ICPS

 Objectives: 

Electronegativity (Unit 7)

• Define & Explain: Electronegativity 

• Trends in electronegativity on the periodic table (across a period; down a group)

  • Use Coulomb’s Law to explain the trends (relate to Zeff) 

Covalent Bonding (Unit 8)

• Define: Covalent Bonding 

• Naming covalent compounds (using prefixes)

• Single, double, and triple bonds 

  • Relate bond order to bond strength to bond length 

• Draw Lewis structures with resonance structures 

• Molecule shapes & geometry using VSEPR 

• 3-D representations of molecules using VSEPR 

• Polar v.s. Nonpolar covalent bonds (based on electronegativity)

  • Relate molecular shape to molecular polarity

Intermolecular Forces (Unit 8)

• Describe Intermolecular Forces (LDF, dipole-dipole, hydrogen bonds)

• Relation to physical properties (boiling/melting points, solubility, etc) 

Electronegativity

Electronegativity Definition 

Definition: The attractive force one atom’s nucleus has for another atom’s valence electrons when covalently bonding 

• Noble gases do not have electronegativity because they don’t like to bond

• Electronegativity only comes into play when atoms are covalently bonding

Trends Across the Periodic Table

Fe=kQ1Q2r2

Q1– charge of the valence electrons of 2nd atom | Q2– Zeff (charge of the nucleus of 1st atom on 2nd atom’s electrons) | r – atomic radius of 1st atom

• Electronegativity is directly proportional to Fe (electrostatic force) because greater force means greater attraction

L → R (across a period)

• Electronegativity Increases 

  • Fe=k(same)(up)(down)  → Fe increases; EN increases

  • Q1 stays the same because electrons’ charge stays constant; 

  • Zeff (Q2) increases because #p increase across a period

  • Atomic radius (r) decreases because Zeff increases

T → B (down a group)

• Electronegativity Decreases 

  • Fe=k(same)(same)(up) → Fe decreases; EN decreases

  • Q1 stays the same

  • Zeff (Q2) stays the same because #p and #inner electrons increase at the same rate down a group

  • Atomic radius (r) increases because more shells 

• Not a periodic trend → only exists when two atoms are covalently bonding (doesn’t exist for 1 atom only)

Covalent Bonding

Covalent Bonding 

Definition: The mutual attraction between 2 atoms’ nuclei for each other’s valence electrons. 

• Only happens between nonmetals 

• The number of covalent bonds an atom forms depends on the number of electrons needed to fill its valence shell

  • Hydrogen & Halogens only form one covalent bond (only needs 1 more electron to fill their valence shell)

  • Bonding only involves valence electrons

Diagram: 

Explanation: 

As atoms approach one another, their electrons start detecting each other, creating slight repulsive forces (too slight to really matter). As the atoms get closer, their positive nuclei attract each other’s electrons. This attractive force decreases potential energy (becomes negative), as shown in the diagram. The bottom of the curve indicates the distance at which the bonded atoms are most stable (lowest energy). The magnitude of the potential energy at that point is how much energy is needed to break the covalent bond. If the atoms move closer, their positive nuclei start to repel. This repulsive force is very strong and dominates the attractive forces, causing a spike in potential energy. Eventually, the repulsive force will break the covalent bond. 

• Negative potential energy indicates that energy must be provided to break the covalent bond

• Positive potential energy indicates that energy must be provided to form the covalent bond (push atoms together)

Naming Covalent Compounds

Only binary covalent compounds (2 types of atoms)

1. 1st element: keeps regular name (e.g Cl→chlorine; O→oxygen)

2. 2nd element: changes to ‘-ide’ name (e.g Cl→chloride; O→oxide)

3. Add prefixes to indicate number of atoms 

   1. *don’t use ‘mono’ before 1st element (keep its original name)

All -ide names

• Only nonmetals (excluding noble gases) covalently bond

• Naming only applies to these elements: 

Hydrogen — Hydride

Carbon — Carbide

Nitrogen — Nitride

Oxygen — Oxide

Fluorine — Fluoride 

Phosphorus — Phosphide 

Sulfur — Sulfide 

Chlorine — Chloride

Selenium — Selenide

Bromine — Bromide

Iodine — Iodide

Prefixes

Examples

N2O Dinitrogen Monoxide

S2Cl2 Disulfur Dichloride 

H2O Dihydrogen Monoxide 

CO2 Carbon Dioxide (don’t include ‘mono’ for 1st element)

Carbon tetrachloride CCl4

Sulfur trioxide SO3

Single, Double, and Triple Bonds

Single Bonds

• One shared pair of electrons (two electrons)

  • (X-X) 

• Lowest energy/strength → longer length

Double Bond 

• Two shared pairs of electrons (four electrons)

  • (X=X) 

Triple Bond 

• Three shared pairs of electrons (six electrons)

  • (X≡X) 

