Atomic Structure - Chemistry PDF Notes

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CALL : 0799288507  ‫مك ـتبــة‬ ‫تـشـكـيــــل‬ ‫دوسيات توجيهي‬ ‫تصوير وثائق‬ ‫للقرطاسية‬ ‫واألدوات المكتبية‬ ‫أسئلة متوقعة ليلة‬ ‫أسئلة وزارية‬ ‫كل امتحان‬ ‫سابقة‬ ‫خلدا ‪ -‬إشارة البنك العربي‬ ‫‪0796117336 :‬‬ ‫‪0776532229 :‬‬ ‫‪06-5532229 :‬‬ ‫‪0796117336‬‬ ‫‪ :‬مكتبة تشكيل‬ ‫‪[email protected] :‬‬ Content Page CHAPTER 1 Atomic Structure 1 - 27 CHAPTER 2 Bonding 28 – 50 CHAPTER 3 Solid and Liquid 51 – 73 CHAPTER 4 Gases 74 – 89 CHAPTER 5 Equilibrium 90 - 109 CHAPTER 6 Acid and Base 110 - 125 CHAPTER 7 Oxidation Reduction 126 - 135 CHAPTER 8 Organic Chemistry 136 – 141 CHAPTER 9 laboratory 142 - 159 CHAPTER 10 Stoichiometry 160 - 182 CHAPTER 11 Type of Reaction 183 – 192 Test 193 – End Mr. Waseem Allabadi Atomic Structure ACT II Atomic structure The atomic models: 1- Dalton model: The atom is the smallest unit in any matter. It can't be subdivided. 2- Thompson model: Discovery of electrons. The experiment: the cathode ray tube. Conclusion: the ray contains negatively charged particles. The particles are called electrons (e-) 3- Rutherford's model: The discovery of nucleus The experiment: gold-foil experiment. Alpha gun (2 4 He nuclei) 1 Mr. Waseem Allabadi Atomic Structure ACT II  Observations: 1- Most of particles passed without deflection 2- Few were deflected. 3- Very few were reflected.  Conclusions: 1- From observation number 1 most of the atom is empty space. 2- From observation number 2 there is a positive center which affected the alpha particles. 3- From observation number 3 this center is concentrated ( the mass of atom is concentrated in the center) Bohr's model of H atom: Electrons move (swim) in a circular path. * (Ground state) very close to the nucleus. * (Excited state) a little far. When electron moves to a higher level (shell) it absorbs energy (quanta) and this is called absorption. When electron moves to a lower level (shell) it emits energy (quanta) and this is called emission. 2 Mr. Waseem Allabadi Atomic Structure ACT II The subatomic particles: The mass of 1 P+ = the mass of 1 n0 1 The mass of 1 e- = 1860 of the P+ mass Why is the mass of the cell is concentrated in the nucleus? Because the mass of e-s is almost zero. Atomic number: is the number of P+ in any atom. P+ is the identity of the atom. Atomic mass: is the mass of the atom = the mass of its nucleus Mass number (nucleus number) = number of P+ + n0 Element symbol: The smallest number is the atomic number Note: number of P + = number of e- Atoms are neutral (number of P+ = number of e-) Find the number of P+, n0, e- 35 17 Cl: 7 3 Li: 27 13 Al: Ions: are atoms after losing or gaining electrons. Ions are two types: 1- Cations + ions 2- Anions – ions Losing e- + ion because P + are more Gaining e- - ion because e- are more Isoelectronic atoms and ions: have the same number of electrons. N3- , O2- , F- , Na+ , Mg2+ , Ne All of the following belong to the same isoelectronic series EXCEPT: (A) S2- (B) P3- (C) Ar (D) Ca2+ 3 Mr. Waseem Allabadi Atomic Structure ACT II Isotopes: the same element (same atomic or proton number) but different number of neutrons Question: isotope 17 35 X has another isotope… (A) 11 23 Na (B) 2 4 He (C) 17 35 Cl (D) 45 131 Xe How to find the average mass of an element by using its isotopes? Isotopes abundance 20 10 Ne 80% 21 10 Ne 15% 22 10 Ne 5% Average mass: (20×80/100) + (21×15/100) + (22×5/100) = Modern atomic models: Quantum numbers: 1- Principle quantum number (n) (shell or level) Indicates A: size B: energy 4 Mr. Waseem Allabadi Atomic Structure ACT II 2- Angular momentum quantum number (l) (Subshell) or (sublevel) Indicates the shape of the subshell Subshell types: 1- S subshell 2- p subshell 3- d subshell 4- f subshell 3- magnetic quantum number (ml): Indicates the orientation of the orbital 5 Mr. Waseem Allabadi Atomic Structure ACT II 4- spin quantum number (ms): Direction of the spin: A: clockwise B: counterclockwise For the all quantum numbers:  Pauli exclusion principle ( we can't find two electrons share the same 4 quantum numbers ) Shell number (n) number of subshells number of orbitals number of electrons 1 s 1 2 electrons 2 s,p 4 8 electrons 3 s,p,d 9 18 electrons 4 s,p,d,f 16 32 electrons Q: 1- How many e- can be held in the 3rd shell (A) 8 (B) 16 (C) 18 (D) 32 2- How many orbitals are there in the third subshell: (A) 2 (B) 5 (C) 9 (D) 18 (E) 32 6 Mr. Waseem Allabadi Atomic Structure ACT II Electron configuration: Electron configuration by using periodic table: 7 Mr. Waseem Allabadi Atomic Structure ACT II Write the electron configuration for the following: 11 Na: 1S2 2S2 2P6 3S1 16 S: 1S2 2S2 2P6 3S2 3P4 Exceptions: Full and half full subshell is more stable. S2 d4 S1 d5 S2 d9 S1 d10 24 Cr: 1s2 2s2 2p6 3s2 3p6 4s1 3d5 29 Cu: 1s2 2s2 2p6 3s2 3p6 4s1 3d10 Noble gases notation: select the noble gas in the period before your elements period then continue the configuration. Fe: [Ar]18 4s23d6 Zr: [Kr]36 5s24d2 As: [Ar]18 4s2 3d104p3 As 3- : [Ar]18 4s23d104p3+3 8 Mr. Waseem Allabadi Atomic Structure ACT II Electron configuration for ions: 1- Write e- configuration of the atom (both ways). 2- Add or remove e- from the last subshell.(has a larger subshell). Na: 1s2 2s2 2p6 3s1 Na +: 1s2 2s2 2p6 Fe 2+: 1s2 2s2 2p6 3s2 3p6 4s2 3d6 Fe: 1s2 2s2 2p6 3s2 3p6 3d6 O: 1s2 2s2 2p4 O2-:1s2 2s2 2p6 9 Mr. Waseem Allabadi Atomic Structure ACT II Unpaired e- : 1- Write the electron configuration. 2- Take the last subshell to be filled. 3- Use hund's rule to fill e- in orbitals (to get the maximum number of unpaired e-) SAT Q: How many unpaired e- are there in Fe atom. Fe: [Ar]18 4s2 3d6 S: [Ne] 10 3s2 3p4 Mg: [Ne] 10 3s2 Ni: [Ar]18 4s2 3d8 Ne: 0 unpaired of electrons. (Noble gases) in p is always full. Ne: 1s2 2s2 2p6 Valence electrons: electrons in the outermost shell. 10 Mr. Waseem Allabadi Atomic Structure ACT II O: 6 valence: 1s2 2s2 2p4 F: 7 valence: 1s2 2s2 2p5 Mg: 2 valence: 1s2 2s2 2p63s2 Notes: 1) number of valence e- detects the chemical properties of elements. 2) Elements in the same group have the same chemical properties. Group 1: alkali metals Group 2: alkaline earth metals Group 3,4,5, and7 are non-metals Group 7: halogens D group: transition metals Group 8: noble gases B, Si, As, Ge, Te, and Sb are metalloids 11 Mr. Waseem Allabadi Atomic Structure ACT II History of the periodic table: Mendeleev’s periodic table grouped elements by their properties. 1. He noticed that when the elements were arranged in order of increasing atomic mass, they shared, in their chemical properties appeared at regular intervals. 2. Mosely arranged elements according to their atomic number. Electron configuration and the periodic table: Lanthanides: 14 elements with atomic number from (58-71), they have the same chemical and physical properties. Actinides: 14 elements with atomic number from (90-103) they have the same chemical and physical properties. Alkali elements: the elements in the first group in the periodic table that have the following characteristic properties: 1. They have Silvery appearance 2. They are soft 3. They are not found as free elements, they combine vigorously 4. They react strongly with water to form H2 and aqueous solution 5. They stored in kerosene 6. They have low melting points. Alkaline earth elements: the elements in the second group in the periodic table that have the following characteristic properties: 1. They contain pair electrons in last shell. 2. They are harder denser and stronger than alkali 3. They are not found as free elements 4. They are less reactive than alkali 5. They have high melting points. Halogens: the element in the seventh group in periodic table that have the following characteristic properties: 1. They are reactive non metals 2. They are found as diatomic molecules (F2, Cl2, Br2, and I2) 3. The first two elements fluorine and chlorine are gases at room temperature, bromine is liquid at room temperature, but iodine is solid. 4. Astatine (At) element is a radioactive element 5. They react with most metals to form salts. 12 Mr. Waseem Allabadi Atomic Structure ACT II Transition metals: the elements in the periodic table that contain un filled d electron configuration and have the following characteristic properties: 1. They are good conductor 2. They have high luster 3. They are less reactive than alkali and some of them are unreactive Periodic trends: 1- Atomic size (radius) (A) For monoatomic elements or noble gases Atomic size is the distance between nucleus and the outermost shell (B) For diatomic elements (N2, O2, Cl2, F2, Br2, I2, H2) Atomic size is the distance between two nuclei As the number of protons increases the attractive forces between e- and nucleus increases. Size gets smaller. What is the difference between the sizes of Li and Be atoms The atomic size down the group means more shielded e- less effective nucleus charge. The e- can move away easily. Size increases more shells bigger size. 13 Mr. Waseem Allabadi Atomic Structure ACT II Size of ions: Parent atom > its cation Na > Na+ Parent atom < its' anion F < F- Size of anion increases to reduce the repulsive forces between the new e- and original electrons. The shell becomes a little bigger. Size of isoelectronic particles. Which of the following particles. Which of the following has the largest and the smallest sizes? O2-, N3-, Mg2+, Na+, F-,Ne Electrons: 10, 10, 10, 10, 10, 10 Protons: 8, 7, 12, 11, 9, 10 The same number of shells More p+ more attractive force. Smaller size. 14 Mr. Waseem Allabadi Atomic Structure ACT II The ionization