Chapter 2- Fall 2011 Bb

2Small Molecules and the

Chemistry of Life

2 Small Molecules and The Chemistry of Life

• 2.1 How Does Atomic Structure Explain the Properties of Matter?

• 2.2 How Do Atoms Bond to Form Molecules?

• 2.3 How Do Atoms Change Partners in Chemical Reactions?

• 2.4 What Makes Water So Important for Life?

2.1 How Does Atomic Structure Explain the Properties of Matter?

All matter is composed of atoms

An Atom is composed of three types of particles

Electrons—negligible mass; negative charge

Protons—have mass; positive charge

Neutrons—have mass; no charge


Neutrons and protons

– packed together in dense core – Atomic Nucleus

Electrons

– located in a cloud or shell around the nucleus

– attracted to nucleus by positively charged protons

Figure 2.1 The Helium Atom


Atoms have volume and mass.

Mass of one proton or one neutron = 1 atomic mass unit (amu) or 1 dalton, or 1.7 × 10–24 grams.

Mass of one electron = 9 × 10–28 (usually ignored).


Element: pure substance containing only one kind of atom.

• Elements are arranged in the periodic table.

• Each element has a unique chemical symbol.


• Each element has a unique chemical symbol.

• Symbol is usually 1 or 2 letters of its name (may be derived from Latin or German)

H = Hydrogen He = Helium

Na = Sodium (Natrium) Fe = Iron (Ferrum)

Figure 2.2 The Periodic Table (Part 1)


About 98% of every living organism is composed of 6 elements:

– Carbon, Hydrogen, Oxygen, Nitrogen, sulfur, and phosphorus

Figure 2.2 The Periodic Table (Part 2)

He from periodic table here


The number of protons identifies an element

Number of protons = atomic number.

– Written as a subscript to left of chemical symbol 2He, 8O, 16S, 26Fe, 11Na

Number of protons and electrons are generally the same – an atom is electrically neutral


All elements except hydrogen have one or more neutrons.

Mass number = number of protons + number of neutrons.

Mass number = mass of atom in daltons.

– Written as a superscript to left of

chemical symbol - 4He, 16O, 32S, 55Fe, 23 Na


Number of Neutrons

– Mass number - Atomic Number = Number of neutrons


Isotopes: forms of an element with different numbers of neutrons, and thus different mass numbers.

Example:

12C has 6 neutrons

13C has 7 neutrons

14C has 8 neutrons


Isotopes of some elements, such as hydrogen, have special names.


Atomic weight: average of mass numbers of isotopes in their normally occurring proportions.

• Atomic weight of carbon = 12.011

• Atomic weight of hydrogen = 1.0079


• Some isotopes are unstable:

• Radioisotopes give off energy in the form of alpha, beta, and gamma radiation from the nucleus.

• This radioactive decay transforms the atom, including changes in the number of protons.


• The released radiation can be used to detect the presence of radioisotopes.

• Radioisotopes can be incorporated into molecules to act as a “tag” or label.

Figure 2.3 Tagging the Brain

Positron Emission Tomography or PET scan


• Radioisotopes are useful in research and in medicine, but radiation can damage cells and tissues.

• It is sometimes used to treat cancer.


• The number and location of electrons determines how atoms will interact.

• Chemical reactions involve changes in the distribution of electrons between atoms.


• Locations of electrons in an atom are described by orbitals.

• Orbital: region where electron is found at least 90% of the time.

• Orbitals have characteristic shapes and orientations, and can be occupied by a maximum of two electrons.

• Orbitals are filled in a specific sequence.

Figure 2.4 Electron Shells and Orbitals


• Orbitals occur in a series called electron shells or energy levels.

• First shell: one s orbital (holds 2 electrons)

• Second shell: 1 s and 3 p orbitals (holds 8 electrons)

• Additional shells: 4 orbitals (8 electrons)


• The further a shell is from the nucleus, the higher the energy level is for an electron occupying that shell

Figure 2.5 Electron Shells Determine the Reactivity of Atoms


• The outermost electron shell (valence shell) determines how the atom behaves.

• If the outermost shell is full, the atom is stable; it won’t react with other atoms.

– Examples: helium, argon and neon

• Reactive atoms have unpaired electrons in their outermost shell.


• Atoms can share electrons, or lose or gain electrons, resulting in atoms bonded together to form molecules.

