1. Chemistry Unit Notes
8th Grade Science

2. Basic Vocabulary Matter: Anything that has mass and volume
Mass: Amount of matter in an object
Weight: Measure of the force of attraction between objects due to mass and gravity
Volume: Amount of space an object takes up
Density: Measurement of how much mass is contained in a given volume

3. More Vocabulary
Atoms: Smallest particle of an element that has all the properties of matter:
Protons- particles in the nucleus with positive charge
Electrons- particles orbiting around nucleus with negative charge
Neutrons- particles in the nucleus with no charge
Elements: Simplest form of a pure substance
Compounds: Two or more elements chemically combined to form a new substance

4. Sub-Atomic Particles Part of Atom Charge Location Mass/Size Electron
- negative
outside nucleus
.0006 amu
(too little to count)
Proton
+ positive
inside nucleus
1 amu
Neutron
no charge

5. Periodic Table

6. Using the Periodic Table
Atomic Number
Equal to # protons = # electrons
Periodic Table is arranged by this number
Symbol
“Shorthand” for the element – Note 2nd letter is always lowercase
Atomic Mass Number
Total AVERAGE mass of Protons + Neutrons + Electrons 17 Cl
35.5

7. Electron Energy Levels
Electrons are arranged in “Shells” around nucleus in predictable locations
Fill “seats” closest to nucleus first (concert – best seats)
“Seats” available
Shell #1 2 electrons
Shell #2 8 electrons
Shell #3 8 electrons
Shell # electrons
Shell # electrons
Shell #6 50 electrons
Ex. Carbon has 6 total electrons so…
Two electrons on first energy level
Four electrons on second energy level
Question: Could we fit more electrons on the second energy level if there were more electrons in carbon??

8. Atomic Structure
Total # of protons and electrons (in a neutral atom)
17 protons in nucleus
17 electrons orbiting nucleus 17 Cl
Element Name
Chlorine
35.5
Total Mass of Nucleus
= 18 neutrons
(Round Atomic Mass)
Notice: electrons follow energy level rules
from previous slide.

9. Atomic Mass – Fractions?
Look at Chlorine (atomic number 17)
Atomic mass of 35.5? I dont’ get it!
Where does the 35.5 come from?
0.5 protons? 0.5 neutrons?  No
Atomic Mass = average number of protons and neutrons in nature

10. More Practice
Determine the name, number of protons, neutrons and electrons for each element shown and draw… 26 8 15 Fe O P 56 16 31

11. Isotopes
An isotope is a variation of an element (same protons) but can have diff. # of neutrons
Ex: carbon (atomic mass = )
Carbon (14) and carbon (12) exist in nature

12. Ions Change in electrons which gives an atom a charge (+ or -)
You can only add or subtract electrons! (protons don’t change)
Ex. Count the number of electrons below…
Carbon ion (-1 charge)
7 electrons (-)
6 protons (+)
Neutral Carbon
6 electrons (-)
6 protons (+)
Carbon ion (+1 charge)
5 electrons (-)
6 protons (+)

13. Valence Electrons An electron on the outermost energy shell of an atom
Important to understand because this is a key factor in how atoms will BOND with each other
Octet rule – stable atom will have 8 electrons in that outer shell
Practice – Valence # of
Chlorine?
Neon?
Nitrogen?
Oxygen?

14. Electron Dot Diagrams
a diagram that represents the # of valence electrons in an atom of an element.
The amount of electrons is displayed by dots around the symbol of the element.
Ex.

15. Types of Chemical Bonds
Ionic- Two elements bond by transferring electrons to create ions that attract together (+ is attracted to - after an electron is transferred)
Covalent- Two elements bond by sharing electrons (strongest bond type)
Metallic- Two metals bond and form a “common electron cloud”. This is a cluster of shared electrons (weakest bond type)

16. Examples of Bonding http://www.youtube.com/watch?v=xTx_DWboEVs

17. Predicting Bonds Ionic Bond = metal to non-metal
Covalent = non-metal to non-metal
Metallic = metal to metal
Do you understand why? HINT: the numbers at the top of the table indicate the # of valence electrons for each column

18. Oxidation Numbers Oxidation numbers are assigned to each element
They represent a predicted “charge” of an atom/ion when it bonds with another element.
(tells us if the atom would prefer give or take electrons, and how many).
They help us to predict what compounds will form when two elements get together.
Oxidation numbers are labeled like this:
Na 1+
O 2-

19. How to Use Oxidation Numbers
Oxidation Number indicates the number of electrons lost, gained or shared when bonding with other atoms.
Ex. Na wants to lose an electron. If an electron is lost, it becomes a +1 charge
SO: oxidation number for Na = 1+
Ex. Cl wants to gain an electron. If an electron is gained, it becomes a -1 charge
SO: oxidation number for Cl = 1-

20. Oxidation Numbers
Each column going down the periodic table has elements with the same oxidation number.

21. Label the oxidation numbers on your periodic table at the top of each column as shown here:
(+/-)

22. Rules for using oxidation numbers to create compounds
Positive ions can only bond with negative ions and vice versa
2. The sum of the oxidation numbers of the atoms in a compound must be zero (the key is to stay balanced)
3. If the oxidation numbers are not equal to zero, then you must add additional elements until they balance at zero.
4. When writing a formula the symbol of the Positive (+) element is followed by the symbol of the negative (-) element.

