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Lecture Presentation
Chapter 2
Atoms and Elements
Catherine MacGowan
Armstrong Atlantic State University
© 2013 Pearson Education, Inc.
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Matter: Its Theories & Laws
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•
•
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Law of the Conservation of Matter
Atomic Theory of Matter
Law of Definite Proportions
Law of Multiple Proportions
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Law of the Conservation of Matter
Lavoisier proposed from his experimental evidence the following law:
Matter is neither created nor destroyed in a chemical
reaction.
•
Total mass of used reactants = Total mass of products produced
• Total number of reactant atoms = Total number of product atoms
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Atomic Theory of Matter
Dalton’s atomic theory of matter proposes that:
– Atoms are small, discrete, indivisible pieces of matter.
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–
–
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All elements are made up of particles called atoms.
An element’s atoms are identical in size, mass, & chemical
properties.
• Scientists did not know about isotopes.
– Isotopes are elemental atoms that differ in their mass
due to different number of neutrons.
Molecules (compounds) are formed when two or more
elements combined.
Molecules are simple whole-number ratios of the combined
elements.
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From Dalton’s atomic theory comes:
Dalton’s atomic theory notes that
“Atoms of different elements can combine to
form molecules in simple whole-number
combinations of atoms.”
•
Law of Definite Proportions
•
Law of Multiple Proportions
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Law of Definite Proportions
Law of definite proportions:
–
For a given compound, the elements always combine in the
same proportion.
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For example:
– Sodium chloride molecule (NaCl) is always a 1:1 ratio of
one sodium atom to chlorine atom.
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A 100.0-g sample of NaCl contains 39.3 g Na & 60.7 g Cl.
Mass Cl =
Mass Na
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60.7 g = 1.54
39.3 g
A 58.44.0-g sample of NaCl contains 22.99 g Na & 35.44 g Cl.
Mass Cl =
Mass Na
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35.44 g
22.99 g
= 1.54
Law of Multiple Proportions
Law of multiple proportions:
•
Two elements A and X can form different compounds
by combining in different proportions.
– These combinations can be represented as a ratio.
•
For example:
– A molecule of carbon dioxide (CO2) has a ratio of 1
C atom to every 2 atoms of oxygen, or 1:2.
–
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A molecule of hydrogen peroxide (CO) has a ratio
of 1 C atom to 1 atom of oxygen, or 1:1.
Atomic Structure
•
Historical Perspective
• Atom Structure
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Atomic Structure: Historic Models
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Greek (philosophical model)
– The ultimate piece of matter is the atom.
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Dalton (solid bullet model)
– Atomic theory of matter
• Elements are composed of discrete particles of
matter called atoms.
• Atoms are tiny, indivisible, indestructible
particles, and each one has a certain mass,
size, and chemical behavior.
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Discovery of the Electron Charge: Mass Ratio
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Thomson’s charge-to-mass experiment
– This experiment investigated the
effect on a cathode ray of placing an
electric field around a tube.
(1) Charged matter is attracted to an
electric field.
(2) Light’s path is not deflected by
an electric field.
–
The cathode rays are made of tiny
particles.
• These particles have a negative
charge because the beam always
deflected toward the (+) plate.
–
The amount of deflection was related
to two factors, the charge and mass
of the particles.
• The charge:mass ratio of these
particles was −1.76 × 108 C/g.
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Atomic Structure: Discovery of the Electron
Discovery of electron:
• Millikan oil drop experiment
– Investigation led to determining the charge of the electron.
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Atomic Structure: Early Discoveries
Discovery of electron:
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Thomson charge-to-mass experiment using cathode
rays
– Thomson believed that these particles were the
ultimate building blocks of matter.
– These cathode ray particles became known as
electrons.
•
Millikan oil drop experiment
–
This experiment showed that the particle had the same amount
of charge as the hydrogen ion, as Thomson observed in his
experiments.
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Atomic Structure: Electron
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Electrons are particles found in all atoms.
–
One of the fundamental pieces of matter
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The electron has a charge of −1.60 × 1019 C.
