Embedded Files
Lewis Structures, Molecular Geometry and Polarity
Lewis Structures, Molecular Geometry and Polarity
Learning Objectives:
Learning Objectives:
Upon completion of this experiment, students will:
(CLO1). Explain basic chemical concepts related to chemical changes, energy, and properties of matter.
(MLO4-3) Apply rules for drawing Lewis structures of compounds.
(MLO4-4) Determine the formal charge on the atoms in a Lewis structure to identify the most likely structure of a compound when more than one Lewis structure can be drawn.
(MLO4-5) Predict the geometry, shape of a molecule, bond angles and polarity using the Valence shell electron pair repulsion model.
(MLO4-6) Describe the difference between electronic geometry and molecular geometry.
(MLO) Apply hybrid orbital theory to predict the hybridization and explain sigma and pi bonding and resonance in molecules.
(MLO) Describe the different between intramolecular and intermolecular
(MLO) Identify the intermolecular forces: dipole-dipole, hydrogen bonding, and dispersion forces present in a given substance.
Utilize a modeling kit to construct models of simple molecules.
Experiment 7 discusses how to represent molecular compounds with Lewis structures, geometry, hybridization and polarity.
The assignments that must be completed are the following:
The assignments that must be completed are the following:
Pre laboratory problems - NONE The student should attempt to draw the Lewis structures for the compounds in the laboratory report before the laboratory session.
Laboratory Report of the experiment using the resources listed below
Conclusion paragraph - NONE
Discussion and Part A: Lewis Structures, Geometry and Polarity
Discussion and Part A: Lewis Structures, Geometry and Polarity
Rules for drawing Lewis structures
Draw the Lewis Structures for each of the molecules or polyatomic ions.
Determine the best structure for the atoms in the structure or the general connection of each atom to the others.
Which atom is considered to be the CENTRAL atom? Is there more than one central atom?
The central atom is most likely the first one written in the chemical formula, because this is generally the weakest atom by electronegativity and therefore this atom tends to share more of its electron density.
Generally there is one central atom and all other atoms are directly bonded to the central atoms. The hydrogen atom, if written first, can not be the central atom because it can only make one bond.
Hydrogen does not have to be bonded directly to the central atom.
Count the number of valence electrons for each atom present in the molecules. For polyatomic ions that have a charge listed, either add or subtract the appropriate number of electrons to the total number of valence electrons.
Assign two electrons where each pair of atoms are connected. This represents a bonding pair of electrons.
Assign the remaining electrons either as non bonding pairs or create double or triple bonds as needed to satisfy the octet for each atom in the structure.
Unless an odd number of valence electrons are counted for the structure, all electrons are represented as pairs. A double bond is represented by 4 electrons being shared between two atoms, while the triple bond is 6 electrons. No two atoms can share 8 electrons between the two atoms, or there is NO quadruple bond. Hydrogen and Helium are the only two atoms that will not attempt to gain an octet because they can only hold 2 electrons. The molecule can have no more or less than the total number of valence electrons counted in "b".
For each central atom within the structure, determine the bond angle.
The bond angle is determined by the geometry of the central atom, which is determined by the number of electron groups around the central atoms.
Electron groups are defined as any of the following: a single non bonded electron, a non bonding electron pair, a single bonding electron pair (single bond), a double bonding pair (double bond) or a triple bonding pair (triple bond). The bond angle is directly related to the electronic geometry of the molecule.
Wayne Breslyn - A video tutorial for how to draw Lewis Structures in five steps. The video covers the basic Lewis structures you'll see in an introductory chemistry class. Get more chemistry help at www.Breslyn.org.
The five steps are:
Find the total valence electrons for the molecule.
Put the least electronegative atom in the center. Note: Hydrogen (H) always goes outside.
Put two electrons between atoms to form a chemical bond.
Complete octets on outside atoms.
If central atom does not have an octet, move electrons from outer atoms to form double or triple bonds.
Lewis Structures are important to learn because they help us predict:
- the shape of a molecule.
- how the molecule might react with other molecules.
- the physical properties of the molecule (like boiling point, surface tension, etc.).
Appendix F-Electronegativities.pdfAppendix F - Periodic Table of electronegativity values.
Example 7.1
Write the Lewis structure for the chloric acid, HClO3.
Step 1. Write the skeletal. structure. The center atom is the weaker atom in the structure. Hydrogen can NEVER be a center atom.
