Elements, Compounds, Mixtures and Atomic Structure
CHEM I – NOTES
Name__________________ Period____
I. Elements, Compounds, and Mixtures
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The science of matter
o How substances interact with each other
o Physical and Chemical properties of substances
What is Matter?
o Anything that has mass and takes up space (has volume)
 EVERYTHING is made of matter.
 If it doesn’t have mass and volume, it’s not there
A. Mixtures
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Mixture- Two or more substances, combined in varying proportions - each retaining its own specific
properties.
The components of a mixture can be separated by physical means, i.e. without the making and breaking
of chemical bonds.
Examples: Air, Milk, wood, salt water, ink, Soda, and Concrete.
1. Heterogeneous Mixture
 Mixture in which the properties and composition are NOT UNIFORM throughout the sample.
Examples: Milk, wood and concrete.
2. Homogeneous Mixture
 Mixture in which the properties and composition ARE UNIFORM throughout the sample.
 Liquid mixtures are termed solutions.
Examples: Air and table salt thoroughly dissolved in water
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B. Pure Substance
A substance with constant composition
Can be classified as either an element or as a compound
Examples: Table salt (sodium chloride, NaCl ), sugar (sucrose, C12H22O11, water (H2O), iron (Fe), copper
(Cu), and oxygen (O2)
1. Element
 A substance that cannot be separated into two or more substances by ordinary chemical (or
physical) means.
 We use the term ordinary chemical means to exclude nuclear reactions. Elements are composed of
only one kind of atom.
 Elements are found on the periodic table.
Examples: Iron (Fe), copper (Cu), and oxygen (O2).
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2. Compounds
 A substance that contains two or more elements, in definite proportion by weight.
 Compounds are composed of more than one kind of atom bonded together.
 The term molecule is often used for the smallest unit of a compound that still retains all of the
properties of the compound.
Examples: Table salt (sodium chloride, NaCl),Sugar(sucrose, C12H22O11), and water (H2O).
C. Physical & Chemical Properties of Matter
1.
Physical:
 Physical properties can be observed or measured without changing the composition of matter.
 Physical properties include appearance, texture, color, odor, melting point, boiling point, density,
solubility, polarity and many others
2. Chemical:
 Chemical properties of matter describes its "Potential" to undergo some chemical change or reaction by
virtue of its composition. What elements, electrons, and bonding are present to give the potential for
chemical change.
 EX: hydrogen has the potential to ignite and explode given the right conditions. This is a chemical
property. Metals in general have they chemical property of reacting with an acid. Zinc reacts with
hydrochloric acid to produce hydrogen gas. This is a chemical property.
D. Physical/Chemical Changes of Matter
1.
Physical:
 Changes in matter that do not alter the matter itself.
 EX: Size, shape, and phase: Solid---liquid---gas (Freezing, Melting, boiling)
2. Chemical:
 Changes that do alter the identity of a substance
 EX: Iron Rusting: 4Fe(s)+3O2(g) 2Fe2O3
 Wood burning and copper turning to brass
3. Observing Chemical Changes
 Watch for the following to establish that a chemical rx has taken place:
o Precipitate (solid formed from solutions).
o Emitted gas.
o Color change
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o
Energy Change (hotter or colder)
 *All chemical RXS have a temp change *
II. Atomic Theory
A. History

Democritus, 400 B.C. proposed that the world was made up of two things:
o Empty space
o Small particles he called atoms (which is tiny indivisible particles)
o His Views were not supported

Isaac Newton and Robert Boyle in the 17th century published articles stating their belief in atomic
structure.
John Dalton in the early 1800s offered the first logical quantitative explanation of atomic structure.
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B. Dalton’s Atomic Theory
1.
All matter is composed of tiny particles called Atoms. These particles could not be broken down into
smaller substances.
