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Ionic and covalent bonding

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Intros
Lessons
  1. Building on bonding
  2. How are atoms joined together?
  3. Ionic bonding.
  4. Properties and strength of ionic bonding.
  5. Covalent bonding.
  6. Polarity and covalent bonds.
  7. Identifying ionic and covalent compounds.
  8. Predicting the formula of ionic and covalent compounds.
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Examples
Lessons
  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
    1. Identify covalent or ionic compounds and predict their formula using valence.
      Predict whether these atoms will form a covalent or ionic compound together, then predict the formula of their compound:

      i) Mg and O
      ii) C and O
      iii) P and Cl
      iv) Na and F
      Topic Notes
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      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.

      Ionic Bonding: Formation and Properties

      Ionic bonding is a fundamental concept in chemistry that involves the transfer of electrons between atoms, typically between metals and non-metals. This type of chemical bond plays a crucial role in the formation of many compounds we encounter in everyday life. To understand ionic bonding, we must first explore the process of electron transfer and the factors that influence it.

      At the heart of ionic bonding is the transfer of electrons from one atom to another. This transfer occurs when atoms with significantly different electronegativities interact. Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond. Metals, which have low electronegativity, tend to lose electrons, while non-metals, with high electronegativity, tend to gain electrons. This difference in electronegativity drives the electron transfer process.

      When a metal atom encounters a non-metal atom, the metal atom gives up one or more of its outermost electrons to the non-metal atom. This electron transfer results in the formation of oppositely charged ions. The metal atom, having lost electrons, becomes a positively charged cation, while the non-metal atom, having gained electrons, becomes a negatively charged anion. These oppositely charged ions are then held together by strong electrostatic forces, forming an ionic bond.

      A classic example of ionic bonding is the formation of sodium chloride (NaCl), commonly known as table salt. In this case, a sodium atom (Na) transfers its single valence electron to a chlorine atom (Cl). The sodium atom becomes a Na+ cation, while the chlorine atom becomes a Cl- anion. These oppositely charged ions are then attracted to each other, forming a stable ionic compound.

      Another example is the formation of magnesium oxide (MgO). Here, a magnesium atom (Mg) loses two electrons, becoming a Mg2+ cation, while an oxygen atom (O) gains two electrons, becoming an O2- anion. The resulting ionic bond between these ions is particularly strong due to the double charge on each ion.

      The strength of ionic bonds is directly related to the magnitude of the charge on the ions and the distance between them. Generally, the greater the charge difference and the smaller the distance between ions, the stronger the ionic bond. This strong electrostatic attraction between ions gives ionic compounds several distinctive properties.

      One of the most notable properties of ionic compounds is their high melting and boiling points. The strong electrostatic forces between ions require a significant amount of energy to overcome, resulting in substances that are typically solid at room temperature and require high temperatures to melt or vaporize. For instance, sodium chloride has a melting point of 801°C (1474°F), far higher than that of most molecular compounds.

      Another characteristic property of ionic compounds is their electrical conductivity in solution or when molten. In their solid state, ionic compounds do not conduct electricity because the ions are fixed in place within the crystal lattice. However, when dissolved in water or melted, the ions become mobile and can carry an electric current. This property is the basis for many applications, including electrolysis and battery technologies.

      Ionic compounds also tend to be brittle and have a tendency to form crystals. The regular arrangement of ions in a crystal lattice gives these compounds their characteristic shapes. When subjected to stress, the alignment of like charges can cause the crystal to shatter along cleavage planes.

      The concept of electronegativity is crucial in predicting the type of bonding that will occur between atoms. Elements with a large difference in electronegativity (typically greater than 1.7 on the Pauling scale) are likely to form ionic bonds. This difference ensures that electron transfer is energetically favorable. Elements with similar electronegativities, on the other hand, are more likely to form covalent bonds by sharing electrons.

      It's important to note that while we often describe ionic bonding as a complete transfer of electrons, in reality, there is rarely a 100% ionic character to any bond. Most bonds exist on a spectrum between purely ionic and purely covalent, with many exhibiting partial ionic and partial covalent character. The degree of ionic character increases with the electronegativity difference between the bonded atoms.

      In conclusion, ionic bonding is a fundamental concept in chemistry that explains the formation of electrical conductivity in solution.

