Chemistry: Atomic Structure

Question 1

Atom

Answer 1

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.

Question 2

John Dalton’s Atomic Theory

Answer 2

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.

Question 3

Protons

Answer 3

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.

Question 4

Nucleus Weight & Volume

Answer 4

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.

Question 5

Neutrons

Answer 5

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.

Question 6

Electrons

Answer 6

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.

Question 7

Ion

Answer 7

An atom that loses or gains an electron and changes its charge.

Question 8

Atomic Weight

Answer 8

The weight in grams of one mole (mol) of a given element and is expressed in terms of g/mol.

Question 9

Mole

Answer 9

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.

Question 10

Isotopes

Answer 10

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.

Question 11

Quantum Theory

Answer 11

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.

Question 12

The Bohr Model

Answer 12

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.

Question 13

Applications of the Bohr Model

Answer 13

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.

Question 14

Atomic Emission Spectra

Answer 14

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.

Question 15

Atomic Absorption Spectra

Answer 15

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.

Question 16

Orbitals

Answer 16

The most important difference between the Bohr model and modern quantum mechanical models is that Bohr’s assumption that electrons follow a circular orbit at a fixed distance from the nucleus is no longer considered valid. Rather, electrons are described as being in a state of rapid motion within regions of space around the nucleus, called orbitals.

An orbital is a representation of the probability of finding an electron within a given region.

Question 17

Heisenberg Uncertainty Principle

Answer 17

It is impossible to determine, with perfect accuracy, the momentum (defined as mass times velocity) and the position of an electron simultaneously. This means that if the momentum of the electron is being measured accurately, its position will change, and vice versa.

Question 18

Quantum Numbers

Answer 18

Modern atomic theory states that any electron in an atom can be completely described by four quantum numbers. n, l, me, and ms. Furthermore, according to the Pauli exclusion principle, no two electrons in a given atom can possess the same set of four quantum numbers.

Question 19

Energy State

Answer 19

The position and energy of an electron described by its quantum numbers is known as its energy state.

Question 20

Principal Quantum Number (n)

Answer 20

This is the quantum number used in Bohr’s model that can theoretically take on any positive integer value. The larger the integer value of n, the higher the energy level and radius of the electron’s orbit. The max number of electrons in energy level n (electron shell n) is 2n^2.

The difference in energy between adjacent shells decreases as the distance from the nucleus increases, since it is related to the expression (1/n2^2)-(1/n1^2).

For example, the energy difference between the 3rd and 4th shells (n=3 to n=4) is less than that between the 2nd and 3rd shells (n=2 to n=3).

Question 21

Azimuthal Quantum Number (l)

Answer 21

Second quantum number. Also called Angular Momentum Quantum Number. This number tells us the “shape” of the orbitals. This number refers to the subshells or sublevels that occur within each principal energy level.

For any given n, the value of l can be any integer in the range of 0 to n-1.

The four subshells corresponding to l=0, 1, 2, and 3 are known as the s, p, d, and f subshells, respectively.

The max number of electrons that can exist within a subshell is given by the equation 4l + 2. The greater the value of l, the greater the energy of the subshell.

However, the energies of subshells from different principal energy levels may overlap. For example, the 4s subshell will have a lower energy than the 3d subshell because its average distance from the nucleus is smaller.

Question 22

Magnetic Quantum Number (ml)

Answer 22

Third quantum number. Describes the orientation of the orbital in space. An orbital is a specific region within a subshell that may contain no more than two electrons. THe magnetic quantum number specifies the particular orbital within a subshell where an electron is highly likely to be found at a given point in time. The possible values are all integers from l to -l, including 0. Therefore, the s subshell, where there is one possible value of ml (0), will contain one orbital; likewise, the p subshell will contain three orbitals, the d subshell will contain five orbitals, and the f subshell will contain seven orbitals. The shape and energy of each orbital are dependent upon the subshell in which the orbital is found. For example, a p subshell has three possible ml values (-1, 0, +1). The three dumbell shaped orbitals are oriented in space around the nucleus along the x, y, and z axes and are often referred to as px, py, and pz.

Question 23

Spin Quantum Number (ms)

Answer 23

Fourth quantum number. The spin of a particle is its intrinsic angular momentum and is a characteristic of a particle, like its charge. In classical mechanics, an object spinning about its axis has an angular momentum; however, this doesn’t apply to the electron. Classical analogies often are inapplicable in the quantum world.

In any case, the two spin orientations are designated +1/2 and -1/2.

Whenever two electrons are in the same orbital, they must have opposite spins. Electrons in different orbitals with the same ms values are said to have parallel spins.

The quantum numbers for the orbitals in the second principal energy level, with their max number of electrons noted in parentheses, are shown in the following table. Electrons with opposite spins in the same orbital are often referred to as paired.

Question 24

Electron Configuration

Answer 24

For a given atom or ion, the pattern by which subshells are filled and the number of electrons within each principal level and subshell are designated by an electron configuration. In electron configuration notation, the first number denotes the principal energy level, the letter designates the subshell, and the superscript gives the number of electrions in that subshell. For example, Sp^4 indicates that there are four electrons in the second (p) subshell of the second principal energy level.

When writing electron configuration of an atom, it’s necessary to remember the order in which the subshells are filled. Subshells are filled from lowest to highest energy, and each subshell will fill completely before electrons begin to enter the next one. The (n+l) rule is used to rank subshells by increasing energy. This rule states that the lower the values of the first and second quantum numbers, the lower the energy of the subshell. If two subshells possess the name (n+l) value, the subshell with the lower n value has a lower energy and will fill first.

Question 25

Orbital FIlling

Answer 25

To determine which subshells are filled, you must know the number of electrons in the atom. In the case of uncharged atoms, the number of electrons equals the atomic number. If the atom is charged, the number of electrons is equal to the atomic number plus the extra electrons if the atom is negative, or the atomic number minus the electrons if the atom is positive.

Question 26

Hund’s Rule

Answer 26

In subshells that contain more than one orbital, such as the 2p subshell with its three orbitals, the orbitals fill according to Hund’s rule.

The rule states that within a given subshell, orbitals are filled such that there are a max number of half-filled orbitals with parallel spins. Electrons “prefer” empty orbitals to half-filled ones, because a pairing energy must be overcome for two electrons carrying repulsive negative charges to exist in the same orbital.

Question 27

Paramagnetic and Diamagnetic

Answer 27

The presence of paired or unpaired electrons affects the chemical and magnetic properties of an atom or molecule. If the material has unpaired electrons, a magnetic field will align the spins of these electrons and weakly attract the atom. These materials are said to be paramagnetic.

Materials that have no unpaired electrons and are slightly repelled by a magnetic field are said to be diamagnetic.