Atomic Structure and Interatomic Bonding PDF

# Atomic Structure and Interatomic Bonding ## Atomic Structure - Fundamental Concepts - Each atom consists of a very small nucleus composed of **protons** and **neutrons** and is encircled by moving **electrons** (quarks, neutrinos, and bosons). - **Electrons** and **protons** are electrically cha...

# Atomic Structure and Interatomic Bonding ## Atomic Structure - Fundamental Concepts - Each atom consists of a very small nucleus composed of **protons** and **neutrons** and is encircled by moving **electrons** (quarks, neutrinos, and bosons). - **Electrons** and **protons** are electrically charged: - **Electrons** are negatively charged. - **Protons** are positively charged. - **Neutrons** are electrically neutral. - Masses for these subatomic particles are extremely small. **Protons** and **neutrons** have approximately the same mass, which is significantly larger than that of an **electron**. - Each chemical element is characterized by the number of **protons** in the nucleus - the **atomic number** (Z). - For an electrically neutral or complete atom, the **atomic number** also equals the number of **electrons**. - This **atomic number** ranges in integral units from 1 for **hydrogen** to 92 for **uranium**, the highest of the naturally occurring elements. - The **atomic mass** (A) of a specific atom may be expressed as the sum of the masses of all atoms of a given element. The number of **neutrons** (N) may be variable. - Thus atoms of some elements have two or more different **atomic masses**, which are called **isotopes**. - The **atomic weight** of an element corresponds to the weighted average of the **atomic masses** of the atom’s naturally occurring **isotopes**. - The **atomic mass unit** (amu) may be used to compute **atomic weight**. ## Electrons in Atoms - Atomic Models - During the late 19th century, classical mechanics failed to explain various phenomena involving **electrons** in solids. This led to the development of **quantum mechanics**: - It is a set of principles governing atomic and subatomic systems. - Understanding **electron** behavior in atoms and solids requires quantum-mechanical concepts. - An early outcome of **quantum mechanics** was the **Bohr atomic model**. This model described **electrons** as revolving in discrete **orbitals** around the **nucleus** with well-defined positions. - A key principle of **quantum mechanics** is the quantization of **electron energy**. This means **electrons** can only occupy specific energy levels and must make "**quantum jumps**" to transition between these levels, absorbing or emitting energy in the process. - The **Bohr model** illustrated these concepts but had limitations in explaining certain **electron behaviors**. - The **Bohr model** was succeeded by the **wave-mechanical model**. This model treats **electrons** as exhibiting both wave-like and particle-like properties. - The **wave-mechanical model** replaces discrete **orbitals** with **probability distributions** or "**electron clouds**". These represent the likelihood of an **electron** being present at various locations around the **nucleus**. ## Quantum Numbers - In **wave mechanics**, every **electron** in an atom is characterized by four parameters called **quantum numbers**. - The size, shape, and spatial orientation of an **electron**'s probability density (or **orbital**) are specified by three of these **quantum numbers**. - **Bohr energy levels** separate into **electron subshells**. **Quantum numbers** dictate the number of states within each **subshell**. - **Shells** are specified by a **principal quantum number** (n), which may take on integral values beginning with unity. **Shells** are sometimes designated by the letters K, L, M, N, O, and so on, which correspond, respectively, to n 1, 2, 3, 4, 5. - The **quantum number** (l) is related to the size of an **electron**'s **orbital** (or its average distance from the **nucleus**). - The **quantum number** (l) designates the **subshell**. Values of I are restricted by the magnitude of n and can take on integer values that range from I=0 to I = (n - 1). | Value of l | Letter Designation | |---|---| | 0 | s | | 1 | p | | 2 | d | | 3 | f | - Each **subshell** is denoted by a lowercase letter - an s, p, d, or f - related to the **quantum number** (l) values. - **Electron orbital shapes** depend on **quantum number** (l). For example, **s orbitals** are spherical and centered on the **nucleus**. There are three **orbitals** for a p **subshell**, each having a nodal surface in the shape of a **dumbbell**. - **Axes** for these three **orbitals** are mutually perpendicular to one another like those of an x-y-z coordinate system. It is convenient to label these **orbitals** px, py, and pz. - The number of **electron orbitals** for each **subshell** is determined by the third (or magnetic) **quantum number**, (m). - The **quantum number** (m) can take on integer values between -l and +l, including 0. When I=0, m can only have a value of 0 because +0 and -0 are the same. This corresponds to an s **subshell**, which can have only one **orbital**. - For l = 1, m can take on values of '-1, 0, and +1, and three p **orbitals** are possible. Similarly, it can be shown that d **subshells** have five **orbitals** and f **subshells** have seven. In the absence of an external magnetic field, all **orbitals** within each **subshell** are identical in energy. However, when a magnetic field is applied, these **subshell states** split, with each **orbital** assuming a slightly different energy. - Associated with each **electron** is a **spin moment**, which must be oriented either up or down. - The fourth **quantum number**, (ms), is related to this **spin moment** and can have two values: +1/2 (for **spin up**) and -1/2 (for **spin down**). - A schematic showing the relative energies of **electrons** for the various shells and subshells is shown. It shows the energy levels from lowest to highest for the s, p, d, and f subshells. - **The principal quantum number** (n) dictates the energy level of an **electron**. The smaller the **principal quantum number**, the lower the energy level, for example, the energy of a 1s state is less than that of a 2s state, which in turn is lower than that of the 3s. - **The quantum number** (l) also influences the energy of an **orbital**. Within each shell, the energy of a **subshell** increases with the value of the **quantum number** (l). For example, the energy of a 3d state is greater than that of a 3p, which is larger than 3s. - There may be overlap in the energy of a state in one shell with states in an adjacent shell. This is especially true of d and f states. For example, the energy of a 3d state is generally greater than that of a 4s. ## Electron configurations - **Electron states** are values of energy that are permitted for **electrons**. - To determine the manner in which these **states** are filled with **electrons**, we use the Pauli exclusion principle. This is another quantum-mechanical concept, which stipulates that each **electron state** can hold no more than two **electrons** that must have opposite **spins**. - Thus, s, p, d, and f **subshells** may each **accomodate**, respectively, a total of 2, 6, 10, and 14 **electrons**. - Not all possible **states** in an atom are filled with **electrons**. - The **electrons** fill up the lowest possible energy **states** in the **electron shells** and **subshells**, two **electrons** (having opposite **spins**) per **state** for most atoms. - When all the **electrons** occupy the lowest possible energies in accord with the foregoing restrictions, an atom is said to be in its ground state. - The **electron configuration** or structure of an atom represents the manner in which these **states** are occupied. - The number of **electrons** in each **subshell** is indicated by a superscript after the shell- subshell designation in the conventional notation. - For example, the **electron configurations** for hydrogen, helium, and sodium, respectively, are 1s¹, 1s², and 1s²2s²2p⁶3s¹. - **Valence electrons** are those that occupy the outermost shell. - **Valence electrons** are extremely important, as they participate in the bonding between atoms to form atomic and molecular aggregates. - Many of the physical and chemical properties of solids are based on these **valence electrons**. - Some atoms have what are termed stable **electron configurations**. In this case, the **states** within the outermost or valence **electron shell** are completely filled. This normally corresponds to the occupation of just the s and p **states** for the outermost shell by a total of eight **electrons**, as in neon, argon, and krypton. One exception is helium, which contains only two 1s **electrons**. - These elements (Ne, Ar, Kr, and He) are the inert, or noble gases, which are virtually unreactive chemically. - Some atoms of the elements that have unfilled valence shells assume stable **electron configurations** by gaining or losing **electrons** to form charged ions or by sharing **electrons** with other atoms. - This is the basis for some chemical reactions and also for atomic bonding in solids. ## The Periodic Table - The **periodic table** classifies all elements according to their **electron configuration**. - The elements are situated, with increasing **atomic number**, in seven horizontal rows called **periods**. - All elements arrayed in a given column or group in the **periodic table** have similar valence **electron structures**, as well as chemical and physical properties. - These properties change gradually, moving horizontally across each **period** and vertically down each column. - The elements positioned in Group 0, the rightmost group, are the inert gases, which have filled **electron shells** and stable **electron configurations**. - Group VIIA and VIA elements are one and two **electrons** deficient, respectively, from having stable structures. - The Group VIIA elements (F, Cl, Br, I, and At) are sometimes termed the halogens. - The alkali and the alkaline earth metals (Li, Na, K, Be, Mg, Ca, etc.) are labeled as Groups IA and IIA, having, respectively, one and two **electrons** in excess of stable structures. - The elements in the three long periods, Groups IIIB through IIB, are termed the transition metals. These have partially