      Covalent Bonding: Sharing Electrons

      Covalent bonding is a fundamental concept in chemistry that explains how non-metal atoms join together to form molecules. Unlike ionic bonding, which involves the transfer of electrons, covalent bonding is characterized by the sharing of electrons between atoms. This sharing of electrons allows atoms to achieve a stable electron configuration, typically by filling their outermost electron shell.

      At the heart of covalent bonding is electron sharing. When two non-metal atoms come close together, their outer electrons begin to interact. Instead of being transferred from one atom to another, these electrons are shared between the atoms, creating a stable bond. This sharing of electrons results in the formation of molecular compounds, which have distinct properties compared to ionic compounds.

      Covalent bonds can be classified into three main types: single bonds, double bonds, and triple bonds. A single bond involves the sharing of one pair of electrons between two atoms. For example, in a hydrogen molecule (H2), each hydrogen atom contributes one electron, forming a single covalent bond. Double bonds involve the sharing of two pairs of electrons, while triple bonds share three pairs.

      To illustrate covalent bonding, let's consider two common examples: carbon dioxide bonding (CO2) and methane (CH4). In carbon dioxide, the central carbon atom forms two double bonds with oxygen atoms. Each oxygen atom shares two pairs of electrons with the carbon atom, resulting in a linear molecule with the formula O=C=O. This double bonding is crucial for CO2's role in the carbon cycle and its impact on climate change.

      Methane (CH4) provides an excellent example of single covalent bonds. In a methane molecule, a central carbon atom forms four single bonds with hydrogen atoms. Each bond represents a shared pair of electrons, resulting in a tetrahedral structure. This arrangement gives methane its unique properties, including its role as a greenhouse gas and its importance in natural gas.

      The formation of covalent bonds leads to the creation of molecules with specific shapes and properties. The number and arrangement of bonds determine the molecule's molecular geometry, which in turn affects its behavior and interactions with other substances. For instance, the bent shape of a water molecule (H2O) is due to the arrangement of its covalent bonds and contributes to water's unique properties, such as its ability to dissolve many substances.

      One notable characteristic of molecular compounds formed by covalent bonding is their relatively low melting and boiling points compared to ionic compounds. This is because the forces between molecules (intermolecular forces) are generally weaker than the strong electrostatic attractions in ionic compounds. For example, water (H2O) has a much lower melting point than sodium chloride (NaCl), an ionic compound.

      The strength of covalent bonds varies depending on factors such as the size of the atoms involved and the number of electron pairs shared. Generally, triple bonds are stronger than double bonds, which are in turn stronger than single bonds. This variation in bond strength influences the reactivity and stability of different molecules.

      Covalent bonding plays a crucial role in organic chemistry, where carbon atoms form the backbone of complex molecules essential for life. The ability of carbon to form up to four strong covalent bonds allows for the creation of diverse organic compounds, from simple hydrocarbons to complex proteins and DNA.

      Understanding covalent bonding is essential for explaining many natural phenomena and technological applications. For instance, the properties of materials like plastics, pharmaceuticals, and advanced materials are directly related to the covalent bonds within their molecular structures. In biochemistry, covalent bonds are fundamental to the structure and function of biomolecules, including proteins, carbohydrates, and nucleic acids.

      In conclusion, covalent bonding, characterized by electron sharing between non-metal atoms, is a cornerstone of molecular chemistry. From simple diatomic molecules to complex organic compounds, covalent bonds shape the world around us. By understanding the principles of single, double, and triple bonds, and how they contribute to molecular structure and properties, we gain insight into the behavior of matter at the molecular level. This knowledge is crucial for advancing fields such as materials science, drug design, and environmental studies, highlighting the far-reaching implications of this fundamental chemical concept.

      Polar Covalent Bonds and Electronegativity

      Polar covalent bonds represent a fascinating middle ground between ionic and pure covalent bonds in the world of chemical bonding. These bonds are characterized by an unequal sharing of electrons between atoms, resulting in a partial charge distribution that gives molecules unique properties. Understanding polar covalent bonds is crucial for grasping the behavior of many common substances and their interactions.

      At the heart of polar covalent bonding lies the concept of electronegativity, which measures 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 to itself. This unequal distribution of electron density creates partial positive and negative charges on the atoms, known as partial charges or delta charges (δ+ and δ-).

      The degree of polarity in a covalent bond depends on the electronegativity difference between the bonded atoms. A larger difference results in a more polar bond, while a smaller difference leads to a less polar or nearly pure covalent bond. This spectrum of polarity allows for a wide range of molecular properties and behaviors.

