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Lesson 4 Ionic and Molecular Compounds

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Course: Science 10 [5 cr] - AB Ed copy 1
Book: Lesson 4 Ionic and Molecular Compounds
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Date: Sunday, 7 September 2025, 6:43 PM

Table of contents


  Introduction


Did you know that approximately 15 000 new substances are added each day to the Chemical Abstracts Service (CAS) registry

BS2.1 Chemist holding a vial
Over the past 50 years, CAS registers, on average, one new substance every 2.5 minutes! Two of the most important concepts in chemistry are becoming familiar with the ways that elements interact to form compounds and knowing how to clearly communicate the name and formula of a compound.

In this section, you will investigate how atoms combine to form compounds. You will look at the different types of compounds and learn how to name and write formulas using IUPAC (International Union of Pure and Applied Chemists) rules. Then you will study the different properties of ionic, molecular, acidic, and basic compounds. 

  Targets

By the end of this section, you will be able to
 
BS2.2 Pen on the periodic table

  1. use the periodic table to explain how and why elements combine to form compounds
    and explain the importance of having a common system for naming compounds
BS2.3 copper on the periodic table

  1. predict formulas and write names for ionic compounds
BS2.4 Molecule

  1. predict formulas and write names for molecular compounds and common acids 
© Wikimedia Commons
BS2.5 Conductivity of various solutions

  1. classify ionic, molecular, acidic, and basic compounds on the basis of their properties; e.g., conductivity, pH, solubility, state, etc.
BS2.6 Skeleton smoking

  1. identify personal and societal issues related to the use of potentially toxic or hazardous compounds

Introduction

Do you know how atoms bond to form compounds?



B4.1 molecule with two hydrogen atoms bonded to an oxygen atom
Compounds are created when elements react together to form chemical bonds. In this lesson, you will investigate why atoms bond and look at two specific types of chemical bonds: covalent bonds, in which non-metal atoms share electrons, and ionic bonds, in which metal atoms and non-metal atoms transfer electrons.

This lesson will give you the knowledge base you need to understand how atoms form ionic and covalent bonds.

  Targets

By the end of this lesson, you will be able to

  • understand the octet rule and its significance in the formation of compounds
  • illustrate how ionic bonds are formed
  • illustrate how covalent bonds are formed
  • identify compounds as ionic or molecular based on the elements they are composed of
  • explain the importance and need of a unified naming system for compounds

  Watch This

Chemical Bonding – Ionic vs. Covalent Bonds @ YouTube RicochetScience


Watch this video to understand why and how atoms bond. Please note that the modified Bohr diagrams in this video do not always show the electrons paired up into orbitals. As well, it would be called a chloride ion, not a chlorine ion. This will help get you in the right mindset for this lesson.

Octet Rule

Why do atoms bond?



B4.2 Modified Bohr diagrams of the first three noble gases
To understand why atoms bond, we will first look at a group of elements that rarely form compounds: the noble gases. Group 18 elements were once known as inert gases, because it was thought that they could not form chemical bonds. Chemists explain the stability (resistance to chemical reaction) of these atoms as resulting from their atomic structures. Specifically, they contain a full valence energy level. With the exception of helium (which only has two valence electrons in the first energy level), having eight valence electrons is considered to be a full and stable (unreactive) valence energy level. Other elements react with each other in order to obtain full outer energy levels, thus making them more stable.

This is called the octet rule.

Elements can form an octet by gaining electrons, losing electrons, or sharing electrons. When an atom gains or loses electrons, it becomes an ion, or charged particle.

  Digging Deeper

© Wikimedia Commons
B4.6 Comparison of atom radius to ion radius

Did you know that cations have a smaller atomic radius than the atom they are formed from? And that anions have a larger atomic radius than the atom they are formed from? https://blausen.com/en/video/loss-and-gain-of-electrons/

Learn More

Non-metal atoms have almost-full valence energy levels and will gain electrons to attain an octet.

As an example:

  • Fluorine has seven valence electrons; so by gaining one electron, it will form an ion with a stable octet.
  • It will no longer be a neutral atom; instead it will be a charged particle called an ion, specifically, an anion.
  • Anions are given unique names:  The element name ending changes to ide (e.g., this example anion is called a fluoride ion).
  • As well, the symbol has the negative charge included (e.g., this example becomes F–).


