the structure of the atom assignment quizlet

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Introduction to solid state chemistry, self-assessment: structure of the atom.

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This self-assessment page completes the Structure of the Atom module, and covers material from the following sessions.

• Session 1: Introduction to Solid State Chemistry
• Session 2: The Periodic Table
• Session 3: Atomic Models: Rutherford & Bohr
• Session 4: Matter/Energy Interactions: Atomic Spectra, Quantum Numbers
• Session 5: Electron Shell Model & Quantum Numbers
• Session 6: Particle-Wave Duality
• Session 7: The Aufbau Principle, Photoelectron Spectroscopy

On this page are a simple weekly quiz and solutions ; relevant exam problems and solutions from the 2009 class; help session videos that review selected solutions to the exam problems; and supplemental exam problems and solutions for further study.

Weekly Quiz and Solutions

This short quiz is given approximately once for every three lecture sessions. You should work through the quiz problems in preparation for the exam problems.

• Quiz problems (PDF)
• Quiz solutions key (PDF)

Exam Problems and Solutions

These exam problems are intended for you to demonstrate your personal mastery of the material, and should be done alone, closed-book, with just a calculator, the two permitted reference tables (periodic table, physical constants), and one 8 1/2" x 11" aid sheet of your own creation.

• Exam problems (PDF)
• Exam solutions key (PDF - 1.9MB)

After you’ve taken the exam, watch the help session videos below for insights into how to approach some of the exam problems.

Exam Help Session Videos

In these videos, 3.091 teaching assistants review some of the exam problems, demonstrating their approach to solutions, and noting some common mistakes made by students.

Clip 1: Exam 1, Problem 1

Clip 2: Exam 1, Problem 2

Clip 3: Exam 1, Problem 4

Clip 4: Exam 1, Problem 6

Clip 5: Exam 2, Problem 2

Supplemental Exam Problems and Solutions

These additional exam problems from prior years’ classes are offered for further study.

• Supplemental exam problems (PDF)
• Supplemental exam solutions key (PDF)

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Learn the Parts of an Atom

The atom is the basic building block of matter. Atoms combine to form pure elements, compounds, and complex forms like computers and phones. Atoms are the smallest particle of matter than cannot be further subdivided using chemical means. In order to understand how atoms interact with each other, you first need to understand the basic parts of an atom. There are 3 main components of atoms:

• Protons – The number of protons in an atom is often denoted by the letter Z. Each proton carries a positive electrical charge. The number of protons determines the type of atom. For example, an element with 1 proton is hydrogen. One with 2 protons is helium. The number of neutrons and electrons don’t have any bearing on the type of atom. In fact, an atom does not even need neutrons or electrons. The most common form of hydrogen consists of a single proton and nothing else. A proton consists of component elementary particles: 2 up quarks and 1 down quark.
• Neutrons – The number of neutrons in an atom is commonly indicated using the letter N. A neutron is about the same size as a proton, but it is electrically neutral. Each neutron consists of 1 up quark and 2 down quarks.
• Electrons – Electrons are extremely tiny, compared with protons or neutrons. The mass of an electron is only 1/1836th that of a proton. Each electron carries a negative electrical charge. Each electron consists of a single elementary particle.

The atomic mass number of an atom is denoted by the symbol A and equals the sum of the number of protons and neutrons or Z + N.

Nucleus and Electron Shell

The parts of each atom are arranged to form a central core or atomic nucleus and an outer set of electron shells.

• Atomic Nucleus – Protons and neutrons collectively are termed nucleons because they bind together to form the nucleus of each atom.
• Electron Cloud – The electron shells describe regions in which an electron is most likely to be located . The innermost shell, the K shell, holds up to 2 electrons. The next shell, the L shell, holds up to 8 electrons. In reality, electrons swiftly spin around the atomic nucleus, but they do not have a well-defined orbit, like the Earth around the Sun. An electron could be just about anywhere (including, briefly, inside the nucleus). The high kinetic energy of the particles forms a sort of outer layer or cloud around an atom. Atoms appear solid, like little balls, because of the motion of electrons. It’s like how a fan blade looks like a solid disc when the fan is in motion.

Opposites Attract

Opposite electrical charges attract each other, so the protons and electrons are drawn toward each other. They don’t meet because the electrons are moving too quickly. It’s sort of like the Moon and the Earth. The Moon is drawn toward the Earth by gravity, but the two bodies don’t crash into each other because of their motion. The Moon is constantly falling around the Earth like electrons are falling around the atomic nucleus.

Ions and Isotopes

Ions and isotopes form when the number of different parts of the atom are changed.

Ions – An ion forms when the number of electrons is different from the number of protons. If there are more protons than electrons, a positively-charged ion called a cation forms. If there are more electrons than protons, the net charge is negative and the species is called an anion . Ions can be formed from single atoms or bound groups of atoms. An ion of a single atom is called an atomic ion.

Isotopes – Isotopes are atoms with the same number of protons, but different numbers of neutrons. There are stable isotopes, where the numbers of protons and neutrons remains the same over time, or radioisotopes. A radioisotope is an unstable or radioactive isotope, which will decay into a more stable isotope, releasing energy and sometimes particles.

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2.5 The Structure of The Atom

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  Learning Objectives

• Write and interpret symbols that depict the atomic number, mass number, and charge of an atom or ion
• Define the atomic mass unit and average atomic mass
• Calculate average atomic mass and isotopic abundance

The development of modern atomic theory revealed much about the inner structure of atoms. It was learned that an atom contains a very small nucleus composed of positively charged protons and uncharged neutrons, surrounded by a much larger volume of space containing negatively charged electrons. The nucleus contains the majority of an atom’s mass because protons and neutrons are much heavier than electrons, whereas electrons occupy almost all of an atom’s volume. The diameter of an atom is on the order of 10 −10 m, whereas the diameter of the nucleus is roughly 10 −15 m—about 100,000 times smaller. For a perspective about their relative sizes, consider this: If the nucleus were the size of a blueberry, the atom would be about the size of a football stadium (Figure \(\PageIndex{1}\)).

