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CHEM 1040 exam full study guide

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These notes cover all the units of CHEM 1040 in-depth including practice exam questions & solutions, diagrams, and easy to understand explanations. It focuses on the tested knowledge from my 2019 exam. Covers many of these topics and more: - Quantum numbers - Trends - Organic Nomenclature - ...

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  • January 14, 2020
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Unit 1 – Atomic Structure
Atoms: consist of protons, neutrons, and electrons; the building blocks of molecules and ions
Isotopes: one, two, or more atoms that have the same atomic number (same protons and
electrons) but different atomic masses (different neutrons and mass). E.g. deuterium and
tritium are isotopes of hydrogen.
Molecules: formed when non-metal atoms combine due to attractive forces that make a
covalent bond such as oxygen, water, and ammonia
Elements: can be atomic (consisting of one atom). E.g. Helium, iron, and calcium. Or molecular
(consisting of two or more atoms). E.g. hydrogen, phosphorous, oxygen, nitrogen, sulfur,
chlorine, bromine, iodine, and fluorine.
Diatomic: indicates there are 2 atoms combined together such as oxygen.
Molar mass: sum of atomic masses of all atoms present in the molecular formula
Ionic compounds: compounds consisting of cations and anions, associated to give an overall
neutral charge. They are generally composed of a combination of metal and non-metal and ions
are held together by electrostatic forces to make an ionic bond.
Empirical formula: relative number of ions in a substance with the smallest possible whole
number subscripts that result in electrical neutrality. E.g. C6H12O6 is CH2O.
Cation: when a neutral atom loss an electron and becomes a positive ion
Anion: when a neutral atom gains an electron and becomes a negative ion

Nomenclature Rules
A. Ionic compounds (metal + non-metal)
1. Binary ionic: name metal (cation) followed by the name of the non-metal (anion),
including suffix ‘ide’.
E.g. Li3N -> lithium nitride

2. Ternary Ionic Compounds: use the name of the cation plus the name of the
polyatomic anion. If transition metals have more than one charge (Fe 2+ or Fe 3+) we
use roman numerals to indicate which version of the metal if used in the compound.
I, II, III, IV, V, VI, VII, VIII, IX.
E.g. NaClO -> sodium hypochlorite OR FeCrO4 -> iron (II) chromate

3. Hydrates: uses prefixes to indicate the number of water molecules attached to the
crystal structure. Mono, Di, Tri, Tetra, Penta, Hexa, Hepta, Octa, Nona, Deca
E.g. CaSO4 + 2H2O -> calcium sulphate dihydrate or MgSO4 + 7H2O -> magnesium
sulphate heptahydrate

B. Molecular Compounds
General Rules and Exceptions
1. Elements further to the left of the periodic table named first and elements closer to
the bottom in a group named first.
2. Oxygen always last unless with fluoride (OF2)
3. Hydrogen last up to group 5 in the periodic table

, 1. Binary Molecular: add suffix ‘ide’ to end of element and use prefixes to indicate the
number of atoms. The a or o at the end of a prefix is dropped when the elements name
begins with a vowel. Mono, Di, Tri, Tetra, Penta, Hexa, Hepta, Octa, Nona, Deca
E.g. P2O5 -> diphosphorous pentoxide (note: the ‘a’ on penta was dropped)

2. Binary Acids: use prefix ‘hydro’ and suffix ‘ic acid’
E.g. HCl -> hydrochloric acid

3. Ternary Oxyacids: modify non-metal (anion) polyatomic ion suffix from ‘ate’ to ‘ic
acid’ and ‘ite’ to ‘ous acid’. Hydrogen in compound is implied in acid.
E.g. H2SO4 -> sulfuric acid OR H2SO3 -> sulfurous acid (note: hydrogen is not named)

Light and Spectroscopy
Spectroscopy: the study of spectra; the dependence of a physical measure to frequency.
Discovered by the study of ‘emission of light’ by atoms and molecules. Used in physical and
analytical chemistry for the identification of substances through the spectrum emitted or
absorbed. Recorded by a spectrometer.
Simple hydrogen emission spectroscopy experiment: Isaac Newton put a narrow beam of light
through a prism and the prism separated the white light into different colours of the rainbow
which shows a continuous spectrum. However, this same experiment was done passing a
current through a tube containing hydrogen gas and it was a line spectrum (not continuous),
instead it shows four lines of colour unique to hydrogen. This is because hydrogen is an excited
gas in which the electrons absorb energy and jump into a higher energy level. When the
electrons fall down to a lower energy level, they emit light.
The spectrum of electromagnetic radiation: a range of all types of electromagnetic radiation,
includes microwaves, infrared light, visible light, ultraviolet light, X-rays, and gamma rays.
Radiation travels at a constant speed in a vacuum.

