UNIT #3: Electrons in Atoms/Periodic Table and Trends

Name: Period: UNIT #3: Electrons in Atoms/Periodic Table and Trends 1. ELECTRON CONFIGURATION Electrons fill the space surrounding an atom s nucleus in a very specific order following the rules listed below: a) Aufbau Principle: Each electron occupies the lowest energy orbital available. The orbitals closest to the nucleus have the lowest energy; the orbitals farthest from the nucleus have the highest energy. Order of increasing energy: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p b) Pauli Exclusion Principle: A maximum of two electrons may occupy a single orbital, but only if the electrons have opposite spins. Each electron in an atom has an associated spin, similar to the way a top spins on its axis. Like a top, an electron can spin in only one of two directions. In an orbital diagram, this is represented by an arrow up for an electron spinning in one direction, and an arrow down for an electron spinning in the opposite direction. c) Hund s Rule: Single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbitals. This is due to the fact that electrons carry like negative charges and thus, repel each other. An electron will pair up with another electron within a given sublevel (s,p,d,f) only when necessary and in doing so, adopts the opposite spin. Key Terms: 1. Principle Energy/Quantum Level: Major energy levels surrounding the nucleus of an atom. Consists of n=1, n=2, n=3, n=4, n=5, n=6, n=7 (corresponding to periods 1 through 7 on the periodic table). 2. Energy Sublevels: Within a principle energy level, electrons occupy sublevels labeled s, p, d or f according to the shape of the atom s orbital. S-orbitals are spherical in shape; p- orbitals are dumbbell shaped; d and f orbitals have varying shapes. 3. Orbitals: Within a sublevel, electrons occupy a specific number of orbitals, each of which contain up to one pair of electrons with opposite spins. The number of orbitals within a sublevel is as follows: S-sublevel: Contains one orbital which contains a maximum of 2 electrons. P-sublevel: Contains three orbitals, each of which contains a maximum of 2 electrons. Maximum number of p-sublevel electrons is six. D-sublevel: Contains five orbitals, each of which contains a maximum of 2 electrons. Maximum number of d-sublevel electrons is ten. F-sublevel: Contains seven orbitals, each of which contains a maximum of 2 electrons. Maximum number of f-sublevel electrons is fourteen. 4. Valence Electrons: Electrons occupying the outermost principle energy level.

Electron Configuration: Denotes the filling of electrons according to the rules listed above. The configurations depict the principle energy level of each electron (coefficient 1 through 7), followed by the sublevel (s,p,d,f), followed by a superscript that represents the number of electrons. NOTE: Electrons filling sublevel d drop one energy level and electrons filling sublevel f drop two energy levels. Order of filling sublevels according to aufbau principle: Period 1 atoms: 1s Period 2 atoms: 1s, 2s, 2p Period 3 atoms: 1s, 2s, 2p, 3s, 3p Period 4 atoms: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p Period 5 atoms: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p Period 6 atoms: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p Period 7 atoms: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p Ex. He: 1s 2 (2 electrons in atom) Ne: 1s 2 2s 2 2p 6 (10 electrons in atom) Ar: 1s 2 2s 2 2p 6 3s 2 3p 6 (18 electrons in atom) Kr: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 (36 electrons in atom) Xe: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 (54 electrons in atom) Rn: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 (86 electrons in atom) NOTE: In these examples, each atom (other than helium) contains 8 valence electrons. This is the stable octet that all other atoms strive to achieve. When atoms become ions, they either lose electrons (metals) or gain electrons (non-metals) to achieve a stable principle energy level similar to their closest noble gas. More examples of neutral atoms versus their corresponding ions: Be 1s 2 2s 2 neutral beryllium atom with 4 electrons Be 2+ 1s 2 beryllium ion with 2 electrons (lost 2) Na 1s 2 2s 2 2p 6 3s 1 neutral atom with 11 electrons Na + 1s 2 2s 2 2p 6 sodium ion with 10 electrons (lost 1) O 1s 2 2s 2 2p 4 neutral oxygen atom with 8 electrons O 2-1s 2 2s 2 2p 6 oxide ion with 10 electrons (gained 2) P 1s 2 2s 2 2p 6 3s 2 3p 3 neutral phosphorous atom with 15 electrons P 3-1s 2 2s 2 2p 6 3s 2 3p 6 phosphide ion with 18 electrons (gained 3) Orbital Diagrams: Denotes each orbital within a sublevel and the electrons occupying those orbitals (indicated by an up arrow or a down arrow ). Electrons fill orbitals singularly at first, then pair as necessary with an opposite spin. Ex. 2p 4 2p 2p 2p 3d 7 _ _ _ 3d 3d 3d 3d 3d 2

