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1.2: Introduction To Chemistry and Matter

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    6973
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    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.

    Electron shell of Sodium
    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).

    Diagram of the maximum number of electrons per shell
    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.

    Period Table of the Elements
    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.

    In the first image, an oxygen atom is shown with six valence electrons. Four of these valence electrons form pairs at the top and right sides of the valence shell. The other two electrons are alone on the bottom and left sides. A hydrogen atom sits next to each the lone electron of the oxygen. Each hydrogen has only one valence electron. An arrow indicates that a reaction takes place. After the reaction, in the second image, each unpaired electron in the oxygen joins an electron from one of the hydrogen atoms so that the valence rings are now connected together. The bond that forms between oxygen and hydrogen can also be represented by a dash.
    Figure \(\PageIndex{4.1}\): Example of a covalent bond - image by OpenStax is licensed under CC BY4.0
    A sodium and a chlorine atom sit side by side. The sodium atom has one valence electron, and the chlorine atom has seven. Six of chlorine’s electrons form pairs at the top, bottom and right sides of the valence shell. The seventh electron sits alone on the left side. The sodium atom transfers its valence electron to chlorine’s valence shell, where it pairs with the unpaired left electron. An arrow indicates a reaction takes place. After the reaction takes place, the sodium becomes a cation with a charge of plus one and an empty valence shell, while the chlorine becomes an anion with a charge of minus one and a full valence shell containing eight electrons.
    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

    1. 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

    1. 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

    1. Hydrogen, H+

    Valence of +1: Need 1 bond

    1. Oxygen, O2-

    Valence of -2: Need 2 bonds

    1. 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.
    Basic Chemical Formulas
    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

    Example of Chemical Formula with a Molecular Ion
    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 “→”

    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 + Al2O3f

    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.

    Physical States of Matter
    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

    This page titled 1.2: Introduction To Chemistry and Matter is shared under a not declared license and was authored, remixed, and/or curated by John Rowe (ZTC Textbooks) .

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