AP WORKSHEET 01a: Elements, Mixtures & Compounds

Transcription

1 AP WORKSHEET 01a: Elements, Mixtures & Compounds 1. Classify each of the following as either, an element, a compound or a mixture. If you classify something as a mixture then also state whether it is a homogeneous or a heterogeneous mixture. (10) (a) Helium (b) Nitrogen (c) Pure water (d) Pure table salt (sodium chloride) (e) Flat (un-carbonated) Coca-Cola (f) Air (g) Fruit cake 2. If three, different samples of a particular compound are found to have exactly the same elemental % by mass, what does that tell us about; (a) the three samples in relation to one another? (2) (b) the purity of the three samples? (2) 3. Visit the following URL: and answer the question found there. (11) 1

2 4. Match the following particulate diagrams to the correct description. In each case a black circle and a red circle represent the atoms of different elements. (4) A. A single, pure, monatomic element. B. A mixture of two elements. C. A single, pure compound. D. A mixture of two compounds. E. A mixture of an element and a compound. 2

3 AP WORKSHEET 01b: Empirical Formula 1. A common oxide of nitrogen contains 25.93% N. Deduce the empirical formula of the oxide. (2) 2. A compound that is usually used as a fertilizer can also be used as a powerful explosive. The compound has the composition 35.00% nitrogen, 59.96% oxygen and the remainder being hydrogen. What is its empirical formula? Given it is ionic, suggest a name for the compound. (3) 3. What are the empirical formulae for these compounds, that both contain five carbon atoms? (2) (a) C 5 H 10 (b) C 5 H A substance has an empirical formula of CH 2 Br and a molar mass of 188 g mol -1. What is the molecular formula of the compound? (1) 1

4 5. The common pain medicine, Advil, contains the active ingredient Ibuprofen that has a molar mass of 206 g mol -1. Ibuprofen contains 75.73% C, 8.74% H, the remainder being oxygen. What are the empirical and molecular formulae for Ibuprofen? (4) 6. The molar mass of the common antibiotic oxytetracycline is found to be 460 g mol -1 and a g sample contains g of carbon, g of H, g of oxygen and the remainder being nitrogen. What is the molecular formula of the oxytetracycline? (4) 2

5 AP WORKSHEET 01c: Isotopes and Mass Spectrometry 1. Many elements have a number of isotopes. (a) Define the term isotope. (2) (b) Complete the following table. (22) Row Isotope symbol Atomic # # Protons # Neutrons Mass # 1 13 C (c) Consider the 2 nd and 4 th row in the table. What three things do they have in common? (3) (d) Consider the 2 nd and 4 th row in the table. Give two differences? (2) (e) Naturally occurring Ni is found to have the following approximate isotopic abundance; 58 Ni 68%, 60 Ni 26%, 62 Ni 4.0% and 61 Ni 2.0% Calculate the average relative atomic mass of Ni to two decimal places. (2) 1

6 2. The results taken from a mass spectrum of chlorine gas show peaks at m/z and m/z (The m/z peaks on a mass spectrum identify the different isotopes of an element that are present in the sample). (a) Given that the relative abundances of Cl and Cl are 77.50% and 22.50% respectively, calculate the average relative atomic mass of chlorine atoms to four significant figures. (2) (b) Suggest all the possible masses of CI 2 molecules that are made when two chlorine atoms bond together. (3) (c) Which of the molecules you have suggested in (b) will be the most abundant? Explain your answer. (2) 2

7 3. Naturally occurring bromine molecules, Br 2 have masses of 158, 160 and 162. They occur in the relative abundances 25.69%, 49.99% and 24.31% respectively. What is the average atomic mass of bromine atoms? What is the relative abundance of 79 Br and 81 Br isotopes? (4) 4. An unknown element Z is analyzed in a mass spectrometer and is found to have the following isotopes with the corresponding relative abundances. Isotope Z 50 Z 52 Z 53 Z 54 Relative abundance (a) Using the axis below, sketch the expected mass spectrum that these data would provide. Label the axes and pay attention to the size of any lines that you draw. (4) (b) Calculate the average atomic mass of Z and identify the element. (3) 3