2. ELEMENTS AND THE PERIODIC TABLE a) An element is a pure substance that cannot be separated into simpler substances by physical or chemical means. b) Each element has a unique chemical name and symbol. The chemical symbol consists of one, two or three letters: the first letter is always capitalized and the remaining letter(s) are always lowercase. c) Seven elements occur in nature as diatomic molecules (2 atoms) because the molecules formed are more stable than the individual atoms. They are Br 2, I 2, N 2, Cl 2, H 2, O 2, F 2. Remember it as BrINClHOF. d) On earth, 91 elements are naturally occurring and their abundance in the universe varies. e) The Periodic Table organizes the elements according to increasing atomic number. 1. Elements are arranged in vertical columns called groups or families. Each group is numbered 1 through 18. 2. Groups 1, 2, 13, 14, 15, 16, 17 and 18 are often referred to as the main group, or representative elements, because they possess a wide range of chemical and physical properties. 3. Groups 3, 4, 5, 6, 7, 8, 9, 10, 11 and 12 are referred to as the transition elements. 4. Elements in the same group have similar chemical and physical properties. 5. Elements are arranged in horizontal rows called periods. Beginning with hydrogen in period 1, there are a total of 7 periods. f) Classification of Elements 1. Metals are elements that are generally shiny when smooth and clean, solid at room temperature, and good conductors of heat and electricity. Most metals are malleable (can be pounded into thin sheets) and ductile (can be drawn into wires). a) Used to transmit electrical power, ex. copper. b) Can be formed into coins, tools, fasteners and wires. c) Group 1 elements (except hydrogen) are known as the alkali metals. d) Group 2 elements are known as the alkaline earth metals. e) Both alkali and alkaline earth metals are chemically reactive, with alkali metals being the more reactive group. f). Groups 3 through 12 elements are divided into 1. transition metals-located in periods 4 through 7. 2. inner transition metals-two sets of inner transition metals, known as the lanthanide and actinide series, appear at the bottom of the periodic table and are usually offset from the numbered periods. These elements are phosphors, substances that emit light when struck by electrons. 2. Nonmetals are elements that are generally gases or brittle, dull-looking solids. They are poor conductors of heat and electricity. The only non-metal that is a liquid at room temperature is bromine. a) Group 17 elements are the halogens. These are the most reactive non-metals. b) Group 18 elements are the noble gases-extremely unreactive due to the most stable and complete electron configuration. 3. Metalloids or semimetals are elements with physical and chemical properties of both metals and nonmetals. a) Located on the right hand side of the periodic table and form a stair-step pattern between the transition metals and the nonmetals. b) Consists of B, Si, Ge, As, Sb, Te and At. 3

3. COMPOUNDS AND LAWS OF DEFINITE/MULTIPLE PROPORTIONS a) A compound is a combination of two or more different elements that are combined chemically. Much of the matter of the universe are compounds; there are approximately 10 million known compounds. Examples are water, table salt, table sugar, aspirin. b) Compounds or elements that occur alone are referred to as pure substances. Compounds or elements that occur in combination with other compounds or elements are referred to as mixtures. 1. Homogenous mixture-one that has a uniform composition throughout and always has a single phase; can be separated by physical means such as distillation (a technique used to separate mixtures based on the differences in the boiling points of the substances) or by evaporation (removing liquid component from solid component); homogenous mixtures are also referred to as solutions. Ex. salt water, sugar water, lemonade, gasoline, steel. 2. Heterogeneous mixture-one that does not have a uniform composition and in which the individual substances remain distinct; can be separated by physical means such as filtration (technique that uses a porous barrier to separate solids from liquids). Ex. sand and water, dirt, Italian salad dressing. c) Law of Definite Proportions 1. Elements making up compounds always combine in definite proportions by mass. Regardless of the amount of a given compound, it is always composed of the same elements in the same proportion by mass. d) Law of Multiple Proportions 1. When different compounds are formed by combinations of the same elements, different masses of one element combine with the same relative mass of the other element in a ratio of small whole numbers. 2. Examples: a) Water is H 2 O: 2 parts hydrogen to 1 part oxygen Hydrogen Peroxide is H 2 O 2: 2 parts hydrogen to 2 parts oxygen Both compounds are comprised of the same elements; however, H 2 O 2 differs from H 2 O in that it has twice as much oxygen. When we compare the mass of oxygen in H 2 O 2 to the mass of oxygen in H 2 O, we get the ratio 2:1. b) Methane is CH 4 ; Carbon = 12amu and Hydrogen = 4amu; C mass : H mass = 12:4 or 3:1 Ethane is C 2 H 6 ; Carbon = 24amu and Hydrogen = 6amu; C mass : H mass = 24:6 or 4:1 4. PERIODIC TABLE TRENDS a) Atomic Radius 1. The radius of an atom is one-half the distance between the nuclei of two atoms of the same element when the atoms are joined; it is comparable to the radius of a circle which is the length of a line from the center of the circle to its edge. 2. Radius decreases as you move across a period. As you move across a period, each successive element has one additional proton in its nucleus; therefore, the positive nuclear pull increases on the negative electrons surrounding the nucleus, causing the radius to decrease. 4