energy (I.E) 1st ionization energy: energy needed to remove 1 e- from a gaseous atom to produce positive ion. X (g) + energy X+ (g) + 1 e- Q: Which of the following atom has the highest first ionization energy? (A) K (B) Ca (C) He (D) Ar Second ionization energy: Energy needed to remove e- from a gaseous positive ion to produce a gaseous + 2 ion. X+ (g) + energy X2+ (g) + 1 e- It needs more energy to move it Second ionization energy > first ionization energy X+ is smaller than X closer to the nucleus harder to remove electrons 15 Mr. Waseem Allabadi Atomic Structure ACT II Size effect: it is harder to remove e- from X+ because the attractive force is larger In general: Group 1 is the highest second ionization energy Group2 is the lowest second ionization energy When group elements lose 1 e- their configuration become like the configuration of noble gases. It is harder to remove a second electron. Q: The element that has the highest second ionization energy is: (A) K (B) He (C) Al (D) Ca How to find the group number from the ionization energy of an element? I.E 1 I.E 2 I.E 3 I.E 4 X 700 930 5230 7500 Y 600 720 1150 7310 The highest jump for X is between I.E 2 and I.E 3 so X is from group 2 The highest jump for Y is between I.E 3 and I.E 4 so X is from group 3 Electron affinity: Energy given or taken when e- is added to gaseous atom to produce gaseous – ion 16 Mr. Waseem Allabadi Atomic Structure ACT II X(g) + 1 e- X- (g) + energy Q: the element that has the highest electron affinity is: (A) Mg (B) Al (C) S (D) F Electronegativity: is the ability of element to attract e- of the bond toward itself Q: the element that has the highest electronegativity is : (A) K (B) Al (C) S (D) F As you move in the periodic table from right to left: 1) Size increases 2) Ionization energy decreases 3) Electron affinity decreases 4) Electronegativity decreases As you move in the periodic table from up to down : 1) Size decreases 2) Ionization energy increases 3) Electron affinity increases 4) Electronegativity increases 17 Mr. Waseem Allabadi Atomic Structure ACT II 18 Mr. Waseem Allabadi Atomic Structure ACT II Questions 1. The Fe³⁺ ion contains how many electrons? A. 23 B. 26 C. 29 D. 31 2. How many neutrons are contained in the nucleus of an atom of 238U92? A. 92 B. 119 C. 146 D. 238 3. The atomic symbol 17X⁻ best represents which of the following? A. A chloride ion with 16 electrons B. A chloride ion with 18 electrons C. A chloride ion with 20 electrons D. A rubidium ion with 36 electrons 4. How many electrons does the chlorine-37 ion ³⁷₁₇Cl⁻ , have ? A. 16 B. 17 C. 18 D. 36 5. Which of the following is the same for both Mg²⁺and F⁻? A. Atomic number B. Mass number C. Number of protons D. Number of electrons 6. The atomic number of an element is always the same as which of the following for its neutral atoms? I. The number of electrons outside the nucleus. II. The number of protons in the nucleus. III. The number of neutrons in the nucleus. A. I only B. I and II only C. I and III only D. I, II, and III 7. Upon removal of a valence electron from a neutral atom, the atom becomes A. A radioactive atom B. An electrically neutral atom C. A positive ion D. A negative ion 8. The chemical behavior of an atom is largely determined by the A. Number of neutrons B. Mass number C. Electronic structure D. Atomic mass 19 Mr. Waseem Allabadi Atomic Structure ACT II 9. Which of the following is true for Na⁺, Mg²⁺ and Al³⁺ ? A. They have the same number of neutrons B. They from insoluble hydroxides C. They have the same nuclear charge D. They have the same total number of electrons 10. Which of the following statements is correct concerning the periodic table? A. Atomic size generally decreases from top to bottom in a group. B. Atomic size generally decreases from left to right in a period. C. Elements in the same period have similar properties. D. The first ionization energy decreases from left to right in a period. 11. In the periodic table, nonmetals with similar chemical characteristics are expected to be in the A. Same horizontal row B. Same vertical column C. Lower part of the table D. Upper part of the table 12. Which of the following elements have valence electrons in p orbitals ? I. Chlorine II. Oxygen III. Magnesium A. I only B. II only C. I and II only D. I, II, and III 13. The S²⁻ ion has the same electron configuration as which of the following atoms? A. Cl B. P C. Ar D. Se 14. Which of the following elements has chemical properties that most closely resemble of Mg ? A. Na B. Al C. K D. Ca 15. A positively charged particle with amass of 1 atomic mass unit A. Proton B. Electron C. Positron D. Alpha particle 16. Appositively charged particle with a mass of 4 atomic mass units E. Proton F. Electron G. Positron H. Alpha particle 20 Mr. Waseem Allabadi Atomic Structure ACT II 17. Which of the following statements concerning the quantum number n is correct ? A. It describes the spin of the electron B. It is principal quantum number C. It indicates the spatial orientation of the atomic orbital D. It indicates the total number of electrons in an atom 18. Which of the following is a transition element? A. Iron B. Carbon C. Potassium D. Radium 19. Properties to be expected for the transition metals include which of the following? I. Variable oxidation states II. Ions that from a variety of colored compounds III. Partially filled d orbital's A. I only B. III only C. I and II only D. I, II, and III 20. In the modern periodic table, the elements are arranged in order of increasing A. Mass number B. Atomic mass C. Atomic radius D. Atomic number 21. Electrons in atoms absorb and emit energy only A. In amounts large enough for ionization B. In discrete amounts called quanta C. In the visible region of the spectrum D. When then atoms are in colored molecules 22. Which of the following is the formula for a sulfide ion ? A. S²⁻ B. SO₄²⁻ C. SO₃²⁻ D. S 23. Is a transition element that can exhibit several oxidation states. A. 1s² 2s² 2p⁶ 3s¹ B. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ C. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d² D. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹º 24. Has the lowest first ionization energy (potential). A. 1s² 2s² 2p⁶ 3s¹ B. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ C. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d² D. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹º 21 Mr. Waseem Allabadi Atomic Structure ACT II 25. Which is a halogen? A. 1s² 2s² 2p⁶ 3s¹ B. 1s² 2s² 2p⁶ 3s² 3p¹ C. 1s² 2s² 2p⁶ 3s² 3p⁵ D. 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹º 4s² 4p⁶ 26. Which is a noble gas? A. s² 2s² 2p⁶ 3s¹ B. 1s² 2s² 2p⁶ 3s² 3p¹ C. 1s² 2s² 2p⁶ 3s² 3p⁵ D. 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹º 4s² 4p⁶ 27. Which is a transition metal? A. 1s² 2s² 2p⁶ 3s¹ B. 1s² 2s² 2p⁶ 3s² 3p¹ C. 1s² 2s² 2p⁶ 3s² 3p⁵ D. 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹º 4s² 4p⁶ Questions 28-30 refer to the following substances at room temperature and 1 atmosphere. A. [He]2s²2p⁴ B. [He] 2s²2p⁵ C. [Ne]3s¹ D. [Ar]4s²3d⁶ 28. Has only two unpaired electrons. 29. Has the lowest first ionization energy 30. Has the highest electronegativity Questions 31 -35 refer to the following electron configurations of neutral atoms. A. 1s² 2s² 2p⁵ B. 1s² 2s² 2p⁴ 3s¹ C. 1s² 2s² 2p⁶ 3s² 3p⁶ D. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ 31. An electron configuration that represents a noble gas 32. An electron configuration that represents a transition element 34 An electron configuration that represents an element with four unpaired electrons in the ground state. 35. An electron configuration that represents an excited electron state Questions 36-39 refer to the elements for which the ground-state electron configurations are shown below. A. 1s² 2s¹ B. 1s² 2s² 2p² C. 1s² 2s² 2p⁵ D. 1s² 2s² 2p⁶ 36 The configuration of the element with the largest second ionization energy. 37. The configuration of the element most likely to form molecules containing four covalent bonds. 38. The configuration of the element that exists as single gaseous atoms at 0˚C and 1 atm. 39. The configuration of the element most likely to form diatomic molecules of the form X₂. 22 Mr. Waseem Allabadi Atomic Structure ACT II Questions 40-43 refer to the following subshells. A. 1s B. 3s C. 3p D. 3d 40 Contains up to ten electrons 41 Contains one pair of electrons in the ground-state electron configuration of the lithium atom 42. Is exactly one-half filled in the ground-state electron configuration of the phosphorus atom 43. Contains the valence electrons in the ground-state electron configuration of the magnesium atom Questions 44-45 refer to the following A. N₂ B. Rn C. He D. Fe 44. Is a radioactive noble gas. 45.Commonly exists as irons with charge +2 or +3. Questions46-49 refer to the following. A. Element with atomic number 3 B. Element with atomic number 10 C. Element with atomic number 38 D. Element with atomic number 53 46. Has the lowest second ionization energy? 47. Commonly exists as a +1 ion in its compounds? 48. Has an S orbital that is half-filled? 49. Is the LEAST reactive of the elements represented above? 