BASIC CHEMISTRY REVIEWIntro 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 DatingMeasuring 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 (CH4)-tetrahedral• 3 valence pairs (NH3)-pyramidal• 2 valence pairs (H2O)-angular• 1 valence pair (HCl)-linear

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 = non-polar 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• 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