Download Atomic Structure

Atoms and the
Periodic Table
Atomic Models
 Democritus (4th century B.C.) first
theorized that matter was made of
particles he called the atom
Atomic Models
 Plum Pudding Model 
(1904) developed
by J.J. Thomson
 Planetary Model 
developed by
Ernest Rutherford
(1911)
Newer Models
 Bohr’s Model (1913)
Developed by
Niels Bohr
 Electron Cloud
Model (1925) 
Protons (p+)
 Positively charged particles found
in nucleus of atom
 Have an electrical charge of +1
 Mass of 1 a.m.u.
 Composed of quarks
 Discovered by Ernest Rutherford
using Gold Foil Experiment
Protons
 The number of protons in an atom
determines its identity
 All oxygen atoms have 8 protons, all
uranium atoms have 92 protons
 If the number of protons change the
identity of the atom changes.
Neutrons (n0)
 Neutral particles found in nucleus
of atom
 Have no electrical charge
 Mass of 1 a.m.u.
 Composed of quarks
 Discovered by James Chadwick
Nucleus
 The nucleus is the positively charged
dense core in the center of the atom
 Houses protons and neutrons
 Contains 99.9% of mass of atom
 Extremely small compared to the
entire size of the atom
Electrons (e-)
 Negatively charged particles found in
electron cloud
 Have an electrical charge of -1
 Constantly moving around outside
nucleus
 Have essentially no mass
 Discovered by J.J. Thomson during
cathode ray experiment
Electrons
 The number of electrons in a neutral
atom is equal to the number of
protons
 Neutral oxygen has 8 protons,
therefore it has 8 electrons
 Neutral lead has 82 protons, therefore,
it has 82 electrons
Valence Electrons
 Electrons in the outermost energy
level of an atom are called valence
electrons
 These are the electrons furthest from
the nucleus
Symbols
 Elements are identified by their
chemical symbols
 Symbols are usually either one capital
letter like C for Carbon, or one capital
and one lowercase letter like Ne for
Neon
Atomic Number (Z)
Whole number shown on periodic table
 Periodic table is arranged by atomic
number
Atomic Number = # of Protons
*Also gives the number of
electrons if the atom is neutral
Atomic Number
Mass Number (A)
 The mass number is the sum of the total
number of protons and neutrons in the
atom
 Mass # = # p+ + # n0
 The mass number is not found
on the Periodic Table
Isotopes
 Isotopes are atoms of the same
element that have different numbers
of neutrons
 All atoms are isotopes
 Each element has isotopes that are
more common than others
Nuclear Symbol
 Isotopes can be designated with their
nuclear symbol
Hyphen Notation
 Isotopes can also be designated using
hyphen notation
Carbon-16
Element Name
Mass Number
Write the Nuclear Symbol and Hyphen
Notation for the Following Isotopes
 Lithium isotope with 3 protons and 4
neutrons
 Sulfur isotope with 17 neutrons
 Lead with 122 neutrons
Ions
 Ions are atoms or groups of atoms
that have a net positive or negative
charge
 The charge results from an unequal
number of electrons and protons
within an atom or group of atoms
Ions
 Anions Ions with more electrons than
protons resulting in a negative charge
 For each extra electron the negative charge
increases by one
 Cations Ions with less electrons than
protons resulting in a positive charge
 For each missing electron the positive charge
increases by one
Ions
 Ions are symbolized with a positive or
negative sign on the upper right side
and number equal to the magnitude of
the charge
 The number one is not included
O
2-
Magnitude
Charge
Ions
 F  9 protons – 10 electrons = -1 charge
 Ca2+  20 protons – 18 electrons = +2
charge
 P3-  15 protons – 18 electrons = -3
charge
Common Ions (Need to
memorize these)