B4.3 a modified Bohr diagram of fluorine gaining an electron to form a fluoride ion


The charge on an anion indicates the number of electrons gained. An anion with a charge of 2– (e.g., O2–) for example, means that that element must have had six valence electrons and gained two to become stable, or attain a full valence.

Metal atoms have almost-empty valence energy levels and will lose electrons to attain an octet.

As an example:

  • Sodium has one valence electron; so by losing one electron, it will form an ion with a stable octet.
  • It will no longer be a neutral atom; instead it will be a charged particle called an ion, specifically, a cation.
  • Cations retain the element name (e.g., this example cation is called a sodium ion).
  • The symbol has the positive charge included (e.g., this example becomes Na+).


             
B4.4 A modified Bohr diagram of sodium losing an electron to form a sodium ion


The charge on a cation indicates the number of electrons lost. A cation with a charge of 2+ (e.g., Ca2+) for example, means that that element must have had two valence electrons and lost them both to become stable, or attain a full valence. 

B4.5 Two fluorine atoms bonding together to form a molecule with a stable octet

Examples

Each example has a video to go with it. To play the video, click on the play icon next to the example.

  1. Using the octet rule and a modified Bohr diagram,
 


Recall:

The atomic number represents the number of protons in an element, and in a neutral atom, the number of protons equals the number of electrons
Recall:

To determine the number of neutrons, round the atomic molar mass to the closest whole number and subtract the number of protons.


Using the periodic table, determine the number of protons, neutrons, and electrons in the atom.

protons = 12
electrons = 12
neutrons = 12
B4.7 Magnesium on the periodic table
Draw the modified Bohr diagram.
B4.8 Modified Bohr diagram of magnesium
Determine if the atom will gain electrons (does it have five or more valence electrons?) or lose valence electrons (does it have three or less valence electrons?).

Mg will lose two valence electrons.
B4.9 Modified Bohr diagram showing Mg losing two electrons
Draw the modified Bohr diagram of the ion, showing that it lost two valence electrons.
B4.10 Modified Bohr diagram showing Mg as an ion
Name the ion. If it is a cation, it will retain the element name.

This ion is called a magnesium ion.
Write the ion symbol. If the ion lost electrons, it will have a positive charge.

protons (+) = 12
electrons (–) = 10
net charge = 2+

The symbol for this ion is Mg2+.
Example 2
  1. Using the octet rule and a modified Bohr diagram,
 


Using the periodic table, determine the number of protons, neutrons, and electrons in the atom.

protons = 7
electrons = 7
neutrons = 7
B4.11 Nitrogen from the periodic table
Draw the modified Bohr diagram.
B4.12 Modified Bohr diagram for nitrogen
Determine if the atom will gain electrons (does it have five or more valence electrons?) or lose valence electrons (does it have three or less valence electrons?)

N will gain three valence electrons.
B4.13 Modified Bohr diagram showing nitrogen gaining three electrons

Draw the modified Bohr diagram of the ion, showing that it gained three valence electrons.
B4.14 Modified Bohr diagram of a nitride ion
Name the ion. If it is an anion, it will have the ending change to “ide.”

This ion is called a nitride ion.
Write the ion symbol. If the ion gained electrons, it will have a negative charge.

protons (+) = 7
electrons (–) = 10
net charge = 3–

This symbol for this ion is N3–.


Example 3
  1. Using the octet rule and a modified Bohr diagram,


 

Using the periodic table, determine the number of protons, neutrons, and electrons in the atom.

protons = 17
electrons = 17
neutrons = 18
B4.15 Chlorine from the periodic table
Draw the modified Bohr diagram.
B4.16 Modified Bohr diagram of chlorine

Determine if the atom will gain electrons (does it have five or more valence electrons?) or lose valence electrons (does it have three or less valence electrons?)

Cl will gain one valence electron.
B4.17 Modified Bohr diagram of chlorine gaining one electron

Draw the modified Bohr diagram of the ion, showing that it gained one valence electron.
B4.18 Modified Bohr diagram of a chloride ion

Name the ion. If it is an anion, it will have the ending change to “ide.”

This ion is called a chloride ion.
Write the ion symbol. If the ion gained electrons, it will have a negative charge.

protons (+) = 17
electrons (–) = 18
net charge = 1–

The symbol for this ion is Cl–.