50. For an atom of cobalt in its ground state, the 3d energy sublevel contains A. 10 electrons B. 9 electrons C. 7 electrons D. 5 electrons 51. Which of the following elements has a valence electron configuration of s²p⁶ ? A. F B. Mg C. Ar D. Fe 52. Which set of orbitals is partially filled in the ground state electron configurations of elements with atomic numbers 21 through 28 A. 2p B. 3s C. 3d D. 4d 53. The electron configuration of sodium is 1s² 2s² 2p⁶ 3s¹, Its characteristic chemical properties result from a change in which part of this configuration? A. 1s² B. 2s² C. 2p⁶ D. 3s¹ 23 Mr. Waseem Allabadi Atomic Structure ACT II 54. Ions of charge 2+ are formed readily from the elements whose atoms have which of the following electronic configuration? I. 1s² 2s¹ II. 1s² 2s² III. 1s² 2s²2p⁶ 3s² A. II only B. III only C. II and III only D. I,II and III 55. Which of the following is the ground-state electron configuration of an alkali metal ? A. 1s² B. 1s² 2s² 2p⁵ C. 1s² 2s² 2p⁶ 3s¹ D. 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹º 4s² 4p² 56. What is the highest-energy orbital occupied by electrons in the ground state of a chlorine atom ? A. 3s B. 2p C. 3p D. 3d 1s²2s²2p⁶3s²3p⁴ 57. Which of the following is true of an atom in the ground state with the electron configuration above ? A. It reacts with sodium to form an ion with a charge of -2. B. It has electrons in sixteen atomic orbitals. C. It has four unpaired electrons in its outer shell. D. It has a low first ionization energy than sodium has. 58. Which of the following configuration is that of an element of the group in the periodic table that contains nitrogen? A. 1s² 2s² 2p⁴ B. 1s² 2s² 2p⁶ 3s² C. 1s² 2s² 2p⁶ 3s² 3p³ D. 1s² 2s² 2p⁶ 3s² 3p⁵ 59. If an atom In the ground state has an outer (valence ) electron configuration of 4s²4p4 , to which family of the periodic table does it belong ? A. The alkali metals B. The alkaline earth metals C. The oxygen family D. The halogens 60. Each of the following ions has ten electrons , which one has the smallest radius ? A. N³⁻ B. O²⁻ C. Na⁺ D. Mg²⁺ 24 Mr. Waseem Allabadi Atomic Structure ACT II 61. The relative sizes of the atomic radii are correctly indicated in which of the following ? I. I > F II. K > Li III. Mg >Ca A. I only B. II only C. I and II only D. I, II, and II 62. Within a family of elements, increasing atomic number generally is accompanied by an increase in which of the following? I. Atomic radius II. Electrongativity III. Ionization energy (potential) A. I only B. I and II only C. I and III only D. I, II, and III 63. Which of the following trends is observed as the atomic number increases in the series of halogens F, Cl, Br, and I? A. The atomic radius increases B. The strength as an oxidizing agent increases. C. The electronegativity increases D. The first ionization energy increases. 64. Which of the following elements has the lowest electronegativity ? A. O B. Al C. Cl D. F 65. Which of the following atoms has the highest first ionization energy ? A. P B. S C. Cl D. Ar 66. Which of the following atoms has the highest first ionization energy ? A. Li B. Na C. K D. Rb 67. The following ions all have the same electron configuration. Which ion is smallest ? A. Se³⁺ B. Ca²⁺ C. K⁺ D. S²⁻ 25 Mr. Waseem Allabadi Atomic Structure ACT II 68. The number of unpaired electrons in a ground state atom of nitrogen is A. 0 B. 1 C. 2 D. 3 69. One isotope of an element contains 13 protons ,13 electrons and 14 neutrons. Another isotope this element is A. ²⁶Mg12 B. ²⁸Al13 C. ²⁷Si14 D. ⁵⁵Fe26 Questions 70-72 Each of the following represents the composition of a particular atom or ion. Electrons Protons Neutrons A. 10 11 12 B. 10 12 11 C. 10 10 11 D. 10 9 9 70. Which is a net charge of -1? 71 Which is an isotope of an element that has a second isotope whose atoms contain 10 electrons, 10 protons, and 12 neutrons? 72. Which is a noble gas atom? Refer to the diagrams below which represent electron transition in the bohr model of the hydrogen atom. The diagrams are roughly to scale 73. The transition to the ground state of the atom 74. The transition that would emit the photon of the greatest energy 75. The transition that requires the greatest input of energy 26 Mr. Waseem Allabadi Atomic Structure ACT II Question 83-85 refer to the following energy diagram. The diagram above is a plot of the energy levels for the electron of the hydrogen atom (roughly to scale) according to the Bohr Theory. The vertical lines represent possible transitions (increases or decreases of energy) that can occurs. 76. The transition from the ground state of hydrogen to first excited state. 77. Of the transition shown the one that involves the LEAST energy 78. The transition that represents the ionization energy (potential) of hydrogen. 27  Mr. Waseem Allabadi Bonding ACT II Bonding The Lewis structure: (The dot structure): (A) For atoms and ions Find the number of valence electrons then represents them as dots around the elements symbol. For molecules (1) Finds the sum of valence e- of all atoms in the molecule. Examples: CH4 = 1(C) + 4(H) = 1(4) + 4(1) = 8e- NH4+ = 1(N)) +4(H)-1= 1(5) +4(1)-1= 8e- CO32- = 1(C) + 3(O) +2 = 24e- CH3OH= 1(C) + 4(H) +1(O) = 1(4) +4(1) +1(6) = 14e- (2) (Select a central atom. (The one that has the lowest electronegativity) (H can't be central) CH4 (Because hydrogen can't be central C is the central atom) NO2 (nitrogen is the central atom) HCN (carbon is the central atom) (3) (Connect the central with the terminal atom by a single bond (- or : ) (4) (Completes the octet of the terminal atoms). (8 around the atom) Exceptions: H 1 e- Be 4 e- B 6 e-- CH4 = 8 e- CCl4= 1(4) + 4(7) = 32 e- (5) put the extra electrons on the central atom 28  Mr. Waseem Allabadi Bonding ACT II Notes: elements in period 3, 4, 5… Can be surrounded by more than 8 e-. H2O 2H+1O = 2(1) +6=8e- It has 2 unshared pair of e- It has 2 shred pair of e- NH3 1(N) + 3(H) = 1(5) +3(1) = 8e- XeF4 8+4(7) = 36e- (6) If the central atom is not octet yet form a double or triple bond with one or more of the terminal atom. CO2 1(C) + 2(O) = 16 e- It will produce a double bond. 29 Mr. Waseem Allabadi Bonding ACT II O2 2(6) = 12 e- It will produce a double bond. N2 2(N) = 2(5) = 10e- It will produce a triple bond C = C > C = C > C-C Strongest weakest Shortest longest CO 4+6=10 e- It will produce triple bond O3 3(O) = 18 e- It has a resonance structures HCN 1+4+5= 10 e- 30 Mr. Waseem Allabadi Bonding ACT II CO32- 4+3(6)+2= 24 e- It has resonance structures also. Lewis structure with unpaired e -: If the sum of valence e- = odd numbers NO2 5+2(6) = 17 e- NO 5+6=11e- Geometry (shaped) of molecules. V: valence S: shell E: electron P: pair R: repulsion 31 Mr. Waseem Allabadi Bonding ACT II Molecules arrange themselves to minimize the repulsive forces between its e- pairs BeCl2 Shape: H2O Shape: BCl3 Shape: NH3 Shape: CH4 Shape: 32 Mr. Waseem Allabadi Bonding ACT II HCN Shape: CCl4 Shape: XeF4 CO32- 33 Mr. Waseem Allabadi Bonding ACT II O3 Note: diatomic molecules have linear structures. O2, HCl, H2, N2 Hydrocarbons: C2H6 C2H4 C2H2 34 Mr. Waseem Allabadi Bonding ACT II Polarity of Bond: 1) Polarity of bonds: The same element (non-polar bond) H-H Different elements (polar bond) H-Cl Polarity of molecules: 1) Lewis structure 2) Study your molecule symmetry Is BeCl2 polar or non- polar? Same particles (same bonds) but opposite directions (non-polar) dipole moment =0 net polarity =0 Is BCl3 polar? Basic structure (with the same terminal atom) Is NH3 polar? Is XeF4 polar? Alcohol polarity? Methanol is polar CH3OH Ethanol is polar CH3CH2OH 35 Mr. Waseem Allabadi Bonding ACT II Bonding (between atoms) 1) Ionic bond 2) Covalent bond 3) Metallic bond The ionic bond: Electrostatic attraction between negative and positive ions. Na:1S22S22P63S1 Na+: 1S22S22P63S23P6 Cl: 1S22S22P63S23P5 Cl-:1S22S22P63S23P6 Na+…..Cl- Atoms loss or gain e- to have the noble gas configuration (stable) How to write formula of an ionic compound? 1) Find charges of elements 2) Cross them without sign Mg2+ Cl1- MgCl2 Note: if the two charges have the same number, don’t cross them. Na1+ Cl1- NaCl Element X reacts with oxygen to form ionic compound X2O. X could be? (A) S (B) Al (C) K (D) Ne 36 Mr. Waseem Allabadi Bonding ACT II Naming of binary ionic compounds: Two types of elements: CaCl2: Leave the name of the first element as it is Add ide to the root of the name Al2O3 Aluminum oxide Ba3N2 Barium nitride Polyatomic ions: Ions with more than 1 atom -1 Name OH1- hydroxide NO31- nitrate CH3COO1- acetate HCO32- bicarbonate MnO41- manganite -2 name: CO32- carbonate SO42- sulfate SO32- sulfite CrO72- dichromate -3 name: PO43- phosphate +1 name: NH41+ ammonium Calcium phosphate: Ca3(PO4)2 1) Write formulas of the following compounds: Potassium sulfate: K2SO4 Aluminum nitrate Al(NO3)3 37 Mr. Waseem Allabadi Bonding ACT II Ammonium phosphate (NH4)3 PO4 2) Which of the following is ionic compound (you can select more than one)? CaI2 SO3 CaSO4 Al(OH)3 Ionic compounds with transition metals: Most of transition metals have more than one charge. Iron: Fe2+ Fe3+ Copper: Cu1+ Cu2+ Cobalt: Co 2+ Co4+ While Ag (silver) has Ag+ only Zn (zinc) has Zn2+ only. To name these compounds you have to add a roman number to the name to show the charge of the metal. CuCl2 Copper (II) chloride CuCl Copper (I) chloride CuO Copper (II) oxide CuSO4 Copper (II) sulphate Fe (NO3)3 Iron (III) nitrate Note: as the difference in electronegativity increases the bond becomes more ionic. (Ionic character) NaF is most ionic than MgF2 MgF2 is most ionic than AlF3 Covalent bond: It is sharing of pair of e- (or more) between 2 nonmetallic atoms. 38 Mr. Waseem Allabadi Bonding ACT II Examples: N-O polar bond N-N nonpolar bond (Could be single double or triple bond) H-O polar bond Naming and formula We used prefix system. Prefix: Number: Mono 1 Di 2 Tri 3 Tetra 4 Penta 5 Hexa 6 Hepta 7 Octa 8 Nona 9 Deca 10 Don’t use mono with the first element Example: CO carbon monoxide CO2 carbon dioxide NCl3 nitrogen trichloride N2O3 dinitrogen trioxide The metallic bond: Attraction between nuclei of a metal and it's free e- ( sea of electrons ) 39 Mr. Waseem Allabadi Bonding ACT II (Among the metal atoms) Properties of ionic compounds: 1- They are solids (at room temperature). 