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The electron has a mass of 9.1 × 10−28 g.
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Atomic Structure: Plum-Pudding Model
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J. J. Thomson (plum-pudding model)
- The atom is composed of a positive cloud of
matter in which electrons are embedded.
• Explains the positive (+), negative (-)
charged behavior of matter
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Rutherford’s Gold Foil
Experiment Setup
Gold foil experiment:
Could not explain Thomson’s plum-pudding atom model.
Led to the discovery of the atom’s nucleus.
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Rutherford & the Nucleus: Gold Foil Experiment
From the gold foil experiment, the following conclusions were proposed:
• The atom contains a tiny, dense center called the nucleus.
• The nucleus has essentially the entire mass of the atom.
– The electrons weigh so little they give practically no mass to the atom.
• The nucleus is positively charged.
– The amount of positive charge balances the negative charge of the
electrons.
– The electrons are dispersed in the empty space of the atom
surrounding the nucleus.
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Atomic Structure: Historic Perspective
•
Rutherford’s model (solar system)
- The atom is mostly empty space with a DENSE center of
mass (nucleus) and circling electrons.
- It proposed that the nucleus had a particle that had the
same amount of charge as an electron but opposite sign.
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These particles are called protons.
-
-
charge = +1.60 × 1019 C
mass = 1.67262 × 10−24 g
Since protons and electrons have the same amount of
charge, for the atom to be neutral, there must be equal
numbers of protons and electrons.
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Atomic Structure: Composition
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Elements
•
Isotopes
• Ions
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Elements
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The number of protons located in an atom’s nucleus determines
the element’s identity.
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The number of protons in the nucleus of an atom is called the
atomic number.
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Elements
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Each element has a unique name and symbol.
– Symbol has either one or two letters
• O (oxygen) or Fe (iron)
– The elements are arranged on the periodic table in order of
their atomic numbers.
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Isotopes
•
Isotopes are elements whose atoms differ in mass only.
– They differ in mass because these elemental atoms have
different number of neutrons.
– They are the same element because they have the same
number of protons (atomic number).
– They are chemically identical.
•
Isotopes are identified by their mass numbers.
–
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Protons + neutrons = mass number
Isotopic symbol
13
Atomic number, Z
Al
Atom symbol
26.981
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Atomic weight
Isotopes & Atomic Mass (Weight)
Problem: How many protons, electrons, and neutrons are in the
following atoms:
protons
32
S
16
65
Cu
29
U-240
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electrons
neutrons
Isotopes & Atomic Mass (Weight)
Problem: How many protons, electrons, and neutrons are in the
following atoms:
protons
electrons
neutrons
S
16
16
16
Cu
29
29
36
U-240
92
92
148
32
16
65
29
NOTE: Neutral atoms will have the same number of protons as electrons.
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Isotopes & Atomic Mass (Weight)
Complete the following table:
Atomic
Mass
Protons Neutrons Electrons Number Number
6
7
42
96
55
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133
Atomic
Symbol
Isotopes & Atomic Mass (Weight)
Complete the following table: Answers
Atomic
Mass
Protons Neutrons Electrons Number Number
6
7
6
6
13
42
54
42
42
96
13
14
13
13
27
55
78
55
55
133
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Atomic
Symbol
Isotopes & Atomic Mass (Weight)
•
What is the connection between an element’s isotopes and its atomic
mass?
– The observed mass of an element is a weighted average of the
weights of all the naturally occurring atoms.
• Average mass = ATOMIC WEIGHT (amu)
•
What information is needed to determine the atomic mass of an
element?
– The natural abundance of each of the element’s isotopes
– The percentage of an element that is one isotope is called the
isotope’s NATURAL ABUNDANCE.
–
Example problem:
• Boron has two isotopes:
– Boron is 19.9% 10B and 80.1% 11B.
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Boron atomic weight:
= 0.199 (10.0 amu) + 0.801 (11.0 amu) = 10.8 amu
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Isotopes & Atomic Mass (Weight)
Problem:
•
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The element silver (Ag) has an atomic mass of 107.868 amu.
