SGCh02 - Chapter Two ATOMS, MOLECULES, AND IONS • • •...

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Unformatted text preview: Chapter Two ATOMS, MOLECULES, AND IONS • • • • • Dalton's Atomic Theory Subatomic Particles and the Structure of the Atom Molecules and Ions Chemical Formulas Naming Inorganic Compounds DALTON'S ATOMIC THEORY STUDY OBJECTIVES 1. 2. 3. State the laws of definite proportions, multiple proportions, and conservation of mass. List and discuss the postulates of Dalton's atomic theory. Describe how Dalton's atomic theory explains these three laws. Atoms. According to John Dalton's atomic theory, which he proposed in 1808, elements are composed of extremely small, indivisible particles called atoms. Dalton assumed that a toms were the smallest unit of an element that can enter into chemical combination. Atoms of the same element all have the same mass, and atoms of different elements have different masses. His theory explained three laws that were known at that time. The atomic theory consists of a set of postulates which can explain the law of definite proportions, the law of multiple proportions, and law of conservation of mass . The postulates are summarized below. D alton's Atomic Theory 1. All matter is made of extremely small, indivisible particles called atoms. An atom is the smallest unit of an element that can enter into chemical combination. Atoms of the same element are identical in size, mass, and chemical properties. 2. Atoms of one element are different from atoms of another element. 3. Compounds are formed when atoms from two or more elements combine. Atoms combine in the ratio of small whole numbers. For a given compound, the number of atoms of one element that combine per atom of the other element is a fixed ratio. The smallest particle that has the properties of the compound is called a molecule. 4. Atoms are the smallest units of chemical change. Chemical change involves the combination, separation, or rearrangement of atoms. Atoms are not destroyed or created in chemical processes. The l aw of definite proportions refers to the elemental constituents of compounds and states that all purified samples of a compound contain its constituent elements combined in the same proportions by mass. For example, all samples of the compound carbon monoxide, no matter what size, show that it is 43% carbon and 57% oxygen by mass. Dalton's proposal was that the smallest particle of carbon monoxide was a molecule consisting of one carbon atom and one oxygen atom. If all carbon atoms had the same mass, and all oxygen 23 Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 2 4 / Atoms, Molecules, and Ions atoms had a mass about 1.3 times greater than carbon atoms, then the composition of carbon monoxide would be exactly as given above. While working on the atomic theory, Dalton discovered the l aw of multiple proportions : When two elements form more than one compound, the various masses of one element combining with a fixed mass of another element are related by small whole-number ratios. This law can be applied to two compounds of nitrogen and oxygen; nitrogen monoxide, and nitrogen dioxide. The mass of oxygen combining with nitrogen in nitrogen dioxide was found to be twice as great as the mass of oxygen combining with nitrogen in nitrogen monoxide. Dalton proposed that the two compounds of N and O were the result of atoms of the two elements combining in different ratios to form two different molecules. When one atom of nitrogen combined with one atom of oxygen, nitrogen monoxide was formed. But when one atom of nitrogen combined with two atoms of oxygen, then nitrogen dioxide was the compound formed. The l aw of conservation of mass refers to an observation that has been made many times: There is no detectable gain or loss of mass during a chemical reaction. The law of conservation of mass is consistent with postulate 4. Atoms are neither created nor destroyed in a chemical reaction. Rather they are merely rearranged to form new combinations of atoms which we would recognize as new compounds. Mass is conserved in chemical reactions because atoms are conserved. EXERCISE 1. Describe briefly in your own words the four postulates of Dalton's atomic theory. SUBATOMIC PARTICLES AND THE STRUCTURE OF THE ATOM STUDY OBJECTIVES 1. 