2- They have high melting and boiling points. 3- They are brittle. ( breaks) 4- They can conduct electricity if they are : (A) Molten; (liquids) at high temperature. (B) Aqueous (dissolved in water) (Because they have free ions) Solid ionic crystal ions are not free to move. Properties of covalent compounds: 1- They could be solids (glass SiO2) Liquids (alcohol, H2O, Br2) Gases (CH4, O2, N2, Cl2, NH3) 2- Some of them conduct electricity. Strong acids(aqueous) graphite ( a form of carbon) Note: sugars dissolve in water but they can't conduct electricity. Because they don't produce ions. Sugar (solid) sugar (aqueous) Properties of metals: 1- They are solids (except mercury (Hg) is a liquid. 2- They have high melting point and boiling point 40 Mr. Waseem Allabadi Bonding ACT II 3- They are malleable ( can be shaped into sheets) 4- They are ductile ( can be shaped into wires ) 5- They are shiny ( lustrous ) 6- They can conduct electricity in any form because they have free electrons. Intermolecular forces: 1- Dipole- dipoleforce. Interaction between polar molecules. Dipole: polar like HCl ……… HCl  Hydrogen bond: a special type of dipole-dipole interaction. A strong attraction between hydrogen atom connected to N or O or F on a molecule and Nor O or F on a nearly molecule. Examples: 1- H2O 2- NH3 3- HF Ammonia dissolves in water because they can form H- bonding with each other. H- Bond is stronger than the normal bonds. The effect of H-bonds on the melting point and boiling point surface tension. 1- Increases the boiling point 2- Increases the melting point 3- Increases the surface tension. Surface tension: the ability of a liquid to minimize its surface area. 2- London forces: (Dispersion or vanderwaals) A weak interaction between noble gas atoms or between nonpolar molecules. Ne…….Ne………Ne Ne……He………Ar H-H…..H-H…….H-H The effect of number of e- and the size of atom on London forces. 41 Mr. Waseem Allabadi Bonding ACT II As number of e- or size of atoms increases, London forces become stronger F-F………F-F London Forces Halogens: the size F2 gas Increases Cl2 gas Br2 gas I2 gas The importance of London forces is to help us to convert gases to liquids under the pressure and low temperature. At room temperature and pressure: O=O ………. O=O (Gas) weak London forces O=O ………O=O (liquid) stronger London forces 42 Mr. Waseem Allabadi Bonding ACT II Practice Questions 1. Which of the following ions can form a compound with aluminum of the type Al₂X₃ ? A. Cl⁻ B. I⁻ C. IO₃⁻ D. SO₄²⁻ 2. Which of the following elements can be correctly represented by X in the chemical formula X₂O₃ ? I. Fe II. B III. Ca A. I only B. II only C. I and II D. II and III 3. All of the following are formulas for compounds that are likely to exist EXCEPT A. AlO₃ B. Li₂O C. MnO₂ D. MgO 4.Element X (an alkaline earth element ) unites with element Z (a halogen ) to form a compound that has what formula ? A. XZ B. X₂Z C. XZ₂ D. X₂Z₃ 5. What is the formula for the compound formed by the reaction between element X (atomic number 13) and element Z (atomic number 17) ? A. XZ₂ B. XZ₃ C. X₂Z D. X₂Z₃ 6. An element X that produces compounds with empirical formulas XCI₅ , XH₃ , and X₂O₅ is most likely which of the following ? A. Aluminum B. Carbon C. Sulfur D. Phosphorus 43 Mr. Waseem Allabadi Bonding ACT II Question 7-8 refer to the following substance at room temperature. A. N₂O B. NO₂ C. SiO₂ D. CaO 7. Is an ionic solid 8. Is a gas in which each molecule has an unpaired electron. 9. The bonding in which of the following compounds exhibits the most ionic character ? A. CsCl B. H₂O C. CCl₄ D. SiC 10. Analysis of beryllium iodide shown that it has the formula BeI₂. The product of the reaction of beryllium and fluorine is A. BeF₂ B. BeF₃ C. Be₂F D. Be₃F 11. Which of the following could be used to illustrate the capability of two elements to combine in different ratios ? A. CCl₄ and CF4 B. O₂ and O₃ C. H₂O and H₂S D. FeO and Fe₂O₃ 12. A crystalline substance that is brittle and that conducts electricity when melted or when dissolved in water is likely to be which of the following ? A. An ionic solid B. A covalent network solid C. A metallic solid D. A polymer 13. which of the following substances contains both ionic and covalent bonds ? A. CaBr₂ B. SO₂ C. NH₄Br D. H₂O 44 Mr. Waseem Allabadi Bonding ACT II 14. Types of chemical bonding present in solid Na₂CO₃ include which of the following ? I. Covalent II. Metallic III. Ionic A. I only B. II only C. III only D. I and II E. I and III 15. Which of the following compounds contains both ionic and covalent bonds ? A. NO₂ B. Na₂CO3 C. CaCI₂ D. H₂O₂ 16. The order of electronegative of certain elements is as follows : F >Cl> Br >> Li > Na >K > Cs. in which of the following substances do the bonds exhibit the greatest ionic character ? A. LiF B. NaBr C. CsF D. NaF Questions 17 -19 refer to the following types of bondsBetween electron orbitals. A. s ___ s B. s ___ p C. p ___ p D. sp³ ___ sp³ 17. Bond in H₂ 18. Bond in HCI 19. Bond between carbon atoms in C₂H₆ 20. Which of the following bonds is the strongest ? A. C-H B. C-C C. C=C D. C=N 21. The Lewis electron-dot structure N₂ is 45 Mr. Waseem Allabadi Bonding ACT II 22. Compounds that have the empirical formula CH include which of the following ? A. I only B. III only C. II and III only D. I , II and III Question 23-25 refer to the following molecules. A. N₂ B. C₂H₆ C. NH₃ D. CO₂ 23. A molecule that has atoms joined by a triple covalent bond. 24. A molecule that is polar. 25. A molecule having a 180º bond angle. 26. Which of the following best describes the type of bonding between iodine atoms a molecule of I₂? A. Ionic bonding B. Metallic bonding C. Hydrogen bonding D. Dispersion ( London ) force interactions 27. Which of the following is the best Lewis structure for carbon dioxide , CO₂ ? (A) (B) (C) (D) (E) K 46 Mr. Waseem Allabadi Bonding ACT II 28. Molecules of substances with low boiling points are usually held together in liquid and solid states by A. London dispersion forces B. Ionic bonding C. Network bonding D. Covalent bonds to each of the nearest neighbors. 29. Which of the following compounds of X and Y , with the structure indicated below , can be polar ? Questions 30 -35 refer to the following bond types. A. Metallic bond B. Ionic bond C. Polar covalent bond D. Nonpolar covalent bond 30. The type of bond that describes the attraction between oppositely charged atoms or groups of atoms in a crystalline solid. 31. The type of bond among the atoms in solid osmium in a crystalline solid. 32. The type of bond within a molecules of phosphorus , P₄. 33. The type of bond within a CCl₄molecule. 34. The total number of electrons used to write the Lewis structure of H₂CO is A. 6 B. 8 C. 12 D. 16 47 Mr. Waseem Allabadi Bonding ACT II Questions 35 -37 refer to the following bond types. A. Single covalent bond B. Double bond C. Triple bond D. Ionic bond in solid 35. Found in N₂ 36. Found in F₂ 37. Present between carbon atoms in C₂H₄- 38. All of the following molecules are linear EXCEPT A. H₂O B. C₂H₂ C. HCN D. HCl 39. In which of the following molecules is the bonding nonpolar ? A. CCl₄ B. Cl₂ C. NH₃ D. HF 40. Which of the following is NOT a valid electron dot formula for a neutral molecule ? 41. What is the molecular geometry of SiCl₄ ? A. Linear B. Trigonal planar C. Square planer D. Tetrahedral 42. Which of the following conversions requires the breaking of covalent bonds ? A. KF(s) → K⁺ + F⁻ B. Ne(g) → Ne⁺(g) + e⁻ C. MgCl₂(s) → MgCl₂(l) D. 2 CO₂(g) → 2 CO(g) + O₂(g) 48 Mr. Waseem Allabadi Bonding ACT II 43. Which of the following molecules is polar ? A. Cl₂ B. C₂H₆ C. NH₃ D. CO₂ E. CCl₄ Questions 44 -45 refer to the following A. Covalent bonding B. Ionic bonding C. Metallic bonding D. Dispersion (London ) forces 44. makes possible the liquefaction of nitrogen gas at a very low temperature 45. Involves the sharing of one or more pairs of electrons between two atoms. 46. The geometric shape of the CCl₄ molecule is A. Square planer B. Linear C. Tetrahedral D. Triangular Questions 47 - 49 refer to the following compounds. A. NH3 B. H₂S C. CH₄ D. CO E. CO₂ 47. Contains a triple bond 48. Contains two double bonds 49. Has one and only one unshared pair of electrons. 50. Of the following , which best describe the shape of a water molecule ? A. Bent B. Tetrahedral C. Linear D. Cyclic 51. All of the following molecules are linear EXCEPT A. H₂O B. C₂H₂ C. HCN D. HCl 52. Which of the following molecules is nonpolar ( has no permanent dipole moment )? A. HCl B. ClF C. CH₂Cl₂ D. CCl₄ 49 Mr. Waseem Allabadi Bonding ACT II 53. the primary intermolecular attraction that makes it possible to liquefy hydrogen gas is called A. London dispersion forces B. Dipole-dipole attraction C. Ionic bonding D. Hydrogen bonding 54. Which of the following is the correct and complete Lewis electron dot diagram for PF₃ ? 55. Which of the following bonds has the shortest length A. Br-Br B. C-C C. C=C D. Si-Si Refer to the following substances at 12 c and 1 atm A. NH 3 B. CH 4 C. LiF D. SO3 E. CS2 56. Which exists as ionic crystals? 57. Which is composed of molecules that are linear and have one or more double bonds? 