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Anatomy & Physiology 2e: 2.2 Chemical Bonds

Anatomy & Physiology 2e
2.2 Chemical Bonds
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table of contents
  1. Cover
  2. Title Page
  3. Copyright
  4. Table Of Contents
  5. Chapter 1. An Introduction to the Human Body
    1. 1.0 Introduction
    2. 1.1 How Structure Determines Function
    3. 1.2 Structural Organization of the Human Body
    4. 1.3 Homeostasis
    5. 1.4 Anatomical Terminology
    6. 1.5 Medical Imaging
  6. Chapter 2. The Chemical Level of Organization
    1. 2.0 Introduction
    2. 2.1 Elements and Atoms: The Building Blocks of Matter
    3. 2.2 Chemical Bonds
    4. 2.3 Chemical Reactions
    5. 2.4 Inorganic Compounds Essential to Human Functioning
    6. 2.5 Organic Compounds Essential to Human Functioning
  7. Chapter 3. The Cellular Level of Organization
    1. 3.0 Introduction
    2. 3.1 The Cell Membrane
    3. 3.2 The Cytoplasm and Cellular Organelles
    4. 3.3 The Nucleus and DNA Replication
    5. 3.4 Protein Synthesis
    6. 3.5 Cell Growth and Division
    7. 3.6 Cellular Differentiation
  8. Chapter 4. The Tissue Level of Organization
    1. 4.0 Introduction
    2. 4.1 Types of Tissues
    3. 4.2 Epithelial Tissue
    4. 4.3 Connective Tissue Supports and Protects
    5. 4.4 Muscle Tissue
    6. 4.5 Nervous Tissue
    7. 4.6 Tissue Injury and Aging
  9. Chapter 5. The Integumentary System
    1. 5.0 Introduction
    2. 5.1 Layers of the Skin
    3. 5.2 Accessory Structures of the Skin
    4. 5.3 Functions of the Integumentary System
    5. 5.4 Diseases, Disorders, and Injuries of the Integumentary System
  10. Chapter 6. Bone Tissue and the Skeletal System
    1. 6.0 Introduction
    2. 6.1 The Functions of the Skeletal System
    3. 6.2 Bone Classification
    4. 6.3 Bone Structure
    5. 6.4 Bone Formation and Development
    6. 6.5 Fractures: Bone Repair
    7. 6.6 Exercise, Nutrition, Hormones, and Bone Tissue
    8. 6.7 Calcium Homeostasis: Interactions of the Skeletal System and Other Organ Systems
  11. Chapter 7. Axial Skeleton
    1. 7.0 Introduction
    2. 7.1 Divisions of the Skeletal System
    3. 7.2 Bone Markings
    4. 7.3 The Skull
    5. 7.4 The Vertebral Column
    6. 7.5 The Thoracic Cage
    7. 7.6 Embryonic Development of the Axial Skeleton
  12. Chapter 8. The Appendicular Skeleton
    1. 8.0 Introduction
    2. 8.1 The Pectoral Girdle
    3. 8.2 Bones of the Upper Limb
    4. 8.3 The Pelvic Girdle and Pelvis
    5. 8.4 Bones of the Lower Limb
    6. 8.5 Development of the Appendicular Skeleton
  13. Chapter 9. Joints
    1. 9.0 Introduction
    2. 9.1 Classification of Joints
    3. 9.2 Fibrous Joints
    4. 9.3 Cartilaginous Joints
    5. 9.4 Synovial Joints
    6. 9.5 Types of Body Movements
    7. 9.6 Anatomy of Selected Synovial Joints
    8. 9.7 Development of Joints
  14. Chapter 10. Muscle Tissue
    1. 10.0 Introduction
    2. 10.1 Overview of Muscle Tissues
    3. 10.2 Skeletal Muscle
    4. 10.3 Muscle Fiber Excitation, Contraction, and Relaxation
    5. 10.4 Nervous System Control of Muscle Tension
    6. 10.5 Types of Muscle Fibers
    7. 10.6 Exercise and Muscle Performance
    8. 10.7 Smooth Muscle Tissue
    9. 10.8 Development and Regeneration of Muscle Tissue
  15. Chapter 11. The Muscular System
    1. 11.0 Introduction
    2. 11.1 Describe the roles of agonists, antagonists and synergists
    3. 11.2 Explain the organization of muscle fascicles and their role in generating force
    4. 11.3 Explain the criteria used to name skeletal muscles
    5. 11.4 Axial Muscles of the Head Neck and Back
    6. 11.5 Axial muscles of the abdominal wall and thorax
    7. 11.6 Muscles of the Pectoral Girdle and Upper Limbs
    8. 11.7 Appendicular Muscles of the Pelvic Girdle and Lower Limbs
  16. Chapter 12. The Nervous System and Nervous Tissue
    1. 12.0 Introduction
    2. 12.1 Structure and Function of the Nervous System
    3. 12.2 Nervous Tissue
    4. 12.3 The Function of Nervous Tissue
    5. 12.4 Communication Between Neurons
    6. 12.5 The Action Potential
  17. Chapter 13. The Peripheral Nervous System
    1. 13.0 Introduction
    2. 13.1 Sensory Receptors
    3. 13.2 Ganglia and Nerves
    4. 13.3 Spinal and Cranial Nerves
    5. 13.4 Relationship of the PNS to the Spinal Cord of the CNS
    6. 13.5 Ventral Horn Output and Reflexes
    7. 13.6 Testing the Spinal Nerves (Sensory and Motor Exams)
    8. 13.7 The Cranial Nerve Exam
  18. Chapter 14. The Central Nervous System
    1. 14.0 Introduction
    2. 14.1 Embryonic Development
    3. 14.2 Blood Flow the meninges and Cerebrospinal Fluid Production and Circulation
    4. 14.3 The Brain and Spinal Cord
    5. 14.4 The Spinal Cord
    6. 14.5 Sensory and Motor Pathways
  19. Chapter 15. The Special Senses
    1. 15.0 Introduction
    2. 15.1 Taste
    3. 15.2 Smell
    4. 15.3 Hearing
