Unit 3: Atomic StructureChemistry 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

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


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.


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


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.


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.


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.


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.


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


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.


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.


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


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)?


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


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

Documents

Unit 3: Atomic StructureChemistry 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