    Example 4
    1. Using the octet rule and a modified Bohr diagram,
     



    Using the periodic table, determine the number of protons, neutrons, and electrons in the atom.

    protons = 13
    electrons = 13
    neutrons = 14
    B4.19 Aluminium from the periodic table

    Draw the modified Bohr diagram.
    B4.20 Modified Bohr diagram of aluminium
    Determine if the atom will gain electrons (does it have five or more valence electrons?) or lose valence electrons (does it have three or less valence electrons?)

    Al will lose 3 valence electrons
    B4.21 Modified Bohr diagram of aluminium losing 3 electrons

    Draw the modified Bohr diagram of the ion, showing that it lost three valence electrons.
    B4. 22 Modified Bohr diagram of the aluminium ion

    Name the ion. If it is a cation, it will retain the element name.

    This ion is called an aluminium ion.
    Write the ion symbol. If the ion lost electrons, it will have a positive charge

    protons (+) = 13
    electrons (–) = 10
    net charge = 3+

    The symbol for this ion is Al3+.




      Watch This

    What Are Ions @ YouTube FuseSchool – Global Education 


    This video looks at why ions form.

    There are a few differences to keep in mind when watching the video. 1. In the modified Bohr diagrams they use, they place the electrons separate in the first energy level. 2. At times, they place two electrons in one orbital. 3. They also use an alternate group numbering system (e.g., group 6 is group 16).
     


        Read This

      Please read pages 33, 34, and 35 to 38 in your Science 10 textbook. Make sure you take notes on your readings to study from later. You should focus on the octet rule and how this relates to why ions form. Remember, if you have any questions or you do not understand something, ask your teacher!

        Practice Questions

      Complete the following practice questions to check your understanding of the concept you just learned. Make sure you write complete answers to the practice questions in your notes. After you have checked your answers, make corrections to your responses (where necessary) to study from.

      1. In your own words describe the octet rule.

        Noble gases are stable and unreactive because they have a full valence energy level, consisting of eight valence electrons. Other elements gain or lose electrons to attain a stable octet.
      2. Fill in the missing information in the following chart.

        Ion Name
        		 Number of
        Protons
        		 Modified Bohr
        Diagram
        		 Number of Electrons
        Lost or Gained
        		 Number of
        Electrons in Ion
        		 Ion Symbol

        		3
        		lose 1



        B4.23 Modified Bohr diagram


        calcium 20
        		lose 2

        sulfide







        		gain 3
        		 18





        		Be2+


        Ion Name
        		 Number of
        Protons
        		 Modified Bohr
        Diagram
        		 Number of Electrons
        Lost or Gained
        		 Number of
        Electrons in Ion
        		 Ion Symbol
        lithium 3
        B4.24 Modified Bohr diagram of lithium ion
        		 lose 1
        		 2 Li+
        oxide
        		 8
        B4.25 modified Bohr diagram of oxygen ion
        		 gain 2
        		 10 O2–
        calcium 20
        B4.26 Modified Bohr diagram of calcium ion
        		 lose 2
        		 18 Ca2+
        sulfide 16
        B4.27 Modified Bohr diagram of sulfur ion
        		 gain 2
        		 18 S2–
        phosphide 15
        B4. 28 Modified Bohr diagram of phosphorus ion
        		 gain 3
        		 18 p3–
        beryllium 4
        B4. 29 Modified Bohr diagram of beryllium ion
        		 lose 2
        		 2 Be2+
      3. What do you notice about the number of electrons for each of the ions in question 2? What elements have the same number of electrons in their atoms?

        Notice that the number of electrons in an ion is always 2, 10, or 18. They are the same as the number of electrons of the noble gases.

      Making Connections

        1. Recall that the periodic table got its name from periodicity, or patterns on it. Fill in the following chart using the answers from the practice questions to help you identify patterns on the periodic table.