    5. 15.4 Equilibrium
    6. 15.5 Vision
  20. Chapter 16. The Autonomic Nervous System
    1. 16.0 Introduction
    2. 16.1 Divisions of the Autonomic Nervous System
    3. 16.2 Autonomic Reflexes and Homeostasis
    4. 16.3 Central Control
    5. 16.4 Drugs that Affect the Autonomic System
  21. Chapter 17. The Endocrine System
    1. 17.0 Introduction
    2. 17.1 An Overview of the Endocrine System
    3. 17.2 Hormones
    4. 17.3 The Pituitary Gland and Hypothalamus
    5. 17.4 The Thyroid Gland
    6. 17.5 The Parathyroid Glands
    7. 17.6 The Adrenal Glands
    8. 17.7 The Pineal Gland
    9. 17.8 Gonadal and Placental Hormones
    10. 17.9 The Pancreas
    11. 17.10 Organs with Secondary Endocrine Functions
    12. 17.11 Development and Aging of the Endocrine System
  22. Chapter 18. The Cardiovascular System: Blood
    1. 18.0 Introduction
    2. 18.1 Functions of Blood
    3. 18.2 Production of the Formed Elements
    4. 18.3 Erythrocytes
    5. 18.4 Leukocytes and Platelets
    6. 18.5 Hemostasis
    7. 18.6 Blood Typing
  23. Chapter 19. The Cardiovascular System: The Heart
    1. 19.0 Introduction
    2. 19.1 Heart Anatomy
    3. 19.2 Cardiac Muscle and Electrical Activity
    4. 19.3 Cardiac Cycle
    5. 19.4 Cardiac Physiology
    6. 19.5 Development of the Heart
  24. Chapter 20. The Cardiovascular System: Blood Vessels and Circulation
    1. 20.0 Introduction
    2. 20.1 Structure and Function of Blood Vessels
    3. 20.2 Blood Flow, Blood Pressure, and Resistance
    4. 20.3 Capillary Exchange
    5. 20.4 Homeostatic Regulation of the Vascular System
    6. 20.5 Circulatory Pathways
    7. 20.6 Development of Blood Vessels and Fetal Circulation
  25. Chapter 21. The Lymphatic and Immune System
    1. 21.0 Introduction
    2. 21.1 Anatomy of the Lymphatic and Immune Systems
    3. 21.2 Barrier Defenses and the Innate Immune Response
    4. 21.3 The Adaptive Immune Response: T lymphocytes and Their Functional Types
    5. 21.4 The Adaptive Immune Response: B-lymphocytes and Antibodies
    6. 21.5 The Immune Response against Pathogens
    7. 21.6 Diseases Associated with Depressed or Overactive Immune Responses
    8. 21.7 Transplantation and Cancer Immunology
  26. Chapter 22. The Respiratory System
    1. 22.0 Introduction
    2. 22.1 Organs and Structures of the Respiratory System
    3. 22.2 The Lungs
    4. 22.3 The Process of Breathing
    5. 22.4 Gas Exchange
    6. 22.5 Transport of Gases
    7. 22.6 Modifications in Respiratory Functions
    8. 22.7 Embryonic Development of the Respiratory System
  27. Chapter 23. The Digestive System
    1. 23.0 Introduction
    2. 23.1 Overview of the Digestive System
    3. 23.2 Digestive System Processes and Regulation
    4. 23.3 The Mouth, Pharynx, and Esophagus
    5. 23.4 The Stomach
    6. 23.5 Accessory Organs in Digestion: The Liver, Pancreas, and Gallbladder
    7. 23.6 The Small and Large Intestines
    8. 23.7 Chemical Digestion and Absorption: A Closer Look
  28. Chapter 24. Metabolism and Nutrition
    1. 24.0 Introduction
    2. 24.1 Overview of Metabolic Reactions
    3. 24.2 Carbohydrate Metabolism
    4. 24.3 Lipid Metabolism
    5. 24.4 Protein Metabolism
    6. 24.5 Metabolic States of the Body
    7. 24.6 Energy and Heat Balance
    8. 24.7 Nutrition and Diet
  29. Chapter 25. The Urinary System
    1. 25.0 Introduction
    2. 25.1 Internal and External Anatomy of the Kidney
    3. 25.2 Microscopic Anatomy of the Kidney: Anatomy of the Nephron
    4. 25.3 Physiology of Urine Formation: Overview
    5. 25.4 Physiology of Urine Formation: Glomerular Filtration
    6. 25.5 Physiology of Urine Formation: Tubular Reabsorption and Secretion
    7. 25.6 Physiology of Urine Formation: Medullary Concentration Gradient
    8. 25.7 Physiology of Urine Formation: Regulation of Fluid Volume and Composition
    9. 25.8 Urine Transport and Elimination
    10. 25.9 The Urinary System and Homeostasis
  30. Chapter 26. Fluid, Electrolyte, and Acid-Base Balance
    1. 26.0 Introduction
    2. 26.1 Body Fluids and Fluid Compartments
    3. 26.2 Water Balance
    4. 26.3 Electrolyte Balance
    5. 26.4 Acid-Base Balance
    6. 26.5 Disorders of Acid-Base Balance
  31. Chapter 27. The Sexual Systems
    1. 27.0 Introduction
    2. 27.1 Anatomy of Sexual Systems
    3. 27.2 Development of Sexual Anatomy
    4. 27.3 Physiology of the Female Sexual System
    5. 27.4 Physiology of the Male Sexual System
    6. 27.5 Physiology of Arousal and Orgasm
  32. Chapter 28. Development and Inheritance
    1. 28.0 Introduction
    2. 28.1 Fertilization
    3. 28.2 Embryonic Development
    4. 28.3 Fetal Development
    5. 28.4 Maternal Changes During Pregnancy, Labor, and Birth
    6. 28.5 Adjustments of the Infant at Birth and Postnatal Stages
    7. 28.6 Lactation
    8. 28.7 Patterns of Inheritance
  33. Creative Commons License
  34. Recommended Citations
  35. Versioning

