1.2: Introduction To Chemistry and Matter

Learning Objectives

• Discuss matter, elements, molecules, and atoms
• Explain electronegativity
• Explain chemical bonding and compounds
• Explain chemical equations and reactions
• Discuss the physical states of matter

Chemistry is an essential aspect of Water Technologies. Technology can be defined as “the application of science to industrial and commercial objectives.” In return, science is “the observation, identification, description, experimental investigation, and theoretical explanation of natural phenomena.” As for chemistry, it is “the science of the composition, structure, properties, and reactions of matter, especially of atomic and molecular systems.”

This chapter presents the fundamentals of chemistry, starting with an introduction to matter and its elemental constituents, i.e., elements, atoms, molecules, and compounds. Molecular arrangements and chemical bonding are then introduced, followed by examples of chemical nomenclature.

Composition of Matter

The universe is made of two things: energy and matter. Matter can be viewed as anything that has a mass and occupies space, i.e., that has a specific volume. Mass is defined as a measurement of the quantity of matter present.

Mass is different than weight: weight is a force that can be calculated by the product of the mass times the acceleration of gravity.

Elements and Atoms

At the center of all matter are elements. Elements are basic substances that cannot be broken down without altering their basic identities; they cannot be further simplified (e.g., hydrogen, H; oxygen, O). An atom is the smallest amount of an element. The center of each atom contains a nucleus made of protons (very small particles with a positive electric charge) and neutrons (very small particles, without electrical charge), with electrons (very small, negatively charged particles) that gravitate around the nucleus. Electrons have an insignificant mass compared to protons and neutrons.

Figure \(\PageIndex{1}\): Electron Shells Gravitating Around Sodium Atom – Image by Greg Robson is licensed under CC BY-SA 2.0

Electrons gravitate around the nucleus of protons and neutrons in layers or shells (Figure 1.2.1). There is only a limited number of electrons per shell, as shown in Figure 1.2.2.

Thus, the maximum number of shells for the element that contains the highest number of electrons is seven (7).

Figure \(\PageIndex{2}\): Maximum Number of Electrons per Shell – Image by COC OER is licensed under CC BY 4.0

Atoms have the same number of protons and electrons, thus, atoms have a neutral electrical charge. However, most atoms tend to gain or lose electrons to complete their last electron shell and obtain a stable electron configuration. Only the Noble Gas (such as helium, neon, and argon) do not tend to gain or lose electrons because their electron configuration is naturally stable. An atom with an unequal number of protons and electrons is called an ion. Atoms that lose an electron(s) become positively charged and are called cation (e.g., sodium, Na+). Atoms that gain an electron(s) become negatively charged and are called anion (e.g., chloride, Cl-). Note that only electrons are gained, lost, or shared because they are readily available; only radioactive compounds can release protons and neutrons.

Cations: Positively charged ions Anions: Negatively charged ions

The Periodic Table illustrates all elements that have been found or were synthesized to date (Figure 1.2.3). In this table, elements are ordered by increasing number of protons (from left to right in each row; rows are called periods) and are grouped in columns (called groups) by electron configuration. For example, chlorine appears as number 17 in the Periodic Table (i.e., its atomic number is 17), which means that it has 17 protons. Chlorine is part of the Halogen Family, which all tend to gain an electron to stabilize their last electron shell. While gaining this electron, they become negatively charged (e.g., chlorine becomes chloride, Cl-). Each column represents a family, e.g., Alkali Metals (Column 1), Alkali Earth Metals (Column 2), Halogens (Column 17), and Noble Gases (Column 18). Some families are named after their first element, e.g., Boron Family (Column 13), Nitrogen Family (Column 15), and Oxygen Family (Column 16). Additional characteristics are presented later.

Figure \(\PageIndex{3}\): Period Table of the Elements - Image is in the public domain

The valence (also called the ionic state or oxidation state) is the number of electrons gained, lost, or shared between atoms. Because all elements of a family share similar electron configuration, they all tend to have the same valence, as follows:

• Noble Gases are stable, and thus, their valence is 0
• Alkali Metals tend to lose one electron, and thus their valence is +1
• Alkali Earth Metals tend to lose two electrons, and thus their valence is +2
• Elements of the Boron Family tend to lose three electrons, and thus their valence is +3
• Halogens tend to gain one electron, and thus their valence is -1
• Elements of the Oxygen Family tend to gain two electrons, and thus their valence is -2
• Elements of the Nitrogen Family tend to gain three electrons, and thus their valence is -3

However, certain elements have multiple valences, with different characteristics based on their valence. For example, ferrous iron, Fe2+, has lost two electrons (it has a valance of +2) and is highly soluble in water. On the other hand, ferric iron, Fe3+ (valence of +3), has lost three electrons and is insoluble in water, i.e., it forms a solid and precipitates. Trivalent chromium (i.e., chromite, also called chromium 3, Cr(III), or Cr3+) has lost three electrons and has a valence of +3. It is an essential element that helps regulate the body’s use of sugar, proteins, and fats. However hexavalent chromium (i.e., chromate, also called chromium 6, Cr(VI), or Cr6+) has lost six electrons (valence of +6) and is toxic to humans.

