Unit 3 Atomic Structure,
Periodic Table, and Moles
I.
Name:_________________________________
Period:___
Atomic Structure: Chapter 2.1
A. ___________Atomic Theory.
a. All matter is made of ____________
b. Atoms are ________________and cannot be divided into ___________________
c. Atoms of one element are __________________, but they are _______________
from atoms of other ________________
B. Subatomic Particles: Actually means ______________ atom.
1. Electrons: This is what makes elements____________________________.
a. Located __________________________________________.
b. Charge is ____________________________.
Indirect Evidence provided information about electrons. Diagram this experiment.
c. Electrons __________________so they can___________________________.
d. Have no ________________________
e. Exist at different______________. The number of ______________can be
found by looking at ____________________________the element
f. The only electrons that can bond with other atoms are the _________________
______________. Called ____________________________.
2. Protons: Protons make elements _______________________ because their positive
charge controls the _______ attractions of an atom, thus controlling it’s ______________
a. Located in the ______________.
b. Charge is ______________.
c. Found on periodic table by looking at the ___________________of an element.
d. Has a mass of __________________________________________
Carbon
12.01
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Example:
Potassium has how many protons? ___________
Potassium (K) has how many electrons? ______________
Potassium has how many energy levels? _____________
Diagram here:
K has how many outer electrons for bonding? _____________
3. Neutrons: This is what adds to the ______________of an atom.
a. Located in the ______________
b. Charge is ______________
d. Total mass of an atom from the _______________ and the ________________
4. Isotopes: All atoms of an element have same number of______________ . But, in
nature, some atoms of the same element have different numbers of______________,
they are called different______________. This causes their masses to be different.
Copy Fig. 2.7, pg.60
a. The average mass of all the isotopes of an element are listed on the periodic tbl.
Copy Fig 2.11, pg. 64
b. When reading the periodic table masses, round to the nearest whole number to
find the most common isotope for that element.
Copy Fig 2.12 & 2.13 , pg. 65 & 66
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Q: How many protons, neutrons, electrons are in the most common isotope of potassium?
How many protons, neutrons, electrons are in the most common isotope of chlorine?
Here is how the Proton was discovered. Diagram below. (Fig 2.9, pg 62)
Obj. 1: Fill in the grid below for each subatomic particle. (Table 2.1, pg 65)
Location in atom
Symbol
Charge
Mass
Protons
Neutrons
Electrons
Obj. 2: Directions: Complete the table for the following isotopes of each element:
Element Symbol
Number of Number of Number of Atomic
Protons
Electrons Neutrons Number
Sodium
Mass
Number
Valence
Electrons
13
47
97
Mercury
85
4
F
X
X
5
19
II. Atomic Models
A. Bohr Model: Uses ______________to show the energy levels. The number of rings
should match the ______________of that element. This model is________________________.
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1. The protons in the nucleus are found by looking at the______________.
2. The neutrons plus the protons must add up to______________.
3. The electrons fill the shells from ______________until they match the number of
______________. They fill in the following order ______________
a. You can check you outer electrons to make sure they match the _____________
b. This model can only be used for the first 20 atoms. After that it ____________
i.e. Calcium 44
i.e. lithium 7
Obj. 3: Draw Bohr Models for the following Isotopes. Include…



# of Electrons in correct orbitals
# of Protons in nucleus
# of Neutrons in nucleus
Sulfur 34
Boron 10
Helium 3
Sulfur 32
Boron 14
Helium 4
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B. Lewis Dot Structure (pg 77): Only shows valence (outer) electrons.
1. Used to show ______________between atoms.
2. Just look at the ________number. This is how many_____you draw for each_______
Try these: Write the symbol then the Lewis dots.
Potassium
Fluorine
Sulfur
Xenon
Phosphorus
Silicon
Aluminum
Beryllium
C. Quantum Mechanical Model (pg 75, 238-240): Out current model to explain the atom.
1. Electrons fly in _________of known ___________based on where they ________fly.
2. Electrons fly in ______________in each cloud.
3. We never know where electrons are, but we can know the ______________of where
they could be at any time.
4. Each pair in a cloud spins in ______________directions.
5. There are the same number of__________________, but they are not stuck in rings.
6. _____________are the tiny particles that make_______________________________.
Two examples
S shaped clouds (group 1A & 2A)
P Shaped Clouds (Groups 3A – 8A)
D. Forces in the Atom
1. Gravity: _________affect b/c particles are so small, and gravity depends on _______
2. Electromagnetic: ________________attraction b/w ________________ (magnet)
3. Strong Force: Holds the__________ together even though + + usually _________
This is the force we break when we blow up______________. Lots of_________.
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Diagram how these forces are shown in the Elegant Universe video.
Obj. 4 Identify forces within the atom.
Strong Force = SF, or Electromagnetic Force = EMF
_______This force would holds the protons together in the nucleus?
_______This is the most Powerful force within the atom?
_______This force holds the electrons near the nucleus instead of flying away?
_______Which force did we overcome in our Flame Test Lab?
II.
Reading trends in the Periodic Table (Ch 3.2)
A. ______________Go down, ______________pass across.
1. Tall groups tell you the number of ______________ electrons in that group.
2. Some common names for special groups
a. Group 1A: __________________________________________
b. Goup 2A: __________________________________________
c. Group 7A: ______________________________________________
d. Group 8A: ______________________________________________
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3. Size (pgs. 256-257): Atoms get _____________ as you go down a Group, but
___________as you go to the right b/c more ______________________ in outer energy level.
4. The ___________identifies ___________to the left, ___________ to the right, and
___________ touch the ladder
a. Short groups are ___________metals, charges may change (are in transition)
b. To remember where metals and non-metals are, just______________________.
5. The ___________will tell you the number of______________________.
III. Bonding and Reactions (Ch. 4.2): The main goal of of chemical bonding and reacting
is _________________________. This is called the ___________ rule.
A.
Two bond types:____________________ & _______________________
1. Ionic Bonds (131 – 133): When a __________ion with a __________charge sticks to
a ______________with a ____________charge. Also called salts.
a. What is an Ion: A atom that has _________ or _________electrons, thus forming
____________.
b. Metals ______their outer electrons. The protons, which are _______then show
their positive charge for that ion. Just look at A group to see how many they lose.
Na
Be
Al
N
c. Non-metals ________ electrons to _________their outer shell. These negative
electrons make the atom become a __________ ion. Just count how many to get 8.
O
Cl
d. Where do these gained and lost electrons come from, or go to?
Show sodium chloride forming (Fig 4.17).
Metal & Nonmetal
Ions in a bond
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Show aluminum oxide forming.
Metal & Nonmetal
Ions in a bond
e. To quickly find the charge of an ion, look at the ________________electrons and
count how many they ________________________.
Q: What charge will the following ions have?
1. sodium
4. magnesium
2. phosphorus
5. carbon
3. bromine
6. sulfur
f. The compound is then balanced by crossing charges, reducing if possible.
Metal Charge
Non-Metal Charge
Compound Formula balanced
Li
+
C

