Unit 3:
Atomic Structure
Chemistry Unit 03: Structure of the Atom p. 1
OVERVIEW
I. ATOMIC THEORY – HISTORICAL DEVELOPMENT ................... 2
A. THE EARLY HISTORY .......................................................................................................... 2 B. EARLY PARTICLE RESEARCH ........................................................................................... 3
II. NUCLEAR CHEMISTRY – COUNTING ATOMS ............................ 6
A. RADIOACTIVITY ................................................................................................................... 6 B. COUNTING ATOMS ............................................................................................................... 8
III. THE MOLE .................................................................................. 10
Chemistry Unit 03: Structure of the Atom p. 2
ATOMIC THEORY – HISTORICAL DEVELOPMENT
A. THE EARLY HISTORY
Although science tries to understand the workings of the universe in a detached and objective fashion, it is
important to keep in mind that science is a still human activity – devised by humans, performed by
humans, and interpreted by humans. As such, what we know about the physical world comes from a
history of human investigation.
For much of western civilization Aristotelian ideals (e.g., the four-elements and that understanding the
natural world involved inductive reasoning; q.v., study guide 01: Introduction to Chemistry) held strong
and alchemy predominated the study of the natural world. However, several important concepts began to
emerge about nature, which, if not exactly contrary to the major paradigm, were inconsistent with it. For
example, the idea of a small, indivisible piece of matter composing all substances began to emerge in the
1600’s in the writings of several prominent scientists, such as Isaac Newton and Robert Boyle, as the only
way to explain certain phenomena. Democritus had first proposed the idea of the indivisible atomos and
described it with the analogy grains of sand making up the desert.
Three important concepts lead John Dalton in 1808 to propose the atomic theory.
1. The French chemist Joseph Louis Proust proposed the Law of constant
composition (a.k.a., law of definite proportions) in 1794: the composition of a pure
compound is always the same. For example, water is always composed of two parts
hydrogen and one part oxygen, wherever it is. Even today, some people believe that
compounds found in nature are somehow different than compounds made in the lab.
Another example is that vitamin C has the same composition whether it is
synthesized in a chemistry lab or isolated from an orange or present in a lime.
2. Antoine Lavoisier, the French scientist and nobleman1, was the first to
unambiguously state the Law of the conservation of mass: the total mass of matter
stays the same in any experiment2.
3. John Dalton, an English schoolteacher, gave us the Law of multiple proportions
(sometimes called Dalton’s Law). Stated one way, if one element can combine with
another element to form more than one compound, then the ratio of the masses of the
elements is in small, whole numbers. This law may be better understood by
explaining with the following example (and using our current understanding of
atomic theory):
a. Hydrogen can combine with oxygen to form either water (H2O) or hydrogen
peroxide (H2O2).
b. In H2O, 8.0 g of oxygen combines with 1.0 g of hydrogen. In H2O2, 16.0 g of
oxygen combines with 1.0 g of hydrogen.
c. Thus, the ratio of oxygen in H2O to that in H2O2 is 16.0 / 8.0, or 2:1.
1 Lavoisier was a controversial figure (e.g., tax collector) and was executed by guillotine during the French revolution. 2 Following the introduction of the atomic age, this has been amended to state that the total mass and energy of matter stays
constant. Matter can be changed into energy according to Einstein’s famous equation E = mc2.
Chemistry Unit 03: Structure of the Atom p. 3
In 1808, John Dalton proposed the atomic theory that is summarized by the following statements:
1. All matter is composed of extremely small particles called atoms.
2. Atoms of a given element are identical in size, mass and other properties; atoms of different
elements differ in size, mass and other properties. (2)
3. Atoms cannot be subdivided, created or destroyed. (3)
4. Atoms of different elements combine in simple whole-whole number ratios to form chemical
compounds, and
5. In chemical reactions, atoms are combined, separated or rearranged.
Notes:
(2) We still believe this if you substitute the works average mass for mass. This is because there are
different isotopes for the same element. E.g., carbon-12 and carbon-14 are the same size, have
the same number of protons (which determine chemical behavior), but differ in atomic mass: C-
12 has 6 protons and 6 neutrons for an atomic mass of 12 amu whereas C-14 has 6 protons but 8
neutrons for an atomic mass of 14 amu.