Column 1 (Li, Na, K, Rb, Cs): +1
Column 2 (Be, Mg, Ca, Sr, Ba): +2
Column 13 (Al, Ga): +3
Column 15 (N, P, As): -3
Column 16 (O, S, Se, Te): -2
Column 17 (F, Cl, Br, I): -1
Average Atomic Mass
 The weighted average of
the naturally occurring
isotopes of an element.
 Found by averaging the
natural abundances of
its isotopes
Calculating Average
Atomic Mass (amu)
If abundance is given as percent
value:
(Mass of Isotope)(% abundance) 

Atomic Mass 
100
 If abundance is given as decimal
value:
Atomic Mass  (Mass of Isotope)(a bundance) 
Average Atomic Mass
Rubidium has two common isotopes, Rb-85 and
Rb-87. If the abundance of 85Rb is 72.2% and
the abundance of 87Rb is 27.8%, what is the
average atomic mass of rubidium?
[(85)(72.2)]  [(87)( 27.8)]

 85.5  86 amu
100
Uranium has three common isotopes. If the
abundance of 234U is 0.0001, the
abundance of 235U is 0.0071, and the
abundance of 238U is .9928, what is the
average atomic mass of uranium?
[( 234  .0001)  (235  .0071)  (238  .9928)]  238.97  238 amu
Columbic Attraction
 There exists an attraction between
oppositely charged particles
 The greater the distance between the
particles the weaker the attraction
Columbic Attraction in the
Atom
 The electrons in the atom are attracted to
the protons
 Electrons closest to the nucleus feel a
stronger attraction force than electrons on
the outermost energy level
 As you move in a row from left to right on
the Periodic Table the number of protons in
an atom increases and so the attractive
force on the outermost electrons increases
Columbic Attraction in the
Atom
 As you move down a column on the
periodic table the distance between
the outermost electrons and the
nucleus is the dominant factor
determining the attractive force
Energy Levels and SubLevels
 Niels Bohr found that electrons
occupy distinct energy levels within
the atom. Ex: 1, 2, 3, 4
 It was later found the electrons also
occupy sublevels within each energy
level. Ex: s, p, d, f
Ground State Orbital
Diagrams
 Represent the electrons within energy
levels and sublevels using arrows
Ground State Orbital
Diagrams
 Pauli Exclusion Principle
 Each orbital can hold TWO electrons
with opposite spins.
Sub levels
Each energy level has a
different amount of
sublevels
1s
2  s, p
3 s, p, d
4 s, p, d, f
5 s, p, d, f
6 s, p, d
7 s, p
 s has 1 orbital
 2 e p has 3 orbitals
 6 e d has 5 orbitals
 10 e f has 7 orbitals
 14 e-
Ground State Orbital
Diagrams
 Aufbau Principle
 Electrons fill the lowest energy
level first.
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d,
5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
Ground State Orbital
Diagrams
 Hund’s Rule
 Within a sublevel, place one e- per orbital
before pairing them.
 The last energy level being filled may not be
completely filled with arrows
WRONG
RIGHT
B. Notation
 Fluorine
 Germanium
Ground State Electron
Configuration
 An electron configuration is a shorthand
description of how electrons are arranged
around the nucleus of an atom.
 Electrons within each energy level and sub
level are represented with numbers in super
script rather than arrows
Ground State Electron
Configuration
 Longhand Configuration
S 16e- 1s2 2s2 2p6 3s2 3p4
Core Electrons
Valence Electrons
 Write the orbtial diagrams and electron
configurations for the following:
 Carbon
 Sulfur
 Strontium
 Iron
Valence Electrons
 Column 1 1
valence e Columns 2-12  2
valence e Column 13  3
valence e Column 14  4
valence e-
 Column 15  5
valence e Column 16  6
valence e Column 17  7
valence e Column 18  8
valence e-
Lewis Symbols
 Also called electron dot diagrams
 Dots represent the valence e Ex: Sodium
 Ex: Chlorine
Lewis Diagram
for Oxygen
Steps to Draw Lewis
Structures
Steps to Draw Lewis
Symbols
1.
Determine how many valence are in the element.
2.
Starting on the Right side of the element draw a
dot to represent a valence electron.
3.
Place one dot on each side of the symbol. One
electron must be drawn around each side of the
element before a second electron can be added
to any side.