2. 3. Name and describe the three subatomic particles of most importance to chemistry and give their location within the atom. Write the symbol of an isotope having been given its mass number and element name. Determine the number of protons, neutrons, and electrons in an atom given its isotopic symbol. Subatomic Particles. In the early 1900s, scientists learned that all atoms are constructed from the same three subatomic particles: the electron, proton, and neutron. Whereas atoms themselves are electrically neutral they were found to consist of electrically charged particles. The electron has a single unit of negative charge; the proton has a single unit of positive charge; and the neutron has no charge. The proton and neutron have essentially the same mass, 1.673 × 10 –24 g, and 1.675 × 10 –24 g, respectively. The electron mass is much smaller: 9.11 × 10 –28 g which is only 1/1836 the mass of a proton. Rutherford's experiment. Rutherford's experiments on alpha particle scattering in 1910 led the way to the nuclear model of the atom. Alpha particles are ejected from certain radioactive elements with very high kinetic energies. Remember that alpha particles have 2 units of positive charge, and a mass 4 times that of a proton. Being smaller than the atoms of most elements they can act as probes of the interior of atoms. Rutherford's experiments lead to the following conclusions and about the structure of atoms. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Atoms, Molecules, and Ions / 2 5 O bservations made by Rutherford when alpha particles were scattered by gold foil: 1. The majority of particles passed through the foil and were either undeflected or only slightly deflected (angle less than 2%). 2. A few particles were deflected (scattered) by more than a few degrees. 3. It was extremely rare, but 1in 20,000 alpha particles were deflected back toward the direction from which they had come. Interpretations made by Rutherford about the nature of the atom: 1. Most of the atom must be empty space. 2. The atom's positive charges are concentrated into a dense central core called the nucleus. 3. The diameter of a typical atom is about 10,000 times greater than the diameter of a typical nucleus. Atomic Number and Mass Number. An atom of one element is distinguished from an atom of another element by its number of protons. The a tomic number , Z , of an element is the number of protons in the atomic nucleus. For example, the atomic number of oxygen is 8; therefore, all oxygen atoms contain eight protons and also eight electrons. The total mass of an atom is determined almost entirely by the number of protons and neutrons. In many cases the mass of the electrons can be neglected because it is so much smaller. The m ass number , A , is the same as the total number of neutrons and protons present in the nucleus of an atom. The number of neutrons in an atom is A – Z . Isotopes. Atoms of a given element that differ in the number of neutrons, and consequently in mass are called i sotopes . For example, there are two atoms of the element lithium, one with a mass number of 6 and another with a mass number of 7. Isotopes can be referred to by their mass numbers as in lithium-6 (pronounced lithium six) and lithium-7. The different mass numbers are the result of different numbers of neutrons per atom. An atom of the isotope lithium-6 contains three protons, three electrons, and three neutrons, whereas an atom of lithium-7 contains three protons, three electrons and four neutrons. Both atoms contain three protons, and so both are isotopes of the element lithium. In nature, elements are found as mixtures of isotopes. Special symbols are used to designate specific isotopes. The general symbol for an isotope is: mass number A ZX symbol of the element atomic number 6 7 The symbols for the isotopes of lithium are 3 Li and 3 Li. Isotopes of an element have similar chemical properties, and form the same types of compounds. The chemical properties of an element depend on the number of protons and electrons in an atom, not the number of neutrons. