This chapter introduces the structure of atoms. As explained in the previous chapter, each element is made of one type of atoms. They determine both the chemical and physical properties of these elements and require a deeper understanding. First, two experimental facts necessary to understand the following chapter are introduced (2.1). This allows the introduction of Dalton's atomic theory (2.2). This rudimentary theory leads the way to more complex topics such as the mass of atoms (2.3), the structure of atoms and their isotopes (2.4 and 2.5). What follows is a discussion of the periodic table (2.6) where relations between different elements are uncovered. Finally, the organization of so called electrons within an atom is discussed (2.7). This leaves us with a basic understanding of the atom and its internal organization. The sections that follow will further enhance this understanding.
2.1 - Chemical Background
Before being able to describe the first atomic theories, two experimental facts known as the Law of Conservation of Mass and Law of Constant Composition have to be stated.
Def 2.1 - Law of Conservation of Mass
Mass is neither created nor destroyed in a chemical reaction. When substances react chemically to create new substances, the total mass of the resulting products is the same as the total mass of the substances that have reacted.
Def 2.2 - Law of Constant Composition
The elements present in a compound are present in fixed and exact proportion by mass, regardless of the compound’s source or method of preparation.
2.2 - Dalton’s Atomic Theory
John Dalton was the first to create a workable general theory of the structure of matter. While this theory is not correct on every particular matter, it provides a basis on which more complex ideas and concepts could be build.
Def 2.3 - Dalton’s Atomic Theory
Dalton’s theory consists of the following propositions:
- All matter is composed of infinitesimally small particles called atoms (believed in Dalton’s time to be indestructible but now known to be composed of even smaller parts).
- The atoms of any one element are chemically identical.
- Atoms of one element are distinguished from those of a different element by the fact that the atoms of the two elements have different masses.
- Compounds are combinations of atoms of different elements and possess properties different from those of their constituent component elements.
- In chemical reactions, atoms are exchanged between starting compounds to form new compounds. Atoms can be neither created nor destroyed.
This is illustrated by the following figure. Atoms (called elementary particles) that are combined into compounds (called compound particles) rearrange to form new compounds. Note that no mass has been lost or gained and that the composition of all compounds is in fact constant.
2.3 - Atomic Masses
In Dalton's atomic theory, the principal difference between elements is the different masses of their atoms. Since one cannot weigh individual atoms to measure their mass, a relative mass scale was proposed. Originally, Dalton assigned a relative mass of 1.0 to the lightest element known; hydrogen. He then assigned masses to other elements by comparing them with hydrogen experimentally. A relative mass scale is still in use today, although the relative masses of atoms are now determined by a technique called mass spectrometry instead of experimentally.
The relative masses of atoms are called the atomic masses of the elements. Atomic masses are now defined relative to the mass of the most common isotope of the element carbon (), whose mass is specified as exactly 12 atomic mass units (, or ).
2.4 - The Structure Of Atoms
It is assumed in Dalton's atomic theory that the atom is featureless and indestructible. This is untrue. An atom is composed of subatomic particles, some of which have an electric charge and some that do not.
Def 2.4 - Subatomic Particles
An atom consists of three types of subatomic particles:
- Protons are particles that are positively charged.
- Electrons are particles that are negatively charged.
- Neutrons are particles that are uncharged.
The charges, atomic mass and location of these subatomic particles can be found in the following table:
Subatomic Particle | Electrical Charge | Mass (amu) | Location |
proton | 1+ | 1.00728 | nucleus |
neutron | 0 | 1.00867 | nucleus |
electron | 1- | 0.005486 | outside nucleus |
The following illustrations shows a view of the atom with its subatomic particles; protons (green), neutrons (yellow) and electrons (red).
It is clear that protons and neutrons account for most of the atom's mass. They are located together in a structure that lies at the center of the atom; the nucleus.
