Unit 5: Bonding

Soap This chapter will introduce the chemistry needed to understand how soap works Section 5.1: Types of Bonds Section 5.2: Drawing Molecules Section 5.3: Compounds in 3D Section 5.4: Polarity of Molecules Section 5.5: Intermolecular Forces Section 5.6: Intermolecular Forces and Properties

Section 1—Types of Bonds

Why do atoms bond? Atoms are most stable when they’re outer shell of electrons is full Atoms bond to fill this outer shell For most atoms, this means having 8 electrons in their valence shell Called the Octet Rule Common exceptions are Hydrogen and Helium which can only hold 2 electrons Called the Duet Rule

Remember Valence Electrons Valence electrons: the electrons found in the highest energy level Short Cut Rule: the group # next to the letter A of the Representative elements represents the valence number

LEWIS DOT DIAGRAMS

Remember, Lewis Dot diagrams show their valence electrons When atoms bond, they have 4 orbitals available (1 “s” and 3 “p’s”). There are 4 places to put electrons Put one in each spot before doubling up! Example: Draw the Lewis Structure for an oxygen atom

What is a Bond & Why does it form? Like glue holding atoms together It’s really forces of attraction between valence electrons holding atoms together They form because it lowers the potential energy of the atoms and creates stability

Three Types of Bonding

Ionic Bonding—Metal + Non-metal Therefore, metals tend to lose their electrons and non-metals gain electrons Metals become cations (positively charged) Non-metals become anions (negatively charged) There is a transfer of electrons The cation & anion are electrostatically attracted because of their charges—forming an ionic bond

One way valence shells become full - - - - - - - - - - - - - Na - Cl - - - - - - - - - - - - - - Sodium has 1 electron in it’s valence shell Chlorine has 7 electrons in it’s valence shell METALS: Some atoms give electrons away to reveal a full level underneath. NONMETALS: Some atoms gain electrons to fill their current valence shell.

One way valence shells become full - - - + - - - - - - - - - - - Na - Cl - - - - - - - - - - - - - - The sodium now is a cation (positive charge) and the chlorine is now an anion (negative charge). These opposite charges are now attracted, which is an ionic bond.

IONIC BONDING Sodium + Chlorine  Sodium Chloride

Transfer electrons in ionic bonding Transfer electrons from metal atoms to non-metal atoms, keeping track of their new charge Example: Draw the Lewis Structure for KCl

Transfer electrons in ionic bonding Transfer electrons from metal atoms to non-metal atoms, keeping track of their new charge Example: Draw the Lewis Structure for KCl K Cl Potassium has 1 electron Chlorine has 7 electrons

Transfer electrons in ionic bonding Transfer electrons from metal atoms to non-metal atoms, keeping track of their new charge Potassium has 1 electron Chlorine has 7 electrons Example: Draw the Lewis Structure for KCl +1 -1 K Cl

Add more atoms if needed If the transfer from one atom to another doesn’t result in full outer shells, add more atoms Example: Draw the Lewis Structure the ionic compound of Barium fluoride

Add more atoms if needed If the transfer from one atom to another doesn’t result in full outer shells, add more atoms Example: Draw the Lewis Structure the ionic compound of Barium fluoride Ba F Barium has 2 electron Fluorine has 7 electrons The fluorine is full, but the Barium isn’t!

Add more atoms if needed If the transfer from one atom to another doesn’t result in full outer shells, add more atoms Example: Draw the Lewis Structure the ionic compound of Barium fluoride Ba F F Barium has 2 electron Fluorine has 7 electrons Add another fluorine atom

Add more atoms if needed If the transfer from one atom to another doesn’t result in full outer shells, add more atoms -1 +2 Example: Draw the Lewis Structure the ionic compound of Barium fluoride Ba F -1 F Barium has 2 electron Fluorine has 7 electrons Now all have full valence shells and the charges are balanced—BaF2!