(3) Atoms can be altered during nuclear reactions – such
as atomic bombs and reactions in the Sun. Dalton
described the atom as a solid, marble-like object.
Relative atomic sizes for the first two periods are
shown.
Figure 1. According to the atomic theory, the mass of carbon dioxide equals the mass of the carbon plus
the mass of the oxygen.
B. EARLY PARTICLE RESEARCH
If all matter is composed of very small atoms, what makes up the atoms? Are they solid, indivisible
structures as was initially believed, or are they themselves, composed of even smaller parts. Because
atoms are so small and we cannot look directly inside to see how they work, indirect methods were used.
1. CATHODE RAY TUBES – AND THE ELECTRON
Cathode ray tubes (CRTs) are glass tubes containing a gas at a very low pressure (Figure 2A). Metal
electrodes are embedded in the glass and when a high voltage (e.g., 10,000 volts) of electricity are passed
to the electrodes the gas glows. This glowing is called fluorescence. The picture in the tube-type of TV
is a cathode ray tube. The tube filled with a gas in a partial vacuum was invented by the German
glassblower Heinrich Geissler in 1857. In the 1880’s, the English scientist William Crookes noted that
the glowing beam originzated from the negatively-charged electrode – called the cathode. Thus, the beam
emitted from the cathode became known as the cathode ray and the tube, the cathode ray tube.
Outside of popular entertainment, CRTs were used extensively for scientific investigation. In 1895, the
Greman physicist Wilhelm Conrad Roentgen discovered that invisible, highly penetrating x-rays were
also produced by the CRT. He made the first x-rays of his wife’s hand. In 1897, Joseph John (J.J.)
Thomson devised an experiment to test the nature of these cathode rays. He noted several important
characteristics of the beam: (a) when the beam hit an internal paddle wheel (Figure 2B), it caused the
Chemistry Unit 03: Structure of the Atom p. 4
paddle wheel to turn and (b) a metal object placed inside the tube caused a clear shadow
on the wall, and (c) the beam was deflected away from the negative pole of a magnet
(Figure 2C) and bent towards the positive end of the magnet (Figure 2D). These
observations led JJ Thomson to conclude:
1. the beam was composed of small particles that had mass, and
2. the particles were negatively charged.
These particles were electrons. Thus, we credit J.J. Thomson with discovering the electron.
Furthermore, Thomson measured the mass/charge ratio of these particles to be ca. 1840 times lighter than
the mass of a hydrogen atom (which we now know contains one proton and one electron).
A. Cathode ray tube without external influencing
magnet.
B. Cathode ray tube with internal paddle wheel.
C. Cathode ray tube with magnet – negative end
(south pole) directed at beam.
D. Cathode ray tube with magnet – positive end
(north pole) directed at beam.
Figure 2. Schematic diagram of a cathode ray tube. Complete the drawing by filling in all four
drawings. (A) Draw the cathode ray unaffected by outside influences. (B) Draw the cathode
ray and its effect on the paddle wheel. Draw the path of the cathode ray influenced by the
negative (south) end of a magnet (C) and by the positive (north) end of a magnet (D).
JJ Thomson modified Dalton’s model of the atom as a solid sphere by proposing the ‘plum-pudding’
model. Much like plum pudding3 (or our oatmeal-raisin cookies), he envisioned the atom to be a sphere
with an overall positive charge (the ‘pudding’ or ‘oatmeal’) and in it, floating about were the negatively
charged electrons (the ‘raisins’). The total charge is neutral so the total negative charge must equal the
total positive charge.
Figure 3. JJ Thomson’s “plum pudding” model of the atom. The
atom is viewed as a slurry of positive charge in which small,
negatively charged electrons are placed much like plum pudding or
oatmeal with raisins.
