Atomic Theory
Particle Model or Kinetic Theory of Matter
States that all matter is made up of atoms.
𝑐𝑟𝑦𝑠𝑡𝑎𝑙𝑙𝑖𝑛𝑒 → 𝑟𝑒𝑔𝑢𝑙𝑎𝑟 𝑠𝑡𝑟𝑢𝑐𝑡𝑢𝑟𝑒
𝑠𝑜𝑙𝑖𝑑𝑠 {
𝑎𝑚𝑜𝑟𝑝ℎ𝑜𝑢𝑠 → 𝑖𝑟𝑟𝑒𝑔𝑢𝑙𝑎𝑟 𝑠𝑡𝑟𝑢𝑐𝑡𝑢𝑟𝑒
Atomic Properties
Ionization Energy – energy required to remove an
𝑘𝐽
electron. Increases LR, decreases UD. Measured in
𝑚𝑜𝑙
Electronegativity – how much an element attracts extra
electrons to itself. Increases LR, decreases UD.
Atomic Radius – decreases LR, increases UD
Atoms
Electrons
Atomic Shells:
K(1), L(2), M(3), N(4)
Electrons orbit the nucleus in discrete energy levels. There
Max electrons in shell n: 2𝑛2 , max 8 electrons in outer
are only several energy levels an electron may have - it
shell for 𝑍 ≤ 20
may not have any values in between.
Max subshells in shell n: 𝑛
Max orbitals in shell n: 𝑛2 (2 electrons per orbital)
Isotopic Symbol of Any Element
Number of subshells = shell number.
𝐴
𝑍𝑋
A = Atomic mass (no. protons + no. neutrons)
Z = Atomic number (no. protons)
X = Element symbol
Atomic Subshells:
Subshell
S
P
D
Electron Capacity
2
6
10
Molecule – formed when a group of atoms covalently bond
Compound – pure substance composed of different
elements in chemical bonds in fixed proportions by mass. Higher S subshell must fill before lower D subshell (4s
Mixture – formed when two or more pure substances are before 3d) unless D subshell can be half or fully filled. In
mixed without reacting
these cases, D subshell fills first, and what remains goes to
S subshell (in this case, 3d before 4s).
Electron Configuration
Describes the arrangement of electrons in energy levels
Ar = relative atomic mass
(shells)
Σ(𝑅𝐼𝑀 × %𝐴𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒)
𝐴𝑟 =
100
Bohr’s theory states that there are a maximum of two
Relative atomic mass is the average weight per atom of a
electrons in the first shell, eight in the second and 18 in the certain element. Takes isotopes into account.
third. However, for the first 20 elements the outer shell
cannot have more than eight electrons.
Orbital Filling Order
Electrons fill from lowest energy state.
When electron absorbs sufficient energy, it jumps to a
higher energy level and becomes unstable. Eventually falls
back to lower energy level by emitting energy as light
(photon). Energy of photon (frequency) is equal to the
difference in energy of two levels. Basis of spectroscopy
Periodic Table
Rows
Columns
Periods
Groups
Elements in the same group have the same valence ∴
similar properties.
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d …
Percentage Composition
Percentage by mass of each element in compound
mass of x in 1 mole of compound
%X =
∗ 100%
molar mass of compound
Hydrated Ionic Compounds
Contain water molecules as part of lattice structure (water
of crystallization). Mass difference between hydrated and
anhydrous compound is known  can find % composition
of water in compound.
Group
General name
1
Alkali Metals
2
Alkaline Earth Metals
Numerical Connections
7
Halogens
1 𝑚𝑜𝑙 = 6.02 × 1023 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒𝑠
𝑔
8
Noble/Inert gases
1 𝑎𝑚𝑢 ≡ 1
Period number = number of shells in atom occupied by
𝑚𝑜𝑙
𝑔
Molar mass = mass in g of one mole of substance.
electrons
𝑚𝑜𝑙
′
𝑛𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒𝑠 = 𝑛𝑚𝑜𝑙𝑒𝑠 × 𝐴𝑣𝑜𝑔𝑎𝑑𝑟𝑜 𝑠 𝑁𝑢𝑚𝑏𝑒𝑟
S2O3 2HPO4 2PO4 3-
Metallic Bonding
When metallic atoms combine, metallic lattice is formed
consisting of cations in sea of delocalized electrons.
Metallic lattice structure influences properties of metals.
Metal atoms have low electronegativity
Ball Bearing Model
Metallic atoms form small areas of perfect close packing
called grains. Grains may be irregularly packed.
Large grains  fewer dislocations  more malleable
Small grains  more dislocations  do not bend easily
Work hardening  creation of more grains by bending
Group 14 and Group 18 elements do not form ions readily Heating  creation of large grains
Alloying  Adding other atoms to lattice to make it less
Ionic compounds are conductive only in liquid or aqueous easy to ‘slide’
state. Very brittle, soluble in water, usually crystalline,
Heat modified metals
high melting/boiling points.