• Octet rule: the tendency of atoms to form stable molecules so that they have eight electrons in their outermost shells (however, hydrogen attains stability when 2 electrons are in its single shell)

2.2 How Do Atoms Bond to Form Molecules?

Chemical bond: attractive force that links atoms together to form molecules.

There are several kinds of chemical bonds.


Covalent bonds: atoms share one or more pairs of electrons, so that the outer shells are filled.

Figure 2.6 Electrons Are Shared in Covalent Bonds


• Compound: a molecule made up of two or more elements bonded together in a fixed ratio.

Examples: – Water: H2O = 2 Hydrogen : 1 Oxygen

– Table salt: NaCl = 1 Sodium : 1 Chlorine

• The molecular weight of a compound is the sum of the atomic weights of all atoms in the molecule.

• Organic Chemistry

– Study of carbon compounds

– Great variety of different compounds

• Versatility of carbon

– Atomic number = 6• Six protons and six electrons

– 4 valence electrons• Forms 4 covalent bonds • Arrangement of bonds is a tetrahedral

shape


Figure 2.7 Covalent Bonding Can Form Compounds

Carbon can form four covalent bonds.

Carbon forms bonds with several elements

Hydrogen(valence = 1)

Oxygen(valence = 2)

Nitrogen(valence = 3)

Carbon(valence = 4)

H O N C

Carbon can form single, double, and even triple bonds with other atoms IF they have the electrons to share…



Covalent bonds are very strong—a lot of energy is required to break them.

At temperatures in which life exists, the covalent bonds of biological molecules are quite stable, but this does not preclude change.


• Orientation of bonds:

• The length, angle, and direction of bonds between any two elements are always the same.

• Example: Methane always forms a tetrahedron.

methane tetrahedron here

The shapes of three simple organic molecules

NameMolecular Formula

Structural Formula

Ball-and-StickModel

Space-FillingModel

(a) Methane

(b) Ethane

(c) Ethene (ethylene)



Covalent bonds can be

• Single—sharing 1 pair of electrons

• Double—sharing 2 pairs of electrons

• Triple—sharing 3 pairs of electrons

C C

C H

N N


Sharing of electrons in a covalent bond is not always equal.

Electronegativity: the attractive force that an atomic nucleus exerts on electrons.

It depends on the number of protons and the distance between the nucleus and electrons.

A polar covalent bond results when electrons are drawn to one nucleus more than to the other, because one atom has more electronegativity.


Figure 2.8 Water’s Covalent Bonds Are Polar

Because of this unequal sharing, the bond has a slight negative charge on one end and a slight positive charge on the other.

The bond is said to be polar and produce polar molecules or polar regions of large molecules.


A nonpolar covalent bond results from the equal sharing of electrons between atoms

This occurs when the atoms are close to each other in electronegativity

Example: O2, H2


When one atom is much more electronegative than the other, a complete transfer of electrons may occur.

This results in two ions with fully paired electrons in their outer shells.


Figure 2.9 Formation of Sodium and Chloride Ions


Ions: electrically charged particles— when atoms lose or gain electrons

Cations—positive

Anions—negative

Ionic bonds are formed by the electrical attraction of positive and negative ions.

Salts (such as NaCl) are ionically bonded compounds.


In a solid, ions are close together and the ionic bond is strong.

In water, the ions are far apart and the attraction is much weaker.

Ions interact with polar molecules (e.g., salts dissolve in water).

Figure 2.10 Water Molecules Surround Ions


Hydrogen bond: weak type of bond

Attraction between the partially negative (δ–) end of one molecule and the partially positive (δ+) hydrogen end of another molecule.

Hydrogen bonds form between water molecules, and are important in the structure of DNA and proteins.


Although individually weak, collectively hydrogen bonds have considerable strength

Hydrogen bonds are important in determining and maintaining the three dimensional structures of DNA and proteins

Figure 2.11 Hydrogen Bonds Can Form Between or Within Molecules


Polar molecules that form hydrogen bonds with water are hydrophylic (“water-loving”).

Nonpolar molecules such as hydrocarbons that interact with each other, but not with water, are hydrophobic (“water-hating”).

Figure 2.12 Hydrophilic and Hydrophobic


van der Waals forces: attractions between nonpolar molecules that are close together.

Individual interactions are brief and weak, but summed over a large molecule, can be substantial.