Electronegativity is the degree of attraction of an element for electrons; it defines an element’s affinity for electrons. Electronegativity determines whether an atom will gain, lose, or share electrons.

Molecules and Compounds

Molecules or compounds result from the combination of two or more atoms that are chemically joined together (or bonded); e.g., oxygen in the air, O2; water, H2O. Atoms will tend to combine in such ways to increase their stability and complete their electron configuration. For some molecules, this means that they will obtain a zero net electrical charge.

Chemical bonds can be grouped into two broad categories, as illustrated in Figures 1.2.4.1 and 1.2.4.2.

Figure \(\PageIndex{4.1}\): Example of a covalent bond - image by OpenStax is licensed under CC BY4.0

Figure \(\PageIndex{4.2}\): Example of an ionic bond – image by OpenStax is licensed under CC BY 4.0

1. In an ionic bond, electrons are transferred from one atom to another. The atom that loses an electron(s) becomes positively charged and is called a cation. Conversely, the atom that gains an electron(s) becomes negatively charged and is called an anion. Examples of ionic bonds are shown in Table 1.2.5.
2. In a covalent bond, electrons are shared between atoms. The electronegativity of each atom will determine the polarity of the resulting molecule:
   1. Homonuclear molecules (i.e., molecules that are composed of only one type of element, such as chlorine, Cl2, or oxygen, O2) are non-polar because each atom has the same electronegativity, or attraction for electrons. Examples are shown in Table 1.2.6.
   2. Heteronuclear molecules are made of different elements, which may have different electronegativities. Because the electrons are not shared equally, the resulting molecule is polarized, i.e., one part of the molecule is slightly more positive and the other part is slightly more negative. Examples are shown in Table 1.2.6, which is found later in this chapter.

Table 1.2.5: Examples of Ionic Bonds
Sodium chloride, NaCl	Sodium tends to lose 1 electron to become Na+
	Chloride tends to gain 1 electron to become Cl-
	Net zero charge: Na+1 + Cl-1 = NaCl0
Sodium oxide, Na2O	Sodium tends to lose 1 electron to become Na+
	Oxygen tends to gain 2 electrons to become O2-
	Net zero charge: (2 x Na+1) + (1 x O2-) = Na2O

Table 1.2.6: Examples of Covalent Bonds
Homonuclear molecules: Equal attraction for the shared electron(s)	Hydrogen, H2: Single covalent bond: H ̶ H
	Chlorine, Cl2: Single covalent bond: Cl ̶ Cl
	Oxygen, O2: Double covalent bond: O = O
	Nitrogen, N2: Triple covalent bond: N ≡ N
Heteronuclear molecule: Unequal attraction for the shared electron(s) Example: Water, H2O	Hydrogen, H+ Valence of +1: Need 1 bond
	Oxygen, O2- Valence of -2: Need 2 bonds
	Net zero charge: (2 x H+1) + (1 x O2-) = H2O

The key differences between ionic and covalent bonds are summarized in Table 1.2.7.

Table 1.2.7: Comparison of Ionic Bonds and Covalent Bonds
Ionic Bonds	Covalent Bonds
Transfer of electron(s) from one atom to another Tend to be inorganic High melting point Often solid at room temperature Good conductor Resulting substance is called a compound	Electrons are shared between atoms Organic compounds Low melting point Solid, liquid or gas at room temperature Poor conductor Resulting substance is called a molecule or molecular compound

Dissociation

Molecules that result from polar covalent bonds (i.e., atoms of different types that share electrons) may breakdown, or dissociate. This is the case for water, H2O, which dissociates into a hydrogen ion, H+, and hydroxide, OH-. The dissociation of water is measured as pH.

Chemical Formulas and Nomenclature

Chemical Formulas

A chemical formula is a shorthand method of describing a chemical substance. There are a number of guidelines that are followed when writing chemical formulas:

• The formula contains the symbols of each element present in the substance, as per the Periodic Table (Figure 1.2.8).
• The formula defines the ratio of the elements present using subscripts to the right of the atom. The subscript indicates that number of atoms of this particular element that are present in the substance. Examples are presented in Figure 1.2.10.
  • Subscripts can also apply to groups of atoms that occur as a unit. In this case, parentheses are used around these groups, and the subscript to the right of the parentheses indicates how many groups are present in the substance. An example is shown in Figure 1.2.10. Note that if there is only 1 atom of a particular element, no subscript is needed.
• The formula typically begins with the cationic atom (positively charged atom), but radicals are kept together. Examples are shown in Table 1.2.9.