2.2 Chemical Bonds

Learning Objectives

By the end of this section, you will be able to:

  • Explain the relationship between molecules and compounds
  • Distinguish between ions, cations, and anions
  • Identify the key difference between ionic and covalent bonds
  • Distinguish between nonpolar and polar covalent bonds
  • Explain how water molecules link via hydrogen bonds

Atoms separated by a great distance cannot link; rather, they must come close enough for the electrons in their valence shells to interact. But do atoms ever actually touch one another? Most physicists would say no, because the negatively charged electrons in their valence shells repel one another. No force within the human body—or anywhere in the natural world—is strong enough to overcome this electrical repulsion. So when you read about atoms linking together or colliding, bear in mind that the atoms are not merging in a physical sense.

Instead, atoms link by forming a chemical bond. A bond is a weak or strong electrical attraction that holds atoms in the same vicinity. The new grouping is typically more stable—less likely to react again—than its component atoms were when they were separate. A more or less stable grouping of two or more atoms held together by chemical bonds is called a molecule. The bonded atoms may be of the same element, as in the case of H2, which is called molecular hydrogen or hydrogen gas. When a molecule is made up of two or more atoms of different elements, it is called a chemical compound. A unit of water, or H2O, is a compound, as is a single molecule of the gas methane, or CH4.