23. Examples of Forming Compounds
Ex. Na (+1) + Cl (-1) = NaCl
Are these oxidation numbers already equal to zero?
If so, you don’t need to add any extra elements to combine them into a compound, so the answer is simply NaCl
Ex. H (+1) + O (-2) = H2O
How many +1 would you need to balance the -2 to zero?
Since you need 2 atoms of the 1+ to balance the 2- to zero the resulting compound would be H2O
In other words: to combine H with O, you MUST have 2 H to balance the oxidation numbers to zero
2+ and = ZERO
Ex. Al (+3) + S (-2) = Al2S3
This one is tricky…we are not even close to balancing + and - to zero.
Because of this we must have more than one Al and more than one S in our final equation.
By using 2 Aluminums instead of just1 we would have 6+
By using 3 sulfers instead of just 1 we would have 6-
Since these are now equal to zero, we combine 2 Aluminums and 3 Sulfers to make Al2S3

24. Chemical vs. Physical Change
Physical Change: A change that can occur without changing the identity of the substance.
Ex. Solid, Liquid, Gas (Phase change)
Chemical Change: Process by which a substance becomes a new and different substance
Ex. Fire

25. Chemical Reactions
Chemical Reaction: a process in which the physical and chemical properties of the original substance change as new substances with different physical and chemical properties are formed

26. Chemical Reaction Basics
H2 + O2 --> H2O
Reactants
Products
Reactants- substance that enters into a reaction
Products- substance that is produced by a chemical reaction

27. Evidence of Chemical Change
EPOCH is an acronym that stands for evidence that a chemical reaction has occurred.
– Effervescence (bubbles and/or gives off gas)
– Precipitate (solid crystals form)
– Odor (change of smell is detected)
– Color change
– Heat (reaction either heats up or cools down)
Does sighting evidence of a chemical reaction mean that a chemical reaction has undoubtedly taken place? E P O C H

28. Types of Reactions Romance Chemistry :)
Synthesis- Marriage/Dating
A + B = AB
Decomposition- Divorce/Breakup
AB= A + B
Single-Replacement- Dance Cut In
A + BC = AC + B
Double-Replacement- Dancing couples switch partners.
AB + CD = AC + BD

29. Cartoon Chemistry
This is an example of synthesis

30. Cartoon Chemistry
This is an example of a decomposition

31. Cartoon Chemistry
This is an example of a single replacement

32. Cartoon Chemistry
This is an example of a double replacement

33. Reaction Types Review…
Match each chemical reaction with one of the reaction types on your chemical cartoons.
Zn + 2HCl  H2 + ZnCl2
N2 + 3H2  2NH3
2KI + Pb(NO3)2  2KNO3 + PbI2
2MgCl  Mg2 + Cl2

34. Conservation of Mass
Atoms cannot be created or destroyed in a chemical reaction.
What goes in must come out.
So we must balance equations to conserve mass.

35. Balancing Equations Rules: Ex. 2H2 + O2 --> 2H2O
We can not add or subtract subscripts from either side of the equation
We can only add coefficients to the front of each compound
Ex H O2 --> 2H2O
H = H = 4
O= O = 2
Before must match After
See “Balancing Act” worksheet for more examples…

36. Solution Chemistry
Mixtures: Matter that consists of two or more substances mixed but not chemically combined
Solutions: Homogeneous Mixture in which one substance is dissolved into another
Solute = Substance that gets dissolved (ex. Kool-Aid powder)
Solvent = Substance that does the dissolving (ex. Water)
Acid: Compound with a pH below 7 that tastes sour and is a proton donor.
Ex. Citrus foods
Base: Compound with a pH above 7 that tastes bitter and is a proton acceptor
Ex. Cleaning Products (soap)

37. Acids and Bases Solutions can be acidic or basic
Acids and Bases have unique properties when dissolved in water
Acids = sour taste
Bases = bitter taste
Indicators are substances that change color when mixed with a solution, which helps to determine if a substance is an acid or a base. (pH paper, Litmus paper, cabbage juice)

38. Acids Proton donors (H+)
Acids contain hydrogen and produce positive ions (H+) when dissolved in water
Acids = good electrolytes
Examples of acids:
Lemon Juice
Citric Acid
Carbonic Acid
HCl

39. Bases Proton acceptors
Bases contain hydroxide ions (OH-) when mixed with water.
Bases = weak electrolytes
Examples of bases:
Ammonia
Soap
Bleach (chlorine)

40. Combining Acids and Bases
-Mixing acids and bases is a balancing act.
(like a teeter-totter)
Acid + Base = neutral (water and salt)

41. Combining Acids and Bases
EXAMPLE:
Acid + Base = neutral (water and salt)
H OH-  HOH Salt
Acid Base water
Ex. HCl NaOH  H2O + NaCl

42. Measuring Acids and Bases
pH scale- used to measure the acidity of a solution.
Measure pH with indicators
pH scale goes from 0 – 14
0 = very acidic
14 = very basic
7 = neutral

43. Acids and Bases