Figure \(\PageIndex{8}\): Examples of Basic Chemical Formulas – Image by John Rowe is licensed under CC BY 4.0

As mentioned above, charged chemical species (called ions) can be composed of two or more atoms that act as a single unit. These can also be called polyatomic ions, or molecular ions. A number of common molecular ions encountered in the water industry are presented in Table 1.2.9.

Table 1.2.9: Common Molecular Ions in the Water Industry
Molecular Ion	Composition
Hydroxide ion (OH-)	1 atom of oxygen, O: -2 1 atom of hydrogen, H: +1 Net charge: -1
Carbonate ion (CO32-)	1 atom of carbon, C: +4 3 atoms of oxygen, O: 3 x (-2) = -6 Net charge: -2
Bicarbonate ion (HCO3-)	1 atom of hydrogen, H: +1 1 atom of carbon, C: +4 3 atoms of oxygen, O: 3 x (-2) = -6 Net charge: -1
Ammonium ion (NH4+)	1 atom of nitrogen, N: -3 4 atoms of hydrogen, H: 4 x (+1) = +4 Net charge: +1
Sulfate ion (SO42-)	1 atom of sulfur, S: +6 4 atoms of oxygen, O: 4 x (-2) = -8 Net charge: -2

Figure \(\PageIndex{10}\): Example of Chemical Formula with a Molecular Ion - Image by John Rowe is licensed under CC BY 4.0

Table 1.2.11: Examples of Chemical Formulas with Radicals
Molecules	Acids	Bases
Sodium bicarbonate: NaHCO3 Calcium bicarbonate: Ca(HCO3)2 Aluminum sulfate: Al2(SO4)3 Ammonium carbonate: (NH4)2CO3	Hydrochloric acid: HCl Sulfuric acid: H2SO4 Nitric acid: HNO3	Sodium hydroxide: NaOH Calcium hydroxide: Ca(OH)2

Chemical Names

The Period Table (Figure 1.2.3) presents all the elements that are known to mankind to this day.

When naming chemical molecules and compounds, the cation name is exactly the same as the element name, as listed in the Periodic Table. The anion name, however, is not exactly the same: the end of the element name ends in –ide. For example, chloride is the anion of the element chlorine.

Binary compounds (i.e., compounds that contain only two elements) are typically made of a metal and a non-metal. If these compounds have predictable valences, they also end in –ide, even if one element has multiple atoms; for example, magnesium chloride, MgCl2.

Metals with variable valences, such as the Transition Metals (Columns 3 through 12 of the Periodic Table, Figure 1.2.3), are more complex because they are followed by a symbol that reflects the valence. This was introduced earlier in Section 2.1.1. For example, trivalent chromium (i.e., chromite) has a valence of +3, and is referred to as Cr(III), or Cr3+; hexavalent chromium (i.e., chromate) has a valence of +6 and is referred to as Cr(VI), or Cr6+.

Naming acids and their derivatives is more complex and may depend on the acid’s oxidation state. Generally, the acids with the highest oxidation state end in –ic; e.g., sulfuric acid (H2SO4), nitric acid (HNO3), or phosphoric acid (H3PO4). Their salts end in –ate; e.g., sulfate (SO42-), nitrate (NO3-), or phosphate (PO42-). Acids with the next lowest oxidation state end in –ous; e.g., sulfurous acid, H2SO3. Their salts end in –ite; e.g., sulfite, SO32-. Acids with the lowest oxidation state begin in hypo– and end in –ous; e.g., hypochlorous acid, HOCl. Their salts begin in hypo– and end in –ite; e.g., hypochlorite ion, OCl-.

Chemical Equations

A chemical change is known as a chemical reaction. A chemical reaction identifies the reactants involved, products formed, and indicate the relative amounts of each substance involved. In a chemical reaction, the reactants are separated from the products by specific symbols, such as “→” or “⇌”. The reactants are separated by “+” signs, and likewise for the products, which are also separated by “+” signs. The reactants and products can also be preceded by coefficients that indicate the relative amount of each substance. A coefficient of “1” is considered where no coefficients are shown. Other symbols can also be included above the yield symbol (i.e., “→”) to provide specific information about the chemical reaction. Likewise, certain symbols sometimes appear after chemical formulas to indicate the physical state of the substance, i.e., whether it is a solid, liquid, or a gas. Table 1.2.12 summarizes the symbols that are typically used in chemical reactions.

