S2

S2.1: The Ionic Model

• giving and taking of electrons to form electrostatic attraction

• positive ion → cation, negative ion → anion

Transition Metals

• When the ionization energies for the 1st, 2nd, 3rd, etc. are close to each other, the element will be able to lose those valence electrons easily

  • results in variable oxidation states depending on the number of valence electrons in the d and s shells

• 

Compound Names

• polyatomic ions

  • nitrate: NO3-

  • sulfate: SO4²-

  • phosphate: PO4³-

  • hydroxide: OH-

  • hydrocarbonate: HCO3-

  • carbonate: CO3²-

  • ammonium: NH4+

Ionic structures and properties

• giant ionic lattice structures

  • made up of very strong electrostatic forces of attraction

    • ionic bonds or lattice enthalpy

• physical properties depend on lattice structures

Lattice Enthalpy (LE)

• energy needed to seperate into constituent ions (hence △H>0)

• △H is a negative enthalpy change

• LE values are a function of the ionic radii and charge

• △H = (Knm)/(Rm+ + Rx-)

  • K - constant

  • n, m - magnitude of charges

  • Rm+, Rx- - ionic radii

• increase in ionic charge = increase in ionic attraction between ions = increase in lattice enthalpy

• increase in ionic radius of on of the ions = decreased attraction between ions = decreased lattice enthalpy

• lattice enthalpy is greater for ions with a larger charge density as they have a small radius and are highly charged

Properties of ionic compounds

• high melting + boiling point

  • due to high LE

• generally soluble in water (polar) but not in non-polar liquids

• good electrical conductivity in liquid/aqueous states

• generally brittle

  • due to crystalline structure

• low volatitilty

  • volatility - tendency of a substance to vaporize

Ionic character and electronegativity

• ionic character is determined by the formula:

  • %ionic character = △Xp/3.2

    • △Xp=electronegativity difference

• bonding continuum

  • Ionic → △Xp>1.8

  • Polar Covalent → 0<△Xp<1.8

  • Covalent → △Xp<1.8

S2.2: The Covalent Model

• atoms sharing electrons

Octet Rule

• tendency of atoms to gain a valence shell consisting of 8 electrons

  • pairs of electrons not involved in the bond are called lone pairs

• ability of two atoms to form a covalent bond is due to similar strength with which they attract valence electrons

• exceptions to the rule

  • applies to small atoms with less than 8 electrons

    • forms an incomplete octet

    • ex: BeCl2, BF3

Bond Strength

• bond length → distance between 2 bonded nuclei

• bond strength = bond enthalpy → energy required to break the bond

  • as bond length increases, bond enthalpy decreases

Coordination Bond

• bond that is formed by both electrons in pair originating from the same atom

  • ex: H3O+

Valence Shell Electron Pair Repulsion (VSEPR)

• because electron pairs in the same valence shell carry the same charge, they repel each other and so spread themselves as far as possible

• electron pair = electron domain

  • all electron locations in valence shell, all single, double, triple = 1 electron domain

• repulsion applies to electron domains (can be single/double/triple or non-bonding pairs)

• total number of electron domains around central atom determines geometrical arrangement of electron domains

• shape of molecules is determined by angles between bonded atoms

• lone pairs and multiple bonds cause slightly more repulsion

  • lone pairs have a higher concentration of charge (not shared)

  • multiple bonds have a higher concentration of charge (more electrons)

  • non-bonding/lone pair > multiple bond > single bond

Structure

• two electron domains

  • linear shape (bond angle of 180°)

• three electron domains

  • triangular planar (bond angle of 120°) → all single bonds

  • bent/V-shaped (120°, 121°, 118°) → one double bond

  • bent/V-shaped (117° bond angle) → one double bond, one lone pair

• four electron domains

  • tetrahedral (109.5°) → all single bonds

  • trigonal pyramid (107°) → one lone pair

  • bent/V-shaped (104.5°) → two lone pairs

Bond Polarity

• polar bonds - differing electronegativities

  • different pulling strength for electrons

• the more electronegative atom exerts a greater pulling power on the shared electrons → gains more “possesion”