3 Which, surprisingly, is not made from plums.
Chemistry Unit 03: Structure of the Atom p. 5
2. OIL-DROP EXPERIMENT – CHARGE AND MASS OF THE ELECTRON
[Millikan’s oil-drop experiment:
o electron charge = –1.6 x 10–19 coulombs
o electron mass = 9.11 x 10–28 grams]
3. GOLD-FOIL EXPERIMENT – AND DISCOVERY OF THE NUCLEUS
In 1911, Ernest Rutherford (a student of JJ Thomson’s) developed an experiment to test
Thomson’s atomic model (Figure 4). He directed a beam of helium nuclei (2 protons & 2
neutrons) at a thin sheet of gold surrounded by a detecting, fluorescent screen (Figure 4).
Based on the ‘plum-pudding’ model of the atom, Rutherford expected that the very
heavier -particles would pass directly through the gold foil. But to his surprise, a few of
the -particles bounced back. His famous remark was it “was if you fired a 15-inch
[artillery] shell at a piece of tissue paper and it came back and hit you.” Rutherford concluded that most
of the mass of the atom is located in a small, dense, positively charged center – the nucleus. He theorized
that the electrons orbit the nucleus much the same way that planets orbit the solar system. The diameter
of the nucleus is ca. 1/10,000th the diameter of the atom4. Hence, most of the atom is empty space!
A. Rutherford’s gold foil experiment.
-particle = a helium nucleus = “+” charge
B. Expected results based on Thomson’s ‘plum-
pudding’ model of the atom.
C. Actual results from the gold foil experiment.
Figure 4. Rutherford’s gold foil experiment showing (A) apparatus, (B) expected results, and (C) actual
experimental results. Fill in the expected and actual results frames.
4. ATOMIC NUMBER – NUMBER OF PROTONS
By studying the x-rays produced the cathode ray tubes with different metal anodes,
Henry Moseley5 (a student of Ernst Rutherford) concluded in 1913 that all of the atoms
of a particular element have the same number of protons. The number of protons is
known as the atomic number (Z). Earlier, when Dmitri Mendeleev had proposed the
4 The nucleus has a density of ca. 2 x 10
8 metric tons/cm
3.
5 When World War I broke out, Moseley enlisted in the Royal Engineers. He fought and died at the Battle of
Gallipoli.
Chemistry Unit 03: Structure of the Atom p. 6
first widely accepted periodic table, he had arranged the elements according to atomic mass, with a few
exceptions. As a result of Moseley’s work, the periodic table is now arranged according to atomic
number. For atoms, because they are neutral, the number of protons equals the number of electrons. Ions
are formed from atoms when electrons are lost or gained, resulting in charged particles. For ions, the
electrical charge = the number of protons – the number of electrons.
5. PARTICLES WITHOUT CHARGE – THE NEUTRON
In 1932, James Chadwick6 identified another particle in the nucleus that had about the same
mass as a proton but had no electrical charge, which he called a neutron. These particles
were produced during a nuclear reaction between beryllium (Be) atoms and alpha particles
(helium nuclei – 2 protons/2 neutrons). This reaction formed carbon (C) atoms and stay
neutrons.
6. PROBLEM – WHAT HOLDS THE ATOM TOGETHER?
Have you ever held two magnets together? Positive end to positive end: they repel; positive end to
negative end: they attract. So, if the protons are in the nucleus, which means a lot of positive charge held
in very close proximity. What keeps these supposedly repelling objects so close together. And, what
keeps the electron, a negatively charged particle, from being so attracted to the positively charged nucleus
that it comes crashing in to the center? This dilemma puzzled scientists for quite some time. As it turns
out, when two protons are in close proximity to each other, there is a strong attraction between them, and
a similar attraction between neutrons. The short-range proton-proton, proton-neutron, and neutron-
neutron forces that hold nuclear particles together are called nuclear forces. As to why the electrons
don’t come crashing into the nucleus was explained by quantum mechanics beginning with Niels Bohr.