It has two isotopes: Ag-109 (108.905 amu) with an abundance of
48.1600% and Ag-107.
Determine the amu for Ag-107.
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Isotopes & Atomic Mass (Weight)
Problem:
The element silver (Ag) has an atomic mass of 107.868 amu. It has
two isotopes: Ag-109 (108.905 amu) with an abundance of 48.1600%
and Ag-107. Determine the amu for Ag-107.
Strategy:
1. Set up an algebraic equation.
2. Solve for x, which is the amu for Ag-107.
Solution:
1.
Set up an algebraic equation.
(0.481600) 108.905 amu + x (0.518400) = 107.868 amu
2.
Solve for x, which is the amu for Ag-107.
52.4490 amu + 0.518400x = 107.868 amu
0.518400x = 55.4180 amu
x = 106.905 amu
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Ions: Charged Atoms
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IONS are atoms or groups of atoms with a positive (+) or
negative (-) charge.
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Taking away an electron from an atom gives a CATION with a
positive charge.
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More protons in nucleus vs. electrons surrounding nucleus
– METALS tend to form cations to achieve the “FULL SHELL” look.
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Adding an electron to an atom gives an ANION with a
negative charge.
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Fewer protons in the nucleus vs. electrons surrounding nucleus
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NONMETALS tend to form anions to achieve the “FULL SHELL”
look.
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Ions: Charged Atoms
Cations
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•
•
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A CATION forms when an atom
loses one or more electrons from
its outer (valence) shell (energy
level).
Anions
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An ANION forms when an atom
gains one or more electrons into its
outer (valence) shell (energy level).
•
Anions are negatively charged
because the atom has fewer
protons (+) than electrons (-).
– F atom has 9 protons & 9
electrons.
– F- ion has 9 protons & 10
electrons.
•
Nonmetal elements tend to form
anions.
•
Example:
Cations are positively charged
because the atom has more protons
(+) than electrons (-).
– Mg atom has 12 protons & 12
electrons.
– Mg2+ ion has 12 protons & 10
electrons.
Metal elements tend to form
cations.
Example:
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Mg  Mg2+ + 2 e-
F + e-  F-
Periodic Table: Overview
• Historic Perspective
•
Organization
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Periodic Table: Historic Perspective
•
Dmitri Mendeleev (1834–1907) developed
the modern periodic table.
– He argued that element properties are
periodic functions of their atomic
weights.
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Periodic law: When the elements are
arranged in order of increasing atomic
mass, certain sets of properties recur
periodically.
– Elements having similar physical &
chemical properties fall within a
column.
•
We now know that element properties are
periodic functions of their ATOMIC
NUMBERS.
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Periodic Table: Organization
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The elements are arranged from left to
right in increasing ATOMIC NUMBER
(number of protons an element has).
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Rows in the periodic table are referred
to as PERIODS.
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Columns in the periodic table are
sometimes referred to as families.
–
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Families because the elements
within the column have similar
physical and chemical properties
The elements, their names, and
symbols are given on the PERIODIC
TABLE.
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Today’s Periodic Table
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Periodic Table:
Metals, Metalloids, & Nonmetals
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Periodic Table: Areas
= Metal
= Metalloid
= Nonmetal
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Periodic Table: Families
= Alkali Metals
= Alkali Earth Metals
= Halogens
= Noble Gases
= Lanthanides
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= Transition Metals
= Actinides
Periodic Table: Metals
Characteristics
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Solids at room
temperature, except Hg
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Reflective surface
– Shiny
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Conduct heat and
electrical current
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Malleable
– Can be shaped
•
Ductile
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Lose electrons and form
cations
•
About 75% of the elements
in the period table are
metals.
