Unit 2: Atoms, Bonds, Ions, and Their Properties
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1 Unit 2: Atoms, Bonds, Ions, and Their Properties Name: Date: Period: SECTION 1: ENGAGE Investigation 1.1: What is an Atom? SECTION 1: EXPLORE Investigation 2.1: Introduction to the Periodic Table SECTION 3: EXPLAIN Investigation 3.1: Atomic Puzzle Investigation 3.2: Understanding an Electron s Behavior Photons Investigation 3.3: Electron Configuration Investigation 3.4: Describing Valence Electrons using Lewis Dot Structures Investigation 3.5: Understanding the formation of bonds and ions SECTION 4: ELABORATE Investigation 4.1: Replacement Reaction Investigation 5.1: Ionic Bonds Investigation 6.1: Covalent Bonds Investigation 7.1: SECTION 5: EVALUATE Investigation 6.1: Test 1
2 Investigation 1.1 What is an Atom? Atoms are the building blocks of matter. To begin our investigation, you will be defining and identifying all the parts of an atom. 1. Hypothesize what you think the following terms mean: Nucleus: Protons: Electrons: Neutrons: 2. In the space below, draw what you think an atom looks like. Please include labels for each of the above terms. 2
3 Background Information: All matter is made up of atoms. The center of the atom is the nucleus which is a cluster of protons and neutrons. The protons have a positive electric charge while the neutrons are electrically neutral. The nucleus makes up almost all of an atom's mass. Whirling at fantastic speeds around the nucleus are smaller and lighter particles called electrons which have a negative charge. 3. Using the background information redefine your terms (incorporate charge when appropriate). Nucleus: Protons: Electrons: Neutrons: Figure 1 4. Compare and contrast your drawing of an atom to the one on the right 3
4 Figure 2 6 Atomic Number = represents the number of Protons as well as the number of Electrons C CARBON 12 Chemical Symbol Chemical Name Atomic Weight = represents the number of Protons and the number of Neutrons added together 5. Using the information below (Table 1) complete table 2. Table Al Aluminum Au Gold Cu Copper Ag Silver U Uranium Element Atomic # Atomic Weight Al Au Cu Ag U Table 2 # of Protons # of Electrons 238 # of Neutrons 4
5 6. Complete the chart using the information provided in Table 3. Table 3 # of Protons # of Electrons Element Atomic # Atomic Weight # of Neutrons Na Cl Fr Pt H 1 1 5
6 2.1 Introduction to the Periodic Table Background: The periodic table is arranged in horizontal rows called periods and vertical columns called groups/ families. The periods have their atomic numbers increase from left to right, while the groups contain elements having similar chemical properties. Each group is represented by a Roman numeral and letter. Letter A represents representative metals, while letter B represents transition metals. In addition, the figure below highlights several other important factors, such as ionization energy, metallic properties, and atomic radius. Organizing the Periodic Table: The periodic table is arranged into several sections, those being Alkali metals, Alkali Earth metals, Transition metals, Metals, Metalloids, Non-metals, Halogens, and Nobel gases. Directions: Using a copy of the period table and eight different colored pencils identify each of the sections listed above. Then answer the questions that follow. 6
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9 Characteristics of Periodic Groups Alkali Metals Alkali Earth Metals Transition Metals Metals Metalloids Non Metals Halogens Noble Gases 9
10 1. Which pair of elements is in the same period? 2. Which pair of elements is in the same group/family? 3. Which element has the smallest atomic number? 4. If the atomic number of element D is 20, then what is the atomic number of element R? (Use your periodic table to answer the remaining questions: 5-10) 5. This element has the lowest atomic number of any Group 16 element. 6. This element has the most protons of any element in Group This element is in the same family as lead, and it has fewer protons than sodium. 8. This element has an atomic number that is one greater than platinum. 9. This element has an atomic number that is double that atomic number of silicon. 10. This element is in Group 1 and has a higher atomic number than chlorine, but a lower atomic number than bromine. 10