2.3 How Do Atoms Change Partners in Chemical Reactions?

Chemical reactions occur when atoms collide with enough energy to combine or change their bonding partners.

In all chemical reactions, matter and energy are neither created nor destroyed.

Energy: capacity to do work or the capacity for change. Energy usually changes form during chemical reactions.

Reactants Reaction Products

2 H2 O2 2 H2O

• All of the atoms in the reactants must be accounted for in the products

– The reaction must be balanced

Reactants Products

In Photosynthesis (summarized reaction):

6 CO2 + 6 H2O C6H12O6 + 6 O2

This is not what happens chemically, but is a balanced equation of the summarized reaction

2.4 What Makes Water So Important for Life?

Water has unique structure and special properties:

• Polar molecule

• Forms hydrogen bonds

• Tetrahedral shape


Ice floats: Each molecule is hydrogen-bonded to four other molecules in a rigid, crystalline structure.

Solid water is less dense than liquid water.

Figure 2.14 Hydrogen Bonding and the Properties of Water


Water has high specific heat—the amount of heat energy required to raise the temperature of 1 gram of water by 1°C.

Water helps moderate climate because of its high heat capacity.


Water has a high heat of vaporization: the amount of heat energy required to change water from a liquid to a gas state.

The basis of evaporative cooling—heat must be absorbed from the environment in contact with the water.


Cohesion: water molecules resist coming apart from one another.

• Helps water move through plants

• Results in surface tension

Figure 2.15 Surface Tension


A solution is a substance (solute) dissolved in a liquid (solvent).

Many important biochemical reactions occur in aqueous solutions.


Solvent - dissolving agent (water)

Solute - substance that is dissolved (sugar)

Aqueous solution - water is the solvent


Mole: The amount of a substance (in grams) that is numerically equal to its molecular weight.

1 mole of sodium ion Na+ = 23 g

1 mole of hydrogen gas H2 = 2 g

1 mole of table sugar (C12H22O11) = 342 g


The number of molecules in 1 mole is constant for all substances.

One mole contains 6.02 × 1023 molecules = Avogadro’s number

A 1 molar solution (1M) is 1 mole of a substance dissolved in water to make 1 liter of solution.


When acids dissolve in water, they release hydrogen ions: H+ (protons).

H+ ions can attach to other molecules and change their properties.

Bases accept H+ ions.


HCl is a strong acid: the dissolution is complete.

ClHHCl


Organic acids have a carboxyl group:

Weak acids: not all the acid molecules dissociate into ions.

HCOOHCOOH


NaOH is a strong base (it accepts H+).

In water:

OHNaNaOH

OHHOH 2


Weak bases:

• Bicarbonate ion

• Ammonia

• Compounds with amino groups

323 COHHHCO

43 NHHNH

32 NHHNH


Acid-base reactions may be reversible.

Ionization of strong acids and bases is irreversible.

Ionization of weak acids and bases is reversible.

HCOOCHCOOHCH 33


Water has a slight tendency to ionize. OHHOH 2


Water acts as both a weak acid and a weak base.

Ionization of water is important to all organisms because of the abundance of water in living systems, and the reactive nature of the H+ ions.


pH = negative log of the molar concentration of H+ ions.

H+ concentration of pure water is 10–7 M, its pH = 7.

Lower pH numbers mean higher H+ concentration, or greater acidity.

HpH log


1 M HCl solution: H+ concentration = 1 M

pH = 0

1 M NaOH solution: H+ concentration = 10–14 M.

pH = 14


A solution with a pH of 7 is neutral

A solution with a pH of less than 7 is acidic (it has more H+ ions than OH- ions)

A solution with a pH of more than 7 is basic

Figure 2.16 pH Values of Some Familiar Substances


Living organisms maintain constant internal conditions (homeostasis).

Buffers help maintain constant pH.

A buffer is a weak acid and its corresponding base.

323 COHHHCO

Bicarbonate ionWeak base

Carbonic acidWeak acid

pH / Buffers• Bicarbonate buffer - blood buffer

• How it works….

H2CO3 HCO3- + H+

Response toa rise in pH

Response toa drop in pH

H+ donor(acid)

H+ acceptor(base)

Hydrogenion

• Carbonic acid and bicarbonate are in equilibrium in human blood

• Human blood is maintained at pH 7.4

Chapter 2- Fall 2011 Bb

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