chemistry
Covalent and ionic atomic sizes: covalent radius, which is defined as one-half the distance between the nuclei of two identical atoms when they are joined by a covalent bond (this measurement is possible because atoms within molecules still retain much of their atomic identity). Ionic radius is the measure used to describe the size of an ion. Ionization energy: Energy required to remove an electron from a gaseous atom or ion. Electron affinity: Electron affinity (the energy associated with forming an anion) is more favorable (exothermic) when electrons are placed into lower energy orbitals, closer to the nucleus. Meaning and periodic trends of. Covalent radius increases as we move down a group because the n level (orbital size) increases. Covalent radius mostly decreases as we move left to right across a period because the effective nuclear charge experienced by the electrons increases, and the electrons are pulled in tighter to the nucleus. Anionic radii are larger than the parent atom, while cationic radii are smaller, because the number of valence electrons has changed while the nuclear charge has remained constant. Ionization energy (the energy associated with forming a cation) decreases down a group and mostly increases across a period because it is easier to remove an electron from a larger, higher energy orbital. Electron affinity (the energy associated with forming an anion) is more favorable (exothermic) when electrons are placed into lower energy orbitals, closer to the nucleus. Therefore, electron affinity becomes increasingly negative as we move left to right across the periodic table and decreases as we move down a group. Cation Formation: Cations are the positive ions formed by the loss of one or more electrons. The most commonly formed cations of the representative elements are those that involve the loss of all the valence electrons.Anions formation: are the negative ions formed from the gain of one or more electrons. When nonmetal atoms gain electrons, they often do so until their outermost principal energy level achieves an octet. Formation of ionic compounds: When atoms form positive ions, their valence electrons are transferred to other atoms that need these electrons to attain full valence shells, and which subsequently form negative ions. Thus the formation of positive and negative ions happens simultaneously. Once positive ions and negative ions are formed, they combine to form ionic compounds. Covalent bonding formation: occurs when pairs of electrons are shared by atoms. Atoms will covalently bond with other atoms in order to gain more stability, which is gained by forming a full electron shell.Molecular compounds formation: When atoms of at least two different elements come together to form chemical bonds, these molecules can be called compounds.
Electronegativity: a measure of the tendency of an atom to attract electrons (or electron density) towards itself. It determines how the shared electrons are distributed between the two atoms in a bond. The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity.
The different types of chemical bonds.
Pure Covalent Bonds: This type of bond forms most frequently between two non- metals.
Polar Covalent Bonds: When there is a greater electronegativity difference than between covalently bonded atoms, tend to occur between non-metals.
Ionic Bonds: for atoms with the largest electronegativity differences (such as metals bonding with nonmetals), the bonding interaction is called ionic, and the valence electrons are typically represented as being transferred from the metal atom to the nonmetal.
The influence exerted by the molecular structure on the way the presence of polar bonds affect the molecular polarity of covalent compounds: It causes something called the dipole moment in which the electrons go are attract to the more electronegative element.
The valence shell electron pair repulsion model and the difference between the concepts of electron pair geometry and molecular structure.
Valence shell electron-pair repulsion theory (VSEPR theory) enables us to predict the molecular structure, including approximate bond angles around a central atom, of a molecule from an examination of the number of bonds and lone electron pairs in its Lewis structure. The VSEPR model assumes that electron pairs in the valence shell of a central atom will adopt an arrangement that minimizes repulsions between these electron pairs by maximizing the distance between them. The electrons in the valence shell of a central atom form either bonding pairs of electrons, located primarily between bonded atoms, or lone pairs. The electrostatic repulsion of these electrons is reduced when the various regions of high electron density assume positions as far from each other as possible.
Valence bond theory describes a covalent bond as the overlap of half-filled atomic orbitals (each containing a single electron) that yield a pair of electrons shared between the two bonded atoms. We say that orbitals on two different atoms overlap when a portion of one orbital and a portion of a second orbital occupy the same region of space.
Sigma bonds: The overlap of two s orbitals (as in H2), the overlap of an s orbital and a p orbital (as in HCl), and the end-to-end overlap of two p orbitals (as in Cl2) all produce sigma bonds (σ bonds).
Pi bonds: A pi bond (π bond) is a type of covalent bond that results from the side-by-side overlap of two p orbitals.
How molecular orbital theory differs from the valence bond theory: Valence bond theory assumes that electrons in a molecule are simply the electrons in the original atomic orbitals, with some used while bonding.
In other words, it does not account for the true distribution of electrons within molecules as molecules, but instead, treats electrons as if they are “localized” on the atoms themselves.
On the other hand, molecular orbital theory accounts for the “delocalization” of electrons, and does not assume that the electrons are in the original atomic orbitals.
Instead, it says that the atomic orbitals combine to form the same number of new molecular orbitals as there were atomic orbitals (conservation of orbitals), and the electron distribution is based on that.
The formation of bonding and antibonding molecular orbitals:
1.Adding electrons to these orbitals creates a force that holds the two nuclei together, so we call these orbitals bonding orbitals.
2.The attractive force between the nuclei and these electrons pulls the two nuclei apart. Hence, these orbitals are called antibonding orbitals.
The explanation of why the diatomic molecule of oxygen is paramagnetic.
when we pour liquid oxygen past a strong magnet,it collects between the poles of the magnet and defies gravity, as in Figure8.1. Such attraction to a magnetic field is called paramagnetism, and it arises in molecules that have unpaired electrons. And yet, the Lewis structure of O2 indicates that all electrons are paired.