1 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. The Lewis structures are useful for visualization, but do not reveal the bent structure for water (105°), the pyramidal shape for ammonia, or the tetrahedral Unit 3 : Chemical Bonding Part 1 : Covalent Bond Chemical compounds are formed by the joining of two or more atoms. A stable compound occurs when the total energy of the combination has lower energy than the separated atoms. The exchange interaction leads to a strong bond for the hydrogen molecule with dissociation energy 4.52 eV at a separation of 0.074 nm. The potential energy of the anti-bonding orbital shown gives some insight into why a third hydrogen atom cannot bond to the two atoms of the hydrogen molecule. It would be in an anti-bonding situation with one of the other hydrogen atoms and would therefore be repelled. We say that the bond in the hydrogen molecule is "saturated" because it cannot accept another bond. http://hyperphysics.phyastr.gsu.edu/hbase/Chemical/bond.html [6 August 2018] Covalent bond: bond in which one or more pairs of electrons are shared by two atoms. In the idealized covalent bond, two atoms share a pair of electrons, closing the shell for each of them. The atoms share a pair of electrons, and that pair is referred to as a bonding pair. The pairs of electrons which do not participate in the bond have traditionally been called "lone pairs". A single bond can be represented by the two dots of the bonding pair, or by a single line which represents that pair. The single line representation for a bond is commonly used in drawing Lewis structures for molecules. A single bond can be represented by the two dots of the bonding pair, or by a single line which represents that pair. The single line representation for a bond is commonly used in drawing Lewis structures for molecules.
2 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Covalent Bond Exercise. Lewis Dot Diagrams of Selected Elements Electron Distributions Into Shells for the First Three Periods Complete the following chart of covalent molecules. The number of Valence electron of the element Lewis Structure Molecular Electron dot formula Dash formula Formula F F O O N N C H N H O F S O N 2 O 2 N 4 O 2 N 5 O
1 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Unit 3 : Chemical Bonding Part 2 : Formal Charges Formal Charges An atom’s formal charge is the electrical charge difference between the valence electron in an isolated atom and the number of electron assigned to that atom in a Lewis structure. Not all atoms within a neutral molecule need be neutral. The location of any charges is often useful for understanding or predicting reactivity. Identifying formal charges helps you keep track of the electrons. It's a good idea to get into the habit of automatically labelling the formal charges on any atoms within a molecule. Knowing the formal charge on a particular atom in a structure is an important part of keeping track of the electrons and is important for establishing and predicting the reactivity. Formal charge equation formally compares the number of valence electrons in an isolated neutral atom with the number of valence electrons around the atom in the molecule: Example molecule of interest. Formal charge on oxygen: Group number = 6 Number of covalent bonds = 2 Number of lone pair electrons = 4 Formal charges for all the different atoms http://www.chem.ucalgary.ca/courses/350/Carey5th/Ch01/c h1-3-2.html#equation [9 September 2018] The formal charge on an atom in a molecule reflects the electron count associated with the atom compared to the isolated neutral atom. If the atom has given away electrons it will be +ve and if it has gained electrons it will be -ve.
2 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Formal Charges Exercise. Complete the following calculate the formal charges of covalent molecules. Covalent Compounds Lewis Structure Formal Charges Covalent Compounds Lewis Structure Formal Charges O3 NCl3 CH2O OCS CO3 2 H2O2 HNO3 CH3COO NO3 CN N2O4 ClO3 H2SO4 OH SO4 2 OF2 HClO3 N2F2
1 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Unit 3 : Chemical Bonding Part 3 : Co-ordinate (Dative covalent) Bonding Co-ordinate (Dative covalent) Bonding A covalent bond is formed by two atoms sharing a pair of electrons. The atoms are held together because the electron pair is attracted by both of the nuclei. In the formation of a simple covalent bond, each atom supplies one electron to the bond - but that doesn't have to be the case. A co-ordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom. Representing co-ordinate bonds In simple diagrams, a co-ordinate bond is shown by an arrow. The reaction between ammonia and hydrogen chloride Dissolving hydrogen chloride in water to make hydrochloric acid Something similar happens. A hydrogen ion (H) is transferred from the chlorine to one of the lone pairs on the oxygen atom. The reaction between ammonia and boron trifluoride, BF3 The lone pair on the nitrogen of an ammonia molecule can be used to overcome that deficiency, and a compound is formed involving a co-ordinate bond. Using lines to represent the bonds, this could be drawn more simply as: https://www.chemguide.co.uk/atoms/bonding/dative.html [9 September 2018]. The H3O ion is variously called the hydroxonium ion, the hydronium ion or the oxonium ion. The arrow points () from the atom donating the lone pair to the atom accepting it.
