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Chapter 2
Atoms
Matter
• Matter is anything that has mass and occupies
space.
• Made up of atoms
• Examples of things that are not matter include
thoughts and emotions.
• Classified as being either a pure substance or a
mixture of two or more pure substances.
Pure Substances
• Pure substances are defined as having a “fixed
composition” - one exact combination of ingredients
that cannot be varied.
• Ingredients of a pure substance are chemically
combined so that they lose their original identity and
cannot be separated out without major chemical
intervention.
Pure Substances
• Elements are substances which consist of only
one type of atom. All the elements are listed on
the periodic. Oxygen
• Compounds are substances which consist of
more than one type of atom combined in a
fixed ratio. Water
Elements
• The “building blocks of nature”. All substances, all
matter is created from elements.
• Elements are given chemical symbols, which are one
or two letter abbreviations derived from the common
or Latin name of the element.
• Chemical symbols are always written with the first
letter capitalized and the second letter (if there is one)
in lower case.
Mixtures
• Mixtures are two or more pure substances mixed in a
“varying ratio”. This means that there is more than
one way to combine the ingredients.
• In addition, the ingredients are physically mixed but
not chemically combined and they can be easily
separated.
• Trail mix, hot chocolate
Homogenous Mixtures
• Homogeneous mixtures (solutions) have a uniform
(same) appearance.
• The ingredients are evenly mixed and typically you
cannot see the individual ingredients.
– Kool aid
• If you take two samples of a homogeneous mixture,
the samples will have the same composition due to
the even mixing.
– Glasses of Kool aid
Heterogeneous mixtures
• Heterogeneous mixtures do not have a uniform
(different) appearance.
• The ingredients do not mix evenly and typically you
can see the individual components.
– Trail mix, Chocolate chip cookies.
• Samples of heterogeneous mixtures do not
necessarily have the same composition.
– Hand fulls of trail mix
Classifying Matter
Oxygen
Water
Kool Aid
Trail Mix
Classifying Matter
• Classify the following substances:
•
•
•
•
•
•
Ice
Blueberry muffin
Vitamin A
Oxygen
Tea
Vegetable soup
John Dalton (1766-1844)
• English chemist & teacher
(started teaching at age 12)
• Experiments led to
discovery that the ratios of the
elements in the compound are
whole number multiples
(law of multiple proportions)
Dalton’s Atomic Theory (1808)
• First modern atomic theory, based on empirical evidence
• Called particles atoms in honor of Democritus
• Based on 3 important laws
Law of Conservation of Matter
Law of Constant Composition
Law of Multiple Proportions
• Dalton proposed an atomic theory for matter that would
explain the existence of all 3 laws.
Postulates of Dalton’s Atomic Theory
1.
2.
3.
4.
5.
All elements are made of tiny, indivisible particles called
atoms.
All atoms of a given element have the same chemical
properties. Conversely, atoms of different elements have
different chemical properties.
In ordinary chemical reactions, no atom of any element
disappears or is changed into an atom of another element
Compounds are formed by the chemical combination of
two or more different kinds of atoms. Ina given compound,
the relative numbers of atoms of each kind of element are
constant, and are most commonly expressed as integers.
A molecule is a tightly bound combination of two or more
atoms that acts as a single unit.
Law of Conservation of Mass
Mass of reactant
100.00g
Mass of products = 93.57 g + 6.43 g
100.00 g
During a chemical reaction, the mass of the reactants before the
reaction is _____ to the mass of the products after the reaction (mass
is “conserved”)
Law of Constant Composition
Wherever found, calcium carbonate is always:
40.04% calcium
12.00% carbon
47.96% oxygen
Monatomic Elements
• Consist of single atoms that are not bonded to
each other.