• bond dipole - two partially seperated opposite electric charges

  • the more electronegative atom becomes partially negative

  • the less electronegative atom becomes partially positive

  • increased electronegativity difference = increased bond polarity

• pure covalent bonds → zero electronegativity difference

• polar bonds introduce some ionic nature to the covalent bonds

• Pure Covalent

  • equal sharing of electrons

  • discrete molecules

• Polar Covalent

  • partial/unequal sharing/transfer of electrons

• Ionic

  • complete transfer of electrons

  • lattice of oppositely charged ions

Molecular Polarity

• depends on polar bonds it contains

• depends on molecular geometry

• non-polar

  • dipoles can cancel out, creating non-polar molecules

• polar

  • if the molecule contains bonds of different polarity or the bonds are not symmetrically arranged, dipoles will not cancel out

  • creates a net dipole (turning force in electric field)

• IR active

  • happens only when an overall dipole moment related to the position and vibration of its bonds is found

Covalent Network Structures

• discrete molecules → finite amount of atoms

• covalent networks → no finite number of atoms

  • single repeating pattern of covalent bonds

• allotropes → different bonding/structural patterns of the same element in the same physical state, causing different chemical and physical properties

Allotropes of Carbon

• Diamond

  • structure: sp3 hybridied and covalently bonded to 4 others tetrahedrally arranged in a regular repetitive pattern (angle → 109.5°)

  • non-conductor of electicity, all electrons are bonded

  • very efficient thermal conducter, better than metals

  • highly transparent, lustrous crystal

  • hardest natural substance, brittle, very high melting point

  • uses: polished for jewelry, tools and machinery for grinding/cutting glass

• Graphite

  • structure: sp2 hybridized and covalently bonded to 3 others, forming hexagons in parallel angles with bond angles of 120° (weak London dispersion forces)

  • good electrical conductor; contains one delocalized electron per atom

  • not a good thermal conductor, unless the heat conducts parallel to crystal layers

  • non-lustrous, grey crystalline solid

  • soft and slippery due to sliding layers, brittle, very high melting point, stable mostly

  • uses: dry lubricant, pencils, electrode rods in electrolysis

• Graphene

  • sp2 hybridized and covalently bonded to 3 others, forming hexagons with bond angles of 120° (single layer → 2D only) honeycomb/chicken wire

  • very good electrical conductor, one delocalized electron per atom

  • best thermal conductivity known

  • almost completely transperent

  • thickness of just one atom (2D) → thinnest material to ever exist, 100x stronger than steel (strongest), very flexible, very high melting point

  • uses: transmission electron microscopy (TEM) grids, photovoltaic cells, touchscreens, high performance electronic devices, etc.

• Fullerene (C60)

  • sp2 hybridized, bonded in a sphere of 60 carbon atoms, consisting of 12 pentagons and 20 hexagons (closed spherical cage)

  • poor conductors of electricity, delocalized electron has little movement

  • very low thermal conductivity

  • black powder

  • very light and strong, reacts with potassium (K) to make superconducting crystalline material, low melting point

  • uses: lubricants, medical, industrial devices for binding specific target molecules; reltaed forms are used to make nanotubes/nanobuds used as capacitators in the electronics indsutry, and catalysts

London Dispersion Forces

• non-polar molecules do not have a permanent dipole

  • electrons behave somewhat like clouds of negative charge, density of the cloud could be greater over one atom at any moment

  • when there is a differing density, the bond will have seperation of charge, creating a weak dipole (temporary/instantaneous dipole)

• creates weak forces of attraction that occur between opposite ends of two temporary dipoles

  • weakest form of intermolecular force

  • strength increases with molecular size (greater number of electrons)

• LDF is the only force that exists for non-polar molecules

• also exists in polar molecules, but is often overlooked for stronger forces

Dipole-dipole attraction

• polar molecules have permanent seperation of charge (electronegativity difference)

  • known as a permanent dipole

  • opposite charges on neighbouring molecules attracting each other

• strength varies on distance and relative orientation of the dipoles

Dipole-induced dipole attraction

• occurs in mixtures with both polar and non-polar molecules

• the permanent dipole from a polar molecule creates a temporary seperation of charge in the non-polar

• act in addition to LDF (non-polar) and dipole-dipole attraction (polar)

• van der Waal’s force: all 3 forces added together

Hydrogen Bonding

• when a molecule contains hydrogen covalently bonded to fluorine, oxygen, or nitrogen (electronegative atoms)

