Chemistry: Atoms First
Julia Burdge & Jason Overby
Chapter 2
Atoms and the
Periodic Table
Kent L. McCorkle
Cosumnes River College
Sacramento, CA
Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
2.2 Subatomic Particles and Atomic Structure
In the late 1800’s, many scientists were doing research involving
radiation, the emission and transmission of energy in the form of
waves.
They commonly used a cathode ray tube, which consists of two
metal plates sealed inside a glass tube from which most of the air
has been evacuated.
Subatomic Particles and Atomic Structure
When metal plates are connected to a high-voltage source, the
negatively charged plate, or cathode, emits an invisible ray.
The cathode ray is drawn to the anode where it passes through a
small hole.
Although invisible, the path is revealed when the ray strikes a
phosphor-coated surface producing a bright light.
Subatomic Particles and Atomic Structure
Researches discovered that like charges repel each other, and
opposite charges attract one another.
J. J. Thomson (1856-1940) noted the rays were repelled by a
plate bearing a negative charge, and attracted to a plate bearing a
positive charge.
Subatomic Particles and Atomic Structure
This prompted him to propose the rays were actually a stream of
negatively charged particles.
These negatively charged particles are called electrons.
By varying the electric field and measuring the degree of
deflection of cathode rays, Thomson determined the charge-tomass ratio of electrons to be 1.76×108 C/g. (C is coulomb, the
derived SI unit of electric charge.)
Subatomic Particles and Atomic Structure
R. A. Millikan (1868-1953) determined the charge on an electron
by examining the motion of tiny oil drops.
The charge was determined to be -1.6022×10-19 C.
Subatomic Particles and Atomic Structure
Knowing the charge, he was then able to use Thomson’s chargeto-mass ratio to determine the mass of an electron.
charge
1.6022  1019 C
28
mass of an electron =
=
=
9.10

10
g
8
charge / mass
1.76  10 C/g
Subatomic Particles and Atomic Structure
Wilhelm Rontgen (1845-1923) discovered X-rays. They were not
deflected by magnetic or electric fields, so they could not consist of
charged particles.
Antoine Becquerel (1852-1908) discovered radioactivity, the
spontaneous emission of radiation.
Radioactive substances, such as uranium, can produce three types of
radiation.
Subatomic Particles and Atomic Structure
Alpha (α) rays consist of positively charged particles, called α
particles.
Beta (β) rays, or β particles, are electrons so they are deflected away
from the negatively charged plate.
Gamma (γ) rays, like X-rays, have no charge and are unaffected by
external electric or magnetic fields.
Subatomic Particles and Atomic Structure
Ernest Rutherford used α
particles to prove the
structure of atoms.
The majority of particles
penetrated the gold foil
undeflected.
Sometimes, α particles were
deflected at a large angle.
Sometimes, α particles
bounced back in the direction
from which they had come.
Subatomic Particles and Atomic Structure
Rutherford proposed a new
model for the atom:
Positive charge is concentrated
in the nucleus.
The nucleus accounts for most
of an atom’s mass and is an
extremely dense central core
within the atom.
 A typical atomic radius is
about 100 pm
 A typical nucleus has a
radius of about 5×10–3 pm
 1 pm = 1×10–12 m
2.1
Atoms First
An atom is the smallest quantity of matter that still retains the
properties of matter.
An element is a substance that cannot be broken down into two or
more simpler substances by any means.
 Examples: gold, oxygen, helium
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What are atoms made of?
Protons: positively charged particles that are housed in the nucleus
of an atom and have significant mass
Neutrons: neutral particles that are housed in the nucleus. They act to
hold the protons in place since like charges repel each
other. Neutrons have significant mass
Electrons have negligible mass, have a negative charge and are allowed
to roam freely in the electron cloud so they take up significant
volume in the atom
Taylor 2010
2.3 Atomic Number, Mass Number, and Isotopes
All atoms can be identified by the number of protons and neutrons
they contain.
The atomic number (Z) is the number of protons in the nucleus.
 Atoms are neutral, so it’s also the number of electrons.
