COVALENT BONDING CHEMICAL BONDING I: LEWIS MODEL. Chapter 7
|
|
- Patrick Stewart
- 5 years ago
- Views:
Transcription
1 Chapter 7 P a g e 1 COVALENT BONDING Covalent Bonds Covalent bonds occur between two or more nonmetals. The two atoms share electrons between them, composing a molecule. Covalently bonded compounds are also called molecular compounds. Structural Formula A structural formula uses lines to represent covalent bonds and shows how atoms in a molecule are connected or bonded to each other. It can also show the geometry of a molecule. The structural formula for H2O2 and CO2 are shown below: CHEMICAL BONDING I: LEWIS MODEL One of the simplest bonding theories is called Lewis theory. Lewis theory emphasizes valence electrons to explain bonding. Using Lewis theory, we can draw models, called Lewis structures. Also known as electron dot structures Lewis structures allow us to predict many properties of molecules. Molecular stability, shape, size, and polarity Gilbert Newton Lewis Why do Atoms Bond? Chemical bonds form because they lower the potential energy between the charged particles that compose atoms. A chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of the separate atoms. To calculate this potential energy, you need to consider the following interactions: Nucleus-to-nucleus repulsions Electron-to-electron repulsions Nucleus-to-electron attractions Valence Electrons and Bonding The column number on the periodic table will tell you how many valence electrons a main group atom has. Valence electrons are held most loosely. Chemical bonding involves the transfer or sharing of electrons between two or more atoms. Because of the two previously listed facts, valence electrons are most important in bonding. Lewis theory focuses on the behavior of the valence electrons.
2 P a g e 2 Lewis Structures of Atoms In a Lewis structure, we represent the valence electrons of maingroup elements as dots surrounding the symbol for the element. Also known as electron dot structures We use the symbol of the element to represent the nucleus and inner electrons. We use dots around the symbol to represent valence electrons. Pair the first two dots for the s orbital electrons. Put one dot on each open side for the first three p electrons. Then, pair the rest of the dots for the remaining p electrons. Lewis Bonding Theory Atoms bond because bonding results in a more stable electron configuration. More stable = lower potential energy Atoms bond together by either transferring or sharing electrons. Usually, this results in all atoms obtaining an outer shell with eight electrons. Octet rule There are some exceptions to this rule: The key to remember is to try to get an electron configuration like a noble gas. Covalent Bonding: Bonding and Lone Pair Electrons Electrons that are shared by atoms are called bonding pairs. Electrons that are not shared by atoms but belong to a particular atom are called lone pairs. Also known as nonbonding pairs. Single Covalent Bonds When two atoms share one pair of electrons, it is called a single covalent bond. Two electrons One atom may use more than one single bond to fulfill its octet. To different atoms H only duet Double Covalent Bonds When two atoms share two pairs of electrons the result is called a double covalent bond. Four electrons
3 P a g e 3 Triple Covalent Bonds When two atoms share three pairs of electrons the result is called a triple covalent bond. Six electrons Covalent Bonding: Model versus Reality Lewis theory predicts that the more electrons two atoms share, the stronger the bond should be. Bond strength is measured by how much energy must be added into the bond to break it in half. In general, triple bonds are stronger than double bonds, and double bonds are stronger than single bonds. However, Lewis theory would predict that double bonds are twice as strong as single bonds; the reality is that they are less than twice as strong. Lewis theory predicts that the more electrons two atoms share, the shorter the bond should be. When comparing bonds to like atoms Bond length is determined by measuring the distance between the nuclei of bonded atoms. In general, triple bonds are shorter than double bonds, and double bonds are shorter than single bonds. Polar Covalent Bond Covalent bonding between unlike atoms results in unequal sharing of the electrons. One atom pulls the electrons in the bond closer to its side. One end of the bond has larger electron density than the other. The result is a polar covalent bond. Bond polarity The end with the larger electron density gets a partial negative charge. The end that is electron deficient gets a partial positive charge. Electronegativity The ability of an atom to attract bonding electrons to itself is called electronegativity. Increases across period (left to right) and decreases down group (top to bottom) Fluorine is the most electronegative element. Francium is the least electronegative element. Noble gas atoms are not assigned values. Opposite of atomic size trend The larger the difference in electronegativity, the more polar the bond. Negative end toward more electronegative atom Electronegativity Difference and Bond Type If the difference in electronegativity between bonded atoms is 0, the bond is pure covalent. Equal sharing If the difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent. If the difference in electronegativity between bonded atoms is 0.4 to 1.9, the bond is polar covalent.
4 P a g e 4 If the difference in electronegativity between bonded atoms is larger than or equal to 2.0, the bond is 100% ionic. Example 6 Classify the bond formed between each pair of atoms as covalent, polar covalent, or ionic. a. Sr and F b. Na and Cl c. N and O d. I and I e. Cs and Br f. P and O
5 P a g e 5 Writing Lewis Structures of Molecular Compounds 1. Write the correct skeletal structure for the molecule. Hydrogen atoms are always terminal. The more electronegative atoms are placed in terminal positions. 2. Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule. For polyatomic ions, consider the charge. Negative ions add electrons Positive ions subtract electrons 3. Distribute the electrons among the atoms, giving octets (or duets in the case of hydrogen) to as many atoms as possible. 4. If any atoms lack an octet, form double or triple bonds as necessary to give them octets. Move lone electron pairs from terminal atoms into the bonding region with the central atom. Note: Attempt these examples on a separate sheet of paper. Example 7 Write the Lewis structures for the following molecules: CO2, CO, NH3, SCl2, CH3SH (C and S central), HCOOH (both O bonded to C), C2H2, Cl2O, N2H2, N3H8 Example 8 Write the Lewis structures for the following ions: BrO, NO 2, O 2 2, NH 4 +, ClO 3, CO 3 2, CO +, NO 3, ClO 4, OCl Resonance and Formal Charge Lewis theory localizes the electrons between the atoms that are bonding together. Extensions of Lewis theory suggest that there is some degree of delocalization of the electrons; we call this concept resonance. Delocalization of charge helps to stabilize the molecule. When there is more than one Lewis structure for a molecule that differ only in the position of the electrons, they are called resonance structures. The actual molecule is a combination of the resonance forms a resonance hybrid. The molecule does not resonate between the two forms, though we often draw it that way. Look for multiple bonds or lone pairs. Example 9 Write the Lewis structures for NO2 and NO3. Include resonance structures.