      Water (H2O) serves as an excellent example of polar covalent bonding. In a water molecule, the oxygen atom is more electronegative than the hydrogen atoms. As a result, the shared electrons spend more time around the oxygen, giving it a partial negative charge while the hydrogen atoms acquire partial positive charges. This uneven distribution of charge creates a dipole moment, making water a polar molecule.

      Another classic example of polar covalent bonding is hydrogen fluoride (HF). The high electronegativity of fluorine compared to hydrogen leads to a significant charge separation, resulting in a strong dipole moment. The polarity of HF contributes to its unique properties, such as its ability to form strong hydrogen bonds.

      The polarity of molecules resulting from polar covalent bonds has profound effects on their properties and behavior. One of the most significant impacts is on solubility. Polar molecules tend to dissolve well in other polar substances, following the "like dissolves like" principle. This explains why water, a polar molecule, is an excellent solvent for many polar substances but not for nonpolar ones like oils.

      Intermolecular forces are also greatly influenced by molecular polarity. Polar molecules can engage in dipole-dipole interactions, where the positive end of one molecule is attracted to the negative end of another. These interactions contribute to the cohesive properties of liquids and affect physical properties such as boiling and melting points.

      In more complex molecules, the overall polarity depends not only on the individual bond polarities but also on the molecular geometry and polarity. Sometimes, polar bonds can be arranged in a way that their dipoles cancel out, resulting in a nonpolar molecule. Carbon dioxide (CO2), for instance, has polar C=O bonds, but its linear geometry results in a nonpolar molecule overall.

      The concept of polar covalent bonds extends beyond simple molecules and plays a crucial role in understanding the behavior of biological macromolecules. Proteins, for example, derive much of their structure and function from the polar and nonpolar interactions between their constituent amino acids. The folding of proteins and their interactions with other molecules are largely governed by these polar interactions.

      In the field of materials science, the polarity of covalent bonds influences the properties of polymers and other advanced materials. Engineers and scientists can manipulate molecular geometry and polarity to create materials with specific characteristics, such as water-repellent coatings or biodegradable plastics.

      Understanding polar covalent bonds and their effects on molecular properties is essential in many areas of chemistry and related fields. From predicting chemical reactions to designing new materials, the concept of bond polarity provides valuable insights into the behavior of matter at the molecular level. As research continues, our understanding of polar covalent bonds and their implications continues to grow, opening new possibilities in fields ranging from medicine to environmental science.

      Predicting Ionic and Covalent Bonds in Compounds

      Understanding how to predict whether a compound will form ionic or covalent bonds is a crucial skill in chemistry. This ability relies heavily on knowledge of the periodic table and the concept of electronegativity. By examining the elements involved in a compound, we can make informed predictions about the nature of the bonds that will form between them.

      The periodic table serves as an invaluable tool in bond prediction. Elements are arranged in a way that reflects their electronic structure and chemical properties. Generally, metals are found on the left side of the table, while non-metals occupy the right side. This arrangement provides our first clue in bond prediction: compounds formed between metals and non-metals are likely to be ionic, while those between non-metals are typically covalent.

      Electronegativity, a measure of an atom's ability to attract electrons in a chemical bond, plays a pivotal role in bond prediction. The electronegativity scale, developed by Linus Pauling, assigns values to elements based on their electron-attracting power. Elements with high electronegativity values, such as fluorine and oxygen, are more likely to form covalent bonds, while those with low values, like sodium and potassium, tend to form ionic bonds.

      To determine bond type based on electronegativity differences, we can follow these guidelines:

      • If the electronegativity difference is less than 0.5, the bond is likely to be non-polar covalent.
      • If the difference is between 0.5 and 1.7, the bond is probably polar covalent.
      • If the difference is greater than 1.7, the bond is likely to be ionic.

      However, it's important to note that these are general guidelines, and there can be exceptions. The bond character exists on a spectrum, with some compounds exhibiting partial ionic and partial covalent character.

      Let's examine some examples to illustrate how we can predict bond types:

      1. Sodium chloride (NaCl): Sodium is a metal, and chlorine is a non-metal. Their electronegativity difference is 2.1. This large difference, combined with the metal-nonmetal pairing, strongly indicates an ionic bond.
      2. Carbon dioxide (CO2): Both carbon and oxygen are non-metals. Their electronegativity difference is 0.89, suggesting polar covalent bonds.
      3. Hydrogen gas (H2): Hydrogen bonds with itself, and the electronegativity difference is zero, indicating a non-polar covalent bond.
      4. Potassium fluoride (KF): Potassium is a metal, and fluorine is a non-metal. Their electronegativity difference is 3.2, clearly indicating an ionic bond.
      5. Water (H2O): Hydrogen and oxygen are both non-metals. The electronegativity difference is 1.24, suggesting polar covalent bonds.

