CH 11 TOPIC 20 MOLECULAR COMPOUNDS COVALENT BONDS 1 You have mastered this topic when you can: 1) define or describe MOLECULAR ELEMENTS and DIATOMIC ELEMENTS. 2) define or describe MOLECULAR COMPOUND and COVALENT BONDING. 3) predict the formation of a covalent bond from the formula of a compound. 4) list the ionic, single, double and triple bonds in order of increasing or decreasing strength. 5) draw ELECTRON DOT DIAGRAMS and STRUCTURAL FORMULAE for simple MOLECULAR COMPOUNDS. MOLECULES I) When two or more non-metal atoms bond together they create a particle called a MOLECULE. H 2(g), N 2(g), O 2(g), H 2 O (l), NH 3(g), CO (g), CO 2(g), SO 2(g), CH 3 OH (l), C 2 H 5 OH (l), C 12 H 22 O 11(s), etc. A) There are two kinds of molecules: MOLECULAR ELEMENTS and MOLECULAR COMPOUNDS. 1) MOLECULAR ELEMENTS. a) DIATOMIC ELEMENTS. MEMORIZE THEM!! i) When in their pure unreacted elemental form, the above elements exist as diatomic molecules. When you read or hear oxygen, you must think O 2 not O, when you read or hear nitrogen, you must think N 2 not N, etc. If you are required to use O, the question will use the phrase oxygen atoms, if you are required to use N, the question will use the phrase nitrogen atoms, etc. 2) MOLECULAR COMPOUNDS. II) WHY DO MOLECULES FORM? A) Molecules form for the same reasons ionic compounds form to create stable octets by filling valence shells of each atom of the molecule. The octet rule states that every atom must have eight electrons in its valence shell to create a stable octet. The exceptions to this rule are hydrogen and helium atoms, and lithium and beryllium ions which require only two valence electrons to fill their valence shells: For hydrogen and helium atoms, lithium and beryllium ions two valence electrons creates a stable octet. MAKING MOLECULAR COMPOUNDS I) COVALENT BONDS A) Molecules are formed when two or more non-metal atoms share one or more pairs of valence electrons in order to fill their valence shells to create a stable octet. B) Consider the diatomic element hydrogen, H 2. H 2 is a very common molecular element and is useful in describing why non-metal atoms create stable octets by sharing valence electrons. If one H atom transferred its lone valence electron to the other H atom, an ionic bond between a cation and an anion would be created as illustrated by the electron dot diagram given below. atoms electron dot diagram H i + i H [H]+ [:H] 1) If this were the case, H 2 would be composed of ions and thus would be an electrolyte. Research has shown that H 2 is a non-electrolyte, which means it is not composed of ions and thus it is not held together by an ionic bond, rather, it is held together by another type of bond. Since H 2 is not held together by a bond formed by a transfer of valence electrons from one atom to another, it must be held together by a bond formed by sharing valence. This type of bond is called a COVALENT BOND.
