Chapter 11: Chemical Bonds: The Formation of Compounds from Atoms

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1 Chapter 11: Chemical Bonds: The Formation of Compounds from Atoms Name: Many of the concepts in this chapter come from the idea that elements are always trying to obtain 8 valence electrons because this is extremely stable. This is called a: or a. 11.1: Periodic Trends in Atomic Properties Metals are on the side of the periodic table. Nonmetals are on the side of the periodic table. Metalloids are located on. Metals tend to (lose/gain) electrons to form (positive/negative) ions called. Metals do this in order to achieve a noble gas configuration. Nonmetals tend to (lose/gain) electrons to form (positive/negative) ions called. Nonmetals do this in order to achieve a noble gas configuration. When metals react with nonmetals, electrons are Determine the # of electrons which must be gained or lost for the following elements to achieve a noble gas configuration: 1. a calcium atom 3. a sodium atom 2. a sulfur atom 4. a fluorine atom Define Atomic Radius: In going from the top to the bottom of a group, the atomic radius of elements. Why does this happen? In going from the left to the right across a period, the atomic radius of elements. Why does this happen? Ionic Radius: Remember that an ion is a charged particle, meaning a neutral atom has either lost or gained electrons. Size of cations vs. neutral atoms: Ex: Ca vs. Ca +2 Size of anions vs. neutral atoms: Ex: S vs. S -2 Define Ionization Energy (IE): What does a higher Ionization Energy value mean? In going from the top to the bottom of a group, the first ionization energy. Why? (Discuss the shielding effect) 1

2 In going from the left to the right across a period, the first ionization energy. Why does this happen? The ionization energy values are the highest for which specific group in the periodic table? Why is this? Label the atomic radius and ionization energy trends on the periodic table below: Define Lewis Structure (Lewis Dot Notation): 11.2: Lewis Structures of Atoms Draw the Lewis Structures for the following elements: 1. Hydrogen 6. Oxygen 2. Beryllium 7. Fluorine 3. Boron 8. Neon 4. Carbon 9. Sodium Nitrogen 10. Chlorine -1 What is the driving force for elements to form a chemical bond to one another? 11.3: The Ionic Bond: Transfer of Electrons from One Atom to Another 2

3 Define Ionic Bond: To predict the formula of an ionic compound, metals lose e- & nonmetals gain e- in order to achieve noble gas configurations. Group 1A metals always form cations. Ex: Group 2A metals always form cations. Ex: Group 3A metals always form cations. Ex: Group 5A elements always form anions. Ex: Group 6A elements always from anions. Ex: Group 7A elements always from anions. Ex: In the discussion in this section, we refer only to representative metals (Groups 1A, 2A, and 3A). The transition metals show more complicated behavior because they can form multiple ions and their formulas are not as easily predictable. 1. Illustrate and explain how sodium and chlorine combine to form sodium chloride, NaCl. 2. Illustrate and explain the formation of the ionic compound formed between the elements magnesium and chlorine. 3. Illustrate and explain the formation of the ionic compound formed between the elements aluminum and fluorine. 4. Illustrate and explain the formation of the ionic compound formed between the elements sodium and oxygen. Define a Covalent Bond: 11.5: The Covalent Bond: Sharing Electrons Illustrate and explain the covalent bonding in a hydrogen molecule. Illustrate and explain the covalent bonding in a chlorine molecule. Define Electronegativity: 11.6: Electronegativity The scale for electronegativity is based on the most electronegative element,, which is assigned a value of. In going from the top to the bottom of a group, the electronegativity of elements. Why is this? 3

4 In going from the left to the right across a period, the electronegativity of elements. Why is this? How do we know whether a bond between 2 atoms is classified as being ionic or covalent? *Check pg. 241: Table 11.5 for Electronegativity values. Draw the electronegativity difference scale, which illustrates the type of bond that s formed between atoms: Define a Dipole: Draw the dipole(s) in the following molecules: 1. HF 2. H HCl What happens when the dipoles in a molecule are pointing in opposite directions? Draw an example of when this occurs. Define a Non-polar Covalent Molecule: Ex: Define a Polar Covalent Molecule: Ex: Classify the following compounds as either ionic or covalent. If they are covalent, are they polar covalent or non-polar covalent? 1. H 2 O 5. HBr 2. NaCl 6. SO 2 3. MgO 7. NH 3 4. Br 2 8. BaCl 2 4