II. NUCLEAR CHEMISTRY – COUNTING ATOMS
A. RADIOACTIVITY
In 1896, Henri Becquerel accidentally discovered that uranium ore gave off some type of rays that had the
ability to expose a photographic place covered with black paper. These rays were later
identified as products of radioactivity in which certain atoms spontaneously break
down by emitting particles and very penetrating rays (similar to x-rays). The study of
radioactivity became the object of intense scientific curiosity. In 1898, Marie (picture
at right) and Pierre Curie discovered the radioactive elements polonium (Po) and
radium (Ra), isolating some 0.1 g from tons of pitchblende ore. From 1896-1905,
Crookes, Becquerel, Rutherford, Soddy, Dorn and others discovered that radioactivity
is produced when atoms of one element are changed into the atoms of a different element – called
transmutation. These changes result from changes in the nucleus – i.e., changing the numbers of protons
and neutrons in a given atom. Radiation affects chemical substances and can cause biological damage
(e.g., aging, birth defects, and cancer). There are three types of radioactive decay:
6 Chadwick had been a prisoner of war in Germany for four years during WWI. After returning to England, he
went to work for his mentor, Ernest Rutherford, during his undergraduate years at Cambridge University.
Chemistry Unit 03: Structure of the Atom p. 7
Table 1. The Three Types of Radiation.
Radiation Identity Charge Mass Stopped by
alpha () helium nucleus
(2p+/2no)
+2 4 few inches of air
piece of paper
beta () electron
-1 neg. several feet of air
mm’s of plastic/light metal
gamma () EM radiation
(high energy photons)
0 0 several inches of lead
source:
http://www.howe.k12.ok.us/~jimaskew/chem/cnucler.htm
Figure 5. Isotope Notation (Nuclear Symbol). The atomic number for lithium is 3. One isotope has a
mass number of 7 (three protons and 4 neutrons). In this case, the isotope is also an ion,
having a +1 charge, or one more proton than electron (i.e., 3 p+ and 2 e–). If there is no charge
written, the charge is assumed to be zero (neutral ion: # p+ = #e–).
Some examples of radioactive decay are as follows:
Table 2. Examples of Radioactive Decay.
Parent
Nuclide
Daughter
Nuclide
Radioactive
Particle Type of Decay
Half-life (t½)
92
238U → Th234
90 + He4
2 -decay 4.51 billion years
C14
6 → N14
7 + e0
1 -decay 5730 years
27
60Co → 27
60Co + -decay 5.27 years
Half-life = time required for amount to decay by ½; e.g., if you start with 100 g, 50 g will remain after 1
half-life, 25 g will remain after 2 half-lives, 12.5 g will remain after 3 half-lives, etc.
-decay: The parent nuclide loses two protons and a mass number of four (two protons and two
neutrons).
-decay: A neutron splits into a proton and an electron. The parent nuclide therefore gains one proton
(from the neutron) but the mass number stays the same (one neutron is lost but one proton is
gained).
-decay: The daughter nuclide has the same number of protons and neutrons as the parent – so it looks
the same on paper, but a lot of energy is lost in the decay process.
Chemistry Unit 03: Structure of the Atom p. 8
B. COUNTING ATOMS
ATOMIC NUMBER. The number of protons (atomic number, Z) in an atom determines the atom’s
identity: e.g., all oxygen atoms have eight protons. If the atom doesn’t have eight protons, it isn’t oxygen.
The number of neutrons differs between isotopes. Nuclide is a general term for any isotope of any
element. For example, there are three known isotopes of hydrogen (Table 2). Two, protium (H-1) and
deuterium (H-2) occur naturally.
Table 3. Three Isotopes of Hydrogen.