  • particular case of dipole-dipole attraction

• the large electronegativity difference between hydrogen and the bonded atoms results in the electron pair being pulled away from hydrogen

• hydrogen now exerts a strong attractive force on a lone pair in the electronegative atom due to its small size and the absence of other electrons to shield the nucleus

• strongest form of intermolecular force

Melting and Boiling Point

• changing state = breaking intermolecular forces

• covalent substances generally have lower MP and BP than ionic substances

  • relatively weak intermolecular forces < electrostatic attraction

  • covalent substances are generally liquid/gas at room temperature

• strength of intermolecular forces increase with molecular size and extent of polarity

Solubility

• non-polar substances are generally dissolvable in non-polar solvents by formation of LDF between solute and solvent

• polar covalent compounds can generally dissolve in water (highly polar H2O) through dipole interactions and hydrogen bonding

• solubility of polar compounds is reduced in larger molecules

  • polar bonds only a small part of the structure

  • non-polar parts reduce solubility

• inability of non-polar groups to associate with water means non-polar substances do not dissolve well in water

• polar substances have low solubility in non-polar solvents

  • they remain together due to dipole-dipole attractions

• giant molecular are generally insoluble in all solvents

  • too much energy required to break the strong covalent bonds

Electrical Conductivity

• covalent compounds do not contain ions, so they cannot conduct electricity in the solid or liquid state

  • some polar covalent molecules (when they can ionize) will conduct electricity

Resonance Structures

• delocalization - tendencty to be shared between more than bonding position

  • delocalized electrons spread out → greater stability for molecule/ion

• delocalization occurs when there is more than one position for a double bond within a molecule

  • two equally valid positions for a double bond

  • ex: expected is 1 single and 1 double bond, in reality is is 2 equal bonds, intermediate in length and strength

• resonance - molecule is a combination of two Lewis formulas

  • electrons from the double bond delocalize and spread themselves equally between both possible bonding positions

    • shown with a dotted line

  • known as a resonance hybrid

• resonance influences bond strengths/lengths which in turn can influence reactivity

Benzene (C6H6)

• six carbon atoms arranged in a hexagonal ring, each bonded to a hydrogen atom in a triangular planar arrangement with bond angle 120°

• true form of benzene is the resonance hybrid

  • circle represents equally spread delocalized electrons

• 1-1 ratio of carbon to hydrogen indicates a high degree of unsaturation, greater than that of alkenes or alkynes

  • does not show characteristic properties

  • no isomers, reluctant to undergo additional reactions

Expanded Octet

• when the central atom is period 3 or lower, sometimes there are more than 8 electrons around the central atom

• d orbitals available in the valence shell have energy values similar to those of the p orbitals

  • promotion of electrons (3p→3d) allows additional electron pairs to form

• causes some elements to expand their octets (5-6 electron domains)

Species with five electron domains

• triangular bipyramidal shape → angles of 90° and 120°

• if one or more domains are non-bonding electrons, they will repel the most

  • one lone pair gives an unsymmetrical tetrahedron or see-saw shape (bond angles <120° and <90°)

• to minimize additional repulsion and bonding domains, the lone pair must be located in an equatorial position (horizontal plane around central atom) instead of an axial osition (above/below horizontal plane)

  • two lone pairs give a T-shaped structure (bond angles <90°)

  • three lone pairs give a linear shape (bond angle 180°)

Species with six electron domains

• octahedral shape with angles of 90°

  • no lone pairs of electrons → symmetrical octahedral shape

  • one lone pair → square pyramidal shape (bond angles slightly less than 90°)

  • two lone pairs → square planar shape (bond angles of 90°)

    • maximizes distance apart by arranging pairs on opposite sides

Molecular Polarity

Formal Charge

• formal charge used to predict a preferred Lewis formula

• treats covalent bonds as if they were purely covalent with equal electron distribution

• FC = V - (1/2 B + L)

  • V = valence, B = bonding, L = lone (number of electrons)

• low FC means less charge transfer has taken place in forming a structure from its atoms

  • generally means most stable → preferred structure

• sum of formal charges for a species must be equal to the charge

Sigma Bond

• when two atomic orbitals combine head-on along the bond axis (imaginary line)

  • overlap of s, p, and hybrid orbitals in different combinations

  • always the bond that forms in a single covalent bond

• electron density is concentrated between the nuclei of the bonded atoms

Pi Bond

• when two p orbitals collide laterally (sideways-on)

