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Mastering Ionic and Covalent Bonding
Dive into the world of chemical bonds! Explore ionic and covalent bonding, understand their formation, and discover how they shape the properties of compounds. Build a strong foundation for advanced chemistry concepts.

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Now Playing:Ionic and covalent bonding – Example 0a
Intros
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  1. Building on bonding
  2. Building on bonding
    How are atoms joined together?
  3. Building on bonding
    Ionic bonding.
Examples
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  1. Identify whether covalent or ionic bonds will form between these elements.
    Will the following pairs of atoms form a covalent or ionic compound? Explain why.

    i) P and O
    ii) Na and O
    iii) C and Cl
    iv) N and Cl
    Introduction to bonding
    Jump to:NotesConceptFAQsPrerequisites
    Notes
    In this lesson, we will learn:
    • The two major types of bonding in chemical compounds
    • To explain why both types exist and when they are likely to form.
    • The varying properties of ionic and covalent bonding.
    • How to predict the formula of ionic and covalent compounds.

    Notes:

    • In Introduction to bonding, we saw the two broad categories of bonding: intermolecular forces (the forces acting between molecules) and intramolecular forces (the bonds within molecules).
      • Molecules of a substance are held together by intramolecular forces - chemical bonds between the atoms that the substance is made of.
      • For example, CO2 is carbon dioxide and it is carbon dioxide because it is made of one carbon atom bonding to two oxygen atoms, sitting between them in the middle of the molecule. If this arrangement changes, CO2 changes into something else.
      We will not use intramolecular forces anymore. They are chemical bonds, or just bonds.

    • One major type of chemical bond is the ionic bond:
      • Ionic bonding is a bonding attraction caused by oppositely charged ions interacting with each other. It occurs between metal and non-metal atoms.
      • To obtain a full outer shell, metal atoms can donate their outer shell electrons to non-metal atoms. Having lost electrons, the metal atom has become a positive ion and the non-metal atom that gained them is now a negative ion.
      • The transfer of electrons is driven by the stability of having a full outer shell and the difference in electronegativity. The non-metal atom will have a much higher electronegativity value than the metal atom.
      • The force of attraction in an ionic bond is the oppositely charged ions interacting with each other. Electrostatic principles apply: a 2+ metal ion with a 2- nonmetal ion is a stronger ionic bond than 1+ with a 1-.
      • Examples of simple ionic compounds are sodium chloride, NaCl, or magnesium oxide, MgO.
      • Ionic bonds form between elements on opposite sides of the periodic table.

    • Periodic trends like atomic radius and the charge of ions will influence ionic bond strength in compounds. These can be explained in terms of electrostatic theory.
      • Ions with a smaller atomic radius enable any oppositely charged ions to come into closer contact, which increases ionic bond strength. In other words, opposite charges make a stronger bond when they are closer to each other.
      • The larger the charge on an ion, the greater the charge difference will be with the oppositely charged ion, which produces a stronger ionic bond. 2+ will attract 2- stronger than 1+ attracts 1-.

      Nuclear charge affects atomic/ionic radius and it helps explain trends seen in isoelectronic ions (ions with the same number of electrons, e.g. N3-, O2-, F-, Na+, Mg2+, Al3+). In a group of these ions, the ion with the greatest nuclear charge has the smallest atomic radius because the increasingly positive nuclear charge attracts the same number of electrons more strongly.

    • Another major type of chemical bond is the covalent bond:
      • Covalent bonding happens when two atoms share a pair of electrons between them. In any electron pair that make a covalent bond, one electron comes from each atom, hence the name co (cooperating, working together) valent.
      • Covalent bonding occurs between non-metal atoms with similar electronegativity. The similar electronegativity is what leads to the electrons being shared, not gained or lost as in an ionic bond. Because covalent bonding involves atoms gaining electrons by sharing them to complete an outer shell, it is very rare to see metal atoms covalently bond.
      • The driving force of a covalent bond forming is that both atoms have more completion of their outer shell.
      • The force of attraction in a covalent bond is the shared electron pair being attracted to the nuclei of both atoms making the covalent bond.
      • Covalent bonding can involve more than one electron pair:
        • A covalent bond with one electron pair is a single bond.
        • A covalent bond with two electron pairs is a double bond and is almost twice as strong as a single bond.
        • A covalent bond with three electron pairs is a triple bond, almost three times as strong as a single bond.
          Well see why its almost twice/three times as strong later.
      • Covalent bonds can vary in length (measured by the distance between the two nuclei), but in general the stronger the bond, the shorter the bond length.

