chapter 2 student summer 2020

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Chapter 2: Homework 2.1 Atomic Theory of Matter The theory that atoms are the fundamental building blocks of matter reemerged in the early nineteenth century, championed by John Dalton. Law of Conservation of Mass : 2.2 The Discovery of Atomic Structure Particles with the same charge repel one another, whereas particles with unlike charges attract one another. The Electron J. J. Thomson (1897) Streams of negatively charged particles were found to emanate from cathode tubes, causing fluorescence. Thomson measured the charge/mass ratio of the electron to be 1.76 ´ 10 8 coulombs/gram (C/g). http://www.youtube.com/watch? v=XU8nMKkzbT8 1

Chapter 2: Homework Millikan Oil-Drop Experiment Robert Millikan (University of Chicago) determined the charge on the electron in 1909. http://www.youtube.com/watch?v=XMfYHag7Liw Radioactivity • Radioactivity • It was first observed by Henri Becquerel. • Marie and Pierre Curie also studied it. • Three types of radiation were discovered by Ernest Rutherford: 1. 2. 3. The Atom, circa 1900 Thomson’s “plum pudding” model a positive sphere of matter with negative electrons imbedded in it. Discovery of the Nucleus Rutherford shot a particles at a thin sheet of gold foil; observed the pattern of scatter of the particles. http://www.youtube.com/watch?v=5pZj0u_XMbc Since some particles were deflected at large angles, Thomson’s model could not be correct. Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom. Most of the volume of the atom is empty space. 2

Chapter 2: Homework 2.3 The Modern View of Atomic Structure • Protons - Rutherford in 1919 • Neutrons - James Chadwick in 1932 • Protons and electrons are the only particles that have a charge • Protons and neutrons have essentially the same mass • mass of an electron is so small we ignore it Symbols of Elements Elements are symbolized by one or two letters. atomic number , Z ,: mass number: Isotopes: Sample Exercise 2.1: How many protons, neutrons, and electrons are in (a) a 138 Ba atom, (b) an atom of phosphorus-31? Sample Exercise 2.2: Give the complete chemical symbol for the atom that contains 82 protons, 82 electrons, and 126 neutrons. 2.4 Atomic Weights Atomic and molecular masses can be measured with great accuracy using a mass spectrometer. http://www.youtube.com/watch?v=OKxRx0ctrl0 Average mass Atomic mass unit : amu – exactly 12 amu in 1 atom of 12 C 3

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Chapter 2: Homework 1 amu = 1.66054 x 10 -24 g or 1g = 6.022x10 23 amu Sample Exercise 2.3: Naturally occurring chlorine is 75.78% 35 Cl (atomic mass 34.969 amu) and 24.22% 37 Cl (atomic mass 36.966 amu). Calculate the atomic weight of chlorine. 2.5 Periodic Table Elements are arranged in order of atomic number • The rows on the periodic chart are periods. • Columns are groups. • Elements in the same group have similar chemical properties. When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities. 4

Chapter 2: Homework Sample Exercise 2.4: Which two of these elements would you expect to show the greatest similarity in chemical and physical properties: B, Ca, F, He, Mg, P? 2.6 Molecules and Molecular Compounds Chemical Formulas: subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound. Diatomic Molecules (2 atoms) 7 naturally occurring: H 2 , O 2 , N 2 , F 2 , Cl 2 , Br 2 , I 2 Molecular compounds Types of Formulas • Empirical formulas: • Molecular formulas: • Structural formulas: Sample Exercise 2.5: Give the empirical formula for diborane , whose molecular formula is B 2 H 6 . 2.7 Ions and Ionic Compounds When atoms lose or gain electrons, they become ions . • Cations are ________ and are formed by elements on the _________ side of the periodic chart. • Anions are ________ and are formed by elements on the __________ side of the periodic chart. Ionic Bonds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals. 5

Chapter 2: Homework Predicting Ionic Charges Try to achieve noble gas configurations Ionic Compounds Ionic compound-made of cations and anions; usually combinations of metals and nonmetals (NaCl) Molecular compounds – generally composed of nonmetals only, (H 2 O) Sample Exercise 2.6: Which of these compounds would you expect to be ionic? Which are molecular? (a) N 2 O _______ (b) Na 2 O_______ (c) CaCl 2 ______ (d) SF 4 _______ (e) CBr 4 ________ (f) FeS ________ (g) P 4 O 6 ________ (h) PbF 2 ______ Writing Formulas Because compounds are electrically neutral, one can determine the formula of a compound this way: • The charge on the cation becomes the subscript on the anion. • The charge on the anion becomes the subscript on the cation. • If these subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor. Sample Exercise 2.7: Write the empirical formula for the compound formed by (a) Na + and PO 4 3– _____________ (b) Zn 2+ and SO 4 2– _____________ (c) Fe 3+ and CO 3 2– ______________ (d) Ca 2+ and O 2- ______________ 6

