Atomic Structure Homework Packet Honors Chemistry Focus Learning Targets Atomic Structure and Quantum Chemistry 1. Give the one main contribution to the development of the atomic model from each of the following scientists: Dalton, Thomson, Rutherford, Chadwick, and Bohr. 2. Identify elements by both name and chemical symbol using a periodic table. 3. Compare protons, electrons, and neutrons in terms of charge, mass, and location in an atom. 4. Use the periodic table to determine the number of protons, electrons, neutrons, and atomic mass for a given element. 5. Define isotope and state how the atomic structure for isotopes of the same element are similar and different. 6. Calculate the average atomic mass from the relative abundances and masses of each isotope. 7. Define valence electrons and determine the number of valence electrons for an atom. 8. Locate rows/periods and groups/families on the periodic table. 9. Draw the Bohr diagram for an atom showing protons and neutrons and the number of electrons in each shell. Draw the Bohr diagram for the ion of an element, showing how the atom establishes a full valence shell. Determine the noble gas that the atom resembles once it forms an ion. 10. Define ion and determine the charge for the ion of an element from the periodic table. Determine the number of electrons for an ion. Give the symbol for the ion. 11. Define cation and anion. 12. Define and give examples of electromagnetic radiation. 13. Define wavelength and frequency and state the units used to measure each quantity. 14. Perform calculations involving wavelength, frequency, and energy, giving answers with the appropriate units and significant figures. 15. Describe the experiment used to show the photoelectric effect and the significance of the findings by defining a photon 16. List the four different kinds of atomic orbitals by their letter designation and state the number of electrons that each can hold. 17. Give orbital notation for a given atom/ion by applying the Aufbau Principle, Hund s Rule, and the Pauli Exclusion Principle. 18. Write electronic configuration for a given atom/ion. 19. Write noble gas configuration for a given atom/ion. 20. Recognize an excited state for a given element. 21. Give two elements that are exceptions to the regular electron configuration rules and write the actual configuration for these elements.
Name: Period: Date: WS#1 The Atomic model of atom Complete the sentences using words from the word bank More than thousand years ago, A Greek philosopher named led a group of scientists now known as atomists. These early Greeks thought that the atom was the possible piece of matter that could be obtained. They guessed that the atom was a small, particle, and that all the atoms were made of the same material. They also thought that different atoms were different shapes and sizes, that the atoms were infinite in number, always and capable of joining together. In 1803, proposed an atomic theory. The theory stated that all elements were made of atoms and that the atoms were and indestructible particles. Dalton s theory also said that atoms of the same element were the same, while atoms of different elements were. The theory also said that were made by joining the atoms of two or more elements together. In 1897, J.J. Thomson discovered a particle even smaller than the atom. He named it the, but today we call it the. As a result of his discovery, Thomson proposed a new atomic. According to Thomson s model, the atom was like a plum it was mostly a thick, positively charged material, with negative electrons scattered about it like in a pudding. In 1908, Ernest Rutherford took an extremely thin sheet of and bombarded it with electrons. Much to his surprise, most of the electrons went right through the foil, and the occasional was seriously deflected. To him, this seemed as likely as a baseball going through a brick wall. He theorized that the gold foil must be mostly empty, or else the electrons would bounce off most of the time. He figured that the atom was made up of a small, dense, positively charged center, called the. In1913, Neil Bohr narrowed down the actual location of the electrons. Bohr s model was similar to Rutherford s in that it had a made up of positively charged material. Bohr went on to propose that the negative particles ( ) orbited the nucleus much like the planets the sun. Today the atomic model is very similar to Bohr s model. The modern atomic model, the Wave model, does not have exact orbits like Bohr did. Instead the modern model has a scatter region surrounding the where an will probably, but not certainly, be found.
