cation • omit the word ion • use a Roman numeral if the cation can have more than one charge 2) Name the anion • omit the word ion Examples: NaCN sodium cyanide FeCl2 iron(II) chloride FeCl3 iron(III) chloride
Worked Example 5.3 Name the following ionic compounds: (a) CaO, (b) Mg3N2, and (c) Fe2S3. Strategy Begin by identifying the cation and anion in each compound, and then combine the names for each, eliminating the word ion. Solution (a) CaO is calcium oxide. (b) Mg3N2 is magnesium nitride. (c) Fe2S3 is iron(III) sulfide. Think About It Be careful not to confuse the subscript in the formula with the charge in the metal ion. In part (c), for example, the subscript on Fe is 2, but this is an iron(III) compound.
Worked Example 5.4 Deduce the formulas of the following ionic compounds: (a) mercury(II) chloride, (b) lead(II) bromide, and (c) potassium nitride. Strategy Identify the ions in each compound, and determine their ratios of combination using the charges on the cation and anion in each. Solution (a) Mercury(II) chloride is a combination of Hg2+ and Cl-. To produce a neutral compound, these two ions must combine in a 1:2 ratio – HgCl2. (b) Lead(II) bromide is a combination of Pb2+ and Cl-. These ions combine in a 1:2 ratio to give PbBr2. (c) Potassium nitride is a combination of K+ and N3-. These ions combine in a 3:1 ratio to give K3N. Think About It Make sure that the charges sum to zero in each compound formula. In part (a), for example, Hg2+ + 2Cl- = (+2) + 2(-1) = 0; in part (b), (+2) + 2(-1) = 0; and in part (c), 3(+1) + (-3) = 0.
Covalent Bonding and Molecules 5.5 When compounds form between elements with similar properties, electrons are not transferred from one element to another but instead are shared in order to give each atom a noble gas configuration. This approach is known as the Lewis theory of bonding, named for it’s proponent, Gilbert Lewis. Lewis theory depicts bond formation in H2 as H∙ + ∙H → H:H This type of arrangement, where two atoms share a pair of electrons, is known as covalent bonding, and the shared pair of electrons constitutes a covalent bond.
Covalent Bonding and Molecules A molecule is a combination of at least two atoms in a specific arrangement held together by chemical forces (chemical bonds). A molecule may be an element or a compound. Different samples of a given compound always contain the same elements in the same ratio. This is known as the law of definite proportions.
Covalent Bonding and Molecules If two elements can for two or more different compounds, the law of multiple proportionstells us that the ratio of masses of one element that combine with a fixed mass of the other element can be expressed in small whole numbers. In addition to carbon dioxide, carbon also combines with oxygen to form carbon monoxide.
Covalent Bonding and Molecules mass ratio of O to C in carbon dioxide mass ratio of O to C in carbon monoxide 2.66 1.33 = 2:1 = The mass ratio of oxygen to carbon in carbon dioxide is 2.66:1, and the ratio of oxygen to carbon in carbon monoxide is 1.33:1. The ratio of two such mass ratios can be expressed as small whole numbers.
Covalent Bonding and Molecules Diatomic moleculescontain two atoms and may be either heteronuclear or homonuclear. Polyatomicmoleculescontain more than two atoms.
Covalent Bonding and Molecules A chemical formula denotes the composition of the substance. A molecular formulashows the exact number of atoms of each element in a molecule. Some elements have two or more distinct forms known as allotropes. • For example, oxygen (O2) and ozone (O3) are allotropes of oxygen. A structural formulashows not only the elemental composition, but also the general arrangements.
Worked Example 5.5 Write the molecular formula of ethanol based on its ball-and-stick model, shown here. Strategy Refer to the labels on the atoms (or see Table 5.3). There are two carbon atoms, six hydrogen atoms, and one oxygen atom, so the subscript on C will be 2 and the subscript on H will be 6, and there will be no subscript on O. Solution C2H6O Think About It Often the molecular formula for a compound such as ethanol (consisting of carbon, hydrogen, and oxygen) is written so that the formula more closely resembles the actual arrangement of atoms in the molecule. Thus, the molecular formula for ethanol is commonly written as C2H5OH.
