Everything deck

Ionic compounds

Formed when electrons are transferred from one atom to another

Forms ions with complete valence shells

The cations and anions are attracted to each other by strong electrostatic forces

Have high melting opints

Naming binary ionic compounds

Cation first, Anion second, add -ide suffix

Covalent bonds

Sharing one or more pairs of electrons so each atoms achieves a noble gas configuration

Coordinate bonds

Both shared electrons come from the same atom

2 Electron domains

Linear, 180°

3 electron doains

Trigonal planar, 120°'

bent or V-shaped (smaller than 120°)

4 electron domains

Tetrahedral, 109.5°

Trigonal pyramidal, 107°

Bent, or V-shaped 105°

Differences in single, double. triple bonds

The more bonds in one, the shorter and stronger

Boiling point

Liquid to gas

All attractive forces between particles are broken

Good indication of the strength of intermolecular forces

Melting point

Crystal structure is broken down, but attractive forces between particles exist

Impurities impact

Impact structure and result in lower melting points

Diamond properties

Each carbon covalently bonded to four others

Giant tetrahedral diamond structure'

109°

1.54/10¹⁰ m long lengths.

All electrons are localized = Doesn’t conduct electricity

All bonds equally strong, so hard to break and high melting point

Graphite

Each Carbon is bonded to three other carbons

Layers of hexagonal rings

120°

1.42/10¹⁰ m

Weak attractive forces between layers = layers slide over each other = lubricant and waxy feeling

Between layers, electrons are delocalized = Good conducter of electricity

Graphene

One atom thick layer of Graphite

Extremely light

Semiconductor

200 times stronger than steel

Fullerenes

Large spheroidal molecules = hollow cage of sixty or more carbon atoms.

E.g. Buckminsterfullerene C₆₀

Hexagons and pentagons → geodesic spherical structure similar to a football

Silicon

1. Giant tetrahedral covalent structure like diamond.

2. hard brittle solid

3. high melting point (4000°C)

4. metalloid

5. semiconductor.

Silicon dioxide (silica)

SiO₂, quartz

Giant tetrahedral covalent structure similar to diamond and silicon.

Very hard and has a high melting point (1600°C)

No delocalized electrons = does not conduct electricity.

unipositive cation

positively charged ion with a charge +1

Bond polarity

More electronegative atoms exert greater attraction for electron pair(s)

One end of the bond will be more electron rich, resulting in bond dipole (polar bond)

Put number behind the atom with positive pole for how many bonds its pulling electrons with. E.g. NH₃ would have Nitrogen with 3§+

Shape can also cause polarity.

Allotropes

Element can exist in different crystalline forms.

Different bonding and structural patterns, thus different chemical and physical properties

Intermolecular forces

Van Der Waals Forces

London Dispersion Forces

Dipole-Dipole Forces

Hydrogen Bonds

Van Der Waals Forces

More inclusive term for intermolecular forces where a temporary dipole occurs between both polar and non-polar molecules

London Dispersion Forces (LDP)

Between NONPOLAR molecules or atoms

Temporary dipoles

• Can induce another dipole in a neighbouring particle

The strength of the dispersion forces increases with increased surface area/molar mass

→ When the carbon chain increases the boiling point increase

→ Branching decreases the surface area and thus the dispersion forces.

Dipole-Dipole Forces

Forces between POLAR molecules

Permanent dipoles

Attracted to each other by electrostatic forces.

Relatively weak, but the attraction is stronger than London dispersion forces.

Hydrogen Bonds

Between POLAR molecules, that have hydrogen is bonded with a small, highly electronegative element

e.g. N-H, O-H or F-H

As the electron from hydrogen is pulled away, the proton attracts a non-bonding pair of electrons from the F, N, or O atom

Strong dipole-dipole attraction

The strength increases with increased number of hydrogen bonds formed.