On the graphic organizer, summarize and illustrate the features of the atom for each model state the problem with each model. Democritus Model- 2000 yrs ago Features: Dalton Model - 1803 Features: Thomson Model - 1897 Features: Diagram: Diagram: Diagram: Problem: Problem: Problem: Rutherford Model - 1908 Features: Bohr Model - 1913 Features: Wave Model - Modern Features: Diagram: Diagram: Diagram: Problem: Problem: Problem:
Atomic Structure 1. Complete the following table Sub atomic Particles Relative charge Relative mass Location in atom Proton Electron Neutron 2. Label the parts of an atom on the diagram below. a. What type of charge does a proton have? b. What type of charge does a neutron have? c. What type of charge does an electron have? WS#2 a. Which two sub atomic particles are located in the nucleus? 3. What is the atomic number and mass number in the diagram above? 4. Which of the following statements are true a) Protons have about the same mass as neutrons b) Protons have about the same mass as electrons c) Protons have twice the mass as neutrons d) Protons have the same magnitude of charge as electrons but are opposite in sign 5. The atomic number tells you the number of in one atom of an element. It also tells you the number of in a neutral atom of that element. The atomic number gives the identity of an element as well as its location on the Periodic Table. No two different elements will have the atomic number. 6. Name the element which has the following numbers of particles. Be specific. (Include charges and mass numbers where possible.) a. 26 electrons, 29 neutrons, 26 protons b. 53 protons, 74 neutrons 7. What is the mass number of an atom with 3 protons, 3 electrons and 4 neutrons? 8. Given the elements name and its mass number give the complete isotopic symbol and the number of neutrons for the following: Lithium-6 Iron-58
9. Complete the table. There is enough information given for each element to determine all missing numbers. mass number (p + n) atomic number (p) A Z X The Atomic number is equal to the number of protons In a neutral atom the number of protons = number of electrons Element Symbol Atomic Number Mass Number Protons Neutrons Electrons Carbon 6 14 O 8 10 Potassium 19 20 19 41 197 79 Αυ Tin Sn 50 68 Zinc 64 30 66 30 68 30 Cobalt Co 27 32 Boron 5 6 10 5 56 Fe 26 26 28
What is an ion? An ion is a charged particle. WS#3 Determining the number of electrons- The number of electrons in an element can change. For a neutral atom, the number of protons is exactly equal to the number of electrons. So the number of electrons is the same as the atomic number. However, it is possible to remove electrons and not change the identity of an element. These are called ions. The charge on the ion tells you the number of electrons. If the charge is positive, subtract that number from the atomic number to get the number of electrons. You have more protons. If the charge is negative, add the amount of charge to the atomic number to get the number of electrons. You have more electrons. Cations are because they electrons or become more Positive. Anions because they gain electrons or become more negative. Fill out the missing data below: Question Ion Mass # Atomic # 1 41 19K 1+ # of Nucleons # of p # of n # of e- Lost or Gained Electrons 2 80 35Br 1-3 25 12Mg 2+ 4 22 25 19 5 138 La 3+ 6 127 53 54 18 7 15 16 8 32 S 2-9 34 45 36 135 10 55Cs 1+ 11 12 13 43 36 55 92 41 36 6 10 5 14 113 49In 3+ 15 56 Fe 2+
Atomic Structure- Bohr s model Atoms can be represented by Bohr diagrams. Bohr diagrams are useful for (i) better understanding the properties of an element, and for (ii) predicting how an atom can combine with others. In this model of the atom, the electrons travel around the nucleus in well-defined circular paths known as or. Steps for Drawing Bohr Diagrams (for use only with the first 20 elements): a. Determine the number of electrons to be drawn:# electrons = b. Draw a circle for the nucleus and write the inside c. Draw shells/levels around the nucleus. The of the element in the Periodic Table is the number of levels. d. Fill the levels with electrons according to the following pattern: e. 1 st level(nearest to the nucleus): filled first, with a maximum of electrons; f. 2 nd level: filled with a maximum of electrons; g. 3 rd level: filled with a maximum of electrons, 1. Draw the Bohrs diagram for the following elements a. Magnesium b. Chlorine c. Nitrogen d. Calcium e. Silicon f. Flourine g. Argon 2. Identify the elements whose Bohr model diagrams are represented below. Write the names of the elements in the space provided (a) (b) (c) (d) (e) (f) a) b) c) d) e) f)