Atoms—and the protons, neutrons, and electrons that compose them—are extremely small. For example, a carbon atom weighs less than 2 \(\times\) 10 −23 g, and an electron has a charge of less than 2 \(\times\) 10 −19 C (coulomb). When describing the properties of tiny objects such as atoms, we use appropriately small units of measure, such as the atomic mass unit (amu) and the fundamental unit of charge (e) . The amu was originally defined based on hydrogen, the lightest element, then later in terms of oxygen. Since 1961, it has been defined with regard to the most abundant isotope of carbon, atoms of which are assigned masses of exactly 12 amu. (This isotope is known as “carbon-12” as will be discussed later in this module.) Thus, one amu is exactly \(1/12\) of the mass of one carbon-12 atom: 1 amu = 1.6605 \(\times\) 10 −24 g. (The Dalton (Da) and the unified atomic mass unit (u) are alternative units that are equivalent to the amu.) The fundamental unit of charge (also called the elementary charge) equals the magnitude of the charge of an electron (e) with e = 1.602 \(\times\) 10 −19 C.

A proton has a mass of 1.0073 amu and a charge of 1+. A neutron is a slightly heavier particle with a mass 1.0087 amu and a charge of zero; as its name suggests, it is neutral. The electron has a charge of 1− and is a much lighter particle with a mass of about 0.00055 amu (it would take about 1800 electrons to equal the mass of one proton. The properties of these fundamental particles are summarized in Table \(\PageIndex{1}\). (An observant student might notice that the sum of an atom’s subatomic particles does not equal the atom’s actual mass: The total mass of six protons, six neutrons, and six electrons is 12.0993 amu, slightly larger than the 12.00 amu of an actual carbon-12 atom. This “missing” mass is known as the mass defect, and you will learn about it in the chapter on nuclear chemistry.)

The number of protons in the nucleus of an atom is its atomic number (\(Z\)) . This is the defining trait of an element: Its value determines the identity of the atom. For example, any atom that contains six protons is the element carbon and has the atomic number 6, regardless of how many neutrons or electrons it may have. A neutral atom must contain the same number of positive and negative charges, so the number of protons equals the number of electrons. Therefore, the atomic number also indicates the number of electrons in an atom. The total number of protons and neutrons in an atom is called its mass number (A) . The number of neutrons is therefore the difference between the mass number and the atomic number: A – Z = number of neutrons.

\[\begin{align*} \ce{atomic\: number\:(Z)\: &= \:number\: of\: protons\\ mass\: number\:(A)\: &= \:number\: of\: protons + number\: of\: neutrons\\ A-Z\: &= \:number\: of\: neutrons} \end{align*} \nonumber \]

Atoms are electrically neutral if they contain the same number of positively charged protons and negatively charged electrons. When the numbers of these subatomic particles are not equal, the atom is electrically charged and is called an ion . The charge of an atom is defined as follows:

Atomic charge = number of protons − number of electrons

As will be discussed in more detail later in this chapter, atoms (and molecules) typically acquire charge by gaining or losing electrons. An atom that gains one or more electrons will exhibit a negative charge and is called an anion . Positively charged atoms called cations are formed when an atom loses one or more electrons. For example, a neutral sodium atom (Z = 11) has 11 electrons. If this atom loses one electron, it will become a cation with a 1+ charge (11 − 10 = 1+). A neutral oxygen atom (Z = 8) has eight electrons, and if it gains two electrons it will become an anion with a 2− charge (8 − 10 = 2−).

Example \(\PageIndex{1}\): Composition of an Atom

Iodine is an essential trace element in our diet; it is needed to produce thyroid hormone. Insufficient iodine in the diet can lead to the development of a goiter, an enlargement of the thyroid gland (Figure \(\PageIndex{2}\)).

The addition of small amounts of iodine to table salt (iodized salt) has essentially eliminated this health concern in the United States, but as much as 40% of the world’s population is still at risk of iodine deficiency. The iodine atoms are added as anions, and each has a 1− charge and a mass number of 127. Determine the numbers of protons, neutrons, and electrons in one of these iodine anions.

The atomic number of iodine (53) tells us that a neutral iodine atom contains 53 protons in its nucleus and 53 electrons outside its nucleus. Because the sum of the numbers of protons and neutrons equals the mass number, 127, the number of neutrons is 74 (127 − 53 = 74). Since the iodine is added as a 1− anion, the number of electrons is 54 [53 – (1–) = 54].

Exercise \(\PageIndex{1}\)

An ion of platinum has a mass number of 195 and contains 74 electrons. How many protons and neutrons does it contain, and what is its charge?

78 protons; 117 neutrons; charge is 4+

Chemical Symbols

A chemical symbol is an abbreviation that we use to indicate an element or an atom of an element. For example, the symbol for mercury is Hg (Figure \(\PageIndex{3}\)). We use the same symbol to indicate one atom of mercury (microscopic domain) or to label a container of many atoms of the element mercury (macroscopic domain).

The symbols for several common elements and their atoms are listed in Table \(\PageIndex{2}\). Some symbols are derived from the common name of the element; others are abbreviations of the name in another language. Symbols have one or two letters, for example, H for hydrogen and Cl for chlorine. To avoid confusion with other notations, only the first letter of a symbol is capitalized. For example, Co is the symbol for the element cobalt, but CO is the notation for the compound carbon monoxide, which contains atoms of the elements carbon (C) and oxygen (O). All known elements and their symbols are in the periodic table.