Table 1.2.12: Typical Symbols Used in Chemical Reactions
Symbol	Meaning	Position
+	Plus, in addition to	Between substances
→	Yields	Between reactants and products
⇌	Equilibrium	Between reactants and products
Δ	Heat	Above the “→”
hν	Light energy	Above the “→”
(s)	Solid	After a chemical formula
(l)	Liquid	After a chemical formula
(g)	Gas	After a chemical formula
(aq)	Aqueous	After a chemical formula
↑	Water escapes	After a chemical formula
↓	Substance precipitates	After a chemical formula

For example, aluminum reacts with iron oxide to form iron and aluminum oxide, according to the following equation:

• Basic chemical equation: Al + Fe₂O₃ → Fe + Al₂O₃
• With proper coefficients: 2 Al + Fe₂O₃ → 2 Fe + Al₂O₃
  • (A coefficient of “1” is considered for Fe₂O₃ and Al₂O₃)

Δ

• Considering the need for heat: 2 Al + Fe₂O₃ → 2 Fe + Al₂O_3f

Δ

• Indication of physical states: 2 Al(s) + Fe₂O₃(s) → 2 Fe(l) + Al₂O₃(s)

Balancing Chemical Equations

Atoms cannot be created or destroyed over the course of chemical reactions. Thus, the number of atoms must be the same for each element and on each side of the equation (i.e., in the reactants and in the products). This is the purpose of balancing chemical equations.

The following steps are involved when balancing chemical equations:

1. Identify the reaction
2. Identify the reactants and products
3. Make sure the chemical formulas are correct
4. Count each atom
5. Use coefficient to balance each substance

Physical States of Matter

Chapter 1 mentioned that “Contaminants in water can be in the gas, dissolved, or solid phase.” These phases can be associated with physical states. The states of matter are distinct forms that matter can take. They are based on how elements are arranged in matter. Matter can take four different states (Figure 1.2.13):

1. Solid: Volume and shape are fixed
2. Liquid: Fixed volume, but variable shape that adapts to its container
3. Gas: Volume and shape are variable
4. Plasma: Variable volume and shape, with electrical charges

The physical states of matter reflect different energy levels: solids have the lowest energy level, followed by liquids. Gases have higher energy level, and plasma have the highest energy level. In water, the first three states (i.e., solid, liquid and gas) are important.

Figure \(\PageIndex{13}\): Physical States of Matter – Image by ElfQrin is licensed under CC BY-SA 4.0

A phase can also be viewed as a homogeneous region of matter. Therefore a homogeneous mixture has only one phase. In comparison, a heterogeneous mixture has more than one phase. It should be noted that substances retain their individual identities within a mixture, and thus, they can be separated using various means. In the water industry, this is done using water treatment processes.

Review Questions

1. Define matter.
2. Define element.
3. Define atom.
4. Define electronegativity.
5. List the types of chemical bonding and define each one.
6. What is a chemical reaction?
7. List the physical states of matter and describe each one.

Chapter Quiz

1. ___________ determines whether an atom will gain, lose, or share electrons.
   1. Covalent bonding
   2. Ionic bonding
   3. Physical states of matter
   4. Electronegativity
2. The physical states of matter reflect different energy levels; ___________ have the lowest energy level.
   1. Solids
   2. Liquids
   3. Gas
   4. Plasma
3. ___________ are basic substances that cannot be broken down without altering their basic identities; they cannot be further simplified.
4. Atoms
5. Molecules
6. Electrons
7. Protons
8. A ___________ is a shorthand method of describing a chemical substance.
   1. Chemical reaction
   2. Reactant
   3. Product
   4. Chemical formula
9. What is defined as the separation of a chemical compound into two or more atoms or simpler compounds?
   1. Synthesis reactions
   2. Combustion reactions
   3. Decomposition reactions
   4. Exchange reactions
10. Identify the product(s): H2SO4 + NaOH → Na2SO4 + H2O
    1. H2SO4
    2. H2SO4 and NaOH
    3. H2O
    4. Na2SO4 and H2O
11. The center of each atom contains a nucleus made of protons (very small particles with a positive electric charge) and neutrons (very small particles, without electrical charge), with electrons (very small, negatively charged particles) that gravitate around the nucleus.
    1. Protons and neutrons
    2. Electrons
    3. Neutrons
    4. Protons
12. ___________ are very small particles with a positive electric charge.
    1. Protons and neutrons
    2. Electrons
    3. Neutrons
    4. Protons
13. ___________ are very small, negatively charged particles.
    1. Protons and neutrons
    2. Electrons
    3. Neutrons
    4. Protons
14. ___________ cannot be created or destroyed over the course of chemical reactions. Thus, the number of ___________ must be the same for each element and on each side of the equation.
    1. Molecules, molecules
    2. Reactants, products
    3. Atoms, atoms
    4. Products, reactants