• Highest energy/strength → shorter length 

Bond Order: number of bonds between a pair of atoms (single, double, triple)

Bond Length: average distance between nuclei of two bonded atoms 

• An average because the atoms are constantly moving & changing distance

• X-coordinate of the bottom of the curve in the diagram above

Bond Energy/Strength: average amount of energy required to break the chemical bond

• An average because it varies depending on the specific molecule the bond is in

Note: can only compare bond orders within the same elements (can’t compare hydrogen triple-bond to nitrogen single-bond because the element changes bond length & strength)

Lewis Structures 

• 2-D representation of the bonding between atoms’ 

  • Only deals with valence electrons (includes bonding pairs & lone pairs)

How to Draw:

1. Count valence electrons in the molecule

2. Identify central atom

   1. Least electronegative atom (least attractive to electrons)

   2. Usually the first atom written 

   3. Never hydrogen or halogens (these can only form 1 covalent bond because their valence shell is 1 away from being full)

3. Create a skeletal structure 

   1. Use lines to indicate bonds; each line represents 1 electron pair

      1. Double bonds have 2 lines; triple bonds have 3 lines

   2. Four atoms around one central atom 

4. Distribute remaining electrons 

   1. Add lone pairs or create double/triple bonds 

   2. Each atom needs a full valence shell (eight electrons, except hydrogen)

      1. Each bond (one line) is 2 electrons 

   3. Total # of electrons on diagram needs to equal total valence electrons

Examples (simple molecules):

H2O

1. Valence electrons: 8 

   1. 2 hydrogen & 1 oxygen

   2. H: 1 valence O: 6 valence 

   3. Add together: 21+16=2+6=8

2. Central Atom: Oxygen (hydrogen cannot be central atom)

3. Skeletal Structure:

4. Remaining Electrons:

SCl2

1. Valence electrons: 20

1. 1 Sulfur & 2 Chlorine

2. S: 6 valence Cl: 7 valence 

3. Add together: 16+27=6+14=20

5. Central Atom: Sulfur (less electronegative)

6. Skeletal Structure:

7. Remaining Electrons:

Lewis Structures for Polyatomic Ions 

Polyatomic Ions—covalently bonded molecules with a charge

• Add brackets around the diagram 

• Write charge on top right corner

Example: 

(element: CO3 | charge: 2-)

Resonance

When there’s a double or triple bond that can be in multiple locations. The actual structure of a molecule with resonance is a hybrid of all the resonance structures. 

Indicating Resonance

• Draw all resonance structures 

• Connect the diagrams with double arrows 

Example:

Element: NO3 Resonance: 3

Reasonable Resonance Structures (Lab): 

Resonance structures are equally likely to occur (electrons involved constantly change positions between the possible structures). 

Reasonable Example:

SO₂ 

• Both oxygens have the same electronegativity (same pull on electrons)

  • One oxygen does not attract electrons more than the other

  • Electrons equally likely to be at either oxygen

• Both resonance structures equally likely to occur

• Bond length: one and a half (electrons constantly switching positions, so bond length is the average)

Unreasonable Example:

CO₂  

• Structures 2 & 3 are unlikely to occur

  • Both oxygens have equal electronegativity — electrons are unlikely to gather near one oxygen 

• Structure 1 is the most likely → unreasonable resonance structurs

  • Electrons likely to spread out equally between two oxygens

VSEPR (Valence Shell Electron Pair Repulsion Theory)

• Theory telling us how electrons repel during bonding

• Used to determine the 3-D geometry & shape of molecular structures

  • Lewis-structure is only 2-D

Theory: 

1. Electron groups (lone pairs & bonding pairs) repel (negative on negative) and will arrange themselves in a way that maximises distance between them. 

• Creates biggest angle possible between electron groups

2. Lone pairs repel more than bonding pairs

• Electrons in lone pairs repel each other more

• Lone pairs repel bonding pairs more

3. Angle of molecular structure varies based on lone pairs

• With lone pairs, the angle between bonding pairs are less than the base angle because lone pairs repel more and take up a greater angle, forcing the bonding pairs closer together (smaller angle)

Lewis v.s. VSEPR

Molecular Geometry

• Examines the location of electrons on the central atom(s)

Steps:

1. Draw Lewis Structure

2. Determine number of electron locations (on central atom)

   1. One bond location (single, double, or triple) is one location

   2. One lone pair is one location

3. Name the geometry depending on the number of electron locations

Types: 

Molecular Shapes

• Examines electron locations relative to each other (#bonding locations v.s. #lone pairs)

Steps:

1. Draw Lewis Structure

2. Determine number of bonding locations v.s. lone pairs

3. Name the shape

Types: 

VSEPR Relation

• VSEPR theory describes the 3-D shape 

  • Linear: no lone pair; bonding pairs repel to create biggest angle (180º); a straight line

  • Bent: lone pair pushes away two bonding pairs, creating a bend

  • Trigonal planar: no lone pair; bonding pairs repel to create biggest angle (120º); exists on one plane

