IB HL Chemistry (first exams 2025) - Structure 2.1, 2.2, 2.3, 2.4

2.1 Ionic Model 2.2 Covalent Model 2.3 Metallic Model 2.4 From Models to Materials

ions

ions

charged particle

cation

positive ion formed by metals when they lose one or more electrons

anion

negative ion formed by non-metals when they gain one or more electron

writing ionic charge

number before sign (1+)

transition elements

elements that have electron configurations which allow them to lose different amounts of electrons, so can form more than one stable ion

polyatomic ions

a covalently bonded set of two or more atoms, or of a metal complex can be considered to behave as a single unit

NH4+

ammonium

NO3-

nitrate

OH-

hydroxide

HCO3-

hydrogen carbonate

CO32-

carbonate

SO42+

sulfate

PO43-

phosphate

ionic compounds

electrostatic attractions between a metal and non-metal

ionic compound properties

high brittleness due to ions of like charge near each other, high melting/boiling points due to strong electrostatic attractions, low volatility due to strong electrostatic attractions, molten/dissolved conduct electricity

lattice enthalpy

the energy needed to convert one mole of ionic solid into gaseous ions infinitely far apart under standard conditions (breaking apart), endothermic

size of lattice enthalpy

controlled by the charges on the ions, their ionic radii, and the packing arrangement of the ions

covalent bonds

two non-metals reacting together to achieve a stable configuration, where a shared pair of electrons is concentrated in the region between the two positively charged nuclei

octet rule

the tendency of covalent bonds to form a stable arrangement of eight electrons in the outer shell

octet rule exceptions

hydrogen only has two, Be and B have less than eight, S and P can have expanded octets (more than eight)

coordinate covalent bonds

form when both electrons in a shared pair originate from the same atom (donated to the other)

single bonds

share one pair of electrons, weakest and longest bond

bond enthalpy

a measure of the energy required to break the bond

double bonds

stronger and shorter than single bonds, longer and weaker than triple bonds

triple bonds

strongest and shortest bond

dipole

a pair of separated equal and opposite electrical charges located on a pair of atoms within a molecule

dipole moment

the product of the charge on a dipole and the distance between the ends; a measure of the polarity of a bond

non-polar molecule

a molecule that has a symmetric distribution of charge and whose individual bond dipoles sum to zero or cancel

polar molecule

a molecule that has an asymmetric distribution of charge: the individual bond dipoles do not sum to zero or cancel

0 - 0.4

pauling scale non-polar covalent

0.4 - 1.8

pauling scale polar covalent

> 1.8

pauling scale ionic

molecular polarity

depends on the polar bonds a molecule contains, and the shape

valence shell electron pair repulsion model

VSEPR model

VSEPR

states that in a small molecule, the pairs of valence electrons are arranged as far apart from each other possible (lone pairs and bonded pairs)

lone pair - lone pair

greatest VSEPR repulsion

lone pair - bonded pair

median VSEPR repulsion

bond pair - bond pair

weakest VSEPR repulsion

linear bonds

two atoms bonded to the central atom, no lone pairs on the central atom, 180 degree bond angle

bent

two atoms bonded to the central atom with one or two lone pairs of electrons on the central atom, 104.5 degrees bond angle

trigonal planar

three atoms bonded to the central atom, no lone pairs of electrons on the central atom, 120 degrees bond angle

trigonal pyramidal

the atoms bonded to the central atom, one lone pair of electrons on the central atom, 107 degrees bond angle

tetrahedral

four atoms bonded to the central atom, no lone pairs of electrons on the central atom, 109.5 degrees bond angle

linear - linear

180, 2 bonding pairs, 0 lone pairs

trigonal planar - trigonal planar

120, 3 bonding pairs, 0 lone pairs

trigonal planar - bent linear

118, 2 bonding pairs, 1 lone pair

tetrahedral - tetrahedral

109.5, 4 bonding pairs, 0 lone pairs

tetrahedral - trigonal pyramid

107, 3 bonding pairs, 1 lone pair

tetrahedral - bent linear

104.5, 2 bonding pairs, 2 lone pairs

octahedral

6 electron domains, angles of 90 degrees and d2sp3 hybridisation

square pyramidal

derivative of the octahedral with one lone pair of electrons on the top or bottom, 90 degrees bond angle

square planar

derivative of the octahedral shape with two lone pairs of electrons, one on top and one on the bottom, 90 degrees bond angle

intermolecular forces

forces between molecules, london/van der waals/dispersion forces, dipole-dipole forces, hydrogen bonding

intramolecular forces

forces within the molecule such as covalent, ionic, and metallic bonding

london/van der waals/dispersion forces

IMF in all covalents, only IMF in non-polar molecules

dipole - dipole forces

IMF only present in polar molecules with permanently charged regions, stronger than dispersion forces, and strength depends on degree of polarity, caused when molecules with permanent dipoles attract each other

hydrogen bonding

strongest IMF, type of dipole - dipole attraction, form when hydrogen bonds to oxygen, nitrogen, or fluorine, strength due to hydrogen small size and large electronegativity of N, O, and Fj

like dissolves like

polar/ionic solvents dissolve polar/ionic solutes, non-polar solvents dissolve non-polar solutes

