Unit 2: Chemical Bonding Chemistry2202

Outline Bohr diagrams Lewis Diagrams Types of Bonding Ionic bonding Covalent bonding (Molecular) Metallic bonding Network covalent bonding

Types of Bonding (cont’d) London Dispersion forces Dipole-Dipole forces Hydrogen Bonding VSEPR Theory (Shapes) Physical Properties

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: Al S C l Be

Lewis Diagrams LD provide a method for keeping track of electrons in atoms, ions, or molecules Also called Electron Dot diagrams the nucleus (p+& n0) and filled energy levels are represented by the element symbol

Lewis Diagrams dots are placed around the element symbol to represent valence electrons

F • • • • • • • Lewis Diagrams lone pair bonding electron lone pair eg. Lewis Diagram for F lone pair • • bonding electron • F lone pair • • • • lone pair

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. Lewis Diagram for C • • C • •

Lewis Diagrams eg. Lewis Diagram for P • • • P • •

Lewis Diagrams eg. Lewis Diagram for Na • Na

Li Be Al Si Mg N B O Lewis Diagrams For each atom draw the Lewis diagram and state the number of lone pairs and number of bonding electrons Li Be Al Si Mg N B O

Lewis Diagrams for Compounds draw the LD for each atom in the compound The atom with the most bonding electrons is the central atom Connect the other atoms using single bonds (1 pair of shared electrons) In some cases there may be double bonds or triple bonds

Lewis Diagrams for Compounds eg. Draw the LD for: PH3 CF4 Cl2O C2H6 C2H4 C2H2

Lewis Diagrams for Compounds eg. Draw the LD for: NH3 SiCl4 N2H4 HCN SI2 CO2 N2H2 CH2O POI CH3OH N2 H2 O2

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 (C3H8) a gas at STP while kerosene (C10H22) a liquid?

Why is graphite soft enough to write with while diamond is the hardest substance known even though both substances are made of pure carbon?

Why can you tell if it is ‘real gold’ or just ‘fool’s gold’ (pyrite) by hitting it with a rock?

‘As Slow As Cold Molasses’ ‘All Because of Bonding’

‘liquids’ @ -30 ºC

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: ionic bonding covalent bonding metallic bonding network covalent bonding

Bonding Intermolecular forces are forces of attraction between molecules. Intermolecular forces include: London Dispersion Forces Dipole-Dipole forces 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 the 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

Ionic Bonding In an ionic crystal the ions pack tightly together. The repeating 3-D distribution of cations and anions is called an ionic crystal lattice.

Ionic Bonding Each anion can be attracted to six or more cations at once. The same is true for the individual cations.

Ionic Bonding

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.

Covalent Bonding Molecular compounds have low melting and boiling points. Exist as distinct molecules.

Covalent Bonding Molecular compounds do not conduct electric current in any form

State at room temperature Property Ionic Molecular Type of elements Force of Attraction Electron movement State at room temperature Metals and nonmetals Non-Metals Positive ions attract negative ions Atoms attract a shared electron pair Electrons move from the metal to the nonmetal Electrons are shared between atoms Solids, liquids, or gas Always solids

Property Ionic Molecular Solubility Conductivity in solid state Conductivity in liquid state Conductivity in solution Soluble or low solubility Soluble or insoluble None None None Conducts Conducts None

Metallic Bonding (p. 171) metals tend to lose valence electrons. valence electrons are loosely held and frequently lost from metal atoms. This results in 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

Metallic Bonding

Metallic Bonding This theory of metallic bonding is called the ‘Sea of Electrons’ Model or ‘Free Electron’ Model

Metallic Bonding This theory accounts for properties of metals electrical conductivity - electric current is the flow of electrons - metals are the only solids in which electrons are free to move solids Attractive forces between positive cations and negative electrons are very strong

Metallic Bonding 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 – Cn carborundum – SiC quartz – SiO2 large molecules with covalent bonding in 3-d each atom is held in place in 3-d 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

1. Network Covalent (Cn ,SiO2 , SiC) MP & BP decreases Strongest 1. Network Covalent (Cn ,SiO2 , SiC) 2. Ionic bonding(metal & nonmetal) 3. Metallic bonding (metals) 4. Molecular (nonmetals) Weakest

Valence Shell Electron Pair Repulsion theory (VSEPR) The shape of molecules is caused by repulsion between valence electron pairs around the atoms in a compound. Repulsion between valence electron pairs force them to move as far away from each other as possible.

