Bonding Notes KAP KEY (1) PDF

## Chemistry KAP Notes: Bonding ### I. Chemical Bonding Overview - **A. Definition** - A chemical bond is a force of attraction that holds atoms or ions together to form a compound. The attractive force is between the nucleus of one atom/ion and the electron cloud of another atom/ion. - **B. Pu...

## Chemistry KAP Notes: Bonding ### I. Chemical Bonding Overview - **A. Definition** - A chemical bond is a force of attraction that holds atoms or ions together to form a compound. The attractive force is between the nucleus of one atom/ion and the electron cloud of another atom/ion. - **B. Purpose of Bonding** - Chemical bonding is driven by the octet rule: atoms will lose or gain valence electrons in order to achieve the electron configuration of the closest noble gas. - Unfilled or partially filled valence orbitals are inherently unstable. ("Unstable" means possessing high potential energy.) Atoms will bond to reduce their potential energy, thereby becoming more stable. - **C. Types of Chemical Bonds** - There are 3 types of chemical bonds: ionic, covalent, and metallic. - Bond type is determined by valence electrons of participating atoms. - **Ionic** bonds involve the transfer of electrons (lose/gain) between a metal and a nonmetal. - **Covalent** bonds involve the sharing of electrons between nonmetals/metalloids. - **Metallic** bonds involve a mobile "electron sea" between metal atoms. ### II. Ionic Bonding - **A. Formation of Ions** - An atom that loses electrons to achieve an octet will form a cation. - Ex: Sodium will lose 1 valence electron to form the Na+ ion. - An atom that gains electrons to achieve an octet will form an anion. - Ex: Chlorine will gain 1 valence electron to form the Cl- ion. - **B. Formation of Ionic Compounds** - Involve a metal ion and a nonmetal ion; may also involve polyatomic ions. - Anions and cations are attracted to each other by electrostatic forces, also known as Coulombic forces. These forces of attraction between oppositely-charged ions are called ionic bonds. - Ex: Formation of sodium chloride: Na+ Cl- = NaCl = NaCl - Ex: Formation of aluminum bromide: Al3+ Br- Br- Br- = Al3+ Br- Br- Br- = AlBr3 - **C. Properties of Ionic Compounds** - A lattice is a 3-D system of points showing the positions of the ions that make up the ionic compound. - The electrostatic (Coulombic) forces between the ions are very strong. Therefore, ionic compounds: - are found in nature as crystalline solids. - have high melting and boiling points. - are hard, brittle, and rigid. - Ionic compounds will conduct electricity when molten or aqueous (dissolved in water). - **Electrolyte:** substance that conducts electricity when molten or in solution ### III. Metallic Bonds - **A. Metallic Bonds:** - Metals characteristically have few electrons in their highest energy levels. - Metals therefore have vacant orbitals (typically d and/or p block orbitals). - The vacant orbitals can be occupied, allowing valence electrons to travel freely throughout the metal. This is described as an "electron sea." - The electrons are delocalized, meaning they do not belong to any one nucleus. - **Metallic** bonding is the chemical bonding that results from the attraction between the nuclei of metal atoms and the surrounding "sea" of electrons. - **B. Properties of Metals:** - The freedom of electrons to roam in metals accounts for the high thermal (heat) and electrical conductivity of metals. (moving charges again). - Metals exhibit a lustrous (shiny) appearance because they can absorb a wide range of light frequencies. - Metals are malleable (they can be hammered into thin sheets). - Metals are ductile (they can be pulled into a wire). ### IV. Nature of Covalent Bonds - **A. Definition/Properties:** - In the covalent bond, electrons are shared by the bonding atoms to achieve octets for the atoms. (valence). - Compounds that consist of covalently bonded atoms are molecular compounds. - The smallest particle of a molecular compound is called a molecule. - The chemical formula for a molecular compound is the molecular formula. - **Nonmetal** elements (and sometimes metalloids) tend to form covalent bonds. - **B. Types of Covalent Bonds:** - **Single** covalent bond - involves 1 shared pair of electrons - Ex: H2O, H-O-H or H:O:H - **Double** covalent bond - involves 2 shared pairs of electrons - Ex: CO2, O=C=O or O::C::O - **Triple** covalent bond - involves 3 shared pairs of electrons - Ex: N2, :N≡N: or :N:::N: ### V. Lewis Structures - **A. Creating Lewis Structures** - Lewis structures are depictions of molecules that show valence electrons as dots. - Shared pairs of electrons (i.e. bonds) are drawn between the atoms sharing them. - Unshared or lone pairs of electrons are represented by dots located on one atom only. - **Drawing Lewis structures for molecular compounds:** 1. Find the "needed" number of electrons (N) for each atom in the compound and add them up. N will generally be 8 for each atom, with the following common exceptions: H=2, Be=4, B=6 2. Find the "available" number of electrons (A) for each atom in the compound and add them up. A is the number of valence electrons for each atom. 3. If the substance is a polyatomic ion, adjust A by using the ion's charge. For example, in nitrate, you would add 1 electron to A because of the -1 charge. In ammonium, you would subtract 1 electron from A because of the +1 charge. 4. Find the "shared" electrons (S) by this formula: S = N - A 5. Draw a "skeleton" of the molecule. In most cases, there is one "different" atom. Put that one in the middle and surround it by the others. 6. Place the shared (S) electrons in between atoms. 7. The S electrons are part of the A electrons. Figure out how many more electrons you need to add with this formula: E = A - S. Place those as unshared pairs in order to give each atom an octet. 