• electron density is concentrated above and below the plane of the bond axis

• only forms within double and triple bonds

• weaker than sigma bonds as electron density is further from nucleus

Sigma and Pi Bonds

• s+s → sigma

• s+p → sigma

• p+p (head-on) → sigma

• hybrid + s → sigma

• hybrid + hybrid → sigma

• p+p (laterally) → pi

Hybridization

• formation of covalent bonds often starts with excitation of the atoms

  • amount of energy put in to achieve this is more than compensated by the extra energy released on forming bonds

• if different orbitals are used in forming covalent bonds, unequal bonds are expected

  • instead, unequal atomic orbitals within an atom mix to form new hybrid atomic orbitals which are identical but different from the original bonds

• hybrid orbirtals have different energies, shapes, and orientation in space from their parent orbitals

  • allows them to form stronger bonds by allowing for greater overlap

• sp³ orbitals

  • 1 s orbital and 2 p orbitals produce 4 sp³ orbitals

  • shape and energy have properties of s and p, but more like p than s

• sp² orbitals

  • 1 s orbital and 2p orbitals produce 3 sp² orbitals

• sp orbitals

  • 1 s orbital and 1 p orbital produce 2 sp orbitals

Carbon → Hybridization

• C: atomic nyumber = 6 (1s²2s²2px12py1)

  • forms 4 covalent bonds, but only has two singly occupied bonding electrons

• excitation occurs (2s → 2p) to change from ground state

• sp³ hybridization

  • orbitals orientate themselves at 109.5°, forming a tetrahedron

  • each hybrid orbital overlaps with an atomic orbital → 4 sigma bonds

• sp² hybridization

  • when carbon forms a double bond

  • orientate themselves at 120°, forming a triangular planar

  • each hybrid orbital overlaps with a neigbouring atomic orbital → 3 sigma bonds

  • as the 2 carbon atoms appproach each other, the p orbitals in each atom that did not hybridize overlap sideways

    • forms a pi bond

    • double bond (C2) → 1 sigma, 1 pi

    • characteristic lobes of electron density above and below the bond axis

• sp hybridization

  • orientate themselves at 180°, giving a linear shape

  • overlap of the two hybrid orbitals with other atomic orbitals → 2 sigma bonds

  • when carbon forms a triple bond

  • C2H2

    • each carbon atom has 2 unhybridized p orbitals that are orientated 90° to each other

      • combines to form 2 pi bonds

    • four lobes of electron density turns into a cylinder of negative charge around the atom, making the molecule susceptible to electrophilic reactants (attracted to electron-dense regions)

Hybridization and Molecular Geometry

• tetrahedral → sp³

• triangular planar → sp²

• linear → sp

• lone pairs can also be used in hybridization

  • non-bonding pairs can also hybridize

    • ex: NH3 → lone pair in N resides in the sp³ orbital

Hybridization and Benzene (C6H6)

• each of the 6 carbon atoms are sp² hybridized

  • forms 3 sigma bonds (120°) → planar shape

  • leaves the unhybridized p electron on each carbon atom

    • dumbbell shape perpendicular to the plane of the ring

    • do not form pi bonds but effectively overlap in both directions

    • spreads themselves evenly to be shared by all 6 carbon atoms

  • forms a delocalized pi electron cloud

    • electron density is concentrated in 2 donut-shaped rings above and below the plane

2.3: The Metallic Model

Metallic Bonding

• metals: low ionization energies so they react by losing valence electrons forming a positive ion

  • metallic character: loss of control over outer shell electrons

• when there is no other element present to accept the electrons and form an ionic compound, the outer electrons are held loosely by the nucleus so they ‘wander off’

  • delocalized electrons

• metal atoms form a regular lattice structure through which electrons move freely

  • metallic bonding: force of electrostatic attraction between lattice of cations and delocalized electrons

Uses of metals

• Good electrical conductivity

  • because of highly mobile delocalized electrons

    • used for electrical circuits (copper)

• Good thermal conductivity

  • because of delocalized electrons and closely packed ions

    • used for pots and pans for cooking

• Malleable (can be shaped under pressure)

  • because of the lack of direction in the movement of delocalized electrons

    • used for machinery

• Ductile (can be drawn out into threads)