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 2 6 / Atoms, Molecules, and Ions _______________________________________________________________________________ EXAMPLE 2.1 Isotopic Symbols The three isotopes of oxygen found in nature are oxygen-16, oxygen-17, and oxygen-18. Write their isotopic symbols. •Method of Solution A The general form of the symbols is Z O, where A is the mass number and Z is the atomic number. The atomic number of oxygen is 8, and so all oxygen atoms contain eight protons. The mass numbers are 16, 17, and 18, respectively. Answer: The isotopic symbols are: 16 8O 17 8O 18 8O _______________________________________________________________________________ EXAMPLE 2.2 Isotopic Symbols How many neutrons are present in the nucleus of each of the oxygen isotopes in the previous example? •Method of Solution The number of neutrons is given by A – Z , the mass number minus the proton number. Answer: For 16 O, there are 16 – 8 = 8 neutrons. For 17 O, there are 17 – 8 = 9 neutrons. For 18 O, there are 18 – 8 = 10 neutrons _______________________________________________________________________________ EXERCISES 2. 3. What evidence did Rutherford find that supported his theory of the atomic nucleus? 4. If the nucleus of an atom was the diameter of a nickel (about 2 cm), what would the diameter be of a such an atom? 5. What is the mass number of a sodium atom that has 13 neutrons? 6. How many neutrons are in an atom of Ag-109? 7. Back What are the three fundamental particles from which atoms are made? What are their electric charges. Which particles are in the atomic nucleus? Write the symbol for each of the following isotopes. a. An atom with Z = 30 and 37 neutrons. b. An atom with Z = 51 and 69 neutrons. Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Atoms, Molecules, and Ions / 2 7 MOLECULES AND IONS STUDY OBJECTIVES 1. 2. 3. Define the terms molecule and ion. Determine the numbers of protons and electrons in given monatomic cations and anions. Learn the formulas of several polyatomic ions. Molecules. Compounds are produced by the combination of atoms of different elements. Compounds fall into two general types depending on how their atoms are held together: molecular and ionic. For a given compound, the number of atoms of one element that combine per atom of the other element is a fixed ratio. Molecules are clusters of atoms that are recognized as single discrete particles. A molecule is the smallest particle that has the properties of the compound. Diatomic molecules contain two atoms, while polyatomic molecules contain more than two atoms. Some representations of simple molecules are given in Figure 2.9 of the text. Compounds Ionic Molecular Ions. Atoms are electrically neutral because they have equal numbers of protons and electrons. However, an atom can acquire a charge by gaining or losing electrons. These processes do not change the identity of an atom because the atomic number (the number of protons) is unaffected. A change in the number of electrons an atom has will alter its electrical charge. Any atom or group of atoms having a net electrical charge is called an ion. Certain atoms (typically metal elements) always tend to lose electrons (1 or 2, some atoms lose 3 electrons). These form positive ions which are called cations. Others atoms (typically those of nonmetal elements) tend to gain 1 or 2 electrons forming negative ions called anions. Take for example, the magnesium atom, which has 12 protons in the nucleus and 12 electrons outside the nucleus. It can lose two of its electrons. The result is a particle called the magnesium ion which has 12 protons and 10 electrons. Its net charge is +2 because in the ion the number of positive charges outnumber the negative charges by two. The symbol for the magnesium ion is Mg2+ . The superscript (2+) represents the net charge of the ion. A few elements form ions with +3 charges. Negative ions are formed when atoms gain extra electrons. For example, the oxygen atom which has 8 protons and 8 electrons can gain two electrons. The result is an oxygen ion which has 8 protons and 10 electrons. The net charge of the ion is –2, because there are two more negative charges than positive charges. The symbol for this ion is O2– . The O 2– ion is called the oxide ion. The names of monatomic anions end in -ide. Polyatomic Ions. In addition to monatomic ions such as Mg2+ and O2– , more complex ions are known. A polyatomic ion is a group of two or more atoms that carries an electrical charge. The hydroxide ion (OH– ) is a simple example. The oxygen atom and hydrogen atom are held together much as they would be in a molecule, but they have gained an extra electron. This gives the ion a –1 charge. Sometimes the net charge on these ions seems unpredicable and so the best thing to do is memorize the names, formula, and charge of the most common polyatomic ions. Table 2.1 lists some important polyatomic ions. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 2 8 / Atoms, Molecules, and Ions Table 2.1 Some Polyatomic Ions Formula Name OH– hydroxide 2– CO3 – HCO3 – ClO 3 – NO3 – NO2 3– PO4 2– SO4 2– SO3 carbonate hydrogen carbonate chlorate nitrate nitrite phosphate sulfate sulfite – Be careful not to confuse a formula such as NO2 with NO2 . The latter formula is a molecular compound as is electrically neutral, whereas the former is a polyatomic ion. The anion does not exist by itself, and will alway be a part of an ionic compound such as NaNO2 . _______________________________________________________________________________ EXAMPLE 2.3 Electrons and Protons in Ions Determine the number of protons and electrons in the ions represented by K+, Fe 3+ , and I– . •Method of Solution The net positive charge of the K+ ion is due to the loss of an electron from a K atom. A K atom has 19 p+ and 19 e– . Therefore, the K+ ion has 19 p+ and 18 e– . The net positive charge of the Fe 3+ ion is due to the loss of three electrons from an Fe atom. An Fe atom has 26 p+ and 26 e– . Therefore, the Fe3+ ion has 26 p+ and 23 e– . The net negative charge of the I– ion is due to a gain of one electron by an I atom. An I atom has 53 p+ and 53 e– . Therefore, the I– ion has 53 p+ and 54 e– . _______________________________________________________________________________ EXERCISES 8. Identify the following as elements or compounds. a. ClF3 b . HCl c. O 3 d. I2 e. Se f. NaI 9. Give an example of: a. a monatomic cation. b. a polyatomic anion 10. Give the number of protons and electrons in each of the following ions. a. Li + b . Sr 2+ c. Fe3+ d. N3– e. Se2– f . Cl – Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Atoms, Molecules, and Ions / 2 9 CHEMICAL FORMULAS STUDY OBJECTIVES 1. 2. Distinguish between molecular and empirical formulas, and write empirical formulas when molecular formulas are given. Write the formulas of inorganic compounds from their names. Molecular Formulas. A molecule of a compound is formed by the combination of atoms of different elements. It is the smallest unit of a compound that has properties of the compound. The molecular formula indicates the numbers of atoms of each element in a molecule of the compound. The formula for ethylene glycol is C2 H6 O2 because there are two carbon atoms, six hydrogen atoms, and two oxygen atoms in a molecule of the compound. All molecules are not necessarily compounds, since some elements exist as molecules. H 2 , N 2 , O 2 , F 2 , Cl 2 , Br2 , and I2 are molecular formulas just as CH 4 , NH 3 , and C 2 H6 O2 are molecular formulas. Figure 2.9 of the text shows some molecular models of several molecular compounds. The e mpirical formula of a compound gives the simplest whole-number ratio of the atoms of the elements making up the compound. They are written by reducing the subscripts in molecular formulas to the smallest whole-number ratio. Benzene, for example, has the molecular formula C6 H6 . Dividing 6 by 6 gives the simplest ratio. The empirical formula is just CH, the simplest ratio of C atoms to H atoms being 1 : 1. The empirical formula of many compounds is the same as the molecular formula. For ammonia (NH 3 ) and methane (CH4 ), for example, the ratios of subscripts cannot be reduced and still remain whole numbers. Ionic compounds consist of individual positive and negative ions. Magnesium chloride (MgCl 2 ), for