Electronegativity: Cl = 3.0, O = 3.5
Step 2. Count the total valence electrons. The total valence electrons = 1 (H, 1s1) + 7 (Cl, 3s2 3p5) + 3 x 6 (O, 2s22p4) = 26, i.e., there are 13 pairs of valence electrons.
Step 3. Draw a single bond between the central atom and each of the attached atoms. Four pairs of electrons are used to form 4 single bonds, and 9 pairs of electrons are left.
Step 4. Complete octets of terminal atoms. The terminal atoms are the H and the O atoms. Eight pairs of electrons are used, and one pair of electrons is left.
Step 5. Assign the leftover electrons to the central atom. The last pair of electrons is placed on the central Cl atom. Since each atom now has an octet (except H), there is no need to go to Step 6 to form multiple bonds.
These are some simple overviews of the concepts in this chapter
Groups of the Periodic Table: https://www.khanacademy.org/science/chemistry/periodic-table-trends-bonding/v/groups-of-the-periodic-table
Valence Electrons: https://www.khanacademy.org/science/chemistry/periodic-table-trends-bonding/v/valence-electrons
Lewis dot diagrams: http://youtu.be/ulyopnxjAZ8
Lewis Structures:
Lewis Structure generating programs
Lewis Structure generating programs
This is an interesting program using the above rules, counting valence electrons and allowing you to build the Lewis structure.
https://www.stolaf.edu/depts/chemistry/courses/toolkits/125/js/lewis/
Another website that contains already completed Lewis stuctures of several compounds.
https://idealcalculator.com/lewis-structure-calculator-lewis-structure-generator/
How to Draw Lewis Structures (with example)
How to Draw Lewis Structures (with example)
chemistNATE Ask the author questions on Facebook: http://www.facebook.com/chemistNATE
How to Draw Lewis Structures (aka Lewis Dot Diagrams)
In this video I explain how to draw simple Lewis Structures, while drawing the Lewis structure for a simple molecule (HCN).
https://youtu.be/5igfwUIKpes?si=5jwZxUbJFQKA_ZiK
C6H12 Isomers: All Possible Isomers with Lewis Structure
Geometry of Molecules Hey there, are you looking for a video that can help determine isomers having formula C6H12 and their lewis structure? If yes, then you are at the right place! Today’s video will share a detailed and step-by-step method to determine the Lewis Structure of every isomer having formula C6H12. Watch this video to find out our straightforward approach to determine the Lewis Structure of these compounds. C6H12 Isomers: All Possible Isomers with Lewis Structure
Exceptions to the Octet Rule and differences in structures resulting in structural and geometric isomers add to the complexity of determining the Lewis structure. Formal charge can be used to help determine the "best" structure for the compound.
Formal charge on a single atom equals the number of valence electrons minus the number of non bonding electrons minus half of the number of bonding electrons that are directly on the single atom.
For example, in the structure of carbon dioxide, the carbon atom is the center atoms and to satisfy the octet rule with the 16 electrons in the structure, the carbon has two double bonds linking each of the oxygen atoms to itself.
O = C = O
O = C = O
If we determine the formal charge of the carbon atom, we would begin with this structure. Knowing carbon has four valence electrons and we can see there are no non bonding electrons directly on the carbon but there are four bonds which contain eight electrons, the formal charge is equal to 4 - 0 - 1/2 (8) = 0. The bonding electrons can also be represented by the actual number of bonds. Eight electrons are in four bonds, therefore, 1/2 of eight bonding electrons equals four bonds.
We can see that in a structure where N is double bonded to a center N which is double bonded to oxygen, N=N=O, the formal charge on the central nitrogen is 5 - 0 - 1/2(8) = +1. Similar structure much different formal charge.
Formal charge and dot structures | Chemical bonds | Chemistry | Khan Academy
Formal charge and dot structures | Chemical bonds | Chemistry | Khan Academy
Khan Academy Organic Chemistry
Definition of formal charge, and how minimization of formal charge can help choose the more stable dot structure.
