HL Chemistry Unit 3

Activation Energy

Minimum of input energy required for a chemical reaction to occur (get reactants to transition state when bonds are broken and able to form products)

Higher activation energy → slower reaction

Catalyst

Speeds up a chemical reaction without being consumed during the reaction

Increases the rate of reaction by decreasing the activation energy

Why are d-block elements good catalysts?

Incomplete d-orbitals and ability to have multiple oxidation states

- Weak bonds with reactants, supplies exterior electrons from 3d and 4s orbitals

- Able to take back electrons after the reaction so that the transition metal remains unchanged

Large surface area absorbs reactants so that they are in closer contact, making the reaction faster

Why are dark bottles used to store hydrogen peroxide? What is the catalyst for this reaction?

Dark bottles protect exterior light, which can provide activation energy, causing a reaction for hydrogen peroxide to decompose

MnO2 catalyst

What is the catalyst of the contact process?

Extracts sulfur, can convert sulfur dioxide to sulfur trioxide, and converts sulfur trioxide into concentrated sulfuric acid

Uses of sulfuric acid: manufacture explosives, acids, dyes, glue, car batteries

Vanadium(V) oxide catalyst

What is the catalyst of the Haber Process?

Converts hydrogen and nitrogen to ammonia, reversible reaction

Uses: manufacture plastics, explosives, textiles, pesticides, dyes, refrigerant gas

Iron catalyst

What is the catalyst in a Heme Group?

Ring-shaped iron-containing molecular component of hemoglobin, found mostly in lungs and tissues

Iron attaches with hemoglobin molecule or myoglobin, needed to bind and transport oxygen throughout body, helps in respiration, detoxification of drugs

What is the catalyst for the conversion of an alkene to an alkane?

Alkane: Carbon and hydrogen compound where each hydrogen is bonded to a carbon and carbon-carbon single bonds hold the atoms together

Alkene: Hydrocarbon with at least one carbon to carbon double bond, more reactive than alkane because of double bond

Nickel is used to convert the alkenes to alkanes when making trans and saturated fats, and margarine

What does Vitamin B-12 (cobalamin) do?

Nerve cell function, red blood cell formation, DNA synthesis

Cobalt in the center, 4 coordination bonds with nitrogen, connected by C-CH3 methylene link and C-H

Acts as a catalyst for methylation reactions and isomerization reactions

What is the catalyst for the Catalytic Converter?

Converts most hydrocarbons, carbon monoxide, and nitrogen oxides into carbon dioxide, nitrogen, and water vapor

Makes pollutants from car less harmful by chemically converting them

3 way: Converts nitrogen oxides to elemental nitrogen and oxygen, carbon monoxide to carbon dioxide, hydrocarbons to carbon dioxide and water

2 way: oxidation of carbon monoxide and oxidation of hydrocarbons to carbon dioxide and water

Pt/Pd catalyst

VSEPR Theory

Valence Shell Electron Pair Repulsion theory

Pure Covalent Bond

Electrons shared evenly between atoms, even charge distribution

Non-polar Covalent Molecule

Dipoles cancel out

Polar Covalent Bond

Uneven distribution of electrons

Polar Covalent Molecule

Dipoles reinforce in one direction

Ion-dipole interaction

Negatively charged ions attract slightly positive atoms in molecule, positive charged ions attract slightly negative → Ionic compounds in polar solvents

Intermolecular Forces of Attraction

Intermolecular (inter → between) force (attractive force between molecules that holds solids and liquid together. tends to be weak between gaseous molecules)

Boiling water ≠ breaking bonds in molecule ⇒ breaking force/bond between molecules, still same substance

Collective referred to as "Van der Waals Forces"

Aggregates

Groups of molecules held together by intermolecular forces to create liquid and solids (gas forces are too weak)

London Dispersion Forces

Caused by instantaneous, temporary, momentary, random dipoles → random motion of electrons (clouds of probability → uneven distribution of charge)

Only intermolecular force between nonpolar molecules and in nonpolar aggregates

All molecules experience London Dispersion Forces

Not all molecules are participating in attraction at the same time → attraction is weak

Reinforced by induced dipoles (temporary dipole on one molecules forces a temporary dipole on an adjacent molecule)