CH 11 TOPIC 20 MOLECULAR COMPOUNDS COVALENT BONDS 2 2) A COVALENT BOND atoms Lewis symbol H i + i H H:H a) The covalent bond created between non-metal atoms results from intramolecular forces, which are the natural attractions between the positive protons in each atom s nucleus and the two shared negative valence electrons. When two H atoms bind together, one valence electron from each atom pair up and locate themselves between each H atom s nucleus. Each atom s positive nucleus is attracted to the negative pair of electrons located between them thus bonding the two H atoms together. b) Covalent bonds are modeled three equally valid ways. MEMORIZE THEM!! i) In ELECTRON DOT DIAGRAMS (LEWIS SYMBOLS), dots around an element s symbol represent unbonded valence electrons; two dots drawn between the symbols represent a single covalent bond. ii) In LEWIS STUCTURES, dots around an element s symbol represent un-bonded valence electrons; dashes between the symbols represent the shared pairs of electrons of the covalent bond. iii) A STRUCTURAL FORMULA is a short hand version of the Lewis structure for a molecular compound using only dashes between the symbols to represent shared pairs of electrons. 3) EXAMPLE PROBLEM: Draw the electron dot diagram, Lewis structure, and structural formula for the diatomic element bromine. Solution: When making the electron dot diagram (EDD) for the diatomic molecule Br 2, the bromine atoms share two valence electrons, written as one pair of dots between the symbols. atoms EDD = Lewis structure = structural formula. a) A single covalent bond holds the bromine atoms together in the molecule. Electron dot diagrams represent a single bond with one pair of dots between the symbols, while Lewis structures and structural formulae represent a single bond with one dash between the symbols. C) Consider the diatomic element oxygen, O 2. When making the electron dot diagram (EDD) for the O 2 molecule, the oxygen atoms share four valence electrons written as two pairs of dots between the symbols. atoms EDD = Lewis structure = structural formula 1) The oxygen atoms in the molecule are held together by a double bond. Electron dot diagrams represent a double bond with two pairs of dots between the symbols, while Lewis structures and structural formulae represent a double bond with two dashes between the symbols. D) Consider the diatomic element nitrogen, N 2. When making the electron dot diagram (EDD) for the N 2 molecule, the nitrogen atoms share six valence electrons written as three pairs of dots STACKED VERTICALLY between the symbols. atoms EDD = Lewis structure = structural formula 1) The nitrogen atoms in the molecule are held together by a triple bond. Electron dot diagrams represent a triple bond with three pairs of dots STACKED VERTICALLY between the symbols, while Lewis structures and structural formulae represent a triple bond with three dashes between the symbols.
CH 11 TOPIC 20 MOLECULAR COMPOUNDS COVALENT BONDS 3 II) BONDING CAPACITY A) BONDING CAPACITY is defined as the number of electrons lost, gained or shared by an atom when it forms a bond. An atom s bonding capacity is used to predict the number of bonds it prefers to form. 1) Bonding capacity is directly related to the number of lone valence electrons around an atom. A lone electron is called a bonding electron because it can pair up with a lone electron from another atom to form a single bond. An atom with one lone electron can form one bond while an atom with two lone electrons can form two bonds. i.e. : F i fluorine has one lone valence electron thus its bonding capacity is 1. i O i oxygen has two lone valence electrons thus its bonding capacity is 2. : N i nitrogen has three lone valence electrons thus its bonding capacity is 3. 2) Table 1: Bonding capacity of some common non-metal elements Atom # of lone electrons # of bonds possible Bonding Capacity carbon & silicon 4 4 4 nitrogen 3 3 3 oxygen 2 2 2 hydrogen & the halogens 1 1 1 a) The information in the above table will help you draw electron dot diagrams, Lewis structures and structural formulae. MEMORIZE IT!! III) DRAWING ELECTRON DOT DIAGRAMS, LEWIS STRUCTURES AND STRUCTURAL FORMULAE FOR MOLECULAR COMPOUNDS. A) USE THESE STEPS TO DRAW ELECTRON DOT DIAGRAMS, LEWIS STRUCTURE AND STRUCTURAL FORMULA FOR SIMPLE MOLECULAR COMPOUNDS. (1) Identify the central atom (usually the element with the largest bonding capacity) and arrange the symbols of the other atoms around it. Atoms around the central atom are called peripheral atoms. (2) Add the number of valence electrons available from each atom. (3) Make bonds by placing a pair of electrons between the central atom and each peripheral atom. (4) Give peripheral atoms stable octets by placing remaining electrons around them ensuring each has four pairs (eight electrons) around it. (5) Ensure every atom has a stable octet by placing unused electrons as needed, or creating double or triple bonds between the central atom and one or more peripheral atom(s) as necessary by moving lone pairs of electrons from one or more peripheral atoms. (6) Draw the Lewis structures by replacing the paired dots in bonds with dashes: 1 pair of electrons = 1 bond = 1 dash = a single bond, 2 pairs of electrons = 2 bonds = 2 dashes = a double bond, 3 pairs of electrons = 3 bonds = 3 dashes = a triple bond. Eliminate dots to draw structural formulae. B) Sample Problems 1 1) Draw the electron dot diagram, Lewis structure and structural formula for H 2 O (l). (1) Identify the central atom (usually the element with the largest bonding capacity) and arrange the symbols of the other atoms around it. Atoms around the central atom are called peripheral atoms (2) Add the number of valence electrons available from each atom. (3) Make bonds by placing a pair of electrons between the central atom and each peripheral atom. Cont. on next page.
