High School Chemistry/Atomic Terminology

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Dalton's Atomic Theory explained a lot about matter, chemicals, and chemical reactions. Nevertheless, it wasn't entirely accurate, because contrary to what Dalton believed, atoms can, in fact, be broken apart into smaller subunits or subatomic particles. One type of subatomic particle found in an atom is the negatively charged electron. Since atoms are neutral, though, they also have to contain positive material. At first, scientists weren't sure exactly what this positive material was, or how it existed in the atom. Thomson thought it was distributed throughout the atom like batter in a plum pudding. Rutherford, however, showed that this was not the case. In his gold foil experiment, Rutherford proved that the positive substance in an atom was concentrated in a small area at the center of the atom, leaving most the rest of the atom as empty space (possibly with a few electrons, or an "electron cloud").

Both Thomson's experiments and Rutherford's experiments answered a lot of questions, but they also raised a lot of questions, and scientists wanted to know more. How were the electrons connected to the rest of the atom? What was the positive material at the center of the atom like? Was it one giant clump of positive mass, or could it be divided into smaller parts as well? In this lesson, we'll look at the atom a little more closely.

Electrons, Protons, and Neutrons[edit]

The atom is composed of three different kinds of subatomic particles. First, there are the electrons, which we've already talked about, and which J. J. Thomson discovered. Electrons have a negative charge. As a result they are attracted to positive objects, and repelled from negative objects, which means that they actually repel each other (Figure 4.9).

Figure 4.9: Electrons repel each other because they are both negatively charged.

Still, most atoms have more than one electron. That's because atoms are big enough to hold many electrons without those electrons ever colliding with each other. As you might expect, the bigger the atom, the more electrons it contains.

Protons are another type of subatomic particle found in atoms. Protons have a positive charge. As a result they are attracted to negative objects, and repelled from positive objects. Again, this means that protons repel each other (Figure 4.10). Unlike electrons, however, which manage to stake out a "territory" and "defend" it from other electrons, protons are bound together by what are termed strong nuclear forces. Therefore, even though they repel each other, protons are forced to group together into one big clump. This clump of protons helps to form the nucleus of the atom. Remember, the nucleus of the atom is the mass of positive charge at the atom's center.

Figure 4.10: Protons repel each other because they are both positively charged. Despite this repulsion, protons are bound together in the atomic nucleus as a result of the strong nuclear force.

Electrons were the first subatomic particles discovered and protons were the second. There's a third kind of subatomic particle, though, known as a neutron, which wasn’t discovered until much later. As you might have already guessed from its name, the neutron is neutral. In other words, it has no charge whatsoever, and is therefore neither attracted to nor repelled from other objects. That's part of the reason why the neutron wasn’t discovered until long after people knew about electrons and protons – because it has no charge, it's really hard to detect. Neutrons are in every atom (with one exception), and they’re bound together with other neutrons and protons in the atomic nucleus. Again, the binding forces that help to keep neutrons fastened into the nucleus are known as strong nuclear forces.

Before we move on, we must discuss how the different types of subatomic particles interact with each other. When it comes to neutrons, the answer is obvious. Since neutrons are neither attracted to, nor repelled from objects, they don't really interact with protons or electrons (beyond being bound into the nucleus with the protons). Protons and electrons, however, do interact. Using what you know about protons and electrons, what do you think will happen when an electron approaches a proton - will the two subatomic particles be attracted to each other, or repelled from each other? Here’s a hint: "opposites attract, likes repel". Electrons and protons have opposite charges (one negative, the other positive), so you'd expect them to be attracted to each other and that's exactly what happens (Figure 4.11).

Figure 4.11: Protons and electrons are attracted to each other because they have opposite charges. Protons are positively charged, while electrons are negatively charged.

Viewers journey inside the atom to appreciate its architectural beauty and grasp how atomic structure determines chemical behavior. Video on Demand – The World of Chemistry – The Atom.

