Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Summary of know/edge ",--_.- ~ .-----~--- " ,ALgo r (JoWr z: SUMMARY OF KNOWLEDGE C Atoms consist Of a nucleus surrounded by one or more electron shells. [2 Atoms contain three sub-atomic particles: electrons, protons and neutrons. [J Protons and neutrons occupy the nucleus of the atom. The electrons move round the nucleus in orbits. 9 Protons are- positively charged and have-a mass almost identical to neutrons. Neutrons are not electrically charged. Electrons carry a negative charge equal in size to that of the proton. D The sub-atomic particles can be distinguished by their behaviour in electric and magnetic fields. w Atoms are electrically neutral due to the presence of equal numbers of protons and electrons. 67 68 A TOMle STRuaURE -. The atomic number (Z) is the number of protons 'in an atom of a chemical element The mass number (A) is the number of protons and neutrons in the nucleus in an atom of a chemical element. ~ A nuclide is an atom with a specific number of protons and neutrons. Nuclides are described by the notation 1X, where Z represents the atomic number, A represents the mass number and X represents the symbol of the chemical element. Positive ions are formed when atoms lose electrons; negative ions are formed when atoms gain electrons. Isotopes of a chemical element have the same atomic number but different mass numbers. Isotopes of a chemical element have the same chemical properties, but slightly different physical properties. A mass spectrometer is used to show the isotopic composition of a chemical element and determine its relative atomic mass. The relative atomic mass is the mass of a weighted average of a sample of atoms compared to one twelfth of the mass of a carbon-12 atom. The relative isotopic mass is the mass of a nuclide compared to one twelfth of the mass of a carbon-12 atom. A mass spectrometer ionizes gaseous atoms and then deflects a beam of its positive ions. A variable magnetic field is used to deflect the ions and bring them to a detector. Varying the magnetic field strength brings ions of higher mass to the detector, creating a mass spectrum. .. The isotopes of some chemical elements are radioactive and release ionizing radiation.,> These isotopes are known as radioisotopes. The three types of ionizing radiation are alpha radiation (helium nuclei), beta radiation (electrons) and gamma radiation (gamma rays). Radioactive decay is an exponential process with a characteristic half-life, which is the time for the rate of radioactive decay to decrease by half. The half-life is independent of the amount of the isotope. i.; Radioisotopes have many uses in medicine (as tracers and radiotherapy) and dating techniques, for example, radiocarbon dating. The uses of radioisotopes depend on their half-life and the radiation they release. Light is a form of energy and can be regarded as having the properties of an electromagnetic wave and a particle. Waves are described by their wavelength, frequency and speed. A continuous spectrum contains all wavelengths from a band of the electromagnetic spectrum. Each line in an emission spectrum of an atom or chemical element corresponds to electrons transitioninq from one energy level to another. During electron transitions electromagnetic radiation is emitted or absorbed. The lines in an emission or line spectrum become closer together as wavelength decreases (or frequency increases). The Bohr model of the atom assumes that electrons rotate in circular orbits around the nucleus. The electrons in each orbit have a fixed amount of potential energy. The frequency of the electromagnetic radiation and the gap between two energy levels is given by Planck's equation: !'!f == hi, where f represents the frequency and h represents Planck's constant. The ground state is the state of an atom in which all electrons have their lowest energies. If electrons are given additional energy and move from a lower to a higher state, then the atom is said to be excited. The Balmer series is caused by excited electrons falling back to the second energy level. The Lyman series is caused by excited electrons falling back to the first energy level. The ionization energy for the hydrogen atom is the energy needed to remove the electron from the ground state of the gaseous atom. ·'1 L. -------··.11 BD I Examination questions 69 Examination questions - a selection • Paper 1 IB questions and IB style questions 01 Which statement is correct about the isotopes of an element? A They have the same mass number. B They have the same numbers of protons and neutrons in the nucleus. e They have more protons than neutrons. D They have the same electron arrangementor confiquration. 