JAMB Chemistry Note: Atomic Structure and Bonding
Table of Contents
- (a) (i) The Concept of Atoms, Molecules, and Ions
- (a) (ii) Atomic Structure, Electron Configuration, Atomic Number, Mass Number, and Isotopes
- (a) (iii) Shapes of s and p Orbitals
- (b) The Periodic Table and Periodicity of Elements
- Chemical Bonding
(a) (i) The Concept of Atoms, Molecules, and Ions
Atoms: An atom is the smallest unit of matter that retains the identity of an element. It consists of three fundamental particles:
- Protons (positively charged)
- Neutrons (neutral)
- Electrons (negatively charged)
The protons and neutrons are located in the nucleus (center of the atom), while electrons are found in electron shells around the nucleus. The number of protons in an atom defines the atomic number, and the sum of protons and neutrons defines the mass number.
Molecules: A molecule is a group of two or more atoms that are chemically bonded together. Molecules can be formed from atoms of the same element (e.g., O₂ for oxygen gas) or from different elements (e.g., H₂O for water). Molecules may be classified as either covalent (where electrons are shared) or ionic (where electrons are transferred).
Ions: An ion is an atom or molecule that has gained or lost one or more electrons, resulting in a net electrical charge. Ions are categorized as:
- Cations: Positively charged ions (e.g., Na⁺, Mg²⁺).
- Anions: Negatively charged ions (e.g., Cl⁻, SO₄²⁻).
The gain or loss of electrons occurs to achieve a more stable electronic configuration, often resembling that of noble gases.
(a) (ii) Atomic Structure, Electron Configuration, Atomic Number, Mass Number, and Isotopes
Atomic Number (Z): The atomic number represents the number of protons in an atom’s nucleus. The atomic number is unique to each element and determines the identity of the element. For instance, hydrogen has an atomic number of 1 (1 proton), while oxygen has an atomic number of 8 (8 protons).
Mass Number (A): The mass number is the total number of protons and neutrons in an atom’s nucleus. Since electrons have a negligible mass, they do not significantly contribute to the mass number.
- For example, Carbon has a mass number of 12 (6 protons + 6 neutrons). The mass number is used to identify isotopes of an element.
Electron Configuration: Electron configuration refers to the arrangement of electrons in the atomic orbitals of an atom. Electrons fill orbitals starting from the lowest energy level, and the electron configuration of an atom can be determined using the Aufbau Principle, Pauli Exclusion Principle, and Hund’s Rule:
- Aufbau Principle: Electrons fill the orbitals in the order of increasing energy.
- Pauli Exclusion Principle: No two electrons can have the same set of quantum numbers.
- Hund’s Rule: Electrons occupy degenerate orbitals (orbitals of the same energy level) singly before pairing.
The electron configuration can be written by denoting the number of electrons in each orbital:
- Example: The electron configuration of Oxygen (O), atomic number 8, is 1s² 2s² 2p⁴, which shows that Oxygen has two electrons in the first shell (1s²), and six electrons in the second shell (2s² 2p⁴).
Example for elements 1 to 20:
- Hydrogen (H): 1s¹
- Helium (He): 1s²
- Lithium (Li): 1s² 2s¹
- Beryllium (Be): 1s² 2s²
- Boron (B): 1s² 2s² 2p¹
- Carbon (C): 1s² 2s² 2p²
- Nitrogen (N): 1s² 2s² 2p³
- Oxygen (O): 1s² 2s² 2p⁴
- Fluorine (F): 1s² 2s² 2p⁵
- Neon (Ne): 1s² 2s² 2p⁶
- Sodium (Na): 1s² 2s² 2p⁶ 3s¹
- Magnesium (Mg): 1s² 2s² 2p⁶ 3s²
- Aluminum (Al): 1s² 2s² 2p⁶ 3s² 3p¹
- Silicon (Si): 1s² 2s² 2p⁶ 3s² 3p²
- Phosphorus (P): 1s² 2s² 2p⁶ 3s² 3p³
- Sulfur (S): 1s² 2s² 2p⁶ 3s² 3p⁴
- Chlorine (Cl): 1s² 2s² 2p⁶ 3s² 3p⁵
- Argon (Ar): 1s² 2s² 2p⁶ 3s² 3p⁶
- Potassium (K): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹
- Calcium (Ca): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
Isotopes: Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons, leading to different mass numbers.