example, consists of individual magnesium and chloride ions. The formula MgCl2 just gives the ratio of Mg2+ ions to Cl– ions in any sample of magnesium chloride. Therefore, magnesium chloride is a compound in which the ratio of Mg 2+ ions to Cl– ions is 1 : 2. See Figure 2.11 of the text which shows the crystal structure of sodium chloride. The formulas of ionic compounds are always the same as their empirical formulas. Ionic compounds consist of positive and negative ions, rather than molecules. Their formulas already give the smallest whole-number ratio of the ions in the compound. Formulas of Ionic Compounds. The names and formulas of the cations and anions in can be used to write formulas of compounds. To write the correct formula for an ionic compound, you must know the ionic charges of the cations and anions. See Table 2.3 of the textbook. The main rule is that all compounds must be electrically neutral. Therefore, the ratio of cations to anions must be such that the sum of the positive charges of cations and the negative charges of anions must be equal to zero. Consider the following examples: 1. 2. Back Potassium bromide. The ions are K+ and Br– . Adding the charges of these ions, we get (+1) + (–1) = 0 The ratio of K+ to Br– must be 1 : 1. The formula is KBr. Potassium sulfide. The ions are K+ and S2– . When the charge on the positive ion is less than the charge on the negative ion, more positive ions will be required to balance the negative charge. If there are two K+ ions for each S2– ion, the compound will be neutral. 2 cations(+1) + 1 anion(–2) = (+2) + (–2) = 0. The formula is K 2 S Notice that the ionic charge is written in parenthesis preceded by the number of ions. Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 3 0 / Atoms, Molecules, and Ions 3. Magnesium sulfate. The ions are Mg 2+ and SO2 – . A compound with one Mg 2+ ion and one SO 2 – ion 4 4 will be neutral. (+2) + (–2) = 0 The formula is MgSO 4 . 4. Aluminum oxide. The ions are Al3+ and O2– . Two Al3+ ions have a charge of +6. It takes three O2– ions to have a charge of –6. 2 cations(+3) + 3 anions(–2) = (+6) + (–6) = 0 The formula is Al 2 O3 . _______________________________________________________________________________ EXAMPLE 2.4 Empirical Formulas What are the empirical formulas of the following compounds? a. B2 H6 b . P 4 O10 •Method of Solution To find the simplest whole-number ratio of atoms divide the smaller subscript into the larger one. a. The simplest ratio of H to B is 6 to 2 which 3 to 1. Therefore the empirical formula is BH3 . b. The simplest ratio of O to P is 10 to 4 which reduces to 5 to 2. Notice that you must keep a whole-number ratio. An atom ratio.of 2.5 to 1 would not be physically reasonable. The empirical formula is P2 O5 . _______________________________________________________________________________ EXAMPLE 2.5 Formulas of Ionic Compounds Write the formulas of the following compounds. a. magnesium chloride b. magnesium oxide c. magnesium phosphate •Method of Solution The formula must represent an electrically neutral grouping of ions. Identify the ions first. a. The ions in magnesium chloride are Mg2+ and Cl– ions. Two chloride ions are needed for each magnesium ion. The formula is MgCl2 . 1 cation(+2) + 2 cations(–1) = 0 b. The ions in magnesium oxide are Mg 2+ and O2– . In this case just one cation and one anion are will give electrical neutrality. (+2) + (–2) = 0 The formula is MgO. c. The ions in magnesium phosphate are Mg2+ and PO4 3– . In this case 2 anions have a –6 charge and it will take 3 cations to have the needed +6 charge. 3 cations (+2) + 2 anions(–3) = 0 The formula is Mg3 (PO4 )2 . _______________________________________________________________________________ EXERCISES 11. Styrene has the molecular formula C8 H8 . What is its empirical formula? 