02 A chemical element with the symbol X has the electron arrangement 2,8,6. Which chemical species is this chemical element most likely to form? A the ion X3+ e the compound HzX [2Wxz-J B the ion X6+ D the compound XFa[X8+8F-] Q3 Which of the following particles contains more electrons than neutrons? I 1H A I only B II only II T~CI- III f§K+ e I and II only D II and III only Standard Level Paper 1, May 00, 06 Q4 What information about the structure of a helium atom can be gained from its emission spectrum? A Most of the mass of the atom is in its nucleus. B A helium atom contains two electrons and two protons. e The electrons in the helium atom are held near the nucleus. D The electrons may exist in any of several energy levels. Q5 An element has the electron arrangement 2,8,6. What is the element? ACe S B P D Ar Q6 Which is an incorrect statement about the atomic emission spectrum of hydrogen? A The frequency of each line depends on the difference in energy between the higher and lower energy levels. B The spectrum consists of several series of lines. e Electronic transitions to the leve~·n = 2 give rise to lines in the visible region. D It is a continuous spectrum. Q7 What is the correct number of each particle in a fluoride ion, 19F-? Protons Neutrons A 9 10 B 9 10 e 9 10 D 9 Electrons. 8 9 10 19 10 Standard Level Paper 1, Nov 03, 05 Q8 What fraction of a radioisotope will remain after three half-lives? A 1/16 B 1/8 e 1/3 D 3/4 Q9 Why was the Bohr theory of the atom developed? A To account for changes in gas volumes with temperature. B To account for the ratios by mass of elements in compounds. To account for the emission or line spectrum of hydrogen atoms. D To account for chemical formulas. e Q10 A particular element consists of two isotopes: 72% of mass number 85 and 28% of mass number 87. What is the expected range of the relative atomic mass? A less than 85 B between 86 and 87 betvveen 85 and 86 D more than 88 e Q11 How many valence electrons (electrons in the outermost shell) are present in the element of atomic number 147 A4 c z 83 D1 Q12 Which one of the following atoms will have the same number of neutrons as an atom of ~~Sr? A §~y e ~§Sr 8 ~fRb D ~Kr Q13 Which statement is correct for the emission spectrum of the hydrogen atom? A The lines converge at lower energies. B The lines are produced when electrons move from lower to higher energy levels. e The lines in the visible region involve electron transitions into the energy level closest to the nucleus. o The line corresponding to the qreatesternission of energy is in the ultraviolet region. Standard Level Paper 1, Nov 03, 06 Q14 Naturally occurring chlorine consists of the isotopes chlorine-35 and chlorine-37. The relative atomic mass of chlorine is 35.5. Which one of the following statements is true? A The chlorine-35 and chlorine-37 atoms are present in equal amounts. 8 The ratio of chlorine-37 atoms to chlorine-35 atoms is 2 : 1. e The ratio of chlorine-37 to chlorine-35 atoms is 37/35. D There are three times as many as chlorine-35 atoms as chlorine-37 atoms. 70 A TOMle STRUaURE Q15 Which statement is correct about a line emission spectrum? A Electrons neither absorb nor release energy as they move from low to high energy levels. B Electrons absorb energy as they move from high to low energy levels. C Electrons release energy as they move from low to high energy levels. DElectrons release energy as they move from high to low energy levels. Standard Level Paper 1, Nov OS, 06 Q16 Which electronic transition within a hydrogen atom requires the greatest energy? A n= 1 ~n=2 B n=3~n=5 Q18 The atomic numbers and mass numbers for four different nuclei are given in the table below. Which two are isotopes? Atomic number II III IV A I and II B II and III Mass number 101 102 102 103 258 258 260 259 .C III and IV 0 I and IV Standard Level Paper 1, Nov 98, 06 