Atomic structure and bonding are fundamental concepts in chemistry that help us understand the behavior of matter at the atomic and molecular levels. In this class note, we will explore the basic principles of atomic structure and bonding, and how they relate to the properties of matter.
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- Atoms, Molecules, and Ions
An atom is the smallest unit of an element that retains its chemical identity. Atoms can combine to form molecules, which are two or more atoms bonded together. Molecules can be made up of atoms of the same element (such as O2) or different elements (such as H2O). Ions are charged particles that can be formed from atoms or molecules by gaining or losing electrons.
- Development of the Atomic Structure
The atomic structure was developed through the contributions of several scientists, including Dalton, Thomson, Rutherford, and Bohr. Dalton proposed that atoms were indivisible and had specific masses, while Thomson discovered the electron and proposed the plum pudding model of the atom. Rutherford discovered the nucleus and proposed the nuclear model of the atom. Bohr proposed the planetary model of the atom, which introduced the concept of energy levels.
- Protons, Neutrons, and Electrons
The atomic number of an element is equal to the number of protons in its nucleus, while the mass number is equal to the sum of the number of protons and neutrons. The number of electrons in an atom is equal to the number of protons in a neutral atom. Isotopes are atoms of the same element that have different numbers of neutrons.
- Arrangement of Electrons
Electrons are arranged in energy levels around the nucleus, with each level having a maximum number of electrons it can hold. The first energy level can hold up to 2 electrons, while the second and third energy levels can hold up to 8 electrons each. The arrangement of electrons in an atom follows the Aufbau principle, which states that electrons fill the lowest energy level first before moving to higher energy levels.
- Bonding
There are three main types of bonding: covalent, ionic, and metallic. Covalent bonding occurs when two atoms share electrons to form a molecule. Ionic bonding occurs when electrons are transferred from one atom to another to form ions that are held together by electrostatic forces. Metallic bonding occurs when electrons are shared by all atoms in a metal.
- Shapes of Molecules
The shape of a molecule is determined by the arrangement of its atoms and the lone pairs of electrons around the central atom. The VSEPR theory predicts the shape of a molecule based on the number of electron pairs around the central atom.
- Nuclear Reactions
Nuclear reactions involve changes in the nucleus of an atom and can be either natural or artificial. Radioactive decay involves the emission of alpha, beta, or gamma particles from the nucleus. The half-life of a radioactive material is the time it takes for half of the material to decay.
- Applications of Radioactivity
Radioactivity has many applications, including in medicine (such as in cancer treatment), energy production (such as in nuclear power plants), and dating archaeological artifacts.
Overall, understanding atomic structure and bonding is crucial to understanding the behavior of matter at the atomic and molecular levels. With these basic principles, we can begin to understand the properties of different elements and compounds, and how they interact with each other.
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Chemistry Keypoints: Atomic Structure And Bonding
(i) Distinguish between atoms, molecules, and ions: Chemistry Keypoints: Atomic Structure And Bonding
Atoms are the basic units of matter that make up all elements. A molecule is a group of two or more atoms held together by chemical bonds. An ion is an atom or molecule that has lost or gained electrons and has a positive or negative charge.
(ii) Identify the contributions of these scientists to the development of the atomic structure:
- John Dalton proposed the Atomic Theory, which stated that all matter is made up of atoms and that each element is made up of atoms of a single type.
- J.J. Thomson discovered the electron and proposed the Plum Pudding Model, which stated that atoms were made up of a positively charged substance with negatively charged electrons scattered throughout.
- Ernest Rutherford conducted the Gold Foil Experiment and proposed the Nuclear Model, which stated that atoms have a small, dense nucleus containing positively charged protons, with electrons orbiting the nucleus.
- Niels Bohr proposed the Planetary Model, which stated that electrons orbit the nucleus in specific energy levels or shells.
(iii) Deduce the number of protons, neutrons, and electrons from the atomic and mass numbers of an atom:
The number of protons in an atom is equal to its atomic number (Z). The number of electrons in a neutral atom is also equal to its atomic number (Z). The number of neutrons can be calculated by subtracting the atomic number from the mass number (A-Z).
(iv) Apply the rules guiding the arrangement of electrons in an atom: The rules guiding the arrangement of electrons in an atom include:
- Aufbau Principle: electrons fill the lowest energy orbitals first
- Pauli Exclusion Principle: each orbital can hold a maximum of two electrons with opposite spin
- Hund’s Rule: when orbitals of equal energy are available, electrons will occupy them singly before pairing up.
(v) Identify common elements exhibiting isotopy:
Isotopes are atoms of the same element with different numbers of neutrons. Some common elements exhibiting isotopy include carbon (C-12, C-13, C-14), hydrogen (H-1, H-2, H-3), and oxygen (O-16, O-17, O-18).
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(vi) Relate isotopy to mass number:
Isotopy is related to the mass number (A) of an atom. Isotopes of the same element have the same atomic number (Z) but different mass numbers because they have different numbers of neutrons. The mass number is equal to the sum of the number of protons and neutrons in an atom.
(vii) Perform simple calculations relating to isotopy:
Isotopes are atoms of the same element that have different numbers of neutrons in their nuclei. This means they also have different mass numbers. The atomic number, which is the number of protons in an atom, remains the same for all isotopes of an element.
To calculate the number of neutrons in an isotope, subtract the atomic number from the mass number:
Number of neutrons = Mass number – Atomic number
For example, to calculate the number of neutrons in carbon-14:
Mass number of carbon-14 = 14 Atomic number of carbon-14 = 6
Number of neutrons = 14 – 6 = 8
Therefore, carbon-14 has 8 neutrons.
(viii) Differentiate between the shapes of the orbitals:
An orbital is a region in an atom where an electron is most likely to be found. The shape of an orbital is determined by the energy of the electron and its position in the atom.
