Unit 11 - Bonding Types of Chemical Bonds Electronegativity Bond Polarity and Dipole Moments Stable Electron Configurations Lewis Structures Lewis Structures with Multiple Bonds Molecular Structure 1

Unit 11 - Bonding Upon completion of this unit, you should be able to do the following: – Predict the types of bonds formed between two atoms and describe the properties of each. – Use electronegativities to predict the percent ionic character of bonds and the polarities of molecules. – Draw Lewis dot structures to represent how atoms share or transfer valence electrons to become more stable. – Explain how multiple bonds can form between the same two atoms. – Explain the concept of resonance when drawing structures and cite examples. – Explain how shared and unshared pairs of electrons determine molecular shape. Predict the shapes and bond angles of simple molecules. Predict molecular polarity from the shape. 2

Types of Chemical Bonds A bond is a force that holds two or more atoms together and makes them function as a unit. In water, the fundamental unit is the H-O-H molecule, which is held together by the two O-H bonds. The strength of a bond is measured by bond energy, the amount of energy required to break the bond. 3

Types of Chemical Bonds Ionic compounds are formed when an atom that loses an electron relatively easily reacts with an atom that accepts an electron. This occurs when a metal reacts with a non-metal. The resulting bonds are called ionic bonds. In an ionic bond, electrons are transferred. 4

Types of Chemical Bonds Consider diatomic hydrogen H – H. When two hydrogen atoms are brought close together, the electrons are equally attracted to both nuclei. When two similar atoms form a bond, the electrons are equally attracted to the nuclei of the two atoms. This is called a covalent bond. In a covalent bond, the electrons are shared by nuclei. 5

Types of Chemical Bonds Ionic bonding and covalent bonding are extremes. Between the extremes are cases where atoms are not so different that electrons are transferred, but different enough that unequal sharing of the electrons results. These bonds are called polar covalent bonds. 6

Types of Chemical Bonds The hydrogen fluoride (HF) molecule contains this type of bond, which produces the following charge distribution. Delta is used to indicate a partial charge. 7

Types of Chemical Bonds 8

Electronegativity The unequal sharing of electrons between two atoms is described by a property called electronegativity, the relative ability of an atom in a molecule to attract shared electrons to itself. The higher the atoms electronegativity value, the closer the shared electrons tend to be to that atom when it forms a bond. Fluorine has the highest electronegativity value at 4.0. Cesium and Francium have the lowest electronegativity value at 0.7 9

10 Electronegativity

The polarity of a bond depends on the difference between the electronegativity values of the atoms forming the bonds. If the atoms have similar electronegativity, values the electrons are shared almost equally and the bond shows little polarity. If the atoms have very different electronegativity values, a very polar bond is formed. 11

Electronegativity In extreme cases, one or more electrons are actually transferred and ions and an ionic bond are formed. Consider NaCl, for example. When a Group 1 metal reacts with a Group 17 element, ions are formed and an ionic substance results. Page 321, example 11.1 Using the electronegativity values given in Figure 11.3, arrange the following bonds in order of increasing polarity: H-H, O-H, Cl-H, S-H, F-H 12

Electronegativity Homework: read pages

Bond Polarity and Dipole Moments As we saw in HF, the molecule has a center of positive charge and a center of negative charge. This is called a dipole moment. The dipolar character of the molecule is often represented by an arrow, pointing towards the negative charge. The tail indicates the center of positive charge. 14

Bond Polarity and Dipole Moments Any diatomic molecule that has a polar bond has a dipole moment. Some polyatomic molecules also have a dipole moment. Because the oxygen in water has a greater electronegativity than the hydrogen atom, the electrons are not shared equally. This results in a charge distribution that causes the molecule to behave as if it has two centers of charge – one positive and one negative. So the water molecule has a dipole moment. 15

Stable Electron Configurations 16 Representative metals form ions by losing enough electrons to attain the configuration of the previous noble gas that occurs before the metal. For example, sodium will lose one electron to attain the configuration of neon. Nonmetals form ions by gaining enough electrons to attain the configuration of the next noble gas. For example, chlorine will add one electron to attain the configuration of argon.