Q19 All isotopes of uranium have the same: I number of protons 1/ number of neutrons III mass number A I only C III only B II only D I and III only Q20 Which is the correct sequence for some of the various stages that typically occur in the analysis of an element during mass spectrometry? A vaporization, electron bombardment, acceleration, deflection, detection B electron bombardment, vaporization, acceleration,deflection, detection C vaporization, electron bombardment, deflection, acceleration, detection o deflection, acceleration, electron bombardment, vaporization, detection and IB style questions Q1 The element bromine exists as the isotopes 79Br . and 81Br,and has a relative atomic mass of 79.90. a Copy and complete the following table t9 show the numbers of sub-atomic particles in the species shown. [3] An atom of 79Br An ion of 81Br Protons Neutrons Electrons b C n=2~n=3 D n=5~n=oo Q17 Which of the following radioisotopes is used in nuclear medicine to image the thyroid gland? A iodine-131 C fluorine-18 B carbon-14 D uranium-235 I Paper 2'18 questions State and explain which of the two isotopes 79Brand 81Bris more common in the element bromine. [1] Standard Level Paper 2, Nov OS, 03 Q2 The element silver has two isotopes, l~~Agand lS1Ag,and a relative atomic mass-of 107.87. a Define the term isotope. [1] b State the number of protons, electrons arid neutrons in 12tAg+. [2] c State the name and the mass number of the isotope relative to which all atomic masses are measured. [1] Q3 A sample of iridium is analysed in a mass spectrometer. The first and last processes in mass spectrometry are vaporization and detection. a i State the names of the second and third processes in the order in which they occur in a mass spectrometer. [2J ii Outline what occurs during the second process. [2] iii State and explain which one of the following ions will undergo the greatest deflection (under the same conditions in a mass spectrometer): 1911r+ or 1931r+ [2] b The sample of iridium is found to have the following composition of stable isotopes: Isotope Relative abundance/% i Ir-191 Ir-193 37.1 62.9 Define the term relative atomic mass. [2] ii Calculate the relative atomic mass of this c sample of iridium, giving your answer to two decimal places. [2] Iridium-192 isa short-lived radioisotope used to treat cancer.Define the term radioisotope and name another radioisotope used in nuclear medicine. [2] Q4 Describe the emission or line spectrum of gaseous hydrogen atoms and explain how this is related to the energy levels in the atom.' [3] •• II II III Ii II II •..•••••.•.•••••••••• ~ •••••.•••••••••••• '-~.....•...• ~.,......:-::'"~-,...,....:_--..:::-::~,.-~ ..~~-.-,.~. Atomic structure 1023 _.::_ .._-::..~:'~ - .•.. -··~.....,.-'.:':"':':::"-· __ -~~_:::-;--~.·7.:::~·--_;:_-_:_-_:::::_:_:_P7_---_:::~~-··-7~-~,,~~_, ------------ 2 Atomic structure Paper 1 IB questions and IB style questions Q1 0 Isotopes are two or more atoms of the same element with different numbers of neutrons (and therefore different relative isotopic masses). They have the same numbers of protons and hence the same number of electrons. 02 C HzS: [2H+ XZ-j Q3 A 1P ii Hi Q4 D 17p 19 P 1e 18 e 18e On 18 n 20 n Emission spectra provide evidence that electrons exist in fixed energy levels within an atom. Q5 C The atomic number of sulfur is 16; it is in group 6 and therefore has six valence electrons. 06 0 A discontinuous spectrum is produced. Q7 C Q8 B Q9 C 9 protons (from periodic table); 10 electrons in atom (due to the single negative charge) and 10 neutrons (19 - 9). 1~.l~.!.~.!. 2 4 8 To account for the emission or line spectrum of hydrogen atoms. 010 C Relative atomic mass is a weighted average. The more abundant isotope has a mass number of 85. 011 A The element is silicon with the electron arrangement 2,8,4. 012 B ~Y: 52 neutrons ~~Rb:52 neutrons ~Sr: 51 neutrons ~Kr: 48 neutrons. 013 0 They converge at high energies; lines are produced when electrons move from higher to lower energies. Q14 0 . . f chlori ReIauve atomic mass 0 onne = (35x3)+(37x1) 3+ 1 =142=355 4 015 0 . Electrons release energy as they move from high to low energy levels: light is released. 016 A -Energy levels rapidly converge with an increase in n. 017 A The thyroxine secreted by the thyroid is an iodinecontaining compound. 