Unit 5 Bonding Theory Notes PDF

Unit 5: Bonding Theory Let’s look at another type of bond! What happens during a covalent bond? Higher electronegativity value = stronger attraction towards electrons. What’s the trend? Tug of War It’s a draw!  Tug of War It’s clear who’s the winner! Electronegativity Scale What’...

Unit 5: Bonding Theory Let’s look at another type of bond! What happens during a covalent bond? Higher electronegativity value = stronger attraction towards electrons. What’s the trend? Tug of War It’s a draw!  Tug of War It’s clear who’s the winner! Electronegativity Scale What’s the rule for naming covalent compounds? Name or write the formula for the following. 1) N2O5 2) arsenic tribromide 3) CSe2 4) dihydrogen monoxide How do we represent how atoms bond? Lewis Dot Structures Gilbert N. Lewis Write the electron configuration (long form). How many electrons are in the last energy level? 1) O 2) P Drawing Lewis Structures 1) Draw Lewis structure for each atom. 2) Determine total number of valence electrons. 3) Arrange structures to show how atoms bond. 4) Complete octet around each atom. Watch out for exceptions: H (1 bond → 2 e-); Be (2 bonds → 4 e-); B (3 bonds → 6 e-). 5) Try double or triple bond if necessary. 6) Change bonded pairs of dots to dash lines. Draw a Lewis structure for the following. 1) H2S 2) CH2F2 Let’s do some modeling! Directions Draw Lewis dot structures for each compound by using the following: ❏ Cards (elements) ❏ Stones (electron dots) Draw a Lewis structure for the following. 1) O2 2) N2 Watch out for exceptions! Draw a Lewis structure for the following. 1) SiO2 – 2) OH 3) BH3 4) SF6 Watch out for less than and expanded octets! Is there another way to draw this? CO2 Formal Charge We use F.C. to decide which structure is most likely. Formal Charge Choose Lewis structure where all F.C. of atoms are zero or close to zero. If there must be a formal charge on an atom, the most electronegative atom should have the negative formal charge. Draw as many Lewis dot structures as possible. Then, use formal charges to identify the most reasonable one. - NCO Let’s build molecules! Count the total number of electron domains (surrounding atoms & lone pairs) around central atom to determine molecular geometry. Draw Lewis structure & determine the molecular shape for the following. 1) NH3 2) BF3 Molecular Models Carbon (C) = Black Hydrogen (H) = Yellow Oxygen (O) = Red Nitrogen (N) = Blue Chlorine (Cl) = Green Bromine (Br) = Orange Iodine (I) / Fluorine (F) = Purple Draw Lewis structure & determine the molecular shape for the following. C2H5OH Dipole moments (uneven charge distribution) arise from differences in electronegativity. How do we determine if a compound is polar or nonpolar? Draw a Lewis structure & determine the molecular shape for the following. 1) CO2 2) H2O Tug of War Even pull! Tetrahedral Trigonal Planar Determine the molecular shape and polarity (polar or nonpolar) for the following. 1) BCl3 2) CH3Br Let’s revisit electron configuration! How do bonds form through orbitals? What kind of properties do ionic & covalent compounds have? Conductivity Chart Scale Red LED Green LED Conductivity 0 Off Off None 1 Dim Off Low 2 Medium Off Medium 3 Bright Dim High 4 Very Bright Medium Very High Ionic Covalent How are ionic compounds able to conduct electricity in water? Draw the symbols to show which atom has the partial negative charge and partial positive charge. Use electronegativity values or trend! Using dash lines, show how water molecules are held together. Think about cohesion! Water molecules are attracted to other water molecules due to hydrogen bonding. DNA is a perfect example of hydrogen bonding! Hydrogen is strongest when bonded to nitrogen (N), oxygen (O), or fluorine (F). Arrange the following in increasing boiling point. Explain your reasoning. BaCl2, H2, CO, HF, Ne Polarizability More electrons results in larger atomic size, which means stronger LDF. Higher number of electrons will increase the chance of forming induced dipole, which will make the interactions stronger. What happens when there’s a lid on the container while water is heated? Remember: A liquid with weak IMF evaporates more easily and has a high vapor pressure.  Substance Boiling Point Bond Length Bond Strength (oC) (A) (kcal/mol) H2 -253 0.75 104.2 N2 -196 1.10 226.8 O2 -182 1.21 118.9 Cl2 -34 1.99 58.0 1) Explain the differences in bond strengths of N2 and O2. 2) Explain why H2 and O2 are gases at room temperature, while water is a liquid at room temperature. 3) Explain the differences in boiling points of O2 and Cl2. Time to design an experiment! Unit 5 Test Ionic & covalent compounds (polar vs nonpolar) Naming/writing ionic & covalent formulas Octet rule Lewis dot structures Formal charges Molecular shapes (VSEPR theory) Hybridization Molecular polarity Ionic & covalent properties Intermolecular forces