CH 11 TOPIC 20 MOLECULAR COMPOUNDS COVALENT BONDS 4 (4) Give peripheral atoms stable octets by placing remaining electrons around them ensuring each has four pairs (eight electrons) around it. If electrons remain, place them around the central atom. (5) Ensure every atom has a stable octet by placing unused electrons as needed, or creating double or triple bonds between the central atom and one or more peripheral atom(s) as necessary by moving lone pairs of electrons from one or more peripheral atoms. Not applicable here (6) Draw the Lewis structures by replacing the paired dots in bonds with dashes: 1 pair of electrons = 1 bond = 1 dash = a single bond, 2 pairs of electrons = 2 bonds = 2 dashes = a double bond, 3 pairs of electrons = 3 bonds = 3 dashes = a triple bond. Eliminate dots to draw structural formulae. electron dot diagram Lewis structure structural formula 2) Draw the electron dot diagram, Lewis structure and structural formula for CO 2(g). (1) (2) = total valence electrons to place = 16 valence electrons to place (3) 4 used, 12 valence electrons to place (4) all 16 valence electrons used (5) (6) (7) electron dot diagram Lewis structure structural formula C) Required Practice 1: Draw the electron dot diagram, Lewis structure and structural formula for each of these compounds. {Answers are on page 10 of these notes.} 1. ClF (g) 2. CBr 4(s) 3. Cl 2 O (g) 4. NH 3(g) 5. OF 2(g) 6. CCl 4(l) 7. SO 2(g) 8. O 3(g) 9. SO (g) 10. CO (g)
CH 11 TOPIC 20 MOLECULAR COMPOUNDS COVALENT BONDS 5 IV) DRAWING ELECTRON DOT DIAGRAMS, LEWIS STRUCTURES & STRUCTURAL FORMULAE FOR POLYATOMIC IONS. A) USE THESE STEPS TO DRAW ELECTRON DOT DIAGRAMS, LEWIS STRUCTURE AND STRUCTURAL FORMULA FOR POLYATOMIC IONS. (1) Identify the central atom (usually the element with the largest bonding capacity) and arrange the symbols of the other atoms around it. Atoms around the central atom are called peripheral atoms. (2) Add the number of valence electrons available from each atom. If the structure is a polyatomic anion, add the ion s charge to the number of electrons available. If the structure is a polyatomic cation subtract the ion s charge from the number of electrons available. (3) Make bonds by placing a pair of electrons between the central atom and each peripheral atom. (4) Give peripheral atoms stable octets by placing remaining electrons around them ensuring each has four pairs (eight electrons) around it. (5) Ensure every atom has a stable octet by placing unused electrons as needed, or creating double or triple bonds between the central atom and one or more peripheral atom(s) as necessary by moving lone pairs of electrons from one or more peripheral atoms. (6) If the structure is a polyatomic ion, place square brackets around it and the appropriate charge outside the upper right corner of the bracket. (7) Draw the Lewis structures by replacing the paired dots in bonds with dashes: 1 pair of electrons = 1 bond = 1 dash = a single bond, 2 pairs of electrons = 2 bonds = 2 dashes = a double bond, 3 pairs of electrons = 3 bonds = 3 dashes = a triple bond. Eliminate dots to draw structural