Complete the counting of valence electrons and Lewis structure for each molecule in all tables of the laboratory report:
Complete the counting of valence electrons and Lewis structure for each molecule in all tables of the laboratory report:
Molecular Shapes: VSEPR (Valence-Shell Electron-Pair Repulsion)
The three-dimensional arrangement of the atoms in a molecule and a polyatomic ion is called its molecular geometry or molecular structure. Lewis structures give no information about the shapes of molecules, which are usually determined experimentally using techniques such as X-ray crystallography. However, the molecular geometry can be predicted using the valence-shell electron-pair repulsion (VSEPR) model with an acceptable Lewis structure for a molecule. The VSEPR model is based on the idea that valence shell electrons, being negatively charged, stay as far apart from each other as possible so that the repulsions between them are minimized. It is the repulsions among pairs of bonding and lone pair electrons that control the angles between bonds from a central atom to other atoms surrounding it. In general the strengths of electron-pair repulsions are:
lone pair – lone pair > lone pair – bond pair > bond pair – bond pair
The following steps can be used for predicting molecular geometry:
Draw the Lewis structure of the molecule or the polyatomic ion.
Determine the number of electron groups around the central atom in the Lewis structure. If there is more than one central atom, determine the electrons groups around each SINGLE atom. A non-bonding electron pair, a single non bonded electron (only in odd electron structures), a bonding pair in a single bond, two bonding pairs in a double, or three bonding pairs in a triple bond are counted as one electron group. If resonance structures exist, you must choose one of them for the application of the VSEPR model. For example: X – A – X, where the line represents a single bond with 2 electrons and represents one electron density, the X atoms only have one electron density and are not of interest, but atom A has two single bond; therefore, two electron densities. These electron densities will attempt to get as far apart as physically possible within the structure of the molecule, because they both contain negatively charged electrons which repel each other. From this model, we know that two electron densities will be on opposite sides of the A molecule as represented and therefore will have an 180o bond angle. Similarly, CO2 which has a bonding structure, O = C = O, would have only two electron densities around the center atom and would also have an 180o bond angle. A compound such as AX5 would have a possible structure in which there are 5 single bonds connecting the X atoms to the center A atom, with the resulting structure having 5 electron densities on atom A. For a compound such as X – A = A – X, we would need to look at each A atom separately, Aleft has 2 electron densities and therefore is linear, Aright would be considered separately and the geometry of the molecule would be described in terms of each A atom.
Identify the most stable electronic arrangement of electron groups as linear, trigonal planar, tetrahedral, trigonal bipyramidal, or octahedral.
If more than one arrangement of bonding pairs and lone pairs is possible, choose the one that will have minimum repulsions between lone pairs. In a trigonal bipyramidal arrangement, repulsion is minimized when every lone pair is in the plane of the triangle, i.e., occupying one or more of the three equatorial positions. In an octahedral arrangement, repulsion is minimized when the lone pairs are on opposite sides of the central atom, i.e., occupying one or two of the two axial positions.
Identify the molecular geometry from the locations of the atoms at the ends of the bonds.
VESPR theory - Geometry: http://youtu.be/nxebQZUVvTg
Tyler DeWitt presentation on Molecular Structure and geometry
VSEPR Theory: Introduction - https://youtu.be/nxebQZUVvTg
VSEPR Theory Part 2: Trigonal Bipyramidal Family - https://youtu.be/cdo6FtSU_k8
VSEPR Theory Part 3: Octahedral Family - https://youtu.be/Qcy-TjJ10xk
VSEPR Theory: Common Mistakes - https://youtu.be/8Tl_bDWCAmo
VSEPR Theory Practice Problems - https://youtu.be/xwgid9YuH58
VSEPR Theory Practice Problems (Advanced) - https://youtu.be/Ip8v87vxSok
VSEPR Theory and Molecular Geometry
VSEPR Theory and Molecular Geometry
Professor Dave Explains - Did you know that geometry was invented by molecules? It's true! Until the first stars went supernova and littered all the elements across the cosmos, everything was simply spheres, from protons to stars. But then, under cooler planetary conditions, atoms came together to form molecules, and in doing so, they created the first shapes the universe had ever seen. Learn about the different shapes that molecules can make!
Molecular Shapes: Basic
Molecular Shapes: Basic
https://phet.colorado.edu/sims/html/molecule-shapes-basics/latest/molecule-shapes-basics_en.html
PhET interactive simulations
you can practice making and looking at simple molecule's electronic and molecular geometry for simple models and real molecules with one center atom
Molecular Shapes Activity
https://phet.colorado.edu/en/simulations/molecule-shapes
PhET interactive simulations
you can practice making and looking at simple molecule's electronic and molecular geometry. This expands to a little more complicated structures and you can build multiple center atom structures.