11 Investigating 3.1 ATOMIC PUZZLE Directions: Use the periodic table to determine the element that each lettered question refers to. Put the element's symbol on the line after each of the clues. When these symbols are arranged one after the other, they will form a word or phrase. When you discover this word or phrase, transfer each symbol into the appropriate lettered box. For each problem, a hint is provided for you. 1. HINT: Produced by the Pancreas? a. Has 49 positive particles b. Has 6 electrons in its third, and outermost shell S c. A greenish, radioactive mineral used in the production of electricity U d. This element's nucleus contains only 4 neutrons e. The atom with 7 electrons a b c d e 2. HINT: What your teacher will "keep" during a test. a. Lightest element in the universe b. It has 7 electrons on its 5 th, and outermost shell i c. Has 16 protons d. Has a mass of 127 e. An element used to form water f. Element with 7 neutrons g. Atomic number 39 h. The gas that our blood absorbs during respiration I. Symbol for uranium a b c d e f g h i 3. HINT: They lived in a submarine in 1966 a. Has four protons b. Element with 125 neutrons, in the same family as fluorine c. Roman numeral for "50" L d. The Actinoid element named after "Albert" a b c d 11
12 4. HINT: A typical Irish breakfast a. Has 104 neutrons b. Symbol for carbon c. Has 7 less electrons than Iron d. Has 39 protons e. Has only a filled inner shell, and half-filled second shell f. An element found in water, and hydrochloric acid g. Has 8 electrons in it's outermost, third shell h. Abbreviation for "Miss" a b c d e f g h 5. HINT: He had an element named after him a. Symbol of the element used in foil b. The element with 6 neutrons and 5 protons B c. Contains 67 more positive particles than Hydrogen d. 20th letter of the alphabet e. Letter used to symbolize a negative subatomic particle _E f. Has a mass that is 31 greater than krypton g. It has 6 electrons on its 3 rd and outermost shell h. Atom whose atomic number is 52 I. Its electron configuration is 2, 8, 18, 18, 3 a b c d e f g h i 6. HINT: Projected onto a screen by a cathode ray tube = TV TUBE a. Symbol for boron b. The element with two electrons in its 7 th shell Ra c. 10th element of the "lanthanoid" series _DY d. It's second, outermost shell has 3 electrons e. Contains a total of 92 protons f. One of the laughing gas elements g. Atomic mass is 12 h. Having no neutrons it is an explosive gas once used in blimps a b c d e f g h 12
13 Investigation 3.2 Laboratory Investigation: Understanding an Electron s Behavior In 1913, the Danish physicist Niels Bohr proposed yet another modification to the theory of atomic structure based on a curious phenomenon called line spectra. Under Bohr's theory, an electron's energy levels (also called electron shells) can be imagined as concentric circles around the nucleus. Normally, electrons exist in the ground state, meaning they occupy the lowest energy level possible (the electron shell closest to the nucleus). When an electron is excited by adding energy to an atom (for example, when it is heated), the electron will absorb energy, "jump" to a higher energy level, and spin in the higher energy level. After a short time, this electron will spontaneously "fall" back to a lower energy level, giving off a quantum of light energy. Key to Bohr's theory was the fact that the electron could only "jump" and "fall" to precise energy levels, thus emitting a limited spectrum of light. The animation linked below simulates this process in a hydrogen atom. Problem: How does the emission of spectra of various types of light appear through a spectroscope? Materials: incandescent light fluorescent light hydrogen light helium light krypton light spectroscope high voltage source Procedure: Part 1 1. Obtain a spectroscope. 2. Look through the spectroscope toward the fluorescent lights in the room. Observe the color produced. 3. Record the color(s) of the spectra in data table 1 (see example in book). 4. Draw the color(s) of the spectra in the box labeled fluorescent light. Record your observations by using colored pencils. 5. Repeat steps 2-4 using an incandescent light bulb. Record all observations. 13
14 Part 2 6. Obtain a spectroscope as you instructor inserts a vacuum tube of an element in the high voltage source. 7. Record the color(s) of the spectra in data table. 8. Draw the color(s) of the spectra in the box labeled fluorescent light. This is the bright line spectrum. Record your observations by using colored pencils. 9. Repeat steps using the other two vacuum tubes of various elements identified in the materials. Observations/Data Table: Type of Light Color of gas as Source visible without a spectroscope Drawing of the continuous and/or brightline spectrum (please used colored pencils for your drawings) Fluorescent Light bulb Incandescent light bulb Helium Hydrogen Krypton Neon Mercury 14