Isotope
Name
Isotope
Notation
Number of
Mass
Number
%
Abundance
in Nature
Diagram
Protons Neutrons
Hydrogen-1 protium 1
1H 1 0 1 99.985
Hydrogen-2 deuterium 1
2 H 1 1 2 0.015
Hydrogen-3 tritium 1
3H 1 2 3 n/a
Table 4. Properties of Subatomic Particles
Particle Symbols Relative Electric
Charge
Mass
Number
Relative Mass
(amu)
Actual Mass
(kg)
Relative
Approximate
Mass
Electron e–; e0
1 –1 0 0.000 5486 9.109 x 10–31 1/2000
Proton p+ +1 1 1.007 276 1.673 x 10–27 1
Neutron n0; n1
0 0 1 1.008 665 1.675 x 10–27 1
MASS NUMBER. Because isotopes differ by the number of neutrons, identifying an isotope requires one to
stipulate the number of neutrons. To do this, it is common to identify the mass number of an isotope,
which is defined as the total number of protons plus neutrons. Mass numbers are whole numbers.
Table 5. Tabulating Isotope Characteristics.
(1)
Element
(2)
Isotope
(3)
Isotope
Symbol
(4)
Atomic
Number
(5)
Mass
Number
Number of (9)
Net
Charge (6)
Protons
(7)
Neutrons
(8)
Electrons
Oxygen Oxygen-16
8
16O 8 16 8 8 8 0
(1) Element Name.
(2) Isotope name is written as: “element name” – “mass number”.
(3) Isotope symbol (nuclear notation) - see Figure 4.
(4) Atomic number = number of protons. Oxygen always has 8 protons – if the atom doesn’t have eight
protons, then it isn’t oxygen. If you haven’t memorized it, look it up on the periodic table.
(5) Mass number = (number of protons) + (number of neutrons). Number of protons (see atomic number,
above).
(6) Number of protons is the atomic number (see above).
(7) A common misconception is that the number of neutrons equals the number of protons. It doesn’t
always. One way to determine the number of neutrons is from the mass number (see above). For
example, for oxygen-18 (in above table), 16 (mass number) - 8 (number of protons) = 8 (number of
neutrons). Another example: how many neutrons are in potassium-40? The atomic number for
Chemistry Unit 03: Structure of the Atom p. 9
potassium is 19 (always, always, always) and the mass number (given in question). 40 – 19 = 21. So,
potassium-40 has 21 neutrons.
(8) Number of electrons. When there is no other information, assume the net charge is zero and that the
number of protons equals the number of electrons. This is always true for atoms. That’s what makes
an atom. When the atom loses or gains an electron(s), it is an ion. For example, oxygen often gains
two electrons, thus having eight protons and ten electrons. (Remember, if the oxygen was to gain two
protons, it would no longer be oxygen but rather it would become neon.) Oxygen with eight protons
(8+) and ten electrons (10–) now has a net charge of –2 and the ionic symbol is O2–. It can be written
as either O2– or O–2. When the charge is one (either 1+ or 1–), the “1” is generally omitted – e.g., H+,
Li+, and Na
+.
ATOMIC MASSES – RELATIVE AND AVERAGE. Atoms are very, very, small. We need a way to assign a
mass to them rather than referring to them using gram or kilogram. By definition, one atomic mass unit
(amu) is exactly the 1/12 the mass of a carbon-12 atom (or 1.660 540 x 10–27 kg). This works out well
because the mass of one carbon-12 atom is 12 amu and the mass of one mole of carbon-12 atoms is 12
grams.
The atomic mass given on the periodic table for a given element (e.g., 12.011) is the average atomic
mass for all of the isotopes of a given element. It is the weighted average of the atomic masses of the
naturally occurring isotopes of the elements. What is the weighted average? Let’s define weighted
average by example. Let’s say you had 100 marbles. Seventy-five are 2.00 g and the remaining 25 are
3.00 g:
75 marbles x
2.00 g = 150 g
1 marble
25 marbles x
3.00 g = 75 g
1 marble
Total mass = 150 g + 75 g = 225 g
Average mass = 225 g / 100 marbles = 2.25 g
Calculating Average Atomic Mass. Let’s calculate the average atomic mass for hydrogen. This is
typically done using only the naturally occurring isotopes. For example, naturally occurring copper
consists of 69.