formulae. B) SAMPLE PROBLEM 2 1) Draw the electron dot diagram, Lewis structure and structural formula for CO 3 2 (aq). (1) (2) = total valence electrons to place = 24 valence electrons to place (3) 6 used, 18 valence electrons to place (4) all 24 valence electrons used (5) (6) (7) electron dot diagram Lewis structure structural formula C) Required Practice 2: Draw the electron dot diagram, Lewis structure and structural formula for each of these compounds. {Answers are on page 11 of these notes.} 2 3 + + 1. ClO 3 (aq) 2. SO 4 (aq) 3. PO 4 (aq) 4. NH 4 (aq) 5. NF 4 (aq)
CH 11 TOPIC 20 MOLECULAR COMPOUNDS COVALENT BONDS 6 A) HYDROCARBONS. 1) Carbon has a bonding capacity of 4, thus it will be the central atom. When a hydrocarbon or a carbohydrate contains more than one carbon atom, the carbon atoms form a central backbone chain of bonded carbon atoms to which the peripheral hydrogen. Hydrogen has a bonding capacity of 1 thus it can only form 1 bond. As a result hydrogen is always at the end of a bond, never in between two bonds, thus it is said to be terminal. 2) USE THESE STEPS TO DRAW ELECTRON DOT DIAGRAMS, LEWIS SYMBOLS AND STRUCTURAL FORMULA FOR HYDROCARBONS. (1) Arrange the symbols of the peripheral atoms around the central carbon atom. If there is more than one carbon atom, place them in a chain with equal numbers of hydrogen atoms around each carbon. (2) Add the number of valence electrons available from each atom. (3) Make bonds by placing one pair of electrons between the atoms in the central chain and between the peripheral atoms and the central atom(s). (4) Make stable octets by placing remaining electrons around any non-hydrogen peripheral atoms then around the central atom or central chain giving each a stable octet. (5) Place remaining electrons in single pairs between the carbon atoms in the backbone chain creating double or triple bonds as necessary until every atom has a stable octet. (6) Draw the Lewis structures by replacing the paired dots in bonds with dashes: 1 pair of electrons = 1 bond = 1 dash = a single bond, 2 pairs of electrons = 2 bonds = 2 dashes = a double bond, 3 pairs of electrons = 3 bonds = 3 dashes = a triple bond. Eliminate dots to draw structural formulae. 3) SAMPLE PROBLEMS 3 a) Draw the electron dot diagram, Lewis structure and structural formula for C 2 H 6(g). (1) Arrange the symbols of the peripheral atoms around the central carbon atom. If there is more than one carbon atom, place them in a chain with equal numbers of hydrogen atoms around each carbon. (2) Add the number of valence electrons available from each atom. (3) Make bonds by placing one pair of electrons between the atoms in the central chain and between the peripheral atoms and the central atom(s). (4) Make stable octets by placing remaining electrons around any non-hydrogen peripheral atoms then around the central atom or central chain giving each a stable octet. (5) Place remaining electrons in single pairs between the carbon atoms in the backbone chain creating double or triple bonds as necessary until every atom has a stable octet. Continued on the next page.