3. Radius increases as you move down a group. As you move down a group, each successive element has an additional energy level surrounding its nucleus and therefore, the radius increases. b) Ionic Radius 1. An ion is an atom or a bonded group of atoms that has a positive charge (due to loss of electrons) or negative charge (due to gaining electrons). 2. When atoms lose electrons to become positive ions, their radius decreases. The loss of valence electrons from the outermost energy level results in an empty valence shell and therefore, the next level down becomes the ion s outermost energy level; therefore, the radius decreases. 3. When atoms gain electrons to become negative ions, their radius increases. The addition of electron(s) to the outermost energy level results in additional repulsive forces between the like-charged electrons. This causes the electrons to move further apart and effectively, increases the ion s radius. c) Ionization Energy 1. Ionization energy is the energy required to remove an electron from a gaseous atom. It is an indication of how strongly the atom s nucleus is pulling on its electrons. A higher ionization energy value means more energy is required to remove an electron, indicating a strong nuclear pull. A lower ionization energy value means less energy is required to remove an electron, indicating a weaker nuclear pull. 2. Ionization energy increases as you move across a period. As the number of protons increases across a period, the nuclear pull increases. 3. Ionization energy decreases as you move down a group. As energy levels are added moving down a group, the valence electrons become farther removed from the nuclear pull and its effect decreases. Also, an increase in the number of electrons between the outermost energy level and the nucleus causes what is termed a shielding effect, that is, the nuclear pull is diminished due to the intervening electrons. d) Electronegativity 1. Electronegativity indicates the ability of an atom to attract electrons in a chemical bond. 2. Electronegativity increases as you move across a period. An increase in the number of protons in the nucleus of each successive atom results in a stronger nuclear pull on the atom s own electrons and on another atom s electrons in a chemical bond. 3. Electronegativity decreases as you move down a group. An increase in the distance between the nucleus and the outermost electrons results in a weaker nuclear pull on the atom s own electrons and on another atom s electrons in a chemical bond. 5

Developing the Concept of Shells, Subshells, Electron Configurations, and More PART I: Discovering how electrons are arranged in an atom 1. Describe the nature of the interaction between protons and electrons in an atom? Consider using some or all of the following terms in your description: attraction, repulsion, neutral, positive, negative, charge, distance, nucleus, force, energy, Coulomb s Law. 2. For each situation below, compare the relative energy necessary to separate positive and negative electrical charges. Compare A to B Compare A to C 3. Consider How many electrons do you see in the picture? How many protons? Which of these electrons is the easiest (requires the least amount of energy) to remove (ionize)? Justify your answer. Compare the energy required to remove the electron from 3 with the energy in 2a 2c The first ionization energy is defined as the minimum energy that must be added to a neutral atom, in the gas phase, to remove an electron from that atom. This definition can be represented by the following chemical equation: energy + A(g) A + (g) + e 4. In the ionization equation above identify which species is at lower energy, A(g) or A + (g) + e? Justify your answer. 5. Explain why energy is required (an endothermic process) to remove the electron in a neutral atom. 6. The value of the first ionization energy for hydrogen is 1312 kj mol -1. energy + H(g) H + (g) + e On the graph on the next page use a short horizontal line to indicate the energy of H(g) and a short horizontal line to indicate the energy of H + (g) + e. Be sure to consider your responses to Q4 and Q5 above.!