- Example: Carbon has two stable isotopes, Carbon-12 (¹²C) and Carbon-14 (¹⁴C). Both isotopes have 6 protons, but Carbon-12 has 6 neutrons, while Carbon-14 has 8 neutrons.
(a) (iii) Shapes of s and p Orbitals
The shape of atomic orbitals is crucial in understanding the distribution of electrons in an atom. Orbitals are defined by quantum numbers and describe the probability distribution of finding an electron in a particular region around the nucleus.
- s Orbitals:
- Shape: s orbitals are spherical in shape, meaning the electron density is evenly distributed around the nucleus. Each energy level (starting from n = 1) has one s orbital.
- Electron Capacity: An s orbital can hold a maximum of 2 electrons.
- p Orbitals:
- Shape: p orbitals are dumbbell-shaped, with a region of high electron density on either side of the nucleus. Each energy level (starting from n = 2) has three p orbitals (designated as px, py, and pz) oriented along the three axes (x, y, and z).
- Electron Capacity: Each p orbital can hold a maximum of 2 electrons, so three p orbitals can hold a total of 6 electrons.
(b) The Periodic Table and Periodicity of Elements
The periodic table organizes elements according to increasing atomic number, and it reveals the periodic patterns in the properties of elements. The periodicity of elements is reflected in the repeated trends in properties like atomic size, ionization energy, and electronegativity.
Families of Elements:
- Alkali Metals (Group 1):
- Elements: Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs), Francium (Fr).
- Characteristics: Alkali metals are highly reactive, especially with water. They have a single electron in their outermost shell, which they readily lose to form cations (e.g., Na⁺).
- Halogens (Group 17):
- Elements: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At).
- Characteristics: Halogens are highly reactive nonmetals with seven electrons in their outermost shell. They readily gain an electron to form anions (e.g., Cl⁻).
- Noble Gases (Group 18):
- Elements: Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), Radon (Rn).
- Characteristics: Noble gases have full outer electron shells, making them chemically inert or very stable. They do not readily form bonds.
- Transition Metals:
- Elements: Scandium (Sc), Titanium (Ti), Vanadium (V), Chromium (Cr), Manganese (Mn), Iron (Fe), Cobalt (Co), Nickel (Ni), Copper (Cu), Zinc (Zn).
- Characteristics: Transition metals have variable oxidation states and are often good conductors of heat and electricity.
Periodic Trends:
- Ionization Energy:
- Ionization energy is the energy required to remove an electron from an atom in the gas phase.
- Across a period: Ionization energy increases because the nuclear charge increases, attracting electrons more strongly, making it harder to remove electrons.
- Down a group: Ionization energy decreases because the outermost electrons are farther from the nucleus, and the shielding effect of inner electrons reduces the nuclear attraction.
- Ionic Radii:
- Ionic radius refers to the size of an ion.
- Across a period: Ionic radii decrease as the nuclear charge increases, pulling electrons closer to the nucleus.
- Down a group: Ionic radii increase because additional electron shells are added, increasing the distance from the nucleus.
- Electron Affinity:
- Electron affinity is the energy change when an electron is added to an atom in the gas phase.
- Across a period: Electron affinity becomes more negative (i.e., energy is released) as elements approach a stable octet configuration.
- Down a group: Electron affinity becomes less negative due to increased atomic size and electron shielding.
- Electronegativity:
- Electronegativity is the tendency of an atom to attract electrons in a chemical bond.
- Across a period: Electronegativity increases because the atoms have more protons and a greater pull on the electrons.
- Down a group: Electronegativity decreases because the atoms are larger, and the outer electrons are farther from the nucleus.