018 B I. 157 neutrons; II. 156 neutrons; III. 158 neutrons and IV. 156 neutrons. 019 A Atoms of isotopes have the same number of protons have the same number of protons but different numbers of neutrons. The mass number is the number of protons and neutrons. 020 A The sample is first vaporized by lowering the surrounding pressure: a mass spectrometer can only accept gaseous atoms or molecules. The molecules or atoms are then subjected to bombardment by high energy electrons which remove a valence electron to form unipositive ions. The ions are then accelerated by an electric field before entering a powerful magnetic field where they experience deflection according to their mass-to-charge ratio. 1024 ANSWERS TO EXAMINA TlON QUESTIONS -------- Paper 2 IS questions and IS style questions An atom of 79Br An ion of8'Br Q1 a b Q2 a b c Q3 a Protons 35 35 Neutrons 44 46 Electrons 35 36 [31 nBr because its relative isotopic is closer to 79.90 [1] Atoms of the. same.element but with different mass numbers and hence different number of neutrons. 47 protons, 46 electrons and 60 neutrons Carbon-12 (12C). !1J III [11 i Ionization, acceleration 12J The gaseous atoms formed from the sample are bombarded with electrons traveliing at high speed with large kinetic energies. !2J iii '"'Ir+ [1] It has the lowest mass-to-charge ratio (or lowest mass since the Charges are identical). [1J The weighted average of relative isotopic masses of all the stable isotopes of an element relative to the mass of one atom of carbon-12. 12J ii (191 x0.371)+(l93 x 0.629) = {70.861 + 121397) = 192.258 = 192.26 (2 d.p.) [2J iii An unstable isotope of an element that decays or disintegrates spontaneously, emitting nudear or ionizing radiation, [1] Cobalt-60, molybdenum-99, chromium-51, iron-59 and iodil1e-125 [1] ii b Q4 The emission spectrum consist of a series of sharp or discrete coloured lines on a black background. [1 J The lines converge together at high energy (high frequency). The lines are generated when excited electrons move from high energy levels to lower energy levels. [1 J light of a particular frequency is released during this process. [1J ---------------- ---~---- 340 A TOMle STRUCTURE TAL·Bol CH--Pf SUMMARY OF KNOWLEDGE 11.. "._ The first ionization energy is the minimum energy required to remove a mole-of electrons from a mole of gaseous atoms to form a moie of gaseous unipositive ions (under standard thermodynamic conditions). The ionization energy of an atom or ion is determined by the following factors: nuclear charge, shielding effect and atomic radius (distance between nucleus and outer electrons). On moving across periods 1, 2 and 3 there is a general increase in first ionization energy. This is due to a large increase in nuclear charge which is accompanied by a small increase in shielding. The nuclear charge is the force that attracts all the electrons to the nucleus. The shielding effect is the effect of shielding the outer electrons from the attraction of the nucleus by the repelling effect of the inner electrons . .-' There-are-two small decreases in first ionization energy in periods 2 and 3. The first is due to a change in sub-shell (sub-level) and loss of shielding and the second is due to enhanced -electron-electron repulsion. These factorsoutweighi:he increase in nuclear charge with increasing atomic number. Electrons can be progressively removed from gaseous atoms and successive ionization energies measured and plotted against number of electrons removed. The graphs offer strong experimental evidence for shells (main energy levels) and sub-shells (sub-levels). The relatively large increases in successive ionization energies correspond to changes in shells; smaller changes correspond to changes in sub-shells or electron pairing. -. Each shell can hold a maximum number of electrons: first shell (two electrons), second shell (eight electrons), third shell (18 electrons) and fourth shell (32 electrons). (The number of electrons in each shell is given by 2n2, where n represents the shell number.) =.::. Each shell (energy level) is composed of one or more sub-shells (sub-levels), The first shell is composed of one sub-shell, the second shell is composed of two sub-shells, the third shell is composed of three sub-shells and the