50 Mr Waseem Allabdi Solid and Liquid ACT II Solid and liquid Solid elements Most elements are solids (at room temperature) Liquid elements Mercury Hg Bromine Br2 Gaseous elements H2, N2, O2, F2, Cl2, and noble gases. Types of bonding in liquids: 1-dipole- dipole 2-hydrogen bonding 3-London forces Types of solids: (A) Crystalline solids (have regular shapes(lattice)) 1) Ionic compounds (high melting point, high boiling point) 2) Molecular solids (covalent)(low melting point ,low boiling point) Examples: S8 (sulfur), P4 (phosphorus) 3) Metallic solids (high melting point and boiling point) 4) Network structure solid (giant structure) like diamond, silica (SiO2) very high melting point and boiling point Note: 1) diamond, graphite, coal, and charcoal are allotropes 2) oxygen and ozone are allotropes also. Allotropes: same element with different structures. (B) Amorphous (no regular shape) Glass, plastic Compounds are three types: 1) Ionic (solids) 2) Covalent (solid SiO2, liquid H2O,alcohol, CCl4,C8H18,Br2 and gas NH3, CH4, CO2, CO 3) Metallic (metals) solids except Hg 51 Mr Waseem Allabdi Solid and Liquid ACT II Phase change: Sublimation Melting Evaporation Solid liquid gas Freezing Condensation Deposition Examples for sublimation: (1) Dry ice (CO2(s)) in atmosphere CO2(s) CO2(g) (2) Iodine crystal I2(s) I2 (g) Purple vapor Heating curve: Heating curve for water: (1) Solid (2) Solid and liquid (3) Liquid (4) Liquid and gas (5) Gas 0 C0 (melting point) 1000 (boiling point) 52 Mr Waseem Allabdi Solid and Liquid ACT II Note: 1) during melting and evaporation the temperature does not increase because the heat is consumed to separate particle. 2) More heat is needed to convert liquid to gas than converting solid to liquid because particles in the gas are separated totally. The boiling point: The temperature at which the vapor pressure of a liquid equals the atmospheric pressure. Note: at 1 atm is called the normal boiling point. (at sea level) The phase diagram (water) HH (A) Solid only (G) Liquid only (H)Gas only (B) Triple point (E) Solid and liquid (C)Solid and gas (F)Liquid and gas (D)Critical point Note: density of ice is lower than density of liquid why? Because ice has open structure (tetrahedral) 53 Mr Waseem Allabdi Solid and Liquid ACT II Solution: contains solvent, and solute. 1) Solvent: the substance that has the most percentage in the solution. 2) Solute: the substance that has the least percentage in the solution. Solubility: The maximum amount of the solute that can be dissolved in 100 mL of a solvent at 25 C0. Types of solutions: (A) Unsaturated (B) Saturated (C) Super saturated Example: Solubility of a substance X is 80 g/ 100 mL H2O at 25 C0 Case1: If the mass of solute < its solubility, unsaturated. Case 2: If the solubility of solute = its solubility, saturated. Case 3: If the solubility of solute > than its solubility, super saturated. Factors that affect the solubility: 1- Temperature: In general, as temperature increases, solubility increases. (with some exceptions). 54 Mr Waseem Allabdi Solid and Liquid ACT II 2- Surface area: As surface area increases solubility increases. Solubility of gases: Solubility of gases increases as temperature decreases or as pressure increases. Like dissolves likes  Polar solvent dissolves polar or ionic substances.  Non- polar solvent dissolves non-polar substances. 55 Mr Waseem Allabdi Solid and Liquid ACT II How can water dissolve ionic compounds? Water molecules surround ions and separate them. Electrolytes: substances that can conduct electricity Strong electrolytes: completely conductors of electricity (A) Metals (free electrons) (B) Aqueous or molten ionic compounds ( free ions) (C) Strong acids HCl, HNO3,…… (free ions) (D) Graphite ( a form of carbon) (free e-) Example: NaCl(s) Na+ (aq) + Cl- (aq) Weak electrolytes: partially conductors of electricity (A) Slightly soluble ionic compounds in water. AgCl(s) Ag+ (aq) + Cl-(aq) (B) Weak acids. Water is not good electrolyte. Non- electrolyte: can't conduct electricity. Example: glass, wood, alcohol, aqueous sugar. Sugar (s) sugar (aq) no ions 56 Mr Waseem Allabdi Solid and Liquid ACT II Calculations: Moles= concentration×volume (solution) Mol= M × L Q: 25 g of NaCl (molar mass =58.5 g/mol) is dissolved in 248 mL H2O to prepare 250 mL solution. The concentration of NaCl is: (A)25/58.5×248/1000 (B) 25/58.5÷248/1000 (C) 25/585×250/1000 (D) 25/58.5÷250/1000 Q: find the moles of CaCl2 (molar mass = 111 g/mol) needed to prepare 300 mL, 0.1 M CaCl2 (A) 3 (B) 0.333 (C) 0.3 (D) 0.03 (E)0.003 Q: find the concentration of MgCl2 solution if 19 g of MgCl2 (molar mass =95 g/mol) is dissolved in water to prepare 2 L solution. (A) 0.1 M (B) 0.01M (C) 0.2M (D) 0.02M (E)0.002M How to find the concentration or moles of an ion in a compound? Example: 0.1M CaCl2 Ca 2+ + 2Cl- 0.1M 0.2M Q: In 0.2M Ca(NO3)2 solution, which of the following concentrations is correct: (A) [Ca2+] = 0.1M (B) [Ca2+] = 0.4M (C) [NO31-] = 0.2M (D) [NO31-] = 0.4M Q: In 0.3 M (NH4)3PO4. The correct concentration is: (A) 0.1 M PO43- (B) 0.3 M PO43- (C) 0.3 M NH4+ (D) 0.6MNH4 Q: if 100 mL, 0.1 M NaCl is mixed with 100 mL 0.1 M CaCl2 and water is added to prepare 1000 mL solution. What is the concentration of Cl- in the final solution? (A) 0.1 M (B) 0.02 M (C) 0.03M (D) 0.3 M Dilution : Adding a solvent to a concentrated solution to decreases its concentration. Moles number will not change. Example: If water is added to 100 mL, 0.1 M NaCl solution to prepare 500 mL solution. Find the final concentration of Cl- in this solution: 57 Mr Waseem Allabdi Solid and Liquid ACT II (A) 0.1 M (B) 0.01 M (C) 0.5 M (D) 0.02 M (E) 0.002 M If water is added to 100 mL, 0.1 M CaCl2 solution to prepare 500 mL solution. The final concentration of Cl- in this solution is: (A) 0.1 M (B) 0.01 M (C) 0.04 M (D) 0.02 M (E) 0.004 M Partially ionized substances: Q: what is the concentration of HF molecules in 0.01 M HF that is 20% ionized? (A) 0.002 (B) 0.008 (C) 0.08 (D) 0.0008 (E) 0.0002 In the previous question what is the concentration of F- or H+ ? Preparation of solutions: How to prepare a solution of exact concentration. 1- Find moles of solute mole= concentration×volume 2- Find mass mass= mole×molar mass 3- Dissolve this mass in a little amount of solvent (completely) 4- Transfer the solution to a volumetric flask and complete to the wanted volume by adding the solvent. How to prepare 1000 mL 0.1 M solution NaCl 1- Mole= concentration×volume Mole= 0.1 × (1000/1000) 2- Mass = moles × molar mass Mass= 0.1×58.5 = 5.85 g 3- Dissolve it completely 4- Transfer it to a volumetric flask Colligative properties: Properties that depend on the number of particles not on their types. - Boiling point elevation - Freezing point depression (increases) - Vapor pressure depression (decreases) - Osmotic pressure increases If a substance added is insoluble it has no effect. 58 Mr Waseem Allabdi Solid and Liquid ACT II Examples: 1- Adding NaCl to the water will increase the boiling point. 2- Adding anti-freeze to the car engine to decreases the freezing point or increases the boiling point of water. 3- Adding salt to ice to melt it. 4- Osmotic pressure. 59 Mr Waseem Allabdi Solid and Liquid ACT II Practice Questions Question 1-3 refer to the figure below, which represents the phase diagram of pure water. For each of the following questions, select the letter corresponding to the appropriate point on the phase diagram. 1. The triple point 2. The critical point 3. The normal boiling point Question 4 - 6 refer to the phase diagram for pure water shown below. 4. A pressure-temperature condition at which water is stable only as a gas. 5. A pressure-temperature condition of water at which there is a dynamic equilibrium between the liquid an vapor phases. 6. A pressure-temperature condition through which water may pass in a transition from the solid phase directly to the vapor phase. Questions 7-9 refer to the phase diagram for pure H₂O , shown below. Each arrow indicates a transition between different sets of conditions. 7. Melting occurs with which transition ? 8. The phase remains the same with which transition ? 9. Condensation occurs with which transition ? 60 Mr Waseem Allabdi Solid and Liquid ACT II Questions 10-12 refer to the following phase diagram for a pure substance. 10. Where are the vapor pressure of the solid and the pressure of the gas equal ? 11. Where does the solid exist in equilibrium with the liquid ? 12. Where does the substance exist only as a gas ? Questions 13-15 refer to the heating curve below for a sample of a pure substance at one atmosphere , energy is added at a constant rate as the substance goes from the solid to the gas phase. 13. What is the melting point of the pure substance ? A. Less than 180 K B. 180 K C. 300 K D. 360 K 14. Over what temperature range is the substance in the liquid phase ? A. 180 to 360 K B. 180 to 860 K C. 360 to 860 K D. 360 to 980 K 15. How much heat is required to vaporize the sample once the liquid has reached its boiling point ? (assume that heat is added at a constant rate of 180 kJ/min.) A. 270 kJ B. 360 kJ C. 540 kJ D. 720 kJ 61 Mr Waseem Allabdi Solid and Liquid ACT II 16. A pure liquid is at temperature above its freezing point. heat energy is removed at a constant rate to bring it to a temperature below the freezing point, if temperature readings are made every 30 seconds, a graph of the measurements would look most like which of the following ? ̊ ? 