    • Covalent bonding requires similar, not equal electronegativity. There can be some variety in this electronegativity gap which creates dipoles, or polar covalent bonds. It is not a black and white issue whether you have a covalent or ionic bond; it is a dial which can be turned from 0 to 100.
      • Start by knowing that covalent bonds occur between metal and nonmetal atoms.
      • The larger the gap in electronegativity between two atoms, the more polar their covalent bond will be. A polar bond means the two atoms making the bond have opposing partial charges (δ \delta+ or δ\delta -). They are not full ions; the atoms are just not sharing the electrons equally because the more electronegative atom has a greater pull on them.
      • Because of the opposite partial charges, we have created something of a north/south pole along the bond. We call this a dipole. It is not an ionic bond, but a polar covalent bond. We can also say this covalent bond has slight ionic character.
      • This slightly ionic character means that the two atoms are slightly attracted to one anothers opposing charges, like an ionic bond - at least more than if there was zero ionic character!
        For this reason, bond polarity tends to make covalent bonds stronger.

      Lets take two examples to show the difference:
      • When carbon bonds with bromine, there is a small gap in electronegativity where bromine pulls the electrons with greater force than carbon. This means there is only a very slight partial charge on the two atoms and the electrons are almost equally shared. The relatively bulky bromine atom also makes the covalent bond long; the carbon and bromine nuclei are relatively far away from the electron pair they share.
        This adds up to a weak C-Br covalent bond with very little polarity.

      • When carbon bonds with fluorine, there is a very large electronegativity gap. Fluorine pulls the electrons toward itself substantially more than carbon does. This creates large partial δ \delta+ and δ \delta- charges on carbon and fluorine respectively. The small size of the fluorine atom also means the atoms can be in closer contact and the attractive forces are stronger.
        This all adds up to a very strong C-F covalent bond with significant ionic character.

      When covalent compounds contain polar bonds, the unequal pull of electrons towards some atoms and away from others creates polar molecules. We say that these molecules have permanent dipoles. These are represented with vector arrows that point from the positive to areas of negative charge.

      Polarity in molecules has a major effect on intermolecular forces and the solubility of a substance, which we will see in Intermolecular forces and Polarity.

    • A gap in electronegativity is the easiest way to predict whether ionic or covalent bonding will occur in a chemical substance since it is electronegativity that drives the tendency for electrons to be gained or lost.
      • When two atoms with a large gap in electronegativity form a bond, the atom with a higher electronegativity is going to pull a lot more electron density (think of the electrons as clouds).

      Covalent and ionic bonds occur within molecules, between the atoms and ions that bond to make a substance what it is. They are both strong forces simple ionic compounds often have melting points of several hundred degrees Celsius, as do some large covalent structures.

    • You can use the valence of an atom to work out the formula of covalent and ionic compounds:
      • For each atom, find the valence (number of unpaired) electrons. This will be the group number up to carbons group. After this, electrons pair up in the outer shell, so it will be 8 minus the group number. This is the valence of the atom. (E.g. N has a valence of 8 - 5 = 3 due to 3 unpaired electrons, H has 1).
      • Cross the valence of each atom with the other this is the number of atoms of the element that will combine to form the compound. (N = 1, H = 3 makes NH3, a covalent compound).
      • Remember to take the lowest whole-number ratio of atoms e.g. in the C2O4 example, this can be simplified to CO2, or carbon dioxide.

    Example 1:
    cross valences example 1


    Example 2:
    cross valences example 2

    • (IB) Some covalent bonds are made by both electrons being donated by the same atom. This is a coordinate covalent bond. They are normally made by lone pairs on atoms like oxygen or nitrogen donating to H+ or a metal atom.
      • For example, in the H2O molecule there are two lone pairs of electrons on oxygen that are not making regular covalent bonds to the hydrogen atoms. These lone pairs can form a covalent bond with a hydrogen ion in solution:

        • H2O + H+   \, →  \, H3O+

        This reaction happens whenever an acid is dissolved in water.
    Concept

    Introduction to Ionic and Covalent Bonding

    Welcome to our lesson on ionic and covalent bonding, two fundamental concepts in chemistry. The introductory video provides a crucial foundation for understanding these chemical bonds. Our objectives are to explore the main types of bonding in compounds, their unique properties, and how to predict their formation. Ionic bonding occurs between metals and non-metals, resulting in the transfer of electrons and the formation of charged ions. Covalent bonding, on the other hand, involves the sharing of electrons between non-metal atoms to form molecules. These bonds play a vital role in determining the characteristics of chemical compounds. By the end of this lesson, you'll be able to distinguish between ionic and covalent bonds, predict their occurrence in various substances, and understand their impact on a compound's properties. This knowledge is essential for grasping more advanced chemical concepts and applications in everyday life.