  • because the metallic bond remaining intact while formation changes

    • used for electric wires and cables

• High melting points

  • because of strong electrostatic forces

    • used for high-speed tools

• Shiny, lustrous appearance

  • because delocalized electrons in metal crystal structure reflect light

    • used for jewelry

• non-directional nature of metallic bonding allows metals to mix with other metals or non-metals in the molten state

  • resulting mixture is an alloy

    • enhances properties of the metallic structure

Metallic bond strength

• determined by

  • number of delocalized electrons

  • charge of the cation

  • radius of the cation

• the greater the electron density and the smaller the cation, the greater the electrostatic attraction

• Across a period

  • increasing melting point

    • greater attraction between ions and delocalized electrons

  • lower degree of reactivity

• Down a group

  • decreasing melting point

    • weaker attraction between ions and delocalized electrons

  • higher degree of reactivity

Transition elements

• elements with an incomplete d-sublevel OR elements that can give rise to cations with an incomplete d-sublevel

• proximity in energy between outer occupied sublevels enables them to delocalize large amounts of d-electrons to form metallic bonds

Transition elements: High melting point

• metals have a large amount of delocalized electrons and a large positive charge on the metal cations which leads to strong metallic bonding → high melting points

• transition metal trends are less evident due to ability of transition elements to delocalize large numbers of electrons and the similar ionic radii

• difficult to predict trends accurately compared to the s-block metals

Transition elements: High electrical conductivity

• metals have a large amount of delocalized electrons which increases their conductivity

  • example: copper (Cu) is used in wires

2.4: Models to Materials

Bonding Triangle

• bonding seen as a continuum (ionic, covalent, metallic) → different bonding types are present to different degrees

  • position on triangle determined from electronegativity

• high electronegativity difference = ionic

• low electronegativity difference = covalent or metallic

• intermediate electronegativity difference = polar covalent

Composite Materials

• mixture between two or more different materials

  • materials have seperate phases (different positions on bonding triangle)

• mixture retains properties of individual materials that compose it

  • example: fibreglass, concrete

Alloys

• produced by adding one metal element to another metal (or carbon) in liquid/molten state so the different atoms can mix

  • in solid, ions of the different metals are scattered through the lattice

    • forms a structure of uniform composition

• metallic bonds are present → delocalized electrons bind the lattice

  • possible due to the non-directional nature of the delocalized electrons and accomodation of the lattice to different sizes of ions

• alloys have properties distinct of component elements (different packing of cations in lattice)

• pure metal → regular arrangment of atoms

  • interrupted in an alloy by different cations

    • more difficult for atoms to ‘slip over each other’ → stronger

• alloy is sronger, more chemically stable, and more resistant to corrosion

Polymers

• monomers (small molecules) are able to react together to form a linked chain held together by covalent bonds, forming a polymer

• polymers are macromolecules → composed of thousands of atoms and so are relatively large compared with other molecules

• nature/properties of a polymer vary with the monomer, length, and amount of branching

• structure is shown as a repeating unit with open bonds on each end

• natural polymers - found naturally (example: protein, starch, DNA)

• synthetic polymers - human-made and non-biodegradable (example: plastics)

Addition Polymers

• addition reaction - a multiple bond in a molecule breaks and creates new bonding positions

  • alkenes/alkynes have double/triple carbon-carbon bonds respectively so they readily undergo addition reactions

    • they can act as monomers and form addition polymers

• %atom economy = molar mass of desired product / molar mass of all reactants x 100

  • addition polymerization reactions do not generate a by-product so it has a 100% atom economy

Condensation Polymers

• condensation reaction - two functional groups react to form a new covalent bond with the release of a small molecule (H2O, HCl, NH3, etc.)

  • A-OH + H-B → A-B + H2O

• to form condensation polymers, monomers must have functional groups (active ends)

  • allows them to form new covalent bonds with neighbours on both sides

  • the functional groups in neighbouring molecules must be able to react together

Carboxylic acid + alcohol → polyester (ester link)

• when one monomer has two carboxylic acid groups (COOH) and the other has two alcohol groups (OH), an ester link forms between them

  • chain extends in both directions → polyester

Carboxylic acid + amine → polyamide (amide link)

• when one monomer has two carboxylic acid groups (COOH) and another monomer has two amine groups (NH2), an amide link forms between them

  • forms a polymer known as polyamide