Transcript CH 2 Notes

The Chemical
Context of Life
Chapter 2
Matter
 Matter consists of chemical elements in pure
form and in combinations called compounds;
living organisms are made of matter.
 Matter -- Anything that takes up space and has
mass.
 Element -- A substance that cannot be broken
down into other substances by chemical
reactions; all matter made of elements.
 Life requires about 25 chemical elements
 96% of living matter is composed of C, O, H,
N.
 Most of remaining 4% is P, S, Ca, K.
 Trace element -- required by organisms in
extremely small quantities: Cu, Fe, I, etc.
Matter cont.
 Compound -- Pure substances made of two or
more elements combined in a fixed ratio.
 Have characterisitics different than the elements
that make them up (emergent property).
 Na and Cl have very different properties from
NaCl.
 Difference between mass and weight:
 Mass -- measure of the amount of matter an object
contains; constant.
 Weight -- measure of how strongly an object is
pulled by earth's gravity; varies.
Nutrient Deficiencies
Atomic structure determines the behavior
of an element
 Atom -- Smallest possible unit of matter that
retains the physical and chemical properties
of its element.
 Subatomic Particles
 1. Neutrons (no charge/neutral; found in
nucleus; ~ 1 amu).
 2. Protons (+1 charge; found in nucleus; ~ 1
amu).
 3. Electrons (-1 charge; electron cloud;
1/2000 amu).
 One amu approx equal to 1.7 x 10-24 g.
Atomic Number and Atomic Weight
 Atomic number = Number of protons in an
atom of a particular element.
 All atoms of an element have the same
atomic number.
 In a neutral atom, # protons = # electrons.
Mass number -- Number of protons and
neutrons in an atom; not the same as an
element's atomic weight.
Examples
 23Mg
Mass number ??
Atomic number ??
12


23
12
 # of protons ??
# of electrons ??
# of neutrons ??

12
12
11
 14C
Mass number ??
Atomic number ??
 6

14
6
 # of protons ??
# of electrons ??
# of neutrons ??

6
6
8
Isotopes
 Isotopes -- Atoms of an element that have the
same atomic number but different mass number;
different number of neutrons.
 Half-life -- Time for 50% of radioactive atoms in a
sample to decay.
 Biological applications of radioactive isotopes
include:
 1. Dating geological strata and fossils.
 Radioactive decay is at a fixed rate; by
comparing the ratio of radioactive and stable
isotope, age can be estimated. in a fossil with
the
 Ratio of Carbon-14 to Carbon-12 is used to
date fossils less than 50,000 years old.
Isotopes cont.
 2. Radioactive tracers
 Chemicals labelled with radioactive isotopes
are used to trace the steps of a biochemical
reaction or to determine the location of a
particular substance within an organism.
 Isotopes of P, N and H were used to
determine DNA structure.
 Used to diagnose disease.
 3. Treatment of cancer
 Can be hazardous to cells.
Energy Levels
 Electrons are directly involved in chemical reactions.
 They have potential energy because of their position
relative to the positively charged nucleus.
 There is a natural tendency for matter to move to the
lowest state of potential energy.
 Different fixed potential energy states for electrons
are called energy levels or electron shells.
 Electrons with lowest potential energy are in energy
levels closest to the nucleus.
 Electrons with greater energy are in energy levels
further from nucleus.
 Electrons may move from one energy level to
another.
Electron Configuration and
Chemical Properties
 Electron configuration -- Distribution of electrons
in an atom's electron shells; determines its
chemical behavior.
 Chemical properties of an atom depend upon the
number of valence electrons (electrons in the
outermost energy level.
 Octet rule -- A valence shell is complete when it
contains 8 electrons (except H and He).
 An atom with an incomplete valence shell is
chemically reactive (tends to form chemical bonds
until it has 8 electrons to fill the valence shell).
 Atoms with the same number of valence electrons
show similar chemical behavior.
Bonding in Molecules
 Chemical bonds -- Attractions that hold molecules together.
 Molecules --Two or more atoms held together by chemical
bonds.
 Covalent bond -- formed between atoms by sharing a pair
of valence electrons; common in organic compounds.
 Single covalent bond -- Bond between atoms formed by
sharing a single pair of valence electrons.
 Double bond -- share two pairs of valence electrons.
 Triple bond -- share three pairs of valence electrons.
 Compound = A pure substance composed of two or more
elements combined in a fixed ratio.
 For example: water (H2O), methane (CH4).
Nonpolar Covalent Bonds
 Electronegativity -- Atom's ability to attract and hold
electrons.
 • The more electronegative an atom, the more
strongly it attracts shared electrons.
 • Scale determined by Linus Pauling:
 O = 3.5; N = 3.0; S and C = 2.5; P and H = 2.1.

 Nonpolar bond -- Covalent bond formed by an equal
sharing of electrons between atoms.
 • Occurs when electronegativity of both atoms is
about the same.
 • Molecules made of one element usually have
nonpolar covalent bonds (H2 and O2).
Polar Covalent Bonds
 Polar bond -- Covalent bond formed by an
unequal sharing of electrons between atoms.
 • Occurs when the atoms involved have
different electronegativities.
 • In water, electrons spend more time around
the oxygen than the hydrogens. This causes
the oxygen atom to have a slight negative
charge and the hydrogens to have a slight
positive charge.
Ionic Bonds
 Ion -- Charged atom or molecule.
 Anion -- An atom that has gained one or more
electrons from another atom; negatively charged.
 Cation -- An atom that has lost one or more
electrons; positively charged.
 Ionic bond -- Bond formed by the electrostatic
attraction after the complete transfer of an electron
from a donor atom to an acceptor.
 Strong bonds in crystals, but fragile bonds in water.
 Ionic compounds are called salts (e.g. NaCl or table
salt).
Biologically important weak
bonds
 Include: Hydrogen bonds; Ionic
bonds in aqueous solutions; Van
der Waals forces.
 Hydrogen bond -- Bond formed
by the charge attraction when a
hydrogen atom covalently
bonded to one electronegative
atom is attracted to another
electronegative atom.
 Van der Waals -- charge
attraction between oppositely
charged portions of polar
molecules.