 Protons determine the identity of an element. For example,
nitrogen’s atomic number is 7, so every nitrogen has 7 protons.
The mass number (A) is the total number of protons and neutrons.
 Protons and neutrons are collectively referred to as nucleons.
Mass number
(number of protons + neutrons)
Atomic number
(number of protons)
A
Z
X
Elemental symbol
Atomic Number, Mass Number, and Isotopes
Most elements have two or more isotopes, atoms that have the
same atomic number (Z) but different mass numbers (A).
1 proton
0 neutrons
1 proton
1 neutron
1 proton
2 neutrons
Isotopes of the same element exhibit similar chemical properties,
forming the same types of compounds and displaying similar
reactivities.
Worked Example 2.1
Determine the numbers of protons, neutrons, and electrons in each of the
35
41
37
following species: (a) 17 Cl, (b) Cl, (c) K, and (d) carbon-14.
Strategy Recall the superscript denotes the mass number (A) and the subscript
denotes the atomic number (Z). If no subscript is shown, the atomic number can
be deduced from the elemental symbol or name. Atoms are neutral so the number
of electrons equals the number of protons.
Solution
(a) Z = 17, so 17 protons
A = 35, so 35 - 17 = 18 neutrons
# of electrons = # of protons, so 17 electrons
(b) Element is chlorine again, so Z must be 17; 17 protons
A = 37, so 37 - 17 = 20 neutrons
17 protons, so 17 electrons
Worked Example 2.1 (cont.)
Determine the numbers of protons, neutrons, and electrons in each of the
35
41
37
following species: (a) 17 Cl, (b) Cl, (c) K, and (d) carbon-14.
Strategy Recall the superscript denotes the mass number (A) and the subscript
denotes the atomic number (Z). If no subscript is shown, the atomic number can
be deduced from the elemental symbol or name. Atoms are neutral so the number
of electrons equals the number of protons.
Solution
(c) Potassium’s atomic number is 19, so 19 protons
A = 41, so 41 - 19 = 22 neutrons
# of electrons = # of protons, so 19 electrons
Worked Example 2.1 (cont.)
Determine the numbers of protons, neutrons, and electrons in each of the
35
41
37
following species: (a) 17 Cl, (b) Cl, (c) K, and (d) carbon-14.
Strategy Recall the superscript denotes the mass number (A) and the subscript
denotes the atomic number (Z). If no subscript is shown, the atomic number can
be deduced from the elemental symbol or name. Atoms are neutral so the number
of electrons equals the number of protons.
Solution
(d) Carbon-14 can also be represented as 14 C
Carbon’s atomic number is 6, so 6 protons
A = 14, so 14 - 6 = 8 neutrons
6 protons, so 6 electrons
Think About It Verify that the number of protons and the number of neutrons
for each example sum to the mass number that is given. In part (a), there are 17
protons and 18 neutrons, which sum to give a mass number of 35, the value
given in the problem. In part (b), 17 protons + 20 neutrons = 37. In part (c),
19 protons + 22 neutrons = 41. In part (d), 6 protons + 8 neutrons = 14.
Do You Understand Isotopes?
14
How many protons, neutrons, and electrons are in 6 C ?
6 protons, 8 (14 - 6) neutrons, 6 electrons
11
How many protons, neutrons, and electrons are in 6 C ?
6 protons, 5 (11 - 6) neutrons, 6 electrons
2.3
Ions
Atoms are neutral—meaning that the number of protons
is equal to the number of electrons
If an atom loses or gains electrons the atom is
No longer neutral but has a charge.
An ion is an atom, or group of atoms, that has a net
positive or negative charge.
cation – ion with a positive charge
If a neutral atom loses one or more electrons
it becomes a cation.
Na
11 protons
11 electrons
Na+
11 protons
10 electrons
anion – ion with a negative charge
If a neutral atom gains one or more electrons
it becomes an anion.