higher

stronger IMF, ________ boiling point

lower

stronger IMF, ________ volatility

dispersion > dipole-dipole > hydrogen bonding

order of volatility in covalent IMFs

dispersion < dipole-dipole < hydrogen bonding

order of melting/boiling point in covalent IMFs

covalents electrical conductivity

generally do not conduct, some giant covalents, some polar covalents

graphite, graphene, diamond, fullerene

allotropes of carbon

diamond

each C atom is sp3 hybridised covalently bonded to 4 others tetrahedrally in a repeating pattern with bond angles of 109.5, density is 3.51 g/cm3, all electrons bonded so no conductivity, high melting point, brittle

graphene

c atom is sp3 hybridized and covalently bonded to three other carbons forming hexagons with bond angles 120, two dimensional single layer in a hexagonal pattern, density 1.5g/cm3, one delocalised electron per atom so conducts electricity, excellent heat conductor, high melting point

graphite

parallel layers of graphene, held together by van der waals (dispersion) forces so they can slide over each other, density is 2.26g/cm3, not a good heat conductor, high melting point, most stable allotrope

fullerene

c atom is sp2 hybridised and covalently bonded in a sphere, closed spherical cage, density is 1.726g/cm3, easily accepts electrons so is a semiconductor, low heat conductivity, low melting point, soluble in benzene

Silicon

group four element with four valence electrons, each Si atom is covalent bonded to four others in a tetrahedral arrangement, results in giant lattice structure

silica/quartz (SiO2)

tetrahedral structure with bonds between Si and O, each Si bonded to 4 oxygen atoms and each O bonded to two Si, strong, insoluble in water, does not conduct electricity or heat, high melting point

formal charge = (no. valence electrons) - ½ (no. bonding electrons) - (no. non-bonding electrons)

formal charge formula extended

FC = V - 1/2B - N

formal charge formula shortened

formal charge

the charge assigned to an atom in a molecule

preferred lewis structure

where the difference on formal charge is closest to zero, the negative charges located on the most electronegative atomd

resonance

a concept used to describe the structures when there are multiple ways to depict the same molecule

benzene C6H6

colourless compound, physical state is liquid with slight aromatic smell, less dense than H2O, immiscible with H2O, miscible with non-polar solvents

ozone O3

bent shape with bond angle of 117 degrees, two resonance structures, double bond with one pi and one sigma bond, bond order is 1.5 which means the length is intermediate and the strength is between a double and single bond

pi bonds

form by the lateral combination of p-orbitals, where the electron density is concentrated on opposite sides of the bond axis (so weaker than sigma), found in double and triple bonds

sigma bonds

strongest type of covalent chemical bond, formed by head on overlapping between atomic orbitals along the bond axis, all single covalent bonds, formed by s/s, s/p, p/p, hybrid/s, hybrid/hybrid

hybridisation

model that describes the changes on the atomic orbitals of an atom when it forms a covalent compound (sp, sp2, sp3)

reasons for hybridisation

VSEPR theory requires molecules to have identical orbitals at the stage of bonding, so the s and p orbitals hybridise to be identical

hybrid orbitals

only found in covalent compounds, equivalent in a compound, no. is equal to the no. of atomic orbitals used to form it, type depends on electron domain geometry

sp3 hybridisation

occurs when three p orbitals and one s orbital hybridise to form four sigma bonds, tetrahedral and has bond angles of 109.5 degrees

sp2 hybridisation

three p orbitals and one s orbital hybridise to form three hybrids and one unhybridised p orbital (which overlaps forming a pi bond), trigonal planar with bond angles of 120 degrees

sp hybridisation

three p orbitals and one s orbital hybridise to form two hybrids and two unhybridised p orbitals (which overlap sideways forming two pi bonds), linear with bond angles 180 degrees

metallic bonds

electrostatic attractions between a lattice of cations and a sea of delocalised electrons

strength of metallic bonding

depends on size (greater positive charge = stronger bond), radius of metal ion (decrease size = stronger bond), no. mobile electrons (more mobile electrons = stronger)

metallic bonding strength/length

delocalised electrons give rise to intermediate bond strength and length

high

if stress is applied to a metal, planes of atoms slide over each other because of delocalised electrons, so metals have ______ malleability

high

when heat is applied to a metal, the kinetic energy of the electrons increases, they move through to cold regions of the lattice, so metals have _______ thermal conductivity

high

if a voltage is applied across the ends of a metal sample, the delocalised electrons flow towards the positive electrode, so metals have _______ electrical conductivity

high

lots of energy required to break metallic bonds, influenced by size of positive charge on ion, radius of ion, no. mobile electrons, so metals have ______ melting points

high melting point

because of the greater number of electrons from the d-sublevel being involved in metallic bonding in addition to the e-electrons, transition metals have stronger metallic bonding, which leads to a _________

good conductors

because of their loosely bound valence electrons, partially filled d orbitals, and closely packed atomic structure, transition metals are ________

alloys

mixtures of a metal/metal or metal/non-metals (solid solutions), occurs when they are mixed together in a molten state and solidify with ions of of the different metals scattered throughout the lattice

substitutional alloy

atoms of one metal are substitutes by atoms of another metal

interstitial alloy

different metal occupies interstitial spaces in the lattice structure

common alloys

steel, stainless steel, brass, bronze, pewter, sterling silver

van arkel - ketelaar triangle

x = (z1 + z2)/2

average electronegativity formula