fourth shell is composed of four sub-shells. "_ Electrons occupy regions or volumes of space called orbitals. Each atomic orbital can accommodate a maximum of two electrons. Electrons can occupy four types of orbitals: s, p, d and f. s orbitals are spherical in shape; p orbitals are 'dumb-bell' shaped and arranged mutually perpendicularly to each other. p, d and f orbitals all have the same energy within the same shell or sub-shell. Orbitals (of the same type) retain the same shape but become larger (and their electron density more diffuse) with an increase in shell number. Summary of know/edge 341 ,~. C The first shell has an s sub-shell (one orbital); the second shell has an s and a psub-shell ::J r: L.. ':J c: ;] (three orbitals); the third shell has an s sub-shell, a p sub-shell and a d sub-shell (five orbitals); and the fourth shell has an s sub-shell, a p sub-shell, a d sub-shell and an f sub-shell (seven orbitals). Electrons are assigned to atomic orbitals of atoms (in their ground state) according to the Aufbau principle, the Pauli exclusion principle and Hund's rule. The Aufbau principle states that electrons occupy atomic orbitals in the order of the energy levels of the orbitals. . Copper and chromium atoms have anomalous electron configurations which violate the Aufbau principle. The chromium atom is 3d54s1 and the copper atom is 3d104si because d sub-shells that are half-filled or fully filled are particularly stable. The Pauli exclusion principle states that only two electrons may occupy the same orbital and that these two electrons must have opposite spins. Hund's rule states when electrons are placed in a set of atomic orbitals with equal energies, the electrons must occupy them singly with parallel spins before they occupy the orbitals in pairs. As a consequence of Hund's rule atoms tend to maximize the number of unpaired electrons. The electron configuration of an atom describes how electrons are distributed among the various orbitals in the various shells and sub-shells and is denoted by an orbital diagram or spdf notation. In orbital notation each orbital is represented by a box and each electron by an arrow. The arrow head indicates the direction of electron spin. In spdf notation the electron configuration is denoted by writing the symbol for the occupied sub-shell and adding a superscript to indicate the number of electrons in that subshell. This can be condensed by describing the core electrons with the nearest noble gas electron configuration. Simple positive ions are formed by removing the appropriate number of outer electrons; simple negative ions are formed by adding the appropriate number of electrons to the outer sub-shell. Excited atoms have one or more electrons promoted to higher energy levels. Many ions formed by non-transition metals obey the octet rule. These ions have a full outer shell of eight electrons. Excited atoms are formed when one or more electrons are promoted to higher energy orbitals. Examination questions - a selection Paper 1 IB questions and IB style questions Q4 An atom of chlorine has the electron configuration [Ne]3s23pS What is the number of orbitals occupied by at least one electron? A 7 B 9 C 13 D 17 Ql What is the electron configuration for an atom with Z = 22? A 1s22s22p63s23p63d4 B 1s22s22p63s23p64s24p2 c 1s22s22p63s23p63d24p2 D 1s22s22p63s23p64s23d2 Higher Level Paper l;"May 03, Q5 Q5 For which of the following Q2 For the species below, which one would require Q6 Which of the following the most energy for the removal A Na+ B F C F- of an electron? D Ar Q3 How many unpaired electrons would the FeZ.•.ion be expected to have? A1 B3 (4 D6 pairs of species are the chemical properties most similar? A tH and tW C ilNa and l§K B l~C and I~N D ~Li'"and ~Be first ionization energy? A 19K B 6C (12Mg D 4He Q7 What is the total number of electrons in d orbitals in a tin atom 50Sn? A 5 B 10 (20 D 0 Q8 Which of the following ionization energy? A Na B F -.:....-,...,.,-...,,---_.