6 P a g e 6 Formal Charge Formal charge is a fictitious charge assigned to each atom in a Lewis structure that helps us to distinguish among competing Lewis structures. In a Lewis structure, we calculate an atom s formal charge, which indicates the charge it would have if all bonding electrons were shared equally between the bonded atom. FC = # valence e [nonbonding e + ½ bonding e ] Sum of all the formal charges in a molecule = 0. In an ion, total equals the charge. Evaluating Resonance Structures Better structures have fewer formal charges. Better structures have smaller formal charges. Better structures have the negative formal charge on the more electronegative atom. Example 10 Write a Lewis structure that obeys the octet rule for each molecule or ion. Include resonance structures if necessary and assign formal charges to each atom. SeO 2, CO 3 2, ClO 3 Example 11 Use formal charge to determine which Lewis structure is better: Example 12 In N2O, nitrogen is the central atom and the oxygen atom is terminal. In OF2, however, oxygen is the central atom. Use formal charge to explain why. Example 13 Draw the Lewis structure (including resonance structures) for methyl azide (CH3N3). For each resonance structure, assign formal charges to all atoms that have formal charge. Exceptions to the Octet Rule There are three exceptions to the Lewis model: i. Odd-electron species ii. Incomplete octets iii. Expanded octet
7 P a g e 7 Odd-Electron Species Molecules and ions with odd number of electrons in their Lewis structures are called free radicals. o They have one unpaired electron. o Radicals are very reactive (unstable) because they want to attain an octet. o They are relatively rare compared to other molecules. For example, nitrogen dioxide, NO2 Incomplete Octets Some elements tend to form incomplete octets. o For example, boron (six electrons instead of eight) and beryllium For BF3, what if we form double bonds to increase the number of electrons around boron? When we assign formal charges to all the atoms, we get the following: What is wrong (if any) with the structure above? One of the ways BF3 can get an octet is through a chemical reaction. It gains electrons to complete its octet. Expanded Octets This means more than 8 electrons, up to 12 (and occasionally 14) electrons. o Elements in the third row (period) and beyond can exhibit expanded octets. o They have energetically accessible d-orbitals to accept the extra electrons. o Consider arsenic pentafluoride and sulfur hexafluoride: Example 14 Write the Lewis structures for XeF2, XeF4, and H3PO4.
8 P a g e 8 Example 15 Write Lewis structures for each molecule or ion. Include resonance structures if necessary and assign formal charges to all atoms. If necessary, expand the octet on the central atom to lower formal charge. a. PO 4 3 b. I 3 c. AsF 6 d. Cl 3 PO Example 16 Draw the Lewis structure for urea, H2NCONH2, one of the compounds responsible for the smell of urine. Does urea contain polar bonds? Which bond (if any) in urea is most polar? [Hint: The central carbon atom is bonded to both nitrogen atoms and to the oxygen atom]. Example 17 Phosgene (Cl2CO) is a poisonous gas used as a chemical weapon during World War I. It is a potential agent for chemical terrorism today. Draw the Lewis structure for phosgene. Include all three resonance forms by alternating the double bond among the three terminal atoms. Which resonance structure is the best? Explain. CHEMICAL BONDING II: VSEPR THEORY VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORY Properties of molecular substances depend on the structure of the molecule. Valence shell electron pair repulsion (VSEPR) theory is a simple model that allows us to account for molecular shape. o Electron groups are defined as lone pairs, single bonds, double bonds, and triple bonds. o VSEPR is based on the idea that electron groups repel one another through coulombic forces. Electron groups around the central atom will be most stable when they are as far apart as possible. We call this VSEPR theory. o Because electrons are negatively charged, they should be most stable when they are separated as much as possible. The resulting geometric arrangement will allow us to predict the shapes and bond angles in the molecule. The Lewis structure predicts the number of valence electron pairs around the central atom(s). o Each lone pair of electrons constitutes one electron group on a central atom. o Each bond constitutes one electron group on a central atom, regardless of whether it is single, double, or triple. Electron Group Geometry There are five basic arrangements of electron groups around a central atom. Based on a maximum of six bonding electron groups Though there may be more than six on very large atoms, it is very rare. Each of these five basic arrangements results in five different basic electron geometries. In order for the molecular shape and bond angles to be a perfect geometric figure, all the electron groups must be bonds, and all the bonds must be equivalent.
9 P a g e 9 For molecules that exhibit resonance, it doesn t matter which resonance form you use since the electron geometry will be the same. Two Electron Groups: Linear Geometry When there are two electron groups around the central atom, they will occupy positions on opposite sides of the central atom. This results in the electron groups taking a linear geometry. The bond angle is 180. Three Electron Groups: Trigonal Planar Geometry When there are three electron groups around the central atom, they will occupy positions in the shape of a triangle around the central atom. This results in the electron groups taking a trigonal planar geometry. The bond angle is 120. Four Electron Groups: Tetrahedral Geometry When there are four electron groups around the central atom, they will occupy positions in the shape of a tetrahedron around the central atom. This results in the electron groups taking a tetrahedral geometry. The bond angle is Five Electron Groups: Trigonal Bipyramidal Geometry When there are five electron groups around the central atom, they will occupy positions in the shape of two tetrahedra that are base to base with the central atom in the center of the shared bases. This results in the electron groups taking a trigonal bipyramidal geometry. The positions above and below the central atom are called the axial positions. The positions in the same base plane as the central atom are called the equatorial positions. The bond angle between equatorial positions is 120. The bond angle between axial and equatorial positions is 90. Octahedral Electron Geometry When there are six electron groups around the central atom, they will occupy positions in the shape of two square-base pyramids that are base to base with the central atom in the center of the shared bases. This results in the electron groups taking an octahedral geometry.
