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Periodic trends: Electronegativity

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Intros
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  1. The periodic trends of electronegativity
  2. Definition of electronegativity.
  3. Trends in electronegativity in the Periodic Table.
  4. Electronegativity and bonding.
  5. Trends in metallic character across a period.
  6. Trends in acid-base properties of oxide compounds.
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Examples
Lessons
  1. Recall and explain the trends in electronegativity in the periodic table.
    Describe and explain the trend in electronegativity across the elements in period 2 of the Periodic Table.
    1. Recall and explain the trends in electronegativity in the periodic table.
      Describe and explain the trend in electronegativity in group 2 of the Periodic Table.
      Topic Notes
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      Introduction to Electronegativity as a Periodic Trend

      Electronegativity is a fundamental periodic trend in chemistry that describes an atom's ability to attract electrons in a chemical bond. Our introduction video provides a comprehensive overview of this crucial concept, helping students grasp its significance in understanding chemical behavior. As we explore electronegativity, we'll discover its vital role in predicting bond polarity, reactivity, and other chemical properties. This trend is closely related to other periodic trends such as atomic radius relationship, ionization energy correlation, and electron affinity. Electronegativity generally increases from left to right across a period and decreases down a group in the periodic table trends. By mastering this concept, students gain valuable insights into chemical bonding, molecular structure, and reactivity. The video serves as an excellent starting point for delving deeper into the fascinating world of electronegativity and its impact on various aspects of chemistry.

      Understanding the atomic radius relationship is essential for predicting how atoms will interact in different chemical environments. Similarly, the ionization energy correlation provides insights into the energy required to remove an electron from an atom, which is crucial for understanding reactivity and bonding. The periodic table trends offer a framework for organizing and predicting the properties of elements, making it a cornerstone of modern chemistry education.

      Definition and Measurement of Electronegativity

      Electronegativity is a fundamental concept in chemistry that plays a crucial role in understanding chemical reactions and molecular properties. At its core, electronegativity is defined as the ability of an atomic nucleus to attract bonding electrons within a molecule. This property is essential in determining how atoms interact and form chemical bonds with one another.

      The concept of electronegativity was first introduced by Linus Pauling in the 1930s. Pauling developed a scale to quantify and compare the electronegativity of different elements, which is now known as the Pauling scale. This scale assigns numerical values to elements based on their ability to attract electrons, with values ranging from 0.7 to 4.0. The Pauling scale has become the most widely used method for measuring electronegativity in chemistry.

      On the Pauling scale, elements with higher electronegativity values have a stronger ability to attract electrons. Fluorine, for example, has the highest electronegativity value of 4.0, making it the most electronegative element. Other highly electronegative elements include oxygen (3.5), nitrogen (3.0), and chlorine (3.0). These elements tend to form strong covalent bonds and are often found in molecules with polar characteristics.

      On the other end of the spectrum, elements with low electronegativity values have a weaker ability to attract electrons. Cesium and francium, both with an electronegativity of 0.7, are the least electronegative elements. Other elements with low electronegativity include potassium (0.8), sodium (0.9), and lithium (1.0). These elements are typically metallic and tend to lose electrons easily, forming ionic bonds with more electronegative elements.

      Electronegativity plays a crucial role in determining the type and strength of chemical reactions formed between atoms. When two atoms with similar electronegativity values bond, they tend to form non-polar covalent bonds, sharing electrons equally. For example, the bond between two carbon atoms in a hydrocarbon molecule is non-polar due to their identical electronegativity values.

      However, when atoms with significantly different electronegativity values bond, they form polar covalent bonds. In these bonds, the more electronegative atom attracts the shared electrons more strongly, creating a partial negative charge on that atom and a partial positive charge on the less electronegative atom. This polarity is responsible for many important molecular properties, such as solubility and intermolecular forces.