- -------. elements has the smallest atoms has the smallest first C Be D CI 342 A TOMle STRuaURE Q9 What is the electron configuration [Ar] ( [Ar]3d5 B [ArJ4s23d4 D [Ar]3d5 B ( [ArJ4s23d3 D energies for two X and Y, are given below. elements, Q10 What is the ground state electron configuration the Fe3+ion? A [ArJ4s13d5 Q20 The first three ionization of C03+? A Element First ionization energy/kJ mol-' of -------- x [Ar]3d5 [Ar]3d3 y Q11 Which set of chemical species is arranged in order of decreasing ionic radius? A Ca2+ > 5r2+ > Ba2+ (K+ 52- B A13+> Mg2+ > Na+ Na+ > (1- > D 02- > F- > Which A following A 12Mg B 2sMn (3SBr D 36Kr for the atoms is incorrect? 1s22s22p63s2 7300 11800 1086 2350 4620 pair of elements Na+ B Na 3Li and 6C ( B 4Be and lSP (Mg2+ D A13+ Q15 How many unpaired electrons would the cobalt 27) in its ground state be expected to have? (4 D3 Q16 Which property changes along with increasing ionization energy? A increasing non-metallic properties increasing atomic radii decreasing electron affinity decreasing nuclear charge A B 4Be (6C orbitals? D 7N Q18 What is the maximum number of electrons that can occupy the 5d sub energy level? A 20 B 10 (32 D 25 Q19 Which of the following first ionization A lBArB elements chemical C Mg2+ species has the D 52- F- (6C (Cu2+ B Ni D Ni2+ series differ from each other mainly in the number electrons? A p electrons B sand p electrons ( p and d and f electrons D d electrons of which Q24 Which is the best periodic correlation atomic number and another A atomic masses B ionic and atomic radii ( first ionization energies D electron configuration atomic type of between property? D 55(S series are the atoms in order of increasing first ionization A Be, Mg, Ca C Be, B, C B 0, F, Ne D Ne, 0, F Q26 What increases in equal steps of one from left to right in the periodic table for the elements to neon? A B has the largest energy? lSP BO and 165 D 2He and 4Be Q23 Elements in the first transition arranged energy? state have two half-filled lSP and Y? Q25 In which of the following Q17 For which element would neutral isolated atoms in the ground X configuration? species would you expect to have the largest radius? A Na+ B N3(Ne D F- B ( D largest radius? A A13+ B A Mn Q14 Which of the following atom (element number electronic configuration AS B2 could represent ,-, Q22 Which atom or ion could have a 3dB electronic [ArJ4s23d4 [Ar]3d104s24p5 [Ar]4s23d104p6 Q13 Which chemical species has the smallest radius? A ---- 520 Q21 Which of the following Q12 Which ground state configuration Second Third ionization ionization energy/kJ mol-' energy/kJ mol-1 the number the number - isotope of occupied of neutrons electron energy levels in the most common ( the number of electrons in the atom D the atomic mass Higher Level Paper 1, May Q27 In which of the following configurations j {He12s22p2 II [He]2s22p3 I only I and II only as, Q6 ground-state electron are unpaired electrons present? III [He]2s2p4 A B lithium c D I, II and III II and III only Examination questions 343 Q28 What is the number of unpaired electrons in the Cr3+ion? A B 2 C 3 0 5 d State which period element C is in and explain your reasoning. [3] e Why is the first ionization energy of element G lower than that of element F? [2] a Q29 A transition metal ion has the electronic configuration X3+= [Ar]3d4. What is the atomic number of element X? A 24 B 22 C 25 0 26 Q30 What is the total number of p orbitals containing one or more electrons in germanium (atomic . number 32)? A 2 B 3 C 5 0 8 Higher Level Paper 1, May 04, Q5 Paper 2 IB questions and IB style questions Q1 The graph below ionization energy Figure 1 refers to Figure 2 refers to group as element shows the variation in first of some chemical elements. the chemical elements and chemical elements in the same C. F Q2 The successive ionization energies of germanium are shown in the following table: ---Ionization energy/ kJ mol-1 1st 2nd 3rd 4th 5th 760 1540 3300 4390 8950 a Identify the sub-level from which the electron is removed when the first ionization energy of germanium is measured. [1] b Write an equation, including state symbols, for the process occurring when measuring the second ionization energy of germanium. [1] c Explain why the difference between the 4th and 5th ionization energies is much greater than the difference between any two other successive values. [2] Higher Level Paper 2, Nov 05, 02 Q3 The graph shows the variation in second ionization energy of the elements silicon to argon. L- 0 E J .Y <, >- e' CD c 2900 2700 / /~ 2500 CD c 2300 ~ 2100 0 x x+1 x+2 x+3 x+4 N 'c .Q Atomic number Figure 1 -0 C 0 o 1900 1700 <D (j) ./ 1500 / Si / ~ P S CI Ar Element >- e'<D a Explain the general increase in second ionization energy from silicon to argon. [3] b Give the detailed electron configurations of the S-;-(g)and CI+(g) ions. [2] C Explain why the increase from sulfur to chlorine is significantly less than the increase to sulfur from phosphorus. [1] c (J) c o ~ .