1. Atoms, Elements & Molecules
Basic Chemistry
Atoms, Elements & Molecules

2. Atoms Around Us
“To understand the very large, we must understand the very small”
Democritus
Atoms
Atomos (indivisible)

3. Atom Anatomy Electrons Protons Neutrons Negative charge
Positive charge
Neutrons
Neutral

4. Atomic Number Each atom has electron orbitals (energy levels) 2n2
If completely filled, the atom is stable
If not completely filled, the atom is unstable
2n2
Elements are arranged according to their atomic number.

5. Electron Dot Diagrams

6. 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

7. Periodic Table
Groups (families)
Periods

8. Periodic Table Metals Nonmetals Group IA = Alkali metals
Group IIA = Alkaline earth metals
Group VIIA = Halogens
Extreme right = Noble gases
Group B = Transition metals

9. Metals
Four characteristics:
Conduction
Reactivity
Chemical
Alloys

10. Happy Noble Gases These elements are in Group VIIIA Helium Neon Argon
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

11. 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.

12. Ions Ions are atoms with either extra electrons, or missing electrons
In other words, the number of electrons are not equal to the number of protons.
Cations
sodium
Anions
chloride

13. 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).

14. Covalent Bonds Shared pairs of electrons Many elements are diatomic
Meaning that they can attach to each other
For example, chlorine atoms: Cl

15. Chemical Reactions Chemical change must occur.
A reaction could include ions, molecules, or pure atoms.
Reaction Rate & Collision Theory
Concentration
Temperature
Pressure

16. Chemical Equations Law of Conservation of Mass Reactants Products
Catalysts
Inhibitors

17. 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

18. Types of Chemical Reactions
Composition (Synthesis)
A + B → AB
Decomposition (Desynthesis)
XY → X + Y
Single Replacement
AB + C → CB + A
Double Replacement
AB + CD → AD + BC
Endothermic
Requires heat energy
Exothermic
Releases heat energy
Formation of a Precipitate

19. Chemical Reaction Properties
Reversible
Equilibrium

20. Acids & Bases Svante Arrhenius (1887)
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
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

21. 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.

22. Acids Monoprotic Diprotic Triprotic Polyprotic Strong Acids Weak Acids
One H+
Diprotic
Two H+
Triprotic
Three H+
Polyprotic
General term for acids that give up more than one H+
Strong Acids
Ionize completely (or nearly completely) in water
HCl (hydrochloric acid)
Weak Acids
Ionize only slightly in water
CH3COOH (acetic acid)

23. Bases Yield hydroxide ions (OH-) in aqueous solutions
Monobasic
One hydroxyl anion
Dibasic
Two hydroxyl anions
Tribasic
Three hydroxyl anions
Polybasic
General term for bases that give up more than one OH-
Strong Bases
Completely ionize
NaOH (sodium hydroxide; lye)
All the bases of Group I and Group II are strong bases
Weak Bases
NH3 (ammonia)

24. pH Scale
pH = -log [H+]

25. 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

26. 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

27. 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

28. Lab Prep (Tomorrow) Salinity & Conductivity SPM filter prep.
Glassware Use
Pipettes
Dilutions of copper II sulfate
Burets
Acids & Base Titration
Nutrients
Prep. of standards for nutrient analysis

29. 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