AP Chemistry Chapter 4: Chemical Bonding 2024-2025 PDF

Chapter 4: Chemical Bonding AP Chemistry 2024-2025 Types of Bonding  Why Bond? Atoms bond to reach their lowest energy, most stable state, which is when they have a complete valence shell of electrons. All atoms, except the very smallest (ex: H), tend to lose, g...

Chapter 4: Chemical Bonding AP Chemistry 2024-2025 Types of Bonding  Why Bond? Atoms bond to reach their lowest energy, most stable state, which is when they have a complete valence shell of electrons. All atoms, except the very smallest (ex: H), tend to lose, gain, or share electrons so that each atom has eight valence electrons. ○ Known as the octet rule. These 8 valence electrons are 2 electrons in the outermost s-subshell  Ionic Bonding Ionic Bond: a bond resulting from the electrostatic attraction of a cation to an anion. ○ Forms between a metal and a nonmetal. ○ Electrons are transferred from the metal to the nonmetal. ○ The metal loses electrons while the nonmetal gains the electrons. ○ Involves elements on opposite sides of the periodic table. Ionic bonding results in the formation of solids with a crystal lattice structure  Electrostatic Potential Energy Electrostatic Potential Energy (Eel): the energy a charged particle has because of its position relative to another charged particle. ○ Also known as coulombic attraction. Eel ∝ Q1Q2 / d The strength of attraction increases with increasing charge of the ions. The strength of attraction increasing with decreasing size (radius) of the ions.  Lattice Energy Lattice Energy: the energy released when one mole of an ionic compound forms from its free ions in the gas phase. Can determine from the Eel equation. You will NOT have to calculate lattice energy, but you do need to know how charge and radius affect lattice energy.  Practice Rank the following ionic compounds in order of increasing coulombic attraction between their ions: KBr, SrBr2, and CsBr  Covalent Bonding Covalent Bond: a bond created by two atoms sharing one or more pairs of electrons. ○ Forms between two nonmetals. ○ Electrons are shared between two atoms. Potential Energy Diagram As two atoms approach each other, there is an attraction between one atom’s nucleus and the electrons of the other. This causes a decrease in potential energy. At the distance of lowest potential energy, the attraction between each nucleus and other atom’s electrons are equal. As we bring the atom’s even closer, we see potential energy increases drastically due to  Potential Energy Diagram Bond Length: the distance between the nuclei of two atoms joined in a bond. ○ Can determine this from the diagram by finding the distance at the lowest potential energy. Bond Energy (bond strength): the energy needed to break one mole of a specific covalent bond in the gas phase. ○ This is the minimum potential  Metallic Bonding Metallic Bond: a bond consisting the nuclei of metal atoms surrounded by a “sea” of mobile shared electrons. ○ Formed between metals ○ The electrons are “delocalized” - not associated with one particular atom or covalent bond. ○ What leads to a metal’s ability to be a good  Electronegativity In Covalent Bonds, valence electrons are shared between two atoms in the bond. But, the electrons are not always shared equally. Electronegativity: a relative measure of an atom’s ability to attract electrons to itself within a bond. ○ This is another periodic trend, which can be explained by Coulomb’s  Electronegativity General Trend: Electronegativity increases from left to right across a period, and decreases as you move down a group. Across a Period: Number of protons increases (Zeff) which causes a greater Coulombic force of attraction between the nucleus and electron. Down a Group: Distance of the valence electrons from the nucleus increases (increasing n), leading to a  Electronegativity & Bond Polarity Fluorine is the most electronegative element. The unequal sharing of electrons causes there to be polarity within the bond through the formation of a dipole. ○ Dipole: a pair of oppositely charged poles separated by a distance. ○ Leads to the formation of partial negative and partial positive charges on the atoms.  Bond Polarity Nonpolar Covalent Bond: a bond characterized by an even distribution of charge; the two atoms equally share the electrons in the bond. ○ This is when we have two of the same atom with the same electronegativity. ○ Bond in Cl2 Polar Covalent Bond: A bond resulting from unequal sharing of pairs of electrons between atoms. ○ Due to an electronegativity difference between the two atoms. ○ Leads to the formation of partial positive and negative charges.  