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Chapter 2: Homework (e) Al 3+ and Cl - _______________ 7

Chapter 2: Homework 2.8 Inorganic Nomenclature • Write the name of the cation. • If the anion is an element, change its ending to - ide ; if the anion is a polyatomic ion, simply write the name of the polyatomic ion. • • If the cation can have more than one possible charge (transition metals), write the charge as a Roman numeral in parentheses • When there are two oxyanions involving the same element: – The one with fewer oxygens ends in - ite . – The one with more oxygens ends in - ate . • Central atoms on the second row can bond to at most three oxygens; those on the third row take up to four. • Charges increase as you go from right to lef . • The one with the second fewest oxygens ends in - ite . • The one with the second most oxygens ends in -ate . • The one with the fewest oxygens has the prefix hypo - and ends in - ite . • The one with the most oxygens has the prefix per - and ends in - ate . Sample Exercise 2.8: The formula for the bromate ion is analogous to that for the chlorate ion. Write the formula for the hypobromite and bromite ions. 8

Chapter 2: Homework Sample Exercise 2.9: Name the ionic compounds (a) K 2 SO 4 ____________________________ (b) Ba(OH) 2 _____________________________ (c) FeCl 3 _____________________________ (d) NH 4 Br _______________________________ (e) Cr 2 O 3 _____________________________(f) Co(NO 3 ) 2 _____________________________ Sample Exercise 2.10: Give the chemical formulas for (a) magnesium sulfate ____________ (b) silver sulfide______________(c) lead(II) nitrate ___________ Acid Nomenclature • If the anion in the acid ends in - ide , change the ending to - ic acid and add the prefix hydro - . – HCl: ___________________ – HBr: ___________________ – HI: ____________________ • If the anion in the acid ends in - ite , change the ending to - ous acid . – HClO: ____________________ – HClO 2 : ___________________ • If the anion in the acid ends in - ate , change the ending to - ic acid . – HClO 3 : ____________________ – HClO 4 : ____________________ Sample Exercise 2.11: Name the acids (a) HCN ____________________________(b) HNO 3 ____________________________ (c) H 2 SO 4 ___________________________(d) H 2 SO 3 ___________________________ Practice Exercise 2.12: Give the chemical formulas for (a) hydrobromic acid _________________________(b) carbonic acid ___________________________ .(c) Hypoiodous acid __________________________ (d) phosphoric acid________________________ Nomenclature of Binary Compounds • The less electronegative atom is usually listed first . (left side of pt) • If both elements are in the same group, the lower one is named first 9

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Chapter 2: Homework • A prefix is used to denote the number of atoms of each element in the compound ( mono - is not used on the first element listed). • The ending on the more electronegative element is changed to - ide . – CO 2 : – CCl 4 : • If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one. N 2 O 5 : Sample Exercise 2.13: Name the compounds (a) SO 2 ________________________________(b) PCl 5 _______________________________ (c) Cl 2 O 3 __________________________________ Sample Exercise 2.14: Give the chemical formulas (a) silicon tetrabromide _________________________ (b) disulfur dichloride ______________________________ 2.9 Nomenclature of Organic Compounds Organic chemistry : study of______________________ alkanes The simplest hydrocarbons (compounds containing only carbon and hydrogen) The first part of the names correspond to the number of carbons ( meth- = 1, eth- = 2, prop- = 3, but- =4, etc.). When a hydrogen in an alkane is replaced with something else (a functional group , like -OH in the compounds above), the name is derived from the name of the alkane. The ending denotes the type of compound. An alcohol ends in - ol . Sample Exercise 2.15: (a) What is the molecular formula of butane, the alkane with four carbons? (b) What are the name and molecular formula of an alcohol derived from butane? 2.1 and 2.2: Atomic Theory and the Discovery of Atomic Structure 10

