Chemistry: Atomic Structure Flashcards
Atom
The basic building block of matter. The smallest unit of a chemical element. Composed of subatomic particles called protons, neutrons, and electrons. The protons and neutrons form the nucleus.
All atoms of an element show similar properties.
John Dalton’s Atomic Theory
All elements are composed of very small particles called atoms. All atoms of a given element are identical in size, mass, and chemical properties. The atoms of one element are different from atoms of all other elements.
All compounds are composed of atoms of more than one element. For any given compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction.
A given chemical reaction involves only the separation, combination, or rearrangement of atoms; it does NOT result in the creation or destruction of atoms.
Protons
Protons carry a single positive charge and have a mass of one atomic mass unit or amu. The atomic number (Z) of an element is equal to the number of protons found in an atom of that element. All atoms of an element have the same Z.
Protons carry the same “quantity of charge” as an electron; however, they have a weight that is 1840 times heavier than that of an electron.
Nucleus Weight & Volume
The mass of the nucleus of an atom comprises almost the entire weight of the atom; but it only occupies 1/10^13 of the volume of the atom.
Neutrons
Carry no charge and have a mass only slightly larger than protons. Different isotopes of one element have different #s of neutrons but same # of protons.
The mass number (A) is = to the total # of protons and neutrons.
Electrons
Carry a charge equal in magnitude but opposite in sign to that of protons. AN electron has a very small mass, 1/1837 the mass of a proton or neutron, which is negligible for most purposes.
Electrons farthest from nucleus are known as valence electrons. Farther the valence electrons are from the nucleus, the weaker the attractive force of the positively charged nucleus and the more likely the valance electrons are to be influenced by other atoms.
The valance electrons and their activity determine reactivity of an atom.
Ion
An atom that loses or gains an electron and changes its charge.
Atomic Weight
The weight in grams of one mole (mol) of a given element and is expressed in terms of g/mol.
Mole
A unit used to count particles and is represented by Avogadro’s number, 6.022 x 10^23 particles.
For example, the atomic weight of carbon is 12.0g/mol, which means that 6.022x10^23 carbon atoms weigh 12.0 g.
Isotopes
For an element, multiple species of atoms with the same number of protons (same atomic number) but different numbers of neutrons (different mass numbers) exist; these are called isotopes of the element.
Isotopes are referred to either by the convention described above or, more commonly, by name of the element followed by the mass number. Like carbon-12 and carbon-14.
Almost all elements exist as a collection of two or more isotopes.
Quantum Theory
Proposed by Max Planck in 1900. Proposed that energy emitted as electromagnetic radiation from matter comes in discrete bundles called quanta. The energy value of a quantum is given by the equation E=hf, where h is a proportionality constant known as Planck’s constant, equal to 6.626x10^-34 Jxs, and f (sometimes designated v) is the frequency of the radiation.
The Bohr Model
In 1913, Niels Bohr developed his model of the electronic structure of the hydrogen atom. He assumed that the hydrogen atom consisted of a central proton around which an electron traveled in circular orbit, and that the centripetal force acting on the electron as it revolved around the nucleus was the electrical force between the positively charged proton and the negatively charged electron.
Bohr’s model used the quantum theory of Planck in conjunction with concepts from classic physics. In classic mechanics, an object, such as an electron, revolving in a circle may assume an infinite # of values for its radium and velocity. Therefore, the angular momentum (L=mvr) and kinetic energy (KE=[mv^2]/2) can take on any value. But, by incorporating the quantum theory, Bohr placed conditions on the value of the angular momentum., which created this equation of angular momentum for an electron.
Angular Momentum = (nh)/(2pi)
Where h is Planck’s constant 6.626x10^-34 Jxs and n is a quantum number that can be any positve integer. Since h, 2, and pi are constants, the angular momentum changes only in discrete amounts with respect to the quantum number.
Bohr then equated the allowed values of the angular momentum to the energy of the electron. He obtained
E= -(Rh)/(n^2)
Where Rh is the Rudverg constant 2.18x10^-18 J/electron. Therefore, like angular momentum, the energy of the electron also changes in discrete amounts WRT the quantum number.
A value of 0 energy is assigned to the state in which the proton and electron were separated completely, no attractive force. So, the electron in any quantisized states in the atom would have a negative energy as a result of the attractive forces. This explains the negative sign in the above equation.
Applications of the Bohr Model
In his model of the structure of hydrogen, Bohr postulated that an electron can exist only in certain fixed-energy states. In terms of quantum theory, the energy of an electron is quantized. Using this model, generalizations concerning the electrons can be made.
The energy of the electron is related to its orbital radius: the small the radius, the lower the energy state of the electron. Smallest orbit is n=1, which is the ground state of the hydrogen electron.
Atomic Emission Spectra
At room temp, majority of atoms in a sample are in the ground state. But electrons can be excited to higher energy levels, by heat or other energy, to yield the excited state of the atom. The lifetime of the excited state is brief, so electrons will return rapidly to ground state, while emitting energy in form of photons. The electromagnetic energy of these photons may be determined using the following equation.
E = hc/wavelength
h is Planck’s constant, c is velocity of light (3.00 x 10^8 m/s), and wavelength is wavelength of the radiation.
When the electrons emit these photons, the quatisized energies of light emitted under these conditions don’t produce a continuous spectrum. Rather, the spectrum is composed of light at specific frequencies and is thus known as a line spectrum, where each line of the emission spectrum corresponds to a specific electronic transition. Because each element can have its electrons excited to different distinct energy levels, each one possesses a unique atomic emission spectrum, which can be used as a fingerprint for the element.
One application of atomic emissions spectroscopy is analysis of stars; the light from a star can be resolved into its component wavelengths, which are then matched to the known line spectra of the elements.
The Bohr model of the hydrogen atom explained the atomic emission spectrum of hydrogen, which is the simplest emission spectrum among all the elements. The group of hydrogen emission lines corresponding to transitions from upper levels n > 2 to n = 2 is known as the Balmer series (four wavelengths in this visible region), while the group corresponding to transitions between upper levels n>1 to n=1 is known as the lyman series (higher energy transitions in the UV region).
When the energy of each frequency of light is observed in the emission spectrum of hydrogen was calculated according to Planck’s quantum theory, the values obtained closely matched those expected from energy-level transitions int eh Bohr model. That is, the energy associated with a change in the quantum number from an initial value ni to a final value nf is equal to the energy of Planck’s emitted photon. Thus:
E=hc/wavelength=-Rh[(1/ni^2)-(2/nf^2)]
and the energy of the emitted photon corresponds to the precise difference in energy between the higher-energy initial state and the lower-energy final state.
Atomic Absorption Spectra
When an electron is excited to a higher energy level, it must absorb energy. The energy absorbed as an electron jumps from an orbital of low energy to one of higher energy is characteristic of that transition. This means that the excitation of electrons in a particular element results in energy absorptions at specific wavelengths. Thus, in addition to an emission spectrum, every element possesses a characteristic absorption spectrum. Not surprisingly, the wavelengths of absorption correspond directly to the wavelengths of emission since the energy difference between levels remains unchanged. Absorption spectra can thus be used in the identification of elements present in a gas phase sample.