Chemical Bonding
Chemical bonding refers to the attractive forces that hold atoms or ions together in molecules, compounds, or crystalline structures. These forces are the result of the interaction between the electrons of atoms, and they play a crucial role in determining the properties of substances, including their physical, chemical, and mechanical characteristics. This note will cover various types of chemical bonds, including electrovalency, covalency, hydrogen bonding, metallic bonding, and coordinate bonding. Special types of bonding forces, such as van der Waals’ forces, will also be discussed.
(c) Chemical Bonding
Chemical bonding can be classified into electrovalent bonds (also known as ionic bonds), covalent bonds, and metallic bonds. In addition, special types of bonds, such as hydrogen bonding, coordinate bonds, and van der Waals’ forces, also play significant roles in the formation and behavior of different compounds.
Electrovalency (Ionic Bonding)
Electrovalency, or ionic bonding, occurs when electrons are transferred from one atom to another, resulting in the formation of oppositely charged ions. This transfer occurs because atoms tend to lose or gain electrons in order to achieve a stable electron configuration, typically similar to that of the nearest noble gas.
- Ionic Bonding: In ionic bonds, an atom with low ionization energy (typically a metal) loses one or more electrons to become a positively charged ion (cation), while an atom with high electron affinity (typically a nonmetal) gains those electrons to become a negatively charged ion (anion). These oppositely charged ions are then held together by strong electrostatic forces.
- Example:
- In sodium chloride (NaCl), sodium (Na) loses one electron to form a cation (Na⁺), and chlorine (Cl) gains the electron to form an anion (Cl⁻). The electrostatic attraction between Na⁺ and Cl⁻ forms the ionic bond, resulting in a neutral compound (NaCl).
- Electron Configuration and Noble Gas Structure:
- Atoms of elements tend to attain the stable electron configuration of noble gases (with a full valence shell of electrons) through bonding. For example, sodium (Na), with an electron configuration of [Ne] 3s¹, loses its one valence electron to attain the stable configuration of neon (Ne), [Ne], while chlorine (Cl), with an electron configuration of [Ne] 3s² 3p⁵, gains the electron to complete its octet, achieving the configuration of argon (Ar), [Ar].
Covalency (Covalent Bonding)
Covalent bonding involves the sharing of electrons between atoms to achieve a stable electron configuration. In covalent bonds, atoms with similar electronegativities (the tendency to attract electrons in a bond) share electrons rather than transferring them.
- Covalent Bonding: Atoms of nonmetals, which have similar electronegativities, will share electrons to achieve a complete valence shell. The shared electrons help each atom reach the stable noble gas configuration.
- Example:
- In a water molecule (H₂O), oxygen shares electrons with two hydrogen atoms. The oxygen atom, which has six valence electrons, shares one electron with each hydrogen atom (which has one valence electron), forming a covalent bond.
- Electron Configuration and Noble Gas Structure:
- The goal of covalent bonding is also to achieve a stable electron configuration. In H₂O, oxygen achieves a full valence shell (similar to neon), and each hydrogen atom achieves the configuration of helium.
Hydrogen Bonding
Hydrogen bonding is a special type of electrovalent bond that occurs when a hydrogen atom, covalently bonded to a highly electronegative atom (like oxygen, nitrogen, or fluorine), is attracted to another electronegative atom in a nearby molecule.
- Mechanism: In a molecule with a highly electronegative atom (such as oxygen in water), the electron pair is pulled closer to the electronegative atom, leaving the hydrogen atom partially positively charged. This positive charge then attracts the lone pair of electrons from another electronegative atom in a neighboring molecule, creating a hydrogen bond.
- Example:
- Water (H₂O) molecules form hydrogen bonds due to the partial positive charge on the hydrogen atoms and the partial negative charge on the oxygen atom. These bonds give water its unique properties, such as its high boiling point and surface tension.
- Importance: Hydrogen bonds play a crucial role in biological systems, such as the structure of DNA and protein folding, and contribute to the physical properties of water.