5 11.7: Lewis Structures of Compounds In writing the Lewis structures for compounds, what is the most important consideration for forming a stable compound? What is the one exception for this? In determining the arrangement of atoms in a compound, usually (but not always), the single atom in the formula (except H) will be the central/middle atom. Steps to draw the Lewis structures of compounds: 1. Count the total # of valence electrons from all of the atoms in the formula. 2. Write the skeletal arrangement of the atoms and connect them with a single covalent bond (2 dots or 1 dash). *Since hydrogen contains only 1 electron, it can form only one covalent bond. 3. Subtract 2 electrons for each single bond used in step 2 from the total # of electrons from step Distribute pairs of electrons (dots) around each atom (except H) to give each atom a noble gas configuration. 5. If there aren t enough electrons to give each atom 8 electrons, change single bonds to double or triple bonds by shifting unbonded pairs of electrons as needed. Check each atom has a noble gas configuration. A double bond counts as 4 electrons for each atom to which it is bonded. Examples: 1. Draw the Lewis structure for water. 2. Draw the Lewis structure for methane. 3. Draw the Lewis structure for carbon dioxide. 4. Draw the Lewis structure for phosphorous tribromide 11.8: Complex Lewis Structures Drawing the Lewis Structures for polyatomic ions (a charged group of atoms sulfate, nitrate, etc.) is a bit different. Examples: 1. Draw the Lewis Structure for a nitrate ion. Define Resonance (as it applies to the nitrate ion s Lewis structures above): 2. Draw the Lewis Structure for a carbonate ion. 11.9: Compounds Containing Polyatomic Ions Most compounds containing polyatomic ions consist of both ionic AND covalent bonds. 5

6 Example: 1. Draw the Lewis Structure for sodium nitrate AND describe the 2 types of bonding in the compound. What happens when sodium nitrate is dissolved in water? 11.10: Molecular Shape The Lewis Structures we have been working with do NOT indicate the actual shape of a molecule. The 3-D arrangement of the atoms within a molecule can give a lot of important info about the chemistry of the molecule. What is the main idea behind the VSEPR model? How do we determine the 3-D shape of a molecule?... By using VSEPR : The Valence Shell Electron Pair Repulsion (VSEPR) Model # of Bonds to Central Atom # of Unpaired Electron Pairs on Central Atom Sketch of the Molecular Structure Model Name of Molecular Structure Bond Angles Examples Predict and draw the shape of the following molecules: 1. CF 4 3. BeI 2 5. AlH 3 2. NF 3 4. SeF 2 6. SO 2 6

7 Periodic Trends Worksheet Name: 1. What do we mean by the atomic radius? 2. Within a group, what happens to the atomic radius as you go down a group? WHY does this happen? 3. Within a period, what happens to the atomic radius as the atomic number increases? WHY does this happen? 4. How are neutral atoms converted into cations? How are neutral atoms converted into anions? 5. Metals usually form what type of ions? Nonmetals usually form what type of ions? 6. When a neutral atom becomes an anion, what happens to its radius? WHY does this happen? 7. When a neutral atom becomes a cation, what happens to its radius? WHY does this happen? 8. For each of the following pairs, circle the atom or ion having the larger radius. a. S or O c. Na +1 or K +1 e. S 2 or O 2 b. Ca or Ca +2 d. Na or K f. F or F 1 9. For each of the following pairs, identify the smaller ion. a. K +1 or Ca +2 c. C +4 or C 4 e. O 2 or F 1 b. F 1 or Cl 1 d. S 2 or F 1 f. Fe +2 or Fe What does the term ionization energy mean? 11. What do we mean by the first, second, and third ionization energies for a particular atom? 12. Why does each successive ionization require more energy than the previous one (meaning 2 nd ionization takes more energy than 1 st )? 13. What is the general trend of ionization energy as you go from left to right across the periodic table? WHY is this? 14. What is the general trend of ionization energy as you go down a group on the periodic table? WHY is this? 15. Which of these elements has the highest first ionization energy: Sn, As, or S? 7

8 Ionic and Covalent Bonding Name: An ionic bond is a bond between a and a. A covalent bond is a bond between two or more. 1.) How are ionic and covalent bonding different (talk about what happens with the electrons)? 2.) What two types of covalent bonds are there? What is the difference between the two? 3.) What does the term electronegativity mean? 4.) What is the trend in electronegativity from the top to bottom of a group? WHY? 5.) What is the trend in electronegativity from the left to right of a period? WHY? 6.) When there is no difference in the electronegativity of atoms in a bond, what type of bond forms? 7.) When there is an electronegativity difference of , what type of bond forms? 8.) When there is an electronegativity difference of , what type of bond forms? 9.) Using the electronegativity table as a guide, classify the bond as covalent (and further as nonpolar or polar) or ionic. 1. O 2 4. HBr 7. BaCl 2 2. NH 3 5. SO 3 8. PCl 5 3. NaF 6. OF 3 9. SrO 10.) Draw the dipole in the following molecules. 1. CO 2. ClF 3 3. CO 2 4. NH 3 5. CCl 4 6. HI 8