Traditionally, the discoverer (or discoverers) of a new element names the element. However, until the name is recognized by the International Union of Pure and Applied Chemistry (   IUPAC ), the recommended name of the new element is based on the Latin word(s) for its atomic number. For example, element 106 was called unnilhexium (Unh), element 107 was called unnilseptium (Uns), and element 108 was called unniloctium (Uno) for several years. These elements are now named after scientists or locations; for example, element 106 is now known as seaborgium (Sg) in honor of Glenn Seaborg, a Nobel Prize winner who was active in the discovery of several heavy elements.

The symbol for a specific isotope of any element is written by placing the mass number as a superscript to the left of the element symbol (Figure \(\PageIndex{4}\)). The atomic number is sometimes written as a subscript preceding the symbol, but since this number defines the element’s identity, as does its symbol, it is often omitted. For example, magnesium exists as a mixture of three isotopes, each with an atomic number of 12 and with mass numbers of 24, 25, and 26, respectively. These isotopes can be identified as 24 Mg, 25 Mg, and 26 Mg. These isotope symbols are read as “element, mass number” and can be symbolized consistent with this reading. For instance, 24 Mg is read as “magnesium 24,” and can be written as “magnesium-24” or “Mg-24.” 25 Mg is read as “magnesium 25,” and can be written as “magnesium-25” or “Mg-25.” All magnesium atoms have 12 protons in their nucleus. They differ only because a 24 Mg atom has 12 neutrons in its nucleus, a 25 Mg atom has 13 neutrons, and a 26 Mg has 14 neutrons.

Information about the naturally occurring isotopes of elements with atomic numbers 1 through 10 is given in Table \(\PageIndex{2}\). Note that in addition to standard names and symbols, the isotopes of hydrogen are often referred to using common names and accompanying symbols. Hydrogen-2, symbolized 2 H, is also called deuterium and sometimes symbolized D. Hydrogen-3, symbolized 3 H, is also called tritium and sometimes symbolized T.

Atomic Mass

Because each proton and each neutron contribute approximately one amu to the mass of an atom, and each electron contributes far less, the atomic mass of a single atom is approximately equal to its mass number (a whole number). However, the average masses of atoms of most elements are not whole numbers because most elements exist naturally as mixtures of two or more isotopes.

The mass of an element shown in a periodic table or listed in a table of atomic masses is a weighted, average mass of all the isotopes present in a naturally occurring sample of that element. This is equal to the sum of each individual isotope’s mass multiplied by its fractional abundance.

\[\mathrm{average\: mass}=\sum_{i}(\mathrm{fractional\: abundance\times isotopic\: mass})_i \nonumber \]

For example, the element boron is composed of two isotopes: About 19.9% of all boron atoms are 10 B with a mass of 10.0129 amu, and the remaining 80.1% are 11 B with a mass of 11.0093 amu. The average atomic mass for boron is calculated to be:

\[\begin{align*} \textrm{boron average mass} &=\mathrm{(0.199\times10.0129\: amu)+(0.801\times11.0093\: amu)}\\ &=\mathrm{1.99\: amu+8.82\: amu}\\ &=\mathrm{10.81\: amu} \end{align*} \nonumber \]

It is important to understand that no single boron atom weighs exactly 10.8 amu; 10.8 amu is the average mass of all boron atoms, and individual boron atoms weigh either approximately 10 amu or 11 amu.

Example \(\PageIndex{2}\): Calculation of Average Atomic Mass

A meteorite found in central Indiana contains traces of the noble gas neon picked up from the solar wind during the meteorite’s trip through the solar system. Analysis of a sample of the gas showed that it consisted of 91.84% 20 Ne (mass 19.9924 amu), 0.47% 21 Ne (mass 20.9940 amu), and 7.69% 22 Ne (mass 21.9914 amu). What is the average mass of the neon in the solar wind?

\[\begin{align*} \mathrm{average\: mass} &=\mathrm{(0.9184\times19.9924\: amu)+(0.0047\times20.9940\: amu)+(0.0769\times21.9914\: amu)}\\ &=\mathrm{(18.36+0.099+1.69)\:amu}\\ &=\mathrm{20.15\: amu} \end{align*} \nonumber \]

The average mass of a neon atom in the solar wind is 20.15 amu. (The average mass of a terrestrial neon atom is 20.1796 amu. This result demonstrates that we may find slight differences in the natural abundance of isotopes, depending on their origin.)

Exercise \(\PageIndex{2}\)

A sample of magnesium is found to contain 78.70% of 24 Mg atoms (mass 23.98 amu), 10.13% of 25 Mg atoms (mass 24.99 amu), and 11.17% of 26 Mg atoms (mass 25.98 amu). Calculate the average mass of a Mg atom.

We can also do variations of this type of calculation, as shown in the next example.

Example \(\PageIndex{3}\): Calculation of Percent Abundance

Naturally occurring chlorine consists of 35 Cl (mass 34.96885 amu) and 37 Cl (mass 36.96590 amu), with an average mass of 35.453 amu. What is the percent composition of Cl in terms of these two isotopes?

The average mass of chlorine is the fraction that is 35 Cl times the mass of 35 Cl plus the fraction that is 37 Cl times the mass of 37 Cl.

\[\mathrm{average\: mass=(fraction\: of\: ^{35}Cl\times mass\: of\: ^{35}Cl)+(fraction\: of\: ^{37}Cl\times mass\: of\: ^{37}Cl)} \nonumber \]

If we let x represent the fraction that is 35 Cl, then the fraction that is 37 Cl is represented by 1.00 − x .