Bond Polarity Practice Problem 4.34: Rank the following bonds from nonpolar to most polar: H–H, H–F, H–Cl, H–Br, H–I  Bond Polarity Practice (Nonpolar) H–H < H–I < H–Br < H–Cl < H–F (Most Polar) 0 < 0.4 < 0.7 < 0.9 < 1.9 Given electronegativity values, you should be able to determine the electronegativity difference to rank polarity. Without electronegativity values, you should be able to use the trends in electronegativity to help you justify differences in polarity. ○ Fluorine is the most electronegative element, so it will cause the largest electronegativity difference. ○ The attached atoms are all halogens, and we know that electronegativity values get smaller as we move down the group. ○ Hydrogen bound to itself will have an electronegativity value of  Chemical Nomenclatu re  Naming Binary Molecular (Covalent) Compounds We use these steps for compounds formed between two nonmetals. Example: Write the name for CCl4 1. Write the names of both elements in the formula. Carbon Chlorine 2. Change the ending of the second element name to -ide. Carbon Chloride 3. Add prefixes to note how many of each type of atom are in the compound. a. The only time a prefix is not used is for mono- on the first element name.  Writing Formulas for Binary Molecular Compounds We use these steps for compounds from two nonmetals. Example: Write the formula for Carbon Tetrachloride 1. Write the element symbol for both elements in the name. CCl 2. If the name of the element has a prefix, include the number of the prefix as a subscript. CCl4  Naming Binary Ionic Compounds We use these steps for compounds formed between a metal and a nonmetal. (opposite sides of the periodic table) Example: Write the name for CaCl2 1. From the formula, write the name of each element. Calcium Chlorine 2. Change the ending of the second element to -ide. a. Note, the metal (cation) is always written first, followed by the nonmetal (anion). Calcium Chloride  Writing Formulas for Binary Ionic Compounds We use these steps for compounds formed between a metal and a nonmetal (opposite sides of the periodic table). Example: Write the Formula for Calcium Chloride 1. Write the atomic symbol for each element in the name. a. Make sure the metal is first, nonmetal is second. CaCl 2. Write the charge for each atom based on the ion they form. Ca2+Cl1- 3. Cross the charges to balance the electrons in the formula.  Ionic Compounds With Transition Metals Many transition metals form multiple cations with different charges. To denote this, we include Roman numerals in the name to symbolize the charge of the cation. Ex: Iron (III) Oxide Lets us know iron has a charge of 3+. Fe3+ O2- Fe2O3 Note: the roman numeral does not tell me the subscript of Iron. To write the name for Fe2O3, you need to determine the charge on the Iron atom and use it as a roman numeral in the name. Since we have three Oxygen atoms, we have a gain of 6 electrons. With only two Fe atoms, each had to have lost 3 electrons (forming a charge of 3+). Inclusion of Polyatomic Ions Polyatomic ions are a charged group of atoms joined by covalent bonds. We will write names/formulas of compounds with polyatomic ions similarly to those of ionic compounds, with a few exceptions. We do not change the ending of the name of the polyatomic ion. If we have to put a subscript on the polyatomic ion in the chemical formula, we must add parentheses around the ion and include the subscript outside of these parentheses. Example: NaNO3 → sodium nitrate Example: Calcium Phosphate → Ca 3(PO4)2  Naming Binary Acids For binary acids (typically H with a halogen in an aqueous solution), Name the acid HCl(aq) 1. Add the prefix hydro- to the front of the second element in the formula. Hydrochlorine 2. Change the last syllable in the name of the second element to -ic and add the word acid. Hydrochloric acid You should know: HF - hydrofluoric acid HCl - hydrochloric acid HBr - hydrobromic acid  Naming of Oxoanions Oxoanions are polyatomic anions that contain at least one non-oxygen central atom bonded to one or more oxygen atoms. For elements that form two oxoanions, the substance with more oxygen ends with -ate and with less oxygen ends with -ite. Ex: NO3 is nitrate; NO2 is nitrite For elements that form more than two oxoanions: - The substance with the most oxygen has the prefix per- and ends with -ate. - The substance with the least oxygen has the prefix  Naming of Oxoacids Oxoacid: oxoanions that are bonded to at least one hydrogen ion (H +). If ending of an oxoanion name is -ate, the name of the oxoacid ends with -ic acid. If ending of an oxoanion name is -ite, the name of the oxoacid ends with -ous acid.  Practice Name: - S2O4 - CuO - K3PO4 - Na2O2 - HNO2 Write the Formula: - Calcium oxide - Sulfur dioxide - Ammonium Sulfate - Nickel (II) hydroxide - Hydrochloric Acid  Practice Solutions Name: - S2O4 Disulfur Tetroxide - CuO Copper (II) Oxide - K3PO4 Potassium Phosphate - Na2O2 Sodium Peroxide - HNO2 Nitrous Acid Write the Formula: - Calcium oxide CaO - Sulfur dioxide SO2  Lewis Structures  Recall: When it comes to chemical bonding, atoms follow the octet rule: The tendency of atoms of main group elements to make bonds by gaining, losing, or sharing electrons to achieve a valence shell containing eight electrons, or four electron pairs.  