2. Atoms of an element were exactly alike and atoms of different elements were unalike.
3. Atoms combine in simple ratios to form compounds. (law of definite and multiple proportions)
4. Atoms cannot be destroyed (only rearranged during a chemical reaction)
C. Research and Revisions of Dalton’s Atomic Theory
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JJ Thompson proposed the Plum Pudding Model in which the negative electrons were held in place by a
random scattering of positive charges.
Thomson work with the cathode ray tube led to the discovery that cathode rays consisted of electrons.
By exposing the ray to a magnetic field and measuring the bending of the ray, he was able to calculate
the ratio of an electron’s charge to its mass.
High voltage electricity is passed into the cathode (negative end). A ray is generated toward the anode
(positive end). When a magnet is placed near the ray, the negative end of magnet would cause the ray to
bend in the opposite direction and the positive end would bend the ray towards the magnet.
Using data from Thompson Robert Millikan obtain the first accurate measurement of the electron
charge
He used data from Thomson, and was able to calculate mass of an electron
Rutherford predicted the presence of neutrons in the nucleus.
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Rutherford's Gold Foil experiment
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Most of the alpha particles are passing through the foil. A few are slightly deflected while even fewer
are greatly deflected.
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Results of Gold Foil Experiment
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o
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1. The atom contains a tiny dense center called the nucleus.
the volume is about 1/10 trillionth the volume of the atom
2. The nucleus is essentially the entire mass of the atom
3. The nucleus is positively charged
the amount of positive charge of the nucleus balances the negative charge of the electrons
4. The electrons move around in the empty space of the atom surrounding the nucleus
D. Early Laws
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Law of conservation of Mass = matter cannot be created nor destroyed; only chemically altered.
A+B=AB.
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Law of Definite Proportions = specific substances always contain elements in the same ratio by mass
Example, mass of sodium to the mass of chlorine is salt is always the same.
Law of Multiple Proportions = ratio of masses of one element that combine with a constant mass of
another element can be expressed in small whole numbers.
o Ex: Tin (11) Oxide = SnO =1:1 Ratio
Water = H2o = 2:1 Ratio
III. Atomic Structure
A. Atomic Structure
Atoms are divisible:
 1. Electrons = Negatively charged Particles
(Millikan and Thomson). Smallest of the subatomic particles (e-) Found on the outside of the central
mass (nucleus)
 2. Protons = Positively charged subatomic particles (Thomson and Rutherford) Slightly smaller in mass
than the neutron. Found in the dense central mass called nucleus.
 3. Neutrons = Neutral, no charge
(Rutherford). The largest of the subatomic particles. Found in the dense central mass called the
nucleus.
Atomic Masses of Subatomic Particles
o Electrons = 0.000549 amu
o Protons = 1.0073 amu
o Neutrons = 1.0087 amu
Where can we find information on protons, neutrons & electrons of atoms? THE PERIODIC TABLE
Symbols are used to represent elements in
the periodic table.
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They are 1 or 2 letter abbreviations
Capitalize the first letter only
 Examples:
C carbon
N nitrogen
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CO cobalt
CA calcium
The number of protons in the nucleus is known as the Atomic number (Z)
B. Isotopes
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Isotopes of an element always have the same number of protons (atomic number) but differ in the
number of neutrons
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Isotopes of an element will have different number of neutrons
Isotopes of an element will have different mass numbers due to more mass of the increased number of
neutrons
o
Example: Isotopes of carbon = C-12 has 6 p+ and 6 N0
o
C- 14 has 6 p+ and 8 NO
C. To Calculate P+, E-, and NO
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Ex:
Atomic numbers and protons always equal one another.
Number of protons equals the number of electron’s (for now).
If you subtract the number of protons from the mass number = number of neutrons
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P+ = 17
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E- = 17
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No = 35-17= 18
D. Atomic Mass & Mass Number
Atomic Mass
 Mass of protons and neutrons
 If in a decimal form is an average of all isotopes
o EX: K = 39.0983
1.
