Presentation on theme: "Unit 2: Chemical Bonding Chemistry2202. Outline Bohr diagrams & Lewis Diagrams Types of Bonding 1. Ionic 2. Covalent (molecular) 3. Metallic 4. Network."— Presentation transcript:
Unit 2: Chemical Bonding Chemistry2202
Outline Bohr diagrams & Lewis Diagrams Types of Bonding 1. Ionic 2. Covalent (molecular) 3. Metallic 4. Network covalent bonding VSEPR Theory (Shapes) Physical Properties 5. London Dispersion forces 6. Dipole-Dipole forces 7. Hydrogen Bonding
Bohr Diagrams (Review) How do we draw a Bohr Diagram for - The F atom? - The F ion? Draw Bohr diagrams for the atom and the ion for the following: AlSC lBe
Lewis Diagrams Lewis Diagrams provide a method for keeping track of electrons in atoms, ions, or molecules. AKA: Electron Dot diagrams the nucleus (p + & n o ), and filled energy levels are represented by the symbol dots are placed around the element symbol to represent valence electrons
Lewis Diagrams eg. Lewis Diagram for F F lone pair bonding electron
Lewis Diagrams lone pair – a pair of electrons not available for bonding bonding electron – a single electron that may be shared with another atom
Lewis Diagrams eg. Draw Lewis Diagrams for: carbon C P Na phosphorus sodium
Lewis Diagrams For each atom draw the Lewis diagram and state the number of lone pairs and number of bonding electrons LiBeAlSi MgNBO
Lewis Diagrams for Compounds To draw the LD for a molecule: draw the LD for each atom in the molecule the atom with the most bonding electrons is the central atom connect the other atoms using single bonds (1 pair of shared electrons) some molecules may have double bonds or triple bonds
Lewis Diagrams for Compounds eg. Draw the LD for: PH 3 CF 4 Cl 2 O C2H6C2H6 C2H4C2H4 C2H2C2H2
Lewis Diagrams for Compounds eg. Draw the LD for: NH 3 SiCl 4 N 2 H 4 HCN SI 2 CO 2 N 2 H 2 CH 2 O POICH 3 OH N 2 H 2 O 2
Lewis Diagrams for Compounds A structural formula shows how the atoms are connected in a molecule. To draw a structural formula: replace the bonded pairs of electrons with short lines omit the lone pairs of electrons
Why is propane (C 3 H 8 ) a gas at STP while kerosene (C 10 H 22 ) a liquid?
Why is graphite soft enough to write with while diamond is the hardest substance? (both are C) Graphite Diamond
Other forms of carbon Nanotube CarbonBuckyball Carbon
How can you which is ‘real gold’ and which is ‘fool’s gold’ (pyrite) by hitting it with a rock?
‘As Slow As Cold Molasses’ ‘All Because of Bonding’
Viscosity of liquids
Malleability, Ductility and Conductivity A single gram of gold can be stretched into a wire 3.2km long. A gram of gold can be flattened into a sheet with an area of 6.7 sq ft. Silver has the highest electrical conductivity of any element and the highest thermal conductivity of any metal.
Bonding Bonding between atoms, ions and molecules determines the physical and chemical properties of substances. Bonding can be divided into two categories: - Intramolecular forces - Intermolecular forces
Bonding Intramolecular forces are forces of attraction between atoms or ions. Intramolecular forces include: 1. ionic bonding 2. covalent bonding 3. metallic bonding 4. network covalent bonding
Bonding Intermolecular forces occur BETWEEN different molecules. Intermolecular forces include: 5. London Dispersion Forces 6. Dipole-Dipole forces 7. Hydrogen Bonding
Ionic and Covalent Bonding ThoughtLab p. 161 Identify #’s 1 - 6
Ionic Bonding Occurs between cations and anions – usually metals and non-metals. An ionic bond is a force of attraction between positive and negative ions. Properties: conduct electricity as liquids and in solution hard crystalline solids high melting points and boiling points brittle Mar. 11
In an ionic crystal the ions pack tightly together. (p. 167) The repeating 3-D distribution of cations and anions is called an ionic crystal lattice. Ionic Bonding Mar. 11
Each anion is attracted to six or more cations at once. The same is true for the individual cations. Ionic Bonding Mar. 11
Ionic Bonding Mar. 11
Covalent Bonding Occurs between non-metals in molecular compounds. Atoms share bonding electrons to become more stable (noble gas structure). A covalent bond is a simultaneous attraction by two atoms for a common pair of valence electrons. Mar. 11
Covalent Bonding Molecular compounds have low melting and boiling points. exist as distinct molecules. Mar. 11
Covalent Bonding do not conduct electric current in any form Mar. 11
PropertyIonicMolecular Type of elements Metals and nonmetals Non-Metals Force of Attraction Positive ions attract negative ions Atoms attract a shared electron pair Electron movement Electrons move from the metal to the nonmetal Electrons are shared between atoms State at room temperature Always solidsSolids, liquids, or gas Mar. 11