The particle theory of light: The amount/type of electromagnetic radiation that is emitted is
proportional to temp. Each wavelength of light contains unique energy in the form of photons,
the energy of a photon is some multiple of an energy quantum called Planck’s Constant.
E.g. Black bodies (427 celsius and below) produce very little radiation at visible wavelengths and
appear black. Black bodies above this temperature begin to produce visible wavelengths
starting at red to orange, yellow, white, then blue.
Planck’s Quantum Constant: Max Planck proposed that electromagnetic radiation comes in
units of defined energy rather than random quantities. He was able to determine the following
equation  = h x V where h is Planck’s constant of one photon 6.63 x 10-34 J s. This equation links
the amount of energy of a photon carries with the frequency of its electromagnetic wave.
(Note: this equation is only the energy for one photon at a given wavelength, for 1 mole you
must multiple the energy () calculated by Avogadro’s number 6.022 x 1023 giving the following
equation E = AN x ›

,Light: a type of electromagnetic radiation with fluctuating electric and magnetic fields described
by wavelength () and frequency (v). The wavelength and frequency can be determined by the
following equation c =  x V; where c is the constant speed/velocity of light = 3.00 x 10^8 ms-1
Wavelength and frequency: distance between corresponding points of two consecutive waves
is the wavelength whereas frequency is the number of waves that pass a fixed place in a given
amount of time, low frequency gives low energy and high frequency gives high energy. These
two correlates with shorter wavelengths producing a higher frequency and longer wavelengths
producing a lower frequency.

Bohr model of the atom: the movement of electrons from one orbital to another requires
absorption or emission
Excitation: the absorption of light causing an electron jumps to a higher energy level due to a
gain of energy (photons)
De-excitation: the emission of light causing an electron drops to a lower energy level due to a
loss of energy (photons)
Development of atomic structure: Broglie introduced the relationship between momentum
and wavelength meaning matter exhibits wave and particle properties. He also discovered that
an atom must have an orbital with a circumference equal to a number multiple of the
wavelength.
Ernest Rutherford: discovered the concept of a nucleus in an atom based on alpha rays. He said
that the entire mass of the atom and all its positive charges were concentration at the nucleus
in the centre.
Erwin Schrodinger: he adapted Rutherford and Planck’s theories to obtain the final atom
structure theory that an electron orbital defines the energy and spatial characteristics of the
electron, he replaced fixed orbitals with probability distributions. Erwin created quantum
numbers (n, L, mL). A specific set of these numbers corresponds to an electron orbital.

Quantum Numbers
1. Principal quantum number (n): determines overall size and indicates the energy level of
an electron in an orbital by a number (1, 2, 3, 4, etc). Increasing values of n correspond
to increasing values of energy for the electron; as n increases the spacing between
energy levels grow smaller.
2. Angular momentum quantum number (L): determines the shape of the electron orbital
with permitted values of 0 to (n-1). 0=s (spherical), 1=p (dumb-bell), 2=d (double dumb-
bells), 3=f (multiple lobes).
3. Magnetic quantum number (mL): determines the orientation and number of existing
orbitals in a subshell ranging from -L to +L.
4. Spin quantum number (mS): limits the number of spin energies for an electron to two
values +1/2 and -1/2. Fill +1/2 first then go back and fill -1/2 if enough electrons are
present.

Diamagnetism: not attracted or slightly repelled. 2 electrons paired have opposite spins
meaning their magnetic properties are cancelled out.

, Paramagnetic: attraction of unpaired electrons. A single electron acts as a magnetic
attraction. The more unpaired electrons the larger magnetic moment.


Aufbau principle: orbitals arrange in order of increasing energy levels. Each electron occupies
the lowest energy available. All orbitals of a given sub-level shell have identical charges.
Pauli exclusion principal: no two electrons can have the same sets of all four quantum
numbers.
Hands rule: electrons are distributed among the orbitals of a sub-shell in such a way as to yield
the maximum number of unpaired electrons. Electrons will fill up an orbital before doubling up.
Spin paired orbital: orbital only contains 1 electron
Full orbital: orbital contains maximum amount of 2 electrons

Periodic Trends
1. Atomic radius: distance between the nucleus and outer shell of atom
a. Left to right -> decreases size due to same energy level and increase nuclear charge
b. Top to bottom -> increase size due to increase energy level
2. Ionisation energy: energy it takes to remove outer electron from atom
a. Left to right -> increases due to increase nuclear charge
b. Top to bottom -> decrease due to increase energy level
3. Electron affinity: energy released when an electron is added to a neutral
atom in the gas state to form a negative ion
a. Left to right -> increases due to increase nuclear charge
c. Top to bottom -> decrease due to increase energy level
4. Electronegativity: power atom has for attracting electrons in a covalent
bond with another atom
a. Left to right -> increases due to increase nuclear charge creating increase attraction
b. Top to bottom -> decreases due to smaller nuclear charge creating less attraction

Lewis Structure: diagram showing electrons of outermost shell in an atom in
which atoms will tend to lose/gain unpaired electrons through ionic bonding to
create a full shell or share electrons by covalent bonding.
Valence Shell: chemical properties are determined by an atom’s valence shell. Atoms aim to
achieve a full octet/stable valence shell either by losing/gaining electrons or sharing electrons.
E.g. lack of reactivity of inert gases is associated with filled outer s and p shells




Lewis structure process
Example: NH4+ (ammonium cation)

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