      In practice, predicting bond types becomes more intuitive with experience. Familiarity with the periodic table and common electronegativity values for elements will greatly enhance your ability to make quick and accurate predictions. Remember that while these guidelines are generally reliable, there can be exceptions, especially for compounds involving transition metals or in cases where multiple bond types are present within the same molecule.

      Understanding bond prediction is crucial for comprehending chemical behavior, reactivity, and properties of compounds. It forms the foundation for more advanced concepts in chemistry, such as molecular geometry, polarity, and intermolecular forces. By mastering the art of bond prediction, chemists can better anticipate how substances will interact, dissolve, and react under various conditions.

      In conclusion, predicting bond types in compounds involves a careful consideration of the elements' positions on the periodic table and their electronegativity values. By applying the guidelines based on electronegativity differences and recognizing patterns in metal-nonmetal combinations, we can make informed predictions about the nature of chemical bonds. This skill is essential for anyone studying chemistry or working in fields that involve material science, pharmaceuticals, or chemical engineering.

      Lewis Structures and Molecular Geometry

      Lewis structures, also known as Lewis dot structures, are fundamental tools in chemistry for representing covalent bonding and electron arrangement in molecules. Named after Gilbert N. Lewis, these diagrams provide a visual representation of how atoms share electrons to form stable compounds. Understanding Lewis structures is crucial for predicting molecular geometry and chemical properties.

      To draw a Lewis structure, we start by identifying the valence electrons of each atom in the molecule. Valence electrons are the outermost electrons that participate in bonding. For main group elements, the number of valence electrons corresponds to the group number in the periodic table. We then arrange these electrons around the atoms, following the octet rule, which states that atoms tend to gain, lose, or share electrons to achieve a stable configuration of eight valence electrons (like noble gases).

      When drawing Lewis structures, we follow these steps:

      1. Count the total number of valence electrons in the molecule or ion
      2. Identify the central atom (usually the least electronegative element)
      3. Connect atoms with single bonds
      4. Distribute remaining electrons as lone pairs
      5. Check for octet rule compliance and adjust if necessary

      For example, in a water molecule (H2O), oxygen has 6 valence electrons, and each hydrogen has 1. The total is 8 electrons. Oxygen is the central atom, connected to two hydrogens by single bonds. The remaining 4 electrons form two lone pairs on oxygen, completing its octet.

      Lewis structures are particularly useful for polyatomic ions, such as carbonate (CO3^2-) or ammonium (NH4^+). In these cases, we consider the overall charge when counting electrons and ensure the structure reflects the ion's properties.

      The concept of molecular geometry is closely related to Lewis structures and is explained by the Valence Shell Electron Pair Repulsion (VSEPR) theory. This theory predicts the three-dimensional shape of molecules based on the arrangement of electron pairs around the central atom. The key principle is that electron pairs repel each other, leading to specific geometric arrangements that minimize repulsion.

      Common molecular shapes include:

      • Linear: Two electron domains around the central atom, such as in carbon dioxide (CO2)
      • Trigonal planar: Three electron domains, as seen in boron trifluoride (BF3)
      • Tetrahedral: Four electron domains, exemplified by methane (CH4)
      • Bent: Three electron domains with lone pairs, like in water (H2O)
      • Trigonal pyramidal: Four electron domains with one lone pair, as in ammonia (NH3)

      The VSEPR theory considers both bonding and non-bonding (lone) electron pairs when predicting geometry. Lone pairs occupy more space than bonding pairs, which can lead to slight deviations from ideal geometries. For instance, the bent shape of water results from the repulsion between two bonding pairs and two lone pairs on the oxygen atom.

      Understanding Lewis structures and molecular geometry is essential for predicting a molecule's polarity, reactivity, and other chemical properties. For example, the tetrahedral shape of methane explains its non-polar nature, while the bent shape of water contributes to its polar character and unique properties like high boiling point and ability to dissolve many substances.

      In more complex molecules, Lewis structures can help identify resonance structures, where electrons are delocalized across multiple atoms. This concept is crucial in organic chemistry for understanding the stability and reactivity of compounds like benzene and carboxylate ions.