You can use this program to visualize structure of simple and very complicated compound both in 2D or more importantly in 3D with rotation
Visualizing Molecular Geometry With 3D Software
Visualizing Molecular Geometry With 3D Software
Professor Dave Explains - We've already learned about VSEPR theory, and how to use it to predict the shapes of various molecules. But we didn't cover all of the molecular geometries for certain hybridizations, and we didn't talk about bond angles, so let's do both of those things now, using some neat 3D software!
Complete the counting of electron groups, electronic and molecular geometry and bond angle for each molecule in all tables of the laboratory report as required:
Complete the counting of electron groups, electronic and molecular geometry and bond angle for each molecule in all tables of the laboratory report as required:
Hybridization
The question then becomes, how does the molecule achieve these geometries with the atomic orbitals of the quantum model (s, p, d and f orbitals) that have very specific geometry. The geometry of the individual atomic orbitals does not match the geometry of the molecules as described by the VSEPR bonding model. The Valence bonding model describes how the individual atomic orbitals can be used to create new mixtures of orbitals that are HYBRIDS of the originals. Therefore, the sp hybrid orbital is created from two atomic orbitals combining together, one s orbital and one p orbital. The hybrid orbitals then follow the rules of the VSEPR model in which the electron orbitals repel each other and move as far apart as possible.
The hybridization of the atomic orbitals is dependent upon the number of bonding orbitals needed and therefore directly related to the geometry of the molecule. If an atom needs to make three bonds, then 3 atomic orbitals need to be combined to form three hybrid molecular orbitals. So a trigonal planar molecule, will use an s, p and p orbital to form three sp2 hybrid orbitals.
Atomic and Molecular orbitals: https://winter.group.shef.ac.uk/orbitron/
Hybridization
Hybridization
Brightstorm
hybridization of the s, p, d orbitals are described and how to apply hybridization to a simple molecule.
Valence Bond Theory, Hybrid Orbitals, and Molecular Orbital Theory
Valence Bond Theory, Hybrid Orbitals, and Molecular Orbital Theory
Professor Dave Explains
Alright, let's be real. Nobody understands molecular orbitals when they first take chemistry. You just pretend you do, and then in your next course you learn them a little better. And then a little better than that. And five years later you sort of understand them. So don't get discouraged! Just watch this for a head start and do your best.
Hybridization Theory
Hybridization Theory
SponholtzProductions
Chapter 1: Why Hybridization Theory was Developed, Why is it Important to Visualize Atoms within a Molecule in Three-Dimensions.
This is a very long and visual video. There are several more chapter in the understanding of hybridization.
Narration: Gordon W. Gribble, PH.D., Dartmouth College
Complete the hybridization for each molecule in all tables of the laboratory report as required:
Complete the hybridization for each molecule in all tables of the laboratory report as required:
Polarity
The polarity of a molecule is based on the electronic geometry and the electronegativity of the atoms in the structure. If two atoms are bonded together in a covalent bond, the electrons are shared between the two atoms of the covalent bond. If the two atoms are sharing the electrons equally because they have the same electronegativities, then the bond is said to be NON - POLAR or the bond does not have a dipole. We can consider this bond to have perfect symmetry. If the atoms are not the same, therefore having different electronegativities, then the electrons will not be shared equally and the bond is said to be POLAR or the bond has a dipole in which one atom carries a positive charge and the other carries a negative charge We can consider this bond to not be symmetrical.
If we extend this idea to a complete molecule that contains more than one bond, the idea of symmetry is best used to determine polarity of the molecule. If all the bonds in a molecule are polar but the molecule is perfectly symmetrical, then the molecule will most likely be non - polar, but if there is limited or no symmetry to the molecule, then the molecule is most likely polar. For a molecule to be polar, there must be polar bonds and a lack of symmetry in structure such that the bond dipoles do not cancel each other out. CCl4 has polar bonds but because of its tetrahedral geometry, is a non-polar molecule. Water has polar bonds and because of its bent geometry is highly polar. In general, polar molecules will have higher boiling points than non-polar molecules with similar molecular mass.
We can also look at the molecular geometry versus the electronic geometry to determine polarity. If the electronic geometry is not the same as the molecular geometry the molecule will most likely be polar. However, if the electronic geometry is the same as the molecular geometry, the individual bonds will determine polarity. What is the difference between the electronic geometry and the molecular geometry? The electronic geometry accounts for ALL electron groups while the molecular geometry only accounts for bonding electron groups such as single, double and triple bonds.