Covalent Bonding and Molecules Molecular substances can also be represented using empirical formulas, the whole-number ratio of elements. While, the molecular formulas tell us the actual number of atoms (the true formula), the empirical formula gives the simplest formula. Molecular formula: N2H4 Empirical formula: NH2 The molecular and empirical formulas are often the same.
Worked Example 5.6 Write the empirical formulas for the following molecules: (a) glucose (C6H12O6), a substance known as blood sugar; (b) adenine (C5H5N5), also known as vitamin B4; and (c) nitrous oxide (N2O), a gas that is used as an anesthetic (“laughing gas”) and as an aerosol propellant for whipped cream. Strategy To write the empirical formula, the subscripts in the molecular formula must be reduced to the smallest possible whole numbers (without altering the relative numbers of atoms).
Worked Example 5.6 (cont.) Solution (a) Dividing each of the subscripts in the molecular formula for glucose by 6, we obtain the empirical formula CH2O. If we had divided the subscripts by 2 or 3, we would have obtained the formulas C3H6O3 and C2H4O2, respectively. Although the ratio of carbon to hydrogen to oxygen atoms in each of these formulas is correct (1:2:1), neither is the simplest formula because the subscripts are not in the smallest possible whole-number ratio. (b) Dividing each subscript by 5, we get the empirical formula CHN. (c) Because the subscripts in the formula for nitrous oxide are already the smallest possible whole numbers, its empirical formula is the same as its molecular formula N2O. Think About It Make sure that the ratio in each empirical formula is the same as that in the corresponding molecular formula and that the subscripts are the smallest possible whole numbers. In part (a), for example, the ratio of C:H:O in the molecular formula is 6:12:6, which is equal to 1:2:1, the ratio expressed in the empirical formula.
Naming Molecular Compounds 5.6 Remember that binary molecular compounds are substances that consist of just two different elements. Nomenclature: 1) Name the first element that appears in the formula. 2) Name the second element that appears in the formula, changing its ending to –ide. Examples: HCl hydrogen chloride HI hydrogen iodide
Naming Molecular Compounds Greek prefixes are used to denote the number of atoms of each element present.
Naming Molecular Compounds The prefix mono- is generally omitted for the first element. For ease of pronunciation, we usually eliminate the last letter of a prefix that ends in “o” or “a” when naming an oxide. Example: N2O5 is dinitrogen pentoxide not dinitrogen pentaoxide
Worked Example 5.7 Name the following binary molecular compounds: (a) NF3 and (b) N2O4. Strategy Each compound will be named using the systematic nomenclature including, where necessary, appropriate Greek prefixes. Solution (a) nitrogen trifluoride (b) dinitrogen tetroxide Think About It Make sure that the prefixes match the subscripts in the molecular formulas and that the word oxide is not preceded immediately by an “a” or an “o”.
Worked Example 5.8 Write the chemical formulas for the following binary molecular compounds: (a) sulfur tetrafluoride and (b) tetraphosphorus decasulfide. Strategy The formula for each compound will be deduced using the systematic nomenclature guidelines. Solution (a) SF4 (b) P4S10 Think About It Double-check that the subscripts in the formulas match the prefixes in the compound names: (a) 4 = tetra and (b) 4 = tetra and 10 = deca.
Compounds Containing Hydrogen The names of molecular compounds containing hydrogen do not usually conform to the systematic nomenclature guidelines. Many are called by the common, nonsystematic names or by names that do not indicate explicitly the number of H atoms present. Examples: B2H6 Diborane SiH4 Silane NH3 Ammonia PH3 Phosphine H2O Water H2S Hydrogen sulfide
Compounds Containing Hydrogen One definition of an acid is a substance that produces hydrogen ions (H+) when dissolved in water. HCl is an example of a binary compound that is an acid when dissolved in water. To name these types of acids: 1) remove the –gen ending from hydrogen 2) change the –ide ending on the second element to –ic. hydrogen chloride → hydrochloric acid
Compounds Containing Hydrogen A compound must contain at least one ionizable hydrogen atomto be an acid upon dissolving.
Organic Compounds Our nomenclature discussion so far has focused on inorganic compounds, generally defined as those without carbon. Organic compounds contain carbon and hydrogen, sometimes in combination with other atoms. Hydrocarbons contain only carbon and hydrogen. The simplest hydrocarbons are called alkanes.
Organic Compounds Many organic compounds contain groups of atoms known as functional groups, which often determine a molecule’s reactivity.