Chemical Bonds PDF

## 4.1 Molecular Forces **Molecular Forces** are the bonding forces **within** or **between** molecules. ### Intramolecular Forces * Intramolecular Forces are the bonds that hold the atoms or ions together **within** the substance and make it up. * They are also called **Chemical Bonds**. * They de...

## 4.1 Molecular Forces **Molecular Forces** are the bonding forces **within** or **between** molecules. ### Intramolecular Forces * Intramolecular Forces are the bonds that hold the atoms or ions together **within** the substance and make it up. * They are also called **Chemical Bonds**. * They determine the chemical properties of substances. * Strong forces * Types: * Ionic bonds * Covalent bonds * Metallic bonds ### Intermolecular Forces * Intermolecular Forces are forces or bonds **between** the particles (atoms, molecules, or ions) of a substance. * They are the forces that hold the particles together in liquids and solids. * They determine the physical properties of substances. * Weak relative to the intramolecular forces * Types: * Dipole-dipole Forces * Ion-Dipole Forces * Hydrogen Bonding * London Dispersion Forces ## 4.2 The Lewis Theory of Chemical Bonding ### An Overview * The attractive forces that hold atoms together in compounds are called **chemical bonds**. ### Why do atoms bond together? * Atoms bond together because the resultant compound is more stable and has lower energy than the separate atoms. * **Bond formation** is an **exothermic reaction** (making bonds always releases energy). Therefore, energy (usually as heat) is always liberated and flows out of the chemical system when a bond forms. * **Breaking bonds** always **absorbs energy** (endothermic reaction). Therefore, energy must be put into the system to break a bond. * **Valence electrons** play a fundamental role in chemical bonding. ### In forming chemical bonds, representative elements apply the **octet rule** or **duet rule**: * **Octet Rule**: Atoms lose, gain, or share their valence electrons until they have 8 electrons in their outer shell and take on the same electron configuration of the nearest noble gas (ns²np⁶). * **Duet Rule**: Some atoms lose, gain, or share their valence electrons until they have 2 in the outer shell and take on the same electron configuration of Helium (He: 1s²). Examples include Hydrogen (H), Lithium (Li), and Beryllium (Be). ## 4.3 Types of chemical Bonding Several types of chemical bonds hold atoms together; three will be discussed: ionic, covalent, and metallic bonds. ### 4.3.1 Ionic Bonds * When atoms lose or gain electrons, they tend to acquire the electron configuration of the nearest noble gas in the periodic table. * When **metals** and **nonmetals** react together, valence electrons are transferred from the metal to the nonmetal atoms forming positive and negative ions. * The electrostatic attraction between the oppositely charged ions gives rise to the **ionic bonds**. * Ionic bonds mostly form **ionic compounds**. **Example:** * The formation of sodium chloride salt by the reaction between sodium metal and chlorine gas [Na] Loses one electron [Na]⁺ + e⁻ (cation) :Cl: Gains one electron :Cl:⁻ (anion) The oppositely charged ions are then electrostatically attracted to each other, forming the ionic bond. ### Characteristics of ionic compounds: * Polar * Have high melting and boiling points * Often soluble in water (Hydrophilic; water lover) * Insoluble in nonpolar solvents like chloroform and benzene. * Do not conduct electricity as solids, but they conduct when molten or dissolved in water. ### 4.3.2 Covalent Bonds * Covalent bonds link nonmetal atoms to form covalent compounds. It is formed by the sharing of valence electrons between the bonded atoms so that each atom acquires the electron configuration of the nearest noble gas. * The simplest way of indicating covalent bonds in molecules is the use of **line-bond structures (Kekulé structures)**, in which a two-electron covalent bond is indicated as a line drawn between atoms. * The number of covalent bonds an atom forms depends on how many additional valence electrons it needs to reach a noble-gas configuration. * **Hydrogen** has one valence electron (1s) and needs one more to reach the helium configuration (1s²), so it forms one bond. * **Carbon** has four valence electrons (2s²2p²) and needs four more to reach the neon configuration (2s²2p⁶), so it forms four bonds. * **Nitrogen** has five valence electrons (2s²2p³), needs 3 more, and forms three bonds. * **Oxygen** has six valence electrons (2s²2p⁴), needs two more, and forms two bonds. * **The halogens** have seven valence electrons, need one more, and form one bond. * Covalent bonding can be **single-covalent**, **double-covalent**, or **triple-covalent** depending