CP Chemistry Mrs. Klingaman Chapter 6: Chemical Bonding Ionic Bonding: Electron Transfer Covalent Bonding: Electron Sharing Name: _______________________________________________ Mods: _____________________________ Chapter 6: Chemical Bonding Reading Guide 6.1 – Introduction to Chemical Bonding (pgs. 165-167) 1. What is a chemical bond? 2. Distinguish between an ionic bond and a covalent/molecular bond: Ionic bond - Covalent bond - 3. Bonding between atoms of different elements is rarely purely ionic or purely covalent; it usually falls somewhere between these two extremes. The degree to which bonding between atoms of two elements is ionic or covalent can be estimated by calculating the difference in the elements’ __________________________. 4. Use Figure 1.2 to fill in the following table. Difference in Electronegativity Bond Type 0.0 - 0.3 0.4 - 1.7 1.7 - 3.3 5. How do polar covalent bonds and nonpolar covalent bonds differ? Nonpolar covalent - Polar covalent - 6.2 – Covalent Bonding and Molecular Compounds (pgs. 168-179) 6. Define the following words: Molecule- Molecular formula- 7. Why do atoms form chemical bonds? Nature favors chemical bonding because atoms have _____________________________________________ when they are bonded to other atoms than they have when they are independent particles. 8. How do atoms form chemical bonds? The attractive and repulsive forces between two atoms play an important role in the formation of a chemical bond. (refer to Figure 2.3) Attractive Forces – Repulsive Forces – The relative strength of attraction and repulsion between the charged particles (nucleus & electrons) depends on the ________________________ separating the atoms. 9. Define the following terms: Bond length- Bond energy- The Nature of Covalent Bonding 10. Atoms tend to form bonds to follow the octet rule. What is the octet rule? 11. Covalent bonding involves only the atom’s outermost, or _____________________ electrons. To keep track of these electrons, it is helpful to use __________________________ notation. 12. Using Lewis Dot Structures to Represent Compounds: Using Lewis Dot Structures, show how two fluorine atoms form a single covalent bond to fulfill the octet rule (pg. 175): The pair of electrons in between the two fluorine atoms represents the ________________ pair of electrons involved in the covalent bond. For this reason, they are also known as the bonding electrons. This pair of electron-dots is often replaced by a long _____________, which represents the covalent bond. An ___________________ pair, or ____________ pair, of electrons are those that are not involved in the bonding and belong exclusively to one atom. For this reason, they are also known as nonbonding electrons. 13. Lewis structures are drawn to show the composition of a molecule, which atoms are bonded to each other, and what type of bond holds them together. There are three types of covalent bonds: Single Bond – Double Bond – ► Example: ethene, C2H4 (pg. 176) Triple Bond – ► Example: nitrogen gas, N2 (pg. 177) 14. Using Figure 2.10 on pg. 177, fill in the table below Bond Bond Length (pm) Bond Energy (kJ/mol) C–C C=C C C When a multiple bond forms, bond length ______________and bond energy ______________. 15. Define resonance and give an example (ozone, O3 - pg. 179) of structures that exhibit resonance: 6.3 – Ionic Bonding and Ionic Compounds (pgs. 180-184) The Nature of Ionic Bonding 16. Define the following terms: Ionic compound- Formula unit- 17. Atoms of metals, such as sodium, readily ________ electrons to form ______________, while atoms of nonmetals, such as chlorine, readily _________ electrons to form ______________. 18. Show the formation of sodium chloride (NaCl) from sodium and chlorine ions (pg. 181) 19. All ionic compounds have an orderly arrangement known as a _______________ lattice structure. This structure arranges the cations and anions so that they surround one another giving the compound its lowest potential energy, thus making it very stable. 20. Define lattice energy. How is this different from bond energy? 21. The strong attraction of (+)/(-) ions is described by the lattice energy (kJ/mol) in Figure 3.6. All lattice energies are negative, which means that energy is released when ionic bonds are formed. This is a good thing, because all atoms and compounds want to be at their lowest possible energy state. In fact, the higher the lattice energy (the greater the energy released) the stronger the ionic bond and the more stable the compound. 22. The forces of attraction between molecules are much _______________________ than the forces of attraction among formula units in ionic bonding. Thus, when comparing the strength of ionic and covalent bonds, _______________________ compounds have a much stronger attraction which gives rise to different properties in the two types of compounds which we will discuss further in class. 