Lewis Structures Lewis Structure: A two-dimensional representation of the bonds and lone pairs of valence electrons in an ionic or molecular compound. For Lewis structures of molecular compounds (sharing electrons, covalent bonds): - Bonding Pairs of Electrons: a pair of electrons that two atoms share. - Single Bond: A bond that results when two atoms share one pair of electrons.  Drawing Lewis Structures Ex: Draw the Lewis Structure for CO 2. Step 1: Determine the total number of valence electrons. - Add together the valence electrons of each atom in the formula. - Group 1 elements have 1 valence electron - Group 2 elements have 2 valence electrons - Group 13 elements have 3 valence electrons - Group 14 elements have 4 valence electrons - Group 15 elements have 5 valence electrons - Group 16 elements have 6 valence electrons - Group 17 elements have 7 valence electrons - Group 18 elements have 8 valence electrons  Drawing Lewis Structures Step 2: Arrange the symbols of the elements to show how their atoms are bonded and then connect them with single bonds (single pairs of bonding electrons). Hints: - The element with the largest bonding capacity will be in the center (most unpaired valence electrons). If more than one, the least electronegative element will be in the center. - Hydrogen will never be a central atom because it can only form one single bond. - Sometimes the formula will suggest the general structure. - Deduct the number of electrons placed as  Drawing Lewis Structures Step 3: Complete the octets of all the atoms bonded to the central atom by adding lone pairs of electrons. - Hydrogen can only form one bond and will have no lone pairs. - Place lone pairs on all of the outer atoms until each has an octet.  Drawing Lewis Structures Step 4: Compare the number of valence electrons in the Lewis structure with the number determined in step 1. - Any unused valence electrons will need to be placed on the central atom (even if doing so gives it more than an octet) until the structure includes all of the valence electrons.  Drawing Lewis Structures Step 5: Complete the octet on the central atom. - If the central atom has an octet, the structure is complete. - If the central atom does not have an octet, create additional bonds by converting one or more lone pairs of electrons on the outer atoms into bonding pairs.  Calculating Formal Charges Formal Charge: the value calculated for an atom in a molecule or polyatomic ion by determining the difference between the number of valence electrons in the free atom and the sum of the lone-pair electrons plus half the electrons in the atom’s bonding pairs.  Using Formal Charges We use formal charges to determine which Lewis structure best represents the bonding in the molecule when more than one option exists. To do so, we use the following criteria: 1. The best structure is the one in which the formal charge on each atom is zero. 2. If no such structure can be drawn, or if the structure is that of a polyatomic ion, the best structure is the one in which most atoms have formal charges equal to zero or as close to zero as possible. 3. Any negative formal charges should be on the atom(s) of the more/most electronegative element. 4. All formal charges should add to equal the charge on the atom (0) or ion.  Resonance Resonance: a characteristic of electron distributions in which two or more equivalent Lewis structures can be drawn for one compound. - The electrons are delocalized. Resonance Structure: one of two or more Lewis structures with the same arrangement of atoms but different arrangements of bonding pairs of electrons. In resonance structures, all bonds are the same (equivalent). - For ozone (to the right), it does not have one single and one double bond. Instead, it contains two equivalent bonds. - The third pair of electrons is really spread out among both bonds. Practice - NO31-  Exceptions to the Octet Rule 1. Atoms in the third period of the periodic table and below can have more than eight valence electrons around the central atom. Ex: SF6, PCl5 2. Atoms, such as B and Be, can form molecules with less than eight valence electrons. a. B tends to form 3 bonds (6 valence electrons). Ex: BF3 b. Be tends to form 2 bonds (4 valence electrons). Ex: BeCl2 3. You are not responsible for odd-electron molecules.  Bond Strength & Length Double Bond: a bond formed when two atoms share two pairs of electrons. Triple Bond: a bond formed when two atoms share three pairs of electrons. Bond order: the number of bonds between atoms (1 for single bond, 2 for double bond, 3 for triple bond). As the bond order increases, the bond strength increases. (weakest) C-C < C=C < C=C (strongest) As the bond order increases, the bond length decreases. (shortest) C=C < C=C < C-C (longest)