Relative Mass and Charge[edit]

Even though electrons, protons, and neutrons are all types of subatomic particles, they are not all the same size. When you compare the masses of electrons, protons and neutrons, what you find is that electrons have an extremely small mass, compared to either protons or neutrons. On the other hand, the masses of protons and neutrons are fairly similar, although technically, the mass of a neutron is slightly larger than the mass of a proton. Because protons and neutrons are so much more massive than electrons, almost all of the atomic mass in any atom comes from the nucleus, which contains all of the neutrons and protons.

Table 4.1 gives the masses of electrons, protons, and neutrons. The second column shows the masses of the three subatomic particles in grams (which is related to the SI unit kilograms according to the relationship 1 kg = 1000 g). The third column, however, shows the masses of the three subatomic particles in "atomic mass units". Atomic mass units (amu) are useful, because, as you can see, the mass of a proton and the mass of a neutron are almost exactly 1.0 in this unit system. We'll discuss atomic mass units in a later section.

Table 4.1: Masses of the Different Subatomic Particles
  Mass in Grams (g) Mass in Atomic Mass Units (amu)
Electron 9.109383×10−28 5.485799095×10−4
Proton 1.6726217×10−24 1.0072764669
Neutron 1.6749273×10−24 1.0086649156

Unfortunately, the numbers in Table 4.1 probably don't give you a very good sense of just how big protons and neutrons are compared to electrons, so here's a comparison that might help. If an electron were the size of a penny, then a proton (or a neutron) would be about the size of a large bowling ball (Figure 4.12).

Figure 4.12: Electrons are much smaller than protons or neutrons. How much smaller? If an electron was the size of a penny, a proton or a neutron would have the mass of a large bowling ball!

In addition to mass, another important property of subatomic particles is their charge. You already know that neutrons are neutral, and thus have no charge at all. Therefore, we say that neutrons have a charge of zero. What about electrons and protons? You know that electrons are negatively charged and protons are positively charged, but what's amazing is that the positive charge on a proton is exactly equal in magnitude (magnitude means "absolute value" or "size when you ignore positive and negative signs") to the negative charge on an electron.

Table 4.2 gives the charges on electrons, protons, and neutrons. The second column shows the charges of the three subatomic particles in the SI unit of Coulombs (a Coulomb is a unit that we use to measure charge, just like a kilogram is a unit that we use to measure mass, and a meter is a unit that we use to measure distance). The third column, however, shows the charges of the three subatomic particles using "elementary charge units"* or "elementary charges". Elementary charge units (e) are appealing, because the charge on a proton and the charge on an electron are exactly 1.0 in this unit system.

Table 4.2: Charges on the Different Subatomic Particles
  Charge in Coulombs (C) Mass in Elementary Charges (e)
Electron −1.6021765×10−19 −1
Proton 1.6021765×10−19 1
Neutron 0 0

Notice that whether you use Coulombs or elementary charge units, when you ignore the positive and negative signs, the charge on the proton and the charge on the electron have the same magnitude.

Previously, you learned that negative and positive charges of equal magnitude cancel each other out. This means that the negative charge on an electron perfectly balances the positive charge on the proton. In other words, a neutral atom must have exactly one electron for every proton. If a neutral atom has 1 proton, it must have 1 electron. If a neutral atom has 2 protons, it must have 2 electrons. If a neutral atom has 10 protons, it must have 10 electrons. You get the idea. In order to be neutral, an atom must have the same number of electrons and protons, but what kinds of numbers are we talking about? That's what we'll look at in the next section.

*Most scientists don't refer to "elementary charges" as a unit. Nevertheless, if you treat elementary charges just like you'd treat any another non-SI unit, like a pound (lb) or a foot (ft), they become a lot easier to understand.

Atomic Number (Z) Identifies the Element[edit]

How do you tell the difference between a bike and a car? What about the difference between a car and a unicycle? Take a look at Figure 4.13.

Torker Unicycle.JPG Lexus LF-A Pic 1.JPG Stevens-s6pro.jpg
Red smart car.jpg Honda 250 CBR.jpg
Figure 4.13: A unicycle, two examples of cars, and two examples of bikes. Can you think of some rule that might allow you to tell all unicycles, cars and bikes apart?