There are four main types of orbitals: s, p, d, and f. The s orbital is spherical in shape, while the p orbital is dumbbell-shaped. The d and f orbitals are more complex in shape.
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(ix) Determine the number of electrons in s and p atomic orbitals:
The s orbital can hold a maximum of 2 electrons, while the p orbital can hold a maximum of 6 electrons. The number of electrons in an atom’s s and p orbitals is determined by its electron configuration, which is the arrangement of electrons in the atom’s energy levels.
(x) Relate atomic number to the position of an element on the periodic table:
The periodic table arranges elements in order of increasing atomic number. Elements with similar chemical and physical properties are grouped together in columns, called groups. The rows of the periodic table are called periods.
(xi) Relate properties of groups of elements on the periodic table:
Elements in the same group on the periodic table have similar chemical and physical properties because they have the same number of valence electrons, which are the outermost electrons involved in chemical bonding. For example, the alkali metals in Group 1 have one valence electron and are highly reactive, while the noble gases in Group 18 have a full outer shell of electrons and are unreactive. Chemistry Keypoints: Atomic Structure And Bonding
(xii) Identify reasons for variation in properties across the period and down the groups:
Properties of elements vary across a period due to the change in the number of electrons in the outermost shell, or valence shell, as you move from left to right across the period. This change affects the element’s reactivity, electronegativity, and ionization energy.
Properties of elements vary down a group due to the change in the number of energy levels, or shells, as you move down the group. The number of valence electrons remains the same within a group, so the chemical properties of the elements in the group are similar. The size of the atoms also increases down a group, which affects the element’s ionization energy and electronegativity.
(xiii) Differentiate between the different types of bonding:
There are three types of chemical bonding: ionic, covalent, and metallic bonding. Ionic bonding occurs between atoms that have a large difference in electronegativity, with one atom donating an electron to another. Covalent bonding occurs between atoms that share electrons to fill their outermost electron shell, and metallic bonding occurs between metal atoms that share their electrons.
(xiv) Deduce bond types based on electron configurations:
The bond type can be determined based on the electronegativity difference between two atoms. If the difference is greater than 1.7, the bond is ionic. If the difference is less than 1.7 but greater than 0.5, the bond is polar covalent. If the difference is less than 0.5, the bond is nonpolar covalent.
(xv) Relate the nature of bonding to properties of compounds:
The nature of the bonding affects the properties of compounds. Ionic compounds have high melting and boiling points, are soluble in water, and conduct electricity when dissolved or melted. Covalent compounds have low melting and boiling points, are often insoluble in water, and do not conduct electricity. Metallic compounds have high melting and boiling points, are insoluble in water, and conduct electricity well.
(xvi) differentiate between the various shapes of molecules:
Molecules can have different shapes depending on the arrangement of atoms and lone pairs of electrons around the central atom. The shapes include linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral.
(xvii) Distinguish between an ordinary chemical reaction and a nuclear reaction:
An ordinary chemical reaction involves a rearrangement of electrons between atoms to form new compounds. The nuclei of the atoms are not affected. In contrast, a nuclear reaction involves changes in the nucleus of an atom, such as nuclear decay, fusion, or fission. Nuclear reactions release much more energy than ordinary chemical reactions.
(xviii) differentiate between natural and artificial radioactivity:
Natural radioactivity is the spontaneous and continuous decay of certain unstable isotopes present in nature, such as uranium, thorium, and potassium. Artificial radioactivity, on the other hand, is the production of radioactive isotopes by human-made nuclear reactions.
(xix) Compare the properties of the different types of nuclear radiation:
There are three types of nuclear radiation: alpha particles, beta particles, and gamma rays. Alpha particles are helium nuclei consisting of two protons and two neutrons, beta particles are high-speed electrons or positrons, and gamma rays are high-energy electromagnetic radiation.
Alpha particles have relatively low penetrating power and can be stopped by a sheet of paper or a few centimeters of air. Beta particles have moderate penetrating power and can be stopped by a sheet of aluminum or a few meters of air. Gamma rays have high penetrating power and can be stopped only by thick lead or concrete walls.
(xx) compute simple calculations on the half-life of radioactive material:
The half-life of a radioactive material is the time taken for half of the radioactive nuclei in a sample to decay. It can be calculated using the formula:
t1/2 = ln(2) / λ
where t1/2 is the half-life, ln(2) is the natural logarithm of 2 (0.693), and λ is the decay constant of the material. For example, if the decay constant of a material is 0.01 per year, its half-life would be:
t1/2 = ln(2) / 0.01 = 69.3 years
(xxii) Identify the various applications of radioactivity: Radioactivity has a wide range of applications in various fields, including:
- Medical diagnosis and treatment: Radioactive isotopes can be used in medical imaging techniques, such as X-rays, CT scans, and PET scans, to diagnose and monitor diseases. Radioactive isotopes can also be used in radiation therapy to treat cancer.
- Industrial applications: Radioactive isotopes can be used in industrial radiography to inspect the integrity of welds and other structures. They can also be used to measure the thickness of materials and the level of liquids in tanks.
- Energy production: Nuclear power plants use radioactive isotopes to generate electricity. The energy released from the fission of uranium atoms is used to heat water and produce steam, which drives turbines to generate electricity.
- Archaeology and geology: Radioactive isotopes can be used to date archaeological and geological specimens. By measuring the number of radioactive isotopes present in a sample and comparing it with the half-life of the isotopes, scientists can determine the age of the sample.
- Food preservation: Irradiation can be used to sterilize food and prolong its shelf life by killing bacteria and insects.
- Environmental monitoring: Radioactive isotopes can be used to trace the movement of pollutants and study the effects of human activities on the environment.