filled d **electron states**, and in some cases one or two **electrons** in the next higher energy shell. - Groups IIIA, IVA, and VA (B, Si, Ge, As, etc.) display characteristics that are intermediate between the metals and nonmetals by virtue of their valence **electron structures**. - A table showing the **periodic table** with atomic number, symbol, and atomic weight is shown. - Most of the elements really come under the metal classification. These are sometimes termed electropositive elements, indicating that they are capable of giving up their few valence **electrons** to become positively charged ions. - The elements situated on the right side of the **periodic table** are electronegative. This means they readily accept **electrons** to form negatively charged ions, or sometimes they share **electrons** with other atoms. - As a general rule, **electronegativity** increases in moving from left to right and from bottom to top. - Atoms are more likely to accept **electrons** if their outer shells are almost full and if they are less “shielded” from (i.e., closer to) the nucleus. - A figure showing the **electronegativity of elements** in the periodic table is shown. ## Atomic Bonding In Solids ## Bonding Forces and Energies - An understanding of many of the physical properties of materials is enhanced by a knowledge of the interatomic forces that bind the atoms together. - The principles of atomic bonding can be illustrated by considering how two isolated atoms interact as they are brought close together from an infinite separation. - At large distances, interactions are negligible because the atoms are too far apart to have an influence on each other. However, at small separation distances, each atom exerts forces on the others. - These forces are of two types, attractive (FA) and repulsive (FR). The magnitude of each depends on the separation or interatomic distance (r). - The origin of an attractive force FA depends on the particular type of bonding that exists between the two atoms. - Repulsive forces arise from interactions between the negatively charged electron clouds of the two atoms. These are important only at small values of r as the outer electron shells of the two atoms begin to overlap. - A graph showing the attractive, repulsive, and net forces between two atoms and their dependence on interatomic separation is shown. - The net force FN between the two atoms is just the sum of both attractive and repulsive components: FN = FA + FR. - Sometimes it is more convenient to work with the potential energies between two atoms instead of forces: E = ∫F dr. - Mathematically, energy is (E) and force is (F). - EN, EA, and ER are, respectively, the net, attractive, and repulsive energies for two isolated and adjacent atoms. - EN = ∫FN dr - EN =∫ FA dr + ∫FR dr - EN = EA + ER - A graph showing the attractive, repulsive, and net potential energies between two atoms and their dependence on interatomic separation is shown. - Although the preceding treatment deals with an ideal situation involving only two atoms, a similar yet more complex condition exists for solid materials because force and energy interactions among atoms must be considered. - Nevertheless, a bonding energy, analogous to E0, may be associated with each atom. The magnitude of this bonding energy and the shape of the energy-versus-interatomic separation curve vary from material to material, and they both depend on the type of atomic bonding. - A number of material properties depend on E0, the curve shape, and bonding type. For example, materials having large bonding energies typically also have high melting temperatures; at room temperature, solid substances are formed for large bonding energies, whereas for small energies, the gaseous state is favored; liquids prevail when the energies are of intermediate magnitude. - Mechanical stiffness (or modulus of elasticity) of a material is dependent on the shape of its force-versus-interatomic separation curve. The slope for a relatively stiff material at the r =r0 position on the curve will be quite steep; slopes are shallower for more flexible materials. - How much a material expands upon heating or contracts upon cooling (i.e., its linear coefficient of thermal expansion) is related to the shape of its E-versus-r curve. - A deep and narrow "trough," which typically occurs for materials having large bonding energies, normally correlates with a low coefficient of thermal expansion and relatively small dimensional alterations for changes in temperature. - A graph showing the Lennard-Jones Interatomic Potential is shown. This graph shows the energy of interaction between two atoms as a function of the distance between their nuclei. - Three different types of primary or chemical bond are found in solids— ionic, covalent, and metallic. - For each type, the bonding necessarily involves the valence electrons; furthermore, the nature of the bond depends on the electron structures of the constituent atoms. - In general, each of these three types of bonding arises from the tendency of the atoms to assume stable electron structures, like those of the inert gases, by completely filling the outermost electron shell. - Secondary