Download Atom Notes Outline - Sewanhaka Central High School District

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ATOMIC STRUCTURE
A. ATOMIC THEORY
Due to contributions of numerous scientists
1. John Dalton: basic unit of matter - tiny particle called the atom.
Dalton’s Theory
A. All matter is composed of extremely small particles called atoms, which cannot be
subdivided, created, or destroyed.
B. All atoms of the same element are identical
C. Atoms of different elements are different
D. In chemical reactions, atoms can be combined to form compounds.
E. In chemical reactions, atoms can be combined, separated, or rearranged, but
never created, destroyed or changed.
Summary:
 Most of Dalton’s Theory was accurate, but some was incorrect to some degree.
2. J.J. Thomson: atom was made up of EVEN SMALLER PARTS
Summary:
 So: If electrons are present in atoms, then what makes up the rest of the atom?
3. Ernest Rutherford: Sought to answer this question.
-Rutherford’s Gold-Foil Experiment (KNOW!!!)
What did Rutherford’s experiment show? (KNOW!!)
1.
2.
4. Neils Bohr : Proposed a model of the atom showing a central dense nucleus with electrons found
in surrounding orbits.
5. The present day model of an atom –-> The WAVE MECHANICAL MODEL
-Central Nucleus, with clouds of electrons around the nucleus.
-Electrons are found in surrounding orbitals
-Orbital:
*Modern model of the atom has evolved over a long period of time through the work of many scientists.
Some other scientists:
1) Planck –said that atoms absorb energy in discrete amounts called quanta
2) Chadwick –Discovered neutrons in the nucleus of an atom
3) Moseley –Coined the term Atomic Number
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B. THE ATOM

Protons, neutrons, and electrons are all inside the atom, called: ___________________
SUBATOMIC PARTICLE COMPARISON (MEMORIZE!)
LOCATION
CHARGE
SYMBOL
MASS
PROTON
NEUTRON
ELECTRON
C. ATOMIC NUMBER:
ATOMIC # = _______________________ (in the nucleus of an atom)
-In a neutral atom the # of protons = the # of electrons
**Atomic # is indicated in bold in the bottom left corner of each elements’ box on The Periodic
Table**
The atomic # for Carbon (C) is ____________
How many protons does a neutral C atom have? _________
How many electrons? ___________
Atomic #
Ex. 2: The element Lithium (Li) has an atomic # of _______
How many protons does a neutral Li atom have? _________
How many electrons? _________
D. MASS NUMBER:
MASS # =
___________________
+ ______________________
# OF NEUTRONS = _______________ - ___________________
Standard Example: Mass #
40
Element Symbol
X
20
Atomic #
# of neutrons = ___________
3
Additional Examples:
1. An atom of Lithium (Li) has a mass # of 7.
Find the following:
Atomic #: _______________
# of protons: ____________
# of neutrons: ____________
# of electrons: ___________
2. Find the following with regard to 39K:
Mass #: __________________
Atomic #: ________________
# of protons: _____________
# of neutrons: ____________
3. Find the following with regard to Nitrogen-14:
Mass #: _______________
Atomic #: ______________
# of protons: ____________
# of neutrons: ____________
# of electrons: _____________
4. An atom has 15 protons and 16 neutrons.
What is its atomic #? _______________
What is its mass # ? ________________
What element is this? _______________
# of electrons: ____________
E. ISOTOPES AND ATOMIC MASS

ISOTOPE - atoms of same element that have SAME _________ but DIFFERENT __________



Isotopes of atoms can be stable or unstable.
Stability of an isotope is based on the # of protons and neutrons in its nucleus.
If the nucleus is unstable, the isotope may spontaneously decay and emit RADIATION.
EXAMPLES:
Hydrogen has 3 isotopes:
Carbon has 2 isotopes:

ATOMIC MASS: WEIGHTED AVERAGE MASS of all the NATURALLY OCCURING
ISOTOPES of that element.
o Represented by a decimal # at the top left hand corner of each element box.
Atomic Mass
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F. CALCULATING AVERAGE ATOMIC MASS
Example: Chlorine has two naturally occurring isotopes – Cl-35 and Cl-37. 76% of the isotopes of Cl are
Cl-35, and 24% of its isotopes are Cl-37. Calculate the average atomic mass of chlorine.
Step
Calculation
Multiply each % by the appropriate mass #
76 x 35 = 2660
24 x 37 = 888
2660 + 888 = 3548
3548 / 100 = 35.48
Add these answers together
Divide by 100
Examples:
Data in various chemical handbooks indicate the relative percentages of these naturally occurring isotopes of
Magnesium (Mg): 24Mg = 79.3% , 25Mg = 10.1%, 26Mg = 10.6%. Given this data, calculate the average
atomic mass for Magnesium.
G. PRINCIPAL ENERGY LEVELS
 Show how far away the electrons are from the nucleus
 There are up to 7 PEL’s that surround the nucleus
 Electrons in _______________________________ have __________________________
 Low PEL # = ___________________________ = ____________________________
 High PEL # = ___________________________ = ____________________________

Diagram of Atom and PEL’s surrounding nucleus:

How many electrons can each energy level hold?
Principle
Max # of electrons
Energy Level (n)
level can hold (2n2)
1
_____electrons
2
_____ electrons
3
_____ electrons
4
_____electrons
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H. ELECTRON CONFIGURATIONS AND VALENCE ELECTRONS

ELECTRON CONFIGURATION:
o Can be found below the atomic # in the bottom left hand corner of each element box.
Electron Configuration
o Examples:
1. Magnesium (Mg): Atomic #: ___________

2. Zinc (Zn): Atomic #: ____________
# of electrons: ________
# of electrons: __________
Electron Config: ____________
Electron Config: ______________
# of PEL’s occupied: ______________
# of PEL’s occupied: __________
VALENCE ELECTRONS:
o participate in chemical reactions with other atoms
o last # in electron configuration
 electron configuration on periodic table shows # of valence electrons a particular
element has
o Examples:


1. How many valence electrons does Mg have?
______

2. How many valence electrons does Zn have?
______
NON-VALENCE ELECTRONS: All the other inner electrons.
I. ELECTRON DOT STRUCTURES (AKA. LEWIS DOT STRUCTURES)
 A diagram that depicts valence electrons ________ around the symbol of the element
 Method:
X

Draw the electron dot structures for the following elements:
1. Helium
3. Sulfur
5. Boron
7. Nitrogen
9. Neon
2. Lithium
4. Calcium
6. Silicon
8. Chlorine
10. Iodine
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J. GROUND STATE VS. EXCITED STATE
 GROUND STATE: low-energy electron configuration, e- fill in from inner PELs to outer PELs
o Periodic table shows ground state e- config
 EXCITED STATE: An atom can become excited when its electrons move to a higher energy level
without fully filling up the lower energy level.
o How do they do this? ___________________________________________________
o Excited electrons move back down to ground state by ___________________ energy.
o Energy released called: ____________________
Examples:
•
•
•
•
Fluorine - ground state e- config (Periodic Table) 2-7
Possible excited state e- configs
1-7-1
2-6-1
1-8
An inner level is left partially filled b/c an e- has jumped to a higher PEL
TOTAL # OF e- REMAINS THE SAME!
1. Which is the electron configuration of an atom in the excited state?
A) 2 – 8 – 2
B) 2 – 7 - 1
C) 2 – 8 – 1
2. Which electron configuration represents an atom in the excited state?
A) 1 – 2
B) 2 – 3
C) 2 – 1
D) 2 – 8 – 3
D) 2 - 7
 BRIGHT LINE SPECTRUM:
Ex: The bright-line spectrum of lithium looks
like this:
So: If you have an unknown element and you want
to see if it is lithium, compare the bright-line
spectrum of the sample you have with the bright line
spectrum of lithium.

FLAME TESTS:
Which sample is lithium?
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K. IONS

IONS ARE _____________________________

Atoms are called ions when they are no longer neutral atoms

Atoms become ions by: __________________________________________

If an atom GAINS ELECTRONS its charge becomes ___________________

If an atom LOSES ELECTRONS its charge becomes ___________________

The charge of an ion is indicated as a superscript! (If there is no superscript, it is a neutral atom.)
How many protons does N+3 have? _________
How many neutrons does N+3 have? ________
How many electrons does N+3 have? ________