Three types of chemical bonds are important in human physiology, because they hold together substances that are used by the body for critical aspects of homeostasis, signaling, and energy production, to name just a few important processes. These are ionic bonds, covalent bonds, and hydrogen bonds.

Ions and Ionic Bonds

Recall that an atom typically has the same number of positively charged protons and negatively charged electrons. As long as this situation remains, the atom is electrically neutral. When an atom participates in a chemical reaction that results in the donation or acceptance of one or more electrons, the atom will then become positively or negatively charged. This happens frequently for most atoms in order to have a full valence shell, as described previously. This can happen either by gaining electrons to fill a shell that is more than half-full, or by giving away electrons to empty a shell that is less than half-full, thereby leaving the next smaller electron shell as the new, full, valence shell. An atom that has an electrical charge—whether positive or negative—is an ion.

External Website

electenergy

Visit this website to learn about electrical energy and the attraction/repulsion of charges. What happens to the charged electroscope when a conductor is moved between its plastic sheets, and why?

Potassium (K), for instance, is an important element in all body cells. Its atomic number is 19 and it has just one electron in its valence shell. This characteristic makes potassium highly likely to participate in chemical reactions in which it donates one electron (it is easier for potassium to donate one electron than to gain seven electrons). The loss will cause the positive charge of potassium’s protons to be more influential than the negative charge of potassium’s electrons. In other words, the resulting potassium ion will be slightly positive. A potassium ion is written K+, indicating that it has lost a single electron. A positively charged ion is known as a cation.

Now consider fluorine (F), a component of bones and teeth. Its atomic number is nine and it has seven electrons in its valence shell. Thus, it is highly likely to bond with other atoms in such a way that fluorine accepts one electron (it is easier for fluorine to gain one electron than to donate seven electrons). When it does, its electrons will outnumber its protons by one and it will have an overall negative charge. The ionized form of fluorine is called fluoride, and is written as F–. A negatively charged ion is known as an anion.

Atoms that have more than one electron to donate or accept will end up with stronger positive or negative charges. A cation that has donated two electrons has a net charge of +2. Using magnesium (Mg) as an example, this can be written as Mg++ or Mg2+. An anion that has accepted two electrons has a net charge of –2. The ionic form of selenium (Se), for example, is typically written Se2–.

The opposite charges of cations and anions exert a moderately strong mutual attraction that keeps the atoms in close proximity forming an ionic bond. An ionic bond is an ongoing, close association between ions of opposite charge. The table salt you sprinkle on your food owes its existence to ionic bonding. As shown in Figure 2.2.1, sodium commonly donates an electron to chlorine, becoming the cation Na+. When chlorine accepts the electron, it becomes the chloride anion, Cl–. With their opposing charges, these two ions strongly attract each other.

The top panel of this figure shows the orbit model of a sodium atom and a chlorine atom and arrows pointing towards the transfer of electrons from sodium to chlorine to form sodium and chlorine ions. The bottom panel shows sodium and chloride ions in a crystal structure.
Figure 2.2.1 – Ionic Bonding: (a) Sodium readily donates the solitary electron in its valence shell to chlorine, which needs only one electron to have a full valence shell. (b) The opposite electrical charges of the resulting sodium cation and chloride anion result in the formation of a bond of attraction called an ionic bond. (c) The attraction of many sodium and chloride ions results in the formation of large groupings called crystals.

Water is an essential component of life because it is able to break the ionic bonds in salts to free the ions. In fact, in biological fluids, most individual atoms exist as ions. These dissolved ions produce electrical charges within the body. The behavior of these ions produces the tracings of heart and brain function observed as waves on an electrocardiogram (EKG or ECG) or an electroencephalogram (EEG). The electrical activity that derives from the interactions of the charged ions is why they are also called electrolytes.

Covalent Bonds

Unlike ionic bonds formed by the attraction between a cation’s positive charge and an anion’s negative charge, molecules formed by a covalent bond which share electrons in a mutually stabilizing relationship. Like next-door neighbors whose kids hang out first at one home and then at the other, the atoms do not lose or gain electrons permanently. Instead, the electrons move back and forth between the elements. Because of the close sharing of pairs of electrons (one electron from each of two atoms), covalent bonds are stronger than ionic bonds.