8 5. Consider the following mass spectrum that was generated from the analysis of an element. 100 Relative Abundance (a) What does the existence of only a single peak in the spectrum suggest about the element? m/z (b) Identify the element. 6. Copper has an atomic mass of amu and has two stable isotopes. Copper-63 has a mass of amu, and copper-65 has a mass of amu. (a) Calculate the percent abundance of each isotope of copper. (b) Sketch the expected mass spectrum of the copper. 4

9 AP WORKSHEET 01d: Quantitative aspects of electrons 1. This question is about breaking covalent bonds. (a) The bond energies of the single bonds between two chlorine atoms within a chlorine molecule and two fluorine atoms within a fluorine molecule are calculated to be 4.02 x J and 2.64 x J respectively. For each bond, calculate the following; (i) The frequency of a photon that could be used to break the bond. (2) (ii) The wavelength of each photon in (i). (2) (b) When an excited electron falls back to its ground state, what can be said of the energy change that occurs when compared to the energy change of the original, promotion process? Explain. (2) 2. Lithium ions give a distinctive red flame test. In one such experiment the energy of this red light is found to have an energy of 3.06 x J. Calculate the wavelength of the light from the lithium ions in nm. (2) 3. Which of the following process will release the greatest amount of energy? Explain your answer. (2) Promoting an electron from n = 1 to n = 6 Promoting an electron from n = 1 to n = 4 An electron falling from n = 6 to n = 2 An electron falling from n = 6 to n = 5 4. When an electron falls from n = 5 to its lowest possible state in the Lyman series, the energy that is released is greater than the energy that is released when an electron falls from n = 5 to its lowest possible state in the Balmer series. Explain. (2) 5. Electron transitions are expected to absorb or emit greater magnitudes of energy in the He + ion than in the hydrogen atom. Why? (2) 1

10 AP WORKSHEET 01e: Orbital filling rules The rules that you have been applying in order to determine the electronic configuration of an atom are summarized below. A. Lowest energy orbitals are filled first. THE AUFBAU PRINCIPLE. B. Orbitals can only contain a maximum of two electrons and when two electrons enter the same orbital they must have opposite spins (+ ½ or ½) so that each electron has a unique set of quantum numbers. (In the electrons in boxes diagram they must be drawn NOT OR ). THE PAULI EXCLUSION PRINCIPLE. C. When orbitals of identical energy (degenerate) are available electrons enter these orbitals singly before any spin pairing takes place. HUNDS RULE. Consider each of the elements listed and the INCORRECT electronic configuration associated with each one. In each case identify which of the above rules or principles (A, B or C) is violated and insert the correct electronic configuration (in a similar format to that of the incorrect configuration). Then add a possible set of quantum numbers for the outer most electron. An example is completed for you. (34) Page 1 of 2

11 INCORRECT CORRECT ELEMENT VIOLATION CONFIGURATION CONFIGURATION N 1s 2 2s 2 2px 2 2py 1 C 1s 2 2s 2 2px 1 2py 1 2pz 1 Al 1s 2 2s 2 2p 6 3p 3 B 1s 2 2s 3 P 1s 2 2s 2 2p 6 3p 5 Mg [Ne] C 1s 2 2s 1 2px 1 2py 1 2pz 1 C 1s 2 2s 2 2px 2 Mn [Ar] 4s 1 3d 6 Ni [Ar] 4s 2 3d xy 2 3d xz 2 3d yz 2 3d z2 2 3d x2-y2 0 Cl [Ne] Sc [Ar] 3d 3 B 1s 2 2s 1 2px 1 2py 1 Na 1s 1 2s 2 2p 6 3s 2 S [Ne] 3s 2 3px 2 3py 2 V [Ar] 3d 5 P [Ne] 3s 2 3px 2 3py 1 Kr [Ar] 4s 2 3d 16 Page 2 of 2