2 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Co-ordinate (Dative covalent) Bonding Exercise. Complete the following Lewis Structure for simplicity, the formal charge have been omitted.. Lewis Structure for simplicity, the formal charge have been omitted. Compound Coordinate covalent bond Compound Coordinate covalent bond SO2 BF3NH3 SO3 NO2 H2SO4 N2O4 H2SO3 HClO3 HClO4 NO2
3 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Lewis Structure for simplicity, the formal charge have been omitted. Compound Coordinate covalent bond Compound Coordinate covalent bond HClO3 NO3 HClO2 CO3 2 HClO PO4 3 H3PO4 CH3COO H3PO3 HNO2 H3PO2 HNO3
1 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Unit 3 : Chemical Bonding Part 4 : Resonance Structures Resonance Structures Resonance structures are a set of two or more Lewis Structures that collectively describe the electronic bonding a single polyatomic species including fractional bonds and fractional charges. Resonance structure are capable of describing delocalized electrons that cannot be expressed by a single Lewis formula with an integer number of covalent bonds. https://chem.libretexts.org/Textbook_Maps/General_Chemistry/Map%3A_Chemistry_- _The_Central_Science_(Brown_et_al.)/08._Basic_Concepts_of_Chemical_Bonding/8.6%3A_ Resonance_Structures [9 September 2018]. Such is the case for ozone (O3), an allotrope of oxygen with a V-shaped structure and an O–O–O angle of 117.5°. Ozone (O3) The central oxygen has only 6 electrons. We must convert one lone pair on a terminal oxygen atom to a bonding pair of electrons—but which one? Depending on which one we choose, we obtain either. Which is correct? In fact, neither is correct. Both predict one O–O single bond and one O=O double bond. As you will learn, if the bonds were of different types (one single and one double, for example), they would have different lengths. It turns out, however, that both O–O bond distances are identical, 127.2 pm, which is shorter than a typical O–O single bond (148 pm) and longer than the O=O double bond in O2 (120.7 pm). Equivalent Lewis dot structures, such as those of ozone, are called resonance structures. The position of the atoms is the same in the various resonance structures of a compound, but the position of the electrons is different. Double-headed arrows link the different resonance structures of a compound: The double-headed arrow indicates that the actual electronic structure is an average of those shown, not that the molecule oscillates between the two structures. The Carbonate (CO3 2−) Ion As with ozone, none of these structures describes the bonding exactly. Each predicts one carbon–oxygen double bond and two carbon–oxygen single bonds, but experimentally all C–O bond lengths are identical. We can write resonance structures (in this case, three of them) for the carbonate ion: The Nitrate (NO3 −) ion Resonance is a mental exercise and method within the Valence Bond Theory of bonding that describes the delocalization of electrons within molecules. These structures are written with a double-headed arrow between them, indicating that none of the Lewis structures accurately describes the bonding but that the actual structure is an average of the individual resonance structures. The actual structure is an average of these three resonance structures.
2 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Resonance structures Exercise. Fill in the blank with each of a resonance structures and a resonance hybrid. Compounds A resonance Structure A resonance hybrid O3 SO2 SO3 H2SO4 H2SO3 CN BF3 C6H6
1 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Unit 3 : Chemical Bonding Part 5 : Naming Covalent Compounds Prefix Names of Covalent Compounds Covalent compounds are named in different ways than are ionic compounds (although there is some overlap). Many of these compounds have common names such as "methane", "ammonia" and "water". However, simple covalent compounds are generally named by using prefixes to indicate how many atoms of each element are shown in the formula. Also, the ending of the last (most negative) element is changed to -ide. Naming Covalent Compounds Naming Binary Covalent Compounds When a pair of elements form more than one type of covalent compound, Greek prefixes are used to indicate how many of each element are in a compound. Some of the Greek prefixes are given in the table below: Prefix Number of Particular Element Prefix Number of Particular Element mono 1 hexa 6 di 2 hepta 7 tri 3 octa 8 tetra 4 nona 9 penta 5 deca 10 Rules for Binary Covalent Compounds 1. The prefix mono is never used for naming the first element of a compound. 2. The final o or a of a prefix is often dropped when the element begins with a vowel. How do you know which element goes first? The element that comes first in the following list "goes" first. B, Si, C, Sb, As, P, N, H, Te, Se, S, I, Br, Cl, O, F https://www.grandinetti.org/naming-covalent-compounds.[12 October 2018]. Naming Acids, Oxyacids and Their Salts If the anion does not contain oxygen, then the acid is named with the prefix hydro- and the suffix -ic. For example, when gaseous HCl is dissolved in H2O, it forms hydrochloric acid. HCN in H2O is hydrocyanic acid. Before we learn the rule for naming oxyacids, let's learn the rules for naming oxyanions. What are oxyanions? They are anions formed from oxygen and a nonmetal. Here are some examples: ClO4 , ClO3 , ClO , 2 ClO , 2 SO , 4 2 SO . 3 There are two rules for naming these: 1. If there are only two members in the same series, then the anion with the least number of oxygens ends in -ite, and the anion with the most ends in -ate. For example, 2 SO3 is sulfite and 2 SO4 is sulfate. 2. When there are more than two oxyanions in a series, hypo- (less than) and per- (more than) are used as prefixes. Here are some examples: ClO is hypochlorite ClO is chlorite 2 ClO is chlorate 3 ClO is perchlorate 4 Finally, H2O, which according to the rules should be called dihydrogen monoxide is always called water, and NH3, or nitrogen trihydride, is always called ammonia.