I
Atomic Mass
126.90
Mass Number
 Mass of protons and neutrons
 Always expressed as a whole number
 Can obtain by rounding atomic mass
o EX: K = 39
Mass Number
127
2. Kr
83.80
84
3. Ca
40.08
40
4. Na
22.98
23
5. Fe
55.85
56
Determining the # of Electrons
o Protons are Positively Charged, Electrons are Negatively charged
o Electroneutrality means that an atom has equal number of positive and negative charges. So…
Protons = Electrons
o Atoms are electrically neutral… but they can gain or lose electrons to become ions
Calculating Electrons - Ions
o Ions- an atom or group of atoms that has an electric charge because it has lost or gained electrons
o Cation – an ion that has a positive charge
o More protons than electrons (gave away on electron) (Example: Li+)
o Anion – an ion that has a negative charge
o More electrons than protons (accepted an electron) (Example: As-3)
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Review
How to calculate subatomic particles:
P = atomic number
N = Mass number – P
E = Protons ± charge
Total # of subatomic particles = P+N+E
D. Average Atomic Mass
o Atomic mass is the weighted average mass of all the atomic masses of the isotopes of that atom.
o Calculating Average Atomic Mass - About 75.5 % of the chlorine found in nature is Cl -35 (17
protons, 18 neutrons) and about 24.5% of the chlorine found in nature is Cl-37 (17 protons, 20
neutrons).
E. Atomic Models
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Old version = Bohr’s
Also known as the planetary atomic model
Describes electron paths as perfect orbits with definite diameters
Good for a visual
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New version = Quantum Theory
Most accepted
Diagrams electrons of a atom based on probability of location at any one time
Bohr’s model:
 Nucleus is in the center of an atom (like the sun) and the electrons orbit the nucleus similar to the
planets.
 Orbits are called shells
o 1st shell = 2 electrons
o 2nd shell = 8 electrons
o 3rd shell =18 electrons
o 4th shell = 32 electrons
Diagram Magnesium’s atomic structure
IV. Quantum Theory
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To better the description of the atomic structure, atoms were exposed to energy (Heat) which made
the electrons go into what is called the excited state (normal = ground state).
When electrons returned to ground state they emitted energy in the form of light.
This method of study is called spectroscopy (spectrum)
Visible light = part of the electromagnetic spectrum between 400-700 nm
Electromagnetic Spectrum
o From crest to crest = frequency which is measured in hertz.
o This therefore can be used to identify elements (Absorption of energy and color emitted is a
fingerprint of an element)
V. Electron Configuration
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Electrons (e-) of atoms are the basis for every chemical reaction.
In quantum theory, electrons exist in orbitals based on probabilities and these orbitals are
arranged within energy levels.
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Quantum numbers specify the properties of atomic orbitals and the properties of electrons in
those orbitals
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Example of Quantum #: 3S²
We will define these numbers & letters.
A. Principle Quantum Number
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Is equal to the number of energy level (N).
The principle quantum # corresponds to the energy levels 1-7 which is the period number (Row) on the
periodic table.
Maximum e- in Energy Levels
o The greatest number of e- in any one level is 2n²
o Calculate the maximum number of electrons that can occupy the 4 th principal quantum number
(period 4).
o Solve: Use 2n²
2(4)²
32 electrons total
B. Sublevels and Orbitals
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An energy level in made up of many energy states called sublevels .
The number of sublevels for each energy level is equal to the value of the principle quantum
EX: One sublevel in energy level one (Period 1)
o Two sublevels in energy level two (Period 2)
o three sublevels in level three (Period 3)
Sublevels
o There are 4 sublevels:
 S
P
D
F
Energy levels and sublevels work together to form an e- cloud.
e- are repelled by one another and move as far apart as possible.
e- clouds take on characteristic shapes called orbitals .
D. Orbital Shapes
s- orbital shape
p- orbital shape
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E. Electrons and Orbitals
•Orbitals overlap and change shape as electrons are added.