12 AP WORKSHEET 01f: Electronic Configuration Summary 1. Give full and abbreviated (noble gas core method) electronic configurations for the following. (8) (a) Br FULL NOBLE GAS CORE (b) Cr FULL NOBLE GAS CORE (c) Fe FULL NOBLE GAS CORE (d) S 2- FULL NOBLE GAS CORE 2. For each of the following sets of orbitals, indicate which orbital is higher in energy. (4) (a) 1s, 2s (b) 2p, 3p (c) 4s, 3d yz (d) 3p x, 3p y, 3p z 3. Indicate the block (s, p or d) in which each of the following elements found. (5) BLOCK (a) Sc (b) P 1

13 (c) Fr (d) Ni (e) As 4. An atom has two electrons with principal quantum number (n) = 1, eight electrons with principal quantum number (n) = 2 and seven electrons with principal quantum number (n) = 3. From these data, supply the following values (if insufficient information is given, say so). (a) The mass number. (2) (b) The atomic number. (1) (c) The electron configuration. (2) 5. Identify the element from the electron configurations of atoms shown below. (3) (a) [Ne] 3s 2 3p 2 (b) [Ar] 4s 2 3d 7 (c) [Xe] 6s 2 6. Give the symbol of the atom or ion represented by the following sets of atomic numbers and electronic configurations. (4) Atomic # Electronic Configuration Symbol of Atom or Ion (a) 8 1s 2 2s 2 2p 4 (b) 11 1s 2 2s 2 2p 6 (c) 14 1s 2 2s 2 2p 6 3s 2 3p 2 (d) 22 1s 2 2s 2 2p 6 3s 2 3p 6 3d 2 2

14 7. Give the electron configurations for the following transition metal ions. (3) (a) Sc 3+ (b) Cr 2+ (c) Ni Consider the element Scandium, atomic # 21. (a) If the electronic configuration of the element were constructed "from scratch", into which orbital (and into which shell) would the final electron be placed? (1) (b) When scandium forms an ion with a charge of +1, from which orbital (and from which shell) would the electron be removed? (1) 9. Of the following species (Sc, Ca 2+, Cl, S 2-, Ti 3+ ), which are isoelectronic? (1) 10. Identify the element that is composed of atoms where the last electron; (5) (a) Enters and fills the 4s sub-shell (b) Enters but does not fill the 4s sub-shell (c) Is the first to enter the 2p sub-shell (d) Is the penultimate to enter the 4p sub-shell (e) Is the second to enter the 4d sub-shell 3

15 11. Write the full electronic configuration for argon. (1) 12. Identify two positive and two negative ions that are isoelectronic with argon. (4) (a) Two Positive ions (b) Two Negative ions 13. Using the electrons in boxes notation complete the electronic configurations of the following elements. (3) 1s 2s 2p 3s 3p 3d 4s 4p Element V 1s 2s 2p 3s 3p 3d 4s 4p Element Ar 1s 2s 2p 3s 3p 3d 4s 4p Element Zn 14. State the number of unpaired electrons in each of the electronic configurations in question 13. (3) # of unpaired electrons (a) V (b) Ar 4

16 (c) Zn 15. How would you expect the magnitude of the energy released in the process, 4th shell 1st shell transition, to vary for a He + ion compared to a Li 2+ ion? Explain your answer. (2) 16. Identify the following atoms as either paramagnetic or diamagnetic. (3) (a) Ga (b) Cr (c) Ni 5

17 AP WORKSHEET 01g: Photoelectron Spectroscopy 1. Consider the simulated PES plot shown below, that is produced by the analysis of the atoms of a single element. All peaks in the PES are shown. Relative number of electrons Energy Increasing (a) Using the plot, suggest the electron configuration of the element and hence identify the element. (2) (b) Which two peaks are likely to represent electrons that are most likely to be removed when these atoms form ions? Explain. (3) (c) Using your answer in (b), identify the mostly likely charge on an ion of this element. Explain (2) (d) Suggest a reason for the huge jump in energy between the peak at 12.1 and the peak at 150. (2) (e) Suggest a reason for the x-axis being labeled with increasing values from right to left. 1