Electrically charged particles repel one another if their charges are the same and attract one another if their charges are opposite. Even though protons and electrons have different masses, the magnitude of their respective charges are the same. That the atom is electrically neutral indicates that the oppositely charged particles are present in equal numbers; the number of protons in an atom is balanced by exactly the same number of electrons. The number of protons in the nucleus is called the atomic number and is unique for each element. Atoms of elements may acquire a charge by gaining or losing electrons (e.g. in reactions with other compounds or elements).
Def 2.5 - Ions
An atom bearing a net electrical charge is called an ion.
- A positively charged ion is called a cation
- A negatively charged ion is called an anion
2.5 - Isotopes
All atoms of a particular element have the same number of protons (i.e. hence the atomic number). However, the number of neutrons in the atoms of that element can vary.
Def 2.6 - Isotopes
An isotope is is a variant of a chemical element that has the same number of protons but a different number of neutrons in its nucleus.
A specific isotope is identified in chemical notation by its mass number. This number is the sum of both protons and neutrons the isotope contains. In most cases, the mass number is written as a superscript. For example, O defines the isotope of oxygen with 17 neutrons in its nucleus. This can also be written as oxygen-17. Evidently, the number of neutrons in an isotope can be easily calculated from the isotope’s mass number and the atomic number of the element.
In nature, any sample of an element will consist of a variety of isotopes of the element. The atomic masses of each element are an average of the atomic masses of the element’s isotopes, based on their natural abundance (i.e. the relative amounts in which they are found in nature).
It is important to distinguish the following concepts:
- Atomic Number ⇒ The number of protons in the nucleus
- Mass Number ⇒ The number of protons and neutrons
- Atomic Mass ⇒ The average of the atomic masses of the isotopes of an element, based on natural abundance
2.6 - The Periodic Table
Def 2.7 - Periodic Law
The properties of elements do not change smoothly and continuously as atomic mass is increased. Instead, these properties are repeated periodically, which is known as the Periodic Law.
This law is embodied in the periodic table (PT). Elements consist of three major categories:
- Metals
- Nonmetals
- Transition elements
In most PTs, each box contains an element, together with their corresponding atomic number. In all versions of the PT, the elements are arranged in such a way that elements with similar chemical properties are aligned in the same vertical column called a group. Each horizontal row is then considered to be a period. A full PT is visualized in the following figure.
- The first column, group I, with the exception of hydrogen, are (alkali) metals, shiny malleable substances that can be easily melted and cast into desired shapes. They are also excellent conductors of heat and electricity.
- The second column, group II, are the alkaline metals. They all have similar chemical properties that can be identified in the same way.
- The elements in the second-to-last column, group VII, are the halogens. These are nonmetals . They cannot be cast into shapes and do not conduct electricity. When they react with other elements, they tend to gain electrons, forming anions.
- The elements in the last column of the PT, group VIII, are called the noble gases (helium, neon, argon, krypton, xenon and radon). They are called noble because they are virtually inert (i.e. under standard conditions, they are chemically unreactive).
Interestingly, the first period only as two members; hydrogen and helium. Hydrogen is not an alkali metal, but does share some properties and is therefor categorized in group I. In the fourth period, a new set of ten elements appears between group II and III. These elements are transition elements. An important property of these elements is that they can form cations of more than a single electrical charge. These transition elements can often be found in biological systems, where electrons are transferred in the course of metabolic processes. There are four series of transition elements, each with increasing complexity from period 4 to 7. Period 6 and 7 contain larger transition groups, each consisting of 14 elements (bottom of the PT) called inner transition elements. Elements with an atomic number between 57 and 70 are called the lanthanide inner transition elements and those between 89 through 102 are called the actinide inner transition elements. They are named after the first element in each series.
The elements within the first two and last five columns are called main-group elements (group I → III). Metals are on the left, while nonmetals are on the right. An intermediate category of elements, called metalloids or semimetals has properties between those of metals and nonmetals.
2.7 - Electron Organization Within The Atom
The visualization of the atom was left with the electrons surrounding the nucleus. The number of electrons is equal to the atom’s atomic number. Until now, it is unclear how these electrons are actually organized. As will become clear, an ordered arrangement of the electrons within atoms further explains the structure of the periodic table and the chemical reactivity of the elements.