      In extreme cases, when the difference in electronegativity between two bonded atoms is very large, the bond may become ionic. Ionic bonds occur when one atom completely transfers its electron to another atom, resulting in the formation of oppositely charged ions. For instance, the bond between sodium (low electronegativity) and chlorine (high electronegativity) in table salt is ionic.

      Understanding electronegativity is essential for predicting and explaining various chemical phenomena. It helps chemists determine the polarity of molecules, which in turn affects properties such as boiling point, melting point, and solubility. Electronegativity also plays a role in understanding reaction mechanisms, acid-base behavior, and the stability of chemical compounds.

      In conclusion, electronegativity is a fundamental concept in chemistry that describes an atom's ability to attract bonding electrons. The Pauling scale provides a quantitative measure of electronegativity, allowing chemists to compare and predict the behavior of different elements in chemical reactions and bond formation. By understanding electronegativity, we gain valuable insights into the nature of chemical bonds, molecular properties, and the behavior of substances in various chemical processes.

      Electronegativity Trend Across a Period

      Electronegativity is a fundamental concept in chemistry that plays a crucial role in understanding the behavior of elements and their compounds. One of the most important trends observed in the periodic table is the increase in electronegativity across a period. This trend has significant implications for chemical bonding, reactivity, and various properties of elements and their compounds.

      As we move from left to right across a period in the periodic table, the electronegativity of elements generally increases. This means that atoms become more capable of attracting electrons in a chemical bond. The primary reason for this trend is the increase in effective nuclear charge and the constant shielding effect experienced by the valence electrons.

      The effective nuclear charge is the net positive charge experienced by the valence electrons due to the attraction from the nucleus and the repulsion from inner-shell electrons. As we progress across a period, the number of protons in the nucleus increases, while the number of inner-shell electrons remains constant. This results in a stronger attraction between the nucleus and the valence electrons, leading to an increase in effective nuclear charge.

      The shielding effect, on the other hand, remains relatively constant across a period. Inner-shell electrons shield the valence electrons from the full nuclear charge, but this shielding doesn't change significantly as we move across a period. The combination of increasing effective nuclear charge and constant shielding effect causes the valence electrons to be held more tightly to the nucleus, resulting in higher electronegativity.

      To illustrate this trend, let's consider period 2 of the periodic table. As we move from lithium (Li) to fluorine (F), we observe a steady increase in electronegativity. Lithium, with its single valence electron, has a relatively low electronegativity of 0.98 on the Pauling scale. Beryllium (Be) has a slightly higher electronegativity of 1.57, followed by boron (B) at 2.04. The trend continues with carbon (C) at 2.55, nitrogen (N) at 3.04, oxygen (O) at 3.44, and finally, fluorine (F) with the highest electronegativity in the period at 3.98.

      This increasing trend in electronegativity has significant effects on the properties of elements and their compounds. Elements with higher electronegativity tend to form more polar covalent bonds when combined with elements of lower electronegativity. For example, the bond between hydrogen and fluorine in hydrogen fluoride (HF) is highly polar due to the large difference in electronegativity between the two elements.

      The electronegativity trend also influences the acid-base properties of compounds. Generally, elements with higher electronegativity form stronger acids when combined with hydrogen. This is why hydrofluoric acid (HF) is a stronger acid than hydrochloric acid (HCl), despite fluorine being smaller than chlorine. The higher electronegativity of fluorine results in a greater polarization of the H-F bond, making it easier for the hydrogen to dissociate in solution.

      Furthermore, the electronegativity trend affects the oxidation states of elements. Elements with higher electronegativity are more likely to gain electrons and form negative ions, while those with lower electronegativity tend to lose electrons and form positive ions. This behavior is evident in the formation of ionic compounds, where metals (with low electronegativity) typically lose electrons to non-metals (with high electronegativity).

      The trend of increasing electronegativity across a period also influences the reactivity of elements. Generally, elements with higher electronegativity are more reactive as non-metals. For instance, fluorine, with its extremely high electronegativity, is one of the most reactive elements in the periodic table. It readily forms compounds with almost all other elements, often through highly exothermic reactions.