Polarity of Molecules
How to determine polarity:
Molecule Polarity
Molecule Polarity
PhET interactives: Molecular Polarity
https://phet.colorado.edu/en/simulation/molecule-polarity
Use this program to explore the simple polarity of 2 atom and 3 atom molecules.
You can use the older Java-flash program to see the polarity a few simple real molecules like carbon dioxide and water.
https://phet.colorado.edu/en/simulation/legacy/molecule-polarity
Use this activity to walk through the program: http://dotterer.weebly.com/uploads/5/6/0/2/56028317/3_polarity_sim_lab.pdf
Predict the polarity for each molecule in all tables of the laboratory report as required:
Predict the polarity for each molecule in all tables of the laboratory report as required:
Polar Molecules Tutorial: How to determine polarity in a molecule
Polar Molecules Tutorial: How to determine polarity in a molecule
Crash Chemistry Academy
This video looks at how to determine polarity in a molecule by understanding how the bond polarities, molecule shape, and outside atoms influence polarity using bond polarity vector addition. This includes a flow chart that guides you through the various decisions needed to determine if a molecule is polar or not.
Polar and Nonpolar molecules
Polar and Nonpolar molecules
This chemistry video tutorial provides a basic introduction into polar and nonpolar molecules.
Intermolecular Forces and Polarity
hydrogen bonding: http://youtu.be/PyC5r2mB4d4
Procedure and Report for Part B: Intermolecular Forces and Polarity
Procedure and Report for Part B: Intermolecular Forces and Polarity
Miscibility and Solubility: In the liquid and solid states, the molecules of a substance are closer to each other caused by an attraction between the molecules due to the bonding, structure and most importantly the polarity of the molecule. These forces are called intermolecular attractions. For covalently bonded substances, the attractions are usually classified into three groupings: dispersion forces (or London forces), ion-dipole attraction, dipole-dipole attractions and a subcategory of the latter, hydrogen bonds. Hydrogens bonded to N, O and F are capable of hydrogen bonding to an N, O or F in another molecule. None of the other elements is electronegative enough to have this capability.
Very briefly, London forces are temporarily induced polarizations that result from the approach of two molecules. The magnitude of these forces depends on the size and shape of the molecules. As polarization of electron clouds are of prime importance in this effect, it would be expected that the strength of the attractions should depend partially on the number of electrons. As the number of electrons correlates with the molecular mass, the boiling point would be expected to correlate with molecular mass if other parameters do not change. We see that the boiling point of the halogen family decreases with decreasing molar mass.
These forces strongly influence the properties of substances. For instance, water and carbon tetrachloride are not miscible, meaning the two liquids do not mix homogeneously. The attractive forces between water are due to the polarity of the water while the forces between carbon tetrachloride molecules are very weak due to the non polar structure of the water and carbon tetrachloride cannot overcome the intermolecular forces between water molecules.
In the gaseous state, molecules behave as though they are independent of each other and do not experience significant intermolecular attractions except at high pressure or very low temperature. To boil a liquid requires that sufficient energy be supplied to overcome the intermolecular attractions. The boiling point should and does generally correlate with the strength of intermolecular attractions.
Intermolecular attractive forces and especially hydrogen bonds are extremely significant in chemistry. They influence boiling points, solubility and chemical reactivity and are responsible for determining the 3 Dimensional structures of proteins and DNA. Molecules with similar polarities will have a greater tendency to dissolve in each other to produce a homogeneous mixture. This is often all the rule of "like dissolves like."
Solubility is defined as the amount of a substance that can dissolve in a volume of the solvent. Generally we speak of the solubility of solids such as NaCl, salt, in water or oxygen gas in water. When two liquids are mixed, we use a different term to describe how they mix. A liquid such as alcohol is said to be miscible in water, while oil is immiscible.
In this experiment, we will investigate the solubility or miscibility of several substances in water, a polar solvent and kerosene, a non polar hydrocarbon.
Extraction is the process of using the solubility of a substance to purify a sample. For example, if we mixed salt and wax, we could use water to extract the salt from the mixture. With mixtures that are already dissolved, the separation is a little more complicated. Consider an immiscible pair of liquids such as vinegar (polar - 95% water, 5% acetic acid) and oil (non-polar). Suppose the vinegar has a small amount of a non-polar solute dissolved in it and you shake the vinegar with some oil. The solute will have two options. As its polarity is closer to that of the oil than of vinegar, the solute will probably shift to the oil. We say that the solute has been extracted from the vinegar by the oil. If the solute had been polar, it would have remained in the vinegar.
Report abuse