To understand electromagnetic radiation and be able to perform calculations involving wavelength, frequency, and energy.
Several properties are used to define waves. Every wave has a wavelength, which is the distance from peak to peak or trough to trough. Wavelength, typically given the symbol λ (lowercase Greek "lambda"), is usually measured in meters. Every wave also has a frequency, which is the number of wavelengths that pass a certain point during a given period of time. Frequency, given the symbol ν (lowercase Greek "nu"), is usually measured in inverse seconds (s−1). Hertz (Hz), another unit of frequency, is equivalent to inverse seconds.
The product of wavelength and frequency is the speed in meters per second (m/s). For light waves, the speed is constant. The speed of light is symbolized by the letter c and is always equal to 2.998×108 m/s in a vacuum; that is,
Another term for "light" is electromagnetic radiation, which encompasses not only visible light but also gamma rays, X-rays, UV rays, infrared rays, microwaves, and radio waves. As you could probably guess, these different kinds of radiation are associated with different energy regimes. Gamma rays have the greatest energy, whereas radio waves have the least energy. The energy (measured in joules) of a photon for a particular kind of light wave is equal to its frequency times a constant called Planck's constant, symbolized h:
These two equations can be combined to give an equation that relates energy to wavelength:
Quantum numbers can be thought of as labels for an electron. Every electron in an atom has a unique set of four quantum numbers.
The principal quantum number n corresponds to the shell in which the electron is located. Thus n can therefore be any integer. For example, an electron in the 2p subshell has a principal quantum number of n=2 because 2p is in the second shell.
The azimuthal or angular momentum quantum number ℓ corresponds to the subshell in which the electron is located. s subshells are coded as 0, p subshells as 1, d as 2, and f as 3. For example, an electron in the 2p subshell has ℓ=1. As a rule, ℓ can have integer values ranging from 0 to n−1.
The magnetic quantum number mℓ corresponds to the orbital in which the electron is located. Instead of 2px, 2py, and 2pz, the three 2p orbitals can be labeled −1, 0, and 1, but not necessarily respectively. As a rule, mℓ can have integer values ranging from −ℓ to +ℓ.
The spin quantum number ms corresponds to the spin of the electron in the orbital. A value of 1/2 means an "up" spin, whereas −1/2 means a "down" spin.
4, 2, -1, -1/2
5, 2, 1, -1/2
4, 2, 1, -1/2
The allowed values for mℓ range from −ℓ to +ℓ and the allowed values for ℓ are integers between zero and n−1. Once you know the value for n, you can determine the acceptable ℓ values, and from there the acceptable mℓ values. The ms values are fixed at either 1/2 or −1/2.
Every electron in an atom is described by a unique set of four quantum numbers: n, ℓ, mℓ, and ms. The principal quantum number, n, identifies the shell in which the electron is found. The angular momentum quantum number, ℓ, indicates the kind of subshell. The magnetic quantum number, mℓ, distinguishes the orbitals within a subshell. The spin quantum number, ms, specifies the electron spin. 1,0,0,1/2
The C atom has two unpaired electrons.
In the Li atom, the 3s, 3p, and 3d orbitals have different energies.
Electrons generally occupy the lowest energy orbital first.
The arrangement of the orbitals in a multielectron atom is different from the arrangement in a single-electron atom owing to the electron-electron repulsions in a multielectron atom. In the case of a single-electron atom, the orbitals in a given principal shell have the same energy. However, in the case of a multielectron atom, the orbitals in a given principal shell have different energies. Electrons occupy the lowest energy orbital first. Each orbital can hold a maximum of two electrons of opposite spins. When more than one orbital of equal energy is available, electrons will first occupy these orbitals singly with parallel spins.
Thus, the C atom has two unpaired electrons in its 2p subshell. The element that follows C is N. It has three unpaired electrons in the 2p subshell. In the N atom, all the three degenerate 2p orbitals are filled with single-electrons each. Thus, it has attained half-filled orbitals. For the next atom, oxygen, the pairing of electrons will occur.
The filling of the electrons in the different orbitals of an atom determines the electron configuration of the atom and indicates the presence or absence of unpaired electrons in the atom.
An atom consists of a small, positively charged nucleus, surrounded by negatively charged electrons. We organize the electrons in a logical manner. As the atomic number increases, electrons are added to the subshells according to their energy. Lower energy subshells fill before higher energy subshells.
The order of filling is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.
The periodic table can be used to help you remember the order.
An orbital-filling diagram shows the number of electrons in each orbital, which are shown in order of energy. The placement of electrons in orbitals follows a certain set of rules.
Lower energy subshells fill before higher energy subshells. The order of filling is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. The periodic table can be used to help you remember this order.
An orbital can hold up to two electrons, which must have opposite spins.
Hund's rule states that if two or more orbitals with the same energy are available, one electron goes in each until all are half full. The electrons in the half-filled orbitals all have the same value of their spin quantum number.
Because the common names of many chemical compounds are not helpful in indicating the chemical nature of the compound (e.g., the common name "table salt" gives no clue that this compound is composed of sodium and chlorine, NaCl), systematic names have been introduced to accurately identify compounds. The naming of simple ionic compounds is relatively straightforward, once you understand the naming system.
Ionic compounds are named according to their cation first, followed by their anion. Once the cation is named, the anion is then named. The naming conventions are listed in the tables to the right.
Finally, although ionic compounds need to have an overall neutral charge, the number of cations or anions are not mentioned in the formula name. For example, AlCl₃ is aluminum chloride, not aluminum trichloride.