Covalent Bonding When two non-metals share electrons 2 Identical Non-metals that share electrons evenly form non-polar covalent bonds 2 Different Non-metals that share electrons un-evenly form polar covalent bonds

COVALENT BONDING NONPOLAR COVALENT BONDING Chlorine + Chlorine  Chlorine gas POLAR COVALENT BONDING Hydrogen + Fluorine  Hydrogen Fluoride + 

Metallic Bonding: composed only of metals Metal atom nuclei are bonded as a network as a pool of electrons that they share together. The electrons are free to move(mobile) throughout the structure—like a sea of electrons

Metallic Bonding

Bonding in Never Purely Ionic or Covalent: Use EN to indicate primary type 0% 5% 50% 100% 0-----.40 ------------------1.69------------------3.30 np polar ionic covalent covalent If the electronegativity difference is: Between 0-.40 NONPOLAR COVALENT Between .41 and 1.69 POLAR COVALENT Greater than 1.7 IONIC

Figure 12.4 The three possible types of bonds. nonpolar Sharing electrons equally polar Sharing electrons unequally ionic Transfer of electrons

CH4 F2 H2O NaF NM/NM M/NM Examples 2.5 – 2.1 = .4 non-polar covalent 4.0 -4.0= 0 nonpolar covalent 3.5 – 2.1 = 1.4 polar covalent 4.0 – .9 = 3.1 ionic

Bond type affects properties There are always exceptions to these generalizations (especially for very small or very big molecules), but overall the pattern is correct

Melting/Boiling Points Of Compounds Ionic Compounds tend to have VERY HIGH melting/boiling points as it’s hard to pull apart those electrostatic attractions of the ionic bond. Strongest forces of attraction. These compounds are found as solids under normal conditions Metals have HIGH/INTERMEDIATE melting/boiling points. Metallic bonds vary in strength. In general, strong forces of attraction. All are found as solids under normal conditions except Hg

Melting/Boiling Points Of Compounds Covalent Compounds Polar Covalent molecules have the next highest melting points. Weak forces of attraction. These compounds are soft solids or liquids under normal conditions Non-polar covalent molecules have the lowest melting/boiling points. Weakest forces of attraction. These compounds are found as liquids or gases

Solubility in Water (polar molecule) Ionic & polar covalent compounds tend to be soluble in water Non-polar compounds & metals tend to be insoluble in water Use this: Like Dissolves Like rule of thumb. Polar dissolves Polar Nonpolar dissolves Nonpolar

Like Dissolves Like

Conductivity of Electricity In order to conduct electricity, charge must be able to move or flow Metals have free-moving electrons—they can conduct electricity in BOTH solid & liquid state Ionic compounds ONLY have free-floating ions when dissolved in water, aqueous form or when melted (liquid state) Covalent compounds never have charges free to move and therefore cannot conduct electricity Exception: POLAR COVALENT Molecules of acids & bases can conduct electricity when dissolved in water

Ionic Compounds ONLY CONDUCT X ☺

Structures: Ionic Compounds(Crystals) Ionic compounds form a lattice framework of positive and negative ions They pack together so that the like-charge repulsions are minimized while the opposite-charge attractions are enhanced. Na+1 Cl-1

Metals: sea of electrons Structures Metals: sea of electrons Covalent compounds: neutral, separate discrete molecules

Visual Representation of 3 Bonding Types

Section 2—Drawing Molecules of Covalent Compounds

Tips for arranging atoms In general, write out the atoms in the same order as they appear in the chemical formula Hydrogen & Halogens (F, Cl, Br, I) can only bond with one other atom Always put them around the outside The least electronegative atom is usually in the middle; Carbon always goes in the middle

Steps to Drawing Lewis Structures 1. Decide how many valence electrons are around each atom 2. Arrange the atoms in a skeletal structure and connect them with a bonding pair of electrons. 3. Place remaining electrons around atoms so they each acquire 8 electrons. Exception is H .

Example: Carbon Tetrachloride Carbon has 4 electrons = 8 electrons Each hydrogen has 1 Example: Draw the Lewis Structure for CH4 Remember, “H” can’t go in the middle… put them around the Carbon! H H C H H

Count electrons around each atom Any electron that is being shared (between two atoms) gets to be counted by both atoms! All atoms are full with 8 valence electrons (except H—can only hold 2) H Example: Draw the Lewis Structure for CH4 Carbon has 8 H C H Each Hydrogen has 2 H All have full valence shells—drawing is correct!