Mg
+
F

K
+
O

Al
+
P

2. Covalent Bonding (136-139): Non-metals _________ electrons to fill outer shells.
Remember, non-metals like to ________ electrons. Since they both want to gain, they
must ___________.
1. Covalent bonds can be modeled simply with Lewis dot structures. These show
only the valence electrons.
H2
O2
N2
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B. Covalent Bonds: Non-metals sharing electrons. Follow these rules for showing bonds.
1. Look at formula, then add up total valence electrons needed for your drawing.
2. Single bond all of the atoms, picking a center atom when possible.
3. Fill every atoms shell (remember most need 8, but Hydrogen only needs 2)
4. Count the atoms, erase some if you have too many. Then move electrons into double
or triple bonds to make every atom happy again.
Cl2
H2O
H2O2
CCl4
CH4
CO2
Ionic or Covalent? Why?
1. NaCl -
5. Al2O3 -
2. H2O2 -
6. CO2 -
3. C6H12O6 -
7. O2 -
4. MgF2 -
8. Li3P
Try these:
N2
F2O
Cl2O2
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CI4
CBr4
CO2
IV. Counting Particles of Matter (Ch 12.1 & 12.2).
A. How can we measure tiny atoms that we cannot see? Mass, counting, volume
1. Chemists use a special unit to measure atoms called the MOLE. A mole represents a
certain amount, similar to a dozen equaling 12.
 1 dozen = 12
 1 ream = 500
 1 score = 20
 1 pair = 2
 1 mole = 602,000,000,000,000,000,000,000 = 6.02 x 1023
B. Molar Masses and the Periodic Table
a. Atomic Masses on the periodic table,
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i. 1 mole of Hydrogen atoms = 6.02 x 1023 atoms of H = mass of 1.01g.
ii. Atoms with more protons have more atomic mass.
H
1.01g

Atomic Mass of hydrogen = the mass of 1 mole of hydrogen atoms.

1.
2.
3.
Try these
What is the mass of one mole of oxygen atoms?
What is the atomic mass of one mole of bromine atoms?
How many atoms is one more of oxygen?
b. Molar Mass and Compounds.
i. Molar mass is the mass of one mole of any substance
ii. An elements molar mass is its atomic number on the periodic table.
iii. Compounds can also be analyzed for molar mass using the periodic table. 10
 H2O – 1 mole of H2O has 6.02 x 10 23 molecules of H2O.
 If you have one mole of H2O, there would be two moles of Hydrogen and
one mole of oxygen since each molecule has two H atoms and one O atom.
 The molar mass of 1 mole of H2O would therefore be
Try These
1. What would be the molar mass of NaCl
2. What would be the molar mass of O2
c. Using Molar Mass as a conversion factor to show your work.
 Remember, a conversion factor is a fraction where the top and bottom are
equal amounts, so the fraction actually equals 1.
 You can put either term on top, depending on what you want as your answer.
 We know how to find the mass of 1 mole of any substance using the periodic
table. Just calculate it, and put it in your fraction with one mole
1. Mole −> Mass
a. What would be the mass of 2 moles of H2O?
b. What is the mass of 4 moles of NaCl?
c. What is the mass of 3 moles of N2 molecules?
d. What is the mass of 3 moles of N atoms?
2. Mass −> Mole
a. How many moles are in 108.12 grams of water?
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*notice that 108.12 grams is exactly six times heavier than one mole of H2O
b. How many moles are in 88.02g of CO2?
c. How many moles are in 175.32 g of NaCl?
Mixed up problems. Try these. Use the examples in your notepack for help.
1. What is the atomic mass of Krypton?
2. How many moles of krypton are in 167.60g of Kr?
3. What is the molar mass of iron oxide, Fe2O3?
4. How many grams would 5 moles of iron oxide weight?
5. How many moles are in 332.94g of CaCl2?
6. What is the molar mass of glucose sugar, C6H12O6?
7. How many grams would 10 moles of glucose have?
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