10 P a g e 10 It is called octahedral because the geometric figure has eight sides. All positions are equivalent. The bond angle is 90. The Effect of Lone Pairs The actual geometry of the molecule may be different from the electron geometry. Lone pair electrons typically exert slightly greater repulsion than bonding electrons, affecting the bond angles. A lone electron pair is more spread out in space than a bonding electron pair because a lone pair is attracted to only one nucleus while a bonding pair is attracted to two nuclei. In general, electron group repulsions vary as follows: Lone pair lone pair > lone pair bonding pair > bonding pair bonding pair Derivatives of the Tetrahedral Geometry When there are four electron groups around the central atom, and one is a lone pair, the result is called a trigonal pyramidal shape, because it is a triangular-base pyramid with the central atom at the apex. When there are four electron groups around the central atom, and two are lone pairs, the result is called a tetrahedral-bent shape. Consider ammonia, NH3: Consider water, H2O:
11 P a g e 11 Derivatives of the Trigonal Bypramidal Geometry Lone pairs on central atoms with five electron groups will occupy the equatorial positions because there is more room. The result is called the seesaw shape (aka distorted tetrahedron). When there are two lone pairs around the central atom, the result is T-shaped. When there are three lone pairs around the central atom, the result is a linear shape. The bond angles between equatorial positions are less than 120. The bond angles between axial and equatorial positions are less than 90. Linear = 180 axial to axial. Seesaw T-Shaped Linear Derivatives of the Octahedral Geometry When there are lone pairs around a central atom with six electron groups, each even number lone pair will take a position opposite the previous lone pair. When one of the six electron groups is a lone pair, the result is called a square pyramid shape. The bond angles between axial and equatorial positions are less than 90. When two of the six electron groups are lone pairs, the result is called a square planar shape. The bond angles between equatorial positions are 90. Square Pyramidal Square Planar Representing Three Dimensional Structures on Paper One of the problems with drawing molecules is trying to show their dimensionality. By convention, the central atom is put in the plane of the paper. Put as many other atoms as possible in the same plane and indicate with a straight line. For atoms in front of the plane, use a solid wedge. For atoms behind the plane, use a hashed wedge.
12 P a g e 12 Example 18 Determine the molecular geometry of NO3. Example 19 Determine the molecular geometry of CCl4. Example 20 Suppose that a molecule with six electron groups were confined to two dimensions and therefore had a hexagonal planar electron geometry. If two of the six groups were lone pairs, were would they be located? a) Positions 1 and 2 b) Positions 1 and 3 c) Positions 1 and 4 Example 21 Predict the geometry and bond angles of PCl3. Example 22 Predict the geometry and bond angles of ICl4. Predicting the Shape of Larger Molecules Many molecules have larger structures with many interior atoms. We can think of them as having multiple central atoms. we describe the shape around each central atom in sequence. Consider the amino acid, glycine:
13 P a g e 13 Example 23 Predict the geometry about each interior atom in methanol (CH3OH) and make a sketch of the molecule. Predicting Polarity of Molecules Draw the Lewis structure, and determine the molecular geometry. Determine whether the bonds in the molecule are polar. o If there are no polar bonds, the molecule is nonpolar. Determine whether the polar bonds add together to give a net dipole moment. Adding Dipole Moments (Vector Addition) 1. The H Cl bond is polar. The bonding electrons are pulled toward the Cl end of the molecule. The net result is a polar molecule.
14 P a g e The O C bond in CO2 is polar. The bonding electrons are pulled equally toward both O ends of the molecule. The net result is a nonpolar molecule. 3. The H O bond in H2O is polar. Both sets of bonding electrons are pulled toward the O end of the molecule. Because the molecule is bent, not linear, the net result is a polar molecule. Example 24 Determine whether NH3 is polar. Example 25 Determine whether CF4 is polar. Example 26 Determine whether CClF3 is polar. Polarity and Solubility in Water Like dissolves like. Polar molecules are attracted to other polar molecules. Because water is a polar molecule, other polar molecules dissolve well in water. And ionic compounds as well. Water and oil do not mix because water molecules are polar and the molecules that compose oil are generally nonpolar. Some molecules have both polar and nonpolar parts for example, soaps.
15 P a g e 15 THE VALENCE BOND (VB) THEORY Valence Bond theory (VB) approaches chemical bonding based on an extension of the quantummechanical model (perturbation theory). When orbitals on atoms interact, they make a bond. These orbitals are hybridized atomic orbitals, a kind of blend or combination of two or more standard atomic orbitals. When two atoms approach each other, the electrons and nucleus of one atom interact with the electrons and nucleus of the other atom. If the energy of the system is lowered because of the interactions, a chemical bond forms. When two atoms with half-filled orbitals approach each other, the half-filled orbitals overlap and the electrons align with opposite spins (spin-pair). This results in a net energy stabilization and hence a chemical bond. A bond can also result from the overlap of a completely filled orbital with an empty orbital. The geometry of the overlapping orbitals determines the shape of the molecule. When two atoms approach each other, the electrons and nucleus of one atom interact with the electrons and nucleus of the other atom. If the interaction lowers the energy (negative interaction energy), a chemical bond forms. If the interaction raises the energy (positive interaction energy), a chemical bond does not form. When the atoms are far apart, the energy is nearly zero because they are not interacting. As they get closer, the energy is lowered (becomes negative) and is minimum at the optimal overlap. At this minimum energy, the overlap has the ideal bond length. If the atoms get too close, there will be mutual repulsion of the two positively charge nuclei. Orbital Diagram for H2S
16 P a g e 16 Hybridization of Atomic Orbitals The overlap of standard half-filled orbitals does not adequately explain the bonding in many other molecules. Consider the bonding between hydrogen and carbon: Based on the electron configurations of carbon and hydrogen, the compound will be CH2 with a bond angle of 90. However, CH2 does not exist! Instead, CH4 is observed. Hybridization is particularly important in carbon, which tends to form four bonds in its compounds and therefore always hybridizes. Hybrid Orbitals The number of standard atomic orbitals combined = the number of hybrid orbitals formed. Combining a 2s with a 2p gives two 2sp hybrid orbitals. H cannot hybridize! Its valence shell has only one orbital. The number and type of standard atomic orbitals combined determines the shape of the hybrid orbitals. The particular kind of hybridization that occurs is the one that yields the lowest overall energy for the molecule. sp 3 Hybridization One s orbital and three p orbitals are mixed. Atom with four electron groups around it. Tetrahedral geometry angles between hybrid orbitals Atom uses hybrid orbitals for all bonds and lone pairs. Example, NH3, CH4, H2O sp 2 Hybridization One s orbital + two p orbitals are mixed. Hybrid orbitals will overlap on axis with orbitals from other atoms. Trigonal planar geometry 120 bond angles between hybrid orbitals.
