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 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)

Hydrogenmolecule (H2)

+ +

+ +

+ +

In each hydrogenatom, the single electronis held in its orbital byits attraction to theproton in the nucleus.

1

When two hydrogenatoms approach eachother, the electron ofeach atom is alsoattracted to the protonin the other nucleus.

2

The two electronsbecome shared in a covalent bond,forming an H2

molecule.

3

Name(molecularformula)

Electron-shell

diagram

Structuralformula

Space-fillingmodel

Methane (CH4). Four hydrogen atoms can satisfy the valence ofone carbonatom, formingmethane.

Water (H2O). Two hydrogenatoms and one oxygen atom arejoined by covalent bonds to produce a molecule of water.

HO

H

H H

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 sodiumatom is transferred to join the 7 valenceelectrons of a chlorine atom.

Each resulting ion has a completedvalence shell. An ionic bond can formbetween the oppositely charged ions.

Na NaCl Cl

+

NaSodium atom(an uncharged

atom)

ClChlorine 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 covalent strong electrons shared equally

electrons0 < 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 theoxygen and apartial positivecharge onthe hydrogens.

H2O

–

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(H2O)

Ammonia(NH3)

OH

H + –

N

HH H

A hydrogenbond results from the attraction between thepartial 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

Hydrogenbonds

+

+

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

• Smaller polar covalent molecules (eg: sugars, alcohols) can dissolve in water, but large nonpolar molecules (eg: oils) do not

• small nonpolar molecules (eg: O2, CO2) 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% H3O+

(aq))

– NaOH(aq) (100% OH-(aq))

• weak acids and bases – ionize only partially when dissolved in water– CH3COOH(aq) (1.3% H3O+

(aq))

– NH3(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)

To Do

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