12. What are the empirical formulas of the following compounds? a. N 2 O4 . b. C 4 H8 c . AlCl3 d. Fe2 O3 e . S 2 F 10 13. Which of the following compounds are likely to be ionic? a. KCl b. CH4 c . AlCl3 d. SO2 e . MgO f. CCl 4 Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Atoms, Molecules, and Ions / 3 1 14. Write formulas of the following binary compounds. a. barium chloride b. magnesium nitride c. iron(III) oxide d. iron(II) fluoride 15. Write the formulas for the following compounds. a. ammonium chloride b. sodium phosphate c. potassium sulfate d. calcium carbonate e. potassium hydrogen carbonate f. magnesium nitrite g. sodium nitrate h. ammonium perchlorate i. strontium hydroxide j. copper(II) cyanide NAMING INORGANIC COMPOUNDS STUDY OBJECTIVES 1. Write the names of several types of inorganic compounds given their formulas. Chemical Names. N omenclature refers to the naming of chemical substances. Chemical nomenclature and formulas form the basis of the language of chemistry. In the textbook, inorganic compounds are divided into four categories: ionic compounds, molecular compounds, acids and bases, and hydrates. Binary compounds are those with only two elements and are the easiest to name. Learning chemical nomenclature is like learning a new language. First, you build up a vocabulary of words, and then you learn rules for putting words together meaningfully. To learn a vocabulary requires strict memorization, and so treat the names of cations and anions as you would words in a new language. Get a package of 3 × 5 cards, and write the name of an ion on one side of each card and its chemical formula on the other side. Table 2.3 in the textbook has 20 cations and 26 anions. If you learn their names and formulas, you will be able to name 572 possible chemical compounds! Ionic Compounds. In binary ionic compounds, the metal ion is positively charged and is named first. The name of a simple cation is the same as the name of the element. The negative ion name appears second, and is derived from the name of the nonmetallic element in which an -ide ending replaces the usual ending of the element name. The following formulas and names illustrate naming binary ionic compounds: NaI CaF 2 SnO sodium iodide calcium fluoride tin oxide Some metals can form more than one kind of cation. This is true of tin, for example, where Sn2+ and Sn4+ are known. These are called tin(II) and tin(IV), respectively. SnO in the preceding list should more correctly be named tin(II) oxide. This distinguishes it from SnO 2 , tin(IV) oxide, which is a different compound. The Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 3 2 / Atoms, Molecules, and Ions transition metals commonly exhibit a variety of ionic charges. The system in which a Roman numeral is used to designate the ionic charge of the cation was developed by Alfred Stock and is called the Stock system. MnO2 Mn2 O3 manganese(IV) oxide manganese(III) oxide In ternary ionic compounds, those consisting of three elements, the metal ion is named first followed by the name of the polyatomic anion. Some examples are: Na2 SO4 Fe(NO2 )2 Mg(NO3 )2 sodium sulfate iron(II) nitrite magnesium nitrate These polyatomic ions are examples of oxoanions. They contain a central atom, such as sulfur or nitrogen, that is bonded to several oxygen atoms. Some elements form several oxoanions that differ only in the number of oxygen atoms. In the case of an element that can have two oxoanions, the name of the anion with more oxygen atoms ends in -ate, and the name of the anion with less oxygen atoms ends in -ite. See nitrate and nitrite in the preceding examples. Molecular Compounds. In binary molecular compounds, the more metallic element is named first, followed by the less metallic elements, whose name ends in -ide. When a pair of elements can form several compounds, questions about which compound is being referred to are removed by adding prefixes to denote the number of atoms of each element. The prefixes taken from Greek are given in Table 2.4 of the textbook. Some examples are: SF 4 SF 6 PCl 3 P 2 Cl 4 sulfur tetrafluoride sulfur hexafluoride phosphorus trichloride diphosphorus tetrachloride Acids and Bases. Binary acids are compounds of hydrogen and a nonmetal. Many hydrogen compounds have completely different properties when they are in the pure gaseous or liquid state than when they are dissolved in water. In the gaseous state, HCl and H2 S, for instance, are molecular compounds called hydrogen chloride and hydrogen sulfide, respectively, whereas water solutions