Factors of strength of LDF

Larger molecules → more mass + electrons → stronger temporary dipoles → more force, stronger LDF

Larger surface area → more points of contact → stronger LDF

Linear > spherical molecules

Dipole-Dipole Interactions

Between polar molecules

Slightly negative end of one molecule is attracted to slightly positive end of an adjacent molecule

Result of permanent dipoles on polar molecules

Ordered fashion → molecules arrange in crystal structure

Hydrogen Bonds

Stronger D-D interaction but can happen (rarely) with instantaneous dipoles

Oxygen dissolves in hydrogen due to hydrogen bond

Only polar molecules → pair of lone electrons on F,O, N (very electronegative) and hydrogen (very electropositive) on different molecule with F,O, N

Hydrogen is attracted to lone pair of different F, O, N molecule

1/10 strength of covalent bond

Larger aggregate → more lone pairs available on F, O, N for hydrogens

Creates 3d crystal structure affect ⇒ why water is more dense frozen than liquid

Small molecule usually gases at room temperature, but strong hydrogen bonds allows it (water) to stay liquid

Intramolecular Force

Forces of attraction within molecules

Covalent bonds

Ionic bond: Chemical bond → forms ionic compounds → does not form molecules → not inter/intramolecular

Ionic Solid Bonding model

Anions and cations

Crystal lattice structure

- Simple cubic

- Body centered cubic

Bond strength

Larger ions → weaker ionic bonds (pack less tightly together than smaller ions )

Atom size increases → bond length increases → decrease in bond strength

Larger magnitude of charge → stronger ionic bond

Smaller interionic distances (more tightly packed) → more stable

Ionic Solid Properties

Brittle, not malleable

Hard → strong uniform forces

High melting/boiling point

Bonds hard to break

Electrical Conductivity as a liquid + aqueous

Free-moving ions, ions have dissociated with each other

Metallic Solid Bonding Model

Electrostatic attraction between cations and an electron cloud

Bond strength

- Radius of ion

- Charge of ion

Metallic Solid Properties

Luster

- Energy from photons absorbed, excites electrons, falls down to emit light

Malleable/ductile

- Non-directional bonds: Cations can slide around without needing to remove any strong bonds

Electrical conductivity as a solid and liquid

- Free-moving and delocalized electrons

- No aqueous, metals do not dissolve in water

Thermal conductivity, transfers heat easily

High melting/boiling point

- Strong metallic bond

Hard → Strong bond

Sonority →Ringing sound when struck

Network solids bonding model

Repeated structure of covalent bonds

3D

- Repeated vespr shapes

- All covalent

- Hard to break

2D

- Covalent bond within layers

- LDF between layers

1D

- Chain are strong covalent bonds

- Weak LDF within chains

Network Solids Properties

High melting/boiling point due to strong covalent bonds

Hardness very high due to the rigid 3D structures

Brittle --> Breaking bonds is the only way to move them

Not conductive except for graphene

- Lack free-moving electrons/charges

Heat conductivity is variable (ex. diamond is very conductive, while silicon dioxide is not)

Allotropes of Carbon

Carbon atoms bonded covalently in a continuous network

- Pure covalent bonds

Crystalline structure

Graphite, graphene, diamond, fullerene, C60 (buckminsterfullerene)

Graphite

Made of layers of graphene

Has delocalized electrons between layers (LDF)

Black but has luster

Opaque

Very soft

'Greasy' feeling

High melting point → strong covalent bonds

Soft → Weak van der waal forces

Electrically conductive

Thermally conductive

Graphene

2D graphite

Single layer of carbon atoms bound in a hexagonal honeycomb lattice

Formed through sp2 hybridization, creating a trigonal planar with 120° angles

Delocalized electron movement through p orbitals in the π bonded network

Lightweight

Transparent

High elasticity and flexibility

Very hard (for a 2D material)

Very thin

Good conductivity

Thermally conductive

Diamond

Atoms arranged in fixed structures, no individual molecules

Overlap of sp3 hybridized carbon atomic orbitals in infinite 3D structure

Graphite can be converted into diamond with high pressure + heat to rearrange bonds