17 P a g e 17 Unhybridized p orbital will overlap sideways, or side by side, with an unhybridized p orbital of another atom. Sigma Bonds and pi Bonds A sigma (σ) bond results when the interacting atomic orbitals point along the axis connecting the two bonding nuclei. Either standard atomic orbitals or hybrids s to s, p to p, hybrid to hybrid, s to hybrid, etc. A pi (π) bond results when the bonding atomic orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei. Between unhybridized parallel p orbitals The interaction between parallel orbitals is not as strong as between orbitals that point at each other; therefore, σ bonds are stronger than π bonds. Bond Rotation Because the orbitals that form the s bond point along the internuclear axis, rotation around that bond does not require breaking the interaction between the orbitals. But, the orbitals that form the p bond interact above and below the internuclear axis, so rotation around the axis requires the breaking of the interaction between the orbitals.
18 P a g e 18 sp Hybridization Atom with two electron groups; for example, C2H2 Linear shape 180 bond angle Atom uses hybrid orbitals for s bonds or lone pairs and uses nonhybridized p orbitals for p bonds Usually will for two s bonds and two p bonds.
19 P a g e 19 sp 3 d Hybridization Atom with five electron groups around it. Trigonal bipyramid electron geometry Seesaw, T-shape, linear 120 and 90 bond angles Use empty d orbitals from valence shell. sp 3 d 2 Hybridization Atom with six electron groups around it Octahedral electron geometry Square pyramid, Square planar 90 bond angles Use empty d orbitals from valence shell to form hybrid. Example 27 Write a hybridization and bonding scheme for a. BrF3 b. XeF4 c. HCN d.
Lecture Presentation. Chapter 10 Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory
Lecture Presentation Chapter 10 Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory Predicting Molecular Geometry 1. Draw the Lewis structure. 2. Determine the number
More informationChapter 7. Chemical Bonding I: Basic Concepts
Chapter 7. Chemical Bonding I: Basic Concepts Chemical bond: is an attractive force that holds 2 atoms together and forms as a result of interactions between electrons found in combining atoms We rarely
More informationChapter 9 Molecular Geometry. Lewis Theory-VSEPR Valence Bond Theory Molecular Orbital Theory
Chapter 9 Molecular Geometry Lewis Theory-VSEPR Valence Bond Theory Molecular Orbital Theory Sulfanilamide Lewis Structures and the Real 3D-Shape of Molecules Lewis Theory of Molecular Shape and Polarity
More informationChapter 10. Structure Determines Properties! Molecular Geometry. Chemical Bonding II
Chapter 10 Chemical Bonding II Structure Determines Properties! Properties of molecular substances depend on the structure of the molecule The structure includes many factors, including: the skeletal arrangement
More information11/14/2014. Chemical Bonding. Richard Philips Feynman, Nobel Laureate in Physics ( )
Chemical Bonding Lewis Theory Valence Bond VSEPR Molecular rbital Theory 1 "...he [his father] knew the difference between knowing the name of something and knowing something" Richard Philips eynman, Nobel
More informationMolecular shape is determined by the number of bonds that form around individual atoms.
Chapter 9 CH 180 Major Concepts: Molecular shape is determined by the number of bonds that form around individual atoms. Sublevels (s, p, d, & f) of separate atoms may overlap and result in hybrid orbitals
More informationChemical Bonding AP Chemistry Ms. Grobsky
Chemical Bonding AP Chemistry Ms. Grobsky What Determines the Type of Bonding in Any Substance? Why do Atoms Bond? The key to answering the first question are found in the electronic structure of the atoms
More informationChapter 9. Chemical Bonding II: Molecular Geometry and Bonding Theories
Chapter 9 Chemical Bonding II: Molecular Geometry and Bonding Theories Topics Molecular Geometry Molecular Geometry and Polarity Valence Bond Theory Hybridization of Atomic Orbitals Hybridization in Molecules
More informationChapter 9. Chemical Bonding I: The Lewis Model. HIV-Protease. Lecture Presentation
Lecture Presentation Chapter 9 Chemical Bonding I: The Lewis Model HIV-Protease HIV-protease is a protein synthesized by the human immunodeficiency virus (HIV). This particular protein is crucial to the
More informationChapter 9. Molecular Geometry and Bonding Theories
Chapter 9. Molecular Geometry and Bonding Theories 9.1 Molecular Shapes Lewis structures give atomic connectivity: they tell us which atoms are physically connected to which atoms. The shape of a molecule
More informationChemical Bonding II. Molecular Geometry Valence Bond Theory Phys./Chem. Properties Quantum Mechanics Sigma & Pi bonds Hybridization MO theory
Chemical Bonding II Molecular Geometry Valence Bond Theory Phys./Chem. Properties Quantum Mechanics Sigma & Pi bonds ybridization MO theory 1 Molecular Geometry 3-D arrangement of atoms 2 VSEPR Valence-shell
More informationCHEMICAL BONDING. Chemical Bonds. Ionic Bonding. Lewis Symbols
CHEMICAL BONDING Chemical Bonds Lewis Symbols Octet Rule whenever possible, valence electrons in covalent compounds distribute so that each main-group element is surrounded by 8 electrons (except hydrogen
More informationHelpful Hints Lewis Structures Octet Rule For Lewis structures of covalent compounds least electronegative
Helpful Hints Lewis Structures Octet Rule Lewis structures are a basic representation of how atoms are arranged in compounds based on bond formation by the valence electrons. A Lewis dot symbol of an atom
More informationChapter 9: Molecular Geometry and Bonding Theories
Chapter 9: Molecular Geometry and Bonding Theories 9.1 Molecular Geometries -Bond angles: angles made by the lines joining the nuclei of the atoms in a molecule -Bond angles determine overall shape of
More informationLewis Dot Structures for Methane, CH 4 The central C atom is bonded by single bonds (-) to 4 individual H atoms
