Zumdahl • Zumdahl • DeCoste World of CHEMISTRY

Chapter 12 Chemical Bonding

12.1 Types of Chemical Bonds Objectives: To learn about ionic and covalent bonds and explain how they are formed. To learn about the polar covalent bond. Copyright © Houghton Mifflin Company

What is a chemical bond? A bond Bond energy is a force that holds groups of two or more atoms together and makes them function as a unit. Bond energy The strength of a bond can be calculated by the energy it takes to break the bond Copyright © Houghton Mifflin Company

Ionic Bonding Ionic bonds form when Ionic compounds an atom loses electrons relatively easily with an atom that has a high affinity for electrons. Ionic compounds are the result of a metal reacting with a non-metal. Copyright © Houghton Mifflin Company

Electrons are shared by the nuclei Covalent Bonding Electrons are shared by the nuclei The electrons are attracted to the two nuclei Copyright © Houghton Mifflin Company

Figure 12.1: The formation of a bond between two identical atoms. Note the electrons tend to reside in the space between the two nuclei. Copyright © Houghton Mifflin Company

Electrons are not shared evenly Polar Covalent Bonds Electrons are not shared evenly One of the atoms has a stronger electron Copyright © Houghton Mifflin Company

Figure 12.2: Probability representations of the electron sharing in HF. Copyright © Houghton Mifflin Company

12.2 Electronegativity Objective: to understand the nature of bonds and their relationship to electronegativity Copyright © Houghton Mifflin Company

Electronegativity The relative ability of an atom in a molecule to attract shared electrons to itself. Electronegativity values can be determined by measuring the bond polarities. Copyright © Houghton Mifflin Company

Figure 12.3: Electronegativity values for selected elements. F has the highest electronegativity of 4.0 The polarity of a bond depends on the differences between the electronegativity values of the atoms forming the bond. Copyright © Houghton Mifflin Company

As a rule: non-metals + Metals = ionic If the differences between the electronegativity of two elements is 2.0 or greater, the bond is considered to be ionic. As a rule: non-metals + Metals = ionic Copyright © Houghton Mifflin Company

Table 12.1 Copyright © Houghton Mifflin Company

What if atoms had the same value of electronegativity values? How would bonding between atoms be affected? What are some differences we would notice? Copyright © Houghton Mifflin Company

Figure 12.4: The three possible types of bonds. Copyright © Houghton Mifflin Company

Example 12.1 Using Electrons to Determine Bond Polarity Arrange the following bonds in order of increasing polarity: H-H O-H Cl-H S-H F-H Copyright © Houghton Mifflin Company

12.3 Bond Polarity and Dipole Moments Objective: To understand bond polarity and how it is related to molecular polarity Copyright © Houghton Mifflin Company

Dipole Moment Dipole moments creates a region of positive charge and a region of negative charge, often indicated with the use of an arrow. Copyright © Houghton Mifflin Company

Figure 12.5: Charge distribution in the water molecule. Copyright © Houghton Mifflin Company

Figure 12.5: Water molecule behaves as if it had a positive and negative end. Copyright © Houghton Mifflin Company

Figure 12.6: Polar water molecules are strongly attracted to positive ions by their negative ends. Copyright © Houghton Mifflin Company

Figure 12.6: Polar water molecules are strongly attracted to negative ions by their positive ends. The ability of water to create dipole moments explains the crucial role that water plays in survival: Salts dissolve in water Water requires a lot of energy to change to a gas. Copyright © Houghton Mifflin Company

How are ionic bonds and covalent bonds different? Focus Questions 12.1 – 12.3 How are ionic bonds and covalent bonds different? How does a polar covalent differ from a nonpolar covalent bond? How do electronegativity values help us to determine the polarity of a bond? For each of the following binary molecules, draw an arrow under the molecule showing its dipole moment (if none exists, write none). Copyright © Houghton Mifflin Company

12.4 Stable Electron Configuration and Charges on Ion Objective: To learn about stable electron configuration. To learn to predict the formulas of ionic compounds. Copyright © Houghton Mifflin Company

Review Trends in the periodic table: certain elements grouped together because they behave similarly there is great similarities within groups, but the differences in behavior between groups is what we will be studying. Copyright © Houghton Mifflin Company

Trends in the periodic table: Table 12.2 Copyright © Houghton Mifflin Company

Electron Configuration of Ions Main group metals form ions with the electron configuration of the previous noble gas . Nonmetals form ions by gaining enough electrons to behave like the noble gas that follows. Copyright © Houghton Mifflin Company

Electron Configuration and Bonding metals group 1,2 &3 and non-metals react to form binary compounds. Valence electrons is completed to form electron configuration of previous noble gas. Two nonmetals react to form a covalent bonds, they share electrons in a way that completes the valence-electron configuration of both atoms. Both nonmetals attain noble gas electron configurations by sharing electrons. Copyright © Houghton Mifflin Company

Predicting Formulas of Ionic Compounds Chemical compounds are always electrically neutral – they have the same number of protons and electrons. To achieve an ionic bond the metal and nonmetal must lose or gain electrons such that they both achieve the configuration of a noble gas. Copyright © Houghton Mifflin Company