          Element Group
          Number
          		 Number of
          Valence Electrons
          		 Ion Charge
          sodium
          potassium
          beryllium
          magnesium
          calcium
          aluminium
          nitrogen
          phosphorus
          oxygen
          sulfur
          flourine
          chlorine

          Element Group
          Number
          		 Number of
          Valence Electrons
          		 Ion Charge
          sodium 1 1 1+
          potassium 1 1 1+
          beryllium 2 2 2+
          magnesium 2 2 2+
          calcium 2 2 2+
          aluminium 13 3 3+
          nitrogen 15 5 3–
          phosphorus 15 5 3–
          oxygen 16 6 2–
          sulfur 16 6 2–
          flourine 17 7 1–
          chlorine 17 7 1–

          Notice that all ions in group 1 have a charge of 1+.
          Notice that all ions in group 2 have a charge of 2+.
          Notice that all metallic ions in group 13 have a charge of 3+.
          Notice that all non-metallic ions in group 15 have a charge of 3–.
          Notice that all non-metallic ions in group 16 have a charge of 2–.
          Notice that all ions in group 17 have a charge of 1–.

        Extending Your Knowledge

        The great thing about the periodic table is that the patterns continue! You can only model the first 20 elements using the modified Bohr diagram, but you can extend your knowledge and understanding of the relationship between group numbers, number of valence electrons, and the octet rule to accurately predict ion charges of elements past the first 20!

        1. Fill in the blanks for each of the following questions. You may need to recall the relationship between group number and the number of valence electrons. Organization of Periodic Table
          1. Arsenic is in group ___ and has ___ valence electrons. This means it needs to ____ ____ electrons in order to have a full octet and it will have an ion charge of ___.

            Arsenic is in group 15 and has five valence electrons. This means it needs to gain three electrons in order to have a full octet and it will have an ion charge of 3–.
          2. Barium is in group ___ and has ___ valence electrons. This means it needs to ____ ____ electrons in order to have a full octet and it will have an ion charge of ___.

            Barium is in group 2 and has two valence electrons. This means it needs to lose two electrons in order to have a full octet and it will have an ion charge of 2+.
          3. Iodine is in group ___ and has ___ valence electrons. This means it needs to ____ ____ electron in order to have a full octet and it will have an ion charge of ___.

            Iodine is in group 17 and has seven valence electrons. This means it needs to gain one electron in order to have a full octet and it will have an ion charge of 1–.

        Forming Ionic Compounds with Ionic Bonds

        Are there spare electrons floating around that non-metals can just grab?



        B4.30 Fluorine atom trying to grab an electron
        As you may have already guessed, there are not spare electrons floating around for non-metal atoms to grab. The electrons that non-metal atoms need to gain come from metal atoms that need to lose electrons to achieve a full valence. There is a transfer of electrons from a metallic atom to a non-metallic atom, thus creating an ion. Next, there will be an electrostatic attraction between cations (+) and anions (–). This attraction is an ionic bond and results in the formation of ionic compounds.



        B4.31 Na giving its electron to F



        © Wikimedia Commons
        B4.32 Transfer of an electron from Na to F


          Watch This

        What Are Ionic Bonds @ Youtube FuseSchool – Global Education 


        This video illustrates the formation of ionic bonds. There are a few differences to keep in mind.

        1. In the modified Bohr diagrams they use, they place the electrons separate in the first energy level.
        2. At times, they place two electrons in one orbital.


         


        Forming Ionic Compounds with Ionic Bonds Continued


        Ionic compounds do not form individual molecules but rather large crystal lattices. These are repeating patterns of positive and negative ions.

        The following model uses just the valence electrons to simplify the diagram.

        Please note

        • the transfer of electrons
        • the attraction between cations and anions
        • The large number of ions involved to create a crystal lattice

        Ionic formulas do not represent the actual number of each element in the large crystal structure, or else the formula for sodium chloride might be «math xmlns=¨http://www.w3.org/1998/Math/MathML¨»«msub»«mi»Na«/mi»«mo»§#8734;«/mo»«/msub»«msub»«mi»Cl«/mi»«mo»§#8734;«/mo»«/msub»«/math»! Ionic formulas are always written as the lowest ratio of cations to anions. In this case, it is a 1:1 ratio; so the formula is NaCl.

        Crystal lattices can come in a variety of shapes and ratios of cations to anions.
        B4.33 the formation of a crystal lattice

        calcium flouride
        CaF2
        		 sodium chloride
        NaCl
        		 lithium nitride
        Li3N
        B4.35 Crystal lattice of CaF2
        B4.36 Crystal lattice of NaCl
        B4.36a Crystal lattice of Li3N

          Watch This

        Giant Ionic Structures/Lattices @ YouTube FuseSchool – Global Education 


        This video illustrates the formation of crystal lattice structures and the properties that result from these structures.
         