CH 11 TOPIC 20 MOLECULAR COMPOUNDS COVALENT BONDS 7 (6) Draw the Lewis structures by replacing the paired dots in bonds with dashes: 1 pair of electrons = 1 bond = 1 dash = a single bond, 2 pairs of electrons = 2 bonds = 2 dashes = a double bond, 3 pairs of electrons = 3 bonds = 3 dashes = a triple bond. Eliminate dots to draw structural formulae. electron dot diagram Lewis structure structural formula b) Draw the electron dot diagram, Lewis structure and structural formula for C 2 H 4(g). (1) (2) = total electrons to place = 12 electrons to place (3) 10 electrons used, 2 electrons to place (4) all 12 electrons used (5) electron dot diagram (6) Lewis structure structural formula
CH 11 TOPIC 20 MOLECULAR COMPOUNDS COVALENT BONDS 8 4) Required Practice 3: Draw electron dot diagrams, Lewis structures and structural formulae for each of these compounds. {Answers are on page 11 & 12 of these notes.} 1. C 3 H 8(l) 2. C 4 H 10(l) 3. C 5 H 12(g) 4. C 2 Cl 6(l) 5. C 2 Cl 4(l) 6. C 2 Cl 2(g) B) Carbohydrates. 1) Carbon has a bonding capacity of 4, thus it will be the central atom. When a hydrocarbon or a carbohydrate contains more than one carbon atom, the carbon atoms form a central backbone chain of bonded carbon atoms to which the peripheral hydrogen and oxygen atoms are bonded. Hydrogen has a bonding capacity of 1 thus it can only form 1 bond. As a result hydrogen is always at the end of a bond, never in between two bonds, thus it is said to be terminal. Oxygen has a bonding capacity of 2 and thus will form 2 bonds. As a result, it is bonded to one carbon atom and one hydrogen atom or it is double bonded to one carbon atom. 2) The molecular formula of a carbohydrate tells you where to place the oxygen atom. In the molecular formula CH 3 OH, the oxygen atom has a hydrogen atom after it. This means it is located between a carbon atom and a hydrogen atom. This arrangement is necessary to because hydrogen can only form 1 bond and thus must be terminal. 3) USE THESE STEPS TO DRAW ELECTRON DOT DIAGRAMS, LEWIS SYMBOLS AND STRUCTURAL FORMULA FOR CARBOHYDRATES. (1) Arrange the symbols of the peripheral atoms around the central carbon atom. If there is more than one carbon atom, place them in a chain with equal numbers of hydrogen atoms around each carbon. (2) Add the number of valence electrons available from each atom. (3) Make bonds by placing one pair of electrons between the atoms in the central chain and between the peripheral atoms and the central atom(s). (4) Make stable octets by placing remaining electrons around any non-hydrogen peripheral atoms then around the central atom or central chain giving each a stable octet. (5) Place remaining electrons in single pairs between the carbon atoms in the backbone chain creating double or triple bonds as necessary until every atom has a stable octet. (6) Draw the Lewis structures by replacing the paired dots in bonds with dashes: 1 pair of electrons = 1 bond = 1 dash = a single bond, 2 pairs of electrons = 2 bonds = 2 dashes = a double bond, 3 pairs of electrons = 3 bonds = 3 dashes = a triple bond. Eliminate dots to draw structural formulae. 4) SAMPLE PROBLEM: Draw the electron dot diagram, Lewis structure and structural formula for CH 3 OH (l). The formula of this compound gives a hint to arrange the atoms. (1) (2) = total electrons to place = 14 electrons to place (3) 10 electrons used, 4 electrons to place Continued on the next page.
CH 11 TOPIC 20 MOLECULAR COMPOUNDS COVALENT BONDS 9 (4) all 14 electrons used (5) electron dot diagram (6) Lewis structure structural formula 5) Required Practice 4: Draw electron dot diagrams, Lewis structures and structural formulae for each of these compounds. {Answers are on page 12 of these notes.} 1. C 2 H 5 OH (l) 2. C 3 H 7 OH (l) 3. HOCH 2 OH (g) 4. HOC 2 H 4 OH (g) 5. C 2 H 3 OH (l) 6. H 2 O 2(l) COVALENT BOND STRENGTH I) BOND STRENGTH A) When a molecule or compound is heated to high temperatures it will remain together, decompose into the smaller molecules, or the ions or atoms it is made of. When a molecule or compound is stable at high temperatures its atoms remain bonded together and the molecule or compound remains intact when it is heated to a high temperature. When a molecule or compound is unstable at high temperatures the bonds between its atoms will break and it will decompose (break apart) into separate individual atoms or ions when it is heated to a high temperature. 1) At temperatures of 801 C or higher, NaCl (l) is unstable and will decompose to form individual Na + (l) and Cl (l) ions. At temperatures of 32 400 C or higher, N 2(g) is unstable and will decompose to form individual N (g) atoms. 3) The stability of a molecule or compound at high temperature is an indication of the strength of the bonds that hold it together. The more stable the molecule or compound, the stronger the bonds that hold its atoms together, which means more energy is required to break them. B) A common property of molecular compounds is their stability at relatively high temperatures at relatively high temperatures molecular compounds do not decompose easily. Research indicates that more energy is required to break a triple bond than a double bond. Similarly, more energy is required to break a double bond than a single bond. This means that triple bonds are stronger than double bonds, which are stronger than single bonds. II) EXPLAINING COVALENT BOND STRENGTH A) A covalent bond results from the attraction that two positive nuclei have for shared pairs of electrons. The more electrons shared by the two atoms, the greater the attraction each nucleus has for the shared electrons within the bond, thus the stronger the bond. A single bond consists of two electrons, a double bond consists of four electrons and a triple bond consists of six electrons. This makes the double bond stronger than a single bond and a triple bond stronger than both the double and single bonds.