H(g) H + (g) + e 7. What does the difference in energy in the lines in your diagram above represent? The values for the first ionization energy for a hydrogen and helium atom are provided in the table below. Atom 1H 2He 3Li Ionization Energy (kj mol 1 ) 1312 2373 8. Based on comparisons you made in Question 2 how would you explain the difference in the values for the first ionization energy for hydrogen and helium? 9. How does your explanation account for the relative charge on hydrogen and helium and the distance of the electron(s) from the nucleus? In the energy diagram below locate (draw a horizontal line) the first ionization energy for hydrogen and the first ionization energy for helium. 10. How does the diagram illustrate the relative ease with which an electron can be removed from each atom? 11. Predict a value for the first ionization energy for lithium. Do not add your prediction to the figure just yet. Justify your prediction (look back at Question 2 if you need guidance).!!

The actual value of the first ionization energy of lithium is 520 kj mol -1. Add this value for to the figure on the previous page. 12. How would you explain the ionization energy for lithium compared to the ionization energy for helium? Compared to hydrogen? 13. Predict the relative value of the energy necessary to remove a second electron (called the second ionization energy) from lithium. Support your prediction with an explanation. 14. Based on the first ionization energies for hydrogen, helium and lithium that you represented in the figure on the previous page, what can you infer about the distance of the electrons from their respective nuclei. The first ionization energies for selected elements from the second period of the periodic table are provided in the table below. Atom 3Li 4Be 6C 7N 9F 10Ne Ionization Energy (kj mol 1 ) 520 899 1086 1302 1681 2081 15. Explain the trend in ionization energies in terms of the charge of the nucleus and the relative location of the electrons. The first ionization energy for the element sodium is given in the following table. Atom 11Na 12Mg 14Si 15P 17Cl 18Ar Ionization Energy (kj mol 1 ) 520 16. Predict the values for the first ionization energy for the other selected third period elements. Explain how you arrived at your predictions. Below is a table containing the electron energies for each of the 18 electrons in an argon atom. The graph of this data is shown. Electron Electron Energy Removed (kj mol 1 ) 1 1521 2 2666! 3 3931! 4 5751! 5 7238! 6 8781! 7 11995! 8 13842! 9 40760! 10 46186! 11 52002! 12 59653! 13 66198! 14 72918! 15 82472! 16 88576! 17 397604! 18 427065! 17. Make observations about the graph in terms of the relative energies of the electrons and their relationship to each other. 18. Based on your responses from the previous questions how many groups (levels or shells) of electrons are shown for Argon? 19. Indicate the number of electrons in each group/level that you identified?!!!!

20. On the graph below draw a horizontal line (to the right of the y-axis) that represents an average energy level for each of the groups of electrons that you identified. Label the levels 1, 2, etc. beginning from the lowest energy level. What do these lines represent? 21. How would you describe the relative energy separation of these energy levels? 22. An electron from which level requires the least amount of energy to remove? The largest amount of energy to remove? Describe the electron structure (location of the electron) of the atom. Consider using some or all of the following terms in your description; nucleus, electron, energy, distance, level, proton, shell, arrangement, attraction, repulsion, positive, negative, charge, location. PART II: Do all electrons in the same level have the same energy? One important conclusion based on the first ionization energy experimental data is that electrons in higher shells require less energy to remove. We have examined experimental data that relates the energy required to remove an electron to the shell the electron occupies. In which shell does an electron require more energy to remove, an electron in the second shell or the fourth shell? An interesting question that cannot be answered from the experimental data of the first ionization energy is Do all electrons in the same shell require the same amount of energy to remove? We CAN answer this question if we look at photoelectron spectroscopy (PES) data for the atoms. In a photoelectron spectroscopy experiment any electron can be ionized when the atom is excited. Like with the first ionization, only one electron is removed from the atom. However in a PES experiment it can be ANY electron, not just the electron that requires the least amount of energy to remove.

Examine the PES spectrum for hydrogen shown in the figure. The label on the y-axis is energy and the units are in megajoules(m J mol 1 ) 1. What does the x-axis depict? Explain. 2. What is the relationship between the photoelectron spectrum and the first ionization energy for hydrogen? Helium is next, but before looking at its photoelectron spectrum answer the following questions: 3. How many electrons does helium have in its first shell? 4. Refer back to Part I of this activity, and obtain the first ionization energy for a helium atom. Can you predict what the PES would look like if a. the same amount of energy is required to remove each of the electrons? b. different amounts of energy are required to remove each electron? Go to back to the previous figure and sketch both scenarios.