      In conclusion, the trend of increasing electronegativity across a period in the periodic table is a fundamental concept in chemistry. It arises from the interplay between increasing effective nuclear charge and constant shielding effect as we move from left to right across a period. This trend has far-reaching consequences for chemical bonding, reactivity, and the properties of elements and their compounds. Understanding this trend is essential for predicting and explaining various chemical phenomena and behaviors observed in nature

      Electronegativity Trend Down a Group

      Electronegativity, a fundamental concept in chemistry, plays a crucial role in understanding chemical bonding and reactivity. One of the most important trends observed in the periodic table is the decrease in electronegativity as we move down a group. This trend is consistent across all groups in the periodic table and has significant implications for chemical behavior and properties.

      The primary reason for this decreasing trend in electronegativity down a group is the increase in atomic size. As we move down a group, each successive element has an additional electron shell, resulting in a larger atomic radius. This increase in size means that the outermost electrons are farther from the nucleus, experiencing less attraction from the positively charged protons. Consequently, these electrons are held less tightly by the atom, leading to a decrease in electronegativity.

      Another crucial factor contributing to this trend is electron shielding, also known as screening effect. As we progress down a group, the number of inner electron shells increases. These inner electrons act as a shield between the nucleus and the outermost electrons, reducing the effective nuclear charge experienced by the valence electrons. This shielding effect further weakens the atom's ability to attract electrons, resulting in lower electronegativity values for elements lower in the group.

      To illustrate this trend, let's consider Group 1 (alkali metals) and Group 17 (halogens) as specific examples. In Group 1, we observe a clear decrease in electronegativity from lithium (Li) to francium (Fr). Lithium, with its smaller atomic size and fewer electron shells, has the highest electronegativity in the group. As we move down to sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and finally francium, the electronegativity values progressively decrease.

      Similarly, in Group 17, we see a parallel trend. Fluorine (F), at the top of the group, has the highest electronegativity not only in the group but in the entire periodic table. Moving down to chlorine (Cl), bromine (Br), iodine (I), and astatine (At), we observe a consistent decrease in electronegativity values. This trend is particularly pronounced in the halogens due to their high overall electronegativity, making the differences more noticeable.

      It's important to note that while electronegativity decreases down a group, it generally increases across a period from left to right. This contrasting trend across periods is due to the increasing number of protons in the nucleus without adding new electron shells. As we move across a period, the atomic size decreases due to the stronger nuclear attraction, resulting in higher electronegativity values.

      The implications of these trends are far-reaching in chemistry. For instance, in Group 1, the decreasing electronegativity down the group correlates with increasing reactivity of the alkali metals with water. Lithium reacts relatively slowly with water, while cesium reacts explosively. In Group 17, the decreasing electronegativity affects the strength of halogen bonds and the acidity of hydrohalic acids, with HF being the weakest acid among them due to fluorine's high electronegativity.

      Understanding these trends is crucial for predicting chemical behavior, bond types, and reactivity. The decrease in electronegativity down a group influences various properties such as ionization energy, electron affinity, and metallic character. It also plays a significant role in determining the nature of chemical bonds formed between elements, affecting everything from the properties of simple compounds to the behavior of complex molecules in biological systems.

      In conclusion, the trend of decreasing electronegativity down a group in the periodic table is a fundamental concept in chemistry, driven primarily by increasing atomic size and electron shielding. This trend, exemplified by groups like the alkali metals and halogens, contrasts with the increasing trend across periods. By understanding these patterns, chemists can better predict and explain a wide range of chemical phenomena, making it an essential tool in the study and application of chemistry across various fields.

      Relationship Between Electronegativity and Metallic Character

      The periodic table trends reveal a fascinating inverse relationship between electronegativity and metallic character, two fundamental properties that greatly influence the behavior of elements in chemical reactions and their physical characteristics. Electronegativity, a measure of an atom's ability to attract electrons in a chemical bond, plays a crucial role in determining the metallic properties of elements across the periodic table.