  • Trigonal pyramidal: lone pair pushes away three bonding pairs, creating pyramid-like shape

  • Tetrahedral: no lone pair; bonding pairs repel to create biggest angle (109.5º)

3-D Representations 

• Depict the molecule’s 3-D shape 

  • Lone pairs are NOT drawn (only bonds)

  • Keeps as many connected atoms on one plane as possible

• Wedges

  • Indicates the bonded atom is on a different plane

  • Different colored wedges: atoms on different planes (only tetrahedrals)

Polarity

Polar Covalent Bonds: 

When 2 bonded atom have differing electronegavities (one atom pulls on electrons harder than the other)

• Electrons spend more time closer to the more electronegative atom

  • Creates partial charge 

  • More electronegative atom: negative partial charge 

  • Less electronegative atom: positive partial charge

• For ICPS: bonded atoms with an electronegativity difference greater than 0.35 create polar bonds 

Dipole Depiction

Nonpolar Covalent Bonds:

When the difference in the electronegativity of 2 bonded atoms is less than or equal to 0.35

• CH (carbon–hydrogen) bonds are not polar 

  • C: 2.55 | O: 2.20 (difference in EN = 0.35)

Molecular Polarity

A molecule is polar if there is an overall partial charge on one end of the molecule (when the dipoles don’t cancel out)

• Depends on molecular shape

Examples

Polar:

Nonpolar: (all CH bonds, which are nonpolar)

Intermolecular Forces

Intermolecular Forces: Forces between molecules

Strength of IMFs LDF (weakest) → Dipole-Dipole → Hydrogen Bond (strongest)

London Dispersion Forces (LDF)

Temporary attractive force between opposite temporary dipoles of molecules. LDF is a relatively weak intermolecular force. 

Temporary Dipoles

• When a molecule’s electrons are dispersed unevenly, causing more of them to gather on one side, creating temporary partial charges/dipoles

  • More electrons=negative partial charge

• Opposite ends of temporary dipoles on two molecules can attract

• More electrons = Greater LDF 

  • The more electrons a molecule has, the more unevenly dispersed they are likely to be, creating greater temporary partial charges

Dipole-Dipole Interactions

Attractive force between opposite dipoles of molecules (positive & negative end); force is stronger & more permanent than LDF

• Molecules need to be polar (needs permanent dipoles that do not cancel out)

Hydrogen Bonding

A type of dipole-dipole interaction that is particularly strong (not actually a bond, but an intermolecular force) 

Molecules that hydrogen bond have strong partial charges because of larger electronegativity differences between atoms in the molecule, making the attractive intermolecular force stronger. 

Molecules must have N–H, O–H, or F–H bond (hence “hydrogen bond”)

• Hydrogen has a low electronegativity (2.20), allowing for greater EN differences with other atoms it covalently bonds with

• Greater EN → stronger partial charges → stronger dipole-dipole force

Physical Properties 

Intermolecular forces bind molecules together. Their strength determines the state of the compound (solid, liquid, gas). 

• Solids: strong IMF; molecules are held tightly together — requires more energy to overcome IMF & break molecules apart

• Gases: weak IMF; molecules held loosely together — requires little energy to overcome IMF & break molecules apart

Phase changes: changing between solid, liquid, and gas states

• Putting in energy to break IMF can make solids become liquid or gas

• Taking away energy can make gases become liquid or solid

Boiling/Melting points

• Stronger IMF = higher boiling/melting points

  • Boiling & melting points are temperatures where substances become gas or liquid 

  • Stronger IMF means more energy is required to break forces & separate molecules → higher boiling/melting points (more heat energy)

Evaporation Rate (IMF Lab)

• A cooling process—the rate at which a liquid becomes gas and escapes into the air (a phase change)

• Can be measured by cooling rate (∆T/∆t) — change in temp/change in time

  • In the evaporation process, energy is provided to break the IMFs holding the compound in its liquid state, converting it to the higher-energy state of vapor (gas) 

  • As liquid evaporates, the higher-energy gas molecules (high energy=high temp) escape into the air, taking the energy with it

  • Lower-energy molecules (low temp) remain, creating a cooling effect 

• Stronger IMF = slower evaporation rate

  • Requires more energy to break IMF & allow liquids to vaporise (become gas)

IMF Lab

R2 — indicates how well the trendline models the data 

• R2 = 1: line of regression perfectly fits the data

  • Variation in the dependent variable is explained by the independent variables 

• R2 = 0: line of regression does not explain the data at all

  • Variations in the dependent variable are not caused by the independent variables at all

• Low R2 — there are other variables (extraneous variables) apart from the considered independent variables influencing the dependent variable 

• High R2 — the independent variables considered are the main factors influencing the dependent variable

Extraneous Variables

Factors that may influence the dependent variable or controlled conditions but are not considered an independent variable in the experiment.