If you had to make a rule to distinguish between a unicycle, a bike, a car, what would it be? You can't use color, because different cars can be different colors and, even worse, a car can be the same color as a bike or unicycle. The same goes for weight. While most cars would weigh more than most bikes, which would weigh more than most unicycles, that isn't always the case. In fact, that the little red "Smart Car" in Figure 4.13 probably weighs less than a large motorbike.

What you really need to distinguish between a car, a bike and a unicycle is a property that is the same within each category, but different between the categories. A good choice would be the number of wheels. All unicycles have one wheel, all bikes have two wheels, and all cars have four wheels. If you count wheels, you will most likely never confuse a unicycle with a bike, or a bike with a car (even a motorbike with a Smart Car!). In other words, if you know the number of wheels, you know which type of vehicle you're dealing with.

Just as we can tell between cars, bikes, and unicycles by counting the number of wheels, scientists can tell between different elements (remember, an element is a specific type of atom) by counting the number of protons. If a vehicle has only one wheel, we know it's a unicycle. If an atom has only one proton, we know it's a hydrogen atom or, said differently, it's an atom of the element hydrogen. Similarly, a vehicle with two wheels is always a bike, just like an atom with two protons is always a helium atom, or an atom of the element helium. When we count four wheels on a vehicle, we know it's a car, and when scientists count four protons in an atom, they know it's a beryllium atom, or an atom of the element beryllium. The list goes on: an atom with three protons is a lithium atom, an atom with five protons is a boron atom, an atom with six protons is a carbon atom… in fact, we have names for atoms containing everything from 1 proton all the way up to 118 protons. So far, the maximum number of protons scientists have been able to pack into a single atom is 118, and thus there are 118 known elements. (On Earth, only atoms with a maximum of 92 protons occur naturally.)

Since an atom of one element can be distinguished from an atom of another element by the number of protons in its nucleus, scientists are always interested in this number, and how this number differs between different elements (Figure 4.14). Therefore, scientists give this number a special name and a special symbol. An element's atomic number (Z) is equal to the number of protons in the nuclei of any of its atoms. The atomic number for hydrogen is Z = 1, because every hydrogen atom has 1 proton. The atomic number for helium is Z = 2 because every helium atom has 2 protons. What's the atomic number of carbon?

Sulfur
Silicon
Gold
Figure 4.14: You can't really distinguish between sulfur and gold based on color because both are yellowish. You could say that gold was shiny, but then how would you tell the difference between gold and silicon? Each element, however, does have a unique number of protons. Sulfur has 16 protons, silicon has 14 protons, and gold has 79 protons.

Mass Number (A) is the Sum of Protons and Neutrons[edit]

In the last section we learned that each type of atom or element has a specific number of protons. This specific number was called the element's atomic number. Of course, since neutral atoms have to have one electron for every proton, an element's atomic number also tells you how many electrons are in a neutral atom of that element. For example, hydrogen has atomic number Z = 1. This means that an atom of hydrogen has one proton, and, if it's neutral, one electron as well. Gold, on the other hand, has atomic number Z = 79, which means that an atom of gold has 79 protons if it's neutral, and 79 electrons as well. So we know the number of protons, and we know the number of electrons, but what about the third type of subatomic particle? What about the number of neutrons in an atom?

The number of neutrons in an atom isn't important for determining atomic number; in fact, it doesn't even tell you which type of atom (or which element) you have. The number of neutrons is important, though, if you want to find a quantity known as the mass number (A). The mass number of any atom is defined as the sum of the protons and neutrons in the atom:

\text{mass number}\ A = (\text{number of protons}) + (\text{number of neutrons})\,\!

An atom's mass number is a very easy to calculate provided you know the number of protons and neutrons in an atom.

Example 1

What is the mass number of an atom that contains 3 protons and 4 neutrons?


Solution:

(\text{number of protons}) =\,\! 3\,\!
(\text{number of neutrons}) =\,\! 4\,\!
\text{mass number},\ A =\,\! (\text{number of protons}) + (\text{number of neutrons})\,\!
A =\,\! (3) + (4)\,\!
A =\,\! 7\,\!
Example 2

What is the mass number of an atom of helium that contains 2 neutrons?