      As we move from left to right across a period in the periodic table trends, electronegativity generally increases. This trend is due to the increasing nuclear charge and decreasing atomic size, which results in a stronger attraction between the nucleus and valence electrons. Conversely, metallic character, which encompasses properties such as electrical conductivity, malleability, and ductility, decreases from left to right across a period.

      This inverse relationship between electronegativity and metallic character is evident in the properties of elements and their compounds. For example, sodium, located on the left side of the periodic table, has low electronegativity and strong metallic character. It readily loses its outermost electron, forming a positive ion and exhibiting typical metallic properties like high conductivity and a lustrous appearance. In contrast, chlorine, positioned on the right side of the table, has high electronegativity and weak metallic character. It tends to gain electrons, forming a negative ion and displaying non-metallic properties such as poor conductivity and a dull appearance.

      The metallic character trend is also observable as we move down a group in the periodic table. As atomic size increases down a group, the attraction between the nucleus and valence electrons weakens, resulting in decreased electronegativity and increased metallic character. This trend explains why cesium, located at the bottom of Group 1, is more reactive and exhibits stronger metallic properties than lithium at the top of the same group.

      The interplay between electronegativity and metallic properties significantly influences the formation and characteristics of chemical compounds. When elements with vastly different electronegativities combine, they tend to form ionic compounds. For instance, the reaction between sodium (low electronegativity) and chlorine (high electronegativity) results in the formation of sodium chloride, an ionic compound with distinct properties such as high melting point and electrical conductivity when dissolved in water.

      On the other hand, when elements with similar electronegativities bond, they form covalent compounds. Carbon and hydrogen, both with intermediate electronegativity values, combine to form various organic compounds with covalent bonds. These compounds often exhibit properties that differ significantly from ionic compounds, such as lower melting points and poor electrical conductivity in their pure form.

      The transition metals, located in the middle of the periodic table, demonstrate how the balance between electronegativity and metallic character can lead to unique properties. These elements have moderate electronegativity values and retain many metallic characteristics while also forming compounds with varying degrees of ionic and covalent bonding. This versatility contributes to their wide range of applications in industry and technology.

      Understanding the inverse relationship between electronegativity and metallic character is crucial for predicting and explaining the behavior of elements and their compounds. This knowledge has practical applications in various fields, including materials science, where it guides the development of new alloys and composites with tailored properties. In chemistry, it helps in predicting reaction outcomes and designing synthetic strategies for creating novel compounds with desired characteristics.

      In conclusion, the inverse relationship between electronegativity and metallic character is a fundamental concept in chemistry that provides valuable insights into the nature of elements and their interactions. By recognizing how these properties change across the periodic table, scientists can better understand and manipulate the behavior of materials at the atomic level, leading to advancements in numerous scientific and technological domains.

      Electronegativity and Chemical Bonding

      Electronegativity plays a crucial role in determining the type of chemical bond formed between atoms. This fundamental concept in chemistry refers to an atom's ability to attract electrons in a chemical bond. The difference in electronegativity between atoms directly influences whether a bond will be ionic, covalent, or polar covalent, each with distinct properties and characteristics.

      Ionic bonds form when there is a large difference in electronegativity between two atoms, typically greater than 1.7 on the Pauling scale. In these bonds, the more electronegative atom completely attracts the electrons from the less electronegative atom, resulting in the formation of ions. For example, sodium chloride (NaCl) exhibits an ionic bond due to the significant electronegativity difference between sodium (0.9) and chlorine (3.0). This electronegativity disparity leads to the transfer of an electron from sodium to chlorine, creating Na+ and Cl- ions. Ionic compounds are characterized by high melting and boiling points, as well as their ability to conduct electricity when dissolved in water or molten.