15 Analysis Questions: 1. White light is made up of all seven parts of visible light. Identify the seven components. 2. Describe the difference you see between the light bulbs (fluorescent/ incandescent) and the other elements that give off light (vacuum tubes). 3. Of the five different types of light spectra you examined, which ones do you think are the MOST similar? Explain. 4. Create a prediction regarding how using an emission spectra would be helpful to scientists? 5. What is the purpose of the high voltage source? 6. Explain the process of how the bright line spectrums are produced in the samples of light that were examined. 7. How might spectral analysis be useful in astronomy? Think about this question carefully before you answer. Please provide a detail explanation of this question. 15
16 Investigation 3.3 Electron Configuration How does electron configuration relate to the periodic table? The periodic table is organized into Periods (rows), Groups 1-18 (columns) and Blocks (s, p, d and f). The periodic table below shows the s, p, d and f view: The arrangement of electrons within the orbitals of an atom is known as the electron configuration. The most stable arrangement is called the ground-state electron configuration. This is the configuration where all of the electrons in an atom reside in the lowest energy orbitals possible. Keeping in mind that each orbital can accommodate a maximum of two electrons, we are able to predict the electron configurations of elements using the periodic table. Basically, the distributions of orbitals can be laid out in the following fashion (read from the bottom up): _ 4s _ 3p _ 3s _ 2p _ 2s _ 1s The bottom energy level is level 1 - it has the lowest energy. Each "_" represents an orbital. You can see that there is 1 orbital for an s subshell. There are 3 orbitals for a p subshell. Each orbital can hold 2 electrons. Therefore, the s subshell can hold 2 electrons and the p can hold 6. As a result, the first energy level can hold 2 electrons (1s = 2), and the second energy level can hold 8 electrons (2s2p = 2 + 6), etc. 16
17 Parts of an electron configuration: Energy level - a number (1, 2, 3 and so on) Sublevel (orbital) - a letter, either s, or p Number of electrons - a superscript number Analogy: The energy level is like a driveway with cars in it, the sublevels are the type of cars in parking lot, and the orbitals are how many seats are in the car. How to write an electron configuration: In a neutral atom, the number of electrons equals the number of protons of the atom. When the electrons fill the orbitals, they occupy the lowest energy orbitals that are available. For example, hydrogen is atomic number 1 (has 1 proton). The one electron that it has occupies the lowest orbital, which is 1s. To write its electron configuration, it would be 1s 1. In an orbital diagram, it would simply be a line with one up arrow in it, which represents the 1s orbital: H: 1s 1 Practice: Write the electron configuration for the following elements- 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 He Li Cl P Ca C O F B Ne 17
18 Investigation 3.4 Describing Valence Electrons using Lewis Dot Structures Background: American chemist Gilbert Lewis developed another method for depicting an atom s electron configuration. Lewis focused on the valence electrons in his representations of elements. Valence electrons are all of the electrons in the outer orbital/shell of an element. 1. Based upon the above information how many valence electrons do the following elements have: a. B b. Mg In a Lewis dot structure the valence electrons are represented as dots placed around the element symbol. 2. Draw what you think the Lewis dot structure looks like for: a. B b. Mg Check your predictions by sharing with the class. 18
19 Draw the Lewis-Dot Structure for the following atoms. Review the following bullets before you begin. Around the element s symbol use dots to represent each valance electron The dots should be spread over the four sides. Dots are not paired until all sides have at least one dot. Place the initial dot above each symbol and then proceed to arrange the rest in a clockwise direction. The number of valance electrons is equal to the group number (Roman numeral). 1. B 2. Ne 3. Li 4. He 5. C 6. P 7. S 8. Mg 9. H 10. F 19