Average Atomic Mass and abundance of Isotopes WS#4 The atomic mass for each element is reported on the periodic table. This number is a weighted average of the masses of each of the isotopes of an element. For example, the atomic mass of carbon is reported as 12.011 amu. Carbon is composed primarily of two isotopes: carbon-12 and carbon-13. The atomic mass is calculated from the relative abundance and the masses for these two isotopes. Using the equation below we can calculate the atomic mass for carbon. Atomic Mass = % isotope 1 mass isotope 1 + % isotope 2 mass isotope 2 +.. Sample problem: Carbon-12 makes up 98.93% of all of the carbon atoms, while carbon-13 is about 1.07% abundant. Since the carbon-12 isotope is more abundant, its mass is weighted more in the calculation of carbon s atomic mass. An example calculation is done below. Isotope % Abundance Mass Carbon-12 98.93% 12.000 amu Carbon-13 1.07% 1 3.003 amu 1. Rubidium is a soft, silvery-white metal that has two common isotopes, 85 Rb and 87 Rb. If the abundance of 85 Rb is 72.2% and the abundance of 87 Rb is 27.8%, what is the average atomic mass of rubidium? ( Ans: 85.56 amu) 2. Uranium is used in nuclear reactors and is a rare element on earth. Uranium has three common isotopes. If the abundance of 234 U is 0.01%, the abundance of 235 U is 0.71%, and the abundance of 238 U is 99.28%, what is the average atomic mass of uranium? (Ans: 237.98 amu) 3. The four isotopes of lead are shown below, each with its percent by mass abundance and the composition of its nucleus. Using the following data, first calculate the mass number of each isotope. Then using that as your mass, calculate the average atomic mass of lead. 82p 82p 82p 82p 122n 124n 125n 126n 1.37% 26.26% 20.82 51.55% Mass #
4. Argon has three naturally occurring isotopes: argon-36, argon-38, and argon-40. Based on argon s reported atomic mass, which isotope exist as the most abundant in nature? Explain. 5. Calculate the atomic mass of silicon. The three silicon isotopes have atomic masses and relative abundances of 27.9769 amu (92.2297%), 28.9765 amu (4.6832%) and 29.9738 amu (3.0872%). 6. Copper used in electric wires comes in two flavors (isotopes): 63 Cu and 65 Cu. 63 Cu has an atomic mass of 62.9298 amu and an abundance of 69.09%. The other isotope, 65 Cu, has an abundance of 30.91%. The average atomic mass between these two isotopes is 63.546 amu. Calculate the actual atomic mass of 65 Cu. ( Ans: 64.9278 amu) 4. Complete the table Isotope Mass (amu) Relative Abundance (%) Neon-20 19.992 90.51 Neon-21 20.994 Neon-22 9.22 Avg. Atomic Mass = Total %: 5. Silver (Atomic weight 107.868) has two naturally-occurring isotopes with isotopic weights of 106.90509 and 108.90470. What is the percentage abundance of the lighter isotope? 6. Chlorine is made up of two isotopes, Cl-35 (34.969 amu) and Cl-37 (36.966 amu). Given chlorine's atomic weight of 35.453 amu, what is the percent abundance of each isotope? 7. Nitrogen is made up of two isotopes, N-14 and N-15. Given nitrogen's atomic weight of 14.007, what is the percent abundance of each isotope?
Calculating Frequency, Wavelength and Energy WS# 5 Wavelength length of a single wave cycle (horizontal arrow double sided arrow) Frequency-# of waves that pass a point in a given amount of time Speed = wavelength x frequency Equations and constants: E = h and E = hc/ E = energy of one photon with a frequency of c = c/ = c = speed of light = 3.0 x 10 8 m/s (meters per second) h = Planck s constant = 6.63 x 10-34 J-s = wavelength in meters = frequency in Hz (waves/s or 1/s or s -1 ) 1. Draw a wave of light with high frequency Draw a wave of light with low frequency 2. If an X-ray machine emits E.M.R with a wavelength of 1.00 x 10-10 meters, what is the frequency? 3. A photon has a frequency ( ) of 2.68 x 10 6 Hz. Calculate its energy. 4. A helium laser emits light with a wavelength of 633 nm. What is the frequency of the light? (Hint: Convert nanometers [nm] to meters by multiplying by 10-9 ) 5. A very bright yellow line in the emission spectrum of sodium has a frequency of 5.10 x 10 14 s -1. Calculate the wavelength of this yellow light. 5.88 x 10-7 m 6. The blue color of the sky results from the scattering of sunlight by air molecules. The blue light has a frequency of about 7.5 x 10 14 Hz. a. Calculate the wavelength, in nm, associated with this radiation