Solubility

Maximum amount of a solute that can dissolve in a specific amount of solvent

"Like tends to dissolve like"

As the carbon chain lengthens in organic materials with polar ends, they become less soluble in water

Polar solvent example

Water

Non-Polar solvent example

Heptane or tetrachloromethane

Solvent for both non-polar and polar compounds

Ethanol

Requirements for conduction

Electrons or ions that are free to move

Covalent compounds (except some like graphite) = Not conductive

Metals = Contain delocalized electrons = Excellent conductors

Molten ionic salts = conduct electricity, but are chemically decomposed in the process

Metallic Bonding

The attraction that two neighbouring positive ions have for the delocalized electrons between them.

Valence electrons detach from atoms to create a closely packed lattice of cations in a "sea" of delocalized electrons.

What does metallic bonding depend on?

The charge of the ions

Radius of the metal ion

Characteristics of metals

Malleable → bent and reshaped under pressure

Ductile → drawn out into a wire

= Because the closely packed layers of cations can slide over each other without breaking more bonds than are made.

Determining factors in metal melting point

Electron density (the size and charge of the ion formed) when the valence electrons are delocalized.

→ Smaller the metal ion formed = stronger the metallic bond and higher melting point.

The way in which the atoms are arranged in the solid metal

Most metals tend to have quite high melting points

Macroscopic properties

Observable, measurable characteristics of matter at a laboratory scale, visible to the naked eye or with low-magnification tools

Microscopic properties

Behavior, structure, and interactions of individual atoms, molecules, or ions, invisible to the naked eye. E.g., molecular geometry, bond energy, intermolecular forces

Three extreme cases for types of bonding

Caesium (metallic)

Fluorine (covalent)

Caesium fluoride (ionic)

Alloys

Homogeneous mixtures that are usually made up of two or more metals

Alters the properties of the metal → Distort the structure of the original metal as the bonding is less directional

Less ductile & less malleable than pure metals

Brass

Principle Metal: Copper

Added Metal: Zinc

Bronze

Principle Metal: Copper

Added Metal: Tin

Solder

Principle Metal: Lead

Added Metal: Tin

Pewter

Principle Metal: Lead

Added Metal: Copper, antimony, bismuth, or lead

Steel

Principal Metal: Iron

Added element: carbon

Extra: Adding chromium makes stainless steel, which has a much increased resistance to corrosion.

Polymers

Consist of monomers

Macromolecules with high molar masses

Subdivided into thermoplastics and thermosets

The longer the chain length of a polymer, the higher the strength and melting point

Cross-linking (or branching) and the orientation of the substituent groups can also affect particular properties

Thermoplastics

Many alkenes polermarize to thermoplastics

Soften when heated

Harden when cooled

Can be remoulded each time they are heated

Thermosetting polymers

E.g. Bakelite, polyurethanes, and vulcanized rubber

Form prepolymers in a soft solid or viscous state which, when cured, irreversibly turn into hardened thermosets

Elastomers

Flexible polymers

E.g. Rubber

Can be deformed under force but return to nearly their original shape once the stress is released.

Addition polymers

Long-chain molecules formed by joining many unsaturated monomer molecules (typically alkenes) together through a chain-reaction that breaks their double bonds

poly(2-methylpropene)

Butyl Rubber

Elastomer

Car tyre inner tubes and cling film

Polystyrene

Contains air trapped in the polymer

Light material

Good thermal insulator

Used in packaging as it has good shock absorbing properties

polymers

Long-chain, high-molecular-weight molecules

• Large molecule

• Covalent bonds 

→ low electrical conductivity / insulators

→ low thermal conductivity

→ low density

→ inert — non-biodegradable

Properties of polymers

• Light (made mostly of light elements C and H) & Durable (long so strong LDPs)

• Longer chain → more LDF — high softening point & decomposition points

• No delocalized electrons → thermal & electrical insulators

• Strong intermolecular forces makes it insoluble in water & many non-polar solvents

• No clear melting point since many contain chains of different lengths → Varying LDP

Environmental considerations

Reduce, Reuse, Recycle

• Takes lots of energy to recycle

• Plastics deteriorate when they are recycled

Natural polymers

Natural sources.