Solution:

(\text{number of protons}) =\,\! 2 \ [\text{Remember that an atom of helium always has 2 protons.}]\,\!
(\text{number of neutrons}) =\,\! 2\,\!
\text{mass number},\ A =\,\! (\text{number of protons}) + (\text{number of neutrons})\,\!
A =\,\! (2) + (2)\,\!
A =\,\! 4\,\!

Why do you think that the "mass number" includes protons and neutrons, but not electrons? You have already learned that the mass of an electron is very, very small compared to the mass of either a proton or a neutron (like the mass of a penny compared to the mass of a bowling ball). Counting the number of protons and neutrons tells scientists about the total mass of an atom, but counting the number of electrons would only confuse things.

Think of it this way – you're asked to lift a box containing some bowling balls and some pennies, but the box has already been taped closed. Now, if you have to decide whether or not to get help lifting the box, which would you prefer to know, the total number of bowling balls and pennies, or the just the total number of bowling balls (Figure 4.15)? Suppose you were told only the number of bowling balls. If you knew that there were 20 bowling balls in the box, you wouldn't lift the box on your own, but if you knew that there was only 1, you probably would, even if that box contained 19 pennies that you didn’t know about. On the other hand, if, instead of being told the number of bowling balls, you were told the number bowling balls and pennies, your decision would be more difficult. What if you were given the number 20? That could mean 20 bowling balls and no pennies, or it could mean 1 bowling ball and 19 pennies. In fact, it could even mean 20 pennies. Unfortunately, you would have no way of knowing what was meant by the number 20. Certainly, you wouldn't choose to lift 20 bowling balls, but lifting 20 pennies would be no problem. Just like you wouldn't care about the number of pennies in the box you were about to lift, scientists don't care about the number of electrons when they calculate the mass number. That's why the mass number is only the sum of the protons and neutrons in the atom.

Figure 4.15: Each of the boxes above contains a total of 10 items. If you had to choose one to lift, though, you'd want to know the number of bowling balls in the each box, not the total number of items in each box. Obviously, you'd rather lift the box with 2 bowling balls than the box with 7 bowling balls.

Isotopes Have Varying Numbers of Neutrons[edit]

If you were reading the last section carefully, you'll already know that you can't use the number of neutrons in an atom to decide which type of atom (or which element) you have. Unlike the number of protons, which is always the same in atoms of the same element, the number of neutrons can be different, even in atoms of the same element. Atoms of the same element, containing the same number of protons, but different numbers of neutrons are known as isotopes. Since the isotopes of any given element all contain the same number of protons, they have the same atomic number (for example, the atomic number of helium is always 2). However, since the isotopes of a given element contain different numbers of neutrons, different isotopes have different mass numbers. The following two examples should help to clarify this point.

Example 3

What is the atomic number (Z) and the mass number of an isotope of lithium containing 3 neutrons? A lithium atom contains 3 protons in its nucleus.


Solution:

\text{atomic number},\ Z =\,\! (\text{number of protons}) = 3\,\!
(\text{number of neutrons}) =\,\! 3\,\!
\text{mass number},\ A =\,\! (\text{number of protons}) + (\text{number of neutrons})\,\!
A =\,\! (3) + (3)\,\!
A =\,\! 6\,\!
Example 4

What is the atomic number (Z) and the mass number of an isotope of lithium containing 4 neutrons? A lithium atom contains 3 protons in its nucleus.


Solution:

\text{atomic number},\ Z =\,\! (\text{number of protons}) = 3\,\!
(\text{number of neutrons}) =\,\! 4\,\!
\text{mass number},\ A =\,\! (\text{number of protons}) + (\text{number of neutrons})\,\!
A =\,\! (3) + (4)\,\!
A =\,\! 7\,\!

Notice that because the lithium atom always has 3 protons, the atomic number for lithium is always Z = 3. The mass number, however, is A = 6 in the isotope with 3 neutrons, and A = 7 in the isotope with 4 neutrons.

In nature, only certain isotopes exist. For instance, lithium exists as an isotope with 3 neutrons, and as an isotope with 4 neutrons, but it doesn't exists as an isotope with 2 neutrons, or as an isotope with 5 neutrons. Scientists can make isotopes of lithium with 2 or 5 neutrons, but they aren't very stable (they fall apart easily), so they don't exist outside of the laboratory.