8. Check: is the total # of dots in the structure = to A? Does every atom have an octet? - **Examples of Lewis structures:** - **Methane (CH4):** - N = 8 + 4(2) = 16 - A = 4 + 4(1) = 8 - S = 16-8 = 8 - E = 8-8 = 0 - **Phosphorus trichloride (PCl3):** - N = 4(8) = 32 - A = 5 + 3(7) = 26 - S = 32-26 = 6 - E = 26-6 = 20 - **Carbonate ion (CO32-):** - N = 4(8) = 32 - A = 4 + 3(6) + 2 = 24 - S = 32-24 = 8 - E = 24-8 = 16 - **Hydrocyanic acid (HCN):** - N = 2 + 2(8) = 18 - A = 1 + 4 + 5 = 10 - S = 18 - 10 = 8 - E = 10 - 8 = 2 - **B. Resonance Structures:** - Resonance structures are structures that occur when it is possible to write 2 or more valid Lewis structures for the same molecule or ion. - Ex: ozone, O3 - Experimental data indicate that the 2 bonds in ozone are the same length, BUT double bonds are shorter than single bonds! So the explanation for this is that the actual bonds are hybrids of those in the 2 resonance structures. The extra electron pair in ozone is delocalized over the two bonding regions. So each bond spends about half the time being single and half being double. - **C. Lewis Structure vs. Structural Formulas:** - **Lewis dot diagram:** shows the pairs of shared electrons as dots between two atoms. There are no lines used; only dots to represent valence electrons. - **Structural formula:** shows the pairs of shared electrons with a straight line to represent each shared pair of electrons. Lone pairs are still depicted as dots. ### VI. VSEPR Theory and Molecular Geometry (shape) - The Valence Shell Electron Pair Repulsion Theory explains the three-dimensional shape of molecules. - The VSEPR theory states that because electron pairs repel each other, the molecular shape results from valence electron pairs positioning themselves as far apart as possible. - **Examples:** - **Methane, CH4**: bond angles 109.5°, not 90°. - **H2O**: has 104.5° bond angle - **NH3**: has 107° bond angles - **Use the Lewis structure for the molecule and the chart on the next page to predict molecular geometry (shape).** ### VII. Polarity of Bonds - **A. Determination of Bond Type** - **To determine type of bond: (not applicable for metallic bonding*)** - **Remember that electronegativity is the measure of an atom's ability to attract bonding electrons. See chart of values below.** - 1) Determine the difference in electronegativity values between the 2 atoms (abs value). - 2) If the difference is: - >1.70: Ionic bond (electrons transferred) - 1.70-0.36: Polar Covalent bond (electrons shared unequally) - <0.36: Nonpolar Covalent bond (electrons shared equally) - *Bonds between 2 metals are always classified as metallic. - **Electronegativities of Some Elements:** (table of elements with their electronegativity values) - **Summary** - A large electronegativity difference leads to an ionic bond. - A moderate electronegativity difference leads to a polar covalent bond. - Little to no electronegativity difference between two atoms leads to a nonpolar covalent bond. - **Examples:** - **Cesium and sulfur:** EN = 1.79, ionic - **Carbon and oxygen:** EN = 0.89, PC - **Carbon and hydrogen:** EN = 0.35, NPC - **Fluorine and fluorine:** EN = 0, NPC - **Zinc and copper (in the alloy brass):** Metallic - **B. Polar vs. Nonpolar Covalent Bonds:** - The bonding pairs of electrons in covalent bonds are located between the nuclei of the atoms sharing the electrons. - When the pull of each nucleus for the electrons is equally strong, the electrons are shared equally and are located on average halfway between the two nuclei. This type of bond is a **nonpolar** covalent bond. - When the pull of one nucleus for the electrons is stronger than the other, the electrons spend more time closer to the more electronegative nucleus. This type of bond is a **polar** covalent bond. - The more electronegative atom acquires a partial **negative** charge. The less electronegative atom therefore acquires a partial **positive** charge. - These partial charges are indicated by the following symbols: - **δ -** for the more electronegative element. - **δ +** for the less electronegative element. - Ex: H2O - The water molecule contains two dipole moments, sometimes just called **dipoles**. The dipole moments can be drawn in with arrow notation. ### VIII. Polarity of Molecules - Note that just because a molecule contains polar bonds, it may or may not be classed as a polar molecule. - The polarity of a molecule will depend on: - **(1) existence of polar bonds** (if none of the bonds are polar, the molecule is nonpolar*) - **(2) shape of the molecule** - **(3) orientation of the polar bonds** - *there are exceptions to this (such as O3), but you will not be tested over them* - **To help determine if the molecule will be polar, look for lines of symmetry. If the molecule is symmetrical in its 3-D shape, then the dipoles will cancel and the molecule is nonpolar. If the molecule is asymmetrical and therefore the dipoles do not cancel, the molecule is polar. Note that water is polar because its dipoles do not cancel.** - **Helpful hint: The following molecular geometries are symmetrical and have a high probability of leading to nonpolar molecules, assuming equivalent dipoles around the central atom:** - Tetrahedral - Linear (triatomic) - Trigonal planar - **For the following examples, (1) determine the molecular geometry; (2) determine the bond polarities and draw in any dipoles that you find; (3) classify the molecule as overall polar or nonpolar. ** - **Examples:** - **HCl:** linear (diatomic) , polar - **CO2:** linear (triatomic), nonpolar - **CF4:** tetrahedral, nonpolar - **NH3:** trigonal pyramidal, polar - **Chart of Molecular Geometry (shape)** (table with information on the number of atoms bonded to the central atom, the number of lone pairs on the central atom, the molecular geometry, the general form and an example)