9 Lewis Structure Worksheet Name: Complete the Lewis structures for each of the following molecules: Simple Molecules 1.) NH 3 2.) SiF 4 3.) NCl 3 CH 4 Polyatomic Ions 1.) Phosphate: PO ) Sulfite: SO ) Nitrite: NO 2-1 Multiple Bonds 1.) HCN 2.) CO 3.) C 2 H 3 Br 4.) C 2 H 2 Incomplete Octet* Elements that don t require 8 valence e- to be stable 1.) BCl 3 2.) BeF 2 Diatomic Molecules 1.) F 2 2.) O 2 3.) N 2 9

10 It s a Chemical Bond Nothin on You Remix [CHORUS] - Ionic bonds, and covalent bonds, Forming and breaking, new molecules making, It's a chemical bond, babehhhhh, A chemical bond, baby, Electron pairs, are they transferred or shared? How do you name them, is polarity there? It's a chemical bond, baby, A chemical bond, baby. [RAP VERSE 2] Im'a talk about covalent bonds now Two nonmetals bond, you wanna know how? Atoms come together, cause they want an octet They share electrons to get eight in their set. Covalent differs from ionic greatly: They have a smaller difference in e-negativity, Percent ionic character is less obviously, What, about polarity? [RAP VERSE 1] - Okay first off we'll start with ionic. They deal with ions, isn't that ironic, Cations, anions, polyatomic, Make a strong bond, like a bomb that's atomic, Typically a metal and a nonmetal bond, To get an octet so that they become strong, They get together and transfer electrons, They share them equally, neither one is conned. [SINGING VERSE 2] Non-polar, beats polar, When it comes to sharing equally. But polar, has a higher, Melting point than non-polar's be. Naming them is easy thankfully, Just put a prefix on the atom please, Rappin' bout bonds, yeah I'm a G. G...G...G [SINGING VERSE 1] - What makes the bonds stronger? And last longer? And have a higher melting point you think? A correct inference, would be the difference, In electronegativity. To name it put the cation first you see, And on the anion put I D E Unless it's polyatomic then let it be Be...be...be... [BRIDGE] Now metallic bonds, are the bond type number three. Electrons aren't stagnant, they move around in a sea. It's senior year with Standish, and I'm lucky with a "C," Can't wait to get out of here, it's AP Chemistry. [CHORUS] 10

11 It s a Periodic Trend Throw it in the Bag Remix [Verse 1 Atomic Radius] Now I know you be thinking bout them periodic trends (Yeah I know you be thinkin bout them periodic trends) And I know you want to know bout dat atomic radius (Yeah I know you want to know bout dat atomic radius) Not a neutron, an electron You be goin down a group and dat atom get bigga The shielding is increasing And as you go across a row that atom don't get thicka [Chorus] Uh-oh, Uh-oh, Uh-oh, its a periodic trend, Uh-oh, Uh-oh, Uh-oh, its a periodic trend, [Verse 2 Ionization Energy] Enough of dat singin, its time fo me to rap, A periodic trend, I don t mean to offend, Its betta den da first, not radius crap, Take an electron, ionization So you got an atom, with a valence e, To take one of em out, you ll need some energy, So its easiest if you only got one, Its gonna be harda if you gots to take a ton. So take a look at the table of the elements, We got a trend for you, so it ll make sense, At the left you got, few electrons, So it s not gonna take much exertion; But at the right, those e s put up a fight, It takes some energy, before they take flight. So from low to high, and from left to right, A slanted arrow, you are gonna write. [Chorus] x2 [Verse 3 Electronegativity] We did radius and ionization energy Now for trend numba three, you are about to see, The ability, to attract an e, the definition of electronegativity. Atoms always want, if possible, To have gases noble, a full orbital, In this state they are most stable, Getting electrons, they re unable. So gases closest to the nobles be, Ones with higher electronegativity, As you go from group eighteen further away, Electronegativity decays. From the bottom of the table to the top, The energy you need gets bigga- it doesn t stop. If you think we aint da greatest, then you pretend, you don t know us Standish (our teacher), it s a periodic trend. [Chorus] x2 [Refrain] Everybody know tha periodic trends, (Ay, know tha periodic trends) Everybody know tha periodic trends, (Ay, know tha periodic trends) Everybody know tha periodic trends, (Ay, know tha periodic trends) 11

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