b. Calculate the energy, in joules of a single photon associated with this frequency. 7. A certain electromagnetic wave has a wavelength of 625m. a. What is the frequency of the wave? What region of the electromagnetic spectrum is it found? 8. The laser used to read information from a compact disk has a wavelength of 780 nm. What is the wavelength of this light?(hint: Convert nanometers [nm] to meters by multiplying by 10-9 ) 9. A mercury lamp emits radiation with a wavelength of 4.36 x 10-7 m. a. What is the color of the light from the mercury lamp b. Calculate the frequency of this radiation 10. Calculate the energy (E) and wavelength ( ) of a photon of light with a frequency ( ) of 6.165 x 10 14 Hz. 11. Calculate the frequency and the energy of blue light that has a wavelength of 400 nm 12. Calculate the wavelength and energy of light that has a frequency of 1.5 x 10 15 Hz. 13. Rank these parts of the electromagnetic spectrum from lowest energy to highest Gamma, Infrared, Microwave, Radio, Visible, Ultraviolet X-ray 14. Rank these parts of the electromagnetic spectrum from lowest frequency to highest Gamma Infrared Microwave Radio Visible Ultraviolet X-ray 15. Rank these parts of the electromagnetic spectrum from shortest wavelength to longest Gamma Infrared Microwave Radio Visible Ultraviolet X-ray 16. What is the relationship between frequency and wavelength? (Direct or Inverse) 17. What is the relationship between frequency and energy? (Direct or Inverse) Answers: (2) 3.00x10 18 s -1 (3) 1.78 x 10-27 J (4) 6.00 Hz (5) 5.88 x 10-7m (6) 4.0 x 10 2 nm, 5.0 x 10-19 J (7) 4.8 x 10 5, radio (8) 3.85 x10 14 s -1 (10) Violet, 6.88 x10 14 s -1 (10) 4.1x10-19 J 4.87 x 10-7 m (11) 7.5 x 10 14 Hz, 4.97 x 10-19 J (12) 2.0 x 10-7 m, 9.95 x 10-19 J (13) Radio wave, Microwave, Infrared, visible, UV, Gamma (14) Radio, Microwave, Infrared, Visible, UV ray, Gamma (15) Radio, Microwave, X-ray, Infrared, Visible, UV (16) Wavelength and frequency are inversely Proportional. (17) As the frequency increases, energy increases.
Electron Energy Level and Electron configuration notes: Bohr s Model: WS#6 In 1913, Bohr introduced his atomic model based on the simplest atom, hydrogen (only 1 electron). Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus. Each electron has a fixed energy = an energy level. Electrons can jump from one energy level to another. Electrons can not be or exist between energy levels. A quantum of energy is the amount of energy needed to move an electron from one energy level to another energy level. The degree to which they move from level to level determines the frequency of light they give off. To move from one level to another, the electron must gain or lose the right amount of energy. The Bohr model was tested with the hydrogen element but failed to explain the energies absorbed and emitted by atoms with more than one electron. Atomic Orbitals: An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron. Region in space where there is 90% probability of finding an electron. Different atomic orbitals are denoted by letters. s-orbitals are spherically shaped. p-orbitals are dumbell shaped. The number of electrons allowed in each of the first four energy levels are shown here. A maximum of 2 electrons per orbital Electron Configuration: The ways in which electrons are arranged in various orbitals around the nuclei of atoms are called electron configurations. Three rules the Aufbau principle, the Pauli exclusion principle, and Hund s rule tell you how to find the electron configurations of atoms. Aufbau Principle: According to the Aufbau principle, electrons occupy the orbitals of lowest energy first. Pauli Exclusion Principle: According to the Pauli exclusion principle, an atomic orbital may describe at most two electrons. To occupy the same orbital, two electrons must have opposite spins; that is, the electron spins must be paired. Hund s Rule: Hund s rule states that electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible. Electron Configurations and the Periodic Table The periodic table can be used as a guide for electron configurations. The period number is the value of n. i.e n=1, n=2, n=3..n=7 Groups 1A and 2A have the s-orbital filled. Groups 3A - 8A have the p-orbital filled. Groups 3B - 2B have the d-orbital filled. The lanthanides and actinides have the f-orbital filled.