Annealed – heated, cooled slowly  larger crystals, softer
Quenched – heated, cooled quickly  smaller crystals
Naming Binary Ionic Compounds
Tempered – quenched then reheated, cooled slowly
Cation first, anion second + ide suffix
(reduces brittleness, retains hardness)
Ionic Bonding
Metals lose electrons to gain ideal configuration
Non-metals gain electrons to obtain ideal configurations
Gain/loss electron depends on electronegativity
For hydrates, use mono, di, tri, tetra, penta, hexa, hepta,
octa, nona, deca prefixes to indicate number of water
molecules in compound
Atomic Size Cations smaller than normal, anions larger
Ionic Formulae Empirical – ratios only. Not ‘molecules’
Polyatomic ions are atoms covalently bonded to form an
anion
Ionic compound sometimes called a salt
Warning: When writing the electron config of an ion,
remember that all the adding/taking of electrons occurs in Precipitation Reactions
Two aqueous ionic compounds react together to form
outer shell. I.e. 𝐶𝑜 → [𝐴𝑟]4𝑠 2 3𝑑 7
insoluble compound. Insoluble compound  precipitate.
But 𝐶𝑜 2+ → [𝐴𝑟]3𝑑 7 NOT [𝐴𝑟]4𝑠 2 3𝑑 5 4s outermost!!!
Chemistry – History Timeline
1789 Antoine Lavoisier – First periodic table
1817 Johann Dobereiner – Law of Triads
1857 William Odling – Classification of elements in 2 dim.
1863 John Newlands – Law of Octaves
1866 Lothar Meyer – Periodic variation of properties
1871 Dmitri Mendeleev – First modern form per. table
1894 John Rayleigh/William Ramsay – Noble gases
1913 Henry Moseley – Atomic Number
1945-61 Glenn Seaborg – Element synthesis
Thiosulfate
Hydrogen Phosphate
Phosphate
Ammonium
Hydroxide
Nitrate
Hydrogen Carbonate
Hydrogen Sulfate
Chlorate
Hydrogen Sulfite
Nitrite
Permanganate
Hypochlorite
Dihydrogen Phosphate
Cyanide
Sulfate
Carbonate
Sulfite
Dichromate
Chromate
NH4+
OHNO3HCO3HSO4ClO3
HSO3NO2MnO4OClH2PO4CNSO4 2CO3 2SO3 2Cr2O7 2CrO4 2-
Alloys: Formed by addition of other substance to metals.
Substitutional Alloy: addition of other atoms about the
same size as the metal atoms
Interstitial Alloy: addition of other atoms with smaller
radius than the metal atoms
Covalent Bonding
Central atom in a covalent molecule is the atom with the
most bonding electrons.
Molecule shapes can be determined using dot diagrams
and VSEPR (Valence Shell Electron Pair Repulsion) Theory.
Linear –if central atom has one bonding pair or
one pair on either side, e.g. H2, CO2
V-Shaped – if central atom has two bonding
pairs and two lone pairs, e.g. H2O
Planar - if central atom has three bonding pairs
and no lone pairs, e.g. C2H4
Pyramidal – if central atom has three bonding
pairs and one lone pair, e.g. NH3
Tetrahedral – if central atom has four bonding
pairs and no lone pairs, e.g. CH4
Naming (Binary Molecular Compounds)
First element in full + Greek prefix to indicate how many of
second element in molecule + Second element + ide suffix
Non-metallic atoms have high electronegativities
Different electronegativities in a covalent bond may result
in a polar covalent bond.
Polar covalent bonds formed if one atom has higher
electronegativity than the other, therefore it ‘hogs’ the
shared electrons.
𝑷𝒐𝒍𝒂𝒓 𝒃𝒐𝒏𝒅𝒔 ≠ 𝑷𝒐𝒍𝒂𝒓 𝒎𝒐𝒍𝒆𝒄𝒖𝒍𝒆
Polar molecule only formed if (a) polar bonds exist (b)
they are not cancelled out in direction.
E.g. NH3 is polar covalent, whereas CCl4 is not.
Intermolecular Forces
Dispersion/Van der Waals/London Forces
Apply to all molecules
Due to attraction of electrons to nuclei of other
atoms
Force magnitude dependent on number of
electrons and molecular size/shape
Dipole-Dipole Interactions
Occur between polar-covalent molecules.
Quite weak compared to dispersion forces, and
occur in addition to them.
Hydrogen Bonding
Special case of dipole-dipole interactions
Occurs when a H is bonded with N, O, F atom
Hydrogen bonds are quite strong.
Examples: 𝐶𝐻4 , 𝐶2 𝐻6 , 𝐶10 𝐻22
1. Determine longest chain of C atoms. If >1, choose one
with most substituent groups.
Homologous series  series of organic compounds in
2. Determine which end is closest to a branch, = bond, or
which each successive member differs by CH2 from previous ≡bond. =/≡ bonds take precedence over branches.