Table 4. Determining Average Atomic Mass
Isotope % Abundance Atomic Mass (amu)
Copper-63 69.17 % 62.929 598
Copper-65 30.83 % 64.927 793
calculation: (0.6917 x 62.929 598) + (0.3083 x 64.927 793) = 63.55 amu
N.B. It is common to round to two decimal places the atomic mass.
Chemistry Unit 03: Structure of the Atom p. 10
III. THE MOLE
THE MOLE. When dealing with most quantities in daily life, the dozen is used as a counting unit.
Twelve pencils is one dozen. One dozen dozen, or 144, is a gross. This works well when the objects are
large enough to be seen. But atoms are very small so to add a dozen atoms of oxygen to two dozen atoms
of hydrogen would yield a dozen molecules of water. But who could pour such a small amount! The SI
unit for amount of a substance is the mole (abbreviated mol). Just as a dozen is a unit of quantity, so is a
mole.
In exactly 12 grams of carbon-12 there is one mole of atoms, which is 6.022 1367 x 1023
atoms.
Generally, we round to two decimals and use 1 mole = 6.02 x 1023
. This number of
particles is known as Avogadro’s number, named after the 19th century Italian scientist
Amedeo Avogadro (his full name was Lorenzo Romano Amedeo Carlo Avogadro, conte
di Quaregna e di Cerreto) whose ideas were crucial in explaining the relationship between
mass and numbers of atoms. In 1811, he hypothesized that equal volumes of gases, at the
same temperature and pressure, equal numbers of particles. For gases, one mole of gas
occupies 22.4 L at 0 oC and 1 atmosphere (atm; approximately the pressure at sea level on
a clear day).
MOLAR MASS. The average atomic mass given on any periodic table is the molar mass, which is the
mass of one mole of a pure substance. For example, the molar mass of carbon is 12.011 g; the molar
mass of just carbon-12 is exactly 12 g. The molar mass of water is (2 x 1.008 g) + (1 x 16.00 g) = 18.02 g
because hydrogen has an average mass of 1.008 g/mol and oxygen has an average mass of 16.00 g, and
there are two hydrogen atoms and one oxygen atom in one molecule of water.
CONVERSIONS: MASS MOLE NUMBER OF PARTICLES. Converting between mass, number of
moles and number of particles (e.g., atoms, molecules) is very common in chemistry. If we want to add
two chemicals together in a quantitative fashion (e.g., two moles of hydrogen with one mole of oxygen),
there is no way to directly measure out one mole. Instead, we would convert the number of moles we
wanted into grams and measure out the mass of the chemicals.
The figure on the next page schematically diagrams the conversions between mass, moles, and number of
molecules. Specific examples are given below.
Chemistry Unit 03: Structure of the Atom p. 11
MASS MOLES
o mass moles: 4.92 g Cu(NO3)2 = ? mol
)Cu(NO mol 378.0)Cu(NO g 5558.187
)Cu(NO mol 1*
1
)Cu(NO g 92.423
23
2323
Cu: 1 * 63.546 = 63.546
N: 2 * 14.0067 = 28.0134
O: 6 * 15.9994 = 95.9964
SUM: 187.5558 g/mol
o moles mass: 7.43 mol NaHCO3 = ? g
NaHCO g 1716.624NaHCO mol 1
NaHCO g 0069.84*
