BASIC CHEMISTRY REVIEW
Intro to Biochemistry
Chemistry Fundamentals
• All living things are made up of matter
• Matter has mass, occupies space and has many
forms
• The atom is composed of a tiny nucleus
containing protons (+), and neutrons (neutral)
surrounded by (-) electrons in orbit
Chemistry Fundamentals
• Mass number = # of protons and neutrons
• Atomic Number = # of protons
Isotopes
• Atoms that have different number of neutrons
are called isotopes. Same atomic number, but
different mass.
• Carbon has 3 isotopes
• 99% of Carbon found in nature is C-12
• It was discovered that the nucleus of some
isotopes break apart (decay) over time
• These are called radioisotopes
Radioisotopes
• Radioisotopes are radioactive, that is, they decay into
smaller atoms, subatomic particles and energy
• All Radioactive isotopes have a half life-the time it takes
for one half of the atoms in a sample to decay.
• Half life for different radioisotopes vary while rate of
decay of a particular isotope is constant.
Radioisotopes
• Benefits-radiometric dating, radioactive tracers
• Tracers can follow chemicals through chemical
reactions and their path as they move through cells
and the body
• Radio labelled molecules can also be used
Carbon Dating
Measuring the ratio of C-12 to C-14 in a dead/fossilized organism
allows one to calculate the time that has elapsed since the
organisms death. C-12 remains constant while C-14 decays
Hazards
• Radiation from decaying radioisotopes is
harmful to living tissues and cells (mutations)
• Dosimeters are used by physicians/scientists
to monitor radiation levels
Chemical Bonding
• Electrons are arranged in energy levels in spaces
around the nuclei called orbitals. The further
away from the nucleus, the more potential
energy electrons have
• An orbital can only accommodate 2 electrons and
has either a spherical (s) shape or a dumbbell
shape (p)
• n=1 (1st energy level) 1s orbital
• n=2 (2nd energy level) 2s orbital, 2p orbitals (x, y,
z)
• n=1 (max. 2 electrons)
• n=2 (max. 8 electrons)
Orbital shapes
Valence electrons
• Located on outermost “s” and “p” orbitals and
determine an atom’s chemical behaviour
• Noble gases have full valence orbitals and are
thus stable (don’t gain, lose, or share
electrons) He, Ne, Ar...
• Other elements attempt to gain, lose, share
electrons to become stable. It is these
interactions that cause chemical bonds and
jump start reactions.
Lewis dot diagrams
Periodic Table
• Vertical columns=group (family)
• Horizontal rows=periods
Ionic vs. Covalent bonding (video)
• Compounds are stable combinations of atoms of
different elements held together by chemical bonds
(intramolecular forces of attraction)
• When atoms lose electrons they become positively
charged (cations), when they gain electrons they
become negatively charged (anions)
• Ionic bond=attraction b/w cations and anions (eg.
NaCl)
• Covalent bond=2 atoms share one or more pairs of
valence electrons (can be single-H2O, double-O2(g) or
triple N2(g) bonds.
• Covalent bonds are usually stronger than ionic
bonds
Electronegativity and Polarity
• Electronegativity is a measure of an atom’s ability to
attract a shared electron pair (in a covalent bond)
• The larger the electronegativity #, the stronger the
atom attracts the electrons.
• Atoms that attract stronger are assigned a negative
charge (-δ)
• Atoms that attract weaker are assigned a positive
charge (+δ)
• “δ” denotes the term “partial” as electrons are
shared
• The difference in electron attraction forms a polar
covalent bond
Table of electronegativities (p.14)
Electronegativity difference
• If ΔEn is zero that means the electron pair are
sharing equally (non-polar covalent bond)
• If ΔEn is greater than zero but less than 1.7, the
bond is polar covalent
• If ΔEn is greater than 1.7 the bond is considered
ionic.
• Atoms from groups 1, 2 and 16, 17 generally form
ionic compounds
Molecular Shape
• A molecule’s function is determined by its
bonds b/w atoms, shape and polarity.
• When covalent bonds are formed,
hybridization occurs (change in the orientation
of the valence electrons)
• The Valence Shell Electron Pair Repulsion
(VSEPR) Theory can predict molecular shape
VSEPR
• B/c electrons are negative, valence electron
pairs repel each other
• 4 valence pairs (CH )-tetrahedral
• 3 valence pairs (NH3)-pyramidal
• 2 valence pairs (H2O)-angular
• 1 valence pair (HCl)-linear
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Molecular Polarity
• Covalent bonds can be polar or non-polar…but the
polarity of a molecule as a whole depends on bond
polarity and molecular shape.
• Symmetrical structure + polar/non-polar bonds =
non-polar molecules
• Asymmetrical structure + non-polar bonds = nonpolar molecules
• Asymmetrical structure + polar bond(s) = polar
molecules
Water
• Life would not exist without water!
• H2O has polar covalent bonds and asymmetrical
shape, making it a highly polar molecule.
• Bonds b/w molecules are called intermolecular
and help determine physical states (solid, liquid,
gas)
• 3 types; London forces, dipole-dipole forces and
Hydrogen bonds (collectively called van der Waals
forces)
Intermolecular bonds
• London dispersion forces are the weakest.
Exist b/w all atoms and molecules (e.g. noble
gases , non-polar molecules)
• Formed by the temporary unequal distribution
of electrons as they move around the nuclei
• This is why He and CH4 are gases at room
temperature
Intermolecular bonds
• Dipole-dipole forces hold polar molecules together.
Stronger than London forces
• H-bonds are the strongest force of attraction.
• Water molecules combine by H-bonds
• Water is considered the universal solvent due to its
polarity (+ve and –ve charges attract other polar
molecules and ions)
• Disassociation = ionic bonds broken
• Salt in water
Water
• Many substances dissolve in water
(soluble-sugar), while others do
not (insoluble-chalk)
• Miscible-liquids that dissolve into
one another (ethanol is miscible in
water)
• Immiscible-liquids that form
separate layers (oil)
• The subscript “aq”, aqueous
means that it is dissolved in water
Water
• Some small non-polar molecules can’t
form H-bonds with water (O2, CO2) are
only slightly soluble.
• That is why hemoglobin (protein
carrier) is needed to transport oxygen
• Non-polar molecules = hydrophobic
• Polar molecules = hydrophilic
Acids, Base and Buffers
• Acids-sour taste, conduct electricity, turn blue litmus red,
pH less than 7. Increase the concentration of H3O+(aq)
• Bases-bitter taste, slippery feel, conduct electricity, turn
red litmus blue, pH greater than 7. Increase the
concentration of OH-(aq)
• Pure water is neutral, pH of 7 (H3O+(aq) = OH-(aq))
• When acids are mixed with bases it is considered a
neutralization reaction (water and salt are produced)
Strong Acids and Bases
• Dependent on the degree they ionize when
dissolved in water
• Strong acids (HCl) and bases (NaOH) ionize
completely
Weak Acids and Bases
• Weak acids (CH3COOH(aq) - acetic acid) and
bases (NH3 (aq) – ammonia) ionize partially in
water.
• They are reversible reactions in a state of
equilibrium
Acid-Base Buffers
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Living systems are sensitive to pH levels
Living cells use buffers to resist significant changes in pH
Carbonic acid-bicarbonate buffer, most common.
When carbon dioxide and water react, they form
carbonic acid, which then ionizes to form bicarbonate
and H+ ions.
• When H+ (aq) ions enter the blood (acidic food), HCO3 (aq)
reacts with it to produce H2CO3 (aq).
• Together they help maintain the pH of blood around 7.4