Chapter 2: Homework 1. (2.15 Brown) How did Rutherford interpret the following observations made during his a particle scattering experiments? a. Most a particles were not appreciably deflected as they passed through the gold foil. b. A few a particles were deflected at very large angles. 2.3 and 2.4: Modern View of Atomic Structure; Atomic Weights 2. (2.20 Brown) Determine whether each of the following statements is true or false. If false, correct the statement to make it true: a. The nucleus has most of the mass and comprises most of the volume of an atom b. Every atom of a given element has the same number of protons c. The number of electrons in a neutral atom equals the number of neutrons in the atom 3. (2.22 Brown) a. Which two of the following are isotopes of the same element: 31 16 X, 31 15 X, 32 16 X? b. What is the identity of the element whose isotopes you have selected? 4. (2.23 Brown) How many protons, neutrons, and electrons are in the following atoms: a. 40 Ar b. 65 Zn c. 70 Ga d. 80 Br e. 184 W f. 243 Am 5. (2.26 Brown) Fill in the gaps in the following table, assuming each column represents a neutral atom Symbol 65 Zn Protons 38 92 Neutrons 58 49 Electrons 38 36 Mass Number 81 235 6. (2.27 Brown) Write the correct symbol, with both superscript and subscript, for each of the following. Use the list of elements inside the front cover as needed: a. The isotope of platinum that contains 118 neutrons b. The isotope of krypton with mass number 84 c. The isotope of arsenic with mass number 75 d. The isotope of magnesium that has an equal number of protons and neutrons 7. (2.32 Brown) Rubidium has two naturally occurring isotopes, rubidium-85 (atomic mass = 84.9118 amu; abundance = 72.15%) and rubidium-87 (atomic mass = 86.9092 amu; abundance = 27.85%). Calculate the atomic weight of rubidium. 8. (2.35 Brown) Naturally occurring magnesium has the following isotopic abundances: 11

Chapter 2: Homework Isotope Abundance Atomic mass (amu) 24 Mg 78.99% 23.98504 25 Mg 10.00% 24.98584 26 Mg 11.01% 25.98259 What is the average atomic mass of mg? 9. (2.87 Brown) The element oxygen has three naturally occurring isotopes, with 8, 9, and 10 neutrons in the nucleus, respectively. a. Write the full chemical symbols for these three isotopes b. Describe the similarities and differences between the three kinds of atoms of oxygen 10. (2.96 Brown) The explosion of an atomic bomb releases many radioactive isotopes, including strontium-90. Considering the location of strontium in the periodic table, suggest a reason for the fact that this isotope is particularly harmful to humans. More book problems: 17, 19, 21, 25, 29, 31, 87, 90; sample exercises 2.1, 2.2, 2.3, 2.4 2.5, 2.6, and 2.7: The Periodic Table; Molecules and Ions 11. (2.38 Brown) Locate each of the following elements in the periodic table; give its name and atomic number, and indicate whether it is a metal, metalloid, or nonmetal: a. Li b. Sc c. Ge d. Yb e. Mn f. Sb g. Xe 12. (2.39 Brown) For each of the following elements, write its chemical symbol, determine the name of the group to which it belongs (Table 2.3), and indicate whether it is a metal, metalloid, or nonmetal: a. Potassium b. Iodine c. Magnesium d. Argon 13. (2.43 Brown) Write the empirical formula corresponding to each of the following molecular formulas a. Al 2 Br 6 b. C 8 H 10 c. C 4 H 8 O 2 d. P 4 O 10 e. C 6 H 4 Cl 2 f. B 3 N 3 H 6 14. (2.45 Brown) How many hydrogen atoms are in each of the following: a. C 2 H 5 OH b. Ca(CH 3 COO) 2 c. (NH 4 ) 3 PO 4 12

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Chapter 2: Homework 15. (2.50 Brown) Fill in the gaps in the following table: Symbol 31 P 3- Protons 34 50 Neutrons 45 69 118 Electrons 46 76 Net charge 2- 3+ 16. (2.53 Brown) Using the periodic table to guide you, predict the chemical formula and name of the compound formed by the following elements: a. Ca and F b. Li and H c. Al and I d. K and S 17. (2.56 Brown) Predict the chemical formulas of the compounds formed by the following pairs of ions: a. Cr 3+ and Br - b. Fe 3+ and O 2- c. Hg 2 2+ and CO 3 2- d. Ca 2+ and ClO 3 - e. NH 4 + and PO4 3- 18. (2.58 Brown) Complete the table by filling in the formula for the ionic compound formed by each pair of cations and anions, as shown for the first pair. Ion Na + Ca 2+ Fe 2+ Al 3+ O 2- Na 2 O NO 3 - SO 4 2- PO 4 3- 19. (2.60 Brown) Which of the following are ionic, and which are molecular? a. PF 3 b. NaI c. SCl 2 d. Ca(NO 3 ) 2 e. FeCl 3 f. AlP g. CoCO 3 h. N 2 O 4 More book problems: 37, 49, 51, 55, 57, 59 sample exercises 2.5, 2.6, 2.7, 2.8, 2.9, 2.10 13