3. Number carbons starting from that end.
one
4. Write down any branches with name + ‘yl’, e.g. ethane 
ethyl, then the longest chain, then any ≡ 𝑜𝑟 = bond.
Reactions of Alkanes
5. In case of 2+ branches on same C atom, indicate # of C
Combustion – react with oxygen to form CO2 and H2O
Substitution – Reaction with Cl or F. One or more Cl/F atoms atom for each branch + names in alphabetical order.
6. In case of 2+ identical branches on different carbon
replace one or more H atoms.
atoms, use di-, tri-, tetra- … prefixes.
Note: CH4 + Cl2 --> CH2Cl2 + H2 is incorrect – only one
substitution per Cl2/F2 molecule!
Addition Polymerisation
Alkenes break C==C bonds to join together and form huge
‘alkane’.
Structural Isomers  Same molecular formula,
different structural formula
Isomers of 𝑪𝟒 𝑯𝟏𝟎 – Butane (First HC with isomers)
Semi Structural Formulae
Carbon-by-carbon  side chains written in brackets after
carbon atom to which they are attached. Repeating units
are also bracketed with an index
𝐶4 𝐻10 → 𝐶𝐻3 (𝐶𝐻2 )2 𝐶𝐻3
Functional Groups
Bond/atom/group of atoms giving a molecule specific
properties. I.e. alkenes  C==C bond, alkanols  –OH
Covalent network – 3D lattice
Covalent layer – 2D layer lattice
Covalent layer lattices (graphite, etc.) conduct
electricity as electrons are free to move between layers
Allotropes: different crystalline or molecular forms of the
same element (i.e. diamond  allotrope of carbon)
Haloalkanes
One or more halogens replaced H atom on alkane.
Haloalkanes≠Hydrocarbons!
Organic Chemistry
Hydrocarbons  Compounds containing only C and H
atoms.
Amines
Have –NH2 functional group. Suffix –amine.
Polymers
Formed from units called monomers
Copolymer – formed from two different monomers
Thermoplastic – will melt/soften when heated
Thermosetting – will char/burn when heated
Covalent Lattices
Organic chemistry is the systematic study of the
compounds of carbon, excluding oxides of carbon,
carbonates, carbides and cyanides. These exceptions plus
the compounds of the remaining elements are considered
to be inorganic.
Reactions
React with alkanols to form esters. Concentrated sulfuric
acid used as catalyst.
Naming Esters
Alkanol name (from which ester is formed) given first,
followed by acid part. –oic ending of acid replaced
with -oate.
Methanol + Ethanoic Acid  Methyl Ethanoate
Alkenes
Contain one double bond in C chain.
Have general formula 𝐶𝑛 𝐻2𝑛
Reactions of Alkenes
Combustion
Addition reactions  hydrogen or halogens attack C==C
double bond and break it.
C—H bonds are nearly pure covalent. Therefore,
hydrocarbons are non-polar. Carbon can single bond to
four other atoms.
Saturated  Single C—C bonds only
Unsaturated  Double/Triple C==C bonds present
Warning: When we say ‘saturated hydrocarbon’, nonhydrocarbons are excluded, such as alkanols, etc.
Aliphatic  straight chain hydrocarbons
Cyclic  ‘ring’ hydrocarbons (cyclo-??)
Alkynes
Contain one triple bond in C chain
Have general formula 𝐶𝑛 𝐻2𝑛−2
Alkanes
Contain only single C—C bonds
Have general formula 𝐶𝑛 𝐻2𝑛+2
Naming Hydrocarbons
Reactions of Alkynes
Combustion & Addition
Alkanols
Carbon chains containing 1+ OH group. End with –ol.
Reactions
Oxidation to produce carboxylic acid
Combustion reactions
Naming
Number of carbon atom to which OH is bonded should be
noted if the C atom is not the first in chain.
Carboxylic Acids
Functional group –COOH at end of chain. End with –oic
acid
C==O bond counts as a double bond – start numbering
from that side, as C==O and –OH group present. THIS
GOES FOR ALL OTHER TYPES (AMINE, ETC)
Linear Polymers
Polymers in which main chain is unbranched. May have
side branches.
Isotactic  Side branches arranged on same side of
polymer. Atactic  Side branches randomly distributed
between sides
Cross Linking
Thermosetting polymers are cross linked. Thermoplastic
polymer chains are linked by heat/chemical to form bonds
between chains.
Small amount crosslinking  rubberlike (elastomer)
Lots of crosslinking  hard/brittle (thermosetting)
Condensation Polymerisation
Formed by head-to-tail joining of monomer units. Usually
involves loss of small molecule. Requires presence of 2
reactive functional groups on each monomer.
Suitable functional groups:
LDPE  Many side branches
HDPE  Few side branches
Surface Interactions
Interface/Surface plane of interaction between
solid/liquid/gas
Surface Energy  Amount of energy required to create a
new surface. Measured in MJ/m2