      As students progress in their chemistry education, they will encounter exceptions to the octet rule, such as expanded octets in compounds like sulfur hexafluoride (SF6) or electron-deficient molecules like boron trifluoride (BF3). These exceptions highlight the limitations of simple Lewis structures and introduce more advanced bonding theories.

      In conclusion, Lewis structures and molecular geometry are foundational concepts in chemistry that provide

      Conclusion: The Importance of Understanding Chemical Bonding

      Chemical bonding, encompassing ionic and covalent bonds, is fundamental to understanding compound properties and molecular structure. The introduction video provides a crucial visual representation of these concepts, making them more accessible. Ionic bonding involves the transfer of electrons between atoms, while covalent bonding involves sharing electrons. These different types of bonds significantly influence the properties of substances, such as melting points, solubility, and reactivity. Understanding chemical bonding is essential for predicting and explaining the behavior of materials in various contexts, from everyday substances to advanced materials in technology. It forms the basis for comprehending chemical reactions and molecular interactions. As you delve deeper into chemistry and materials science, the knowledge of chemical bonding will prove invaluable. We encourage you to explore further applications of this topic, as it plays a vital role in numerous scientific and technological advancements, from drug design to the development of new materials.

      Building on Bonding: How are Atoms Joined Together?

      Step 1: Introduction to Bonding

      In this lesson, we will explore the two main types of bonding that join atoms together to form compounds and molecules. These bonds are essential for creating the larger, more complex structures that make up the substances we encounter in everyday life. Our primary focus will be on understanding the nature of these bonds, their properties, and the conditions under which they form.

      Step 2: Objectives of the Lesson

      The objectives of this lesson are to understand the two main types of bonding in compounds: ionic and covalent. We will delve into why both types of bonds exist, their properties, and the factors that influence their formation. Additionally, we will learn how to predict the formulas of ionic and covalent compounds based on our understanding of these bonds.

      Step 3: Predicting Bond Types

      By the end of this lesson, you should be able to look at two atoms and, based on your knowledge of ionic and covalent bonding, predict which type of bond is likely to form between them. This skill is crucial for understanding the behavior of different substances and their chemical properties.

      Step 4: Understanding Bonding

      Bonding, in a general sense, refers to the forces that bring atoms together. Just as people can bond over shared interests, atoms bond through attractive forces. These forces can be categorized into two main types: intramolecular forces and intermolecular forces.

      Step 5: Intramolecular Forces

      Intramolecular forces are the attractive forces that hold atoms together within a molecule. These forces are responsible for the chemical identity of a substance. For example, a molecule of carbon dioxide (CO2) consists of one carbon atom and two oxygen atoms held together by intramolecular forces. If these bonds are broken, the molecule no longer exists as CO2.

      Step 6: Importance of Intramolecular Forces

      Intramolecular forces are crucial because they define the substance. Breaking these bonds results in a chemical reaction that changes the substance's identity. For instance, breaking the bonds in CO2 would result in separate carbon and oxygen atoms, no longer forming the compound CO2.

      Step 7: Intermolecular Forces

      Intermolecular forces, on the other hand, are the forces between molecules. These forces are responsible for the physical states of substances (solid, liquid, gas). Overcoming intermolecular forces changes the state of the substance but does not alter its chemical identity. For example, CO2 can exist as a gas, liquid, or solid, but it remains CO2 in all states.

      Step 8: Focus on Chemical Bonds

      For the purpose of this lesson, we will focus on intramolecular forces, which are also known as chemical bonds. These include ionic and covalent bonds. Understanding these bonds will help us predict the behavior and properties of different compounds.

      Step 9: Introduction to Ionic Bonding

      We will begin our detailed exploration with ionic bonding. Ionic bonds form when atoms transfer electrons from one to another, resulting in the formation of positively and negatively charged ions. These oppositely charged ions attract each other, creating a strong bond.

      Step 10: Introduction to Covalent Bonding

      Next, we will examine covalent bonding. Covalent bonds form when atoms share electrons, creating a stable balance of attractive and repulsive forces between them. This type of bonding is common in organic compounds and many other substances.

      Step 11: Summary and Application

      By understanding the principles of ionic and covalent bonding, we can predict the properties and behaviors of various compounds. This knowledge is fundamental to the study of chemistry and helps us understand the molecular world around us.

      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.

      Prerequisite Topics for Understanding Ionic and Covalent Bonding

      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.

      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.