PropertyIonicMolecular SolubilitySoluble or low solubility Soluble or insoluble Conductivity in solid state None Conductivity in liquid state ConductsNone Conductivity in solution ConductsNone Mar. 11
Na Metallic Bonding (p. 171) Na Na Na Na Na Na Na
Na + Metallic Bonding (p. 171) Na + Na + Na + Na + Na + Na + Na +
Metallic Bonding (p. 171) metals tend to lose valence electrons. valence electrons are loosely held and frequently lost from metal atoms. this produces in positive metal ions surrounded by freely moving valence electrons. metallic bonding is the force of attraction between the positive metal ions and the mobile or delocalised valence electrons
This theory of metallic bonding is called the ‘Sea of Electrons’ Model or ‘Free Electron’ Model
Metallic Bonding Metallic bonding theory accounts for properties of metals 1. electrical conductivity - electric current is the flow of electrons - metals are the only solids in which electrons are free to move 2. solids - attractive forces between positive cations and negative electrons are very strong
Metallic Bonding 3. malleability and ductility - metals can be hammered into thin sheets(malleable) or drawn into thin wires(ductile). - metallic bonding is non-directional such that layers of metal atoms slide past each other under pressure.
Network Covalent Bonding (p. 199) occurs in 3 compounds (memorize these) diamond – C n carborundum – SiC quartz – SiO 2 produces large molecules with covalent bonding in 3d each atom is held in place in 3d by a network of other atoms
Network Covalent bonding Properties: the highest melting and boiling points the hardest substances brittle do not conduct electric current in any form
Valence Shell Electron Pair Repulsion or VSEPR theory The shape of molecules is determined by the arrangement of valence electron pairs around the atoms in a compound. The shapes are the result of REPULSION between pairs of valence electrons. Valence electron pairs move as far away from each other as possible.
There are 5 shapes that can be determined by the # of bonds and # of lone pairs on the central atom. Valence Shell Electron Pair Repulsion or VSEPR theory
1.Tetrahedral (4 bonds; 0 lone pairs)
2. Pyramidal (3 bonds; 1 lone pair)
3. V-shaped (2 bonds; 2 lone pairs)
4. Trigonal Planar (3 bonds; 0 lone pairs)
5. Linear (2 bonds; 0 lone pairs)
For each molecule below draw the Lewis diagram and the shape diagram. 1 central atom HOClH 2 SeH 2 SiO NBr 3 CHCl 3 SiH 4 PBr 3 HCN I 2 2 central atoms C 2 F 4 C 2 H 6 CH 3 OH C 2 H 2 H 2 O 2
Electronegativity (EN - p. 174) EN is a measure of the attraction that an atom has for shared electrons. A higher EN means a stronger attraction or electrostatic pull on valence electrons EN values increase as you move: - from left to right in a period - up in a group or family
Electronegativity & Covalent Bonds 1. polar covalent bond - a bond between atoms with different EN - the shared electron pair is attracted more strongly to the atom with the higher EN ClH δ− δ+
2. nonpolar covalent bond - occurs between atoms with same EN - the shared electron pair is attracted more strongly to the atom with the higher EN - the separation of charge or bond dipole is shown using an arrow pointing toward the more electronegative atom. - the Greek letter delta (δ) indicates ‘partial’ charges Complete: #’s 7 – 9 on p.178 Electronegativity & Covalent Bonds
Electronegativity and Ionic Bonds Because the EN of metals is so low, metals lose electrons to form cations Nonmetals gain electrons to form anions because their EN is relatively high When ions form, the resulting electrostatic force is an ionic bond
Electronegativity and Covalent Bonds Atoms in covalent compounds can either have the same EN eg. Cl 2, PH 3, NCl 3 OR different EN eg. HCl
Electronegativity and Covalent Bonds Atoms with the same EN have the same attraction for shared valence electrons. Covalent bonds resulting from equal sharing of the bonding electron pairs are called Nonpolar Covalent Bonds Atoms with different EN attract the shared valence electron pair at different strengths. (higher EN has a stronger attraction for the shared electron pair)
Electronegativity and Covalent Bonds eg. HCl Cl has a higher EN the bonding electron pair is pulled closer to the chlorine atom this produces slight positive and negative charges within the bond these charges are referred to as “partial charges” and are denoted with the Greek letter delta (δ).