20 Investigation 3.5: Understanding the formation of bonds and ions. The octet rule says that atoms tend to gain, lose or share electrons so as to have eight electrons in their outer electron shell. When atoms form ions they seek to obtain a stable electron configuration (8 valance electrons Lewis-dot structures). As a result, they will either attempt to gain electrons and or lose electrons to become stable. In either situation they will do whatever is the easiest. The protons (+) in the nucleus of an atom remain unchanged by ordinary chemical reactions, but atoms readily gain or lose electrons (-). When electrons are removed from or added to a neutral (non-charged) atom, a charged particle called an ion is formed. If the atom gains electron(s), its net charge becomes negative. A negative ion is called an anion. If the atom loses electron(s), its net charge becomes positive. A positive ion is called a cation. Gaining or losing electrons allows every energy level in an ion to hold only a "stable" number of electrons, namely, 2, 8, 18 or 32. How many electrons are in the following ion a. Al 3+ b. Te 2- c. Si 4+ d. Sb 3- e. Cs 1+ f. He g. C 4- h. Sr 2+ i. F 1- j. Bi 3- II. According to the Octet Rule, what do these ions need to do in order to become stable? Indicate how many electrons they would gain or lose. a. Calcium: Ca b. Boron: B c. Chlorine: Cl d. Indium: In f. Phosphorus: P g. Sodium: Na h. Tin: Sn i. Argon: Ar e. Sulfur: S j. Hydrogen: H III. Are the following atoms anions (a negative ion), cations, both, or neither. a. Magnesium: d. Neon: b. Gallium: e. Carbon: c. Rubidium: f. Iodine: 20
21 Investigation 4.1: Replacement Reaction Procedure: Safety- During this lab you will be wearing googles, gloves and should make sure that you do not spill any solutions on yourself as there is a strong possibility of staining clothes. 1. Obtain a piece of metal from your teacher and mass it. Record this mass in your notes. 2. Obtain a beaker for the mystery solution. 3. Obtain 0.9 grams of the mystery powder and 50 milliliters of water from the tap and dissolve the mystery powder in the water using a stirring rod. To avoid stains, be sure to rinse the stirring rod with distilled water before setting it down. 4. Place the piece of metal in the mystery solution, making sure not to splash any of the solution. 5. Make note of any reaction that takes place in the beaker. 6. After 10 minutes, gently shake the metal to remove the accumulated solid. Then place the metal on a paper towel on your lab bench to allow it to dry. 7. Once dry, mass the metal again and record this value in your notes. Filtering Procedure: 1. Obtain a piece of filter paper and initial the paper with a pencil. Then, mass the filter paper, recording this in your data table. 2. Fold the filter paper using the following visuals as a guide: 3. Place the filter paper in a funnel and place another beaker below the funnel. 4. Pour the solution from your beaker through the filter paper, allowing all of the solution to transfer to the bottom beaker. 5. Take your filter paper and let it dry. Once dry, mass the filter paper and precipitate. Subtract the mass of the filter paper to get a final value for amount of precipitate. Record all data in your data table. 21
22 Analysis Questions: 1) Describe what happened in the beaker when you placed the metal in the mystery solution? 2) Provide quantitative evidence to support what you described in the answer above. 3) Hypothesize what you think occurred when the metal and mystery solution were combined. 22
23 Investigation 5.1 Ionic Bonds In ionic bonding, electrons are completely transferred from one atom to another. In the process of either losing or gaining negatively charged electrons, the reacting atoms form ions. The oppositely charged ions are attracted to each other by electrostatic forces, which are the basis of the ionic bond. The following bullets characterize ionic bonding: Transfer of electrons Form between metals and nonmetals Are crystalline solids Dissolve easily in water Conduct electricity in solution Complete the chart for each element. Element # of Protons # of Electrons # of Valance Electrons Sodium Chlorine Beryllium Fluorine Lithium Oxygen Potassium Directions: Write the symbol for each element Using a different colored pen/pencil create a Lewis dot structure for each element Draw an arrow(s) indicating the transfer of the electron(s) Determine the charge of each ion and write the formula Make sure the sum of the bond is zero 1. Potassium + Fluorine 2. Magnesium + Iodine 23
24 3. Sodium + Oxygen 4. Sodium + Chlorine 5. Calcium + Chlorine 6. Aluminum + Chlorine 24