Valence Shell Electron Pair Repulsion theory (VSEPR) To determine molecular shapes, count the # of bonds and # of lone pairs on the central atom(s). We will examine 5 molecular shapes

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. HOCl H2Se H2O2 NBr3 C2F4 C2H6 CHCl3 CH3OH PBr3 I2 SiH4 HCN SiH2O C2H2

Electronegativity & Covalent Bonds Electronegativity - EN - p. 174 EN measures 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

Increases

Electronegativity & Covalent Bonds polar covalent bond a bond between atoms with different EN the bonding pair is attracted more strongly to the atom with the higher EN δ− δ+ Cl H bond dipole

Electronegativity & Covalent Bonds nonpolar covalent bond a bond between atoms with the same EN the bonding pair is shared equally between the atoms Complete: #’s 7 – 9 on p.178

Electronegativity & Covalent Bonds polar molecule - a molecule in which the bond dipoles do not cancel each other - a polar molecule has a molecular dipole that points toward the more electronegative end of the molecule. eg. HCN

Electronegativity & Covalent Bonds nonpolar molecule - a molecule in which the bond dipoles cancel each other OR - there are no bond dipoles eg. CO2 PH3

Electronegativity & Covalent Bonds To determine whether a molecule is polar: - draw the Lewis diagram & shape diagram - draw the bond dipoles & determine whether they cancel

Intermolecular Forces

1. Network Covalent (Cn ,SiO2 , SiC) MP & BP decreases Strongest 1. Network Covalent (Cn ,SiO2 , SiC) 2. Ionic bonding(metal & nonmetal) 3. Metallic bonding (metals) 4. Molecular (nonmetals) Weakest

To compare mp and bp in covalent compounds you must use: - London Dispersion forces (p. 204) (in all molecules) - Dipole-Dipole forces (pp. 202, 203) (in polar molecules) - Hydrogen Bonding (pp. 205, 206) (when H is bonded to N, O, or F)

Intermolecular Forces (p. 202)

Intermolecular Forces Covalent compounds have low mp and bp because attractive forces between molecules are very weak. Intermolecular forces were studied extensively by the Dutch physicist Johannes van der Waals In his honor, two types of intermolecular force are called Van der Waals forces.

Intermolecular Forces Intermolecular forces can be used to account for the physical properties of covalent compounds.

Intermolecular Forces

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.

1. London Dispersion Forces The strength of these forces depends on: the number of electrons more electrons produce stronger LD forces that result in higher mp and bp eg. CH4 is a gas at room temperature. C8H18 is a liquid at room temperature. C25H52 is a solid at room temperature. Account for the difference.

1. London Dispersion Forces Two molecules that have the same number of electrons are isoelectronic eg. C2H6 and CH3F

1. London Dispersion Forces b) shape of the molecule molecules that “fit together” better will experience stronger LD forces eg. Cl2 vaporizes at -35 ºC while C4H10 vaporizes at -1 ºC. Use bonding to account for the difference.

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 CH3F or C2H6 ?

p. 202

3. Hydrogen Bonds - a special type of dipole-dipole force (about 10 times stronger) - only occurs between molecules that contain a H atom which is directly bonded to F, O, or N ie. the molecule contains at least one H-F, H-O, or H-N covalent bond.

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 C3H8 C2H5OH C2H5F

p. 206

NOTE: To compare covalent compounds you must use: - London Dispersion forces (all molecules) - Dipole-Dipole forces (polar molecules) - Hydrogen Bonding (H bonded to N, O, or F)

Alchem worksheet pp. G32, 33

p. 226 #13 Omit parts g), j) – o), q), u), & v) - Answers on p. 815 for #13 - Incorrect answers c), d), & s)

p. 210

Intermolecular Forces 1. Use intermolecular forces to explain the following: a) Ar boils at -186 °C and F2 boils at -188 °C . b) Kr boils at -152 °C and HBr boils at -67 °C. c) Cl2 boils at -35 °C and C2H5Cl 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 3. Which substance in each pair has the higher boiling point. Justify your answers. (a) SiC or KCl (b) RbBr or C6H12O6 (c) C3H8 or C2H5OH (d) C4H10 or C2H5Cl

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 other compounds.

p. 226 #’s 13 & 14

Electronegativity Electronegativity is a result of the space between the nucleus and the electrons As the number of protons in the nucleus increases, the attractive force on the electrons increases, pulling them closer to the nucleus

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 the EN of nonmetals is relatively high

Electronegativity and Ionic Bonds 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. Cl2 , PH3, NCl3 different EN eg. HCl

Electronegativity and Covalent Bonds Atoms that have different EN attract the shared pair of valence electrons at different strengths The atom with the higher EN exerts a stronger attraction on the shared electron pair eg. H2O

Electronegativity and Covalent Bonds Since the oxygen atom has a higher EN the bonding electrons will be pulled closer to the oxygen atom This results in 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 oxygen atom will be slightly negative, and around the hydrogens will be slightly positive

Electronegativity and Covalent Bonds The symbol, δ+ represents a partial positive charge (less than +1) and δ− represents a partial negative charge (less than −1). Since the bond is polarized into a positive area and a negative area the bond has a “bond dipole”.

Electronegativity and Covalent Bonds The arrow points to the atom with the higher EN. p.178

Electronegativity and Covalent Bonds Covalent bonds resulting from unequal (electronegativities) sharing of bonding electron pairs are called Polar Covalent Bonds

Electronegativity Homework #’s 7, 8, & 9 - p. 178 #’s 1, 2, & 3 - p. 180

Bond Energy (pp. 179-180) 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.)