Atomic Spectra
Normally, light can be separated or resolved into its component colors or frequencies, called the visible spectrum. This is a series of colors that continuously change and merge into one another. In contrast, when light is emitted by excited or energized atoms or individual elements that have been vaporized in a flame, it does not produce a continuous spectrum but a series of sharply defined and separated lines of different colors.
Def 2.8 - Atomic Emission Spectra
An atomic emission spectrum is the pattern of light wavelengths emitted by an element when its atoms are excited, revealing unique spectral lines that identify the element.
Radiation is absorbed by an atom when that radiation corresponds to the same specific frequency of light emitted from the atom when it is excited in a flame. This way, the emission of light of a specific frequency by excited atoms of an element is complemented by the absorption of radiation of that same frequency when the atoms of the element in an unexcited state are exposed to that frequency of light.
Def 2.9 - Atomic Absorption Spectra
An atomic absorption spectrum is the pattern of wavelengths of light absorbed by atoms in their ground state, used to identify and quantify elements based on the specific wavelengths absorbed.
Atomic spectra can be key to understanding chemical reactivity. This key lies in the relation between the color (frequency) of light and its energy.
Electromagnetic Radiation & Energy
Def 2.10 - (Electromagnetic) Radiation
Radiation describes the transfer of energy from one point in space to another. Heat and visible light are small parts of the full spectrum of electromagnetic (EM) radiation, which has its origin in the oscillation/vibration of charged particles.
The frequency (i.e. how many times per second a vibration is completed) is measured or calculated in cycles per second (cps). The SI unit of frequency is Hertz (Hz).
Type of Radiation | Frequency (Hz) | Energy relative to visible light |
Cosmic rays | 100,000,000 | |
-rays | 1,000,000 | |
X-rays | 10,000 | |
Ultraviolet light | 10 | |
Visible light | 1 | |
Infrared light | 0.1 | |
Microwave radiation | 0.0001 | |
FM radio | 0.0000001 | |
AM radio | 0.000000001 |
The higher the frequency of EM radiation, the higher the energy (this is visible in the table above).
Atomic Energy States
Light consists of small packages called photons. For light to be emitted, energy must be absorbed. The absorption of energy raises an atom from a stable, low energy state (ground state) to a higher-energy state (excited state).
Consequently, the lines of atomic spectra are the result of light emitted when an atom returns from an excited state to its ground state. EM radiation energies that can be absorbed by small systems such as atoms and electrons come in small packages called quanta, the size of which depends on the frequency of the radiation. This means that the energy of atoms can only be increased in discrete units (jumps). If the size of a quantum of energy striking an atom is equal to the energy difference between two of the atom’s energy states, the atom will absorb the energy and enter an excited state. No absorption will take place if the quantum is larger or smaller than the energy difference. In this case, the atom remains in the ground state.
This allows for a new model of the atom, one where atomic energy levels are pictured as a miniature solar system. Electrons (planets) move around the nucleus (which represents the sun). Because the electrons are limited to certain values of energy, they must remain at fixed distances in paths about the nucleus and cannot occupy intermediate positions in these paths. This was an attractive model because it worked for hydrogen, but failed for other elements.
2.8 - The Quantum Mechanical Atom
The modern theory that describes the properties of atoms and subatomic particles is called quantum mechanics. This theory requires to abandon the ability locate an electron with any precision to be abandoned. The best that can be achieved is to estimate the probability of finding an electron in a given region of space. Quantum mechanics can show that electrons are organized within the atom into shells, subshells and orbitals. The relation to one another depends on the energy state of the atom, which is defined by its principal quantum number.
- Shell: The shell is identified by a principal quantum number (1, 2, 3, …, ) that specifies the energy level or energy state of the shell. The higher the quantum number, the greater the energy and the farther the shell electrons are from the nucleus.
- Subshells: Locations within a shell, identified by a lowercase letter s, p, d and f.
- Orbital: The region of space within a subshell that has the highest probability of containing an electron.