2 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Finally, here are the rules for naming acids of oxyanions. If the anion name ends in -ate, then the acid name ends in -ic or -ric. If the anion name ends in -ite, then the acid name ends in -ous. Here are examples of the last three rules: Acid Anion Acid Name Acid Anion Acid Name HClO hypochlorite hypochlorous acid HClO3 chlorate chloric acid HClO2 chlorite chlorous acid HClO4 perchlorate perchloric acid Fill in the blank with each of a covalent structure and naming compounds. Element Covalent Compounds Naming Covalent Compounds Molecular formula Lewis structure H S Be F B F C Cl C O C 2O N H N O
3 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Element Covalent Compounds Naming Covalent Compounds Molecular formula Lewis structure 2N O N 2O 2N 4O P Cl P Cl
1 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Unit 3 : Chemical Bonding Part 6 : Lewis structure of several common oxoacids Lewis structure of several common oxoacids An oxoacid (sometimes called an oxyacid) is an acid that contains oxygen. To be more specific, an oxoacid is an acid that: 1. contains oxygen 2. contains at least one other element 3. has at least one hydrogen atom bonded to oxygen 4. forms an ion by the loss of one or more protons in solution. All oxoacids have the acidic hydrogen bound to an oxygen atom, so bond strength (length) is not a factor, similar to binary nonmetal acids; instead, the main determining factor for an oxacid’s relative strength has to do with the central atom’s electronegativity (X), as well as the number of O atoms around that central atom. https://courses.lumenlearning.com/introchem/chapter/oxoacids/[10 November 2018]. The central atom in the oxoacids is sp3 hybridized. Every oxoacid has essentially one X-OH bond. Whereas most oxoacids have X=O bonds present in them. This double bond between oxygen and halogen is d pi-pi in nature. In the series of oxoacids, the first member possesses high acidic strength. This is only due to high electronegativity and small size of the halogen atom. The acidic strength increases with increase in the oxidation number of halogens. https://byjus.com/chemistry/oxoacids-of-halogens/[10 November 2018]. Fill in the blank with each of a covalent structure and naming compounds. Oxoacids Lewis Structure for simplicity, the formal charge have been omitted. Name of oxoacids Lewis Structure Lewis Structure of some Common oxoacids HNO2 HNO3 H3PO4
2 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Oxoacids Lewis Structure for simplicity, the formal charge have been omitted. Name of oxoacids Lewis Structure Lewis Structure of some Common oxoacids H3PO3 H3PO2 H2SO4 H2SO3 HClO4 HClO3
3 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Oxoacids Lewis Structure for simplicity, the formal charge have been omitted. Name of oxoacids Lewis Structure Lewis Structure of some Common oxoacids HClO2 HClO HBrO3 HIO3 H2CO3
1 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Unit 3 : Chemical Bonding Part 7.1 : Molecular Geometry Predicting the Shapes of Molecules There is no direct relationship between the formula of a compound and the shape of its molecules. The shapes of these molecules can be predicted from their Lewis structures, however, with a model developed about 30 years ago, known as the valence-shell electron-pair repulsion (VSEPR) theory. The VSEPR theory assumes that each atom in a molecule will achieve a geometry that minimizes the repulsion between electrons in the valence shell of that atom. The five compounds shown in the figure below can be used to demonstrate how the VSEPR theory can be applied to simple molecules. There are only two places in the valence shell of the central atom in BeF2 where electrons can be found. Repulsion between these pairs of electrons can be minimized by arranging them so that they point in opposite directions. Thus, the VSEPR theory predicts that BeF2 should be a linear molecule, with a 180o angle between the two Be-F bonds. There are three places on the central atom in boron trifluoride (BF3) where valence electrons can be found. Repulsion between these electrons can be minimized by arranging them toward the corners of an equilateral triangle. The VSEPR theory therefore predicts a trigonal planar geometry for the BF3 molecule, with a F-B-F bond angle of 120o. BeF2 and BF3 are both two-dimensional molecules, in which the atoms lie