Cl
17 protons
17 electrons
Cl-
17 protons
18 electrons
A monatomic ion contains only one atom
Na+, Cl-, Ca2+, O2-, Al3+, N3-
A polyatomic ion contains more than one atom
OH-, CN-, NH4+, NO3-
Nuclear Symbol of Ions
A C
X
Z
X = element symbol
A = atomic number
Z = mass number
C = charge of Ion
Number of Protons = Z
Numbers of Neutrons = A – Z
Number of electrons = Z - C
Do You Understand Ions?
How many protons, neutrons and electrons are in
27
13
Al
3+
?
Protons = 13
Neutrons = 27-13 = 14
Electrons = 13 – 3 = 10
78 2How many protons, neutrons and electrons are in 34Se ?
Protons = 34
Neutrons = 78 – 34 = 44
Electrons = 34 – (-2) = 36
Taylor 2012
Micro World
atoms & molecules
Macro World
grams
Atomic mass is the mass of an atom in
atomic mass units (amu)
By definition:
1 atom 12C “weighs” 12 amu
On this scale
1H
= 1.008 amu
16O
= 16.00 amu
3.1
2.4 Average Atomic Mass
Atomic mass is the mass of an atom in atomic mass units (amu).
1 amu = 1/12 the mass of a carbon-12 atom
The average atomic mass on the periodic table represents the
average mass of the naturally occurring mixture of isotopes.
Isotope
Isotopic mass (amu)
Natural
abundance (%)
12C
12.00000
98.93
13C
13.003355
1.07
Average mass (C) = (0.9893)(12.00000 amu) + (0.0107)(13.003355 amu)
= 12.01 amu
Worked Example 2.2
Oxygen is the most abundant element in both Earth’s crust and the human body.
17
16
The atomic masses of its three stable isotopes, 8O (99.757 percent), 8O (0.038
percent),188O (0.205 percent), are 15.9949, 16.9991, and 17.9992 amu,
respectively. Calculate the average atomic mass of oxygen using the relative
abundances given in parentheses.
Strategy Each isotope contributes to the average atomic mass based on its
relative abundance. Multiplying the mass of each isotope by its fractional
abundance (percent value divided by 100) will give its contribution to the
average atomic mass.
Solution
(0.99757)(15.9949 amu) + (0.00038)(16.9991 amu) + (0.00205)(17.992 amu)
= 15.9994 amu
Think About It The average atomic mass should be closest to the atomic mass
of the most abundant isotope (in this case, oxygen-16) and, to four significant
figures, should be the same number that appears in the periodic table on the
inside front cover of your textbook (in this case, 16.00 amu).
Natural lithium is:
7.42% 6Li (6.015 amu)
92.58% 7Li (7.016 amu)
Average atomic mass of lithium:
(7.42 x 6.015) + (92.58 x 7.016)
100
= 6.941 amu
3.1
Average atomic mass (6.941)
Names associated with an amount
Can you think of any more?????
The Mole…..
The mole (mol) is the amount of a
substance that
contains as many elementary
entities as there
are atoms in exactly 12.00 grams
of 12C
1 mol = NA = 6.0221367 x 1023
Avogadro’s number (NA)
eggs
Molar mass is the mass of 1 mole of shoes in grams
marbles
atoms
1 mole 12C atoms = 6.022 x 1023 atoms = 12.00 g
1 12C atom = 12.00 amu
1 mole 12C atoms = 12.00 g 12C
1 mole lithium atoms = 6.941 g of Li
For any element
atomic mass (amu) = molar mass (grams/mol)
3.2
2.6 The Mole and Molar Mass
The mole is defined as the amount of a substance that contains as
many elementary entities as there are atoms in exactly 12 g of
carbon-12.
This experimentally determined number is called Avogadro’s
number (NA).
NA = 6.0221415 x 1023
We normally round this to 6.022×1023.
1 mole = 6.022×1023, just like 1 dozen = 12 or 1 gross = 144.
Worked Example 2.3
Calcium is the most abundant metal in the human body. A typical human body
contains roughly 30 moles of calcium. Determine (a) the number of Ca atoms in
30.00 moles of calcium and (b) the number of moles of calcium in a sample
containing 1.00×1020 Ca atoms.