Def 2.11 - Atomic Orbital
The region of space in which an electron is most likely to be found is called an atomic orbital.
Atomic orbitals are best visualized as clouds surrounding the nucleus. The size and shape depends on the atom’s energy state; the greater the energy, the larger and more complex the orbitals. Different orbitals are visualized in the following figure (source).
The surface of an orbital encloses the region of space in which there is a high likelihood or probability of finding a given electron. The larger the principal quantum number, the further away electrons are from the nucleus, meaning their orbitals occupy a larger volume (the orbitals do retain their shape). Since orbitals are restricted to certain positions in the space around an atoms nucleus, the connections that they create between atoms gives rise to unique shapes and consequently the characteristic properties of molecules of different substances. Different properties such as number of subshells and orbital types can be found in the following tables.
Principal quantum number (Shell) | Number of subshells | Orbital types |
1 | 1 | s |
2 | 2 | s, p |
3 | 3 | s, p, d |
4 | 4 | s, p, d, f |
Subshell / Orbital type | Number of orbitals |
s | 1 |
p | 3 |
d | 5 |
f | 7 |
For example, an atom in principal quantum state 1 has one s orbital available for electrons. An atom in principal quantum state 2 has one s and three p orbitals available for electrons, and so forth. Note that principal quantum number is always equal to the number of subshells.
The final property electrons possess is spin. A single electron within an orbital is called an unpaired electron or a lone electron. An orbital can only contain two electrons when their spin is opposite. The electrons are then said to be spin-paired. This implies that an orbital can never contain more than two electrons!
Subshell | Number of orbitals | Electrons per orbital | Total electrons / subshell |
s | 1 | 2 | 2 |
p | 3 | 2 | 6 |
d | 5 | 2 | 10 |
f | 7 | 2 | 14 |
2.9 - Atomic Structure and Periodicity
Modern quantum mechanical theory explains the arrangement of elements in the periodic table, as discussed in section 2.6. To better interpret how the periodic table is constructed, the Aufbau procedure can be used. This procedure starts at hydrogen and adds electrons one at a time around an atomic nucleus. To keep the atom neutral, the atomic number (and hence the number of protons) must increase simultaneously every time a new electron is added. The result of the Aufbau procedure is a complete description of the electron organization of the atom.
Def 2.12 - Electron Configuration
The complete description of the electron organization of the atom is called the electron configuration.
Aufbau can be accomplished using the following set of rules:
- The larger the principal quantum number, the greater the number of subshells (see section 2.8). For example, an electron described as 3d is in the third shell in a d subshell.
- Each subshell has a unique number of orbitals.
- The order in which electrons enter available shells and subshells depends on the orbital energy levels. The lowest energy orbitals (i.e. those closest to the nucleus) are filled before the higher energy orbitals.
- There can be no more than two electrons per orbital. The two electrons can only occupy the same orbital if they are spin-paired. This is known as Pauli’s Exclusion Principle.
- In a subshell with more than one orbital (i.e. p, d and f orbitals) electrons entering that subshell will not spin-pair until every orbital in the subshell contains one electron. This is known as Hund’s Rule. All electrons have the same electrical charge and tend to repel each other. This repulsive force can be reduced if the electrons stay as far apart as possible (i.e. by occupying different orbitals when possible).
Example 2.1 - Aufbau Example
In this example the electron configuration of element 16 will be constructed using the Aufbau rules. According to the periodic table, the element with atomic number 16 (i.e. number of protons) is Sulfur. Therefor, to construct a neutral atom, the configuration should consist of 16 electrons.
Shell 1
To start, the first shell can have one orbital of type s, indicating that the first electron will occupy this space. The second electron will spin-pair with the first (according to Pauli’s exclusion principle and Hund’s rule). Since there is only space for one s orbital in the first shell, the second shell will start to be occupied. 14 electrons are left to configure at this step.