Atomic Mass is a Calculated Value[edit]

Of course, this whole discussion of isotopes brings us back to Dalton's Atomic Theory. According to Dalton, atoms of a given element are identical. But if atoms of a given element can have different numbers of neutrons, then they can have different masses as well! How did Dalton miss this? It turns out that elements found in nature always exist as constant uniform mixtures of their naturally occurring isotopes. In other words, a piece of lithium always contains both types of naturally occurring lithium (the type with 3 neutrons and the type with 4 neutrons). Moreover, it always contains the two in the same relative amounts (or "relative abundances"). In a chunk of lithium, 93% will always be lithium with 4 neutrons, while the remaining 7% will always be lithium with 3 neutrons.

Unfortunately, Dalton always experimented with large chunks of an element – chunks that contained all of the naturally occurring isotopes of that element. As a result, when he performed his measurements, he was actually observing the averaged properties of all the different isotopes in the sample. Luckily, aside from having different masses, most other properties of different isotopes are similar. As a result, the fact that atoms of a given element aren't, strictly speaking, identical, isn't all that important for most chemistry problems.

Knowing about the different isotopes is important, however, when it comes to calculating atomic mass. The atomic mass of an element is the average mass of the masses of its isotopes and their relative percentages, and is typically given in "atomic mass units" (u). (Remember that an "atomic mass unit" is a convenient unit to use when studying atoms, because a proton is almost exactly 1.0 u). You can calculate the atomic mass of an element provided you know the relative abundances the element's naturally occurring isotopes, and the masses of those different isotopes. The examples below show how this is done.

Example 5

Boron has two naturally occurring isotopes. In a sample of boron, 20% of the atoms are B-10, which is an isotope of boron with 5 neutrons and a mass of 10 amu. The other 80% of the atoms are B-11, which is an isotope of boron with 6 neutrons and a mass of 11 amu. What is the atomic mass of boron?


Solution:

To do this problem, we will calculate 20% of the mass of B-10, which is how much the B-10 isotope contributes to the "average boron atom". We will also calculate 80% of the mass of B-11, which is how much the B-11 isotope contributes to the "average boron atom".

Step One: Convert the percentages given in the question into their decimal forms by dividing each by 100:

\text{Decimal form of } 20% = \frac{20}{100} = 0.20
\text{Decimal form of } 80% = \frac{80}{100} = 0.80

Step Two: Multiply the mass of each isotope by its relative abundance (percentage) in decimal form:

20% \text{ of the mass of } B - 10 - 0.20 \times 10\,\text{amu} = 2.0\,\text{amu}\,\!
80% \text{ of the mass of } B - 11 - 0.80 \times 11\,\text{amu} = 8.8\,\text{amu}\,\!

Step Three: Find the total mass of the "average atom" by adding together the contributions from the different isotopes:

\text{Total mass of average atom} = 2.0\,\text{u} + 8.8\,\text{u} = 10.8\,\text{u}\,\!

The mass of an average boron atom, and thus boron's atomic mass, is 10.8 u.

Example 6

Neon has three naturally occurring isotopes. In a sample of neon, 90.92% of the atoms are Ne−20, which is an isotope of neon with 10 neutrons and a mass of 19.99 u. Another 0.3% of the atoms are Ne−21, which is an isotope of neon with 11 neutrons and a mass of 20.99 u. The final 8.85% of the atoms are Ne−22, which is an isotope of neon with 12 neutrons and a mass of 21.99 u. What is the atomic mass of neon?


Solution:

To do this problem, we will calculate 90.9% of the mass of Ne − 20, which is how much Ne − 20 contributes to the "average neon atom". We will also calculate 0.3% of the mass of Ne − 21 and 8.8% of the mass of Ne − 22, which are how much the Ne − 21 isotope and the Ne − 22 isotope contribute to the "average neon atom" respectively.