Nonpolar Covalent Bonds

Figure 2.2.2 shows several common types of covalent bonds. Notice that the two covalently bonded atoms typically share just one or two electron pairs, though larger sharings are possible. The important concept to take from this is that in covalent bonds, electrons in the outermost valence shell are shared to fill the valence shells of both atoms, ultimately stabilizing both of the atoms involved. In a single covalent bond, a single electron is shared between two atoms, while in a double covalent bond, two pairs of electrons are shared between two atoms. There are even triple covalent bonds, where three atoms are shared.

The top panel in this figure shows two hydrogen atoms sharing two electrons. The middle panel shows two oxygen atoms sharing four electrons, and the bottom panel shows two oxygen atoms and one carbon atom sharing 2 pairs of electrons each.
Figure 2.2.2 Covalent Bonding

You can see that the covalent bonds shown in Figure 2.2.2 are balanced. The sharing of the negative electrons is relatively equal, as is the electrical pull of the positive protons in the nucleus of the atoms involved. This is why covalently bonded molecules that are electrically balanced in this way are described as nonpolar; that is, no region of the molecule is either more positive or more negative than any other.

Polar Covalent Bonds

Groups of legislators with completely opposite views on a particular issue are often described as “polarized” by news writers. In chemistry, a polar molecule is a molecule that contains regions that have opposite electrical charges. Polar molecules occur when atoms share electrons unequally, in polar covalent bonds.

The most familiar example of a polar molecule is water (Figure 2.2.3). The molecule has three parts: one atom of oxygen, the nucleus of which contains eight protons, and two hydrogen atoms, whose nuclei each contain only one proton. Since every proton exerts an identical positive charge, a nucleus that contains eight protons exerts a charge eight times greater than a nucleus that contains one proton. This means that the negatively charged electrons present in the water molecule are more strongly attracted to the oxygen nucleus than to the hydrogen nuclei. Each hydrogen atom’s single negative electron, therefore, migrates toward the oxygen atom, making the oxygen end of their bond slightly more negative than the hydrogen end of their bond.

This figure shows the structure of a water molecule. The top panel shows two oxygen atoms and one hydrogen atom with electrons in orbit and the shared electrons. The middle panel shows a three-dimensional model of a water molecule and the bottom panel shows the structural formula for water.
Figure 2.2.3 Polar Covalent Bonds in a Water Molecule

What is true for the bonds is true for the water molecule as a whole; that is, the oxygen region has a slightly negative charge and the regions of the hydrogen atoms have a slightly positive charge. These charges are often referred to as “partial charges” because the strength of the charge is less than one full electron, as would occur in an ionic bond. As shown in Figure 2.2.3, regions of weak polarity are indicated with the Greek letter delta (∂) and a plus (+) or minus (–) sign.

Even though a single water molecule is unimaginably tiny, it has mass, and the opposing electrical charges on the molecule pull that mass in such a way that it creates a shape somewhat like a triangular tent (see Figure 2.2.3b). This dipole, with the positive charges at one end formed by the hydrogen atoms at the “bottom” of the tent and the negative charge at the opposite end (the oxygen atom at the “top” of the tent) makes the charged regions highly likely to interact with charged regions of other polar molecules. For human physiology, the resulting bond, formed by water, is one of the most important—the hydrogen bond.

Hydrogen Bonds

A hydrogen bond is formed when a weakly positive hydrogen atom already bonded to one electronegative atom (for example, the oxygen in the water molecule) is attracted to another electronegative atom from another molecule. In other words, hydrogen bonds always include hydrogen that is already part of a polar molecule.

The most common example of hydrogen bonding in the natural world occurs between molecules of water. It happens before your eyes whenever two raindrops merge into a larger bead, or a creek spills into a river. Hydrogen bonding occurs because the weakly negative oxygen atom in one water molecule is attracted to the weakly positive hydrogen atoms of two other water molecules (Figure 2.2.4).

This figure shows three water molecules and the hydrogen bonds between them.
Figure 2.2.4 – Hydrogen Bonds between Water Molecules: Notice that the bonds occur between the weakly positive charge on the hydrogen atoms and the weakly negative charge on the oxygen atoms. Hydrogen bonds are relatively weak, and therefore are indicated with a dotted (rather than a solid) line.

Notice that the bonds occur between the weakly positive charge on the hydrogen atoms and the weakly negative charge on the oxygen atoms. Hydrogen bonds are relatively weak, and therefore are indicated with a dotted (rather than a solid) line.

