Oxygen - Periodic Table (2024)

What is Oxygen

Oxygenis a chemical element with atomic number8which means there are 8 protons and 8 electrons in the atomic structure. Thechemical symbolfor Oxygen isO.

Oxygen is a colourless, odourless reactive gas, the chemical element of atomic number 8 and the life-supporting component of the air. It is a member of the chalcogen group on the periodic table, a highly reactive nonmetal, and an oxidizing agent that readily forms oxides with most elements as well as with other compounds. By mass, oxygen is the third-most abundant element in the universe, after hydrogen and helium.

Oxygen – Properties

Atomic Number of Oxygen

Oxygenis a chemical element with atomic number8which means there are 8 protons and 8 electrons in the atomic structure. Thechemical symbolfor Oxygen isO.

The atomconsist of a small but massivenucleussurrounded by a cloud of rapidly movingelectrons. The nucleus is composed ofprotons andneutrons. Total number of protons in the nucleus is called theatomic numberof the atom and is given thesymbol Z. The total electrical charge of the nucleus is therefore +Ze, where e (elementary charge) equals to1,602 x 10^-19coulombs. In a neutral atom there are as many electrons as protons moving about nucleus. It is the electrons that are responsible for the chemical bavavior of atoms, and which identify the various chemical elements.

See also: Atomic Number – Does it conserve in a nuclear reaction?

Atomic Number and Chemical Properties

Every solid, liquid, gas, and plasma is composed of neutral or ionized atoms. Thechemical properties of the atomare determined by the number of protons, in fact, by number and arrangement of electrons. The configuration of these electrons follows from the principles of quantum mechanics. The number of electrons in each element’s electron shells, particularly the outermost valence shell, is the primary factor in determining its chemical bonding behavior. In the periodic table, the elements are listed in order of increasing atomic number Z.

It is thePauli exclusion principlethat requires the electrons in an atom to occupy different energy levels instead of them all condensing in the ground state. The ordering of the electrons in the ground state of multielectron atoms, starts with the lowest energy state (ground state) and moves progressively from there up the energy scale until each of the atom’s electrons has been assigned a unique set of quantum numbers. This fact has key implications for the building up of the periodic table of elements.

Neutron Number and Mass Number of Oxygen

Mass numbers of typical isotopes of Oxygen are 16; 17; 18.

The total number of neutrons in the nucleus of an atom is called theneutronnumberof the atom and is given thesymbol N. Neutronnumber plusatomic numberequals atomic mass number:N+Z=A. The difference between the neutron number and the atomic number is known as theneutron excess: D = N – Z = A – 2Z.

Neutron number is rarely written explicitly in nuclide symbol notation, but appears as a subscript to the right of the element symbol. Nuclides that have the same neutron number but a different proton number are called isotones. The various species of atoms whose nuclei contain particular numbers of protons and neutrons are callednuclides. Each nuclide is denoted by chemical symbol of the element (this specifies Z) with tha atomic mass number as supescript. Therefore, we cannot determine the neutron number of uranium, for example. We can determine the neutron number of certain isotope. For example, the neutron number of uranium-238 is 238-92=146.

Neutron and Mass Numbers and Nuclear Properties

Properties of atomic nuclei (atomic mass, nuclear cross-sections) are determined by the number of protons and number of neutrons(neutron number). It must be noted, especially nuclear cross-sections may vary by many orders from nuclide with the neutron number N to nuclide with the neutron number N+1. For example, actinides with odd neutron number are usually fissile (fissionable with slow neutrons) while actinides with even neutron number are usually not fissile (but are fissionable with fast neutrons). Heavy nuclei with an even number of protons and an even number of neutrons are (due to Pauli exclusion principle) very stable thanks to the occurrence of ‘paired spin’. On the other hand, nuclei with an odd number of protons and neutrons are mostly unstable.