Chapter 10 (Hill/Petrucci/McCreary/Perry Bonding Theory and Molecular Structure This chapter deals with two additional approaches chemists use to describe chemical bonding: valence-shell electron pair
More informationChapter 10: Chemical Bonding II: Molecular Shapes; VSEPR, Valence Bond and Molecular Orbital Theories
C h e m i s t r y 1 A : C h a p t e r 1 0 P a g e 1 Chapter 10: Chemical Bonding II: Molecular Shapes; VSEPR, Valence Bond and Molecular Orbital Theories Homework: Read Chapter 10: Work out sample/practice
More informationChapter 9. Molecular Geometry and Bonding Theories
Chapter 9. Molecular Geometry and Bonding Theories PART I Molecular Shapes Lewis structures give atomic connectivity: they tell us which atoms are physically connected to which atoms. The shape of a molecule
More informationChapter 10 Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory
10.1 Artificial Sweeteners: Fooled by Molecular Shape 425 10.2 VSEPR Theory: The Five Basic Shapes 426 10.3 VSEPR Theory: The Effect of Lone Pairs 430 10.4 VSEPR Theory: Predicting Molecular Geometries
More informationBonding. Honors Chemistry 412 Chapter 6
Bonding Honors Chemistry 412 Chapter 6 Chemical Bond Mutual attraction between the nuclei and valence electrons of different atoms that binds them together. Types of Bonds Ionic Bonds Force of attraction
More informationAdapted from CHM 130 Maricopa County, AZ Molecular Geometry and Lewis Dot Formulas Introduction
Adapted from CHM 130 Maricopa County, AZ Molecular Geometry and Lewis Dot Formulas Introduction A chemical bond is an intramolecular (within the molecule) force holding two or more atoms together. Covalent
More informationbond energy- energy required to break a chemical bond -We can measure bond energy to determine strength of interaction
bond energy- energy required to break a chemical bond -We can measure bond energy to determine strength of interaction ionic compound- a metal reacts with a nonmetal Ionic bonds form when an atom that
More informationCHEMISTRY. Chapter 10 Theories of Bonding and Structure. The Molecular Nature of Matter. Jespersen Brady Hyslop SIXTH EDITION
CHEMISTRY The Molecular Nature of Matter SIXTH EDITION Jespersen Brady Hyslop Chapter 10 Theories of Bonding and Structure Copyright 2012 by John Wiley & Sons, Inc. Molecular Structures Molecules containing
More informationChapter 9 Molecular Geometry and Bonding Theories
Lecture Presentation Chapter 9 Geometry James F. Kirby Quinnipiac University Hamden, CT Shapes Lewis Structures show bonding and lone pairs, but do not denote shape. However, we use Lewis Structures to
More informationFill in the chart below to determine the valence electrons of elements 3-10
Chemistry 11 Atomic Theory IV Name: Date: Block: 1. Lewis Diagrams 2. VSEPR Lewis Diagrams Lewis diagrams show the bonding between atoms of a molecule. Only the outermost electrons of an atom (called electrons)
More informationChapter 7 Chemical Bonding and Molecular Structure
Chapter 7 Chemical Bonding and Molecular Structure Three Types of Chemical Bonding (1) Ionic: formed by electron transfer (2) Covalent: formed by electron sharing (3) Metallic: attraction between metal
More informationChapter 9. Molecular Geometry and Bonding Theories
9.1 Molecular Shapes Read Sec. 9.1 and 9.2, then complete the Sample and Practice Exercises in these sections. Sample Exercise 9.1 (p. 347) Use the VSEPR model to predict the molecular geometries of a)
More informationCHEMISTRY 112 LECTURE EXAM II Material
CHEMISTRY 112 LECTURE EXAM II Material Part I Chemical Bonding I Lewis Theory Chapter 9 pages 376-386 A. Drawing electron dot structures HOW TO: 1. Write e- dot structure for the individual atoms. 2. a)
More informationName Unit Three MC Practice March 15, 2017
Unit Three: Bonding & Molecular Geometry Name Unit Three MC Practice March 15, 2017 1. What is the hybridization of the oxygen atom in water? a) sp b) sp 2 c) sp 3 d) It is not hybridized 2. When a double
More informationStructure and Bonding of Organic Molecules
Chem 220 Notes Page 1 Structure and Bonding of Organic Molecules I. Types of Chemical Bonds A. Why do atoms forms bonds? Atoms want to have the same number of electrons as the nearest noble gas atom (noble
More informationCHEM 110 Exam 2 - Practice Test 1 - Solutions
CHEM 110 Exam 2 - Practice Test 1 - Solutions 1D 1 has a triple bond. 2 has a double bond. 3 and 4 have single bonds. The stronger the bond, the shorter the length. 2A A 1:1 ratio means there must be the
More informationTest Bank for Introductory Chemistry Essentials 5th Edition by Tro
Test Bank for Introductory Chemistry Essentials 5th Edition by Tro Sample Introductory Chemistry, 5e (Tro) Chapter 10 Chemical Bonding 10.1 True/False Questions 1) Bonding theories are used to predict
More informationPeriodic Trends. Homework: Lewis Theory. Elements of his theory:
Periodic Trends There are various trends on the periodic table that need to be understood to explain chemical bonding. These include: Atomic/Ionic Radius Ionization Energy Electronegativity Electron Affinity
More informationChemistry: The Central Science. Chapter 9: Molecular Geometry and Bonding Theory
Chemistry: The Central Science Chapter 9: Molecular Geometry and Bonding Theory The shape and size of a molecule of a particular substance, together with the strength and polarity of its bonds, largely
More informationHonors Chemistry Unit 6 ( )
Honors Chemistry Unit 6 (2017-2018) Lewis Dot Structures VSEPR Structures 1 We are learning to: 1. Represent compounds with Lewis structures. 2. Apply the VSEPR theory to determine the molecular geometry
More informationChapter 9. Molecular Geometry and Bonding Theories
Chapter 9 Molecular Geometry and Bonding Theories MOLECULAR SHAPES 2 Molecular Shapes Lewis Structures show bonding and lone pairs do not denote shape Use Lewis Structures to determine shapes Molecular
More informationChapter 8: Bonding. Section 8.1: Lewis Dot Symbols
Chapter 8: Bonding Section 8.1: Lewis Dot Symbols The Lewis electron dot symbol is named after Gilbert Lewis. In the Lewis dot symbol, the element symbol represents the nucleus and the inner electrons.