18 2. Consider the simulated PES plot shown below, that is produced by the analysis of the atoms of a single element. All peaks in the PES are shown. Relative number of electrons Energy Increasing (a) Write the electron configuration and identify the element. (2) (b) The plot is divided into three separate areas on the x-axis. Why is the axis divided in this manner? (2) (c) What would be the charge on an ion formed from this atom? Justify your answer. (2) (d) What is the significance of three of the peaks having the same height? (2) (e) The peaks at 1.25 & 2.44, as well as the peaks at 20.2 & 26.8, are relatively close to one another but have different energies? Explain why they are of the same magnitude but slightly different. (2) 2

19 3. Consider a PES plot for carbon atoms. (a) How many peaks would you expect in the PES for carbon? Explain. (2) (b) What would be the relative heights of the peaks that you have identified in (a)? Explain your answer carefully. (2) (c) How would you expect the height of the 2p peak in carbon s PES to compare to the height of the 2p peak in nitrogen s PES? Explain. (2) 3

20 AP WORKSHEET 01h: Ionization Energy 1. Explain each of the following observations. (a) Sodium has a lower first-ionization energy than lithium. (2) (b) Oxygen has a lower first-ionization energy than nitrogen. (2) 2. Consider the ionization energy plot shown below. Explain each of the following. (a) There is a general increase in the first ionization energy from sodium to argon. (2) (b) Boron has a lower first ionization energy than beryllium. (2) 1

21 2. continued: (c) The first ionization energy of neon (atomic number 10) is significantly higher than that of argon (atomic number 18) but significantly lower than the first ionization energy of helium (atomic number 2), despite all three elements being in the same group. (2) (d) Helium has the highest first ionization of all the elements shown. (2) 2

22 3. Consider the ionization energies of elements X and Y shown below in kjmol -1. X and Y are in the same period of the periodic table and are adjacent to one another in the table. 1 st 2 nd 3 rd 4 th 5 th 6 th 7 th 8 th 9 th X Y (a) In which group would one find element X? Explain. (2) (b) Does element X lie to the right or the left of element Y in the periodic table? Explain. (2) (c) Which is the first period on the periodic table that these elements could be in? Explain. (2) (d) Why are the second ionization energies of both elements larger than their respective first ionization energies? (2) 3

23 (e) It is found that Y has the largest first ionization energy in the period that it is found. What does this tell us about Y? (2) (f) It is found that element Q, which is in the same period as X and Y but lies to the left of element X in the periodic table, only has values for its first four ionization energies. Suggest a reason for this observation. (2) 4. (a) Define first ionization. (2) (b) Write an equation to show the second ionization energy of calcium. (2) 4

24 AP WORKSHEET 01i: Atomic and Ionic Size 1. In each of the following pairs, pick the larger species. Explain you answer in each case. (6) (a) Cu and Cu 2+ (b) F and F - (c) Na and K 2. Identify and explain the trend in atomic size for the following transitions in the periodic table. (4) (a) Moving vertically from Ar to He (b) Moving horizontally from Na to Ar 3. Only one of the following statements is correct. Which one? (1) (a) All cations are larger than their corresponding atoms (b) All anions are smaller than their corresponding atoms (c) Atomic size increases on transitioning from left to right across period 2 of the periodic table (d) The most common ion of chlorine is smaller than a chlorine atom (e) The most common ion of strontium is larger than a strontium atom (f) The most common potassium ion is larger than the most common sodium ion (g) The ions most commonly formed by group 16 elements are smaller than their corresponding atoms 1

25 4. Consider the plot below that shows atomic and ionic radii of the most commonly formed ion (in units of pm) for selected elements, plotted against atomic number. (a) Which color represents the plot for atomic radii? Explain your answer by using any element as an example. (2) (b) What do the elements that have smaller ionic radii than their corresponding atomic radii have in common? (2) (c) Suggest a reason for the absence of comparative atomic and ionic radii data for elements with atomic numbers of 2, 10 and 18. (2) (d) Identify the element with atomic number 19, identify the formula of the ion that it commonly forms, and convert the radii of both the atom and the ion to units of cm. (2) (e) What common feature can be identified for all of the non-metals on the plot? (2) (f) What accounts for the sharp increase in height of the blue lines that occurs at elements with atomic numbers 3, 11 and 19 respectively? (2) 2