Step One: Convert the percentages given in the question into their decimal forms by dividing each by 100:

\text{Decimal form of } 90.92% = \frac{90.92}{100} = 0.9092
\text{Decimal form of } 0.30% = \frac{0.30}{100} = 0.0030
\text{Decimal form of } 8.85% = \frac{8.85}{100} = 0.0885

Step Two: Multiply the mass of each isotope by its relative abundance (percentage) in decimal form:

90.92% \text{ of the mass of Ne} - 20 = 0.909 \times 20.00 = 18.18\,\text{amu}\,\!
0.3% \text{ of the mass of Ne} - 21 = 0.003 \times 21.00 = 0.063\,\text{amu}\,\!
8.85% \text{ of the mass of Ne} - 22 = 0.088 \times 22.00 = 1.93\,\text{amu}\,\!

Step Three: Find the total mass of the "average atom" by adding together the contributions from the different isotopes:

\text{Total mass of average atom} = 18.18\,\text{amu} + 0.06\,\text{amu} + 1.93\,\text{amu} = 20.17\,\text{amu}\,\!

The mass of an average neon atom, and thus neon's atomic mass, is 20.17 amu.

Atomic Information in the Periodic Table[edit]

Most scientists don't want to have to calculate the atomic mass of an element every time they do an experiment. Nor do they want to memorize the number of protons, or the atomic number, of each of the 118 elements that have been discovered. As a result, this information is stored in the Periodic Table. Figure 4.16 shows a Periodic Table that contains more detail than the Periodic Table you saw back in Chapter 1.

Group → 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
↓ Period
1 1
H
1.01

2
He
4.00
2 3
Li
6.94
4
Be
9.01

5
B
10.1
6
C
12.0
7
N
14.0
8
O
16.0
9
F
19.0
10
Ne
20.2
3 11
Na
23.0
12
Mg
24.3

13
Al
27.0
14
Si
28.1
15
P
31.0
16
S
32.1
17
Cl
35.5
18
Ar
40.0
4 19
K
39.1
20
Ca
40.1
21
Sc
45.0
22
Ti
47.9
23
V
50.9
24
Cr
52.0
25
Mn
54.9
26
Fe
55.9
27
Co
58.9
28
Ni
58.7
29
Cu
63.6
30
Zn
65.4
31
Ga
69.7
32
Ge
72.6
33
As
74.9
34
Se
79.0
35
Br
79.9
36
Kr
83.8
5 37
Rb
85.5
38
Sr
87.6
39
Y
88.9
40
Zr
91.2
41
Nb
92.9
42
Mo
95.9
43
Tc
98
44
Ru
101
45
Rh
103
46
Pd
106
47
Ag
108
48
Cd
112
49
In
115
50
Sn
119
51
Sb
122
52
Te
128
53
I
127
54
Xe
131
6 55
Cs
133
56
Ba
137
57
La
139
72
Hf
179
73
Ta
181
74
W
184
75
Re
186
76
Os
190
77
Ir
192
78
Pt
195
79
Au
197
80
Hg
201
81
Tl
204
82
Pb
207
83
Bi
209
84
Po
209
85
At
210
86
Rn
222
7 87
Fr
233
88
Ra
226
89
Ac
227
104
Rf
263
105
Db
262
106
Sg
266
107
Bh
264
108
Hs
269
109
Mt
268
110
Ds
272
111
Rg
272
112
Cn
277
113
Uut
284
114
Uuq
289
115
Uup
288
116
Uuh
292
117
Uus
292
118
Uuo
294
Lanthanides 58
Ce
140
59
Pr
141
60
Nd
144
61
Pm
145
62
Sm
150
63
Eu
152
64
Gd
157
65
Tb
159
66
Dy
163
67
Ho
165
68
Er
167
69
Tm
169
70
Yb
173
71
Lu
175
Actinides 90
Th
232
91
Pa
231
92
U
238
93
Np
237
94
Pu
244
95
Am
243
96
Cm
247
97
Bk
247
98
Cf
251
99
Es
252
100
Fm
257
101
Md
258
102
No
259
103
Lr
260

Natural occurrence

  • solid borders: at least one isotope is older than the Earth (Primordial elements)
  • dashed borders: at least one isotope naturally arises from decay of other chemical elements and no isotopes are older than the earth
  • dotted borders: only artificially made isotopes (synthetic elements)

Figure 4.16: A periodic table showing both the atomic number (Z) of each element and the mass number (A) of each element.