Hardest naturally occurring substance (10 on Mohs Hardness Scale)

Used in cutting tools (e.g. diamond bit drills)

Refractive index of 2.42

High density of 3.50-3.53 g/cm3

Brittle

Varies widely in colour (colourless to yellow and pinks)

Good electrical insulator and good heat conductor

Fullerene

Hexagons and pentagons that form a spherical molecule, attaching a carbon atom to each vertex

Forms hollow "cage-life" structures

Carbon atoms are present in the sp2 hybridization form, linked by covalent bonds (double or single bonds)

Held together by LDF(not strong)

Slippery and excellent lubricants as weak intermolecular forces allow fullerene to slide past each other

Brittle/Soft due to weak Van der Waals interaction

Dark needle-like crystals

Lower melting point than graphene

Low water solubility

Not conductive

C60 (buckminsterfullerene)

Most common fullerene

Has 60 carbons with 20 hexagons and 12 pentagons

Has 90 covalent bonds between them, 60 single bonds and 30 double bonds

Highly structured and symmetrical

Polar molecular solids bond model

Held together by LDF and dipole dipole

- Same changes repel but isn't as significant as opposite attraction

Organised in lattice structure where oppositely charged ends are closest together

Shape affect strength of the bonds

- Solids has stronger bonds → Regular geometric shapes

Polar molecules solids properties

Soluble in polar solvents

Very soft, not often in solid form

Low melting and boiling point

Still higher than non polar

Low vapor pressure and volatility

Insulators

Non-polar molecular solids bond model

Symmetrical electron distribution + held together by LDF (induced dipoles) and strong covalent bonds within molecule

Stretched out shapes → more surface area → more LDF

Non-polar molecular solids properties

Very soft → weak forces

Low melting/boiling point

Weak forces/induced and instantaneous dipoles

High volatility/vapor pressure

Weaker forces/attraction, require less energy to overcome

Hard to undergo chemical change → Non-polar

Low electrical conductivity/thermal conductivity

No free moving electrons

Flexible/malleable/soft/brittle

Soluble in non-polar solvents

Low melting and boiling points

Hybridization

Combination of two or more orbitals (l) in the same energy level (n)

Ensures that electron pairs are at the same energy to interact in VSEPR models

Has energy intermediate between orbitals it came from + leaves last orbital unhybridized

Transition Metals

All are found in d-block (but not all d-block are transition)

Multiple oxidation states (commonly +2)

Can have magnetic properties

- Unpaired d-electrons (spin in same direction)

Atom of 1+ ions with incomplete d-orbitals

Make good catalysts

Transition metals can make "complex ions" with one or more "ligands"

Ligand

Molecule or ion with lone pairs of electrons that can be donated to a coordination bond

mono-, bi-, tri-, tetra-, dentate → how many "teeth" a ligand has → treat as 1 lone pair of electron per ligand

Complex Ion

Metal ion covalently bonded to one or more ligands

Why are transition metal complex ions coloured?

The metal ions will absorb specific colors (color → λ → f → ΔE → Δd) from white light

"Empty spaces" in higher energy split d-orbitals allowing for excitation of lower d orbitals

"Absorption of light is due to electrons in excited/promoted to higher energy levels"

The substance will appear to be coloured the complementary color of the one "removed"/absorbed

Factors that affect the colour of transition metal complex ions

Type of ion: Different metals have different numbers of protons and number of electrons between nucleus and d-orbitals (shielding)

- Changes how attractive metal nucleus is to d-orbitals

Type of ligand: More electronegative ligand will pull on d-orbitals with more forces than a less electronegative ligand

Oxidation state of the metal: Changes number of electrons in the d-orbitals

- Smaller change in colour

Number of ligands (can't change without changing type of ligand): More ligands will pull on d-orbitals with more combined force and vice versa

Orbital Overlap

Orbitals that have been hybridized overlap end-to-end or axially in sigma bonds

Axial Overlap/Head-On/End-to End

Forms sigma bonds

Parallel Overlap

Forms pi bonds

Common ligands

aqua (H2O)

hydroxy (OH-)

amine (NH3)

chloro (Cl-)

cyano (CN-)