26 (g) Make a prediction about the relative heights of the blue line and red line if data were added to the plot for the element with an atomic number of 15. Explain. (2) (h) The element with atomic number 1 has a red line that is significantly taller than its blue line. Under what circumstance would the red line be shorter than the blue line for this element? (2) (i) If data were added to the plot for the element with atomic number 7, which would be taller, the blue or the red line? Explain. (2) 3

27 AP WORKSHEET 01j: Periodicity Summary 1. Complete the table. (5) Element Charge on most common ion Rb Cs Ga At Se 2. Define Ionization Energy. (2) 3. Using the metal magnesium as an example, write two separate equations to show the first and second ionization energy of magnesium. (Remember state symbols are important as they from part of the definition). (4) First Ionization Second Ionization 4. Which of the following elements (one from each pair) would you expect to have the highest first ionization energy? Explain your answers. (4) Ca or Be Na or Ar 1

28 5. Consider the table of the first four ionization energies for element A shown below. Ionization Energy in kj mol -1 1 st 2 nd 3 rd 4 th (a) In which group does A appear on the periodic table? (1) (b) Predict the formula of the compound that A forms with fluorine. (1) (c) What is the minimum number of electrons that A must have? (1) 6. Arrange the following species in order of increasing size. Rb +, Y 3+, Br -, Kr, Sr 2+ and Se 2-. (1) SMALLEST LARGEST 7. Are there any atoms for which the second ionization energy is greater than the first? Explain your answer. (2) 8. Is it possible for two different atoms to be isoelectronic? If so give examples. (2) 9. Is it possible for two different anions to be isoelectronic? If so give examples. (2) 10. Define electron affinity. (2) 11. Write an equation to summarize the process of second electron affinity of oxygen. (Remember state symbols are important as they from part of the definition). (2) 2

29 12. Consider the table of ionization energies for element X shown below. Ionization Energy in kj/mol 1 st 2 nd 3 rd 4 th 5 th 6 th (a) In which group will X be found? (1) (b) Explain your answer to 12(a). (2) (c) Predict the formula of X s bromide. (1) 13. Explain carefully why rubidium tends only to form a +1 ion? (2) 14. Explain carefully why elements in the same group react in similar ways? (1) 15. Use the data below in order to predict the boiling point of Radon. (1) Noble Gas Boiling Point in K Helium 4.21 Neon 27.1 Argon 87.3 Krypton 120. Xenon 165 3

30 16. How would expect the sizes of the hydrogen ion and the hydride ion to compare with that of the hydrogen atom? (3) 17. How would expect the sizes of the hydrogen ion and the hydride ion to compare with that of the helium atom? (3) 18. Identify any isoelectronic species in the following list; Fe 2+, Sc 3+, Ca 2+, F -, Co 2+, Co 3+, Sr 2+, Cu +, Zn 2+ and Al 3+. (4) 19. Is it possible for two different cations to be isoelectronic? If so give examples. (2) 20. Arrange the following atoms into order of increasing first ionization energy. Sr, Cs, S, F and As. (1) LOWEST HIGHEST 21. Write equations to show the reaction of potassium with water and calcium with water. (2) 22. Write equations to illustrate the following. (3) (a) The basic nature of rubidium oxide. (b) The acidic nature of an oxide of phosphorus. (c) The reaction of calcium oxide with water. 23. What do you understand by the term, shielding? (2) 4

31 AP WORKSHEET 01k: Moles and Avogadro s Number Question 1 A sample of Ge is found to contain 9.70 x atoms of Ge. (a) How many moles of Ge atoms are in the sample? (1) (b) What is the mass of the sample? (1) Question 2 (a) How many W atoms are found in 0.43 moles of pure W? (1) (b) What is the mass of the W? Question 3 (a) What is the mass, in grams, of moles of Sn? (1) (b) How many Sn atoms are in the sample? (1) Page 1 of 3