E.g. starch, lignin, silk, DNA, proteins…

Synthetic polymers

Made in laboratories.

E.g. plastics; PE, PET, kevlar, and synthetic textiles; nylon…

Polymerization

Chemical process that combines small molecules into long-chain, high-molecular-weight molecules.

Method of Polymerization

Polyaddition

Polyaddition

Double bonds will break and the monomer will form two covalent bonds with other monomers

Monomer

reactant

Repeating unit

Naming of polymers

Prefix poly- and the name of the monomer (e.g. ethene = polyethene)

Amphoteric

Can be both an acid and a base

acyclic

chain

cyclic

ring

heterocyclic

ring with other atoms

Saturated cmpound

Only single C-C bonds

Unsaturated compound

One C=C or C=-C bond

Alipahatic

No benzene ring

Polyunstaurated compound

Multiple C=C or C=-C bond

homologous series

series of organic compounds with the same general formula, but differ by one unit

- similar chemical properties but different physical

BP / MP Factors

• Length (longer -> more LDF)

• Branching (straight -> more LDF)

• Functional group (-OH forms hydrogen bonds, others from dipole-dipole)

Stereochemical formula

3D formula of molecules bonds

Bond INTO plane

Dashed wedge

Bond OUT OF plane

Solid wedge

Structural isomer

Compunds with same molecular but different structural formula

Chain isomers

Different branching

Functional group isomer

Different functional group

Position isomer

Functional group is at different position

primary alcohol

OH bonded to C bonded to one chain

Secondary alcohol

OH bonded to C bonded to two chains

Tertiary alcohol

OH bonded to C bonded to three chains

Primary, secondary, tertiary amide

N is bonded to x amout of carbon chains

Homolytic Fission

Each of the two atoms forming the bond retains one of the electrons, resulting in the formation of two free radicals

Heterolytic fission

Both of the electrons in the shared pair go to one of the atoms, resulting in a negative ion and a positive ion

Covalent bonds break

homyltically or heterolytically

Free radicals

Signified with a dot

Contain an unpaired electron

Are very reactive

Alkanes reactivity

Relatively unreactive due to strong C-C and C-H bonds

Combust with oxygen

Substitution with Halogens in UV light

Mechanism of chlorination of methane

1. Initation

   1. UV breaks the bond between two halogens as they are weaker than C-C or C-H bonds

   2. Creates two free radicals

2. Propagation

   1. Highly reactive Cl radical comes in contact with Methane, it produces H-Cl and a methyl radical

   2. Highly reactive methyl radical reacts with Cl₂ to produce chloromethane and sanother Cl radical

   3. Innitiates chain reaction

3. Termination

   1. Two radicals react together to form a non-radical

   2. Cl radicals can still react with chloromethane to produce dichloromethane… etc.

Nucleophile

Electron rich species with min. 1 lone pair (attracted by +) which can create a coordination covalent bond

e.g. OH, water, ammonia

Nucleophilic substitution reactions

Heterolytic fission

(R-X) → R + X(-)

R + X(-) + Nu -> (R-Nu) + X(-)

Electrophiles

"Electron poor" species which can accept a coordinate bond

- Neutral but molecules or positive charge

Examples of electrophiles

H⁺, Cl⁺, NO²⁺, and CH³⁺

Hydrogen halides

Electrophilic addition reactions of alkenes

Unsaturated → Saturated

Leb test for presence of alkene

Add bromine, solution goes from yellow-ish to clear due to the electrohilic addition reaction

Hydration

Alkene + Water to form Alcohol

electrophilic addition

Double bond breaks to form 2 new bonds

1. halogens and alkenes

2. hydrogenhalades and alkenes

3. hydration (addition of water)

Corrosive