          Interactive Activity

        Ionic Bonding Tutorial @ ChemThink


        Work through slides 1 to 27. At times, you will need to interact with simulations to continue the slideshow. https://quick.adlc.ca/IonicBonding

        Click on the procedure tab to continue.
        1. Click on the play icon to open the tutorial.
        2. Read the content on slides 1 and 2.

          What happens when you drag the negative ions close together?

          They repel each other.
        3. Read slide 3.

          What happens when you drag the negative ion close to the positive ion?

          They attract each other.
        4. Read slides 4 to 19.

          Describe the structure on slide 19.

          The structure is an alternating arrangement of small red spheres and larger blue spheres, set up in a three-dimensional array.
        5. Read slides 20 and 21.

          Count the number of chloride ions and sodium ions.

          sodium ions = 3
          chloride ions = 3
        6. Read slides 22 to 24. On Slide 24, rotate the crystal.

          Count the number of calcium ions and fluoride ions. What is the formula for the ionic crystal?

          calcium ions = 4
          fluoride ions = 8
          formula = CaF2
        7. Read slides 25 to 27.

        Read This

        Please read pages 34, 35, and 40 to 41 in your Science 10 textbook. Make sure you take notes on your readings to study from later. You should focus on how electrons are transferred to form ionic bonds. Remember, if you have any questions or you do not understand something, ask your teacher!

          Practice Questions

        Complete the following practice questions to check your understanding of the concept you just learned. Make sure you write complete answers to the practice questions in your notes. After you have checked your answers, make corrections to your responses (where necessary) to study from.

        Fill in the blanks for each of the following questions.

        1. Ionic bonds are formed between a __ and a ___.

          2. Ionic bonds are formed between a metal and a non-metal.

        2. An ionic bond is formed by the __ of ___.

          An ionic bond is formed by the transfer of electrons.
        3. A metal atom __ electrons to form a ___, while a non-metal ___ electrons to become an ____.

          A metal atom loses electrons to form a cation, while a non-metal gains electrons to become an anion.
        4. The two oppositely charged ions ___ each other to form an ___ compound.

          The two oppositely charged ions attract each other to form an ionic compound.
        5. Using your own words, explain the formation of an ionic compound, starting with sodium atoms and chlorine atoms.

          Your answer should be a variation of the following. Remember that elements gain or lose electrons when forming compounds so that they achieve a full outer shell. Let's now look at the reaction between sodium and chlorine. A sodium atom loses one electron to achieve a full outer shell, and chlorine gains one electron to complete a full outer shell. So when a sodium atom reacts with a chlorine atom, the sodium atom loses its one electron to chlorine. The two ions formed are a sodium ion, Na+, and a chlorine ion, Cl–. The two ions have opposite charges, and they attract one another. The force of attraction between them is an electrostatic one. This type of attraction is strong. It is called an ionic, or electrovalent, bond.
        6. Explain why Mg25F52 would never be a chemical formula.

          Metals and non-metals join in an ionic structure that is expressed in lowest terms. MgF2 is the correct formula.

        Forming Molecular Compounds with Covalent Bonds

        Did you know that only molecular compounds are composed of molecules?


        Ionic compounds are composed of large three-dimensional arrays of alternating ions, while molecular compounds are composed of individual discrete molecules. These individual molecules are formed when two non-metals bond together. They both need to gain electrons to complete their octets. The only way this can be achieved is if they share their valence electrons. The bond formed by sharing electrons is called a covalent bond.


        B4.35 Two fluorine atoms sharing pair of electrons


        There are a vast number of compounds that exist as molecules: from a relatively simple water molecule to more complex structures such as vitamin B9 (folic acid).


        B4.37 Sharing electrons in a water molecule



        © Wikipedia
        B4.38 Vitamin B9


        What these compounds have in common is the sharing of valence electrons to create stable octets.

        Study these two examples to strengthen your understanding of the covalent bond.