17. According to the graph above , which of the liquids has a normal boiling point closest to 50 C A. Acetone B. Chloroform C. Ethyl alcohol D. Benzene 18. Which of the following groups contains substances that are all good conductors of electric current in the solid state ? A. Na₂O, C (diamond), Na B. Na, Cu, Au C. Na, Na₂O , Na₂CO₃ D. C (diamond) , I₂, S₂ 19. Which of the following is NOT a good conductor of electricity? A. 0.1 M strontium nitrate B. 0.1 M hydrochloride acid C. Molten potassium chloride D. Molten sulfur E. Solid chromium metal 62 Mr Waseem Allabdi Solid and Liquid ACT II Questions 20-22 refer to the following elements. A. Chlorine B. Sulfur C. Bromine D. Carbon 20. Is a nonmetal that is a liquid at 25˚C and 1 atm. 21. Is a good electrical conductor in one of its solid forms at room temperature and a atm. 22. Ground-state atoms of this element contain electrons is d orbital's. 23. A 1 M solution of sodium chloride conducts an electric current better than 1 M solution of sugar does, because A. Sugar contains carbon , but sodium chloride does not B. Sodium chloride is a network solid, but sugar is not C. Sugar molecules have water molecules tightly bound to them , but sodium chloride molecules do not D. In solution sodium chloride is dissociated into ions, but sugar is not dissociated 24. Characteristics associated with solid metals include which of the following ? I. High thermal conductivity II. High electrical conductivity III. Lustrous appearance A. II only B. I and II only C. I and III only D. I, II, and III Refer to the following A. Metallic solids B. Ionic solids C. Covalent network solids D. Amorphous solids 25. Have no regular arrangements of particles 26. Are good conductors of heat and electricity 27. Generally do not conduct electricity until melted 28. which of the following solids can be expects decompose most readily on heating ? A. CaCO₃ B. KNO₃ C. SiO₂ D. NaCl I₂(s) → I₂ (g) 29. The process that occurs in the change shown above is usually called A. Condensation B. Sublimation C. Boiling D. Melting 63 Mr Waseem Allabdi Solid and Liquid ACT II 30. The energy required to vaporize one mole of a pure solid is defined as the heat of A. Combustion B. Fusion C. Formation D. Sublimation 31. Which of the following statements about a liquid that evaporates readily at room temperature is correct? A. It has a high vapor pressure B. It is considered to be nonvolatile C. It would make its container feel warm to the touch D. It should be stored in an open container 32. Which of the following is a list of three substances that are all liquids at 25˚C and 1 atm ? A. Hg, Br₂ , H₂O B. Hg, Na , Fe C. Fe₂O₃ , CO₂ D. Cl₂ , HCl , H₂O 34. Of the following, the metal that reacts most readily with water at room temperature is A. Rb B. Fe C. Ag D. Pb 35. which of the following compounds is expected to have the highest melting point ? A. NaF B. NaCl C. NaBr D. Nal 36. All of the following statements describing water are correct EXCEPT: A. It can exhibit hydrogen bonding B. It is more dense at 4˚C than at 0˚C C. It is a polar substance. D. It is a strong electrolyte 37. Correct statements about a solution of sodium chloride include which of the following ? I. it has a normal boiling point identical with that of pure water II. It has a greater density than pure water at the same temperature III. It has a lower freezing point than pure water A. I only B. III only C. II and III only D. I, II, and III 38. When equal masses of each of the following substances absorb the same quantity of energy , which substance of any , will show the greatest temperature change ? A. Ethanol , specific heat = 2.46 J/(g˚C) B. Carbon (graphite), specific heat = 0.720 J/(g˚C) C. Carbon ( diamond) , specific heat = 0.502 J/(g˚C) D. Mercury specific heat = 0.139 J/(g˚C) 64 Mr Waseem Allabdi Solid and Liquid ACT II 39. A baker contains some crystals of NaCl in equilibrium with a saturated solution of NaCl at constant temperature. If additional crystals of NaCl are added, which of the following occurs ? A. The crystals dissolve completely. B. Heat is evolved C. The density of the solution increases. D. No change occurs in the concentration of the solution. 40.when NaCl is dissolved in pure water at 25˚C and 1 atm , the resulting solution , compared to pure water , has a A. Lower density B. Higher freezing point C. Higher electrical conductivity D. Higher vapor pressure 41. When crystals of solid KClO₄ are added to an unsaturated solution of KClO₄ they dissolve. All resulting solution EXCEPT A. A saturated solution has been produced B. The boiling point increases C. The freezing point decreases D. The pH of the solution remains unchanged 42. Which of the following substances is a liquid at room temperature ? A. F₂ B. Cl₂ C. Br₂ D. H₂ 43. The graph above shows the solubility curves of four crystalline solids ( W, X, Y, and Z ) in water over the temperature range from 10˚C to 100˚C. solubilities are in a grams per 100 milliliters of water. These data shown that A. The last soluble of the four solid is W B. The solid that is most soluble at 15˚C is X C. The saturated solution that changes least in concentration on cooling from 100˚C to 10˚C is that of Z D. Solid W is infinitely soluble at 100˚C 44. In a series of experiments, hydrogen is added to a constant amount ( 1mole ) of oxygen and allowed to react to form water. If the number of moles of water formed is plotted against the number of moles of hydrogen added, which of the following graph results ? 65 Mr Waseem Allabdi Solid and Liquid ACT II 45. the graph above was developed in a group of experiments in which varying amounts of reactant X were added to a fixed amount of reactant Y. the mass of product formed in each experiment was the measured. The statements concerning this experiment include which of the following ? I. When 1 g of X is added , all of X reacts. II. When 5 g of X is added , X is present in excess. III. When 2 g of X is added , both X and Y react completely. A. I only B. I and III only C. II and III only D. I, I, and III 46. Which of the following statements is the concerning a saturated solution of a salt at a constant temperature? A. The concentrations of salt and solvent are usually equal. B. The amount of dissolved salt is constant C. Addition of solid salt shifts the equilibrium , which results in an increase in the amount of dissolved salt D. At the same temperature , a saturated solution of any other salt has the same concentration. 47. The boiling points of NH₃, H₂O and HF are all higher than would be expected on the basis of their molecular masses because of the A. Low kinetic energy of the molecules B. high ionization energies of N, O, and F C. high potential energy of the molecules D. hydrogen bonding between the molecules 48. If 100 grams of sucrose raises the boiling point of 250 grams of water by 0.60˚C , then 200 grams of sucrose will raise the boiling point of 500 grams of water by A. 0.15˚C B. 0.30˚C C. 0.60˚C D. 2.40˚C 49. The density of a5.00 cm by 2.00 cm by 10.0 cm rectangular solid with a mass of 156 g is most correctly reported as A. 0.64 g/cm³ B. 0.641 g/cm³ C. 1.56 g/cm³ D. 1.560 g/cm³ 66 Mr Waseem Allabdi Solid and Liquid ACT II 50. The mass and volume measurements of a sample of aluminum at 25˚C are 75 g and 28.0 mL , respectively. Which of the following is the correct value for the density of aluminum and uses the appropriate number of significant digits ? A. 0.37 g/ml B. 0.370 g/ml C. 2.12 g/ml D. 2.70 g/ml 51. Figure I above shown a 10 mL graduated cylinder partially filled with water , Figure II above shown the same graduated cylinder after the addition of a piece of metal with mass of 60.0 grams. The density of the metal is approximately A. 0.33 g/mol B. 2 g/mol C. 1.5 g/mol D. 3 g/mol 67 Mr Waseem Allabdi Solid and Liquid ACT II Molarity 1. What is the number of moles of NaOH (molar mass 44. g/mol) present in 250 mL of 0.10 M NaOH(aq) ? A. 0.025 mol B. 0.10 mol C. 1.0mol D. 4.0 mol E. 25 mol 2. How many grams of RbNO₃ (molar mass 147 g/mol) are required to produce 0.500 L of a 0.200 M RbNO₃ solutions? A. 73.5 g B. 29.6 g C. 14.7 g D. 2.96 g E. 1.47 g 3. A 4 g sample of sodium hydroxide (molar mass 40 g/mol) is dissolved in sufficient water to make 200 mL of solution. The molarity of the solution is A. 0.002 M B. 0.01M C. 0.5 M D. 0.8 M E. 20 M 4. What number of moles of AgNO₃ is contained in 50. mL of a 0.50 M solution ? A. 0.025 mol B. 0.050 mol C. 0.10 mol D. 0.25 mol E. 0.50 mol 5. A 21 g sample of NaF(s) (molar mass 42 g/mol ) is dissolved in enough water to yield 2.0 L of solution. what is the molar concentration of Na⁺(aq) ? A. 0.010 M B. 0.050 M C. 0.10 M D. 0.25 M E. 0.50 M 6. What is the molarity of CaBr₂ (molar mass 200 grams) in a solution prepared by dissolving 12 grams of CaBr₂ in sufficient water to make 0.50 liter of solution ? A. 0.020 M B. 0.10 M C. 0.12 M D. 0.20 M E. 0.24 M 68 Mr Waseem Allabdi Solid and Liquid ACT II 7. A 0.050 M solution contains 29 g of NaCl (molar mass 58 g/mol) , What is the volume of the solution ? A. 0.50 L B. 1.0 L C. 2.5 L D. 5.0 L E. 10. L 8. A student wishes to prepare a 0.1000 M solution of vitamin C, C₆H₈O₆ (molar mass 176.14 g/mol). the student should dissolve 17.614 g of vitamin C in A. 982.4 mL of water B. 1000 mL of water C. 1000 g of water D. Enough water to make 1.000 mL of solution E. Enough water to make 1.000 g of solution 2 Ag⁺(aq) + SO₄²⁻(aq) → Ag₂SO₄ (s) 9. A mixture of AgNO₃(aq) and Na₂SO₄(aq) is prepared so that at the instant of mixing before any reaction occurs.[Ag⁺] = [SO₄²⁻] = 0.6 M if the reaction represented by the equation above goes to completion , what concentration of SO₄²⁻(aq) remains unreacted ? A. 0. 