Examine the PES for helium and compare it to your prediction from the previous question. 5. Explain the relative energy of the peak(s) and the number of electrons represented by each peak in the PES for helium and for hydrogen. 6. For lithium a. How many electrons does lithium have? b. What shells (levels) do those electrons occupy? 7. Predict what you expect the PES for lithium to look like. (Note: you do not have to predict the exact energies of each electron, you can make a reasonable estimate based on the first ionization energies for lithium and helium - refer back to Part I of this activity. Look at this PES and compare it to the prediction you made in the previous question. 8. For each peak in the PES of lithium, identify the shell the electrons represented by that peak occupy. Be sure to comment about the relative energy of the peak(s) and the number of electrons for each peak for Li.) The next element in the Periodic Table is beryllium. 9. How many electrons does beryllium have and what shells do those electrons occupy? 10. For the PES for beryllium predict a. how many peaks b. the number of electron for each peak c. estimate the relative energies.

The next element in the Periodic Table is boron. 11. How many electrons does boron have and what shells do those electrons occupy? 12. For the PES for boron predict a. how many peaks b. the number of electron(s) for each peak c. estimate the relative energies Below is the PES for boron. 13. Briefly describe how to interpret the PES for boron. 14. Predict what changes in the PES you would expect to see going across period 2 of the periodic table, from carbon to neon? Look at the PES for these second period elements.

Below is the PES for the period 2 elements from boron to neon. 15. Answer the following questions after looking at the PES for hydrogen through neon. a. Would you agree or disagree with the following statement? Explain your answer. The electrons in the second shell all have the same energy. b. How many subshells are found in the second shell? c. How many subshells are found in the first shell? d. How many electrons are in each subshell in the second shell? In the first shell? e. Moving systematically from lithium to neon; i. How many electrons are in the first shell? ii. What happens to the energy required to remove an electron in the first shell moving from left to right in the second period? Support your observation with an explanation. iii. What happens to the energy of the electrons in the outer most shell? 16. Look at the PES for the elements in the third period (sodium argon) and describe your observations. Any surprises? Explain. A notation has been agreed upon for writing an electron configuration to identify the location of the shell and subshell of each electron in an atom. Shells are labeled with a number; 1, 2, 3, etc. and subshell are labeled with letters; s, p, d, and f. Every shell contains an s subshell. 17. Write the complete electron configuration for the first ten elements in the periodic table? Look at the PES for potassium, calcium and scandium. 18. Explain what happens in the PES for scandium that has not occurred in any element prior. 19. If one electron is removed from scandium, which electron (identify the shell and subshell) requires the least amount of energy to remove?

Name: Date: Period: Bohr Model Worksheet Directions Draw the Bohr Models showing all the electrons in each energy level. 1. Magnesium compounds are used in the production of uranium for nuclear reactors. Draw the Bohr model for magnesium. Niels Bohr 2. Sodium is found in salts that can be used to seed clouds to increase rainfall. Draw the Bohr model for sodium. Page 1 of 4 2004 High School Technology Initiative (HSTI) Educational Materials: The ATOM: Structure Name: Date: Period: 3. Neon is often found in lasers. Draw the Bohr model for neon. 4. Argon gas can be found in Geiger counters which are used to detect radiation. Draw the Bohr model for argon. Page 2 of 4 2004 High School Technology Initiative (HSTI) Educational Materials: The ATOM: Structure

Name: Date: Period: 5. Aluminum alloys are used in airplane construction due to their low density. Draw the Bohr model for aluminum. 6. Oxygen is often added to rocket fuel as an oxidizer. Draw the Bohr model for oxygen. Page 3 of 4 2004 High School Technology Initiative (HSTI) Educational Materials: The ATOM: Structure Name: Date: Period: 7. Lithium can be found in Mount Palomar s 200-inch telescopic mirror. Draw the Bohr model for lithium. 8. Sulfur dioxide is often used at water treatment facilities to dechlorinate water. Draw the Bohr model for sulfur. Page 4 of 4 2004 High School Technology Initiative (HSTI) Educational Materials: The ATOM: Structure

Electron Configurations Worksheet Write the complete ground state electron configurations and orbital notations for the following: # of e Element (atom) e - configuration Orbital Notations/ diagrams 1) lithium 2) oxygen 3) calcium 4) nitrogen 5) potassium 6) chlorine 7) hydrogen 8) copper 9) neon 10) phosphorous Write the abbreviated ground state electron configurations, noble gas configuration, for the following: # of electrons Element Electron Configuration 11) helium 12) nitrogen 13) chlorine 14) iron 15) zinc 16) barium 17) bromine 18) magnesium 19) fluorine 20) aluminum Page 1 of 8