Neutron and Atomic Numbers and Nuclear Stability

Nuclear stabilityis a concept that helps to identify the stability of an isotope. To identify the stability of an isotope it is needed to find the ratio ofneutronsto protons. To determine the stability of an isotope you can use the ratio neutron/proton (N/Z). Also to help understand this concept there is a chart of the nuclides, known as a Segre chart. This chart shows a plot of the known nuclides as a function of their atomic and neutron numbers. It can be observed from the chart that there aremore neutrons than protonsin nuclides withZ greaterthan about 20 (Calcium). Theseextra neutronsare necessary for stability of the heavier nuclei. The excess neutrons act somewhat like nuclear glue.Only two stable nuclides have fewer neutrons than protons: hydrogen-1 and helium-3.

Atomic nuclei consist of protons and neutrons, which attract each other throughthe nuclear force, while protons repel each other viathe electric forcedue to their positive charge. These two forces compete, leading to various stability of nuclei. There are only certain combinations of neutrons and protons, which formsstable nuclei.

Neutrons stabilize the nucleus, because they attract each other and protons , which helps offset the electrical repulsion between protons. As a result, as the number of protons increases,an increasing ratio of neutrons to protons is neededto form a stable nucleus. If there are too many or too few neutrons for a given number of protons, the resulting nucleus is not stable and it undergoesradioactive decay.Unstable isotopesdecay through various radioactive decay pathways, most commonly alpha decay, beta decay, gamma decay or electron capture. Many other rare types of decay, such as spontaneous fission or neutron emission are known.

Atomic Mass of Oxygen

Atomic mass of Oxygen is 15.9994 u.

The atomic mass is the mass of an atom. The atomic mass or relative isotopic mass refers to the mass of a single particle, and therefore is tied to a certain specific isotope of an element. The atomic mass is carried by the atomic nucleus, which occupies only about 10^-12of the total volume of the atom or less, but it contains all the positive charge and at least 99.95% of the total mass of the atom. Note that, each element may contain more isotopes, therefore this resulting atomic mass is calculated from naturally-occuring isotopes and their abundance.

The size and mass of atoms are so small that the use of normal measuring units, while possible, is often inconvenient. Units of measure have been defined for mass and energy on the atomic scale to make measurements more convenient to express. The unit of measure for mass is the atomic mass unit (amu). One atomic mass unit is equal to 1.66 x 10^-24grams. One unified atomic mass unit isapproximatelythe mass of one nucleon (either a single proton or neutron) and is numerically equivalent to 1 g/mol.

For¹²C the atomic mass is exactly 12u, since the atomic mass unit is defined from it. For other isotopes, the isotopic mass usually differs and is usually within 0.1 u of the mass number. For example,⁶³Cu(29 protons and 34 neutrons) has a mass number of 63 and an isotopic mass in itsnuclear ground state is 62.91367 u.

There are two reasons for the difference between mass number and isotopic mass, known as themass defect:

1. Theneutron isslightlyheavierthan theproton. This increases the mass of nuclei with more neutrons than protons relative to the atomic mass unit scale based on¹²C with equal numbers of protons and neutrons.
2. Thenuclear binding energyvaries between nuclei. A nucleus with greater binding energy has a lower total energy, and therefore alower massaccording toEinstein’s mass-energy equivalencerelationE=mc². For⁶³Cuthe atomic mass is less than 63 so this must be the dominant factor.

Note that, it was found therest mass of an atomic nucleusis measurably smaller thanthe sum of the rest masses of its constituentprotons,neutronsand electrons. Mass was no longer considered unchangeable in the closed system. The difference is a measure of thenuclear binding energywhich holds the nucleus together. According to the Einstein relationship (E=mc²), this binding energy is proportional to thismass differenceand it is known as themass defect.

See also: Atomic Mass Number – Does it conserve in a nuclear reaction?

Atomic Radius of Oxygen

The atomic radius of Oxygen atom is 66pm (covalent radius).

It must be noted, atoms lack a well-defined outer boundary. The atomic radius of a chemical element is a measure of the distance out to which the electron cloud extends from the nucleus. However, this assumes the atom to exhibit a spherical shape, which is only obeyed for atoms in vacuum or free space. Therefore, there are various non-equivalent definitions of atomic radius.