• Helium
• Krypton
• Neon
Diatomic Elements
• Contain 2 atoms bonded to each other in a
molecule. Usually exist as gasses and will not
exist as “free atoms”
• Oxygen gas: O2
• Hydrogen gas:
• Nitrogen gas:
• Fluorine gas:
• Chlorine gas:
• Bromine gas:
• Iodine gas:
Polyatomic Elements
• Contain many of the same atoms bonded to
each other
• Ozone
• Figure 2.5 page 37
The Atom
• Made up of 3 subatomic particles
– Proton
– Electron
– Neutron
Subatomic Particles
Particle
Electron
Proton
Neutron
Location
Relative
electrical
charge
Actual
Mass (g)
Relative
Mass
(amu)
__1__
1840
In space
around the
nucleus
In the
nucleus
9.11 x 10-28
1.673 x 10-24
1.0
In the
nucleus
1.675 x 10-24
1.0
Rutherford’s Nuclear Model
of the Atom
• The atom is mostly empty space (99.99+%)
• The positive charge is centrally located in a dense nucleus
(most of the mass of the atom)
• The nucleus is approximately 100,000 times smaller than the
atom
• The electrons are located at a distance away from the nucleus
and avoid being pulled into the nucleus by orbiting at rapid
speeds
Rutherford’s Atomic Model
(“planetary model”)
________
e-
_________
__________
Diameter = 10-15 m
Diameter of atom = 10-10 m
• Carbon-12
– ___ protons
– ___ neutrons
– ___ electrons
Mass Number
• The mass number is the number of _____ and
______ in an atom.
• The mass of an atom comes from the ______
and _______, because the _______ had
“negligible” mass.
• Problem 2.2 page 39
Atomic Number (Z)
• The atomic number is the number of __________ in
an atom.
• It is unique to each type of atom. For example, all
carbon atoms have ____ protons. Any atom with six
protons is a ______ atom.
• The atomic number also gives us the number of
_______ in an atom.
• Atoms are charge-neutral, which means that every
proton in the nucleus is balanced by an electron in the
electron cloud.
Atomic Number
Carbon has an atomic number of __
Chlorine has an atomic number of __
Sodium has an atomic number of __
me™ and a
Problem 2.3 and 2.4 page 40
ed) decompressor
Problem
Element
Symbol
Atomic
Number
N
Mass
Number
Protons
Neutrons
38
50
15
Calcium
42
14
56
16
138
Electrons
What is an isotope?
Atoms with the same number of protons but different
numbers of neutrons are called isotopes.
A= mass number
Z = atomic number
Isotopes differ only in numbers of neutrons
Isoptopes
• It is possible for atoms of the same element to have
different numbers of neutrons.
• The identity of the element is based on the number of
protons
• The neutrons are charge-neutral, changing the number
of neutrons only changes the mass of the atom.
Measuring atomic masses
Isotope
196Hg
% found
0.15
198Hg
9.97
199Hg
16.87
200Hg
23.10
201Hg
13.18
202Hg
29.86
204Hg
6.87
Magnesium in Nature
Magnesium has three naturally-occurring isotopes:
24Mg
78.99 %
25Mg
10.00 %
26Mg
11.01 %
1. How many protons, neutrons & electrons in
each isotope?
2. What is the principal isotope of Mg
Bromine in Nature
Bromine has two naturally-occurring isotopes:
79Br
50.69 %
81Br
49.31 %
1. How many protons & neutrons in each?
2. In 10,000 atoms of Br, how much of each isotope
would be present?
Problem 2.5 page 41
Atomic Mass
• The mass of an atom is based on the average mass of
all the isotopes of an element.
• Elements have a naturally occurring abundance of
each isotope.
• These naturally occurring abundances are used with
the isotopic mass to calculate the average atomic
mass of an element.
Atomic Mass Calculations
• The atomic masses of each element are shown
on the periodic table.
• When you do these problems, you can always
check your work against the periodic table.
• Your results should be close but will probably
not be exactly the same.
Magnesium has three naturally-occurring isotopes:
24Mg
78.99 %
25Mg
10.00 %
26Mg
11.01 %
1. What is the average atomic mass of Mg?
Bromine has two naturally-occurring isotopes:
79Br
81Br
50.69 %
49.31 %
2. What is the average atomic mass of Mg?
Atomic Mass Calculations
• Copper has two isotopes: 63Cu with a mass of 63.01
amu and an abundance of 69% and 65Cu with a mass
of 65.01 amu and an abundance of 31%. Calculate
the atomic mass of copper.