Stable Electron Configurations 17

Stable Electron Configurations 18 In observing millions of stable compounds, chemists have observed that in almost all stable compounds, all of the atoms have achieved a noble gas configuration.

Stable Electron Configurations 19 When a non-metal and a Group 1, 2 or 3 metal react to form a binary ionic bond, the ions form so that the non-metal completes the valence-electron configuration of the next noble gas and the metal empties the valence orbitals to achieve the configuration of the previous noble gas.

Stable Electron Configurations 20 When two non-metals react to form a covalent bond, they share electrons in a way that completes the valence-electron configuration of both atoms.

Predicting Formulas of Ionic Compounds 21 Consider calcium and oxygen Ca [Ar] 4s 2 O [He] 2s 2 2p 4 Electronegativity of O is 3.5, Ca is 1.0 so electrons transfer to O. O + 2e - → O 2- [He] 2s 2 2p 4 + 2e - → [He] 2s 2 2p 6 or [Ne] Ca → Ca e - [Ar] 4s 2 → [Ar] + 2e -

Stable Electron Configurations 22

Lewis Structures 23 Bonding involves just the valence electrons of atoms. The Lewis structure is a representation of a molecule that shows how the valence electrons are arranged among the atoms in the molecule. Dots are used to represent the valence electrons. Examples: KBr

Lewis Structures 24 Duet rule – hydrogen Octet rule – elements that are surrounded by 8 electrons. Electrons that are shared form bonds. See F 2 example in Lab 17.

Lewis Structures 25 In class practice Draw the Lewis structure for the water molecule. Draw the Lewis structure for HCl. Draw the Lewis structures for C 2 H 6, C 2 H 4 and C 2 H 2.

Lewis Structures 26 A single bond involves two atoms sharing one electron pair. A double bond involves two atoms sharing two electron pairs. A triple bond involves two atoms sharing three electron pairs.

Lewis Structures 27 A molecule shows resonance if more than one Lewis structure can be drawn. Consider CO 2. It has 16 valence electrons. How can we arrange 16 electrons so that all three atoms have satisfy the octet rule. You can also have two other Lewis structures

Lewis Structures 28 Consider CO 2. It has 16 valence electrons. How can we arrange 16 electrons so that all three atoms have satisfy the octet rule. You can also have two other Lewis structures A molecule shows resonance if more than one Lewis structure can be drawn for the molecule. HW – page 352; problems 58, 60, 62 and 64

Molecular Structures and VSEPR 29 LinearCO 2 BentH 2 O Trigonal planarBF 3 TetrahedralCH 4 Trigonal pyramidNH 3

Molecular Structures and VSEPR 30 Bond angle is often specified to more precisely describe the molecular structure. See below. – Linear = 180 o – Trigonal planar = 120 o – Tetrahedral = o – Trigonal pyramid = 107 o

Valence Shell Electron Pair Repulsion (VSEPR) 31 Rules for Using VSEPR Model 1.Two pairs of electrons on a central atom in a molecule are always placed 180 o apart. This is a linear arrangement of pairs. 2.Three pairs of electrons on a central atom in a molecule are always placed 120 apart in the same plane as the central atom. This is a trigonal planar arrangement of the pairs. 3.Four pairs of electrons on a central atom are always placed apart. This is a tetrahedral arrangement of electron pairs.

Valence Shell Electron Pair Repulsion (VSEPR) 32 Rules for Using VSEPR Model 4.When every pair of electrons on the central atom is shared with another atom, the molecular structure has the same name as the arrangement of electron pairs. 5.When one or more of the electron pairs around a central atom are unshared (lone pairs), the name for the molecular structure is different from that of the arrangement of electron pairs.