• Van der Waals radius.In principle, Vana der Waals radius is half the minimum distance between the nuclei of two atoms of the element that are not bound to the same molecule.
• Ionic radius. An ionic radius is one-half the distance between the nuclei of two ions in an ionic bond.
• Covalent radius. Covalent radius is the nominal radius of the atoms of an element when covalently bound to other atoms.
• Metallic radius. A metallic radius is one-half the distance between the nuclei of two adjacent atoms in a crystalline structure, when joined to other atoms by metallic bonds.

On the periodic table of the elements, atomic radius tends to increase when moving down columns, but decrease when moving across rows (left to right). Consequently, the smallest atom is helium with a radius of 32 pm, while one of the largest is caesium at 225 pm. The atomic radii decrease across the periodic table because as the atomic number increases, the number of protons increases across the period, but the extra electrons are only added to the same quantum shell. Therefore, the effective nuclear charge towards the outermost electrons increases, drawing the outermost electrons closer. As a result, the electron cloud contracts and the atomic radius decreases.

The volume of an atomis about15 orders of magnitudelargerthan the volume of a nucleus. Foruranium atom, theVan der Waals radiusis about186 pm = 1.86 ×10⁻¹⁰m. The Van der Waals radius, r_w, of an atom is the radius of an imaginary hard sphere representing the distance of closest approach for another atom. Assuming spherical shape, theuraniumatom have volume of about 26.9 ×10⁻³⁰m³. But this “huge” space is occupied primarilyby electrons, because thenucleus occupies only about1721×10⁻⁴⁵m³of space. These electrons together weigh only a fraction (let say 0.05%) of entire atom.

It may seem, that the space and in fact the matter isempty,but it is not. Due to thequantum nature of electrons, the electrons are not point particles, they are smeared out over the whole atom. The classical description cannot be used to describe things on the atomic scale. On the atomic scale, physicists have found that quantum mechanics describes things very well on that scale. Particle locations in quantum mechanics are not at an exact position, they are described by aprobability density function. Therefore the space in an atom (between electrons and an atomic nucleus) is not empty, but it is filled by a probability density function of electrons (usually known as “electron cloud“).

Densityof Oxygen

Densityof Oxygen is1.429g/cm³.

Typical densities of various substances are at atmospheric pressure.

Densityis defined as themass per unit volume. It is anintensive property, which is mathematically defined as mass divided by volume:

ρ = m/V

In words, the density (ρ) of a substance is the total mass (m) of that substance divided by the total volume (V) occupied by that substance. The standard SI unit iskilograms per cubic meter(kg/m³). The Standard English unit ispounds mass per cubic foot(lbm/ft³).

Density – Atomic Mass and Atomic Number Density

Since the density (ρ) of a substance is the total mass (m) of that substance divided by the total volume (V) occupied by that substance, it is obvious, the density of a substance strongly depends on its atomic mass and also on the atomic number density(N; atoms/cm³),

• Atomic Weight. The atomic mass is carried by the atomic nucleus, which occupies only about 10^-12of the total volume of the atom or less, but it contains all the positive charge and at least 99.95% of the total mass of the atom. Therefore it is determined by the mass number (number of protons and neutrons).
• Atomic Number Density. The atomic number density(N; atoms/cm³), which is associated with atomic radii, is the number of atoms of a given type per unit volume (V; cm³) of the material. The atomic number density (N; atoms/cm³) of a pure material havingatomic or molecular weight (M; grams/mol) and thematerial density(⍴; gram/cm³) is easily computed from the following equation using Avogadro’s number (N_A= 6.022×10²³ atoms or molecules per mole):

Sincenucleons(protonsandneutrons) make up most of the mass of ordinary atoms, the density of normal matter tends to be limited by how closely we can pack these nucleons and depends on the internal atomic structure of a substance. Thedensest materialfound on earth is themetal osmium, but its density pales by comparison to the densities of exotic astronomical objects such as whitedwarf starsandneutron stars.

If we include man made elements, the densest so far isHassium.Hassiumis a chemical element with symbolHs and atomic number 108. It is a synthetic element (first synthesised at Hasse in Germany) and radioactive. The most stable known isotope,²⁶⁹Hs, has a half-life of approximately 9.7 seconds. It has an estimated density of 40.7 x 10³kg/m³. The density of Hassium results from itshigh atomic weightand from the significant decrease inionic radiiof the elements in the lanthanide series, known aslanthanide and actinide contraction.