Strategy Use Avogardo’s constant, 1 mole = 6.022×1023, to convert from
moles to atoms and from atoms to moles.
Solution
23 Ca atoms
6.022×10
(a) 30.00 mol Ca ×
1 mol Ca
(b) 1.00×1020 Ca atoms ×
= 1.807×1025 Ca atoms
1 mol Ca
6.022×1023 Ca atoms
= 1.66×10-4 mol Ca
Think About It Make sure that units cancel properly in each solution and that
the result makes sense. In part (a), for example, the number of moles (30) is
greater than one, so the number of atoms is greater than Avogadro’s number. In
part (b), the number of atoms (1×1020) is less than Avogadro’s number, so
there is less than a mole of substance.
The Mole
One mole each of some
familiar substances:
Helium (in balloon)
Water
Aluminum
Sugar (Sucrose)
Copper
Salt (Sodium Chloride)
Molar Mass
The molar mass of a substance is the mass in grams of one mole of
the substance.
By definition, the mass of a mole of carbon-12 is exactly 12 g.
 Mass of 1 carbon-12 atom: exactly 12 amu
 Mass of 1 mole of carbon-12: exactly 12 g
Although molar mass specifies the mass of one mole, making the
units (g), we usually express molar masses in units of grams per
mole (g/mol) to facilitate cancellation of units in calculations.
Worked Example 2.4
Determine (a) the number of moles of C in 25.00 g of carbon, (b) the number of
moles of He in 10.50 g of helium, and (c) the number of moles of Na in 15.75 g
of sodium.
Strategy Molar mass of an element is numerically equal to its average atomic
mass. Use the molar mass for each element to convert from mass to moles.
Think About It Always double-check unit cancellations in
Setup (a) The molar mass of carbon is 12.01 g/mol. (b) The molar mass of
problems such as these–errors are common when molar
helium is 4.003 g/mol. (c) The molar mass of sodium is 22.99 g/mol.
mass is used as a conversion factor. Also make sure that
Solution
the results make
sense. For example, in the case of part (c),
1 mol C
(a) 25.00
g C ×smaller than the
= 2.082
mol
C corresponds to less
a mass
molar
mass
12.01 g C
than a mole.
1 mol He
(b) 10.50 g He ×
= 2.623 mol He
4.003 g He
1 mol Na
(c) 15.75 g Na ×
22.99 g Na
= 0.6851 mol Na
Interconverting Mass, Moles, and Number of Atoms
Molar mass is the conversion factor from mass to moles, and vice
versa.
Avogadro’s constant converts from moles to atoms.
Worked Example 2.5
Determine (a) the number of C atoms in 0.515 g of carbon, and (b) the mass of
helium that contains 6.89×1018 He atoms.
Strategy Use the conversions depicted in the previous slide to convert (a) from
grams to moles to atoms and (b) from atoms to moles to grams.
Think About It A ballpark estimate can help you prevent
Setup (a) The molar mass of carbon is 12.01 g/mol. (b) The molar mass of
common errors. For example, the mass in part (a) is smaller
helium is 4.003 g/mol. NA = 6.022×1023
than the molar mass of carbon. Therefore, you should
expect a number of atoms smaller than Avogadro’s number.
Solution
23 C atoms
1 mol
C
6.022 in
×10
Likewise,
the
number
of
atoms
part
(b) is smaller
than 22 C atoms
(a)0.515 g C ×
= 2.58×10
×
g C Therefore, you
1 molshould
C
Avogadro’s12.01
number.
expect a mass
of helium smaller than the molar mass of helium.
1 mol He
4.003 g He
18
(b)6.89×10 He atoms ×
6.022 ×1023 He atoms × 1 mol He
= 4.58×10-5 g He
How many atoms are in
0.551 g of potassium (K) ?
1 mol K = 39.10 g K
1 mol K = 6.022 x 1023 atoms K
1 mol K
6.022 x 1023 atoms K
0.551 g K x
x
=
1 mol K
39.10 g K
8.49 x 1021 atoms K
What is the mass of 1.21  1020 atoms of sulfur
(in mg) ?
6.44 mg