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Chemistry Unit Notes
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
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
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
inside nucleus 1 amu
Periodic Table
Using the Periodic Table
17

Atomic Number
– Equal to # protons = # electrons
– Periodic Table is arranged by this number
Cl
35.5

Symbol
– “Shorthand” for the element – Note 2nd letter is
always lowercase

Atomic Mass Number
– Total AVERAGE mass of Protons + Neutrons +
Electrons
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
Shell
Shell
Shell
Shell
Shell
#1
#2
#3
#4
#5
#6
2 electrons
8 electrons
8 electrons
18 electrons
32 electrons
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??
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
36 - 17 = 18 neutrons
(Round Atomic Mass)
Notice: electrons follow energy level rules
from previous slide.
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
More Practice

Determine the name, number of protons,
neutrons and electrons for each element
shown and draw…
15
8
26
P
O
Fe
31
16
56
Isotopes

An isotope is a variation of an element
(same protons) but can have diff. # of
neutrons

Ex: carbon (atomic mass = 12.011)
– Carbon (14) and carbon (12) exist in
nature
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 (+)
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?
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.


http://www.fordhamprep.org/gcurran/sho/sho/le
ssons/lesson38.htm
Types of Chemical Bonds

Ionic-

Covalent-

Metallic-
Two elements bond by transferring electrons to create ions
that attract together (+ is attracted to - after an electron is transferred)
bond type)
Two elements bond by sharing electrons (strongest
Two metals bond and form a “common electron cloud”.
This is a cluster of shared electrons (weakest bond type)
Examples of Bonding
http://www.youtube.com/watch?v=xTx_DWboEVs
http://www.youtube.com/watch?v=1wpDicW_MQQ
http://www.youtube.com/watch?v=QqjcCvzWwww
http://lc.brooklyn.cuny.edu/smarttutor/core3_22/Bonds.html
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
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-
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-
Oxidation Numbers

Each column going down the periodic table
has elements with the same oxidation
number.

Label the oxidation numbers on your periodic table at the top of each
column as shown here:
1+ 2+
3+ 4(+/-) 3- 2- 1- 0
Rules for using oxidation
numbers to create compounds
1. 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.
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 2- = 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 6Since these are now equal to zero, we combine 2 Aluminums and 3 Sulfers to make Al 2S3
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
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
Chemical Reaction Basics
H2 + O2 --> H2O
Reactants
Products
Reactants- substance that enters into a reaction
Products- substance that is produced by a chemical reaction
Evidence of Chemical Change
 EPOCH is an acronym that stands for evidence that a
chemical reaction has occurred.
E–
P –
O–
C –
H –

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?
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
Cartoon Chemistry
This is an example of synthesis
Cartoon Chemistry
This is an example of a decomposition
Cartoon Chemistry
This is an example of a single replacement
Cartoon Chemistry
This is an example of a double replacement
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
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.

Balancing Equations

Rules:
– 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.
2H2 + O2 --> 2H2O
H=4
O=2
Before
H=4
O=2
must match
After
See “Balancing Act” worksheet for more examples…
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)
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)
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
Bases
Proton acceptors
 Bases contain hydroxide ions (OH-) when
mixed with water.
 Bases = weak electrolytes
 Examples of bases:

– Ammonia
– Soap
– Bleach (chlorine)
Combining Acids and Bases
-Mixing acids and bases is a balancing act.
(like a teeter-totter)
Acid + Base = neutral (water and salt)
Combining Acids and Bases
EXAMPLE:
Acid + Base = neutral (water and salt)
H+
Acid
+ OH-  HOH + Salt
Base
water
Ex. HCl +
NaOH

H2O + NaCl
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

Acids and Bases