        Ammonia: In each molecule of ammonia, NH3, one nitrogen atom shares electrons with three hydrogen atoms. This way, they all attain full valence energy levels. Note that since the first energy level is the valence for hydrogen, it only needs two electrons instead of the usual eight.


        B4.39 Nitrogen atom sharing electrons with three hydrogen atoms


        Carbon tetrafluoride: Its formula is CF4. In each molecule, one carbon atom shares electrons with four fluorine atoms.


        B4.40 Carbon atom sharing electrons with four fluorine atoms

          Digging Deeper


        B4.36 Two oxygen atoms sharing two pairs of electrons

        B4.36a Two nitrogen atoms sharing three pairs of electrons


        Two atoms can share more than just one pair of electrons. For example, two oxygen atoms can bond and create stable octets by sharing four electrons; this is called a double bond: O=O.

        Two nitrogen atoms share six electrons to create an octet. This is a triple bond: «math xmlns=¨http://www.w3.org/1998/Math/MathML¨»«mi mathvariant=¨normal¨ mathcolor=¨#FFFFFF¨»N«/mi»«mo mathcolor=¨#FFFFFF¨»§#8801;«/mo»«mi mathvariant=¨normal¨ mathcolor=¨#FFFFFF¨»N«/mi»«/math».

          Interactive Activity

        Covalent Bonding Tutorial @ ChemThink

        This interactive slideshow will model the attractions in molecules. Only work through slides 1 to 13 of the 35 slides.

        Click the procedure tab to continue. 

        1. Click on the play icon to open the tutorial. The tutorial can also be accessed at https://quick.adlc.ca/attraction
        2. Read slides 1 and 2.
        3. Move the atoms toward one another. Continue moving them around, slowly.

          Describe what happens when one atom is brought closer to the other.

          The second atom becomes attracted and moves toward the first atom, but it always remains the same distance apart

          Describe what happens when you continue to move the one atom around slowly.

          The second atom “follows along” or is attracted to the first atom and remains the same distance apart, if you move it slowly.

          Why do you think this happens?

          There is an attraction between the two atoms. The electrons of one atom are attracted to the nucleus of the neighbouring atom and vice versa.
        4. Read slides 3 to 13.

          Read This

        Please read pages 46 to 48 in your Science 10 textbook. Make sure you take notes on your readings to study from later. You should focus on the formation of covalent bonds by non-metallic atoms. Remember, if you have any questions or you do not understand something, ask your teacher!

          Practice Questions

        Complete the following practice questions to check your understanding of the concept you just learned. Make sure you write complete answers to the practice questions in your notes. After you have checked your answers, make corrections to your responses (where necessary) to study from.

        1. Covalent bonds are formed by the _____ of electrons.

          Covalent bonds are formed by the sharing of electrons.
        2. What types of elements form covalent bonds?

          non-metallic
        3. Using modified Bohr diagrams, model the covalent bonds that will form in a molecule of oxygen difluoride (OF2).


          B4.41 Oxygen atom sharing electrons with two fluorine atoms

        Why the Names?

        How do we clearly and unmistakably identify compounds?

        B4.42 Spotlight on a stool
        The phrase “in the limelight” means being the centre of attention and was first coined in the 1800s. At that time, a process was invented to spotlight actors on a stage. The process involved heating a piece of quicklime in a flame of burning oxygen and hydrogen. An intense white light was produced that was used to spotlight key actors on the stage.

        You probably understand and recognized the substances hydrogen (H) and oxygen (O) but are at a loss for what quicklime is. This was a common problem in the early days of chemistry. Compounds were not named in a systematic way but instead named randomly by their discoverer. The naming was perhaps based on the location of the discovery, the use of the compound, the method of preparation, the appearance of the substance, or a myriad of other reasons. These random names lead to confusion about the identity of the compounds. For example, if the materials list for a chemistry lab included lye, muriatic acid, oil of vitriol, quicksilver, limestone, saltpeter, and slaked lime would you have any idea what chemicals you were working with and thus what precautions you should take?

          Digging Deeper

        © Wikimedia Commons
        B4.43 Caustic soda

        Visit this website that lists many common names of substances. Their scientific and chemical formulas are included. https://www.spectro.in/Chemical-Compounds.html

        Learn More

        Common names do not indicate the composition of the substance. For example, table salt gives no indication of the constituent elements, whereas the systemic name, sodium chloride, does. Starting in the 1800s, many chemists supported the idea of and put forth recommendations for a naming system. In 1919, the International Union of Pure and Applied Chemists (IUPAC) was established and a set of guidelines were developed for naming substances.

