Chapter 2—Chemical Context

of Life

Atoms, Elements, Compounds,

and Molecules

Hierarchy of Biological Order

Emergent Properties

Figure 2.2

I. Chemical Elements &

Compounds

• Element

– Cannot be broken down to other substances

by chemical reactions – Examples: carbon (C), sodium (Na), oxygen (O)

• Compound

– Substance containing 2 or more elements

combined in a fixed ratio – Examples: H2O, NaCl, C6H12O6 (emergent properties)

Which 4 are the most common elements in the human body?

Life requires ~25 chemical elements

• Four elements make up 96% of living

matter: • carbon (C) • hydrogen (H)

• oxygen (O) • nitrogen (N)

• Four more elements make up most of

remaining 4%: • phosphorus (P) • calcium (Ca)

• sulfur (S) • potassium (K)

• Trace elements (<0.01%)

II. Atoms & Molecules

• Atomic structure determines the behavior

of an element

– Atom

• Smallest unit of matter that retains the properties

of an element

– C (atom) vs. C (element)

Subatomic Particles

Particle Charge Location Mass

(amu/dalton)

proton + nucleus 1

neutron 0 nucleus

1

electron - Cloud outside

nucleus 0

(negligible)

Atomic nucleus vs. cell nucleus?

Simplified Model of a Helium (He)

Atom

Atoms are mostly empty space—

(nucleus = golf ball, electron cloud = 1 km)

Figure 2.5

Atomic Number and Mass

• Atomic number

– # of protons in nucleus

of an atom

– Also = # of electrons

– Unique for a particular

atom

• Mass number

– The sum of protons +

neutrons in nucleus

Isotopes

• How are isotopes different than ‘regular’

atoms?

– Isotope

• An atom with more neutrons than usual

(larger mass)

• Behaves the same in chemical reactions

– Examples: carbon-13, carbon-14 (99% = carbon-12)

– Why is the atomic mass of carbon 12.011, not 12?

Use of Radioactive (unstable) Isotopes

Substances are ‘labeled’

with isotopes in order to:

- follow metabolic

processes

- find their locations within

cells

- to use as diagnostic

tools in medicine

Figure 2.6

Electron Energy Levels

• Electrons have potential energy due to position in relation to nucleus

• Electrons exist only at fixed levels of potential energy (electron shells)

Figure 2.9

Electron Energy Levels

Electron energy levels (shells) have different states of potential energy

(Higher levels have more energy)

Ball on a

staircase…

Electron Configurations & Chemical

Properties

•Chemical behavior/bonding of an atom depends on # of electrons in

its outermost shell (valence shell)

•Atoms with same # of valence electrons behave similar chemically

(F & Cl) (O & S)

•Atoms with completed

valence shell are

unreactive (noble

gases) (Ne & Ar)

Figure 2.10

Chemical Reactivity

• Atoms tend to complete a partially filled

valence shell

or

• empty a partially filled valence shell

– This tendency drives chemical

reactions…and creates bonds

Electron Orbitals

Orbital = 3-dimensional space where an electron is found 90% of the time

Rule—no more than 2 electrons per orbital

One electron shell/level may contain multiple orbitals

Figure 2.11

Strangers getting

on a bus….

Atoms combine by chemical

bonding

• Chemical Bonds

– Attraction between 2 atoms due to:

• sharing of outer shell electrons (covalent bonds)

– or

• The presence of opposite charges on the atoms

(ionic bonds)

– Bonded atoms gain complete outer electron

shells

Covalent Bonding forms MOLECULES

•Covalent Bond = 2

atoms sharing a pair

of valence electrons

•Valence = bonding

capacity of an atom

(# of unpaired e-)

(Single, double, &

triple bonds

possible) Figure 2.12

Nonpolar Covalent Bonds

• Electronegativity

– Attraction of an atom for the electrons in a covalent bond

• The more electronegative, the more strongly it pulls

• Nonpolar covalent bond

– electrons are shared equally between atoms (equal tug of war)

– Examples: O2, H2, CH4

Type of Bonding?

Water Molecule

(polar covalent bonds)

Polar covalent

bond = electrons

are not shared

equally between

atoms

(e- spend more

time closer to the

more

electronegative

atom)

i.e. H2O

Figure 2.13

Type of Bonding?

Ionic Bonding

Transfer of electron from one atom to another causes ions to form

cation—ion with positive charge

anion—ion with negative charge

Opposite charges attract = ionic bond

Figure 2.14

Ex. NaCl

Ionic Compounds (salts)

Why is an ionic

compound not

called a

molecule?

(no definite size

or number of

atoms, only a

ratio of

elements)

Example: MgCl2

Figure 2.15

Type of Bonding?

Hydrogen Bonding

(weak chemical bond)

A hydrogen atom (+)

from one molecule is

attracted to an

electronegative atom (-)

in another molecule

Attraction = hydrogen

bond

Figure 2.16

Van der Waals interactions

weak attractions between molecules or parts

of molecules due to localized charge

fluctuations

Due to random chance

Molecules/atoms must be very close together

The function of a molecule is related to

its shape

Specific molecular

shapes allow for

molecule to molecule &

cell to cell

communication

(lock & key)

Figure 2.17

Molecular Shape & Brain Chemistry

Molecular Shape & Brain Chemistry

Figure 2.18

Molecular Mimics

Figure 2.19

Chemical Reactions—making and

breaking chemical bonds

Law of Conservation of Mass—same # of each atom on both sides

• Most reactions are reversible • Example: 3H2 + N2 ↔ 2NH3

• Chemical Equilibrium

– the point at which the rate of the forward

reaction equals the rate of the reverse

reaction

– Concentrations of products/reactants stop

changing