1
NaHCO mol 43.73
3
33
Na: 1 * 22.9898 = 22.9898
H: 1 * 1.00794 = 1.00794
C: 1 * 12.011 = 12.011
O: 3 * 15.9994 = 47.9984
SUM: 84.0069 g/mol
MOLES MOLECULES
o moles molecules: 1.37 mol MgCl2 = ? molecules
MgCl molecules 1025.8MgCl mol 1
MgCl molecules 6.02x10*
1
MgCl mol .3712
23
2
2
23
2 x
o molecules moles: 2.51 x 1024
molecules N2O5 = ? mol
ON mol 17.4ON molecules 6.02x10
ON mol 1*
1
ON molecules .51x10252
52
23
5252
24
Chemistry Unit 03: Structure of the Atom p. 12
MASS MOLECULES
o mass molecules: 4.92 g Cu(NO3)2 = ? molecules
)Cu(NO molecules 1028.2)Cu(NO mol 1
)Cu(NO molecules 1002.6*
)Cu(NO g 5558.187
)Cu(NO mol 1*
1
)Cu(NO g 92.423
23
23
23
23
23
2323 xx
Cu: 1 * 63.546 = 63.546
N: 2 * 14.0067 = 28.0134
O: 6 * 15.9994 = 95.9964
SUM: 187.5558 g/mol
o molecules mass: 2.51 x 1024
molecules N2O5 = ? g
ON g .450ON mol 1
ON 108.0105*
ON molecules 6.02x10
ON mol 1*
1
ON molecules .51x10252
52
52
52
23
5252
24
N: 2 * 14.00674 = 28.0135
O: 5 * 15.9994 = 79.9970
SUM: 108.0105 g/mol
Amount of Element in
MOLES
Mass of Element in
GRAMS
NUMBER OF FORMULA
UNITS (e.g., MOLECULES)
x Molar Mass
/ Molar Mass / 6.02 x 1023
x 6.02 x 1023
Problems:
1. What is the mass (g) of 3.60 moles of the element copper?
2. What is the mass in grams of 0.375 mol of iron?
3. A chemist produced 11.9 g of aluminum. How many moles were produced?
4. How many moles of gold are in 3.60 x 10–10 g?
5. How many moles of silver are in 3.01 x 1023
atoms?
6. How many moles of lead are in 1.50 x 1012
atoms?
7. What is the mass in grams of 1.20 x 108 atoms of copper?
8. What is the mass in grams of 3.01 x 1014
molecules of water?
9. What is the mass in grams of 9.0 moles of sucrose (C12H22O11)?
Chemistry Unit 03: Structure of the Atom p. 2
ANSWERS:
1. What is the mass (g) of 3.60 moles of the element copper?
3.6 mol Cu 63.546 g Cu = 228.766 g Cu
1 1 mol Cu
2. What is the mass in grams of 0.375 mol of iron?
0.375 mol Fe 55.847 g Fe = 20.943 g Fe
1 1 mol Fe
3. A chemist produced 11.9 g of aluminum. How many moles were produced?
11.9 g Al 1 mol Al = 0.441 mol Al
1 26.98154 g Al
4. How many moles of gold are in 3.60 x 10–10 g?
3.60E-10 g Au 1 mol Au = 1.83E-12 mol Au
1 196.9665 g Au
5. How many moles of silver are in 3.01 x 1023
atoms?
3.01E+23 atoms Au 1 mol Au = 5.00E-01 mol Au
1 6.02E+23 atoms Au
6. How many moles of lead are in 1.50 x 1012
atoms?
1.50E+12 atoms Pb 1 mol Pb = 2.49E-12 mol Pb
1 6.02E+23 atoms Pb
7. What is the mass in grams of 1.20 x 108 atoms of copper?
1.20E+08 atoms Cu 1 mol Cu 63.546 g = 1.27E-
14 g
1 6.02E+23 atoms Cu 1 mol Cu
8. What is the mass in grams of 3.01 x 1014
molecules of water?
3.01E+14
molecules
H2O 1 mol H2O 18.01528 g = 9.01E-
09 g
1 6.02E+23
molecules
H2O 1 mol H2O
Chemistry Unit 03: Structure of the Atom p. 3
9. What is the mass in grams of 9.0 moles of sucrose (C12H22O11)?
9.0
mol
C12H11O11 342.3001 g = 3080.70 g
1 1
mol
C12H11O11
C: 12 * 12.011 = 144.132
H: 22 * 1.00794 = 22.17468
O: 11 * 15.9994 = 175.9934
342.3001 g/1 mol