1M B. 0.2 M C. 0.3 M D. 0.4 M E. 0.5 M BaCl₂(aq) + 2 AgNO₃(aq) → 2 AgCl(s) + Ba(NO₃)₂(aq) 10. What volume of 0.5 M AgNO₃(aq) solution is needed to react completely with 0.2 mol of BaCl₂(aq) according to the equation above? A. 100 mL B. 200 mL C. 400 mL D. 600 mL E. 800 mL 11. When 0.050 mol of NaCI(s) and 0.050 mol of CaCl₂(s) are mixed with enough water to make 1.000 mL of solution , what is the concentration of Cl⁻(aq) in the solution ? A. 0.10 M B. 0.15 M C. 0.20 M D. 0.25 M E. 0.30 M 12. When 1.0 L of 1.0 M NaOH is diluted to 2.0 L with distilled water , the Na⁺ concentration in the final solution is A. 4.0 M B. 2.0 M C. 1.0 M D. 0.50 M E. 0.20 M 69 Mr Waseem Allabdi Solid and Liquid ACT II 13. What is the concentration of HF molecules in a 0.01 M HF solution that is 20 percent ionized ? A. 0.002 M B. 0.008 M C. 0.01M D. 0.08 M E. 0.2 M 14. Which of the following concentrations of ions is correct for a 0.2 M solution of Ca(NO₃)₂ ? A. 0.1 M Ca²⁺ B. 0.4 M Ca²⁺ C. 0.1 NO₃⁻ D. 0.2NO₃⁻ E. 0.4NO₃⁻ 15. A student mixed 50 mL. of 2.0 M HCI with 50 mL of 2.0 M NaOH₂ both initially at 20.0˚C , and observed an increase of t Celsius degrees in the student were to repeat this procedure using 25 mL each of the same solutions , the approximate increase in the temperature for this trail would be how many Celsius degrees ? A. ¼ t B. ½ t C. 1 t D. 2 t E. 3 t 16. If 100. mL of 0.50 M NaOH(aq) exactly neutralizes 50. mL of a solution of HCl(aq) , the molarity of the HCl(aq) is A. 0.10 M B. 0.25 M C. 0.50 M D. 1.0 M E. 2.0 M 17. In standardizing a solution of sodium hydroxide, a student found that 30. mL of the solution exactly neutralized 20. mL of 0.3 M hydrochloric acid. The molarity of the sodium hydroxide solution was A. 0. 1M B. 0.2 M C. 0.3 M D. 0.4 M E. 0.5 M 18. When 50.0 mL of 5.00 M NaOH is mixed with 150. mL of 2.00 M NaOH, 200. mL of solution is produced. The molarity of the resulting solution is A. 1.75 M B. 2.75 M C. 3.50 M D. 7.00 M E. 11.0 M 70 Mr Waseem Allabdi Solid and Liquid ACT II 19. if 30.0 g of NaOH(s) (molar mass 40.0 g/mol) added to 70.0 mL of H₂O(l) so that the volume of the solution is 75.3 mL what is the molarity of the NaOH ? A.. 30.0. M 40.0 X 0.0700 B.. 30.0. M 40.0 X 0.0753 C.. 30.0. M 40.0 X 753 D. 30.0 X 40.0 M 70.0 E. 30.0 X 40.0 M 0.0753 20. In order to prepare 2.00 L of a 1.00 M aqueous solution of K₂SO₄ (molar mass 174 g/mol), what mass of K₂SO₄ is needed ? A. 696 g B. 348 g C. 174 g D. 87.0 g E. 43.5 g 21. How many moles of Na₂CO₃ (molar mass 106 g/mol) are contained in 20.0 mL of 1.50 M Na₂CO₃ ? A. 0.0150 mol B. 0.0300 mol C. 0.300 mol D. 0.500 mol E. 1.50 mol 22. Which of the following amounts of 10. M HCl and water combine to form 1.0 M HCl ? ( assume volumes are additive ) 10. M HCl WATER A. 1.0 L 10.0 L B. 1.0 L 9.0 L C. 1.0 L 1.0 L D. 1.0 L 10.0 mol E. 0.10 kg 1.0 kg 23. a solution is prepared by adding 100. mL of 1.00 M HCl to 100. mL of 2.00 M NaCl and diluting to a total final volume of 1.00 L. the concentration of Cl⁻(aq) in the final solution is A. 0.100 M B. 0.200 M C. 0.300 M D. 2.00 M E. 3.00M 24. What volume of 15.0 M H₃PO₄ is needed to prepare 1.00 L of a 3.00 M H₃PO₄ solution ? A. 20.0 mL B. 45.0 mL C. 200. mL D. 450. mL E. 500. mL 71 Mr Waseem Allabdi Solid and Liquid ACT II 25. Which of the following is observed when a small piece of solid zinc is added to an open beaker containing 100 mL of 1 M HCl(aq) ? A. A white precipitate forms and the temperature rises. B. A green precipitate forms and a gas is evolved. C. The zinc disappears and a gas is evolved. D. The solution turns green but no new solid forms. E. Only an increase in the temperature of the solution is observed. 26. What is the mass of Ba(NO₃)₂(s) (molar mass 261 g/mol) required to prepare 500 mL of 0.100 M Ba(NO₃)₂(aq) ?? A. 13.1 g B. 26.1 g C. 52.2 g D. 131 g E. 261 g 27. When 100. mL of 0.10 M Kl solution is mixed with 100. mL of 0.10 M Pb(NO₃)₂ solution , what number of moles of solid PbI₂ is found in the precipitate ? A. 0.040 mol B. 0.030 mol C. 0.020 mol D. 0.010 mol E. 0.0050 mol 28. in an acid-base titration, 30.0 mL of 0.10 M KOH(aq) was required to titrate 60.0 mL of a solution containing an unknown monoprotic acid. What is the molarity of the acid ? A. 0.018 M B. 0.050 M C. 0.18 M D. 0.20 M E. 0.50 M 29. What is the molar concentration of Cl⁻(aq) in 0.12 M MgCl2(aq) ? A. 0.020 M B. 0.040 M C. 0.060 M D. 0.12 M E. 0.24 M 2 SO2 (g) + O2(g) ⇄ 2 SO3(g) 30. If all measurements are made under the same conditions, what volume of SO3 (g) is produced when 40 mL of SO2 (g) and 20 mL of O2 (g) react completely according to the equation above ? A. 10 mL B. 20 mL C. 40 mL D. 60 mL E. 80 mL 31. If 100. mL of each of the following aqueous solutions is mixed with 100. mL samples of 1.0 M NaOH(aq), the greatest amount of heat is liberated by A. 0.010 M H2SO4 B. 0.010 M HCl C. 0.10 M HCl D. 0.10 M H2SO4 E. 1.0 M HCl 72 Mr Waseem Allabdi Solid and Liquid ACT II 32. A solution is prepared by adding 100 milliliters of 1.00 molar HCl to 100 milliliters of 2.00 molar NaCl and diluting to a total final volume of 1.00 liter. The concentration of Cl⁻ in the final solution is A. 0.100 M B. 0.200 M C. 0.300 M D. 2.00 M E. 3.00M 33. When 50. mL of 1.5 M NaCl(aq) is diluted with pure water to a final volume of 150. mL , what is the molarity of the resulting solution ? A. 0.10 M B. 0.50 M C. 1.5 M D. 4.5 M E. 5.0 M 34. The molarity of solution X is to be determined by titration procedure. To carry out this procedure , all of the following must be known EXCEPT the A. Equation for the chemical reaction that occurs during the titration B. Volume of solution X that is used. C. Mass of solution X that is used. D. Volume of the solution that reacts with X E. Molarity of the solution that reacts with X 35. How much water should be added to a 50 ml sample of 1.0 M CaCl2(aq) prepare a 0.40 M solution of CaCl2 (aq) ? (Assume volumes are additive.) A. 12 mL B. 20. mL C. 50 ml D. 75 ml E. 125 ml 36. what mass of NaOH is required to make 250 ml of a 0.010 M solution of NaOH (molar mass 40g/mol) A. 0.10g B. 0.20g C. 0.30g D. 0.40g E. 0.50g 73 Mr. Waseem Allabadi Gas ACT II Gas Gases: Clean air contains 79% N2, 20% O2, and 1% other gases: noble gases, carbon dioxide, and water vapor. Atmospheric pressure: Is the pressure produced by the air on the earth. What is the effect of altitude on the atmospheric pressure? Space Atmosphere p atm < 1 atm Above sea level P atm Sea level P atm = 1 atm Below Sea level P atm > 1 atm As altitude increases atmospheric pressure decreases. 74 Mr. Waseem Allabadi Gas ACT II Units of pressure: Pressure = force (N/m2) pascal Pa Area Pascal is the SI unit 1- Atm : atmospheric 2- mm Hg: millimeter mercury. 3- torr: pressure. 1 atm = 760 mm Hg 1 atm= 760 torr 1 torr= 1 mmHg  How to measure the atmospheric pressure in the lab: The manometer: It is used to measure the pressure of an enclosed gas. Open end manometer: 1) P gas > P atm P gas = p atm + h 2) P gas < P atm 3) P gas = P atm P gas = P atm 75 Mr. Waseem Allabadi Gas ACT II All the three cases: Closed end manometer: P gas = h 76 Mr. Waseem Allabadi Gas ACT II The gas law: 1) Boyle's law: P and V ( T is constant) As the pressure increases the volume decreases. P1V1=P2V2 Q: 2 L of gas under 8 atm is allowed to expand to 4 L the new pressure is : (A) 1 atm (B) 2 atm (C) 3 atm (D) 4 atm 2) Charle's law: T and V ( P is constant) As T increases V increase. As the temperature increases, the speed of particles increases, and the kinetic energy increases Stronger collision between particles and walls of the container. Temperature should be in kelvin not C0 K= C0 + 273 T1 =T2 V1 V2 Q: 1 L of a gas at 25C0 was heated to 40 C0. The new volume is: (A) Less than 1 L (C) more than 2 L (D) between 3 and 4 (E) between 1 and 2 L 77 Mr. Waseem Allabadi Gas ACT II 3) Gay- lussac's law: (Constant volume (closed rigid container)) P and T As T increases speed increases kinetic energy increases collision strength increases and because the volume is constant, pressure will increase.  Standard temperature and pressure. T: 273 K P: 1 atm (760mmHg) Combined gas law: P1V1 = P2V2 T1 T2 Q: The pressure of 2 L of a gas at 27 C0 is 740 mmHg. What is the volume of this gas at standard temperature and pressure? (A) V= 2 × 27 ×740 273 760 (B) V= 2 × 273 ×740 300 760 (D) V= 2 × 273 × 760 300 740 (E) V= 2 × 273 × 760 300 740 Q: An ideal gas initially at 400 K and 2 atm occupies a volume of 1 L. when the gas is compressed to a volume of 0.5 L while its temperature is increased to 800 K what is the final pressure of the gas? (A) 1 atm (C) 4 atm (D) 8 atm (E) 16 atm 78 Mr. Waseem Allabadi Gas ACT II Dalton's law: Gas mixtures don't react with each other. P total= P A+ P B + P C+…… Total pressure = partial pressure of gas A + partial pressure of gas B+ partial pressure of gas C+ ……. P A = P total. X A XA : mole fraction of gas A XA= mol A (mol A + mol B + mol C) total mol Q: a sealed vessel contains 0.040 mole of gaseous He and 0.060 mole of gaseous Ne at constant temperature and a total pressure of 10 3 mm Hg. What is the partial pressure of He gas? (A) 40 mmHg (C) 100 mm Hg (D) 400 mm Hg (E) 600 mmHg Q: a closed flask contains 2 mol of O2(g). 