Notice that each box still contains the symbol (a capital letter or a capital letter followed by a lower case letter) for one of the elements, but now there are two new numbers that have been added to each square, one number above the element's symbol, and another number below the element's symbol.

The number above the element's symbol in each square is the element's atomic number. Just as you learned previously, hydrogen (symbol H) has atomic number Z = 1, helium (symbol He) has atomic number Z = 2, lithium (symbol Li) has atomic number Z = 3, beryllium (symbol Be) has atomic number Z = 4, boron (symbol B) has atomic number Z = 5, and carbon (symbol C) has atomic number Z = 6. The number below the element's symbol in each square is the element's atomic mass. Notice that atomic mass of boron (symbol B) is 10.8, which is what we calculated in example 5, and the atomic mass of neon (symbol Ne) is 20.18, which is what we calculated in example 6. Observe how compactly the Periodic Table stores and presents a large amount of information about each element. Take time to notice that not all Periodic Tables have the atomic number above the element's symbol and the mass number below it. If you are ever confused, remember that the atomic number (Z) should always be the smaller of the two, while the atomic mass should always be the larger of the two. (The average mass must include both the number of protons (Z) and the average number of neutrons).

Lesson Summary[edit]

  • Electrons are a type of subatomic particle with a negative charge. As a result, electrons repel each other, but are attracted to protons.
  • Protons are a type of subatomic particle with a positive charge. As a result, protons repel each other, but are attracted to electrons. Protons are bound together in an atom's nucleus as a result of the strong nuclear force.
  • Neutrons are a type of subatomic particle with no charge (they’re neutral). Like protons, neutrons are bound into the atom's nucleus as a result of the strong nuclear force.
  • Protons and neutrons have approximately the same mass, but they are both much more massive than electrons (approximately 2,000 times as massive as an electron).
  • The positive charge on a proton is equal in magnitude ("size when you ignore positive and negative signs") to the negative charge on an electron. As a result, a neutral atom must have an equal number of protons and electrons.
  • Each element has a unique number of protons. An element's atomic number (Z) is equal to the number of protons in the nuclei of any of its atoms.
  • The mass number (A) of an atom is the sum of the protons and neutrons in the atom: mass number A = (number of protons) + (number of neutrons)
  • Isotopes are atoms of the same element (same number of protons) that have different numbers of neutrons in their atomic nuclei.
  • An element's atomic mass is the average mass of one atom of that element. An element's atomic mass can be calculated provided the relative abundances of the element's naturally occurring isotopes, and the masses of those isotopes are known.
  • The periodic table is a convenient way to summarize information about the different elements. In addition to the element's symbol, most periodic tables will also contain the element’s atomic number (Z), and element's atomic mass.

Review Questions[edit]