Electron Configuration Elements (atoms) and Ions Write the electron configuration and orbital notations for the following Atoms and ions: Element / Ions F Atomic number # of e - Electron Configuration F 1- O O -2 Na Na 1+ Ca Ca +2 Page 2 of 8

Al 3+ Al N N 3- S 2- Cl 1- K 1+ S Br 1- Mg 2+ Page 3 of 8

Electron Configuration Practice Directions: Write and draw the electron configurations of each of the following atoms. Example: Co : 27 e - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 Co 1s 2s 2p 2p 2p 3s 3p 3p 3p 3d 3d 3d 3d 3d 4s 1. Scandium: 2. Gallium: 3. Silver: 4. Argon: 5. Nitrogen: 6. Lithium: 7. Sulfur: Page 4 of 8

Electron Position and Configuration Position: Draw the Electron Position of each of the following atoms. Example: He: 1. Li 3. O 2. C 4. Ar Directions: Draw the electron configurations of each of the following atoms. Example: F 1s 2s 2p 2p 2p 1. Chlorine: 5. Sodium: 2. Nitrogen: 6. Potassium: 3. Aluminum: 7. Sulfur: 4. Oxygen: 8. Calcium Page 5 of 8

Electron Configuration Practice In the space below, write the expanded electron configurations (ex. = 1s 2 2s 1 ) of the following elements: 1) Sodium 2) potassium 3) chlorine 4) bromine 5) oxygen In the space below, write the abbreviated electron configurations (ex. Li= [He]2s 1 ) of the following elements: 6) manganese 7) silver 8) nitrogen 9) sulfur 10) argon In the space below, write the orbital notation (arrows) of the following elements: 11) manganese 12) silver 13) nitrogen 14) sulfur 15) argon Determine what elements are denoted by the following electron configurations: 16) 1s 2 2s 2 2p 6 3s 2 3p 4 17) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 18) [Kr] 5s 2 4d 10 5p 3 19) [Xe] 6s 2 4f 14 5d 6 20) [Rn] 7s 2 5f 11 Determine which of the following electron configurations are not valid: 21) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 4d 10 4p 5 22) 1s 2 2s 2 2p 6 3s 3 3d 5 23) [Ra] 7s 2 5f 8 24) [Kr] 5s 2 4d 10 5p 5 25) [Xe] Page 6 of 8

Electrons, Valence, and Lewis Dot Structures Chem 544/545 Dr. Brielmann Name Period 1. How many electrons are present in: Helium (He) Carbon (C) Neon (Ne) Sodium (Na) Zinc (Zn) 2. How many valence electrons are present in: Helium (He) Carbon (C) Neon (Ne) Sodium (Na) Potassium (K) Fluorine (F) Chlorine Bromine 3. Draw Lewis Dot Structures for the following elements: Helium (He) Carbon (C) Neon (Ne) Sodium (Na) Ne 4. Correct the following Lewis Dot Structures: Oxygen Nitrogen Beryllium Fluorine O N Be F 5. Fill in the following table: number of electrons: Carbon Carbon anion Carbon cation C - + C number of valence electrons Lewis structure Page 7 of 8

S-C-5-3_Periodic Trends Worksheet and KEY 10. For each of the following, circle or highlight the correct element that best matches the statement on the right. Li Si S metal N P As smallest ionization energy K Ca Sc largest atomic mass S Cl Ar member of the halogen family Al Si P greatest electron affinity Ga Al Si largest atomic radius V Nb Ta largest atomic number Te I Xe member of noble gases Si Ge Sn 4 energy levels Li Be B member of alkali metals As Se Br 6 valence electrons H Li Na nonmetal Hg Tl Pb member of transition metals Na Mg Al electron distribution ending in s 2 p 1 Pb Bi Po metalloid B C N gas at room temperature Ca Sc Ti electron distribution ending in s 2 d 2 Source: http://www.gpb.org/files/pdfs/gpbclassroom/chemistry/periodictabletrendswkst.pdf

Unit 3 Note Quiz Questions Unit 3.2: Electron Configuration 1. a 5. a 2. a 6. a 3. z 7. A 8. A 4. 9. A

10. Unit 3.3: Periodic Trends 1. a 2. 3. A

4. a 5. a 6. a 7. a 8. a

9. a 10.