Transcript Basic Chemistry

Chemistry
Elements, Atoms and Molecules
Why Chemistry?

Nature is not neatly packaged into the individual
life sciences.

Biology is a multidisciplinary science, drawing on
the insights from other sciences.
Essential Elements

96.3% of living matter
O = 65%, C = 18.5%, H = 9.5%, N = 3.3%

3.7%
Ca, P, K, S, Na, Cl, Mg

trace elements (<0.01% each) required in minute
amounts
B, Cr, Co, Cu, F, I, Fe, Mn, Mo, Se, Si, Sn, V, Zn

Some trace elements, like iron (Fe) are
required by all organisms.


other trace elements are
required only by some species.
for example, a daily intake
of 0.15 milligrams of iodine
is required for normal
activity of the human
thyroid gland.
Atoms


Subatomic particles:
atomic nucleus



protons (p +)
neutrons (n)
cloud around the nucleus

electrons (e -)
Electrostatic Charges

The attractions between the positive charges in
the nucleus and the negative charges of the
electrons keep the electrons in the vicinity of the
nucleus.
Mass of Subatomic Particles

1 dalton = 1.67 X 10-24 g
p = 1 dalton
n = 1 dalton
e = 1/2000 dalton

* Dalton = a.m.u. (atomic mass unit)



Relationships


Atomic N° =
mass N° =



# of protons
#p+#n
2
4
Atomic weight = a measure of its mass, can be
approximated by the mass number.
For example, He has a mass number of 4 and an
estimated atomic weight of 4 daltons.
More precisely, its atomic weight is 4.003 daltons
Radioactive Isotope



Most isotopes are stable; they do not tend to lose
particles.
 Both 12C and 13C are stable isotopes.
Some isotopes are unstable and decay spontaneously,
emitting particles and energy.
 14C is a one of these unstable or radioactive isotopes.
In its decay, an neutron is converted to a proton and
electron.
 This converts 14C to 14N, changing the identity of that
atom.
 Half life 14C = 5600 yrs.
 Half life 40K = 1.3 billion yrs.
Radioactive Isotope Uses

Used for
 geological dating
 Treatment


to trace atoms in metabolism
Diagnosis research
Energy Levels or
Electron Shells

Electrons of an
atom may vary in
the amount of
energy that they
possess.

The farther electrons are from the nucleus,
the more potential energy they have.
Electron Shells

Outermost shell (valence shell) determines
behavior of atom


electrons in the outermost shell (valence
shell)


Valence electrons
Atom’s valence
number of covalent bonds that it can form


Octet Rule
valence shell complete with 8 electrons (except H,
He)
Electron Shells
Electron Orbitals and Shells

Electron orbitals



space around nucleus where electrons are most likely to
be found (90% of the time)
can be shaped differently
Shells -energy levels within an orbital



the first shell, closest to the nucleus, has the lowest
potential energy
electrons in outer shells have more potential energy
electrons can only change their position if they absorb
or release a quantity of energy that matches the
difference in potential energy between the two levels
Electron Orbitals and Shells
Superimposed orbitals
Covalent Bond
Sharing of electron pairs




single bonds
double bonds
triple bonds
covalent bonds


nonpolar covalent bonds


electrons shared equally
polar covalent bonds

electrons shared unequally
Single Covalent Bond

the sharing of a pair of valence electrons by two
atoms
Molecular formula
Electron
distribution
Structural
formula
Space Model
Double Covalent Bond

sharing two pairs of valence electrons
Molecular formula
Electron
distribution
Structural
formula
Space
model
Nonpolar Covalent Bond

Carbon and hydrogen have similar
electronegativities
Molecular formula
Electron
formula
Structural
model
Space distribution
Polar Covalent Bond


oxygen has a much higher electronegativity than does
hydrogen.
Compounds with a polar
covalent bond have regions
that have a partial negative
charge near the strongly
electronegative atom and a
partial positive charge near
the weakly electronegative
atom.
Ions


Atoms that have gained or lost electrons
Cations and Anions


lost electron = cation
gain electron = anion
11p + 11e + 11n
17p + 17e + 17n
Na+ = 11p + 10e + 11n
Cl - = 17p + 18e + 17n
Activity: atoms




1. Match the following….isotope, cation, or
anion
When an atom loses an electron you form…
When an atom gains an electron you form…
When an atom changes the number of
neutrons you form…



When an atom loses an electron? ---cation
When an atom gains an electron? ---anion
When an atom changes the number of
neutrons you form… ---isotope
Activity: atoms
 39K




with atomic number 19
What is the proton number?
Neutron number?
Number of electrons?
Mass?
 39K




with atomic number 19
What is the proton number? 19
Neutron number? 20
Number of electrons? 19
Mass? 39
Hydrogen Bond


A slightly positive H atom of a polar covalent
bond in one molecule is attracted to a slightly
negative atom of a polar covalent bond in
another molecule
For example, ammonia
molecules and water
molecules link together
with weak hydrogen bonds
Van der Waals Interactions

Even molecules with nonpolar covalent bonds can
have partially negative and positive regions.



Because electrons are constantly in motion, there can be
periods when they accumulate by chance in one area of a
molecule
This created ever-changing regions of negative and
positive charge within a molecule.
Molecules or atoms in close proximity can be
attracted by these fleeting charge differences,
creating van der Waals interactions
Chemical Reactions
Reactants
+
2 H2
products
=
O2
2 H2O
Redox Reaction
(topic under cellular respiration and photosynthesis)




Chemical reaction involving the transfer of 1 or
more electrons from one reactant to another
Redox = oxidation-reduction reaction
Red = reduction, gain of electrons by a substance
ox = oxidation, loss of electrons by a substance
Redox Reaction
(cellular respiration)
C6H12O6
Reducing
agent
+
6O2
Oxidizing
agent
6CO2 +
6H2O
Oxidized
molecule
Reduced
molecule
Chemical equilibrium
The End