      Covalent bonds, on the other hand, occur when the electronegativity difference between atoms is small, typically less than 0.4. In these bonds, atoms share electrons rather than transferring them completely. Molecules like hydrogen (H2) and methane (CH4) exemplify covalent bonding. In H2, both hydrogen atoms have the same electronegativity (2.2), resulting in an equal sharing of electrons. Covalent compounds often exist as gases, liquids, or low-melting-point solids at room temperature and generally do not conduct electricity.

      Polar covalent bonds represent an intermediate case, forming when the electronegativity difference falls between 0.4 and 1.7. In these bonds, electrons are shared unequally, with the more electronegative atom having a slightly stronger attraction for the shared electrons. This uneven distribution of electron density creates partial positive and negative charges within the molecule. Water (H2O) is a classic example of a polar covalent molecule. The electronegativity difference between oxygen (3.5) and hydrogen (2.2) results in a slight negative charge on the oxygen atom and slight positive charges on the hydrogen atoms. This polarity gives water its unique properties, such as its ability to dissolve many substances and its high boiling point compared to similarly sized molecules.

      The influence of electronegativity on chemical bonding extends beyond just determining bond types. It also affects various properties of compounds, including solubility, reactivity, and intermolecular forces. For instance, the polarity induced by electronegativity differences in water molecules leads to hydrogen bonding, which is responsible for water's high surface tension and its ability to exist as a liquid over a wide temperature range. In organic chemistry, the electronegativity of different functional groups influences the reactivity and properties of molecules, playing a crucial role in understanding and predicting chemical reactions.

      Understanding the relationship between electronegativity and chemical bonding is essential for predicting and explaining the behavior of substances in various chemical and physical processes. From the formation of simple diatomic molecules to complex biological systems, electronegativity differences shape the interactions between atoms and molecules, ultimately determining the macroscopic properties of materials we observe in everyday life. As chemists continue to explore and manipulate these fundamental principles, new materials with tailored properties can be designed, leading to advancements in fields ranging from materials science to pharmaceutical development.

      Conclusion

      Understanding electronegativity trends in the periodic table is crucial for predicting chemical behavior and properties. As we move from left to right across a period, electronegativity generally increases due to the growing nuclear charge and shrinking atomic size. Conversely, electronegativity decreases as we move down a group, with valence electrons becoming more distant from the nucleus. Fluorine, located in the top-right corner, is the most electronegative element. These trends help explain bond polarity, reactivity, and various chemical phenomena. By grasping these patterns, chemists can anticipate how elements will interact and form compounds. To solidify your understanding of electronegativity trends and their significance in chemistry, we encourage you to rewatch the introduction video. This comprehensive overview will reinforce key concepts and provide a strong foundation for further study of chemical behavior in the periodic table.

      The Periodic Trends of Electronegativity

      Definition of electronegativity.

      Step 1: Introduction to Electronegativity

      Electronegativity is a fundamental concept in chemistry that describes the ability of an atom to attract bonding electrons to its outer electron shell. This property is crucial for understanding how atoms interact and form bonds with each other. In this lesson, we will explore the definition of electronegativity, how it is measured, and its significance in chemical bonding.

      Step 2: Definition of Electronegativity

      Electronegativity is defined as the ability of an atomic nucleus to attract bonding electrons to its outer electron shell. Although we often refer to the atom as a whole, it is the nucleus that plays a central role in this process. The positively charged nucleus attracts negatively charged electrons, influencing the atom's ability to form bonds.

      Step 3: Measuring Electronegativity

      Electronegativity is measured using the Pauling scale, named after the chemist Linus Pauling. The scale ranges from 0 to 4, with fluorine having the highest electronegativity value of 4.0 and francium having the lowest value of 0.7. This scale allows us to compare the electronegativity of different elements and predict their behavior in chemical reactions.