(The fraction that is 35 Cl + the fraction that is 37 Cl must add up to 1, so the fraction of 37 Cl must equal 1.00 − the fraction of 35 Cl.)

Substituting this into the average mass equation, we have:

\[\begin{align*} \mathrm{35.453\: amu} &=(x\times 34.96885\: \ce{amu})+[(1.00-x)\times 36.96590\: \ce{amu}]\\ 35.453 &=34.96885x+36.96590-36.96590x\\ 1.99705x &=1.513\\ x&=\dfrac{1.513}{1.99705}=0.7576 \end{align*} \nonumber \]

So solving yields: x = 0.7576, which means that 1.00 − 0.7576 = 0.2424. Therefore, chlorine consists of 75.76% 35 Cl and 24.24% 37 Cl.

Exercise \(\PageIndex{3}\)

Naturally occurring copper consists of 63 Cu (mass 62.9296 amu) and 65 Cu (mass 64.9278 amu), with an average mass of 63.546 amu. What is the percent composition of Cu in terms of these two isotopes?

69.15% Cu-63 and 30.85% Cu-65

The occurrence and natural abundances of isotopes can be experimentally determined using an instrument called a mass spectrometer. Mass spectrometry (MS) is widely used in chemistry, forensics, medicine, environmental science, and many other fields to analyze and help identify the substances in a sample of material. In a typical mass spectrometer (Figure \(\PageIndex{5}\)), the sample is vaporized and exposed to a high-energy electron beam that causes the sample’s atoms (or molecules) to become electrically charged, typically by losing one or more electrons. These cations then pass through a (variable) electric or magnetic field that deflects each cation’s path to an extent that depends on both its mass and charge (similar to how the path of a large steel ball bearing rolling past a magnet is deflected to a lesser extent that that of a small steel BB). The ions are detected, and a plot of the relative number of ions generated versus their mass-to-charge ratios (a mass spectrum ) is made. The height of each vertical feature or peak in a mass spectrum is proportional to the fraction of cations with the specified mass-to-charge ratio. Since its initial use during the development of modern atomic theory,   MS has evolved to become a powerful tool for chemical analysis in a wide range of applications.

Video \(\PageIndex{1}\) : Watch this video from the Royal Society for Chemistry for a brief description of the rudiments of mass spectrometry.

An atom consists of a small, positively charged nucleus surrounded by electrons. The nucleus contains protons and neutrons; its diameter is about 100,000 times smaller than that of the atom. The mass of one atom is usually expressed in atomic mass units (amu), which is referred to as the atomic mass. An amu is defined as exactly \(1/12\) of the mass of a carbon-12 atom and is equal to 1.6605 \(\times\) 10 −24 g.

Protons are relatively heavy particles with a charge of 1+ and a mass of 1.0073 amu. Neutrons are relatively heavy particles with no charge and a mass of 1.0087 amu. Electrons are light particles with a charge of 1− and a mass of 0.00055 amu. The number of protons in the nucleus is called the atomic number (Z) and is the property that defines an atom’s elemental identity. The sum of the numbers of protons and neutrons in the nucleus is called the mass number and, expressed in amu, is approximately equal to the mass of the atom. An atom is neutral when it contains equal numbers of electrons and protons.

Isotopes of an element are atoms with the same atomic number but different mass numbers; isotopes of an element, therefore, differ from each other only in the number of neutrons within the nucleus. When a naturally occurring element is composed of several isotopes, the atomic mass of the element represents the average of the masses of the isotopes involved. A chemical symbol identifies the atoms in a substance using symbols, which are one-, two-, or three-letter abbreviations for the atoms.

Key Equations

• \(\mathrm{average\: mass}=\sum_{i}(\mathrm{fractional\: abundance \times isotopic\: mass})_i\)

Paul Flowers (University of North Carolina - Pembroke), Klaus Theopold (University of Delaware) and Richard Langley (Stephen F. Austin State University) with contributing authors.  Textbook content produced by OpenStax College is licensed under a Creative Commons Attribution License 4.0 license. Download for free at http://cnx.org/contents/[email protected] ).

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• 1. Multiple Choice Edit 30 seconds 1 pt What carries no charge  Neutrons Electrons Protons
• 2. Multiple Choice Edit 30 seconds 1 pt The nucleus of an atom is    charged. positively (+) negatively (-) variable not
• 3. Multiple Choice Edit 30 seconds 1 pt A particle that moves around the nucleus is a(n)   . proton ion neutron electron
• 4. Multiple Choice Edit 30 seconds 1 pt The __________ contains the protons and neutrons of an atom, and is most of the mass of an atom Nucleus Atomic Number Atomic Mass Electrons
• 5. Multiple Choice Edit 30 seconds 1 pt The particles that are found in the nucleus of an atom are neutrons and electrons electrons only protons and neutrons protons and electrons
• 6. Multiple Choice Edit 30 seconds 1 pt Which of the following is a negatively charged subatomic particle? Electron Proton Neutron Quark
• 7. Multiple Choice Edit 30 seconds 1 pt All matter is made of what? energy  atoms  electrons  compounds 
• 8. Multiple Choice Edit 30 seconds 1 pt What is the unit used to measure weighted average atomic mass? amu gram angstrom nanogrm
• 9. Multiple Choice Edit 30 seconds 1 pt Contributes basically no mass to an atom. protons neutrons electrons protons and neutrons
• 23. Multiple Choice Edit 30 seconds 1 pt Matter is anything that ____________________. has mass takes up space has mass and takes up space has energy
• 24. Multiple Choice Edit 30 seconds 1 pt What is the outer shell of an atom called? Electron cloud Shell All of the above Orbital 
• 25. Multiple Choice Edit 30 seconds 1 pt Subatomic particle with a positive charge. proton neutron electron

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2.1 What is Organic Chemistry?

