Chemistry I – Unit 2 Notes
Te x tbook Correlation
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Unit 2 Consists of the following sections in the text:
The Elements
A . An element is:
B . Elements can exist as pure substances or as parts of compounds.
Examples:
Symbols for the Elements
Chemical symbols are:
* Be forewarned, the one or two letter chemical symbol for many elements is not the same as the first one or two letters in the element name.
Examples:
Dalton’s Atomic Theory
John Dalton (1766-1844) – English Scientist page
Dalton’s Atomic Theory
1. Elements are made of tiny particles called atoms.
2. All atoms of a given element are identical (we now know that this is not exactly true)
3. The atoms of a given element are different from those of any other elements.
4. Atoms of one element can combine with atoms of other elements to form compounds.
A given compound always has the same relative numbers and types of atoms.
Examples:
5. Atoms are indivisible in chemical processes. That is, atoms are not created or destroyed in chemical reactions. A chemical reaction simply changes the way the atoms are grouped together.

The Structure of the Atom
A. History of the Structure of the Atom
1. J. J. Thomson’s Plum Pudding Model (1890’s) described atoms as being:
2. Ernest Rutherford’s “Nuclear” Model (1911) of the atom described the structure of atoms as follows: a. Gold Foil Experiment

Introduction to the Modern Concept of Atomic Structure
A. From Rutherford’s gold foil experiment we learned that.
1. Atoms have a central positive nuclear charge
2. Atoms are 99.99% empty space.
B. Subatomic Particles
Mass Charge Particle
Electron
Proton
Neutron
“Box” on the Periodic Table:
Location in the Atom

Mass Number = Number of Protons + Number of Neutrons
Also, for a neutrally charged atom,
Atomic Number = Number of Protons = Number of Electrons,
*the atomic number determines the identity of an element
Use the equations above and your periodic table to fill in the missing values in the table below:
Element Atomic # Mass # # of Protons # of Electrons # of Neutrons
H 1 1
S
8 16
16
9 19
39 19
Isotopes
A. Atoms of the same element always have the same number of ___________________.
*the number of protons an atom of a certain element contains is given by its
_____________________________________.
B. Isotopes are:
C. The sum of the number of protons and the number of neutrons contain in an atom is know as the atom’s ______________________________________.
D. Isotope Notation:
1. Isotopes are often are symbolized by an element symbol with superscripts and
Subscripts denoting the isotope’s mass number and atomic number.
Examples:
2. Isotopes are can also be written as the element’s name, a dash, and the mass number of the isotope.
Examples:

Silicon-28
Silicon-29
Silicon-30
Example 2
Isotope name
Iron-54
Iron-56
Iron-57
Iron-58
Weighted Average Atomic Mass
A. What is a weighted average? Why do we need one?
An element can exist in a number of forms, called isotopes. Isotopes are forms of the same atom that vary in mass.
For example, there are two different types (isotopes) of copper atoms. One type of copper atoms weighs in at 62.93 amu, the other has a mass of 64.94 amu. The lighter isotope is more common with 69.09% of the naturally occurring copper having a mass of 62.93 amu per atom. The remainder of the atoms, 30.91 %, have a mass of 64.94 amu.
To find the Average Atomic Mass of an atom, we take into account all of the isotopes that exist and the percentage of each type. The calculation of the average atomic mass is weighted average .
B. How to calculate a weighted average atomic mass
Take the sum of the products of each isotope’s mass and its corresponding relative abundance as a decimal (take percent and move decimal 2 places left).
Example 1
Isotope name Isotope mass (amu) Relative Abundance
27.98
28.98
29.97
92.21
4.70
3.09
Isotope abundance
5.90%
91.72%
2.10%
0.280%
Isotope mass (amu)
53.94
55.93
56.94
57.93

Rutherford’s Atom
A. His model leaves many questions about electrons unanswered
1. How are the electrons arranged?
2. How do they move?
3. Since the nucleus and the electrons are oppositely charged, why doesn’t the atom collapse?
Electromagnetic Radiation – energy transmitted as a wave
A. Parts of a wave
1. Wavelength is –
2. Frequency is –
B. Types of EM Radiation and wavelength.
Emission of Energy by Atoms
A.
Electrons surrounding an atom can absorb a discrete packet of energy called a
___________________________ to become “excited”.

B. When excited electrons loose that extra energy they fall back into their
_______________________. The release of energy by the electron results in emission of a photon of a certain wavelength. Each element has its own unique spectrum of wavelengths that are released.
Examples:
The Energy Levels of Hydrogen
A. Electrons can only absorb quantized amounts of energy…this means
___________________________________________________________
B. With only one electron, a hydrogen atom is the simplest way to view what can happen when electrons get excited.
C. How can Hydrogen produce photons with 4 different energies (colors) when it only has 1 electron to excite? (See Figure on pg. )

Electronic Transitions in the Bohr Model for the Hydrogen Atom
The Bohr Model of the Atom
A. Bohr’s model of the atom included the following main points.
1. Central nucleus made up of ______________ and _______________.
2. Electrons were restricted to circular orbits.
The Wave Mechanical Model of the Atom (Quantum Mechanics Model)
A.
This model of the atom is our most current and up to date model.
B. The biggest difference between this model and Bohr’s model is “Orbits vs. Orbitals”.
1. An Orbit is –
2. An Orbital is –

(a) The Probability Distribution for the Hydrogen 1s Orbital in Three-Dimensional Space
(b) The Probability of Find the Electron at Points Along a Line Drawn From the Nucleus
Outward in Any Direction for the Hydrogen 1s Orbital
The Hydrogen Orbitals
A. Within each principal energy level there can be one or more orbitals.
1. “s” orbitals:
2. “p” orbitals:
3. “d” orbitals:
B. Each principle energy level is a little larger and further away from the nucleus than the last and contains more orbitals than the last.

Two representations of the Hydrogen 1s, 2s, and 3s Orbitals (a) The Electron Probability
Distribution (b) The Surface Contains 90% of the Total Electron Probability (the Size of the Oribital, by Definition)
Organization of Principle Energy Levels, Orbitals, and Sublevels.
A. Principle energy levels contain orbitals , which in turn contain sublevels .
B. Pauli exclusion principle:
The Orders of the Energies of the Orbitals in the First Three Levels of Polyelectronic
Atoms

Electron Arrangements in the First 18 Atoms on the Periodic Table
A. Electron Arrangement = Electron Configuration
B. Box Diagram
1. Principal Energy Level
2. Type of orbital
3. Hund’s Rule
4. Valence Electrons
5. Core Electrons