Chapter 2: Homework 2.8 and 2.9: Naming Inorganic Compounds; Organic Molecules 20. (2.61 Brown) Give the chemical formula for a. Chlorite ion b. Chloride ion c. Chlorate ion d. Perchlorate ion e. Hypochlorite ion 21. (2.64 Brown) Give the names and charges of the cation and anion in each of the following compounds: a. CuS b. Ag 2 SO 4 c. Al(ClO 3 ) 3 d. Co(OH) 2 e. PbCO 3 22. (2.66 Brown) Name the following ionic compounds a. KCN b. NaBrO 2 c. Sr(OH) 2 d. CoS e. Fe 2 (CO 3 ) 3 f. Cr(NO 3 ) 3 g. (NH 4 ) 2 SO 3 h. KMnO 4 i. Ag 2 Cr 2 O 7 23. (2.68 Brown) Give the chemical formula for each of the following ionic compounds a. Sodium phosphate b. Zinc nitrate c. Barium bromate d. Iron(II) perchlorate e. Cobalt(II) carbonate f. Chromium(III) acetate g. Potassium sulfate 24. (2.70 Brown) Provide the name or chemical formula, as appropriate, for each of the following acids: a. Hydroiodic acid b. Chloric acid c. Nitrous acid d. H 2 CO 3 e. HClO 4 f. CH 3 COOH 25. (2.71 Brown) Give the name or chemical formula, as appropriate, for each of the following binary molecular substances: a. SF 6 b. IF 5 c. XeO 3 d. Dinitrogen tetroxide e. Hydrogen cyanide f. Tetraphosphorus hexasulfide 26. (2.77 Brown) a. What is a functional group? 14

Chapter 2: Homework b. What functional group characterizes an alcohol? 27. (2.104 Brown) Many familiar substances have common, unsystematic names. For each of the following, give the correct systematic name: a. Saltpeter, KNO 3 b. Soda ash, Na 2 CO 3 c. Lime, CaO d. Muriatic acid, HCl e. Epsom salts, MgSO 4 f. Milk of magnesia, Mg(OH) 2 More book problems: 63, 65, 69, 73, 75, 104 sample exercises 2.11, 2.12, 2.13, 2.14, 2.15, 2.16, Answers 1. a. Most of the volume of an atom is empty space in which electrons move. Most alpha particles passed through this space. The path of the massive alpha particle would not be significantly altered by interaction with a “puny” electron b. Most of the mass of an atom is contained in a very small, dense area called the nucleus. The few alpha particles that hit the massive, positively charge gold nuclei were strongly repelled and essentially deflected back in the direction they came from. c. The Be nuclei have much smaller volume and positive charge than the Au nuclei; the charge repulsion between alpha particles and the Be nuclei will be less, and there will be 15

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Chapter 2: Homework fewer direct hits because the Be nuclei have an even smaller volume than the Au nuclei. Fewer alpha particles will be scattered in general and fewer will be strongly back scattered 2. a. False. The nucleus has most of the mass but occupies very little of the volume of an atom b. True c. False. The number of electrons in a neutral atom is equal to the number of protons in the atom 3. a. 31 16 X and 32 16 X are isotopes of the same element, because they have identical atomic numbers b. Sulfur, S, atomic number = 16 4. a. Ar 18p, 22n, 18e b. Zn 30p, 35n, 30e c. Ga 31p, 39n, 31e d. Br 35p, 45n, 35e e. W 74p, 110n, 74e f. Am 95p, 148n, 95e 5. Symbol 65 Zn 96 Sr 87 Sr 36 Kr 235 U Protons 30 38 38 36 92 Neutrons 35 58 49 45 143 Electrons 30 38 38 36 92 Mass Number 65 96 87 81 235 6. a. 196 78 Pt b. 84 36 Kr c. 75 33 As d. 24 12 Mg 7. 85.47 amu 8. 24.31 amu 9. a. 16 8 O, 17 8 O, 18 8 O b. All isotopes are atoms of the same element, oxygen, with the same atomic number (Z=8), 8 protons in the nucleus and 8 electrons. Elements with similar electron 16