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Periodic Table: Nonmetals
Characteristics
• Can be found in all three states (gas,
liquid, & solid) of matter
• Poor conductors of heat & electricity
• Solids are brittle
• Gain electrons to become anions
• Except for H, found mostly in the upper
right of the periodic table
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Periodic Table: Metalloids
Characteristics
• Can exhibit the properties of metals
and/or of nonmetals
• Known as semiconductors
• Poor conductors of heat
• Solids at room temperature
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The Mole:
The “Chemist’s Dozen”
• Definition of Mole
•
Molar Mass of a Compound
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What is a mole?
Chemistry is quantitative in nature;
– Its unit is the mole.
•
The mole as unit vs. “dozen” as an unit
– The unit “dozen” is associated with 12 units.
– The unit MOLE is associated with 6.02 x 1023
units.
•
There is Avogadro’s number of units in every
mole.
– 6.02 x 1023 units is known as Avogadro’s
number.
– 1 mole = 6.02 x 1023 units
•
1 mole of Cu atoms has
– an atomic mass of 63.55 g, which is
– 6.022 × 1023 Cu atoms, which is approximately
22 pennies.
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How large is a mole?
Problem:
You have inherited ¼ mole of pennies. If you spent 1 x 1012 dollars every
second, how many decades would it take for you to spend your
inheritance?
Solution:
¼ mol x (6.02 x 1023 cents/1 mol) x ($1/100 cents)
= 1.51 x 1021 dollars
1.51 x 1021 dollars x (1 sec/ 1 x 1012 dollars) x (1 min/60 sec) x
(1 hr /60 min) x (1 day/24 hrs) x (1 yr/365 days) x (1 decade/10 yrs)
= 4.79 decades
Every mole contains 6.02 x 1023 units—that is a lot of pieces!!!
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Mole & Mass Relationship
An element’s molar mass in grams per mole is
numerically equal to the element’s atomic mass in
atomic mass units.
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Mass-to-Mole Conversions
To go from mass to units, you must go through the MOLE.
[ Insert the scheme found at the top of pg 62 here]
Element:
Mass (g) / (1 mol/atomic mass (element)) = mole of element
Mass (g) x (1 mol/atomic mass) x (6.02 x 1023 atoms/1 mole atoms) = # atoms
Molecule/Compound:
Mass (g) compound/molecular mass (compound)
= mole of compound
Mass (g) x (1 mol/molecular mass) x (6.02 x 1023 molecules/1 mol molecules)
= molecules
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Mass-to-Mole Conversions
How many Mg atoms are in 0.200 g?
Mg has an atomic mass of 24.3050 g/mol.
Need to convert 0.200 g Mg to moles of Mg
0.200 g Mg x (1 mol/24.31 g) = 8.23 x 10-3 mol Mg
Now that you know the moles of Mg, you can determine the
number of Mg atoms.
8.23 x 10-3 mol Mg x (6.02 x 1023 atoms/1 mol Mg)
= 4.95 x 1021 atoms Mg
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Mole-to-Mass Conversions
Problem: How many grams of CO2 are in 6.75 x 1030 molecules of
CO2?
Strategy:
1. Need to know the molecular mass of CO2.
2. Convert molecules of CO2 to moles.
3. Convert moles of CO2 to grams.
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Mole-to-Mass Conversions
Problem: How many grams of CO2 are in 6.75 x 1030 molecules of
CO2?
1. CO2 has a molecular mass of 44.0g/mol.
There is 1 mole carbon in every mole of CO2.
1 mole carbon = 12.0 g
There are 2 moles oxygen in every mole of CO2.
1 mole oxygen = 16.0 g
Therefore, the molecular mass of CO2 is:
1 x 12.0 g =
12.0 g
from 1 mole C
2 x 16.0 g =
32.0 g
from 2 mole O
44.0 g
mass of 1 mole CO2
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Mole-to-Mass Conversions
Problem: How many grams of CO2 are in 6.75 x 1022 molecules of
CO2?
2. Determine the moles of CO2.
6.75 x 1022 molecules CO2 x (1 mole CO2/6.02 x 1023 molecules)
= 1.12 x 10-1 moles CO2
3. Determine the grams of CO2.
1.12 x 10-1 moles CO2 x (44.0 g/1 mol CO2) = 4.93 g
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