of these compounds contain hydrogen ions (H+) and anions and have slightly different names. A cids are substances that release H+ ions when they dissolve in water. In naming binary acids, the word hydrogen is replaced by hydro-, and the -ide ending of the anion name is replaced with -ic acid. For example: HCl(g) HCl(aq) H2 S(g) H2 S(aq) hydrogen chloride hydrochloric acid hydrogen sulfide hydrosulfuric acid The oxoacids are compounds containing hydrogen and oxoanions. The acids of nitrate and sulfate are: HNO3 H2 SO4 nitric acid sulfuric acid Note that the -ate ending of the anion name has been changed to -ic acid. For oxoanions with fewer oxygen atoms, the -ite ending of the anion name is changed to -ous acid for the oxyacid. HNO2 H2 SO3 Back Forward nitrous acid sulfurous acid Main Menu TOC Study Guide TOC Textbook Website MHHE Website Atoms, Molecules, and Ions / 3 3 Bases are substances that release hydroxide ions (OH– ) in aqueous solution. The metal hydroxides are named as if they were binary compounds and employ the -ide ending. LiOH Ca(OH)2 lithium hydroxide calcium hydroxide ___________________________________________________________________________ EXAMPLE 2.6 Naming Compounds Name the following compounds according to the Stock system: a. CuBr b. CuSO 4 •Method of Solution In the Stock system the symbol of the metallic element in a compound is followed by a Roman numeral derived from the charge of the cation. a. Assuming bromide to be –1 as in the Br– ion, Cu must be a +1 ion in order to have a neutral compound. Therefore, CuBr is named copper(I) bromide. b. In this case copper must be a +2 ion and the compound is copper(II) sulfate. ____________________________________________________________________ EXERCISES 16. Name the following compounds: a. K 3 N b. Ag 2 CO3 c. Mg(OH)2 f. Fe(NO3 )2 g . CaSO 4 ·2H2 O d. NaCN 17. Name the following compounds: a. PCl5 b . SO 3 c . P 4 O10 d. N 2 O e. NH4 I e. NO2 18. Name the following acids: a. HNO3 b. HNO2 c. HBr d. HCN e. H2 S _______________________________________________________________________________ CONCEPTUAL QUESTIONS 1. 2. What is the difference between an atom and a molecule? 3. Do the symbols O3 and 3O signify the same thing? Explain. 4. Back How does the discovery of isotopes affect Dalton's postulate that atoms of the same element all have identical mass. Why are the chemical formulas of ionic compounds the same as the empirical formulas? Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 3 4 / Atoms, Molecules, and Ions PRACTICE TEST 1. Complete the following isotope table: Number of Number of Number of Mass Name Symbol Protons Electrons Neutrons Number _________________________________________________________________________ 23 Na Sodium 11 _____ 12 23 40 Ar _______ ______ _____ 22 _____ 75 As Arsenic ______ _____ _____ _____ Lead _____ ______ _____ _____ _____ ______ _____ ______ _____ 20 _____ _________________________________________________________________________ 56 2. a. What is the total number of subatomic particles (protons, neutrons, and electrons) in an atom of 26 Fe? b. What is the mass number of a copper atom that has 35 neutrons? 3. Give the number of protons, neutrons, and electrons in atoms of the following. 17 1 07 2 22 a. 8 O b . 4 7 A g c. 8 6 R n 4. Which of the following are isotopes of element X? 46 20 43 46 20 X , 46 X , 20 X , 43 X 5. Take the mass of an atom to be the sum of the masses of its protons, neutrons, and electrons. Determine 12 the percentages by mass of the protons, the neutrons, and the electrons in a 6 C atom. 6. Which of the following molecules are forms of a pure element? P 4 , He, N2 , O 3 , N 2 O3 7. What is the empirical formula of each of the following compounds? a. C 6 H8 O6 b. C 2 H2 c. Hg 2 Cl 2 d. H2 O2 e. C 2 H2 O4 f. MgCl2 8. Write the formulas for the following compounds. a. calcium hypochlorite b. mercury(II) sulfate c. barium sulfite d. zinc oxide e. dinitrogen oxide f. sodium carbonate g. copper(II) sulfide h. lead(IV) oxide 9. Name the following compounds: a. Na2 HPO4 b. HI(gas) c. P 4 O6 d. LiNO3 f. Sr(NO2 )2 g. NaHCO3 h. K 2 SO3 10. Name the following acids: a. H 2 SO3 b. HClOc. HClO4 Back Forward Main Menu TOC e. HI(solution) i. Na 3 PO4 d. H3 PO4 j. Al(OH)3 e. HCN Study Guide TOC Textbook Website MHHE Website Atoms, Molecules, and Ions / 3 5 ANSWERS Exercises 2. Protons, electrons, and neutrons. Protons have one unit of positive charge, and electrons have one unit of negative charge. Neutrons have no charge. Protons and neutrons are found in the nucleus. 