Table 12.3 Copyright © Houghton Mifflin Company

12.5 Ionic Bonding and Structures of Ionic Compounds Objectives: To learn about ionic structures To understand factors governing ionic size Copyright © Houghton Mifflin Company

Ionic Compounds are Very stable The strong bonds in these compounds results from the attractions between the oppositely charged cations and anions. Large amounts of energy are required to ‘take them apart’ e.g. the melting point of NaCl is approx. 800°C When metals and nonmetals react, the resulting compounds are stabel Large amounts of energy are required to take them apart Copyright © Houghton Mifflin Company

Figure 12.9: Relative sizes of some ions and their parent atoms. Notice that when a metal loses all of its valence electrons to form a cation, it gets much smaller . *size is given in picometers 1012m Copyright © Houghton Mifflin Company

Figure 12.8: Ions as packed spheres. The structure of ionic compounds results from the packing of ions as hard spheres. Note: the cation is always considerable and smaller than the anion. This allows the cations to fill the empty space between the anions creating a compact compound Copyright © Houghton Mifflin Company

Figure 12.8: Positions (centers) of the ions. This structure shows the position (centers) of the ions. The spherical ions are backed in a way to maximize the ionic attractions. Copyright © Houghton Mifflin Company

Ionic Compounds Containing Polyatomic Ions Polyatomic atoms: charged species composed of several atoms. The individual polyatomic ions (e.g. NH4+) are held together by covalent bonds – all of the atoms behave as a unit. The reaction between NH4+ with NO3- forms the ionic compound NH4NO3, (Ammonium Nitrate) when dissolved in water NH4NO3 it behaves as a strong electrolyte. Copyright © Houghton Mifflin Company

Why does oxygen form an O2- ion and not an O3- ion? Focus Questions 12.4 – 12.5 Why do metals lose electrons to form ions? When does a metal stop losing ions? Why does oxygen form an O2- ion and not an O3- ion? Copyright © Houghton Mifflin Company

Write the electron configurations for the pairs of atoms given below: predict the formula for an ionic compound formed from these elements: Mg, S K,Cl Cs, F Ba, Br Copyright © Houghton Mifflin Company

4. Why is aluminum foam useful in making cars more fuel-efficient? Aluminum foam is lighter and stronger than steel. The weight of a car is directly related to its fuel economy. The lighter the car, the more fuel efficient. Copyright © Houghton Mifflin Company

5. Why are cations smaller than their parent atoms 5. Why are cations smaller than their parent atoms? Why are anions larger? 6. How do polyatomic anions differ from simple anions? Copyright © Houghton Mifflin Company

Objective: To write Lewis Structures Copyright © Houghton Mifflin Company

Bonding involves just the valence electrons of atoms. Ionic bonds Valence shell electrons are transferred Covalent Bonds Valence shell electrons are shared. Copyright © Houghton Mifflin Company

Lewis Structure Lewis structures show how the valence electrons are arranged among the atoms in the molecule. The most important requirement for the formation of a stable compound is that the atoms achieve noble gas electron configurations. Copyright © Houghton Mifflin Company

Writing Lewis Structures Include only the valence electrons Use dots to represent an electron. Writing Lewis structures for ionic bond vs covalent bonds (pg. 371-372) When writing Lewis Structures only the valence electrons are considered, that is the core electrons are not shown. Copyright © Houghton Mifflin Company

Rules of Lewis Structures Duet Rule Hydrogen forms stable molecules where it shares two electrons Octet rule The second row of nonmetals (C – F) follow the octet rule – eight electrons are required to fill these orbitals. Bonding pairs – valence electrons Lone pairs/ unshared electrons Copyright © Houghton Mifflin Company

Guild lines for Writing Lewis Structures We must include all the valence electrons from all atoms. The total number of electrons available is the sum of all the valence electrons from all the atoms in the molecule. Atoms that are bonded to each other share one or more pairs of electrons The electrons are arranged so that each atom is surrounded by enough electrons to fill the valence orbital Copyright © Houghton Mifflin Company

Practice pg. 374-375 Copyright © Houghton Mifflin Company

12.7 Lewis Structures of Molecules with Multiple Bonds Objective: To learn how to write Lewis Structures for molecules with multiple bonds Copyright © Houghton Mifflin Company

Writing Lewis Structure is trial and error The total number of valence electrons must be displayed. The octet rule says that eight electrons are required to fill the orbital. To achieve the octet rule single, double and triple bonds occur. Copyright © Houghton Mifflin Company

Bonding Patterns Single bonds Double Bonds Triple Bonds Involves two atoms sharing one electron pair Double Bonds Involves two atoms sharing two pair of electrons Triple Bonds Involves two atoms sharing three pairs of electrons. Copyright © Houghton Mifflin Company

Resonance A molecule shows resonance when more than one Lewis structure can be drawn for the molecule. e.g. CO2 (pg. 376) Start out with single bonds between atoms and add multiple bonds as needed. Example 12.3, 12.4 and Self check 12.4 Copyright © Houghton Mifflin Company