Shell 2
The second shell contains one s and three p orbitals (, and ). The s orbital will be occupied first similar to the first shell. Only 12 electrons are now left. Next, according to Hund’s rule, the , and orbitals will first receive three electrons in this order. The next three electrons will then spin-pair respectively on the different p orbitals. The second shell is now saturated. At this point, 6 electrons are left to be configured.
Shell 3
The third shell consists of one s, three p and five d orbitals. Similarly to the second shell, first the only s orbital will be filled, leaving 4 electrons to be configured. Of these 4, the first three will occupy one p orbital each (according to Hund’s rule). Next, using Pauli’s exclusion principle, the final electron will spin-pair with one of these p orbitals, saturating only one of them.
Electron Configuration
To construct the electron configuration, we follow the following principles:
- The first shell is completely occupied →
- The second shell is also completely occupied →
- The third shell is special, since the and orbitals are only filled with one electron, the full configuration for this shell becomes →
The final electron configuration (in full) then becomes:
This method quickly becomes cumbersome because of the space necessary to occupy the notation. The electron configuration can be written more compactly by exploiting the fact that every period in the PT ends with a noble gas, which is an element whose outer shell contains the maximum number of electrons that the shell can hold. For example, neon () has atomic number 10 and thus has 10 electrons. The electron configuration for neon is indeed , where all the p orbitals are fully saturated. The bracketed name of the noble gas can be used to represent its full electron configuration. This allows to shorten the notation for all elements in the period following the noble gas.
Example 2.2 - Short Electron Configuration Notation
Take the element potassium, which has the atomic number 19. Its full electron configuration is
The part that is bracketed is called the argon core of the potassium atom. The argon core can be denoted as . The resulting electron configuration for potassium can thus be expressed as
In section 2.6, three distinct families of elements or groups were discussed; the alkali metals (Group I), the halogens (Group VII) and the noble gases (Group VIII). Their electron configurations are seen in the following table.
Alkali Metals (Group I) | Halogens (Group VII) | Noble Gases (Group VIII) |
Li ⇒ | F ⇒ | Ne ⇒ |
Na ⇒ | Cl ⇒ | Ar ⇒ |
K ⇒ | Br ⇒ | Kr ⇒ |
Rb ⇒ | I ⇒ | Xe ⇒ |
Cs ⇒ | At ⇒ | Rn ⇒ |
Notice that each noble gas has eight electrons in its outer shell, each halogen has seven and each alkali metal has one. Thus, the group an element belongs to indicates the number of electrons in its outer shell. This fact can be summarized by saying that the number of outer-shell electrons is the same as the main group element’s group number.
This fact is easy to verify in the first three periods, but is more complex once the fourth period is entered. This is where the inner 3d subshell begins to fill to form the transition elements after the outer 4s subshell has filled.
2.10 - Atomic Structure, Periodicity and Chemical Reactivity
The method described in the previous section to construct the electron configuration can be used to see the resulting electron configurations of chemical reactions. Take the following reaction as an example:
When this reaction is written in terms of electron configuration, the sodium atom gains an electron (making it a sodium cation) and the chloride atom loses an electron (making it a chloride anion).
Notice that the chemical reaction obtains the following result, which is a general result: the reaction leads to a noble-gas configuration in the atoms’ outer electron shells. Because the noble gases are inert and chemically stable, it can be concluded that a filled outermost electron shell is the most stable configuration that an atom can have. This known as the octet rule (Helium is an exception to this, since it consists of only two electrons).
Def 2.13 - Octet Rule, Valence Shell/Electrons
A filled outermost electron shell is the most stable electron configuration an electron can have. This is known as the octet rule. In an atom
- the outermost electron shell (in a given period) is called the valence shell.
- the electrons in the outermost electron shell are called the valence electrons.
To properly express this, Lewis symbols can be used. The nucleus and inner-shell electrons are represented by the element’s symbol. Dots around it represent the valence-shell electrons. An example for the first three periods is given in the following figure.
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Chapter 3 - Molecules & Chemical Bonds
Sources
- “General, Organic, and Biochemistry” (2nd Edition) - Blei I. and Odian G.
No sources other than the syllabus where used to write this chapter.