Electronegativity and Covalent Bonds The region around the chlorine atom will be slightly negative The region around the hydrogen will be slightly positive.
Electronegativity and Covalent Bonds Because the bond is polarized into a positive area and a negative area the bond has a “bond dipole”. an arrow points to the atom with the higher EN. Covalent bonds resulting from unequal sharing of bonded electron pairs are Polar Covalent Bonds.
Electronegativity and Covalent Bonds eg. H 2 O
eg. HF Electronegativity and Covalent Bonds
Electronegativity Homework p. 178 #’s 7, 8, & 9 p. 180 #’s 1, 2, & 3
Bond Energy (pp ) 1. Describe the forces of attraction and repulsion present in all bonds. 2. What is bond length? 3. Define bond energy. 4. Which type of bond has the most energy? 5. How can bond energy be used to predict whether a reaction is endothermic or exothermic?
Test Outline Bohr Diagrams (atoms & ions) Lewis Diagrams (Electron Dot) Ion Formation Ionic Bonding, Structures & Properties Covalent Bonding, Structures & Properties
Test Outline Metallic Bonding Theory& Properties Network Covalent Bonding & Properties Electronegativity Bond Dipoles & Polar Molecules VSEPR Theory LD, DD, & H-bonding Predicting properties (bp, mp, etc.)
Molecular Dipoles The vector sum of all the bond dipoles in a molecule is a Molecular Dipole A Polar Molecule has a molecular dipole that points toward the more electronegative end of the molecule. eg. H 2 O
Molecular Dipoles Nonpolar molecules DO NOT have molecular dipoles. This occurs when: - bond dipoles cancel - there are no bond dipoles To determine whether a molecule is polar: - draw the LD and the shape diagram - draw the bond dipoles and determine whether they cancel eg. CO 2 eg. PH 3
Molecular Dipoles See Handout #1
Intermolecular Forces Mar. 31
Strongest bonds; Highest mp and bp 1. Network Covalent (C n SiO 2 SiC) 2. Ionic bonding(metal & nonmetal) 3. Metallic bonding (metals) 4. Molecular (nonmetals) Weakest bonds; Lowest mp and bp - Intermolecular forces present Mar. 31
To compare mp and bp in covalent compounds you must use: - London Dispersion forces (p. 204) (all molecules) - Dipole-Dipole forces (pp. 202, 203) (polar molecules) - Hydrogen Bonding (pp. 205, 206) (H bonded to N, O, or F) Mar. 31
Intermolecular Forces (p. 202) Mar. 31
Intermolecular Forces Covalent compounds have low mp and bp because forces between molecules in covalent compounds are very weak. Intermolecular forces were studied by the Dutch physicist Johannes van der Waals In his honor, two types of intermolecular force are called Van der Waals forces. Intermolecular forces can be used to explain physical properties of covalent compounds. Mar. 31
1. London Dispersion Forces LD forces exist in ALL molecular elements & compounds. The positive charges in one molecule attract the negative charges in a second molecule. The temporary dipoles caused by electron movement in one molecule attract the temporary dipoles of another molecule. Apr. 1
1. London Dispersion Forces The strength of these forces depends on: a)the number of electrons more electrons produce stronger LD forces that result in higher mp and bp eg.CH 4 is a gas at room temperature. C 8 H 18 is a liquid at room temperature. C 25 H 52 is a solid at room temperature. Account for the difference. Apr. 1
1. London Dispersion Forces Two molecules that have the same number of electrons are isoelectronic eg. C 2 H 6 and CH 3 F Apr. 1
1. London Dispersion Forces b)shape of the molecule molecules that “fit together” better will experience stronger LD forces eg. Cl 2 vaporizes at -35 ºC while C 4 H 10 vaporizes at -1 ºC. Use bonding to account for the difference. Apr. 1
2. Dipole-dipole Forces -occur between polar molecules - the δ+ end of one polar molecule is attracted to the δ- end of another polar molecule (& vice-versa) eg. Which has the higher boiling point; CH 3 F or C 2 H 6 ? Apr. 5
p. 202 Apr. 5 In the liquid state, polar molecules are oriented such that oppositely charged ends of the molecules are close to each other.