32 Question 4 (a) How many moles of Ca are in 2.03 g of Ca? (1) (b) How many Ca atoms in the sample? Question 5 (a) 5.00 moles of a binary, group 2 oxide are found to have a mass of 521 g. Identify the group 2 metal. (2) (b) How many ions are in the sample? (2) Question 6 (a) How many Ta atoms are found in a g sample of Ta? (2) (b) How many moles of Ta atoms are in the sample? (1) Page 2 of 3

33 Question 7 (a) What is the mass of 8.11 x atoms of Sulfur? (2) (b) How many moles of S are in the sample? (1) Question 8 What mass of Cu atoms have the same number of atoms as there are in a 4.21g sample of Si? (2) Page 3 of 3

34 AP WORKSHEET 01l: Titrations 1. Limestone can be considered to be essentially pure calcium carbonate. Acid rain can be considered to be sulfuric acid. When limestone buildings are eroded by acid rain the reaction can be summarized by the chemical equation below. CaCO 3(s) + H 2 SO 4(aq) CaSO 4(aq) + CO 2(g) + H 2 O (l) (a) This reaction can be replicated in the lab by performing a titration. If 7.21 g of CaCO 3 solid is added to 1.20 L of mol L -1 sulfuric acid, how many moles of acid will remain after reaction? (2) (b) What volume of mol L -1 KOH must be added to neutralize this excess acid? (2) 2KOH (aq) + H 2 SO 4(aq) K 2 SO 4(aq) + 2H 2 O (l) 1

35 2. In a very similar reaction to the one carried out in question #1, some pure magnesium carbonate was added to 145. ml of 1.00 mol L -1 HCl. When the reaction had finished, the solution was acidic ml of mol L -1 Na 2 CO 3 solution was required to neutralize the excess acid. What mass of magnesium carbonate was originally used? (4) MgCO 3(s) + 2HCl (aq) MgCl 2(aq) + CO 2(g) + H 2 O (l) Na 2 CO 3(s) + 2HCl (aq) 2NaCl (aq) + CO 2(g) + H 2 O (l) 2

36 g of a mixture of sodium carbonate and potassium bromide was dissolved and made up to the mark in a 250. ml volumetric flask ml portions of this solution were neutralized by an average of 24.7 ml of mol L -1 HCl. Find the percentage of sodium carbonate in the original 5.00 g sample of the mixture. (4) (Potassium bromide does not react with HCl) Na 2 CO 3(aq) + 2HCl (aq) 2NaCl (aq) + H 2 O (l) + CO 2(g) 3

37 AP WORKSHEET 01n: Mass Spectrometry 1. How many peaks would be observed in the mass spectrum of O 2 +, given that there are three, commonly occurring isotopes of O with mass numbers of 16, 17 and 18? (2) 2. Bromine has two isotopes, Br 79 and Br 81. The isotopes occur in a 50:50 (1:1) ratio. Given that the mass spectrum of bromine contains peaks for both bromine atoms and diatomic bromine molecules, predict the number of peaks in the spectrum. What would be the relative height of the atomic peaks? What would be the relative height of the molecular peaks? Assume that z = +1 in each case. (4) 3. A sample of carbon atoms were chemically combined with a sample of oxygen atoms to yield the compound carbon dioxide. The sample of oxygen atoms was artificially manufactured to have a 3:1 ratio of O 16 and O 18, Assuming carbon to have only a single isotope C 12, predict the mass spectrum that the resulting sample of carbon dioxide would produce in terms of the number of peaks. Which peak would be the smallest? Explain. Assume the mass spectrum only includes peaks for the compound carbon dioxide and that z = +1 in each case. (4) 4. Neon atoms are known to produce a mass spectrum that consists of three peaks at m/z values of 20, 21 and 22, with the relative abundance of the peaks found to be in the ratio 112:0.21:11.1. Assuming the z value of the species causing the peaks to be +1 in each case, calculate the average atomic mass of neon based on these data. (2) 1