•Each orbital can only hold 2 electrons.
Orbitals and Electrons per Sublevel
Prinicipal Quantum
Number (n)
1
2
3
4
Sublevel
# of Orbitals
# of Electrons per
Orbital
S
1
2
S
P
S P
D
S P
DF
1
3
13
5
13
57
26
26
10
26
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F. Distribution of Electrons
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Atoms are electronically neutral.
There is an electron for every proton in the nucleus.
The larger the atom, the larger the electron cloud.
o Pauli Exclusion Principle: only two e- can occupy the same orbital due to opposite electronic spin.
Electron Filling Diagram
1s_______
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G. Electron Configuration
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WRITE the Electron Configuration
o C___________________
o Kr__________________
o Ca__________________
o
o
Fe__________________
Hg__________________
DRAW the electron configuration for carbon.
o Carbon has 6 e- (same as protons)
o Start with lowest energy level and place one electron in each orbital. Spins must be in same
direction within orbitals of the same energy level.
o If there are remaining e-, pair up singles in same energy level before moving to next highest
energy level.
Draw the arrows
Carbon’s electron config. is: 1S²2S²2P²
•Superscripts total the number of electrons.
2+2+2=6
*Notice that you can write the electron configuration based on the orbital diagram.
*When asked to _____________________________, use __________________________. When asked to
_______________, use ________________________________________.
G. Nobel Gas Electron Configuration
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Electron Configuration demonstrates a periodic trend, so you can write shorthand electron configuration
using the electron configuration of the noble gases in group 18 of the periodic table.
Noble gases have stable configuration.
*Label energy levels, sublevel blocks, and electrons on the blank periodic table.
When writing shorthand e- config for an element, refer to the noble gas in the energy level (period) just
above the element.
Write the symbol of the noble gas in brackets.
Write out the remaining e-config based on the energy filling diagram
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EX: Na
Step 1: Na is the period 3 so refer to the noble gas in period 2 which is neon.
Step 2: Write the Ne in brackets (Ne)
Step 3: Now write remaining electrons in standard form 3S¹
Step 4: Combine (Ne)3S¹
EX: Br
Step 1: Br is in period 4 so refer to noble gas from period 3 which is argon.
Step 2: Write in brackets. (AR)
Step 3: Write remaining electrons.
4S²3D¹º4P5
Step 4: Combine to form: (Ar)4S²3D¹º4P5
*Check your work: Add the number of electrons from the noble gas (18) to the subscripts of the remaining econfig (17). 18+17=35 which is the electrons for Br.
Now try:
1. I ______________________________
3. Na ________________________
2. Kr _____________________________
4. Cu ________________________
H. Electron Configuration with Ions
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When we write the electron configuration of a positive ion, we remove one electron
for each positive charge:
Na
→
Na+
1S²2S²2P63S¹→
1S²2S²2P6
When we write the electron configuration of a negative ion, we add one electron for
each negative charge:
O
→
O21S²2S²2P4 → 1S²2S²2P6
Now try:
1.
Ca+2 _______________________________________
2.
Fe-3 _______________________________________
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Chem I – Demonstration
Chemical Change
Label as Discussed:
1.
Mg (s) + energy + ________ (g)
MgO (s)
What signified a chemical change?______________________________________________
2.
Pb(NO3)2 (aq) + KI (aq)
PbI2 (s) + __________(aq)
What signified a chemical change?______________________________________________
3.
HC2H3O2 (aq) +
CaCO3 (s)
Ca(C2H3O2) (s) + ________(g)
What signified a chemical change?______________________________________________
4.
HC2H3O2 (aq) +
Na2CO3 (s)
Na(C2H3O2)(aq) + H2O (l) + ______(g)
What signified a chemical change?______________________________________________
5.
C12H22O11 (s) +
H2SO4 (aq)
H2O (g) + acid + ______(s)
What signified a chemical change?______________________________________________
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