    FAQs

    1. What is the main difference between ionic and covalent bonding?

      Ionic bonding involves the transfer of electrons between atoms, typically from a metal to a non-metal, resulting in oppositely charged ions held together by electrostatic forces. Covalent bonding, on the other hand, involves the sharing of electrons between non-metal atoms to form molecules. Ionic compounds generally have higher melting points and are often soluble in water, while covalent compounds typically have lower melting points and may be less soluble in water.

    2. How can I predict whether a compound will form ionic or covalent bonds?

      To predict bond types, consider the electronegativity difference between the atoms involved. If the difference is greater than 1.7, the bond is likely ionic. For differences between 0.5 and 1.7, expect polar covalent bonds. Differences less than 0.5 usually result in non-polar covalent bonds. Additionally, bonds between metals and non-metals are typically ionic, while those between non-metals are usually covalent.

    3. What are Lewis structures and why are they important?

      Lewis structures are diagrams that show the bonding between atoms in a molecule and the lone pairs of electrons around them. They are important because they help visualize the arrangement of valence electrons, predict molecular geometry, and understand chemical bonding. Lewis structures are crucial for determining the shape of molecules, which in turn affects their properties and reactivity.

    4. How does electronegativity affect bond polarity?

      Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond. When two atoms with different electronegativities form a covalent bond, the more electronegative atom pulls the shared electrons closer, creating a partial negative charge on itself and a partial positive charge on the other atom. This uneven distribution of electron density results in a polar covalent bond. The greater the electronegativity difference, the more polar the bond becomes.

    5. What is the relationship between molecular geometry and polarity?

      Molecular geometry plays a crucial role in determining overall molecular polarity. While individual bonds may be polar, the overall molecule can be non-polar if the bond dipoles cancel each other out due to symmetrical arrangement. For example, carbon dioxide (CO2) has polar C=O bonds, but its linear geometry results in a non-polar molecule. Understanding molecular geometry is essential for predicting polarity, which affects properties like solubility and boiling point.

    Prerequisites

    To fully grasp the concepts of ionic and covalent bonding, it's crucial to have a solid foundation in several key areas of chemistry. Understanding these prerequisite topics will significantly enhance your ability to comprehend the intricacies of chemical bonding and molecular structures.

    One fundamental prerequisite is the history and development of the periodic table. Knowing how elements are arranged and their properties is essential for predicting bonding behavior. The periodic table arrangement provides insights into electron configurations and valence electrons, which are critical in determining the type of bonds elements will form.

    Another important concept to master is molecular geometry and VSEPR theory. This knowledge allows you to predict the three-dimensional shape of molecules, which directly influences their properties and reactivity. Understanding molecular geometry prediction is crucial when studying ionic and covalent bonding, as it helps explain how atoms orient themselves in space and how this affects bond strength and polarity.

    Closely related to bonding is the study of intermolecular forces. These forces play a significant role in determining the physical properties of substances and how molecules interact with each other. Recognizing the different types of intermolecular forces and their relative strengths is essential for understanding why certain compounds form ionic or covalent bonds and how these bonds influence the substance's behavior.

    Lastly, familiarity with conjugation and resonance structures is valuable when delving deeper into covalent bonding. Resonance structures help explain the stability of certain molecules and ions, particularly in organic compounds. This concept is crucial for understanding electron delocalization and its effects on bond strength and reactivity in covalently bonded molecules.

    By mastering these prerequisite topics, you'll be well-equipped to tackle the complexities of ionic and covalent bonding. The periodic table provides the framework for understanding element properties, while molecular geometry helps visualize bond arrangements. Intermolecular forces explain interactions between molecules, and resonance structures offer insights into electron behavior within covalent bonds. Together, these concepts form a comprehensive foundation for exploring the fascinating world of chemical bonding.

    Remember, chemistry is a subject where concepts build upon each other. Taking the time to thoroughly understand these prerequisites will not only make learning about ionic and covalent bonding easier but will also enhance your overall comprehension of chemistry. As you progress in your studies, you'll find that these fundamental concepts continue to be relevant and applicable across various areas of chemistry and related sciences.