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Atomic Structure Notes
Name ____________________
I. Atom History
A. John Dalton (1766 – 1844)
1. He revisited and revised Democritus’ theory.
2. Dalton’s Theory (1803)
____a. All matter is composed of tiny indivisible particles called _________.
____b. All atoms of an element are _____________.
____c. Atoms cannot be ________, _________, or ___________.
____d. Atoms of different elements combine chemically to form ___________. Ex.
____e. Atoms of 1 element ___________________ into a new element during chemical
reactions. (Alchemists of the time were trying to change Pb into Au.)
B. JJ Thomson – 1897
1. Experiment - _____________________
a. To determine mass of single cathode ray particle.
b. Found it to be ____________ charged w/ mass less than hydrogen atom.
2. Discovered the _________________.
3. He believed the atom was a big _____________ _____________ “blob” with little
negatively charged ___________________ floating around in it.
4. Called his model the “___________________________” model.
5. Which part of Dalton’s theory did Thomson disprove?
_____________________________________
C. Ernest Rutherford – 1911
1. Experiment - ________________________
a. Thought alpha particles (+) would pass right through ______ _____.
b. Some particles actually bounced back or were greatly deflected.
2. Discovered the ___________________.
3. Concluded that the atom:
a. was mostly _____________ through which electrons move
b. contained a _______, dense, ______________ charged area (_________)
in its center
D. James Chadwick – 1932
1. Problem – Rutherford’s model couldn’t account for the total mass of atom. It was
always heavier than predicted.
2. He proposed existence of __________________.
a. has mass nearly equal to the mass of a _________________
b. carries no charge (_______________)
E. Niels Bohr – early 1900’s
1. Problem – Since oppositely charged particles attract each other, why didn’t the positive
nucleus draw the negative electrons into it, causing the atom to collapse?
2. He proposed that the electrons are arranged in 7circular _______________
around the nucleus
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II. Subatomic Particles and their Properties
A. The Pieces
Particle
Discoverer
Location
Charge
Symbol
Relative Mass
Actual Mass
B. The Whole
1. Inside nucleus
a. Composed of ________________ & __________________
b. Overall ______________ charge
c. Very dense (Extremely small % of total volume of atom, BUT 99.97% of its _________)
2. Outside nucleus
a. 99.9% of atom is this empty space through which the _____________ travel.
b. Overall __________________ charge
C. How they fit together
1. Electrons are held within the atom due to their attraction to the nucleus.
2. Neutrons help to stabilize the nucleus
a. they act as “pillows” to separate the positive protons
b. too many neutrons can make the atom radioactive
3. Since all atoms are electrically neutral, ______________ = ________________
III. How Atoms Differ:
A. Atomic Number
1. Henry Moseley (1912) – discovered that atoms of 2 different elements contain different
numbers of ____________________.
2. Atomic Number – The number of ________________ in an element
a. How many protons does a gold atom have?_______ An aluminum atom?________
b. How many electrons are in a gold atom?_________ An aluminum atom?________
B. Isotopes – Atoms of the _________ element with different number of _______________.
1. Disproved Dalton’s Theory (Dalton said all atoms of the same element were alike.)
2. All atoms of the same element MUST have same # of ______________, but not necessarily
the same # of neutrons.
3. Example – Potassium (There are 3 different kinds of K atoms.)
Potassium – 39
93.25%
Potassium – 40
6.73%
Potassium – 41
0.12%
Protons
Neutrons
Electrons
4. Two ways to write the symbol for an isotope:
potassium – 39
or
Mass #
39
19
K
Atomic #
2
5. Mass Number = ______________ + ________________
(The atomic # identifies the element, the mass number identifies the isotope.)
6. Isotope Similarities
a. Chemically and physically alike (# of electrons determines the chemical behavior)
b. Same # of ________________ and ___________________ (atoms)
7. Isotope Differences
a. Different # of ______________________
b. Different _________________
c. Different ______________ ________________
8. Examples: Fill in the chart for the atoms listed below:
Element
Atomic #
Mass # # Protons # Neutrons
# Electrons
Neon
22
Calcium
26
8
17
26
31
Symbol
64
124
Zn
80
C. Atomic Mass – The weighted average mass of the _______________ of that element.
Cl
1. Expressed in AMU’s
17
35.45
2. AMU = atomic mass unit (1amu = 1/12 the mass of a carbon – 12 atom)
a. proton = 1 amu
b. neutron = 1 amu
c. electron = 1/1840 amu
3. Usually not a whole number since it is a weighted __________________ of the isotopes.
4. Average Atomic Mass = (% abundance)(mass) + (% abundance)(mass) + ...
Isotope 1
Isotope 2
a. Ex1 There are 2 isotopes of chlorine:
35
Cl (75.77% abundance) and 37Cl (24.23%abundance)
b. Ex2 There are 2 isotopes of copper.
Copper-63 has an abundance of 69.1% and Copper-65 has an abundance of 30.9%.
Calculate copper’s average atomic mass.
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IV. Ions
A. Ion – An atom or group of atoms that have a ___________ or ____________ charge.
They are formed when an atom ____________ or ____________ electrons.
****Elements want to have a _______ _____________ _____________ (____________ __________), to
become ___________________.
1. Cation – An atom with a _____________ charge.
a. Metals tend to form positive ions by losing electrons
b. Ex: Sodium-23 atom
vs
Sodium-23 ion
2. Anion – An atom with a ________________ charge.
a. Non-metals tend to form negative ions by gaining electrons.
b. Ex: Fluorine-19 atom
vs.
Fluoride-35 ion
Draw a Bohr model of the following elements:
Helium
Boron
3. Summary:
p+ = e- (___________)
4. Ex: Fill in the chart below:
Name
p+
no
Boron
atom
Oxygen
e-
p+ > e- (_____________)
Atomic #
Mass #
Charge
p+ < e- (________)
Atom,
cation, or
anion
symbol
6
20
17
21
+2
35
-1
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Chapter 5: The Periodic Table
I: Periodic Table arrangement:

Elements ordered by ______________________________

Elements with similar ___________ fall into ______________
A. Groups or ____________
 Designated by numbers ________________
 ______________________
 Group number (ones) represents # of valence electrons – _____________________________

Similar ____________________ and ______________________________________
 Which of the following elements act like carbon?
Li, Na, Al, Si, P, Ga, Ge, Cl F
B. Periods
 Designated by numbers ______________
 _______________________________________
 Number of ____________________________
*** You are more like your _____________ than you are like your ____________
C. Blocks of Elements – designated by _____________________________
D. Zig-Zag Line

Separates ___________________ from ____________________ (label them)
Metals
 __________________________________________
 ___________________________________________
 ___________________________________________
Non-metals
 __________________________________________
 __________________________________________
 __________________________________________
Metalloids
 found along and to either side of the zig-zag line, except ________
 some properties of ____________ and some properties of ___________________
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V. Periodic Table Groups
Group 1: __________________________ -
Group 2: __________________________-
Group 3-12: _______________________ -
Group 17: _____________________ -
Group 18: _______________________ -
VI. Periodic Table Trends
A. Oxidation Number
Ion – an atom (or group of combined atoms) that __________________________ because of the
_______________________________ Ex:
Oxidation Number - the ___________ on an _____
 atoms _______ or _________ electrons to get ____
electrons in their _____________________ (octet rule –
atoms always want a ______ valence shell).
Group
#
Metals
1
2
13
14
Nonmetals
15
16
17
18
Valence
e-
Oxidation
Number
Ion
Ex.
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B. Atomic Radii –
Ex: Li / C has the larger atomic radius
 Ex: Ba / Mg has the larger atomic radius
Which element has the largest atomic radius?
C. Ionic Radii – distance from the center of the nucleus to the outer edge of the ion
 Cations - ______ electrons, less e- means protons are now pulling fewer eresulting is a __________ radius.
 Anions - _______ electrons, more e- means protons are now pulling more e- and
also more e- are repelling each other pushing them out. Results in a _______
radius.
D. Ionization Energy - the energy required to __________________________ from an atom (kJ/mol)

________________________________________

________________________________________________________________________
 Ex: Ca / Br has a higher IE
 Ex: Na / Cs has a higher IE
 Which element has the highest IE?
IE for Period 2 Elements (kJ/mol)
Period 2
Li
Be
B
520
900
800
C
1090
N
1400
O
1310
F
1680
Ne
2080
E. Electronegativity – _________________________, a measure of an atom’s ability to ___________
 Ex: Be / O has a higher electronegativity
 Ex: N / Sb has a higher electronegativity
What is the most electronegative element?
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F. Reactivity – refers to how likely an atom is ________________________________, depends on ...
- how easily electrons can be __________________________
- how easily electrons can be __________________________
What is the most reactive metal?
What is the most reactive nonmetal?
Recap:
a) Largest Atomic Radii =
b) Largest Ionic Radii =
Smallest Ionic Radii =
c) Highest Ionization Energy =
d) Highest Electronegativity =
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