Basic Chemistry
Atoms, Elements & Molecules
Atoms Around Us
“To understand the very large, we must
understand the very small”
Democritus
 Atoms

Atomos (indivisible)
Atom Anatomy

Electrons


Protons


Negative charge
Positive charge
Neutrons

Neutral
Atomic Number

Each atom has electron orbitals (energy
levels)




If completely filled, the atom is stable
If not completely filled, the atom is unstable
2n2
Elements are arranged according to their
atomic number.
Electron Dot Diagrams
Periodic Table

Dimitri Ivanovich Mendelèev (1834 – 1907)

Organized elements in order of increasing atomic
weight

Atomic weight is the average mass of the atoms
in a representative sample of an element.

Isotopes

Atoms with the same number of protons, but different
number of neutrons
Periodic Table
Groups (families)
Periods
Periodic Table







Metals
Nonmetals
Group IA = Alkali metals
Group IIA = Alkaline earth metals
Group VIIA = Halogens
Extreme right = Noble gases
Group B = Transition metals
Metals

1.
2.
3.
4.
Four characteristics:
Conduction
Reactivity
Chemical
Alloys
Happy Noble Gases

These elements are in Group VIIIA

8 electrons in the outer orbital





The fact that their outer orbitals are full means they
are quite happy not reacting with other elements
Helium
Neon
Argon
Xenon
Transition Metals



Advanced orbital rules
In general, they are
elements in which the
inner electron energy
levels are being filled.
In other words, they
are able to use the
two outermost orbitals
to bond with other
elements.
Ions

Ions are atoms with either extra electrons, or
missing electrons


Cations


In other words, the number of electrons are not
equal to the number of protons.
sodium
Anions

chloride
Ionic Bonds

Generally speaking, those elements on the
left hand side of the table react with those
elements on the far right (excluding the noble
gases) to form stable crystalline solids.

Metals give up electrons to elements on the
right (nonmetals).
Covalent Bonds

Shared pairs of electrons

Many elements are diatomic


Meaning that they can attach to each other
For example, chlorine atoms:
Cl
Chemical Reactions



Chemical change must occur.
A reaction could include ions, molecules, or
pure atoms.
Reaction Rate & Collision Theory



Concentration
Temperature
Pressure
Chemical Equations

Law of Conservation of Mass

Reactants
Products
Catalysts
Inhibitors



Balancing Chemical Equations

A silver spoon tarnishes. The silver reacts
with sulfur in the air to make silver sulfide, the
black material we call tarnish:
___ Ag + ___S → ___Ag2S
Types of Chemical Reactions

Composition
(Synthesis)




AB + C → CB + A
Double Replacement


Requires heat energy

Exothermic

Releases heat energy
XY → X + Y
Single Replacement

Endothermic
A + B → AB
Decomposition
(Desynthesis)


AB + CD → AD + BC

Formation of a
Precipitate
Chemical Reaction Properties


Reversible
Equilibrium
Acids & Bases
Svante Arrhenius (1887)
ACIDS
1.
2.
3.
4.
Turn indicator dye litmus
from blue to red
React with active metals
such as zinc, iron, and tin,
dissolving the metal and
producing hydrogen gas
Taste sour, if diluted
enough to be tasted safely
React with certain
compounds called alkalis
or bases to form water and
compounds called salts
BASES
1.
2.
3.
4.
Turn the indicator dye
litmus from red to blue
Feel slippery or soapy on
the skin
Taste bitter
React with acids to form
water and salts
Acids

Arrhenius proposed that these characteristic
properties of acids are actually properties of
the hydrogen ion (H+), and that acids are
compounds that yield H+ in aqueous
solutions.


Slightly modified today
Hydronium ion (H3O+)

For simplification, we’ll stick with the H+
terminology.
Acids

Monoprotic


One H+
Two


H+
Triprotic
Three H+
Strong Acids

Diprotic




Weak Acids



Polyprotic

General term for acids
that give up more than
one H+
Ionize completely (or
nearly completely) in
water
HCl (hydrochloric acid)
Ionize only slightly in
water
CH3COOH (acetic acid)
Bases
Yield hydroxide ions (OH-) in aqueous solutions

Monobasic



Two hydroxyl anions
Tribasic


One hydroxyl anion
Strong Bases
Dibasic



Three hydroxyl anions
Polybasic

General term for bases
that give up more than
one OH-

Completely ionize
 NaOH (sodium
hydroxide; lye)
 All the bases of Group I
and Group II are strong
bases
Weak Bases

NH3 (ammonia)
pH Scale

pH = -log [H+]
Brønsted-Lowry Acid-Base
Theory

By the 1920’s chemists were working with
solvents other than water.

Acid


Proton (H+) donor
Base

Proton (H+) acceptor
Acid-Base Titrations

Method used to
determine just how much
acid (or base) there is in
a solution of unknown
concentration

Buret

A piece of laboratory
glassware designed to
deliver known amounts of
liquid into another
container
A Word About Moles….

A mole used in chemistry is something like
the dozen we use every day.

A mole simply means that you have 6.02 x
1023 of whatever you’re talking about.


Avogardo’s number
Molarity is defined as the number of moles of
solute divided by the number of liters of
solution

Molarity (M) = moles of solute
liters of solution
Lab Prep (Tomorrow)



Salinity & Conductivity
SPM filter prep.
Glassware Use

Pipettes

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Burets


Dilutions of copper II sulfate
Acids & Base Titration
Nutrients

Prep. of standards for nutrient analysis
Lab Prep (Next Week - Sierra)

Field Trip to collect water for nutrient
analyses



Salinity, DO, and pH will be recorded on site
Nutrients measured the week after in the lab
Watershed Readings

North Carolina Division of Water Quality