More informationWold of Chemistry Notes for Students [Chapter 12, page 1] Chapter 12 Chemical Bonding
Wold of Chemistry Notes for Students [Chapter 12, page 1] Chapter 12 Chemical Bonding 1) The History of the Development of the Period Table (Not in the book!) Similarities between the chemical and physical
More informationChapter 8. Chemical Bonding: Basic Concepts
Chapter 8. Chemical Bonding: Basic Concepts Chemical bond: is an attractive force that holds 2 atoms together and forms as a result of interactions between electrons found in combining atoms We rarely
More informationLewis Dot Formulas and Molecular Shapes
Lewis Dot Formulas and Molecular Shapes Introduction A chemical bond is an intramolecular (within the molecule) force holding two or more atoms together. Covalent chemical bonds are formed by valence electrons
More informationValence Bond Theory - Description
Bonding and Molecular Structure - PART 2 - Valence Bond Theory and Hybridization 1. Understand and be able to describe the Valence Bond Theory description of covalent bond formation. 2. Understand and
More informationCh 6 Chemical Bonding
Ch 6 Chemical Bonding What you should learn in this section (objectives): Define chemical bond Explain why most atoms form chemical bonds Describe ionic and covalent bonding Explain why most chemical bonding
More informationChemical Bonding Chapter 8
Chemical Bonding Chapter 8 Get your Clicker, 2 magnets, goggles and your handouts Nov 15 6:15 PM Recall that: Ionic-Involves the transfer of electrons - forms between a metal and a nonmetal Covalent-Involves
More informationChapter 9. Molecular Geometries and Bonding Theories. Lecture Presentation. John D. Bookstaver St. Charles Community College Cottleville, MO
Lecture Presentation Chapter 9 Theories John D. Bookstaver St. Charles Community College Cottleville, MO Shapes The shape of a molecule plays an important role in its reactivity. By noting the number of
More informationEXAM II Material. Part I Chemical Bonding I Lewis Theory Chapter 9 pages A. Drawing electron dot structures HOW TO:
CHEMISTRY 112 LECTURE EXAM II Material Part I Chemical Bonding I Lewis Theory Chapter 9 pages 376-386 A. Drawing electron dot structures HOW TO: 1. Write e- dot structure for the individual atoms. 2. a)
More informationChapter 9. Molecular Geometries and Bonding Theories. Lecture Presentation. John D. Bookstaver St. Charles Community College Cottleville, MO
Lecture Presentation Chapter 9 Theories John D. Bookstaver St. Charles Community College Cottleville, MO Shapes The shape of a molecule plays an important role in its reactivity. By noting the number of
More informationChapter 9. Lewis Theory-VSEPR Valence Bond Theory Molecular Orbital Theory
Chapter 9 Lewis Theory-VSEPR Valence Bond Theory Molecular Orbital Theory Problems with Lewis Theory Lewis theory generally predicts trends in properties, but does not give good numerical predictions.
More informationAP Chemistry. Unit #7. Chemical Bonding & Molecular Shape. Zumdahl Chapters 8 & 9 TYPES OF BONDING BONDING. Discrete molecules formed
AP Chemistry Unit #7 Chemical Bonding & Molecular Shape Zumdahl Chapters 8 & 9 TYPES OF BONDING BONDING INTRA (Within (inside) compounds) STRONG INTER (Interactions between the molecules of a compound)
More informationExperiment 21 Lewis structures and VSEPR Theory
Experiment 21 Lewis structures and VSEPR Theory Introduction 1. Lewis Structures and Formal Charge LG.N. Lewis, at the University of California at Berkeley devised a simple way to understand the nature
More informationCHEMISTRY Matter and Change Section 8.1 The Covalent Bond
CHEMISTRY Matter and Change Section Chapter 8: Covalent Bonding CHAPTER 8 Table Of Contents Section 8.2 Section 8.3 Section 8.4 Section 8.5 Naming Molecules Molecular Structures Molecular Shapes Electronegativity
More informationCovalent Compounds: Bonding Theories and Molecular Structure
CHM 123 Chapter 8 Covalent Compounds: Bonding Theories and Molecular Structure 8.1 Molecular shapes and VSEPR theory VSEPR theory proposes that the geometric arrangement of terminal atoms, or groups of
More informationEx. 1) F F bond in F = 0 < % covalent, no transfer of electrons
#60 Notes Unit 8: Bonding Ch. Bonding I. Bond Character Bonds are usually combinations of ionic and covalent character. The electronegativity difference is used to determine a bond s character. Electronegativity
More informationThe Shapes of Molecules. Chemistry II
The Shapes of Molecules Chemistry II Lewis Structures DEFINITIN: A structure of a molecule showing how the valence electrons are arranged. 1) nly the valence electrons appear in a Lewis structure. 2) The
More informationChapter 9 Molecular Geometry and Bonding Theories
Chapter 9 Molecular Geometry and Bonding Theories molecular shapes the VSEPR model molecular shape and molecular polarity covalent bonding and orbital overlap hybrid orbitals multiple bonds 9.1 Molecular
More informationChapter 9 The Shapes of Molecules Cocaine
Chapter 9 The Shapes of Molecules 1 Cocaine 10.1 Depicting Molecules & Ions with Lewis Structures 2 Number of Covalent Bonds 3 The number of covalent bonds can be determined from the number of electrons
More information8.1 Types of Chemical Bonds List and define three types of bonding. chapter 8 Bonding General Concepts.notebook. September 10, 2015
chapter 8 Bonding General Concepts.notebook Chapter 8: Bonding: General Concepts Mar 13 11:15 AM 8.1 Types of Chemical Bonds List and define three types of bonding. Bonds are forces that hold groups of
More informationVSEPR. Ch10. Valence Shell Electron Pair Repulsion theory allows you to predict molecular shape. Lewis Dot theory extended to 3 dimensions.