25 Investigation 6.1Covalent Bonds Covalent Bonding Covalent bonding occurs when atoms share electrons. As opposed to ionic bonding in which a complete transfer of electrons occurs, covalent bonding occurs when two (or more) elements share electrons. Covalent bonding occurs because the atoms in the compound have a similar tendency for electrons (generally to gain electrons). The following bullets characterize covalent bonding: Share electrons Generally occur between two nonmetals Complete the chart for each element. Element # of Protons # of Electrons # of Valance Electrons Hydrogen Oxygen Chlorine Carbon Fluorine Helium Lithium # of Electrons to Fill Outer Shell Directions: Write the symbol for each element Using a different colored pen/pencil create a Lewis dot structure for each element Rearrange the electrons to pair up electrons from each atom. Draw circles to show the sharing of electrons Write the chemical formula for each molecule 1. Hydrogen + Hydrogen (Diatomic Element) 2. Hydrogen + Oxygen 25
26 3. Chlorine + Chlorine (Diatomic Element) 4. Oxygen + Oxygen (Diatomic Element) 5. Carbon + Oxygen 6. Carbon + Hydrogen 26
27 Investigation 7.1: Lab: Ionic vs. Covalent Bonds Hypothesis: (Based on the properties of ionic and covalent compounds make a hypothesis as to which type of bond each compound will have) Procedure: Compound Hypothesis Compound Hypothesis Gelatin Magnesium Oxide Sugar Detergent Salt Cornstarch Baking Soda Copper Chloride 1. You will need safety glasses, lab apron, and gloves. 2. Label 8 pieces of paper towel (good size) with the names of the compounds. 3. Place 5 grams of the compounds on each piece of paper (do ONE at a time). 4. Observe and WRITE the description of the compounds on the data chart. 5. Place 200 ml. of water in a beaker. 6. Add about ½ of the compound to the water and stir with a stirring rod for 10 seconds and observe. If the compound dissolves write SOLUBLE in the data chart under Solubility. If it does not dissolve, write INSOLUBLE. 7. Place the leaders of the conductivity tester (as shown below) into the mixture. Observe whether or not the bulb lights. If it does, write YES under Conductivity in the data chart. If it does not light up write NO. 8. Rinse and dry the beaker. 9. Repeat the steps for each of the compounds. 10. Be sure to follow the teacher s directions for how to dispose of the mixtures. DO NOT just dump them into the sink. 27
28 Data: Compound Description Solubility Conductivity Bond Gelatin Sugar Salt Baking Soda Magnesium Oxide Detergent Cornstarch Copper Chloride Analysis Questions: 1. Some characteristics of many ionic compounds include solubility in water and the ability to conduct electricity. On the basis of these two properties, which compounds appear to have ionic bonds (please indicate in last column Type of Bond in data chart)? 2. Water solutions of covalent compounds do not conduct electricity. Based on this property, which compounds that you tested would you classify as covalent compounds (please indicate in last column Type of Bond in data chart)? 3. Explain how ionic and covalent compounds are different. USE EVIDENCE FROM YOUR LAB. 4. Did all of the compounds that conducted electricity show the same amount of conductivity? How can you tell (from the lab )? 5. Solid table salt does not conduct electricity. Why do you think dissolving table salt in water allows the salt to conduct electricity? 28
29 Unit 2 Study Guide-Atoms, Bonds, ions, and their properties Key Terms: You must be able to define and explain each of the following terms below Atom Molecule Element Proton Neutron Electron Nucleus Ionic bonding Covalent bonding Isotope Ion Cation Anion Lewis Dot Atomic number Atomic weight group/family Period Octet Rule Valence Electron Application: You must be able to evaluate and utilize the periodic table in order to: Decipher the number of proton, electrons, neutrons, atomic number, and atomic weight. Evaluate whether the element is a Alkali Metal, Alkali Earth Metal, Transition Metal, Metal, Metalloide, Non-Metal, Halogen, or a Noble gas In terms of characteristics, be able to identify and explain the similarities and difference between and among the eight different groups Situate the electrons into their respective orbital using the electron configuration model and the Lewis-Dot model. Interpret and explain ions according to the Octet Rule, example Ca 20, protons 20 electrons, and 20 neutrons. According to the Octet Rule it is simpler for Ca to give away 2 electrons than gain eight. Therefore, Ca becomes Ca 2+ with electron configuration of Determine whether an ion is a cation, anion, or neither. Draw and depict a covalent and ionic bond Graphing/ Short Response: Be able to organize and synthesize data 29
30 Unit Three Key Terms Directions: Define the following terms using either your book or the internet. Please attempt to make the definition clear, concise, and simple. 1. Atom 2. Molecule 3. Element 4. Proton 5. Neutron 6. Electron 7. Nucleus 8. Ionic bonding 30
31 9. Covalent bonding 10. Ion 11. Cation 12. Anion 31
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