Bonding Pair Pair of electrons shared by two atoms…they form the “bond” H C Bonding pair

Try These CH3I PCl3

What if they’re not all full after that? Multiple Covalent Bonds Are needed when there is not enough electrons to complete an octet To satisfy: move lone pair in between atoms to satisfy the duet/octet rule

Examples with Double Bonds Draw the Lewis Structure for CH2O H C O Remember that hydrogen atoms can’t go in the middle!

Draw the Lewis Structure for CH2O The two hydrogen atoms are full But the carbon and oxygen only have 7 each! H Example: Draw the Lewis Structure for CH2O H C O

Draw the Lewis Structure for CH2O But they each have a single, unshared electron. They could share those with each other! H Example: Draw the Lewis Structure for CH2O H C O

Draw the Lewis Structure for CH2O Now the carbon and oxygen both have a full valence! H Example: Draw the Lewis Structure for CH2O H C O

Double Bonds & Lone Pairs Double bonds are when 2 pairs of electrons are shared between the same two atoms Lone pairs are a pair of electrons not shared—only one atom “counts” them H C O Lone pair Double Bond

And when a double bond isn’t enough… Sometimes forming a double bond still isn’t enough to have all the valence shells full Example: Draw the Lewis Structure for C2H2

Draw the Lewis Structure for C2H2 … Example: Draw the Lewis Structure for C2H2 H C C H Remember that hydrogen atoms can’t go in the middle!

Draw the Lewis Structure for C2H2 … Each carbon atom only has 7 electrons…not full Example: Draw the Lewis Structure for C2H2 H C C H

Draw the Lewis Structure for C2H2 But they each have an un-paired electron left! Example: Draw the Lewis Structure for C2H2 H C C H

Draw the Lewis Structure for C2H2 Now they each have 8 electrons! Example: Draw the Lewis Structure for C2H2 H C C H

Triple Bonds A Triple Bond occurs when two atoms share 3 pairs of electrons H C Triple Bond

TRY These: HCN CO2

Properties of multiple bonds Single Bond Double Bond Triple Bond Longer & Weaker bonds (Lower Bond Energy: takes less energy to break) Shorter (atoms closer together & Stronger bonds (Higher Bond Energy: takes more energy to break)

Bond Dissociation Energy

Draw the Lewis Structure for CO3-2 Polyatomic Ions They are a group of atoms bonded together that have an overall charge Example: Draw the Lewis Structure for CO3-2

Drawing Polyatomic Ions Example: Draw the Lewis Structure for CO3-2 O C O O When there’s a single atom of one element, put it in the middle

Drawing Polyatomic Ions None of the atoms have full valence shells…they all have 7! The carbon can double bond with one of the oxygen atoms Example: Draw the Lewis Structure for CO3-2 O C O O

Drawing polyatomic Ions Now the Carbon and the one oxygen have 8…but the other two oxygen atoms still only have 7 Example: Draw the Lewis Structure for CO3-2 O C O O This is a polyatomic ion with a charge of “-2”…that means we get to “add” 2 electrons!

Drawing Polyatomic Ions Now the Carbon and the one oxygen have 8…but the other two oxygen atoms still only have 7 -2 Example: Draw the Lewis Structure for CO3-2 O C O O This is a polyatomic ion with a charge of “-2”…that means we get to “add” 2 electrons!

Covalent bond within…ionic bond between Polyatomic ions have a covalent bond within themselves… But can ionic bond with other ions Na O C O O Covalent bonds within Na

Covalent bond within…ionic bond between Polyatomic ions have a covalent bond within themselves… But an ionic bond with other ions +1 -2 Na O C O Ionic bond with other ions O Covalent bonds within +1 Na

Section 3—Molecules in 3D

Bonds repel each other Bonds are pairs electrons. Electrons are negatively charged Negative charges repel other negative charges SO bonds repel each other Molecules arrange themselves in 3-D so that the bonds are as far apart as possible

Valence Shell Electron Pair Repulsion Theory (VSEPR Theory) Outer shell of electrons involved in bonding Bonds are made of electron pairs Those electron pairs repel each other Attempts to explain behavior This theory attempts to explain the 3-D shape of molecules.

Coding for a Shape A code can help you connect the molecule to its shape. “A” stands for the central atom “B” stands for the number of bonding atoms off the central atom. “E” stands for the number of lone pair coming off the CENTRAL atom.

