Biochemistry Basics Section 1.1

Subatomic Particles and the Atom Protons (+ charge) and neutrons (neutral) – found in the nucleus Electrons (- charge) – Surround the nucleus in a “cloud” or orbital Orbital – the 3D space where an electron is found 90% of the time – Each orbital can only fit only 2 electrons

Bonding – Covalent Bonds Atoms bond through interaction of their valence (outer orbital) electrons Covalent bond – electrons are shared between atoms and the valence orbitals overlap Hydrogen atoms (2 H) Hydrogen molecule (H 2 ) In each hydrogen atom, the single electron is held in its orbital by its attraction to the proton in the nucleus. 1 When two hydrogen atoms approach each other, the electron of each atom is also attracted to the proton in the other nucleus. 2 The two electrons become shared in a covalent bond, forming an H 2 molecule. 3

Name (molecular formula) Electron- shell diagram Structural formula Space- filling model Methane (CH 4 ). Four hydrogen atoms can satisfy the valence of one carbon atom, forming methane. Water (H 2 O). Two hydrogen atoms and one oxygen atom are joined by covalent bonds to produce a molecule of water. H O H HH H H C

Ionic Bonds In some cases, atoms strip electrons away from their bonding partners Ionic bond – electrons are transferred from one atom to the other, resulting in a negative ion (anion) and a positive ion (cation), which are electrostatically attracted to each other

Cl – Chloride ion (an anion) – The lone valence electron of a sodium atom is transferred to join the 7 valence electrons of a chlorine atom. Each resulting ion has a completed valence shell. An ionic bond can form between the oppositely charged ions. Na Cl + Na Sodium atom (an uncharged atom) Cl Chlorine atom (an uncharged atom) Na + Sodium on (a cation) Sodium chloride (NaCl)

Covalent bonds are stronger than ionic bonds Covalent and Ionic bonds are intramolecular forces of attraction because they are within molecules

Polarity Electronegativity – Is the attraction of an atom for electrons The more electronegative an atom – The more strongly it pulls electrons toward itself The smaller the atom – the more electronegative

to determine the type of bond between two atoms, calculate the difference between their electronegativity values =0 covalentstrongelectrons shared equally electrons 0 < x < 1.7 polar covalent partially shared >= 1.7 ionic weak electrons not (extreme polarity) shared the greater their difference in electronegativity, the greater the polarity of that substance

Polar Covalent Bond – electrons are shared unequally between atoms of different electronegativity; electrons are closer to the atom with the higher value This results in a partial negative charge on the oxygen and a partial positive charge on the hydrogens. H2OH2O –– O H H ++ ++ Because oxygen (O) is more electronegative than hydrogen (H), shared electrons are pulled more toward oxygen.

Intermolecular Forces intermolecular forces of attraction exist between molecules London forces – form when the electrons of one molecule are attracted to the positive nuclei of neighbouring molecules; holds large nonpolar molecules together; very weak

hydrogen bonds – form when the slightly negative O or N that is bonded to a slightly positive H is attracted to the slightly positive H of a neighbouring molecule; strongest  – –  + + Water (H 2 O) Ammonia (NH 3 ) O H H  + +  – – N H H H A hydrogen bond results from the attraction between the partial positive charge on the hydrogen atom of water and the partial negative charge on the nitrogen atom of ammonia. ++ ++ Figure 2.15

dipole-dipole forces – form when the slightly negative end of a polar molecule is attracted to the slightly positive end of a neighbouring polar molecule; stronger – Occurs because electrons are in constant motion and may accumulate by chance on one part of the molecule. The result is “hot spots” of positive and negative charge.

Water highly polar because of asymmetrical shape and polar covalent bond The polarity of water molecules results in hydrogen boding Hydrogen bonds + + H H + +  – –  – –  – –  – – Figure 3.2

“Like Dissolves Like” ionic compounds dissolve in water because the ions separate

However, molecules do not need to be ionic to dissolve in water polar covalent molecules (eg: sugars, alcohols) can dissolve in water, but large nonpolar molecules (eg: oils) do not small nonpolar molecules (eg: O 2, CO 2 ) are slightly soluble and need soluble protein molecules to carry them (eg: hemoglobin transports oxygen through the blood)

hydrophilic – “water-loving;” dissolves in water – e.g. polar or ionic molecules, carbohydrates, salts hydrophobic – “water-fearing;” does not dissolve in water – e.g. non-polar molecules, lipids

Acids and Bases acid – donates H + to water; pH 0-7 base –donates OH - to water (or H3O); pH 7-14 neutralization reaction – the reaction of an acid and a base to produce water and a salt (ionic compound)

Strong and Weak Acids/Bases strong acids and bases – ionize completely when dissolved in water – HCl (aq) (100% H 3 O + (aq) ) – NaOH (aq) (100% OH - (aq) ) weak acids and bases – ionize only partially when dissolved in water – CH 3 COOH (aq) (1.3%  H 3 O + (aq) ) – NH 3(aq) (10%  OH - (aq) )

Buffers The internal pH of most living cells must remain close to pH 7 Buffers – Are substances that minimize changes in the concentrations of hydrogen and hydroxide ions in a solution – Can donate H+ ions or remove H+ ions when required – E.g. carbonic acid creates bicarbonate ions (base) and hydrogen ions (acid) (reversible reaction)

Functional Groups Functional groups – Are reactive clusters of atoms attached to the carbon backbone of organic molecules GroupChemical Formula Structural FormulaFound In hydroxyl—OHalcohols (eg: ethanol) carboxyl—COOHacids (eg: vinegar) amino—NH 2 bases (eg: ammonia)

sulfhydryl—SHrubber phosphate—PO 4 ATP Carbonyl (aldehydes) (keytones) —COH —CO— aldehydes (eg: formaldehyde) ketones (eg: acetone)

To Do Section 1.1 Questions – Pg. 23 #1, 2, 4, 6-8, 12, 14, 15