23. Define polyatomic ion and give one example: 6.4 – Metallic Bonding (pgs. 185-186) 24. Define the following terms: Metallic bonding- Malleability- Ductility- 25. The amount of energy as heat required to vaporize a metal is a measure of the strength of the metallic bonds that hold it together. This is known as the metal’s _______________________ ___________________________________. Generally, the greater the metal’s nuclear charge and the more electrons the metal has, the stronger the metallic bonds (refer to Figure 4.3). Summary of Chemical Bonds and Their Properties Ionic Bond: a nonmetal ____________ electrons from a metal due to a __________________ difference in electronegativity, leading to an attraction between the resulting cation and anion. Example: Polar Covalent Bond: an __________________ difference in electronegativity leads to an unequal sharing of electrons between two ___________________. Example: Nonpolar Covalent Bond: two ____________________ with ________________ electronegativity or only a ______________ difference in electronegativity equally share electrons. Example: Metallic Bond: results from the attraction between ___________ atoms and the surrounding sea of __________________. Example: Compound Melting Point Conductivity of Solid Solubility in Water (polar solvent) Conductivity in Water Solubility in Hexane (nonpolar solvent) Hardness Brittleness Ionic Bond Polar Covalent Bond Nonpolar Covalent Bond Metallic Bond Sodium Chloride Sucrose Stearic Acid (NaCl-table salt) (C12H22O11-table sugar) (CH3(CH2)16COOH-wax) Aluminum (Al) and Mercury (Hg) Determining the Type of Chemical Bond Bond types based on differences in electronegativity: Ionic > 1.7 Polar Covalent = 0.4 – 1.7 Nonpolar Covalent < 0.4 Directions: Determine the electronegativity difference between the atoms listed in the table below and then determine the corresponding type of chemical bond. Atoms that are bonding 1 Hydrogen and oxygen 2 Magnesium and sulfur 3 Strontium and fluorine 4 Carbon and oxygen 5 Magnesium and nitrogen 6 Nitrogen and oxygen 7 Chlorine and bromine 8 Hydrogen and nitrogen 9 Lithium and oxygen 10 Oxygen and fluorine 11 Bromine and oxygen 12 Hydrogen and iodine 13 Magnesium and hydrogen 14 Hydrogen and hydrogen Electronegativity Difference Bond Type Electron Dot Structure & Ions 1) Draw electron dot structures for the following atoms: a. Magnesium (Mg) e. Oxygen (O) b. Bromine (Br) f. c. Krypton (Kr) g. Lithium (Li) d. Silicon (Si) h. Phosphorous (P) Aluminum (Al) 2) Fill in the table below by stating if the atom will lose or gain electrons to form an ion, how many electrons it will lose/gain, what the resulting ion symbol will be, and if it is a cation or an anion. Atom Lose or Gain? How Many Electrons? Ex) Nitrogen gain 3 Magnesium Bromine Krypton Silicon Oxygen Aluminum Lithium Phosphorous Ion Symbol Cation or Anion? anion The Formation of Ions from Atoms Answer the following questions about CALCIUM: Is calcium a metal or a nonmetal? _____________________ How many total electrons does a calcium atom have? ____________ How many valence electrons does a calcium atom have? _____________ Draw the electron dot structure for a calcium atom in the space below: Calcium will ___________ ____ electrons to become an ion. Draw the symbol for a calcium ion in the space below: Is the calcium ion a cation or an anion? ______________________ The calcium ion has the same electron configuration as ____________________. Answer the following questions about SULFUR: Is sulfur a metal or a nonmetal? _____________________ How many total electrons does a sulfur atom have? ______________ How many valence electrons does a sulfur atom have? _____________ Draw the electron dot structure for a sulfur atom in the space below: Sulfur will ___________ ____ electrons to become an ion. Draw the symbol for a sulfide ion in the space below: Is the sulfide ion a cation or an anion? ______________________ The sulfide ion has the same electron configuration as ____________________. Valence Electrons Review 1) Define valence electrons- 2) Fill in the following table: Element Barium Sulfur Aluminum Bromine Sodium Carbon Lithium Neon Group # Block (s, p, d, f) Period # # of Valence Electrons Ion Symbol Lewis Structures – Ionic Compounds Directions: Draw Lewis structures for the following ionic compounds showing the transfer of electrons and the ions formed. 1) SrCl2 2) AlF3 3) BaI2 4) LiBr 5) Al2S3 6) CaO Lewis Structures – Ionic Compounds Continued Directions: Draw Lewis structures for the following ionic compounds showing the transfer of electrons and the ions formed. 7) K3P 8) Mg3N2 9) Na2O 10) FrF 11) Be3As2 12) MgSe Rules for Drawing Lewis Dot Structures 1) Determine the number of atoms of each element present in the molecule 2) Sum the valence electrons from all atoms For an ANION add one electron to the total for each negative charge For a CATION subtract one electron from the total for each positive charge 3) Determine the central atom in the molecule and attach all other atoms to it The central atom is the least electronegative element in the compound excluding hydrogen – hydrogen can NEVER be central! (Note: typically, the central atom is the one written first in the molecular formula) Write the symbols for all the other atoms around the central atom. Connect the atoms to the central with a single bond (a dash). Keep track of the electrons being used. Each single bond made uses 2 electrons. Chemical formulas are often written in the order in which the atoms are connected in the molecule (ex: HCN carbon is the central atom) 4) Complete the octets around all the outer atoms bonded to the central Note: Hydrogen atoms may only have a single electron pair around them Keep track of the electrons being used to complete the octets 5) Place any leftover electrons on the central atom, even if doing so results in more than an octet of electrons around the atom. Note: Expanded octets occur when more than 8 