1 mol of N2 (g). If the partial pressure of O2 (g) in the flask in 200 mm Hg. What is the partial pressure of N2 (g) in the flask? (A) 100 mmHg (B) 200 mmHg (D) 560 mmHg (E) 760 mmHg Ideal gases and kinetic molecular theory: Ideal gas: no attractive and repulsive forces. At high volume low pressure, and at high temperature. Q: Deviation from ideal gas: (A) At high temperature and pressure (B) At high temperature and low pressure (C) At low temperature and pressure (D) At low temperature and high pressure Kinetic molecular theory: 1) as temperature increases speed of molecules increases 2) As molecular mass increases speed decreases H2 O2 2×1=2 g/mol 2×16= 32g/mol Faster slower 79 Mr. Waseem Allabadi Gas ACT II 3) Different gases at the same conditions have the same kinetic energy. Ke= 1/2 m V2 (V is velocity) H2 O2 Faster slower Smaller mass larger mass 4) Gas particle have different speeds at the temperature due to collisions 5) Gas particles move in random way due to collision 6) Collisions between gas particles are elastic ( no loss or gain of energy ) 7) Gases consist of discrete particles( atoms (noble gases), molecules H2,O2,N2, and Cl2) 8) No attraction no repulsion between the gas particles. 9) Volume of gases is negligible compare to the volume of container. Ideal gas law: PV= nRT P: pressure ( in atm) ( mmHg/760) V: volume ( in L) (mL/1000) n: number of moles mol= mass/ molar mass Example: Molar mass of NH3 (1N+3H)= 1×14+3×1=17 g/mol R: gas constant (0.082) T: temperature in K= (C0+273) Q: one mol of an ideal gas is at a pressure of 800 mm Hg and a temperature of 127 C0. The volume of the gas is represented by which of the following expression: (A) V=800/760 1×0.082×400 (B) V=1×0.082×400 760/800 (C) V=1×0.082×127 800/760 (D) V=1×0.082×12.7 760/800 80 Mr. Waseem Allabadi Gas ACT II Density of gases: Gases have low densities in comparison to liquids or solids. Unit of gas density is g/L Density= mass/ volume Density(gas)= P×molar mass R×T or if the gas under STP conditions ( 0 C0 273K, 1 atm ) Density = molar mass/ 22.4 Q: the density of C2H4 gas at 0 C0 and 1 atm: (A) 0.79 (B) 1.30 (C) 1.70 (D) 2.50 Avogadro's law of gases: Different gases with the same volume and under the same temperature and pressure have the same number of moles ( or molecules ). Example: O2 PV = n O2 R T N2 PV = n N2 R T At the same P,T,and V n O2 = n N2 Q: 3 L of O2 gas sample contains 8.02×10 23 molecules under the same temperature and pressure. 3 L of N2 contains how many molecules: (A) 6.02×10 23 (B) 8.04×10 23 (C) 12.04×10 23 (D) 24.08×10 23 (E) 3.01×10 23 Collection of gases in the lab. 1) By water displacement ( or over water ) 81 Mr. Waseem Allabadi Gas ACT II Used to collect insoluble gases ( H2, N2, O2, CH4) can be collected ( HCl, SO2, SO3,NH3,H2S ) can't be collected because they are soluble gases CO2 can be collected but some of it dissolve in water. 2) By air displacement: 1) Downward delivery (upward displacement) When the density of gas is lower than the density of air. The density of air = 1.29 g/L Molar mass of air = 29 g/ mol 2) Upward delivery (downward displacement) When the density of gas is heavier than the density of air. H2 2g/mol < 29 g/mol CH4 16g/mol < 29 g/mol NH3 17g/mol < 29 g/mol SO2 64g/mol > 29g/mol 3) By using gas syringe To collect any type of gases Diffusion and effusion: Diffusion through the space or a second substance from a high concentration to a low concentration. Effusion into an evacuated chamber. As the molar mass increases its diffusion rate decreases. Q: the gas that has the highest rate of diffusion is: (A) O2 (B) N2 (C) H2 (D) He 82 Mr. Waseem Allabadi Gas ACT II Practice Questions 1. A rise in temperature causes an increase in all of the following EXCEPT the A. Vapor pressure of a liquid. B. Volume of a gas sample held at constant pressure. C. Average kinetic energy of the molecules in gas sample. D. Mass of the molecules in a gas sample. 2. In order to determine accurately the volume of a gas collected over water , the gas must be A. More dense than air B. More dense than water C. Insoluble in water D. At room temperature 3. which of the following is NOT consistent with the kinetic theory of gases ? A. Collisions between gas molecules result in a net decrease in kinetic energy. B. Gases consist of discrete particles (atoms or molecules ). C. Molecular motion is affected by temperature. D. Gas molecules travel in straight lines between collisions. 4.which of the following is NOT a part of the kinetic-molecular theory of gase ? A. At any instant all molecules have the same speed. B. Pressure is a result of collisions between molecules and the walls of container. C. There are no attractive forces among the molecules of an ideal gas. D. At a pressure of 1 atm , molecules in a gas are widely separated. 5. According to the kinetic molecular theory , which of the following statements concerning the molecules in a flask containing air at room temperature and 1 atm is correct ? A. They move at constant speed B. They have the same mass number C. They move in one direction D. They collide with one another as well as with the walls of the flask 6. Increasing the temperature of a gas in a rigid closed container increases which of the following ? I. The pressure of the gas II. The average speed of the gas molecules III. The mass of the gas A. I only B. I and II only C. II and III only D. I, II, and III 7. In general, when are deviations from ideal gas behavior greatest ? A. When both pressure and temperature are high B. When both pressure and temperature are low C. When pressure is high and temperature is low D. When pressure is low and temperature is high 8. For equal volumes of CH₄ (g) (molar mass 16.0 g/mol) and of O₂(g) (molar mass 32.0 g/mol ) kept at the same temperature and pressure , which of the following statements is correct ? A. The two gases have the same density. B. The number of CH₄ (g) molecules equals the number of O₂(g) molecules. C. The average kinetic energy of the CH₄ (g) molecules is less than that of the O₂(g) molecules. D. The average speed of the CH₄ (g) molecules is less than that of the O₂(g) molecules 83 Mr. Waseem Allabadi Gas ACT II 9. Equal volumes of different gases at the same pressure and temperature have the same A. Mass B. Number of neutrons C. Number of electrons D. Number of molecules 10. When two gases are mixed at constant temperature , the total pressure of the mixture may depend on which of the following ? I. The number of moles of each gas II. The size of the container III. Any reaction that occurs between the gases. A. I only B. II only C. II and III only D. I, II, and III 11. a freshly opened can of carbonated beverage fizzes because A. An acid is reacting with dissolved solids B. The solution is boiling C. Oxygen from the air reacts with dissolved carbon D. Dissolved gas escapes when the pressure is reduced 12. of the following the conditions under which the molar volume of gaseous helium is greatest are A. 237 K and 1 atm B. 300 K and 1 atm C. 400 K and 1 atm D. 500 K and 2 atm 13. which of the following gases has the highest rate of diffusion at a given temperature and pressure ? A. CH₄ B. C₂H₄ C. C₂H₆ D. C₃H₈ 14. A 1 L sample of which of the following gases would have the same mass as 1 L of nitrous oxide, N₂O , under the same conditions of temperature and pressure ? A. C₃H₈ B. C₃H₆ C. CH₃OHC₃ D. CO 15.at given temperature and pressure , which of the following gases has the highest rate of effusion through a pinhole ?? A. H₂ B. O₂ C. O₃ D. H₂O E. H₂O₂ 84 Mr. Waseem Allabadi Gas ACT II 16. at 25˚C , the average speed of the molecules is greatest for which of the following gases ? A. CH₄ B. C₄H₁₀ C. F₂ D. Ar 17. under the same conditions , which of the following gases diffuses most slowly ? A. CH B. C₂H₄ C. C₂H₆ D. CO₂ 19. A sealed vessel contains 0.040 mole of gaseous He and 0.060 mole of gaseous Ne at constant temperature and a total pressure of 10³ mm Hg. What is the partial pressure of the He gas ? A. 40 mm Hg B. 60 mm Hg C. 100 mm Hg D. 400 mm Hg 20. A glass vessel contains 2.0 mol of N₂(g) and 3.0 mol of Ar(g) , the total pressure is 1.5 atm. What is the partial pressure of the Ar(g) ? A. 0.2 atm B. 0.6 atm C. 0.9 atm D. 1.0 atm 21. When 3.0 mol of helium gas and 2.0 mol of neon gas at a temperature of 10. ˚C are added to an empty flask, the total pressure in the flask is 10. atm. The partial pressure of the neon gas is A. 0.20 atm B. 0.50 atm C. 2.0 atm D. 4.0 atm 22. A closed flask contains 2.00 mol of O₂(g) 1.00 mol of N₂(g). if the partial pressure O₂(g) in the flask is 200. mm Hg. when partial pressure of N₂(g) in the flask ? A. 100. mm Hg B. 200. mm Hg C. 300. mm Hg D. 560. mm Hg 23. A gas mixture at a total pressure of 600. mm Hg contains 0.500 mol of nitrogen , 0.300 mol of argon , and 0.200 mol of oxygen. The partial pressure of oxygen is A. 120. mm Hg B. 180. mm Hg C. 200. mm Hg D. 300. mm Hg 24. A vessel contains 0.10 mol of oxygen gas and 0.20 mol of nitrogen gas. The total pressure inside the vessel is 1.5 atm. Which of the following is correct ? A. The partial pressure of oxygen is 0.75 atm B. The partial pressure of nitrogen is 0.75 atm C. The partial pressure of oxygen is 0.50 atm D. The partial pressure of nitrogen is 0.50 atm 85 Mr. Waseem Allabadi Gas ACT II 25. At constant temperature, the relation between the pressure P and the volume V of a fixed mass of an ideal gas is given by which of the following expressions ? A. PV = a constant B. VP = a constant C. ½ PV² = a constant D. PV = a constant 26. A certain amount of an ideal gas at 1 atm is contained in a closed rigid vessel at 20˚C (293 K). if the temperature is increased to 200˚C (473 K ), the new pressure of the gas will be A. Greater than 2 atm B. Greater than 1 atm, but less than 2 atm C. Equal to 1 atm D. Greater than ½ atm , byt less than 1 atm 27. A 2.0 g sample of a gas has a volume of 1.0 L at 27˚C and 1 atm, the molar mass of the gas is closest to A. 4.0 g/mol B. 12 g/mol C. 22 g/mol D. 49 g/mol 28.one mole of an ideal gas is at pressure of 800 mm Hg and a temperature of 127˚C. the volume of the gas is represented by which of the follow