2.2 elements, atoms, and the periodic table, elements and abundance, atomic theory, subatomic particles, protons determine the identity of an element, isotopes and atomic mass, electrons and the periodic table of the elements, features of the periodic table, 2.3 chapter summary, 2.4 references.

Have you ever wondered why some plants can be used to make medicines while others are toxic and can kill you? Or why some foods are thought of as healthy while others are bad for you?  Or how beverages like beer, cider and wine are made? This course is designed to introduce the reader to fundamental concepts in Organic Chemistry using consumer products, technologies and services as model systems to teach these core concepts and show how organic chemistry is an integrated part of everyday life.

Organic chemistry is a growing subset of chemistry. To put it simply, it is the study of all carbon-based compounds; their structure, properties, and reactions and their use in synthesis. It is the chemistry of life and includes all substances that have been derived from living systems.  The application of organic chemistry today can be seen everywhere you look, from the plastic making up components of your computer, to nylon which make up your clothes, to macromolecules and cells that make up your very body! Organic chemistry has expanded our world of knowledge and it is an essential part of the fields of medicine, biochemistry, biology, industry, nanotechnology, rocket science, and many more!

To begin our discussions of organic chemistry, we need to first take a look at chemical elements and understand how they interact to form chemical compounds.

(Back to the Top)

An element is a substance that cannot be broken down into simpler chemical substances. There are about 90 naturally occurring elements known on Earth. Using technology, scientists have been able to create nearly 30 additional elements that are not readily found in nature. Today, chemistry recognizes a total of 118 elements which are all represented on a standard chart of the elements, called the Periodic Table of Elements (Figure 2.1). Each element is represented by a one or two letter code, where the first letter is always capitalized and, if a second letter is present, it is written in lowercase. For example, the symbol for Hydrogen is H, and the symbol for carbon is C.  Some of the elements have seemingly strange letter codes, such as sodium which is Na. These letter codes are derived from latin terminology. For example, the symbol for sodium (Na) is derived from the latin word, natrium, which means sodium carbonate.

Figure 2.1: Elements. Some examples of pure elements include  (A) Bismuth, Bi, a heavy metal is used as a replacement for lead and in some medicines, like pepto-bismol, the antidiarrheal and (B) Strontium, Sr, a major component in fireworks.  (C)  All of the elements that have been discovered are represented on the Periodic Table of Elements, which provides an elegant mechanism for not only displaying the elements, but describing many of their characteristics.

The elements vary widely in abundance. In the universe as a whole, the most common element is hydrogen (about 90%), followed by helium (most of the remaining 10%). All other elements are present in relatively minuscule amounts, as far as we can detect. On the planet Earth, however, the situation is rather different. Oxygen makes up 46.1% of the mass of Earth’s crust (the relatively thin layer of rock forming Earth’s surface), mostly in combination with other elements, while silicon makes up 28.5%. Hydrogen, the most abundant element in the universe, makes up only 0.14% of Earth’s crust. Table 2.1 “Elemental Composition of Earth” lists the relative abundances of elements on Earth as a whole and in Earth’s crust. Table 2.2 “Elemental Composition of a Human Body” lists the relative abundances of elements in the human body. If you compare Table 2.1 “Elemental Composition of Earth” and Table 2.2 “Elemental Composition of a Human Body”, you will find disparities between the percentage of each element in the human body and on Earth. Oxygen has the highest percentage in both cases, but carbon, the element with the second highest percentage in the body, is relatively rare on Earth and does not even appear as a separate entry in Table 2.1 “Elemental Composition of Earth”; carbon is part of the 0.174% representing “other” elements. How does the human body concentrate so many apparently rare elements?

The relative amounts of elements in the body have less to do with their abundances on Earth than with their availability in a form we can assimilate. We obtain oxygen from the air we breathe and the water we drink. We also obtain hydrogen from water. On the other hand, although carbon is present in the atmosphere as carbon dioxide, and about 80% of the atmosphere is nitrogen, we obtain those two elements from the food we eat, not the air we breathe.

The modern atomic theory, proposed about 1803 by the English chemist John Dalton, is a fundamental concept that states that all elements are composed of atoms. An atom is the smallest part of an element that maintains the identity of that element. Individual atoms are extremely small; even the largest atom has an approximate diameter of only 5.4 × 10 −10 m. With that size, it takes over 18 million of these atoms, lined up side by side, to equal the width of your little finger (about 1 cm).

Most elements in their pure form exist as individual atoms. For example, a macroscopic chunk of iron metal is composed, microscopically, of individual iron atoms. Some elements, however, exist as groups of atoms called molecules. Several important elements exist as two-atom combinations and are called diatomic molecules. In representing a diatomic molecule, we use the symbol of the element and include the subscript 2 to indicate that two atoms of that element are joined together. The elements that exist as diatomic molecules are hydrogen (H 2 ), oxygen (O 2 ), nitrogen (N 2 ), fluorine (F 2 ), chlorine (Cl 2 ), bromine (Br 2 ), and iodine (I 2 ).

There have been several minor but important modifications to Dalton’s atomic theory. For one thing, Dalton considered atoms to be indivisible. We know now that atoms not only can be divided but also are composed of three different kinds of particles with their own properties that are different from the chemical properties of atoms.

The first subatomic particle was identified in 1897 and called the electron. It is an extremely tiny particle, with a mass of about 9.109 × 10 −31 kg. Experiments with magnetic fields showed that the electron has a negative electrical charge.