on the number of pairs of electrons shared. #### Lone-pair electrons (nonbonding electrons) * Valence electrons that are not used for bonding are called **lone-pair electrons**, or **nonbonding electrons**. ### Electronegativity Difference and Bond Type * The electronegativity difference (AEN) of two atoms determines their bond type: Ionic or Covalent. Covalent bonds can be either polar or nonpolar. * **Nonpolar Covalent Bond:** Non-metal atoms have no or little difference in electronegativity (EN). Electron pairs are shared equally. (ΔEN ≤ 0.5) * **Polar Covalent Bond:** Non-metal atoms have a significant difference in electronegativity (EN). Electron pairs are shared unequally. They are drawn closer to the atom with higher EN. (0.5 < ΔEN < 1.7) * **Ionic Bonds:** Metal and non-metal atoms with larger differences in electronegativity (EN). Electrons are completely transferred from metal to nonmetal atoms. (ΔEN > 1.7) ### 4.3.3 Metallic Bonds * Metallic bond occurs among METAL atoms. * In a metal, the atoms are packed closely together in a regular arrangement called a lattice. * Each atom of the metal crystal loses all valence electrons forming a pool of electrons (sea of electrons) in which positively charged metal ions are held together by electrostatic attraction of the pool electrons and the positive ions. * The outer shell electrons are free to move throughout the metal lattice, (delocalized electrons or mobile electrons). They are not associated with any one particular atom or bond. * Metallic bonding is strong because the ions are held together by the strong electrostatic attraction between their positive charges and the negative charges of the delocalized electrons. This electrostatic attraction acts in all directions. * The strength of metallic bonding increases with: * Increasing the number of mobile electrons per atom. * Increasing positive charge on the ions in the metal lattice * Decreasing the size of metal ions in the lattice ## 4.4 Intermolecular Forces * Types of intermolecular forces include: * Dipole-dipole Forces * Ion-Dipole Forces * Dispersion Forces * Hydrogen Bonding ### Study Tips **Intermolecular Forces and Boiling Point:** * The greater the intermolecular force between the particles of a substance, the higher the boiling point of that substance. * This is because more heat energy is required to overcome the intermolecular forces to change the liquid into a gas. ### 4.4.1 Dipole-Dipole Interactions * Permanent dipole-dipole interactions occur between polar covalent molecules. * A polar covalent molecule has a partially positively charged end (δ+) and a partially negatively charged end (δ-). * When polar covalent molecules come close to one another, the opposite charges of the dipoles tend to attract one another. * The polarity of molecules depends on the difference in electronegativity (EN) between bonded atoms. * The more polar a molecule, the more its dipole-dipole forces. * Thus, polar covalent molecules tend to have higher boiling points (B.P) than non-polar species of the same molecular weight. As the molecule polarity increases, the boiling point temperature increases. **Example:** * The boiling point of polar CO (M.wt=28) (-192°C) is four degrees higher than that of non-polar N₂ (M.wt=28) (-196 °C). ### 4.4.2 Ion-Dipole Forces * Occurs between ions and polar solvents. * For example, when an ionic compound like table salt (NaCl) is dissolved in a polar solvent like water. * Ion-dipole force results from the electrostatic attraction between an ion and a neutral molecule that has a dipole (polar covalent compound). * A positive ion (cation) attracts the partially negative end of a neutral polar molecule. * A negative ion (anion) attracts the partially positive end of a neutral polar molecule. ### 4.4.3 Hydrogen Bonds * A hydrogen bond occurs between partially positive hydrogen (H) of one molecule and partially negative electronegative atom of another molecule, mainly Nitrogen, Oxygen, and Fluorine atoms (NO. F) * **Note:** The electronegativities of Nitrogen, Oxygen, and Fluorine are 3, 3.5, and 4 respectively. #### Types of Hydrogen Bonding: * Hydrogen bonds are of two types, Intermolecular and Intramolecular. * **1. Intermolecular hydrogen bonds (Between different molecules):** * The classic example of intermolecular hydrogen bonds is the hydrogen bonds between water molecules. * Each water molecule can form four hydrogen bonds with the other four water molecules. * Water has a high boiling point temperature, with respect to its tiny mass, because of the intermolecular hydrogen bonds between its molecules. * Other Examples of Intermolecular Hydrogen Bonds include: * Intermolecular Hydrogen Bonds in Methanol * Intermolecular Hydrogen Bonds in Acetic Acid * Intermolecular Hydrogen Bonds in Water * **2. Intramolecular hydrogen bonds (Inside the same molecule):** * The intramolecular hydrogen bond is formed between a hydrogen atom and a highly electronegative atom (N, O, or F) inside the same molecule. * Examples of Intramolecular hydrogen bonds * o-Nitrophenol * Salicylic acid * Salicylaldehyde ### 4.4.4 Dispersion Forces * Dispersion forces are also known as London forces Dipole-induced dipole forces, or Van der Waals forces. * They are the weakest intermolecular attractive forces that are important only over extremely short distances. * They are the only type of forces present in non-polar molecules such as CO₂, O₂, N₂, Cl₂, Br₂, H₂ and monatomic species such as noble gases. London forces are also responsible for the condensation of these substances. * **Dispersion forces result** from the attraction of the positively charged nucleus of one atom for the electron cloud of an atom in the nearby molecules. This induces a temporary dipole between the atoms or molecules. The Temporary (momentary) dipole in one molecule INDUCES a similar dipole in a neighboring molecule. Thus, the resulting temporary dipoles cause weak attraction among the molecules. ### Strength of Dispersion Forces * London dispersion force is the weakest intermolecular force. However, its strength depends on **molecular polarizability**; the greater the polarizability of molecules, the stronger the dispersion forces between them. #### Polarizability * Polarizability increases with increasing numbers of electrons and sizes of molecules. * Polarizability in turn depends on following two factors: * **1. Molecular Shape or Surface area:** * The elongated molecule is more easily polarized than the compact molecule. (See image on Page 16) * **2. Molecular Size or Weight:** * Greater molecular size or weight causes larger, more polarizable electron clouds. ### Strength Order of Intermolecular Forces * Among the four types of intermolecular forces mentioned before, the ion-dipole force is the strongest. But, of course, it is a weak force compared to intramolecular forces (ionic and covalent bonds). The order of strength from strongest to weakest is: 1. Ion - Dipole forces 2. Hydrogen Bonds 3. Dipole - Dipole forces 4. London forces ## Exercise and Practice 1. A covalent bond is formed by the .............. * a. Transfer of electrons from the metal to non-metal. * b. Sharing of electrons between non-metal * **c. Transfer of electrons from non-metal to another non-metal** * d. Sharing of electrons between metals. 2. In a covalent compound, the bonding is between * a. metals and non-metals * b. non-metals and non-metal * **c. metals and metals** * d. aqueous solutions 3. A triple covalent bond is present in * a. O₂ * **b. N₂** * c. H₂ * d. Cl₂ 4. What is the bond type between the atoms of chlorine molecules? * a. Ionic bond * b. Hydrogen bond * c. Polar covalent bond * **d. Non-polar covalent bond** 5. Which of the following bonds is characterized by unequal sharing of the electron pairs? * a. lonic bonds * b. Nonpolar covalent bonds * **c. Polar covalent bonds** * d. Hydrogen bonds. 6. Which of the following is the weakest intermolecular attractive force? * a. Ion dipole forces * b. Hydrogen bond * c. Dipole-dipole force * **d. Dispersion forces** 7. What is the type of intermolecular forces in an aqueous solution of NaCl? * a. Dispersion forces * **b. Ion-dipole forces** * c. Hydrogen bonds * d. Dipole-dipole forces ## Homework Activity 1. Calculate the electronegativity values for the following bonds, and then name the type of bond. * (Electronegativities: Br = 2.8, Cl = 3.0, F = 4.0, H = 2.1 and I = 2.5) | Bonds | ΔEN | Bond Type | |---|---|---| | Br-F | 4.0 – 2.8 = 1.2 | Polar Covalent | | Cl-Cl | 3.0 – 3.0 = 0 | Non-Polar Covalent | | H-Cl | 3.0 – 2.1 = 0.9 | Polar Covalent | | I-Cl | 3.0 – 2.5 = 0.5 | Non-Polar Covalent | 2. What is the ionic bond, and how is it formed? Give examples. 3. Explain how covalent bonds are formed. 4. Mention two examples of molecules having nonpolar covalent bonds. 5. Mention two examples of compounds having polar covalent bonds. 6. What are the weakest intermolecular forces? 7. What are the strongest intermolecular forces? 8. What is the difference between lone pair electrons and the shared bond pair? 9. What type of bond is formed between Chlorine (Cl) and Potassium (K)? Electronegativities of Chlorine and Potassium are 3 and 0.8, respectively. 10. What are the main intermolecular forces that work in the following? * i. Water molecules (H₂O): Hydrogen bonding. * ii. aqueous sodium chloride (NaCl) solution: Ion-dipole interactions. * iii. Non-polar substances like CO₂, Br₂, and He: London dispersion forces * iv. Gaseous hydrogen chloride (HCl). Dipole-dipole interactions.