electrons surround an atom. This exception to the octet rule is allowed for any atom in the 3 rd row of the periodic table and after! Keep track of the electrons – make sure the total number of valance electrons available were used in the Lewis structure 6) If there are not enough electrons to give the central atom an octet, multiple bonds are needed. Remove one or more of the nonbonding pairs of electrons on one of the outer atoms and draw a double bond (2nd dash) connecting the outer atom to the central. If need be, a triple bond (3rd dash) may be formed by removing another nonbonding pair from the same outer atom. Keep track of the electrons – make sure the total number of valance electrons available were used in the Lewis structure Drawing Lewis Structures – Examples Lewis Dot Structures: 1) Methane: CH4 4) Sulfite ion: SO32- 2) Hydrochloric Acid: HCl 5) Hydrocyanic Acid: HCN 3) Ammonia: NH3 6) Methanal (aka: formaldehyde): H2CO Resonance Structures: multiple Lewis structures are used to describe molecules which have double or triples bonds that can be moved to different sides of the central atom and still stay bonded to an identical outer atom. 7) Nitrite ion: NO2 – Exceptions to the Octet Rule: Less than an Octet central atoms which have FEWER than 8 electrons (seen with B & Be) 8) BF3 Expanded Octet central atoms which have MORE than 8 electrons (this is allowed for central atoms found in the 3rd period and after) 9) SBr6 10) XeI4 Lewis Structures – Covalent Compounds Directions: Draw Lewis structures for the following molecules and polyatomic ions. Remember to consider exceptions to the octet rule, double/triple bonds, ion charges, and resonance structures. 1) CBr4 4) NBr3 2) SiS2 5) CO32– 3) SF4 6) PCl5 7) N2 10) H2S 8) CN– 11) CCl2O 9) NH4 + 12) BCl3 Chapter 6 Review 1) Identify and define the three major types of chemical bonding. ___________________________________________________________________________ ___________________________________________________________________________ ___________________________________________________________________________ ___________________________________________________________________________ ___________________________________________________________________________ ___________________________________________________________________________ 2) In a covalent bond, how is the octet rule fulfilled so that every atom in a molecule has 8 valence electrons? ___________________________________________________________________________ ___________________________________________________________________________ ___________________________________________________________________________ 3) In an ionic bond, how is the octet rule fulfilled? ___________________________________________________________________________ ___________________________________________________________________________ ___________________________________________________________________________ 4) In writing Lewis structures, how do you know when a multiple bond is needed? ___________________________________________________________________________ ___________________________________________________________________________ ___________________________________________________________________________ 5) In general, how do ionic and covalent compounds compare in terms of melting points and boiling points? What does this tell you about the relative strength of ionic vs. covalent bonds? ___________________________________________________________________________ ___________________________________________________________________________ ___________________________________________________________________________ ___________________________________________________________________________ 6) How do the properties of metals differ from those of both ionic and molecular compounds? ___________________________________________________________________________ ___________________________________________________________________________ ___________________________________________________________________________ ___________________________________________________________________________ 7) Complete the table below by determining what type of bond (ionic, polar covalent, or nonpolar covalent) will form between each set of atoms. Bonded Atoms Electronegativity Difference Bond Type More Electronegative Atom H and I S and O K and Br Si and Cl K and Cl Se and S C and H 8) List the bonding pairs from Question 8 in order of increasing ionic character. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ 9) The lattice energy of sodium chloride, NaCl is -787.5 kJ/mol. The lattice energy for potassium chloride, KCl, is -715 kJ/mol. In which compound is the bonding between ions stronger? Why? __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ 10) Use electron dot notation to illustrate the number of valence electrons in each of the atoms below. a. Iodine b. Strontium c. Carbon d. Nitrogen Drawing Lewis Diagrams: Covalent vs. Ionic 1. K2O Type:___________________ 2. OF2 Type:___________________ 3. PO43- Type:__________________ 4. BeS Type:___________________ 5. AlBr3 Type:__________________ 6. CH2O Type:__________________ 7. IBr5 Type:____________________ 8. CaF2 Type:__________________ 9. CH2F2 Type:__________________ 10. H2S Type:____________________ 11. Ba3N2 Type:__________________ 12. SiO2 Type:___________________ **Challenge** -- NO31- Type: _________________________