Metallic Bonding
Metallic bonding occurs between metal atoms, where electrons are not shared or transferred between individual atoms but instead form a “sea” of delocalized electrons that move freely around positively charged metal ions. This type of bonding explains many of the physical properties of metals, such as electrical conductivity and malleability.
- Mechanism: In a metal, atoms release some of their electrons into a common pool of delocalized electrons, which are free to move through the metal lattice. The metal cations (positive ions) are held together by the electrostatic attraction to the delocalized electrons.
- Example:
- In a piece of copper (Cu), the copper atoms release their valence electrons to form a sea of electrons that can move freely, allowing copper to conduct electricity. This metallic bond gives copper its high electrical conductivity.
- Properties:
- Metals are usually good conductors of heat and electricity, malleable (they can be hammered into thin sheets), and ductile (they can be drawn into wires) due to the flexibility of the electron sea.
Coordinate Bond (Dative Covalent Bond)
A coordinate bond (also called a dative covalent bond) is a type of covalent bond where both electrons in the shared pair come from the same atom. This type of bond is commonly found in coordination complexes, where a central metal atom or ion is surrounded by ligands.
- Mechanism: In a coordinate bond, one atom donates a lone pair of electrons to another atom that does not have an electron to share. This forms a bond where both electrons come from the donor atom.
- Example:
- In the complex ion [Fe(CN)₆]³⁻, the iron ion (Fe³⁺) forms coordinate bonds with six cyanide ions (CN⁻). The cyanide ions each donate a lone pair of electrons to the central iron atom, forming a stable complex.
- [Fe(CN)₆]³⁻: In this complex, the cyanide ions (CN⁻) are donating lone pairs to the Fe³⁺ ion, creating coordinate bonds.
- [Cu(NH₃)₄]²⁺: In this complex, four ammonia (NH₃) molecules donate lone pairs of electrons to the copper (Cu²⁺) ion, forming a dative covalent bond.
- [Ag(NH₃)₂]⁺: Similarly, in this complex, two ammonia molecules donate lone pairs to the silver (Ag⁺) ion, forming coordinate bonds.
Van der Waals Forces
Van der Waals forces are weak, non-covalent interactions between molecules or atoms. These forces are responsible for the condensation of gases into liquids and the formation of solid molecular compounds.
There are three types of van der Waals forces:
- Dispersion Forces (London Forces): These are the weakest type of van der Waals forces and arise from temporary fluctuations in the electron distribution around atoms, creating instantaneous dipoles that attract neighboring molecules.
- Example: Dispersion forces are significant in noble gases like helium (He), where the weak interaction between the atoms leads to condensation at very low temperatures.
- Dipole-Dipole Forces: These occur between molecules that have permanent dipoles (positive and negative ends). The positive end of one molecule is attracted to the negative end of another molecule.
- Example: In hydrogen chloride (HCl), the positive hydrogen end of one molecule is attracted to the negative chlorine end of another molecule.
- Hydrogen Bonding: This special type of dipole-dipole interaction, as discussed earlier, occurs when hydrogen is covalently bonded to a highly electronegative atom like oxygen or nitrogen.
Summary of Chemical Bond Types
- Electrovalency (Ionic Bonding): Atoms transfer electrons to form oppositely charged ions, leading to strong electrostatic attraction.
- Covalency: Atoms share electrons to achieve a stable electron configuration, forming covalent bonds.
- Hydrogen Bonding: Special interactions between hydrogen atoms and electronegative atoms (like oxygen, nitrogen, or fluorine).
- Metallic Bonding: Delocalized electrons form a “sea” of electrons around metal ions, giving metals their properties like electrical conductivity and malleability.
- Coordinate Bond: A type of covalent bond where both electrons in the bond come from the same atom, commonly found in coordination complexes.
- Van der Waals Forces: Weak intermolecular forces that include dispersion forces, dipole-dipole forces, and hydrogen bonds.
These different types of bonds contribute to the wide range of chemical compounds and materials found in nature, each with its own unique set of physical and chemical properties.