      Step 4: Electronegativity in the Periodic Table

      When examining the periodic table, we observe that electronegativity increases as we move from left to right across a period and decreases as we move down a group. For example, fluorine, located in the top right corner of the periodic table, has the highest electronegativity, while francium, located in the bottom left corner, has the lowest. This trend is due to the increasing positive charge of the nucleus and the decreasing distance between the nucleus and the outer electron shell as we move across a period.

      Step 5: Importance of Electronegativity in Bonding

      Electronegativity plays a crucial role in determining the type of bonds that atoms can form. Atoms with high electronegativity, such as fluorine, are more likely to attract electrons and form ionic bonds. In contrast, atoms with lower electronegativity, such as carbon, are more likely to share electrons and form covalent bonds. The difference in electronegativity between two atoms can also lead to the formation of polar covalent bonds, where electrons are shared unequally.

      Step 6: Application of Electronegativity

      Understanding electronegativity allows us to predict the behavior of elements in chemical reactions and their ability to form bonds. For instance, fluorine's high electronegativity enables it to attract an additional electron and form a fluoride ion, which can then participate in ionic bonding. Similarly, the difference in electronegativity between carbon and fluorine can lead to the formation of strong polar covalent bonds.

      Step 7: Conclusion

      In summary, electronegativity is a key concept in chemistry that describes an atom's ability to attract bonding electrons. It is measured using the Pauling scale and varies across the periodic table. Electronegativity influences the type of bonds that atoms can form and is essential for understanding chemical interactions. By mastering this concept, we can better predict and explain the behavior of elements in various chemical contexts.

      FAQs

      Here are some frequently asked questions about electronegativity trends in the periodic table:

      1. Why does electronegativity increase across a period?

      Electronegativity increases across a period due to increasing effective nuclear charge and decreasing atomic size. As you move from left to right, the number of protons increases while the number of electron shells remains constant. This results in a stronger attraction between the nucleus and valence electrons, leading to higher electronegativity.

      2. Why does electronegativity decrease down a group?

      Electronegativity decreases down a group because of increasing atomic size and electron shielding. As you move down a group, each element has an additional electron shell, placing valence electrons farther from the nucleus. The inner electrons also shield the outer electrons from the nuclear charge, resulting in a weaker attraction and lower electronegativity.

      3. What is the order of electronegativity on the periodic table?

      The general order of electronegativity on the periodic table is: Fluorine > Oxygen > Nitrogen > Chlorine > Bromine > Carbon > Sulfur > Iodine > Hydrogen > Metals. Electronegativity generally increases from bottom-left to top-right on the periodic table, with fluorine being the most electronegative element.

      4. How does electronegativity affect chemical bonding?

      Electronegativity differences between atoms determine the type of chemical bond formed. Large differences (>1.7) result in ionic bonds, small differences (<0.4) form covalent bonds, and intermediate differences (0.4-1.7) create polar covalent bonds. These differences influence properties such as bond polarity, reactivity, and molecular behavior.

      5. What is the relationship between electronegativity and metallic character?

      Electronegativity and metallic character have an inverse relationship. As electronegativity increases, metallic character decreases. This is why metals, which have low electronegativity, are found on the left side of the periodic table, while non-metals with high electronegativity are on the right. This trend affects properties such as conductivity, malleability, and reactivity of elements.

      Prerequisite Topics

      Understanding the concept of electronegativity in periodic trends is crucial for mastering chemistry, but it's essential to recognize that this topic doesn't exist in isolation. To fully grasp the intricacies of electronegativity, students must first have a solid foundation in related concepts. Two key prerequisite topics that play a vital role in comprehending electronegativity are atomic radius and ionization energy.

      The atomic radius relationship is fundamental to understanding electronegativity. As we move across a period in the periodic table, the atomic radius generally decreases. This trend directly influences an atom's ability to attract electrons, which is the essence of electronegativity. Atoms with smaller radii tend to have higher electronegativity because their valence electrons are closer to the nucleus, experiencing a stronger attractive force.