By 1920, experimental evidence indicated the existence of a second particle. A proton has the same amount of charge as an electron, but its charge is positive, not negative. Another major difference between a proton and an electron is mass. Although still incredibly small, the mass of a proton is 1.673 × 10 −27 kg, which is almost 2,000 times greater than the mass of an electron. Because opposite charges attract each other (while ‘like’ charges repel each other), protons attract electrons (and vice versa).

Finally, additional experiments pointed to the existence of a third particle, called the neutron. Evidence produced in 1932 established the existence of the neutron, a particle with about the same mass as a proton but with no electrical charge.

We understand now that all atoms can be broken down into subatomic particles: protons, neutrons, and electrons. Table 2.3 “Properties of the Subatomic Particles” lists some of their important characteristics and the symbols used to represent each particle.  Experiment have shown that protons and neutrons are concentrated in a central region of each atom called the nucleus (plural, nuclei). Electrons are outside the nucleus and orbit about it because they are attracted to the positive charge in the nucleus. Most of the mass of an atom is in the nucleus, while the orbiting electrons account for an atom’s size. As a result, an atom consists largely of empty space. (Figure 2.4 and 2.5).

Fig 2.4  The anatomy of an atom.  The protons and neutrons of an atom are found clustered at the center of the atom in a structure called the nucleus.  The electrons orbit the nucleus of the atom within an electron cloud, or the empty space that surrounds the atom’s nucleus.  Note that most of the area of an atom is taken up by the empty space of the electron cloud.

Source: https://upload.wikimedia.org/wikipedia/commons/2/24/Figure_02_01_01.jpg

Fig 2.5 The path of the electron in a hydrogen atom.  Electrons are not in discrete orbits like planets around the sun. Instead there is a probability that an electron may occupy a certain space within the electron cloud (a) The darker the color, the higher the probability that the hydrogen’s one electron will be at that point at any given time. (b) Similarly, the more crowded the dots, the higher the probability that hydrogen’s one electron will be at that point.  In both diagrams, the nucleus is in the center of the diagram.

As it turns out, the number of protons that an atom holds in its nucleus is the key determining feature for its chemical properties.  In short, an element is defined by the number of protons found in its nucleus. The proton number within an element is also called its Atomic Number and is represented by the mathematical term, Z (Fig 2.6). If you refer back to the Periodic Table of Elements shown in figure 2.1, you will see that the periodic table is organized by the number of protons that an element contains. Thus, as you read across each row of the Periodic Table (left to right), each element increases by one proton (or one Atomic Number, Z ).

Fig 2.6 Structure of the Periodic Table. Each element on the periodic table is represented by the atomic symbol (Cu for Copper), the Atomic Number in the upper lefthand corner, and the Atomic Mass in the righthand corner.

Isotopes, Allotropes, and Atomic Mass

How many neutrons are in atoms of a particular element? At first it was thought that the number of neutrons in a nucleus was also characteristic of an element. However, it was found that atoms of the same element can have different numbers of neutrons. Atoms of the same element that have different numbers of neutrons are called isotopes (Fig. 2.7). For example, 99% of the carbon atoms on Earth have 6 neutrons and 6 protons in their nuclei; about 1% of the carbon atoms have 7 neutrons and 6 protons in their nuclei. Naturally occurring carbon on Earth, therefore, is actually a mixture of isotopes, albeit a mixture that is 99% carbon with 6 neutrons in each nucleus. Isotope composition has proven to be a useful method for dating many rock layers and fossils.

Fig 2.7 Isotopes of Hydrogen. All hydrogen atoms have one proton and one electron. However, they can differ in the number of neutrons. (a) Most hydrogen atoms onlycontain one p+ and one e- and no neutrons (b) A small amount of hydrogen exists as the isotope deuterium, which has one proton and one neutron in its nucleus, and (c) an even smaller amount contains one proton and two neutrons in its nucleus and is termed Tritium. Note that Tritium is unstable isotope and will breakdown over time. Thus, Tritium is a radioactive element.

Most elements exist as mixtures of isotopes. In fact, there are currently over 3,500 isotopes known for all the elements. When scientists discuss individual isotopes, they need an efficient way to specify the number of neutrons in any particular nucleus. The atomic mass (A) of an atom is the sum of the numbers of protons and neutrons in the nucleus (Fig. 2.6). Given the atomic mass for a nucleus (and knowing the atomic number, Z , of that particular atom), you can determine the number of neutrons by subtracting the atomic number from the atomic mass.

A simple way of indicating the mass number of a particular isotope is to list it as a superscript on the left side of an element’s symbol. Atomic numbers are often listed as a subscript on the left side of an element’s symbol. Thus, we might see

which indicates a particular isotope of copper. The 29 is the atomic number, Z , (which is the same for all copper atoms), while the 63 is the atomic mass (A) of the isotope. To determine the number of neutrons in this isotope, we subtract 29 from 63: 63 − 29 = 34, so there are 34 neutrons in this atom.

Allotropes of an element are different and separate from the term isotope and should not be confused. Some chemical elements can form more than one type of structural lattice, these different structural lattices are known as allotropes .  This is the case for phosphorus as shown in Figure 2.2. White or yellow phosphorus forms when four phosphorus atoms align in a tetrahedral conformation (Fig 2.8). The other crystal lattices of phosphorus are more complex and can be formed by exposing phosphorus to different temperatures and pressures.  For example, the cage-like lattice of red phosphorus can be formed by heating white phosphorus over 280 o C (Fig 2.8). Note that allotropic changes affect how the atoms of the element interact with one another to form a 3-dimensional structure.  They do not alter the sample with regard to the atomic isotope forms that are present, and DO NOT alter or affect the atomic mass ( A ) of the element.