What shapes do molecules form? Linear 2 bonds, no lone pairs OR any 2 atom molecule AB2 BeCl2 NO CODE: a 2 atom molecule has no central atom Linear HCl 180°

What shapes do molecules form? Trigonal planar 3 bonds, no lone pairs AB3 BF3 Indicates a bond going away from you 120° Indicates a bond coming out at you

What shapes do molecules form? Tetrahedral 4 bonds, no lone pairs CH4 AB4 109.5°

What shapes do molecules form? Trigonal bipyramidal 5 bonds, no lone pairs PCl5 AB5 90° and 120°

What shapes do molecules form? Octahedral 6 bonds, no lone pairs SF6 AB6 90°

How is shape affected by Lone Pairs? Lone pairs are electrons, too…they must be taken into account when determining molecule shape since they repel the other bonds as well. But only take into account lone pairs around the CENTRAL atom, not the outside atoms!

What shapes do molecules form? Bent 2 bonds, 1 lone pair SO2 AB2E Bent 2 bonds, 2 lone pairs H2O AB2E2

What shapes do molecules form? NH3 Trigonal pyramidal 3 bonds, 1 lone pair AB3E

Lone Pairs take up more space Lone pairs aren’t “controlled” by a nucleus (positive charge) on both sides, but only on one side. This means they “spread out” more than a bonding pair. They distort the angle of the molecule’s bonds away from the lone pair.

Example of angle distortion 105° O 109.5° C

Rotating Molecular Shapes http://intro.chem.okstate.edu/1314F00/Lecture/Chapter10/VSEPR.html

Section 4—Polarity of Molecules

Showing Partial Charges of POLAR BONDS Determine which element in the bond is more electronegative. Use the symbol “-” to express the more electronegative atom and the symbol “+” for the least electronegative atom Use of an arrow pointing towards the partial negative atom with a “plus” tail at the partial positive atom. The arrow represents DIPOLES which are: forces created due to a separation of opposite charge!( a difference in charged ends) C O H + - C O H

Let’s Practice: If the bond is polar, draw the polarity arrow C – H O—Cl F—F P—N

If the bond is polar, draw the polarity arrow Let’s Practice C – H O—Cl F—F P—N 2.5 – 2.1 = 0.4 non-polar*exception 3.5 – 3.0 = 0.5 polar 4.0 – 4.0 = 0.0 non-polar 3.0 – 2.1 = - 0.9 polar Example: If the bond is polar, draw the polarity arrow

Predicting Bond Polarity Two atoms sharing equally: Draw N2 Each nitrogen atom has an electronegativity of 3.0. They pull evenly on the shared electrons. The electrons are not closer to one or the other of the atoms. This is a non-polar covalent bond. Most compounds that contain ONLY non-polar covalent bonds are NONPOLAR MOLECULES!

Atoms sharing unequally: Draw H2S Electronegativities: H = 2.1 sulfur = 2.5 The sulfur pulls on the electrons slightly more, pulling them slightly towards the sulfur. This is a polar covalent bond Asymmetrical (lop-sided) compounds that contain polar covalent bonds are POLAR MOLECULES

Sharing unevenly: Draw CH2O Electronegativities: H = 2.1 C = 2.5 O = 3.5 The carbon-hydrogen difference isn’t great enough to create partial charges : It’s actually a NON POLAR bond: **Exception to the rule But the oxygen atoms pulls significantly harder on the electrons than the carbon does. This does create a polar covalent bond. Asymmetrical (lop-sided) compounds that contain polar covalent bonds are POLAR MOLECULES

Sharing unevenly: Draw BCl3 Electronegativities: B = 2.0 Cl = 3.0 The chlorine pulls on the electrons slightly more than boron. This is a polar covalent bond Symmetrical compounds that contain polar covalent bonds are NONPOLAR MOLECULES

How do Dipoles Cancel? Dipoles must move in equal but opposite directions in order for the forces to cancel The molecule is classified as NONPOLAR.

Polar Bonds versus Polar Molecules Not every molecule with a polar bond is polar itself If the polar bonds form Dipoles that cancel out then the molecule is overall non-polar. The dipoles cancel out. No net dipole The dipoles do not cancel out. Net dipole This one is hard to tell!