Exceptions to the Octet Rule Boron tends to form compounds in which the boron atom has fewer than eight electrons around it. Beryllium also forms molecules that are electron deficient. Any molecule that contains an odd number of electrons does not conform to our rules for Lewis Structure. Copyright © Houghton Mifflin Company

Why is bonding primarily on the octet rule? Why not a sextet rule? 12.6 -12.7 Focus Questions Why is bonding primarily on the octet rule? Why not a sextet rule? What is the difference between a bonding pair of electrons and a lone pair of electrons? Why do pairs of atoms share pairs (or multiples of pairs) of electrons? Why not share odd numbers? Copyright © Houghton Mifflin Company

12.6 -12.7 Focus Questions Cont., 4. For each molecule a. give the sum of the valence electrons for all atoms. b. draw Lewis structure c. circle the octet (or duet) for each 1. CIF 2. Br2 3. H2O 4. O2 Copyright © Houghton Mifflin Company

Objective: to understand molecular structure and bond angles 12.8 Molecular Structures Objective: to understand molecular structure and bond angles Copyright © Houghton Mifflin Company

Molecular Structure and Geometric Shape The three dimensional arrangement of the atoms in a molecule Notice the shape: it is bent. The angle is approx. 105°. Copyright © Houghton Mifflin Company

The different structures Bent Linear Trigonal planar tetrahedral Copyright © Houghton Mifflin Company

Examples of each shape Linear : 180° C O O Copyright © Houghton Mifflin Company

BF3 : four atoms in the same plane with 120° bond angles Trigonal Planar BF3 : four atoms in the same plane with 120° bond angles F B F F Copyright © Houghton Mifflin Company

Figure 12.12: Molecular structure of methane. Tetrahedral structure Copyright © Houghton Mifflin Company

12.9 Molecular Structure: The VSEPR Model Objective: To learn to predict molecular geometry from the numbers of electron pairs. Copyright © Houghton Mifflin Company

Valence Shell Electron Pair Repulsion model VSEPR Valence Shell Electron Pair Repulsion model Used to predict the molecular structure of molecules formed from NONMETALS. The structures around a given atom is determined by minimizing repulsions between electrons (that is bonding pairs and non-bonding pairs are as far apart as possible). Copyright © Houghton Mifflin Company

Steps for predicting moleluar structure using the VSEPR model Draw the Lewis structure for the molecule Count the electron airs and arrange them in the way that minimizes repulsion (put the pairs as far apart as possible) Determine the positions of the atoms from the way the electron pairs are shared. Determine the name of the molecular structure from the possible position of the atoms. Copyright © Houghton Mifflin Company

LINEAR Arrangements Whenever two pairs of electrons are present around an atom, they should always be placed at an angle of 180° to each other to give a linear arrangement. Copyright © Houghton Mifflin Company

NH3 Ammonia is used as fertilizer and as a household cleaner NH3 Ammonia is used as fertilizer and as a household cleaner . Predict the structure using VSEPR. Copyright © Houghton Mifflin Company

Figure 12.13: Tetrahedral arrangement of electron pairs. NH3 has four pairs of electrons. Three shared pairs and one lone pair. Copyright © Houghton Mifflin Company

Figure 12.13: Hydrogen atoms occupy only three corners of the tetrahedron. The three Hydrogen atoms share electron pairs Copyright © Houghton Mifflin Company

Figure 12.13: The NH3 molecule has the trigonal pyramid structure. The placement of electron pairs determines the structure, but the name is based on the position of the atoms. Copyright © Houghton Mifflin Company

Figure 12.14: Tetrahedral arrangement of four electron pairs around oxygen. Copyright © Houghton Mifflin Company

Figure 12.14: Two electron pairs shared between oxygen and hydrogen atoms. Copyright © Houghton Mifflin Company

Figure 12.14: V-shaped molecular structure of the water molecule. Copyright © Houghton Mifflin Company

2 PAIRS OF ELECTRONS ON A CENTRAL ATOM = 180° or Linear RULES FOR PREDICING 2 PAIRS OF ELECTRONS ON A CENTRAL ATOM = 180° or Linear 3 pairs of electrons around a central atom 120° or trigonal planar 4 pairs of electrons around a central atom is 109.5 ° apart or tetrahedral Unshared electrons change the name (table 12.4) Copyright © Houghton Mifflin Company

Table 12.4 Copyright © Houghton Mifflin Company

12.10 Molecular Structure: Molecules with Double Bonds Objective: to learn to apply VESPR model to molecules with double bonds. Copyright © Houghton Mifflin Company

Focus Questions 12.8 – 12.10 How does drawing a Lewis structure for a molecule help in determining its molecular shape? When is the molecular structure for a molecule the same as the arrangement of the electron pairs? What causes the name for the molecular structure to be different from the name for the arrangement of electron pairs? Copyright © Houghton Mifflin Company

Focus Questions Cont., 4. If double bonds contain four electrons instead of two, why should they be treated as if they were the same as a single bonds when determining molecular structure? Copyright © Houghton Mifflin Company