      Similarly, the ionization energy correlation is closely tied to electronegativity. Ionization energy, which is the energy required to remove an electron from an atom, typically increases across a period. This trend aligns with electronegativity because atoms that require more energy to lose an electron are often more effective at attracting electrons in chemical bonds.

      By understanding these prerequisite topics, students can develop a more comprehensive view of periodic trends. The interplay between atomic radius, ionization energy, and electronegativity becomes clearer, allowing for a deeper appreciation of how these properties change across the periodic table. For instance, elements with small atomic radii and high ionization energies often exhibit high electronegativity, as seen in elements like fluorine and oxygen.

      Moreover, grasping these concepts helps in predicting and explaining chemical behavior. The ability to connect atomic radius trends with electronegativity can aid in understanding bond polarity and molecular geometry. Similarly, knowing how ionization energy relates to electronegativity can provide insights into the reactivity of elements and the stability of compounds.

      In conclusion, while studying electronegativity, it's crucial to revisit and reinforce your understanding of atomic radius and ionization energy. These prerequisite topics form the foundation upon which a robust comprehension of electronegativity can be built. By mastering these interconnected concepts, students will be better equipped to tackle more advanced chemistry topics and develop a holistic understanding of periodic trends.

      In this lesson, we will learn:
      • The definition of electronegativity and how it is measured.
      • To apply our understanding of electrostatic principles to the periodic trends in electronegativity.
      • To predict the electronegativity of elements compared to each other.
      • How metallic and acid-base properties change across a period.

      Notes:

      • As seen in Periodic trends: Atomic radius, chemists have found, through experimenting, some principles of electrostatic forces forces that exist because charged particles attract or repel each other. Some principles are:
        • #1: Oppositely charged particles attract each other, while particles of like charge repel each other.
        • #2: The greater the charge difference of two particles, the greater their force of attraction (for example, the attractive force between a 2+ ion and a 2- ion is stronger than the attractive force between a 1+ and a 1- ion).
        • #3: Attractive forces between oppositely charge particles decrease with distance.
        • #4: Repulsive forces between like-charged particles decrease with distance.

      • Electronegativity is the ability of an atom (specifically the nucleus) to attract bonding electrons to its outer electron shell. It is measured using the Pauling scale fluorine is highest at 4.0 on the scale, the most electronegative element, whilst francium is the lowest at 0.7 and is the least electronegative element.

      • Electrostatic theory explains the trend in electronegativity in the Periodic Table, in the Periodic Table both across a period and down a group:
        • As you go across the elements in a period, each element has an increased effective nuclear charge, or Zeff attracting its outer shell electrons. Effective nuclear charge is the positive charge that the outer shell electrons effectively feel from the nucleus. To find this, subtract the number of inner electrons from the number of protons in the atom.
          • In lithium for example, Zeff = 3 - 2 = +1.
            This equation means that two of lithiums 3 protons are cancelled out by the two inner shell electrons and only one is available for the one outer shell electron. Using electrostatic theory: Li is a Zeff 1+ nucleus attracting 1- of electron charge.
          • Moving across to beryllium, Zeff = 4 - 2 = 2+.
            Using electrostatic theory: Be is a 2+ Zeff nucleus attracting 2 outer shell electrons or a 2- charge..

          Increasing Zeff means the nucleus attracts and holds its outer shell electrons with increasingly greater force. It also better attracts other electrons to complete the outer shell.
          Therefore, as you go across the period, it is easier for atoms to attract bonding electrons into their outer shell. This means electronegativity is higher.

        • As you go down a group of the periodic table, each element has an extra inner electron shell as well as increased nuclear charge. However, the increased nuclear charge is cancelled out by the extra inner shell, so Zeff is unchanged.

          • The extra inner shell of electrons also causes shielding of the nucleus. This is where the outer electrons are repelled away from the nucleus because of the like-charge inner shell electrons. In other words, positive attracts negative but not if there is a big wave of negative charge between them!
          This means that going down the group, nuclei lose their ability to attract and hold onto bonding electrons. As a result, electronegativity decreases as you go down a group in the Periodic Table.