Chapter 2: Homework arrangements have similar chemical properties. Since the 3 isotopes all have 8 electrons, we expect their electron arrangements to be the same and their chemical properties to be very similar, perhaps identical. Each has a different number of neutrons (8,9, or 10), a different mass number (16,17,or 18) and thus a different atomic mass. 10. Strontium is an alkaline earth metal, similar in chemical properties to calcium and magnesium. Calcium is ubiquitous in biological organisms, humans included. It is a vital nutrient required for formation and maintenance of healthy bones and teeth. As such, there are efficient pathways for calcium uptake and distribution in the body, pathways that are also available to chemically similar strontium. Harmful strontium imitates calcium and then behaves badly when the body tries to use it as it uses calcium. 11. a. Lithium, 3, metal b. Scandium, 21, metal c. Germanium, 32, metalloid d. Ytterbium, 70, metal e. Manganese, 25, metal f. Antimony, 51, metalloid g. Xenon, 54, nonmetal 12. a. K, alkali metal, metal b. I, halogen, nonmetal c. Mg, alkaline earth metal, metal d. Ar, noble gas, nonmetal 13. a. AlBr 3 b. C 4 H 5 c. C 2 H 4 O d. P 2 O 5 e. C 3 H 2 Cl f. BNH 2 14. a. 6 b. 6 c. 12 15. Symbol 31 P 3- 79 Se 2- 119 Sn 4+ 197 Au 3+ Protons 15 34 50 79 Neutrons 16 45 69 118 Electrons 18 36 46 76 17

Chapter 2: Homework Net charge 3- 2- 4+ 3+ 16. a. CaF 2 , calcium fluoride b. LiH, lithium hydride c. AlI 3 , aluminum iodide d. K 2 S, potassium sulfide 17. a. CrBr 3 b. Fe 2 O 3 c. Hg 2 CO 3 d. Ca(ClO 3 ) 2 e. (NH 4 ) 3 PO 4 18. Na + Ca 2+ Fe 2+ Al 3+ O 2- Na 2 O CaO FeO Al 2 O 3 NO 3 - NaNO 3 Ca(NO 3 ) 2 Fe(NO 3 ) 2 Al(NO 3 ) 3 SO 4 2- Na 2 SO 4 CaSO 4 FeSO 4 Al 2 (SO 4 ) 3 PO 4 3- Na 3 PO 4 Ca 3 (PO 4 ) 2 Fe 3 (PO 4 ) 2 AlPO 4 19. Molecular: a, c, h Ionic: b, d, e, f, g 20. a. Chlorite, ClO 2 - b. Chloride, Cl - c. Chlorate, ClO 3 - d. Perchlorate, ClO 4 - e. Hypochlorite, ClO - 21. a. Copper, 2+; sulfide, 2- b. Silver, 1+; sulfate, 2- c. Aluminum, 3+; chlorate, 1- d. Cobalt, 2+; hydroxide, 1- e. Lead, 2+; carbonate, 2- 22. a. Potassium cyanide b. Sodium bromite c. Strontium hydroxide d. Cobalt(II) sulfide e. Iron(III) carbonate f. Chromium(III) nitrate g. Ammonium sulfite h. Sodium dihydrogen phosphate i. Potassium permanganate j. Silver dichromate 23. 18

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Chapter 2: Homework a. Na 3 PO 4 b. Zn(NO 3 ) 2 c. Ba(BrO 3 ) 2 d. Fe(ClO 4 ) 2 e. CoCO 3 f. Cr(CH 3 COO) 3 g. K 2 SO 4 24. a. HI b. HClO 3 c. HNO 2 d. Carbonic acid e. Perchloric acid f. Acetic acid 25. a. Sulfur hexafluoride b. Iodine pentafluoride c. Xenon trioxide d. N 2 O 4 e. HCN f. P 4 S 6 26. a. Functional groups are groups of specific atoms that are constant from one molecule to the next. For example, the alcohol functional group is an –OH. Whenever a molecule is called an alcohol, it contains the –OH group. b. –OH 27. a. Potassium nitrate b. Sodium carbonate c. Calcium oxide d. Hydrochloric acid e. Magnesium sulfate f. Magnesium hydroxide 19

chapter 2 student summer 2020

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