Different allotropes of different elements can have different physical and chemical properties and are thus, still important to consider. For example, oxygen has two different allotropes with the dominant allotrope being the diatomic form of oxygen, O 2 .  However, oxygen can also exist as O 3 , ozone. In the lower atmosphere, ozone is produced as a by-product in automobile exhaust, and other industrial processes where it contributes to pollution. It has a very pungent smell and is a very powerful oxidant.  It can cause damage to mucous membranes and respiratory tissues in animals.  Exposure to ozone has been linked to premature death, asthma, bronchitis, heart attacks and other cardiopulmonary diseases.  In the upper atmosphere, it is created by natural electrical discharges and exists at very low concentrations. The presence of ozone in the upper atmosphere is critically important as it intercepts very damaging ultraviolet radiation from the sun, preventing it from reaching the Earth’s surface.

Figure 2.8 Allotropes of Phosphorus.   (A) White phosphorus exists as a (B) tetrahedral form of phosphorus, whereas (C) red phosphorus has a more (D) cage-like crystal lattice.  (E) The different elemental forms of phosphorus can be created by treating samples of white phosphorus with increasing temperature and pressure.

Source: https://en.wikipedia.org/wiki/Allotropes_of_phosphorus

Remember that electrons are 2000 times smaller than protons and yet each one contains an equal, but opposing charge. Electrons have a negative charge while protons have a positive charge.  Interestingly, when elements exist in their elemental form, as shown on the periodic table, the number of electrons housed in an atom is equal to the number protons. Therefore, the electric charge of an element cancels itself out and the overall charge of the atom is zero.

Electrons are the mobile part of the atom.  They move and orbit the nucleus of the atom in the electron cloud, the term used for the space around the nucleus. However, they do not move around in random patterns.  Electrons have addresses, or defined orbital spins, within the electron cloud, much the same way our apartment buildings have addresses within our cities. To find the address of an electron, you need to know a little bit about the organization of the electron cloud (…or the city that the electron lives in).

The electron cloud of an atom is divided into layers, called shells, much the way an onion has layers when you peel it. However, it is incorrect to think of a shell as a single layer without thickness and depth to it. A shell has 3-dimensional space within it that contains a wide variety of ‘apartments’ or spaces for the electrons to occupy.  Thus, the shell, or n number, is only the first part of an electron’s address within an atom.  It would be similar to only knowing the neighborhood where your friend lives.  If you only know the neighborhood, it will be difficult to find your friend if you want to take them to dinner.

There are a total of 7 shells (or layers) that an atom can have to house it’s electrons.  If an atom is small, it may only have 1 or 2 shells.  Only very large atoms have all 7 layers. After this point, adding an 8th shell appears to make the atom too unstable to exist…at least we have never found atoms containing an 8th shell!  In the periodic table (Fig. 2.9), you will notice that there are a total of 7 rows on the periodic table (note that the Lanthanide and Actinide rows of elements are generally shown below the main table to make them fit onto one page, but they really belong in the middle of rows 6 and 7 on the periodic table, according to their atomic numbers). Each of these rows represents an electron shell.  Thus, as atoms get larger and house more electrons, they acquire additional shells, up to 7.

Fig 2.9 Structure of the Periodic Table. Each element on the periodic table is represented by the atomic symbol (Cu for Copper), the Atomic Number in the upper lefthand corner, and the Atomic Mass in the righthand corner.

Source: Robson, G.(2006) Wikipedia.  https://en.wikipedia.org/wiki/Electron_shell

Within this textbook, we are not concerned with learning the addresses of all the electrons, but we are very interested about the electrons that are nearest to the surface of the atom, or the ones that are in the outer shell of the atom.  The electrons that are closest to the surface of the atom are the most reactive and are integral in forming bonds between the atoms.  These electrons are said to be housed in the atom’s, valence shell, or the electron shell that is the farthest away from the nucleus of the atom. (or nearest to the surface of the atom).

Elements that have similar chemical properties are grouped in columns called groups (or families). As well as being numbered, some of these groups have names—for example, alkali metals (the first column of elements), alkaline earth metals (the second column of elements), halogens (the next-to-last column of elements), and noble gases (the last column of elements).

Each row of elements on the periodic table is called a period. Periods have different lengths; the first period has only 2 elements (hydrogen and helium), while the second and third periods have 8 elements each. The fourth and fifth periods have 18 elements each, and later periods are so long that a segment from each is removed and placed beneath the main body of the table.

Certain elemental properties become apparent in a survey of the periodic table as a whole. Every element can be classified as either a metal, a nonmetal, or a semimetal, as shown in Figure 2.10 “Types of Elements”. A metal is a substance that is shiny, typically (but not always) silvery in color, and an excellent conductor of electricity and heat. Metals are also malleable (they can be beaten into thin sheets) and ductile (they can be drawn into thin wires). A nonmetal is typically dull and a poor conductor of electricity and heat. Solid nonmetals are also very brittle. As shown in Figure 2.7 “Types of Elements”, metals occupy the left three-fourths of the periodic table, while nonmetals (except for hydrogen) are clustered in the upper right-hand corner of the periodic table. The elements with properties intermediate between those of Another way to categorize the elements of the periodic table is shown in Figure 2.11 “Special Names for Sections of the Periodic Table”. The first two columns on the left and the last six columns on the right are called the main group elements. The ten-column block between these columns contains the transition metals. The two rows beneath the main body of the periodic table contain the inner transition metals. The elements in these two rows are also referred to as, respectively, the lanthanide metals and the actinide metals (Fig 2.11).

Fig 2.10. Types of Elements. Elements are either metals, nonmetals, or semimetals. Each group is located in a different part of the periodic table.