3. Hydrogen Bonds -a special type of dipole-dipole force (about 10 times stronger) - only occurs BETWEEN MOLECULES that contain H directly bonded to F, O, or N ie. the molecule contains at least one H-F, H-O, or H-N covalent bond. Apr. 5
3. Hydrogen Bonds - the hydrogen bond occurs between the H atom of one molecule and the N, O, or F of a second molecule. eg. Arrange these from highest to lowest boiling point; C 3 H 8 C 2 H 5 OHC 2 H 5 F Apr. 5
p. 206 Apr. 5
NOTE: To compare mp and bp in covalent compounds you must use: - London Dispersion forces (all molecules) - Dipole-Dipole forces (polar molecules) - Hydrogen Bonding (H bonded to N, O, or F)
WorkSheet: Bonding #4 p. 225 #’s 9 & 10 p. 226 #’s 12 – 14,
Intermolecular Forces 1. Use intermolecular forces to explain the following: a) Ar boils at -186 °C and F 2 boils at -188 °C. b) Kr boils at -152 °C and HBr boils at -67 °C. c) Cl 2 boils at -35 °C and C 2 H 5 Cl boils at 13 °C. 2. Examine the graph on p. 210: a) Account for the increase in boiling point for the hydrogen compounds of the Group IV elements. b) Why is the trend different for the hydrogen compounds of the Group V, VI, and VII elements? c) Why are the boiling points of the Group IVA compounds consistently lower than the others.
3.Which substance in each pair has the higher boiling point. Justify your answers. (a)SiC or KCl (b)RbBr or C 6 H 12 O 6 (c)C 3 H 8 or C 2 H 5 OH (d)C 4 H 10 or C 2 H 5 Cl
Summary Strongest - Network Covalent - Ionic - Metallic Weakest - Covalent ↣ LD forces (all molecules) ↣ DD forces (polar molecules) ↣ H-Bonding (H bonded to N, O, or F)
Ion-Dipole Forces An ion-dipole force is the force of attraction between an ion and a polar molecule (a dipole).
Ion-Dipole Forces NaCl dissolves in water because the attractions between the Na + and Cl - ions and the partial charges on the H 2 O molecules are strong enough to overcome the forces that bind the ions together.
Dispersion (London) Forces Bond vibrations, which are part of the normal condition of a non-polar molecule, cause momentary, uneven distribution of charge; a non-polar becomes slightly polar for an instant, and continues to do so in a random but constant basis.
Dispersion (London) Forces At the instant that one non-polar molecule is in a slightly polar condition, it is capable of inducing a dipole in a nearby molecule This force of attraction is called a dispersion force.
Dispersion (London) Forces Two factors affect the magnitude of dispersion forces 1. # of electrons in the molecule - Vibrations within larger molecules with more electrons than smaller molecules can easily cause an uneven distribution of charge. - dispersion forces between these larger molecules are thus stronger, which raises the boiling point for larger molecules.
Dispersion (London) Forces 2. The shape of the molecule: - molecules with spherical shape have a smaller surface area than a straight chain molecule with the same number of electrons - substances with spherical moleculular shape will have weaker dispersion forces and a lower boiling point. - London dispersion forces are responsible for the formation and stabilization of the biological membranes surrounding every living cell.
Hydrogen Bonding - In order to form a hydrogen bond, a hydrogen atom must be bonded to a highly electronegative atom such as oxygen, nitrogen, or fluorine. - These bonds are very polar, and since hydrogen has no other electrons, the positive proton, H +, is exposed and can become strongly attracted to the negative end of a nearby dipole. - A hydrogen bond is an electrostatic attraction between the nucleus of a hydrogen atom, bonded to fluorine, oxygen, or nitrogen and the negative end of a dipole nearby.
Hydrogen Bonding … H H O H O H δ+ δ+ δ+ δ+ δ−δ− H
In biological systems, these polar bonds are often parts of much larger molecules (ie. N H bonds and C O bonds found in biological molecules)
Hydrogen Bonding in Water Hydrogen bonds between the hydrogen atoms in one water molecule and the oxygen atom in another account for many unique properties of water. H O H H O H δ+ δ+ δ+ δ+ δ−δ− … H
Hydrogen Bonding in Water In liquid water, each water molecule is hydrogen bonded to at least four other water molecules. The large number of bonds between water molecules makes the net attractive force quite strong
Hydrogen Bonding in Water the strong attractive forces are responsible for the relatively high boiling point of water. The water molecules are farther apart in ice then they are in liquid water making ice less dense than liquid water.
Hydrogen Bonding in Water Hydrogen bonds force water molecules into the special hexagonal, crystalline structure of ice when the temperature is below 4 degrees celcius.