or physical forces and energies are also found in many solid materials; they are weaker than the primary ones but nonetheless influence the physical properties of some materials. ## Primary Interatomic Bonds - **Ionic Bonding:** Ionic bonding is perhaps the easiest to describe and visualize. It is always found in compounds composed of both metallic and nonmetallic elements, elements situated at the horizontal extremities of the periodic table. Atoms of a metallic element easily give up their valence electrons to the nonmetallic atoms. In the process, all the atoms acquire stable or inert gas configurations (i.e., completely filled orbital shells). They also acquire an electrical charge and become ions. - **Sodium chloride** (NaCl) is the classic ionic material. A sodium atom can assume the electron structure of neon (and a net single positive charge with a reduction in size) by a transfer of its one valence 3s electron to a chlorine atom. After such a transfer, the chlorine ion acquires a net negative charge, an electron configuration identical to that of argon. It is also larger than the chlorine atom. - A diagram showing the transfer of an electron from a sodium atom to a chlorine atom is shown. - The attractive bonding forces are coulombic. This means that positive and negative ions, by virtue of their net electrical charge, attract one another. - A diagram showing the ionic bonding in sodium chloride is shown. - A table showing the bonding energies and melting temperatures for various ionic materials is shown. - **Covalent Bonding:** Covalent bonding, is found in materials whose atoms have small differences in electronegativity—that is, that lie near one another in the periodic table. For these materials, stable electron configurations are assumed by the sharing of electrons between adjacent atoms. Two covalently bonded atoms will each contribute at least one electron to the bond, and the shared electrons may be considered to belong to both atoms. - A diagram showing the formation of a covalent bond between two hydrogen atoms is shown. - Many nonmetallic elemental molecules (e.g., Cl2, F2), as well as molecules containing dissimilar atoms, such as CH4, H2O, HNO3, and HF, are covalently bonded. - This type of bonding is found in elemental solids such as diamond (carbon), silicon, and germanium and other solid compounds composed of elements that are located on the right side of the periodic table, such as gallium arsenide (GaAs), indium antimonide (InSb), and silicon carbide (SiC). - Covalent bonds may be very strong, as in diamond, which is very hard and has a very high melting temperature, 3550Cor they may be very weak, as with bismuth, which melts at about 270C. - A table showing the bonding energies and melting temperatures for various covalent materials is shown. - Bonding energies and melting temperatures for a few covalently bonded materials are presented in Table 2.3. Inasmuch as electrons participating in covalent bonds are tightly bound to the bonding atoms, most covalently bonded materials are electrical insulators, or, in some cases, semiconductors. - Mechanical behaviors of these materials vary widely: some are relatively strong, others are weak; some fail in a brittle manner, whereas others experience significant amounts of deformation before failure. It is difficult to predict the mechanical properties of covalently bonded materials on the basis of their bonding characteristics. - **Metallic Bonding:** The metallic bond: - Is found in metals and their alloys. - Is a result of the electrostatic attraction between the positively charged metal ions and the negatively charged free electrons. - The free electrons are able to move freely throughout the metal, giving it good electrical and thermal conductivity. - Metals are generally strong and ductile because the free electrons can easily move around the metal ions, allowing the metal to deform without breaking. - A table showing the bonding energies and melting temperatures for various metallic materials is shown. - **Van der Waals Bonding:** - Is a weak intermolecular force that arises from temporary fluctuations in the electron distribution around atoms and molecules. - Though weak, van der Waals forces are responsible for the condensation of gases into liquids and the solidification of liquids into solids. - A table showing the bonding energies and melting temperatures for various materials bonded by van der Waals forces is shown. - **Hydrogen Bonding:** - Is a type of intermolecular force that occurs when a hydrogen atom is bonded to a highly electronegative atom, such as oxygen or nitrogen. - The hydrogen atom is attracted to the electronegative atom in an adjacent molecule, forming a weak bond. - Hydrogen bonding is responsible for the high boiling point of water and the unique properties of DNA. - A table showing the bonding energies and melting temperatures for various materials bonded by hydrogen bonding is shown. ## Returning the exercise to Moodle - Create two practice exercises from lecture slides or videos. - Write the assignment and solution on the ppt slides. - Return the assignment to Moodle by 30.1.2025. - Be inventive. ## Thanks