Water molecules also strongly attract other types of charged molecules as well as ions. This explains why “table salt,” for example, actually is a molecule called a “salt” in chemistry; it consists of equal numbers of positively-charged sodium (Na+) and negatively-charged chloride (Cl–), dissolves so readily in water, in this case, forming dipole-ion bonds between the water and the electrically-charged ions (electrolytes). Water molecules also repel molecules with nonpolar covalent bonds, like fats, lipids, and oils. You can demonstrate this with a simple kitchen experiment: pour a teaspoon of vegetable oil, a compound formed by nonpolar covalent bonds, into a glass of water. Instead of instantly dissolving in the water, the oil forms a distinct bead because the polar water molecules repel the nonpolar oil.

Chapter Review

Each moment of life, atoms of oxygen, carbon, hydrogen, and the other elements of the human body are making and breaking chemical bonds. Ions are charged atoms that form when an atom donates or accepts one or more negatively charged electrons. Cations (ions with a positive charge) are attracted to anions (ions with a negative charge). This attraction is called an ionic bond. In covalent bonds, the participating atoms do not lose or gain electrons, but share them. Molecules with nonpolar covalent bonds are electrically balanced, and have a linear three-dimensional shape. Molecules with polar covalent bonds have “poles”—regions of weakly positive and negative charge—and have a triangular three-dimensional shape. An atom of oxygen and two atoms of hydrogen form water molecules by means of polar covalent bonds. Hydrogen bonds link hydrogen atoms already participating in polar covalent bonds to anions or electronegative regions of other polar molecules. Hydrogen bonds link water molecules, resulting in the properties of water that are important to living things.

Interactive Link Questions

Visit this website to learn about electrical energy and the attraction/repulsion of charges. What happens to the charged electroscope when a conductor is moved between its plastic sheets, and why?

The plastic sheets jump to the nail (the conductor), because the conductor takes on electrons from the electroscope, reducing the repellant force of the two sheets.

Review Questions

An interactive H5P element has been excluded from this version of the text. You can view it online here:
https://open.oregonstate.education/aandp/?p=64#h5p-20

An interactive H5P element has been excluded from this version of the text. You can view it online here:
https://open.oregonstate.education/aandp/?p=64#h5p-21

An interactive H5P element has been excluded from this version of the text. You can view it online here:
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An interactive H5P element has been excluded from this version of the text. You can view it online here:
https://open.oregonstate.education/aandp/?p=64#h5p-23

An interactive H5P element has been excluded from this version of the text. You can view it online here:
https://open.oregonstate.education/aandp/?p=64#h5p-24

Critical Thinking Questions

Explain why CH4 is one of the most common molecules found in nature. Are the bonds between the atoms ionic or covalent?

A carbon atom has four electrons in its valence shell. According to the octet rule, it will readily participate in chemical reactions that result in its valence shell having eight electrons. Hydrogen, with one electron, will complete its valence shell with two. Electron sharing between an atom of carbon and four atoms of hydrogen meets the requirements of all atoms. The bonds are covalent because the electrons are shared. Although hydrogen often participates in ionic bonds, carbon does not because it is highly unlikely to donate or accept four electrons.

In a hurry one day, you merely rinse your lunch dishes with water. As you are drying your salad bowl, you notice that it still has an oily film. Why was the water alone not effective in cleaning the bowl?

Water is a polar molecule. It has a region of weakly positive charge and a region of weakly negative charge. These regions are attracted to ions as well as to other polar molecules. Oils are nonpolar and are repelled by water.

Could two atoms of oxygen engage in ionic bonding? Why or why not?

Identical atoms have identical electronegativity and cannot form ionic bonds. Oxygen, for example, has six electrons in its valence shell. Neither donating nor accepting the valence shell electrons of the other will result in the oxygen atoms completing their valence shells. Two atoms of the same element always form covalent bonds.

Annotate

Next chapter
2.3 Chemical Reactions
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Anatomy and Physiology
Copyright © 2019 by Lindsay M. Biga, Sierra Dawson, Amy Harwell, Robin Hopkins, Joel Kaufmann, Mike LeMaster, Philip Matern, Katie Morrison-Graham, Devon Quick & Jon Runyeon

Anatomy & Physiology by Lindsay M. Biga, Sierra Dawson, Amy Harwell, Robin Hopkins, Joel Kaufmann, Mike LeMaster, Philip Matern, Katie Morrison-Graham, Devon Quick & Jon Runyeon is licensed under a Creative Commons Attribution-ShareAlike 4.0 International License, except where otherwise noted.

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