Energy Levels and Electron Configuration Worksheet: WS#6 1. Explain the difference between a Bohr orbit and a quantum mechanical orbital. 2. What is the difference between the ground state of an atom and an excited state of an atom? 3. List the four possible subshells in the quantum mechanical model, the number of orbitals in each subshell, and the maximum number of electrons that can be contained in each subshell. Subshell Orbitals Electrons s p d f 4. What is the Pauli exclusion principle? Why is it important when writing electron configurations. 5. What is Hund s rule? Why is it important when writing orbital diagrams. 6. Write full electron configurations for each of the following elements. Element # e - Long hand Electron Configuration Ar Ge Kr Sr Fr Po Pb Cf Rn In Fm
7. Draw the full orbital diagrams for each of the following elements and indicate the number of unpaired electrons in each. Element # e - Orbital notation N O Ar Ga Ca Zr 8. Write the short hand (noble gas configuration) for the following elements. Element # e - Orbital notation N Ba Zn Co I Cf Th Os 9. Write electron configurations for each of the following ions. What do all of the electron configurations have in common? F - P 3- Li + Al 3+ 9. Explain what is wrong with each of the following electron configurations and write the correct configuration based on the number of electrons. a) 1s 4 2s 4 2p 12 b) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 c) 1s 2 2p 6 3s 2 d) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 4d 10 4p 3
Review-Atomic Structure 1. Identify the main points of John Dalton s Atomic Theory. 2. Which one of his points was later proven incorrect? 3. Explain/ illustrate his Plum Pudding Model of the atom. 4. What did Ernest Rutherford contribute to the atomic model? 5. State the law of definite and multiple proportion. 6. What did Bohr contribute to the atom? 7. Illustrate Bohr s model of the atom. Draw an example of Bohr s model 8. What did Schrodinger contribute to the current atom? 9. Which two subatomic particles contribute charge? 10. Which two subatomic particles contribute mass? 11. Do these subatomic particles contribute mass equally? 12. What is an atom s atomic number? 13. What would happen if gold (Au) lost an electron? 14. What is an atom s mass number? 15. Complete the table below these neutral atoms. Element # of protons # of neutrons # of electrons Fe-58 Zn 38 Ca-42 20 Mn 25 16. What is an isotope? 17. For each of the following ions, indicate the total number of protons and electrons in the ion. Ion Protons Electrons Neutons Atomic # Cu +2 13 18. How many protons, neutrons, and electrons are present in the 91 Zr +4 ion 40 19. How many protons, neutrons, and electrons are present in the 140 Ce +3 ion? 58 20. How many protons, neutrons, and electrons are present in the 79 Se -2 ion? 34 Mass # 21. The abundance of bromine-79 (78.9183 amu) is equal to 50.69% and the abundance of bromine 81 (80.9163 amu) is equal to 49.31%. What is the average atomic mass of bromine? 35 Rb +1 37 O -2 56 138 P -3 15 31
22. The atomic weight of Thallium is 204.3833 amu. The masses for the two stable isotopes are 202.9723amu for thallium-203 and 204.9744amu for thallium-205. Calculate the percent abundance of each isotope 23. Calculate the atomic mass of an element if 60.4% of the atoms have a mass of 68.9257 amu and the rest have a mass of 70.9249 amu. Identify the element in the periodic 24. Bohr Diagrams Forming Ions In the Bohr-Rutherford model of the atom, charged electrons orbit around a nucleus. The nucleus contains and, and makes up most of the mass of the atom. Having a full valence shell makes the atom more 25. Draw the following Bohr Diagrams, and calculate the ionic charge. Sodium (Na) # protons = # electrons = Ionic Charge = Sodium Ion (Na 1+ ) # protons = # electrons = Ionic Charge = Fluorine (F) # protons = # electrons = Ionic Charge = Fluorine Ion (F 1- ) # protons = # electrons = Ionic Charge = 26. Complete the chart Element Br P Al Ca F O Ba K Se N Nearest Kr Noble gas Resulting ion Br- 27. How many electrons can occupy an s orbital? 28. How many electrons and orbitals are in third energy level? 29. Write a ground state electron configuration and the abbreviated(noble gas) configuration Uranium Phosphrous Fe 3+ Radium
Tellurium 30. Draw the orbital notation and identify the number of unpaired electrons. Nitrogen Potassium Manganese 31. Complete the following chart: Element Atomic Number of e - in each Number Energy Level O Na S K Electron Configuration 32. Which one is not valid? State which rule has been violated. a. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 4d 10 4p 5 b. 1s 2 2s 2 2p 6 3s 3 3d 5 c. [Ra] 7s 2 5f 8 Number of e- lost or gained Charge on Ion 33. Isoelectronic species are different elements (different Z s) that have the same electron configurations. Which of these are isoelectronic? (a) Li +,H -, He (b) Ca 2+, Ne, S 2-34. The electron configuration of Chromium and Copper are given below. Chromium: Expected - [Ar] 3d 4 4s 2 Actual - Why? Two half-filled orbitals have more stability and lower energy than one full and one partially full Copper: Expected - [Ar] 3d 8 4s 2 Actual - Why? A full d orbital and unfilled s orbital have more stability and lower energy than one full and one partially full 35. Which has a greater frequency red or yellow? 36. Which has greater energy, a wavelength of 674nm or 480nm? 37. Find the color of light whose frequency is 5.21 x 10 14 cycles/sec 38. How many Joules of energy are there in one photon of yellow light whose wavelength is 620nm?