Ch10 VSEPR Valence Shell Electron Pair Repulsion theory allows you to predict molecular shape. Lewis Dot theory extended to 3 dimensions. version 1.5 Nick DeMello, PhD. 2007-2016 Valence Shell Electron
More informationCh 10 Chemical Bonding, Lewis Structures for Ionic & Covalent Compounds, and Predicting Shapes of Molecules
Fructose Water Ch 10 Chemical Bonding, Lewis Structures for Ionic & Covalent Compounds, and Predicting Shapes of Molecules Carbon Dioxide Ammonia Title and Highlight TN Ch 10.1 Topic: EQ: Right Side NOTES
More informationChapter 8: Covalent Bonding. Chapter 8
: Covalent Bonding Bonding Ionic Bonding - attracted to each other, but not fully committed Covalent Bonding - fully committed, and shares everything Two methods to gain or lose valence electrons: Transfer
More informationChapter 4. Molecular Structure and Orbitals
Chapter 4 Molecular Structure and Orbitals Chapter 4 Table of Contents (4.1) (4.2) (4.3) (4.4) (4.5) (4.6) (4.7) Molecular structure: The VSEPR model Bond polarity and dipole moments Hybridization and
More informationChapter 8. Chemical Bonding: Basic Concepts
Chapter 8. Chemical Bonding: Basic Concepts Chemical bond: is an attractive force that holds 2 atoms together and forms as a result of interactions between electrons found in combining atoms We rarely
More informationChapter 7. Ionic & Covalent Bonds
Chapter 7 Ionic & Covalent Bonds Ionic Compounds Covalent Compounds 7.1 EN difference and bond character >1.7 = ionic 0.4 1.7 = polar covalent 1.7 Electrons not shared at
More informationChapter 9: Molecular Geometries and Bonding Theories Learning Outcomes: Predict the three-dimensional shapes of molecules using the VSEPR model.
Chapter 9: Molecular Geometries and Bonding Theories Learning Outcomes: Predict the three-dimensional shapes of molecules using the VSEPR model. Determine whether a molecule is polar or nonpolar based
More informationUnit Six --- Ionic and Covalent Bonds
Unit Six --- Ionic and Covalent Bonds Electron Configuration in Ionic Bonding Ionic Bonds Bonding in Metals Valence Electrons Electrons in the highest occupied energy level of an element s atoms Examples
More informationChapter 10. Geometry
Chapter 10 Molec cular Geometry 1 CHAPTER OUTLINE Molecular Geometry Molecular Polarity VSEPR Model Summary of Molecular Shapes Hybridization Molecular Orbital Theory Bond Angles 2 MOLECULAR GEOMETRY Molecular
More informationMolecular Geometry and intermolecular forces. Unit 4 Chapter 9 and 11.2
1 Molecular Geometry and intermolecular forces Unit 4 Chapter 9 and 11.2 2 Unit 4.1 Chapter 9.1-9.3 3 Review of bonding Ionic compound (metal/nonmetal) creates a lattice Formula doesn t tell the exact
More informationChemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals
Chemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals 1 Chemical Bonding II Molecular Geometry (10.1) Dipole Moments (10.2) Valence Bond Theory (10.3) Hybridization of Atomic Orbitals
More informationMolecular shape is only discussed when there are three or more atoms connected (diatomic shape is obvious).
Chapter 10 Molecular Geometry (Ch9 Jespersen, Ch10 Chang) The arrangement of the atoms of a molecule in space is the molecular geometry. This is what gives the molecules their shape. Molecular shape is
More informationChapter 6. Preview. Objectives. Molecular Compounds
Section 2 Covalent Bonding and Molecular Compounds Preview Objectives Molecular Compounds Formation of a Covalent Bond Characteristics of the Covalent Bond The Octet Rule Electron-Dot Notation Lewis Structures
More informationEXPERIMENT #13 Lewis Structures and Molecular Geometry
OBJECTIVES: EXPERIMENT #13 s and Draw Lewis structures of atoms, ions, and molecules Build models of linear, trigonal planar tetrahedral, trigonal bipyramidal, and octahedral arrangements of electron pairs
More informationChapter 8 Covalent Boding
Chapter 8 Covalent Boding Molecules & Molecular Compounds In nature, matter takes many forms. The noble gases exist as atoms. They are monatomic; monatomic they consist of single atoms. Hydrogen chloride
More information10-1. The Shapes of Molecules, chapter 10
10-1 The Shapes of Molecules, chapter 10 The Shapes of Molecules; Goals 10.1 Depicting Molecules and Ions with Lewis Structures 10.2 Valence-Shell Electron-Pair Repulsion (VSEPR) Theory 10.3 Molecular
More informationNotes: Covalent Bonding
Name Chemistry Pre-AP Notes: Covalent Bonding Period The main focus of this unit is on the covalent bond; however, we will briefly treat the ionic and metallic bond as well. I. Chemical Bonding Overview
More informationLewis Theory of Shapes and Polarities of Molecules
Lewis Theory of Shapes and Polarities of Molecules Sulfanilamide Lewis Structures and the Real 3D-Shape of Molecules Molecular Shape or Geometry The way in which atoms of a molecule are arranged in space
More informationMolecular Models: The shape of simple molecules and ions
Molecular Models: The shape of simple molecules and ions Background The shape of a molecule is very important when investigating its properties and reactivity. For example, compare CO 2 and SO 2. Carbon
More informationChapters 9&10 Structure and Bonding Theories
Chapters 9&10 Structure and Bonding Theories Ionic Radii Ions, just like atoms, follow a periodic trend in their radii. The metal ions in a given period are smaller than the non-metal ions in the same
More informationDownloaded from
Points to Remember Class: XI Chapter Name: Chemical Bonding and Molecular Structure Top Concepts 1. The attractive force which holds together the constituent particles (atoms, ions or molecules) in chemical
More informationClass XI Chapter 4 Chemical Bonding and Molecular Structure Chemistry
Class XI Chapter 4 Chemical Bonding and Molecular Structure Chemistry Question 4.1: Explain the formation of a chemical bond. A chemical bond is defined as an attractive force that holds the constituents
More informationClass XI Chapter 4 Chemical Bonding and Molecular Structure Chemistry
Class XI Chapter 4 Chemical Bonding and Molecular Structure Chemistry Question 4.1: Explain the formation of a chemical bond. A chemical bond is defined as an attractive force that holds the constituents
More informationSubtopic 4.2 MOLECULAR SHAPE AND POLARITY
Subtopic 4.2 MOLECULAR SHAPE AND POLARITY 1 LEARNING OUTCOMES (covalent bonding) 1. Draw the Lewis structure of covalent molecules (octet rule such as NH 3, CCl 4, H 2 O, CO 2, N 2 O 4, and exception to
More informationWhat is a Bond? Chapter 8. Ionic Bonding. Coulomb's Law. What about covalent compounds?