CH 11 TOPIC 20 MOLECULAR COMPOUNDS COVALENT BONDS 10 Required Practice 1 from page 4 ANSWERS TO THE REQUIRED PRACTICE 1. : Cl : F : : Br : 2. : Br : C : Br : 3. : Cl : O : Cl : 4. H : N : Br : H : H : Br : : Cl F : : Br C Br : : Cl O Cl : H N H : Br : H Br Cl F Br C Br Cl O Cl H N H Br H 5. : F : O : F : 6. : Cl : Cl : : C : Cl : Cl : : Cl : : 7. : O : S :: O 8. : O : O :: O : F O F : : Cl C Cl : : O S = O : O O = O : Cl : Cl F O F Cl C Cl O S = O O O = O Cl 9. S :: O S = O S = O 10. : C O : : C O : C O
CH 11 TOPIC 20 MOLECULAR COMPOUNDS COVALENT BONDS 11 Required Practice 2 from page 5 1. : O : Cl : O : O : 1 2. : O : S : O : 2 : O : 3 H 1+ : O : 3. : O : P : O : 4. H : N : : O : : O : H : H : O : 2 : O : 3 H 1+ : O Cl O : 1 : O S O : : O P O : H N H : O : : O : : O : H O 2 O 3 H 1+ O Cl O 1 O S O O P O H N H O O O H : F : 1+ 5. : F : N : F : : F : : F : 1+ : F N F : : F : F 1+ F N F F Required Practice 3 from page 8 H H H H H H H H H H H H 1. H : C : C : C : H 2. H : C : C : C : C : H 3. H : C : C : C : C : C : H H H H H H H H H H H H H H H H H H H H H H H H H H C C C H H C C C C H H C C C C C H H H H H H H H H H H H H H H H H H H H H H H H H H C C C H H C C C C H H C C C C C H H H H H H H H H H H H H
CH 11 TOPIC 20 MOLECULAR COMPOUNDS COVALENT BONDS 12 : Cl : : Cl : : Cl : : Cl : 4. : Cl : C : C : Cl : 5. C :: C : Cl : : Cl : : Cl : : Cl : 6. : Cl : C C : Cl : : Cl :: Cl : : Cl :: Cl : : Cl C C Cl : C = C : Cl : Cl :: Cl : : Cl :: Cl : Cl Cl Cl Cl C C Cl : Cl C C Cl C = C Cl C C Cl Cl Cl Cl Cl Required Practice 4 from page 9 H H H H H H 1. H : C : C : O : H 2. H : C : C : C : O : H 3. H : O : C : O : H H H H H H H H H H H H H H C C O H H C C C O H H O C O H H H H H H H H H H H H H H C C O H H C C C O H H O C O H H H H H H H H H H H 4. H : O : C : C : O : H 5. C :: C : O : H 6. H : O : O : H H H H H H H H H O C C O H C = C O H H O O H H H H H H H H H O C C O H C = C O H H O O H H H H