3. The evidence was related to the direction of scattering of the positively charged alpha particles by gold atoms. The fact that 1 in 10,000 alpha particles was deflected "backward" meant that all the positive charge of the gold atom must be concentrated into a very small part of the atom. 4. 200 m 5. 24 6. 62 67 1 20 7. a. 30 Z n b. 51 S b 8. a. compound b. compound c. element d. element e. element f. compound 2– 9. a. Fe 3+ iron(III) ion b. CO3 carbonate ion 10. a. 3p, 2e b. 38p, 36e c. 26p, 23e d. 7p, 10e e. 34p, 36e f. 17p, 18e 11. CH 12. a. NO2 b . CH 2 c . AlCl3 d. Fe2 O3 e . SF 5 13. a, c, e. 14. a. BaCl2 b . Mg 3 N2 c . Fe 2 O3 d. FeF2 15. a. NH4 Cl b. Na3 PO4 c. K2 SO4 d. CaCO3 e. KHCO3 f. Mg(NO2 )2 g. NaNO3 h. NH4 ClO 4 i. Sr(OH)2 j. Cu(CN)2 16. a. potassium nitride b. silver carbonate c. magnesium hydroxide d. sodium cyanide e. ammonium iodide f. iron(II) nitrate g. calcium sulfate dihydrate 17. a. phosphorus pentachloride b. sulfur trioxide c. tertaphosphorus decoxide d. dinitrogen monoxide e. nitrogen dioxide 18. a. nitric acid b. nitrous acid c. hydrobromic acid d. hydrocyanic acid e. hydrosulfuric acid Conceptual Questions 1. 2. 3. 4. Back With regard to atomic mass there is no real difference because the percent abundances turn out to be essentially constant anywhere on Earth. This means that the average atomic mass of an element is constant and that as far as mass is concerned the atoms in a sample of an element act on average as if they have the same mass. An atom is the smallest particle of an element. A molecule is a particle that is made of two or more atoms. When the atoms in an molecule are identical we have a molecular form of an element. When the atoms are different we have the smallest particle of a compound. No. The symbol O3 represents a molecule of ozone, a molecular form of oxygen. 3O represents 3 separate oxygen atoms. The empirical formula of a compound gives the simplest whole-number ratio of the atoms of the elements making up the compound. Ionic compounds consist of positive and negative ions in the ratio needed to give an electrically neutral substance. The subscrips in the formula represent the simplest ratio of positive to negative ions. Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 3 6 / Atoms, Molecules, and Ions Practice Test 1. _______________________________________________________________________________ Number of Number of Number of Mass Name Symbol Protons Electrons Neutrons Number _______________________________________________________________________________ 23 Na Sodium 11 11 12 23 40 Ar Argon 18 18 22 40 75 As Arsenic 33 33 42 75 206 Pb Lead 82 82 124 206* 40 Ca Calcium 20 20 20 40* _________________________________________________________________________________ *Arbitrarily chosen isotope; others are acceptable 2. 3. 4. 5. 6. 7. a. 108 b. 64 a. 8p, 9n, 8e b. 47 p, 60n, 47e c. 86p, 136, 86e 46 43 20 X & 20 X 50% proton mass, 50% neutron mass, & 2.8 × 10 –2 % electron mass P 4 , He, N2 , O 3 a. C3 H4 O3 b. CH c. HgCl d. HO e CHO 2 f. MgCl2 8. a. Ca(OCl)2 b. HgSO4 c. BaSO3 d. ZnO e. N 2 O f. Na 2 CO3 g. CuS h. PbO2 9. a. sodium hydrogen phosphate d. lithium nitrate g. sodium hydrogen carbonate j. aluminum hydroxide b. hydrogen iodide e. hydroiodic acid h. potassium sulfite b. hypochlorous acid e. hydrocyanic acid 10. a. sulfurous acid d. phosphoric acid c. tetraphosphorus hexoxide f. strontium nitrite i. sodium phosphate c. perchloric acid ________________________________________________________________________________ Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website ...
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This note was uploaded on 09/15/2009 for the course CHEM 102 taught by Professor Bastos during the Spring '08 term at Adelphi.

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