The Importance of VSEPR in Predicting Polarity. Shape is important. All molecules must be drawn in the correct shape to see the proper canceling of dipoles to determine if its polar or nonpolar. Water drawn this way shows all the dipoles canceling out. O H But water drawn in the correct VSEPR structure, bent, shows the dipoles don’t cancel out! Net dipole H O H

Practice: Draw the molecule NH3 Example: Is NH3 a polar molecule?

N H Example: Is NH3 a polar molecule? Electronegativities: N = 3.0 Difference = 0.9 Polar bonds VSEPR shape = Trigonal pyramidal Yes, NH3 is polar Net dipole

Draw the molecule for dihydrogen monoxide, H2O. Is it polar or non-polar? What shape?

O H Net dipole Yes, H2O is polar Is water polar or non-polar? What shape? Electronegativities: H = 2.1 O = 3.0 Difference = .9 Polar bonds VSEPR shape = bent

Draw the molecule CO2 Example: Is CO2 a polar molecule?

Draw the molecule of carbon dioxide, CO2 Example: Is CO2 a polar molecule? Electronegativities: C = 2.5 O = 3.5 Difference = 1.0 Polar bonds VSEPR shape = linear No CO2 is nonpolar Dipole cancels

Draw the molecule for carbon tetrachloride, CCl4. Electronegativities: C = 2.5 H = 2.1 Difference = .4 Nonpolar bonds VSEPR shape = tetrahedral H H C H H No Net dipole Yes, CH4 is nonpolar

Let’s make this Simple: Nonpolar bonds= Nonpolar molecule. Polar bonds with a lone pair on the central atom = most likely a polar molecule Polar bonds & no lone pair on central atom & all terminal atoms are the same= nonpolar molecule. If terminal atoms are different, its polar.

Section 5—Intermolecular Forces http://www.wisc-online.com/Objects/ViewObject.aspx?ID=GCH6804

Intramolecular Forces- versus Inter-molecular Forces So far this chapter has been discussing “Intramolecular Forces” Intramolecular forces = forces within the molecule AKA: chemical bonds

Intra-Force Intra-Force

Breaking Intramolecular forces Breaking of intramolecular forces (within the molecule) is a chemical change Example: 2 H2 + O2  2 H2O Bonds are broken within the molecules and new bonds are formed to form new molecules Requires a larger amount of energy to break than an intermolecular force

Inter-molecular Forces Intermolecular forces = forces between separate molecules

Intra-Force Inter-Force

Breaking Intermolecular forces Breaking of intermolecular forces (between separate molecules) is a physical change Breaking glass & Boiling water are examples Example: 2H2O(l)  2 H2O(g) Does not require as much energy to break compared to an intramolecular force

London Dispersion Forces (LDF): are the primary force between nonpolar molecules but are found in all molecules! Electrons move around the nuclei & They can momentarily “gang up” on one side This lop-sidedness of electrons creates a partial negative charge in one area and a partial positive charge in another. + Positively charged nucleus - Negatively charged electron + - + - Electrons are fairly evenly dispersed. + - As electrons move, they “gang up” on one side.

London Dispersion Forces (LDF) Once the electrons have “ganged up”, a temporary dipole is created & the molecule is now temporarily polar. The positive end of one temporarily polar molecule can be attracted to the negative end of another molecule. + - + - LDF nonpolar nonpolar

London Dispersion Forces (LDF)

Strength of London Dispersion Forces (LDF) Since, the electrons will move again, returning the molecule back to non-polar The polarity was temporary, therefore the molecule cannot always form LDF. London Dispersion Forces are the weakest IMF because molecules can’t form it all the time, only temporarily

Strength of London Dispersion Forces (LDF) All molecules have electrons…all molecules can have London Dispersion Forces Larger molecules (greater molar mass) have stronger London Dispersion Forces than smaller molecules. Larger molecules have more electrons. The more electrons that gang-up, the larger the partial negative charge.