Fig 2.11. Special Names for Sections of the Periodic Table. Some sections of the periodic table have special names. For example, the elements lithium, sodium, potassium, rubidium, cesium, and francium are collectively known as alkali metals. Note that the main group elements do not include the transition metals.

The periodic table is organized on the basis of similarities in elemental properties, but what explains these similarities? It turns out that the arrangement of the columns or families in the Periodic Table reflects how subshells are filled with electrons. Of note, elements in the same column share the same valence shell electron configuration. For example, all elements in the first column have a single electron in their valence shells.  This last observation is crucial. Chemistry is largely the result of interactions between the valence electrons of different atoms. Thus, atoms that have the same valence shell electron configuration will have similar chemistry (Fig 2.12).

Fig 2.12. Number of Valence Shell Electrons. The placement of elements on the periodic table corresponds with the number of valence electrons housed in that element.  Families (columns) on the periodic table all contain the same number of valence shell electrons, which gives them similar chemical properties and reactivities. You can easily count across the main group elements to see the increasing number of electrons in the valence shell.  All of the transition metals have 2 e- in their valence shell, although they also contain an inner orbital subshell that is very close to the valence shell. This gives some of these metals different levels of reactivity. Note that the maximum number of valence shell electrons possible is 8, and that is obtained only by the Noble Gases.

Fig 2.13. Role of iron in oxygen transportation. The hemoglobin protein makes up about 95% of the dry content of the red blood cell and each hemoglobin protein can bind and carry four molecules of oxygen (O 2 ).

Adapted from: https://en.wikipedia.org/wiki/Hemoglobin and https://en.wikipedia.org/wiki/Capillary

Chapter 2 materials have been adapted from the following creative commons resources unless otherwise noted:

1. Organic Chemistry Portal. WikiUniversity. Available at:  https://en.wikiversity.org/wiki/Portal:Organic_chemistry

2. Anonymous. (2012) Introduction to Chemistry: General, Organic, and Biological (V1.0). Published under Creative Commons by-nc-sa 3.0. Available at: http://2012books.lardbucket.org/books/introduction-to-chemistry-general-organic-and-biological/index.html

3. Poulsen, T. (2010) Introduction to Chemistry. Published under Creative Commons by-nc-sa 3.0. Available at: http://openedgroup.org/books/Chemistry.pdf

• CH105: Chapter 1 – Measurements in Chemistry
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• CH105: Chapter 10 – Compounds with Sulfur, Phosphorous, and Nitrogen
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1.2 Structural Organization of the Human Body

Learning objectives.

By the end of this section, you will be able to:

• Describe the structure of the human body in terms of six levels of organization
• List the eleven organ systems of the human body and identify at least one organ and one major function of each

Before you begin to study the different structures and functions of the human body, it is helpful to consider its basic architecture; that is, how its smallest parts are assembled into larger structures. It is convenient to consider the structures of the body in terms of fundamental levels of organization that increase in complexity: subatomic particles, atoms, molecules, organelles, cells, tissues, organs, organ systems, organisms and biosphere ( Figure 1.3 ).

The Levels of Organization

To study the chemical level of organization, scientists consider the simplest building blocks of matter: subatomic particles, atoms and molecules. All matter in the universe is composed of one or more unique pure substances called elements, familiar examples of which are hydrogen, oxygen, carbon, nitrogen, calcium, and iron. The smallest unit of any of these pure substances (elements) is an atom. Atoms are made up of subatomic particles such as the proton, electron and neutron. Two or more atoms combine to form a molecule, such as the water molecules, proteins, and sugars found in living things. Molecules are the chemical building blocks of all body structures.

A cell is the smallest independently functioning unit of a living organism. Even bacteria, which are extremely small, independently-living organisms, have a cellular structure. Each bacterium is a single cell. All living structures of human anatomy contain cells, and almost all functions of human physiology are performed in cells or are initiated by cells.

A human cell typically consists of flexible membranes that enclose cytoplasm, a water-based cellular fluid together with a variety of tiny functioning units called organelles . In humans, as in all organisms, cells perform all functions of life. A tissue is a group of many similar cells (though sometimes composed of a few related types) that work together to perform a specific function. An organ is an anatomically distinct structure of the body composed of two or more tissue types. Each organ performs one or more specific physiological functions. An organ system is a group of organs that work together to perform major functions or meet physiological needs of the body.

This book covers eleven distinct organ systems in the human body ( Figure 1.4 and Figure 1.5 ). Assigning organs to organ systems can be imprecise since organs that “belong” to one system can also have functions integral to another system. In fact, most organs contribute to more than one system.

In this book and throughout your studies of biological sciences, you will often read descriptions related to similarities and differences among biological structures, processes, and health related to a person's biological sex. People often use the words "female" and "male" to describe two different concepts: our sense of gender identity, and our biological sex as determined by our chromosomes, hormones, organs, and other physical characteristics. For some people, gender identity is different from biological sex or their sex assigned at birth. Throughout this book, "female" and "male" refer to sex only, and the typical anatomy and physiology of XX and XY individuals is discussed.

The organism level is the highest level of organization. An organism is a living being that has a cellular structure and that can independently perform all physiologic functions necessary for life. In multicellular organisms, including humans, all cells, tissues, organs, and organ systems of the body work together to maintain the life and health of the organism.

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• Authors: J. Gordon Betts, Kelly A. Young, James A. Wise, Eddie Johnson, Brandon Poe, Dean H. Kruse, Oksana Korol, Jody E. Johnson, Mark Womble, Peter DeSaix
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• Book title: Anatomy and Physiology 2e
• Publication date: Apr 20, 2022
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