        • Remember: Noble gases are not given an electronegativity value because their atoms generally do not form bonds and since they already have full outer shells, they do not attract electrons to complete full outer shells!

      • The trends in electronegativity mean fluorine is the most electronegative element. The effect of electron shielding down a group is more influential than the effect of increased nuclear charge across a period, so oxygen is the second most electronegative element (around 3.5 on the Pauling scale), followed by chlorine (around 3.0).

      • The difference in electronegativity of atoms affects how different atoms bond with one another and can lead to substances of varying properties. This is a very important part of chemistry which the next chapter will look at – bonding between atoms and the properties of compounds they make as a result!>.

      • Electronegativity affects the metallic character of an element (metals have very low electronegativity) and as such, metallic character across a period changes. From left to right across a period, metallic character decreases. For example, in period 3:
        • Sodium, magnesium and aluminium are metals;
        • Silicon is a metalloid (it displays properties of both metals and nonmetals);
        • Phosphorus, sulfur, chlorine and argon are all nonmetals.

      • There is a sliding scale of metallic character across the period which affects several properties. One of these is the properties of oxide compounds. Using period 3 as an example:
        • Sodium oxide (Na2O), magnesium oxide (MgO) and aluminium oxide (Al2O3) are metal oxides with an ionic structure. As such, they usually have high melting points. When dissolved in water, metal oxides form basic solutions with a pH of over 7. The respective equations for sodium oxide and magnesium oxide reacting with water are:

          • Na2O + H2O   \, →  \, 2NaOH

          MgO + H2O   \, →  \, Mg(OH)2

          Sodium hydroxide is a strong base, which is why sodium oxide in water creates a basic solution. As you progress left to right, the metal oxides become less basic in nature.

        • Aluminium oxide is amphoteric; it acts a base reacting with acids and acts as an acid when reacting with bases, producing an aluminium salt both times. It also does not react with water in the way magnesium and sodium oxide does. Its reaction with acid is:

          • Al2O3 + 6HCl   \, →  \, 2AlCl3 + 3H2O

        • As silicon is a metalloid, silicon dioxide has very different properties to metal oxides like Na2O or MgO, in fact it is very weakly acidic. SiO2 has a giant covalent structure and does not react with water as the structure cannot be broken down by the interactions with H2O molecules. As it is weakly acidic, it will react with some metal oxides to produce silicon salts.


        • Phosphorus, sulfur and chlorine oxides are simple covalent molecules; comparing to metal oxides they have much lower melting and boiling points. When dissolved in water the nonmetal oxides form acidic solutions with a pH below 7. There are multiple oxides of phosphorus, sodium and chlorine but they all react to produce an acidic species with water.
          The equation for phosphorus (V) oxide and water is:

          • P4O10 + 6H2O   \, →  \, 4H3PO4


        • Sulfur trioxide reacts to produce sulfuric acid, a reaction which is run on an industrial scale in the contact process:

          • SO3 + H2O   \, →  \, H2SO4


        • In the same manner, chlorine (VII) oxide reacts with water to produce perchloric acid which is a very strong acid.

          • Cl2O7 + H2O   \, →  \, 2HClO4

          Chlorine also has the chlorine (I) oxide form, Cl2O which reacts with water to form a weaker acid, HOCl.

        • This trend of acidic nonmetal oxides is not only in period 3. Oxides of nitrogen like nitrogen dioxide, NO2, react with water to produce nitric acid, HNO3 and the weak acid nitrous acid, HNO2:

          • 2 NO2 + H2O   \, →  \, HNO3 + HNO2

          The nitrous acid breaks down over time into more nitric acid and NO gas. Like those in period 3, nitric acid is a strong acid and will produce very acidic solutions.

          From all of this, you can determine that going from left to right across a period, oxide compounds go from being basic to acidic in nature.