London Forces explain why Chlorine is a gas, Bromine is a liquid and Iodine is a solid! Chlorine Gas = 34 e- Bromine Liquid = 70 e- GREATER # ELECTRONS, Greater Molar Mass, STRONGER FORCES Iodine Solid = 106 e-

Dipole- Dipole Forces: primary force between polar molecules They have permanent permanent dipoles. The positive area of one polar molecule can be attracted to the negative area of another molecule. The partial positive & negative poles are shown as + and -

Strength of Dipole Forces In general, Dipole forces are stronger than London Dispersion Forces Polar molecules always have a partial separation of charge. They always have the ability to form attractions with opposite charges + - + - Dipole force polar polar

Dipole-Dipole Forces

Hydrogen Bonding A special dipole force between polar molecules that contain a hydrogen atom connecting with F, O, or N of another molecule. (H is FON) A very strong dipole forms since F, O, and N are all very small, highly electronegative atoms.

Hydrogen Bond N H N H Hydrogen bond

Hydrogen Bonds The ladder rungs in a DNA molecule are hydrogen bonds between the base pairs, (AT and GC).

Special Case: Carbon Allotropes: Diamond vs. Graphite Diamond: Hard Tetrahedral-Special : Network Covalent Bonds Graphite: soft Strong Sheets of carbon rings but weak forces holding the sheets together

NETWORK COVALENT BONDS special covalent compounds compounds that contain only carbon (diamond, graphite) or silicon compounds (silicon dioxide- quartz) super strong bonding super high melting points http://www.youtube.com/watch?v=fuinLNKkknI

Rank the forces of attraction in order of weakest to strongest Rank the Intramolecular Forces: Ionic, Covalent, and Metallic Covalent< Metallic < Ionic Rank the Intermolecular Forces: Dipole, London Dispersion, Hydrogen bonding London Dispersion forces< Dipole- Dipole forces< Hydrogen bonding Rank ALL the Forces: London Dispersion forces< Dipole- Dipole forces< Hydrogen bonding <Covalent< Metallic< Ionic , NETWORK COVALENT

Tutorial: must be in Mozilla http://www.wisc-online.com/Objects/ViewObject.aspx?ID=GCH6804 Wisconsin online :intermolecular forces

Section.6—Intermolecular Forces & Properties

IMF’s and Properties The number and strength of the intermolecular forces affects the properties of the substance. Energy is needed to break IMF’s Energy is released when new IMF’s are formed

IMF’s and Changes in State From solid to liquid, some IMF’s are broken. From liquid to gas, the remainder of the IMF’s are broken Breaking IMF’s requires energy = ENDOTHERMIC PROCESS The stronger the IMF’s are, the higher the melting and boiling point

The Uniqueness of Water Water is a very small molecule In general small molecules have low melting and boiling points Based on it’s size, water should be a gas under normal conditions However, water is polar and can form dipole interactions and hydrogen bonding, It’s boiling point is much higher as a result!

Boiling Point of Polar Molecules

What is More Viscous? Molasses or Water?

Viscosity Viscosity is the resistance to flow Molasses is much more viscous than water Larger molecules and molecules with strong IMF’s become inter-twined and “stick” together more More sticking together = higher viscosity An increase in temperature will help break the IMF’s and make a substance less viscous

Solubility Solute: the substance that is dissolved Solvent: the is doing the dissolving

Solubility: the ability to dissolve Solvent, water (polar) + - - + Solute, sugar (polar) Water particles break IMF’s between other water molecules allowing them to spread out to form new IMF’s with the sugar molecules. - + - + - +

Solubility Solvent, water (polar) + - - + Solute, sugar (polar) As new IMF’s are formed, the solvent “carries off” the solute—this is “dissolving” - +

Solubility If the energy needed to break old IMF’s is much greater than the energy released when the new ones are formed, the process won’t occur An exception to this is if more energy is added somehow (such as heating)

Like Dissolves Like Polar solvents dissolve polar solutes Nonpolar solvents dissolve nonpolar solutes Polar solvents can also dissolve ionic compounds because of the charged ends of both

Therefore, oil and water don’t mix! Oil & Water Water is polar and can hydrogen bond, Oil is non-polar. Water can only form London Dispersion with the oil. That doesn’t release much energy Much more energy is required to break apart the water than is released when water and oil combine. Water has London Dispersion, Dipole forces and hydrogen bonding. That takes a lot of energy to break Therefore, oil and water don’t mix!

Surface Tension Surface tension is the resistance of a liquid to spread out. Due to higher IMF’s in the liquid, the more the molecules “stick” together, the less they want to spread out. Higher IMF’s =higher the surface tension.

Surface Tension of Water metal paper clip on water water forms “beads”