Chapter 8 What is a Bond? A force that holds atoms together. Why? We will look at it in terms of energy. Bond energy- the energy required to break a bond. Why are compounds formed? Because it gives the
More informationIntroduction to Chemical Bonding
Chemical Bonding Introduction to Chemical Bonding Chemical bond! is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together Why are most
More informationC H E M 1 CHEM 101-GENERAL CHEMISTRY CHAPTER 7 CHEMICAL BONDING & MOLECULAR STRUCTURE INSTR : FİLİZ ALSHANABLEH
C H E M 1 CHEM 101-GENERAL CHEMISTRY CHAPTER 7 CHEMICAL BONDING & MOLECULAR STRUCTURE 0 1 INSTR : FİLİZ ALSHANABLEH CHAPTER 7 CHEMICAL BONDING & MOLECULAR STRUCTURE The Ionic Bond Formation of Ions The
More informationBonding. Polar Vs. Nonpolar Covalent Bonds. Ionic or Covalent? Identifying Bond Types. Solutions + -
Chemical Bond Mutual attraction between the nuclei and valence electrons of different atoms that binds them together. Bonding onors Chemistry 412 Chapter 6 Types of Bonds Ionic Bonds Force of attraction
More informationInstant download Test bank for Chemistry The Central Science 10th Edition by Brown, LeMay, Bursten CLICK HERE
Chemistry, 10e (Brown) Chapter 9, Molecular Geometry and Bonding Theories Instant download Test bank for Chemistry The Central Science 10th Edition by Brown, LeMay, Bursten CLICK HERE http://testbankair.com/download/test-bank-for-chemistry-the-central-science-10th-edition-by-brown-lemay-bursten/
More informationActivity Hybrid Atomic Orbitals
Activity 201 8 Hybrid Atomic Orbitals Directions: This Guided Learning Activity (GLA) discusses Hybrid Atomic Orbitals, which are the basis for Valence Bond Theory. Part A introduces σ- and π-bonds. Part
More informationCovalent Bonding. In nature, only the noble gas elements exist as uncombined atoms. All other elements need to lose or gain electrons
In nature, only the noble gas elements exist as uncombined atoms. They are monatomic - consist of single atoms. All other elements need to lose or gain electrons To form ionic compounds Some elements share
More informationCHEMISTRY XL-14A CHEMICAL BONDS
CHEMISTRY XL-14A CHEMICAL BONDS July 16, 2011 Robert Iafe Office Hours 2 July 18-July 22 Monday: 2:00pm in Room MS-B 3114 Tuesday-Thursday: 3:00pm in Room MS-B 3114 Chapter 2 Overview 3 Ionic Bonds Covalent
More informationChapters 8 and 9. Octet Rule Breakers Shapes
Chapters 8 and 9 Octet Rule Breakers Shapes Bond Energies Bond Energy (review): The energy needed to break one mole of covalent bonds in the gas phase Breaking bonds consumes energy; forming bonds releases
More informationChapter 10. VSEPR Model: Geometries
Chapter 10 Molecular Geometry VSEPR Model: Geometries Valence Shell Electron Pair Repulsion Theory Electron pairs repel and get as far apart as possible Example: Water Four electron pairs Farthest apart
More informationChemical Bonding. Section 1 Introduction to Chemical Bonding. Section 2 Covalent Bonding and Molecular Compounds
Chemical Bonding Table of Contents Section 1 Introduction to Chemical Bonding Section 2 Covalent Bonding and Molecular Compounds Section 3 Ionic Bonding and Ionic Compounds Section 4 Metallic Bonding Section
More informationFind the difference in electronegativity between the hydrogen and chlorine atoms
Answers Questions 16.2 Molecular polarity 1. Write a dot diagram for the HCl molecule. Find the difference in electronegativity between the hydrogen and chlorine atoms Difference in electronegativity =
More informationChapter 9. and Bonding Theories
Chemistry, The Central Science, 11th edition Theodore L. Brown, H. Eugene LeMay, Jr., and Bruce E. Bursten Chapter 9 Theories John D. Bookstaver St. Charles Community College Cottleville, MO Shapes The
More informationPART 3 Chemical Bonds, Valence Bond Method, and Molecular Shapes. Reference: Chapter 9 10 in textbook
PART 3 Chemical Bonds, Valence Bond Method, and Molecular Shapes Reference: Chapter 9 10 in textbook 1 Valence Electrons Valence ae Electron Define: the outer shell electrons Important for determination
More informationShapes of Molecules and Hybridization
Shapes of Molecules and Hybridization A. Molecular Geometry Lewis structures provide us with the number and types of bonds around a central atom, as well as any NB electron pairs. They do not tell us the
More informationLecture outline: Section 9. theory 2. Valence bond theory 3. Molecular orbital theory. S. Ensign, Chem. 1210
Lecture outline: Section 9 Molecular l geometry and bonding theories 1. Valence shell electron pair repulsion theory 2. Valence bond theory 3. Molecular orbital theory 1 Ionic bonding Covalent bonding
More informationChapter 8. Molecular Shapes. Valence Shell Electron Pair Repulsion Theory (VSEPR) What Determines the Shape of a Molecule?
PowerPoint to accompany Molecular Shapes Chapter 8 Molecular Geometry and Bonding Theories Figure 8.2 The shape of a molecule plays an important role in its reactivity. By noting the number of bonding
More informationChapter 6. Chemical Bonding
Chapter 6 Chemical Bonding Section 6.1 Intro to Chemical Bonding 6.1 Objectives Define chemical bond. Explain why most atoms form chemical bonds. Describe ionic and covalent bonding. Explain why most chemical
More informationCHAPTER 12: CHEMICAL BONDING
CHAPTER 12: CHEMICAL BONDING Problems: 1-26, 27c, 28, 33-34, 35b, 36(a-c), 37(a,b,d), 38a, 39-40, 41-42(a,c), 43-58, 67-74 12.1 THE CHEMICAL BOND CONCEPT chemical bond: what holds atoms or ions together
More informationWhat Do Molecules Look Like?
What Do Molecules Look Like? The Lewis Dot Structure approach provides some insight into molecular structure in terms of bonding, but what about 3D geometry? Recall that we have two types of electron pairs:
More informationCHM 151LL: Geometry of Covalent Compounds
CM 151LL: Geometry of Covalent Compounds Introduction Octet Rule A Lewis structure (or electrondot formula) is a twodimensional structural formula showing the arrangement of electrons around atoms in covalently
More information