• In 100 g of copper, how many grams of 63Cu?
Dmitri Mendeleev (1834-1907)
• Chemists began trying to design
a system of classifying the
elements based on their properties
• When Mendeleev arranged the
elements by increasing atomic
weight, he noticed a periodic
repetition in atomic properties
(e.g. density, melting point, etc.)
Mendeleev’s Table
• Mendeleev left blanks in his table, predicting
elements (and their properties) not yet discovered
• Led to search for these elements that, when
discovered, had the properties Mendeleev had
predicted
• Powerful predictive properties of the table led to its
acceptance
The Periodic Table
• The periodic table is a systematic arrangement
of all the elements which exist in:
– Nature
– contrived laboratory conditions
– In our minds (?)
The Periodic Table
• Groups: Vertical columns
• Periods: Horizontal rows
• Elements in the same groups have similar
chemical properties.
The Periodic Table
• Metals: To ______ of jagged line (except hydrogen)
• Nonmetals: To _____ of jagged line (except hydrogen)
• Metalloids: _______ jagged line (except aluminum
which is a metal)
• Alkali metals: Group ____
• Alkaline earth metals: Group ____
• Halogens: Group ____
• Noble gases: Group ____
• Transition metals: All Group ____
Metals, Nonmetals, and Metalloids
• Metals are shiny, ductile and malleable, and are good
conductors of heat and electricity. They have _____
densities and _____ melting points.
• Nonmetals are dull and brittle. They are poor
conductors of heat and electricity and have relatively
_____ densities and _____ melting points.
• Metalloids are shiny and brittle. They are poor
conductors of heat and electricity, and have _____
densities and melting points (although not as _____
as metals).
Using the Periodic Table
Electron Energy Levels
• Electrons are constantly moving
• The shape of the electron path is called the orbital
– Sphere shapes are called s orbitals
– Figure 8 shapes are called p orbitals
• The distance between electrons and the nucleus is
called the energy level
– Closest to the nucleus is called n=1
– Energy levels go to infinity
• Practice worksheet from your Chem 121 packet
Electron Energy Levels
• Electrons are free to move between energy levels,
like going up and down floors in a building
– Moving up in energy levels requires atom to absorb energy
– Moving down in energy levels requires atom to release
energy
• Energy absorbed and released during energy level
transitions is called photons
• Photons are more commonly known as “light” and
“radiation”
Valence Electrons
• We now know that periodic trends are the result of
similarities in the electrons in an atom. More
specifically, periodic trends are the results of
similarity in valence electrons, which are the
electrons located in the highest energy level (furthest
from the nucleus).
• Valence electrons perform the chemistry in a
chemical reaction
Valence Electrons
• All atoms in Group 1A have one valence electron,
Group 2A have 2 valence electrons, Group 3A have 3
valence electrons etc…
• Valence electrons are typically represented in one of
two ways: electron configuration (not covered in this
class) and Lewis dot symbols (also called electron-dot
symbols).
Lewis Dot Symbols (Structures)
• Lewis dot symbols show only the valence
electrons in an atom.
Periodic Trends
• The periodic table arranges elements according
to similarities in properties, such as reactivity
or physical state.
• These similarities are called periodic trends.
Periodic Trends
• The periodic table was originally designed
based completely on observations of periodic
trends.
• The chemists who created the periodic table
were even able to correctly predict physical
properties and the reactivity of elements that
had not yet been discovered.
Periodic Properties
• Atomic Radius is the volume of the electron
clouds in atoms
• Increases going down each group
• Decrease going across from left to right
Atomic Radius
Increasing
Atomic
Radius
Decreasing
Atomic Radius
Periodic Properties
• Ionization Energy is the energy needed to
remove the least tightly bound electron from
an atom in the gaseous state. When this
happens, the atom becomes an ion
• Increases going across from left to right
• Decreases going down a group
Ionization Energy
Ionization Energy Increases
Ionization
Energy
Decreases