in the same plane. If we place the same restriction on methane (CH4), we would get a square-planar geometry in which the H-CH bond angle is 90o. If we let this system expand into three dimensions, however, we end up with a tetrahedral molecule in which the H-C-H bond angle is 109o28'. Repulsion between the five pairs of valence electrons on the phosphorus atom in PF5 can be minimized by distributing these electrons toward the corners of a trigonal bipyramid. Three of the positions in a trigonal bipyramid are labeled equatorial because they lie along the equator of the molecule. The other two are axial because they lie along an axis perpendicular to the equatorial plane. The angle between the three equatorial positions is 120o, while the angle between an axial and an equatorial position is 90o. There are six places on the central atom in SF6 where valence electrons can be found. The repulsion between these electrons can be minimized by distributing them toward the corners of an octahedron. The term octahedron literally means "eight sides," but it is the six corners, or vertices, that interest us. To imagine the geometry of an SF6 molecule, locate fluorine atoms on opposite sides of the sulfur atom along the X, Y, and Z axes of an XYZ coordinate system. http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/vsepr.html,[20 November 2018].
2 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Incorporating Double and Triple Bonds Into the VSEPR Theory Compounds that contain double and triple bonds raise an important point: The geometry around an atom is determined by the number of places in the valence shell of an atom where electrons can be found, not the number of pairs of valence electrons. Consider the Lewis structures of carbon dioxide (CO2) and the carbonate (CO3 2-) ion, for example. There are four pairs of bonding electrons on the carbon atom in CO2, but only two places where these electrons can be found. (There are electrons in the C=O double bond on the left and electrons in the double bond on the right.) The force of repulsion between these electrons is minimized when the two C=O double bonds are placed on opposite sides of the carbon atom. The VSEPR theory therefore predicts that CO2 will be a linear molecule, just like BeF2, with a bond angle of 180o. The Lewis structure of the carbonate ion also suggests a total of four pairs of valence electrons on the central atom. But these electrons are concentrated in three places: The two C-O single bonds and the C=O double bond. Repulsions between these electrons are minimized when the three oxygen atoms are arranged toward the corners of an equilateral triangle. The CO3 2- ion should therefore have a trigonalplanar geometry, just like BF3, with a 120o bond angle. The Role of Nonbonding Electrons in the VSEPR Theory The valence electrons on the central atom in both NH3 and H2O should be distributed toward the corners of a tetrahedron, as shown in the figure below. Our goal, however, isn't predicting the distribution of valence electrons. It is to use this distribution of electrons to predict the shape of the molecule. Until now, the two have been the same. Once we include nonbonding electrons, that is no longer true. https://courses.lumenlearning.co m/boundlesschemistry/chapter/moleculargeometry/,[20 November 2018]. Molecular geometry describes the threedimensional arrangement of atoms in a molecule. Molecular geometry may be predicted using VSEPR and Lewis structures and verified using spectroscopy and diffraction.
3 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Draw the Lewis and VSEPR structures for the following compounds and label them with their geometry. Molecular formula Lewis Structure Electron and Molecular Geometry on Central Atom Bonding Regions Lone Pairs VSEPR Molecular Geometry BeF2 CO2 CO N2 BF3 2 CO3 O3 SO2 O2
4 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Molecular formula Lewis Structure Electron and Molecular Geometry on Central Atom Bonding Regions Lone Pairs VSEPR Molecular Geometry CH4 2 SO4 NH3 H O3 H O2 2 ICl OH PF5 SF4 4 IF
5 Chemistry for World-Class Standard School: Created by ; Mr.Suthat Chanprakhon ; since 2018 Name……………………………………………..…Id………………..Class………..Ordinal……….. Molecular formula Lewis Structure Electron and Molecular Geometry on Central Atom Bonding Regions Lone Pairs VSEPR Molecular Geometry ClF3 3 I XeF2 SF6 PF6 2 SiF6 BrF5 2 SbCl5 XeF4 4 ICl