~ c .Q x-8 x x+8 Atomic number Figure 2 a -: Define the term first ionization energy of an element. [2] b Element C, with atomic number x, is in group 2 of the periodic table. Justify this using all the information from Figure 1. [4] C Explain the trend in the first ionization energy as shown in Figure 2. [3] Atomic structure 103s--------------------------------------------------------- 12 Atomic structure Q12 B The electron configuration for a manganese atom is 4s23d5• Q13 0 Na 2,8,1 (11 protons), Na+2,8 (11 protons), Mg'+ 2,8 (12 protons) and AP+ 2, 8 (13 protons). The radii decrease due to the increasing effectivenudear charge. Q14 B Na+2,8 (11 protons), N3--2,8(7 protons), F- 2,8 (9 protons), Ne 2,8 (10 protons). The electrons in the nitride ion are experiencing the lowest effective nuclear charge. Q15 D A cobalt atom has the following electron configuration: 1 s'2s'2p63s'3p63d74s'. The 3d sub-shell has two spin pairs and three unpaired electrons. Q16 A Ionization energies generally increase across periods. This results in a decrease in metallic behaviour, that is, a decrease in tendency to form cations (positive ions). Q17 C A carbon atom has the ronowing electron configuration 1522£221". The p sub-shell has two unpaired electrons in accordance with Hund's rule. Q18 All d sub-energy levels have 10 orbitals and hence can hold to a maximum of 10 electrons. B Q19 A Ionization energies decrease down a group and generally increase across a period. Ihe electrons in argon are experiencing the highest effective nuclear charge. Q20 A X is in group 1 and Y is in group 2 or 4. Q21 0 Al3+2,8 (13 protons), Mg'+ 2,8 (12 protons), F- 2,8 (9 protons), 52- 2,8,8 (16 protons). The electrons in the sulfide ion are experiencing the lowest effective nudear charge. Paper 1 IB questions and IB style questions Q1 D Q2 Titanium has an atomic number of 22. Its electron configuration is 1s'2s'2p63s'3p64s'3cf2. A The removal of an electron from a sodium ion would involve the removal of a core electron from the second shell close to the nucleus. This electron wouid be strongly held and hence would have a large ionization energy. This electron would have the highest effective nuclear charge. Q3 C The Fel+ion has the configuration [Ar]3d6. Hence there is one spin pair and four unpaired electrons. 04 B 1s sub-shell: 1 occupied orbital; 2s sub-shell: 1 occupied ,. orbital; 2p sub-shell: 3 occupied orbitals; 3s sub-shell: 1 occupied orbital; 3p sub-shell: 3 occupied orbitals: two filled and one half-filled. 05 C The two elements are both in the same group (group 1). Elements in groups, especially those at the extremes of the periodic table, have very similar chemical properties. Q6 A Ionization energies decrease down a group and generalty increase across a period. Q7 C A tin atom has the following electron configuration: 1s22s'2p63s'3p63dl04s24p64dl°5s'5p>. Q8 A Ionization energies decrease down a group and generally increase across a period. The valence electron in sodium is experiencing the lowest effective nuclear charge. Q9 D Tne electron configuration of the cobalt(IIf}'ion is: W2s'2p63s23p63d6. The 4s electrons are lost from a transition metal atom before the 3d electrons. Ql0 C The-electron configuration of the iron(m)ion is 1s22s22pG3s'3p63d5. The 4s electrons are lost from a transition metal atom before the 3d electrons. Q11 D The electron arrangements of the species are all 2,8, that is. they are ISO electronic. However, the atomic (proton) numbers are 8, 9 and 11. The effective nuclear charge increases left to right and hence radii decrease. Q22 0 Ni 3dE4s2;Mn 3d54s2; CUlT 3d9; Ni2+3dS Q23 0 Elements in the first transition series (with the exception of copper and chromium) have a 3&4s2 configuration. Q24 D The atomic number is the number of protons in the nucleus. The proton number determines the number of electrons and hence the electron configuration. Q25 B Oxygen 2,6 (8 protons), fluorine 2.7 (9 protons), neon 2,8 (10 protons). Effective nuclear charge is increasing from left to right Q26 C From lithium to neon the atomic number (proton number) increases by one. This is accompanied by the addition of an extra electron. Q27 C I has two unpaired 2p electrons; II has three unpaired 2p electrons and III has two unpaired 2p electrons. Q28 C The chromium(m) ion has the electron configuration 3d3. The chromium atom has the electron configuration 3d54s1. Q29 C X will have the electron configuration is [Ar]3d54s'. Transition metal ions ionise via loss of 45 and then 3d electrons. The total number of electrons in the atom is 18 + 5 + 2 = 25. Q30 A germanium atom has the following configuration: Js22s22p€3s23p54sz 3d104p2.Three 2p orbitals, three 3p orbitals and two unfilled 4p orbitals. 0 Paper 2 IB questions and IB style questions Q1 a The first ionization energy is the minimum energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of unipositive gaseous ions (under standard thermodynamic conditions). [2] 1036 ANSWERS TO EXAMINA nON QUESTIONS 7_-=-:",:"""'-'~_-+7:;-:~ ~~~-_-. -:------:-_-~----~~_~.-_--_-. ~--.~---__"-=----_--. ---~~~--.,-~~"""---- .------------ -----_--:--:-:--: -~-- -~--~.- ---- ---'-- ~----..----~---------- b c In a group 2 element (C) the electron (for the first ionization energy) is removed from a spin pair in a s subshell. For D, a group 3 element, the electron (for the first ionization energy) is removed from a 3p sub-shell further away from the nucleus. The 3p sub-shell also experiences more shielding. Hence there is a decrease in first ionization energy when moving from C to D. The increase in nuclear charge accounts for the increase from E (np2 to F np3). In addition C cannot be in group 5 (the other group after which there is a slight decrease) because there is a steady rise for the next three elements (0 to E to F) indicating the filling of a p sub-shelll. [4} As you move down group 2 the first ionization energies decrease. As you descend the group the nudear charge increases due to the presence of additional protons, but the shielding effect progressively increases due to the presence of an extra electron shell as you move from one period to the next. The effect of the extra protons is compensated for by the effect of the extra electrons. The only factor left is the additional distance between the outer electron and the nucleus: That lowers the first ionization d e Q2 i ii Hi Q3 a energy. 13J (is in period 3 since in period 2 the group 2 element would have the higher ionization energy while in period 4 the group 2 element would have a lower ionization energy. This can be accounted for by the progressively increasing average distance the valence electron is from the nucleus. C cannat be in the first period (He) as p-orbrtals are being filled, it cannot be in the fourth period (Ca) as (x + 8) would- not bring you back to group 2 again. [3] There is a slight decrease from F to G due to the presence of a spin pair in the np sub-shell for G. The resulting electron-electron repulsion is greater than the effect of the increase in nuclear charge and reduces the ionization energy of G+. [2} 4p (1] Ge+(g) ~ Ge'+(gl+ e[1] 5th electron removed from 3rd energy level and 4th electron from 4th energy level. The attraction by the protons in the nucleus is greater since the electrons are closer to the nucleus. [2} Across a period the atomic number increases progressively by one unit, thus resulting in an additional proton being added to the nucleus and thus an increase in the attraction for all the electrons, including the outermost electron. 11] Since the number of electron shells for all elements in period 3 is constant, namely, three, there is very little increase in shielding. [1 J Hence, the increase in nuclear charge across the period will bring all the electrons progressively closer to the nucleus, causing the removal of the second electron from the atom to require progressively more energy. {1 J b The electron configurations of the S+and (1+ ions are 1SZ2S22p63s23p3 and 1s22s22p63s23p'. [2J c The (1+ ion has a spin pair of electrons in the 3p sub-shell. This electron-electron repulsion counteracts, to some extent, the increase in nuclear charge. Hence, the increase in ionization energy is relatively small. {1J