        Naming compounds using a fixed set of rules is important so that all compounds are recognized globally by the same name to avoid confusion and misunderstanding. In Science 10, it is expected that you learn the rules and use IUPAC names, instead of common names, for most compounds. There are a few compounds that have widely recognized and used common names, such as water (H2O), ammonia (NH3), and table sugar (C12H22O11).

          Read This

        Please read pages 40 and 41 in your Science 10 textbook. Make sure you take notes on your readings to study from later. You should focus on the confusion that is created by the use of common compound names. Remember, if you have any questions or you do not understand something, ask your teacher!

          Practice Questions

        Complete the following practice questions to check your understanding of the concept you just learned. Make sure you write complete answers to the practice questions in your notes. After you have checked your answers, make corrections to your responses (where necessary) to study from.

        1. Why was there a need to create a unified set of rules for naming compounds?

          There was a need for unified naming rules to ensure that chemical names are unambiguous (i.e., a name can only refer to one substance.) These standards make it easier to use chemicals; otherwise, scientists everywhere could use different names for the same compound. That could make for some very dangerous situations in chemistry laboratories.
        2. Research! What is the IUPAC name and chemical formula of quicklime?

          calcium oxide, CaO
        3. Research! What is the IUPAC name and chemical formula of lye?

          Lye, commonly used in soapmaking, is used to refer to sodium hydroxide (NaOH) or potassium hydroxide (KOH)

        Conclusion

        Atoms transfer or share electrons to obtain an electron structure with a full valence.

        B4.43 Modified Bohr diagram of neon
        You learned previously that noble gases are chemically stable because of their electronic structure. Noble gases all have a full valence; this is known as the octet rule. Other elements will gain or lose electrons to obtain a full valence. The resulting particles are charged and called ions. Positive ions are cations, and negative ions are anions.

        You also learned that ionic bonds are formed between metals and non-metals by the transfer of electrons and that covalent bonds are formed between non-metallic atoms by sharing electrons. Both of these bonds form to create more stable electron configurations. The atoms involved will end with the same electron arrangement as the nearest noble gas.

          Watch This

        How Atoms Bond @YouTube TED-Ed 


        Watch this video as a review on ionic and covalent bonding.


          Interactive Activity

        Ionic and Covalent Bonding @AACT


        Please work through the following examples to solidify your understanding of chemical bonding. Please note that only the valence electrons are used in these diagrams (electron dot diagrams), as only the valence electrons are involved in chemical bonding.

        Click on the procedure 1 tab to continue.

        1. Click on the play icon to open the interactive. The interactive can also be accessed at https://quick.adlc.ca/periodic-table2
        2. Choose elements K and F from the periodic table to bond.

          What type of bond is this combination likely to form?

          ionic bond
        3. Looking at the number of valence electrons, determine the number of atoms of each element.

          K = 1
          F = 1
        4. Click “Submit Answer.”
        5. Click the reset button: 
        6. Click on the procedure 2 tab to continue
        1. Choose elements N and Cl from the periodic table to bond.

          What type of bond is this combination likely to form?

          covalent bond
        2. Looking at the number of valence electrons, determine the number of atoms of each element.

          N = 1
          Cl = 3
        3. Click “Submit Answer.”
        4. Click the reset button:
        5. Click on the procedure 3 tab to continue

        1. Choose elements Na and S from the periodic table to bond.

          What type of bond is this combination likely to form?

          ionic bond
        2. Looking at the number of valence electrons, determine the number of atoms of each element.

          Na = 2
          S = 1
        3. Click “Submit Answer.”
        4. Click the reset button:
        5. Click on the procedure 4 tab to continue
        1. Choose elements S and F from the periodic table to bond.

          What type of bond is this combination likely to form?

          covalent bond
        2. Looking at the number of valence electrons, determine the number of atoms of each element.

          S = 1
          F = 2
        3.  Click “Submit Answer.”



        2.3 Assignment

        Unit 2 Formative Assessment


        It is now time to complete 2.3 Assignment. Click on the button below to go to the assignment page.

        2.3 Assignment