  1. Decide whether each of the following statements is true or false.
    (a) The nucleus of an atom contains all of the protons in the atom.
    (b) The nucleus of an atom contains all of the neutrons in the atom.
    (c) The nucleus of an atom contains all of the electrons in the atom.
    (d) Neutral atoms of a given element must contain the same number of neutrons.
    (e) Neutral atoms of a given element must contain the same number of electrons.
  2. Match the subatomic property with its description.
    (a) electron i. has an atomic charge of +1e
    (b) neutron ii. has a mass of 9.109383×10−28 grams
    (c) proton iii. is neither attracted to, nor repelled from charged objects
  3. Arrange the electron, proton, and neutron in order of decreasing mass.
  4. Decide whether each of the following statements is true or false.
    (a) An element's atomic number is equal to the number of protons in the nuclei of any of its atoms.
    (b) The symbol for an element's atomic number is (A).
    (c) A neutral atom with Z = 4 must have 4 electrons.
    (d) A neutral atom with A = 4 must have 4 electrons.
    (e) An atom with 7 protons and 7 neutrons will have A = 14.
    (f) An atom with 7 protons and 7 neutrons will have Z = 14.
    (g) A neutral atom with 7 electrons and 7 neutrons will have A = 14.
  5. Use the periodic table to find the symbol for the element with:
    (a) 44 electrons in a neutral atom
    (b) 30 protons
    (c) Z = 36
    (d) an atomic mass of 14.007 amu
  6. When will the mass number (A) of an atom be…
    (a) bigger than the atomic number (Z) of the atom?
    (b) smaller than the atomic number (Z) of the atom?
    (c) equal to the atomic number (Z) of the atom?
  7. Column One contains data for 5 different elements. Column Two contains data for the same 5 elements, however different isotopes of those elements. Match the columns by connecting isotopes of the same element.
    Table 4.3
    Column One Column Two
    a. an atom with 2 protons and 1 neutron i. a C (carbon) atom with 6 neutrons
    b. a Be (beryllium) atom with 5 neutrons ii. an atom with 2 protons and 2 neutrons
    c. an atom with Z = 6 and A = 13 iii. an atom with Z = 7 and A = 15
    d. an atom with 1 proton and A = 1 iv. an atom with A = 2 and 1 neutron
    e. an atom with Z = 7 and 7 neutrons v. an atom with Z = 4 and 6 neutrons
  8. Match the following isotopes with their respective mass numbers.
    (a) an atom with Z = 17 and 18 neutrons i. 35
    (b) an H atom with no neutrons ii. 4
    (c) A He atom with 2 neutrons iii. 1
    (d) an atom with Z = 11 and 11 neutrons iv. 23
    (e) an atom with 11 neutrons and 12 protons v. 22
  9. Match the following isotopes with their respective atomic numbers.
    (a) a B (boron) atom with A = 10 i. 8
    (b) an atom with A = 10 and 6 neutrons ii. 2
    (c) an atom with 3 protons and 3 neutrons iii. 3
    (d) an oxygen atom iv. 4
    (e) an atom with A = 4 and 2 neutrons v. 5
  10. Answer the following questions:
    (a) What's the mass number of an atom that contains 13 protons and 13 neutrons?
    (b) What's the mass number of an atom that contains 24 protons and 30 neutrons?
  11. Answer the following questions:
    (a) What's the mass number of the isotope of manganese (Mn) containing 28 neutrons?
    (b) What's the mass number of the isotope of calcium (Ca) containing 20 neutrons?
  12. Answer the following questions:
    (a) What's the atomic number of an atom that has 30 neutrons, and a mass number of A = 70?
    (b) What's the atomic number of an atom with 14 neutrons, if the mass number of the atom is A = 28?
  13. Answer the following questions:
    (a) What's the mass number of a neutral atom that contains 7 protons and 7 neutrons?
    (b) What's the mass number of a neutral atom that contains 7 electrons and 7 neutrons?
    (c) What's the mass number of a neutral atom that contains 5 protons, 5 electrons and 6 neutrons?
    (d) What's the mass number of a neutral atom that contains 3 electrons and 4 neutrons
  14. Answer the following questions:
    (a) What element has 32 neutrons in an atom with mass number A = 58?
    (b) What element has 10 neutrons in an atom with mass number A = 19?
  15. Copper has two naturally occurring isotopes. 69.15% of copper atoms are Cu − 63 and have a mass of 62.93 amu. The other 30.85% of copper atoms are Cu − 65 and have a mass of 64.93 amu. What is the atomic mass of copper?

Vocabulary[edit]

atomic mass units (amu)
A unit used to measure the masses of small quantities like protons, neutrons, electrons and atoms. It is useful, because the mass of a proton is very close to 1.0 amu.
atomic number (Z)
An element's atomic number is equal to the number of protons in the nuclei of any of its atoms.
electron
A type of subatomic particle with a negative charge.
elementary charge (e)
The magnitude of charge on one electron or one proton. You can treat elementary charges as a unit of charge.
mass number (A)
The mass number of an atom is the sum of the protons and neutrons in the atom.
neutron
A type of subatomic particle with no charge. Neutrons are found in the nucleus of an atom.
proton
A type of subatomic particle with a positive charge. Protons are found in the nucleus of an atom.
strong nuclear force
The force that holds protons and neutrons together in the nucleus of the atom. The strong nuclear force is strong enough to overcome the repulsion between protons.


Further Understanding of the Atom · The Bohr Model of the Atom