Density – Pressure and Temperature

The density of a material varies with temperature and pressure. This variation is typically small for solids and liquids but much greater for gases. Most materials expand when their temperatures increase. Rising temperatures make the liquid expand in a liquid-in-tube thermometer and bend bimetallic strips. As a result of this expansion, the density of most materials decreases. This effect is caused by a decrease in the atomic number density. This dependence is usually expressed by the coefficient of linear or volume expansion.

Increasing the pressure on an material (especially for liquids or gases) decreases the volume of the object and thus increases its density via the atomic number density. Compressibility (also known as the coefficient of compressibility is a measure of the relative volume change of a fluid or solid as a response to a pressure (or mean stress) change.

See also:What is Density

See also:Densest Materials of the Earth

Electron Configuration and Oxidation States of Oxygen

Electron configuration of Oxygenis [He] 2s22p4.

Possible oxidation states are -2.

Electron Configuration

The periodic table is a tabular display of the chemical elements organized on the basis of their atomic numbers, electron configurations, and chemical properties. The electron configuration is the distribution of electrons of an atom or molecule (or other physical structure) in atomic or molecular orbitals. Knowledge of the electron configuration of different atoms is useful in understanding the structure of the periodic table of elements.

Every solid, liquid, gas, and plasma is composed of neutral or ionized atoms. Thechemical properties of the atomare determined by the number of protons, in fact, by number and arrangement of electrons. The configuration of these electrons follows from the principles of quantum mechanics. The number of electrons in each element’s electron shells, particularly the outermost valence shell, is the primary factor in determining its chemical bonding behavior. In the periodic table, the elements are listed in order of increasing atomic number Z.

It is thePauli exclusion principlethat requires the electrons in an atom to occupy different energy levels instead of them all condensing in the ground state. The ordering of the electrons in the ground state of multielectron atoms, starts with the lowest energy state (ground state) and moves progressively from there up the energy scale until each of the atom’s electrons has been assigned a unique set of quantum numbers. This fact has key implications for the building up of the periodic table of elements.

The first two columns on the left side of the periodic table are where thessubshells are being occupied. Because of this, the first two rows of the periodic table are labeled thes block. Similarly, thep blockare the right-most six columns of the periodic table, thed blockis the middle 10 columns of the periodic table, while thef blockis the 14-column section that is normally depicted as detached from the main body of the periodic table. It could be part of the main body, but then the periodic table would be rather long and cumbersome.

For atoms with many electrons, this notation can become lengthy and so an abbreviated notation is used. The electron configuration can be visualized as the core electrons, equivalent to the noble gas of the preceding period, and the valence electrons (e.g. [Xe] 6s2 for barium).

Oxidation States

Oxidation states are typically represented by integers which may be positive, zero, or negative. Most elements have more than one possible oxidation state. For example, carbon has nine possible integer oxidation states from −4 to +4.

The current IUPAC Gold Book definition of oxidation state is:

“Oxidation state of an atom is the charge of this atom after ionic approximation of its heteronuclear bonds…”

and the term oxidation number is nearly synonymous. An element that is not combined with any other different elements has an oxidation state of 0. Oxidation state 0 occurs for all elements – it is simply the element in its elemental form. An atom of an element in a compound will have a positive oxidation state if it has had electrons removed. Similarly, adding electrons results in a negative oxidation state. We have also distinguish between the possible and common oxidation states of every element. For example, silicon has nine possible integer oxidation states from −4 to +4, but only -4, 0 and +4 are common oxidation states.

Electron Affinity and Electronegativityof Oxygen

Electron affinity of Oxygen is 141 kJ/mol.

Electronegativity of Oxygen is3.44.

Electron Affinity

In chemistry andatomic physics, theelectron affinityof an atom or molecule is defined as:

the change in energy (in kJ/mole) of a neutral atom or molecule (in the gaseous phase) when an electron is added to the atom to form a negative ion.

X + e^–→ X^–+ energy Affinity = – ∆H

In other words, it can be expressed as the neutral atom’slikelihood of gaining an electron. Note that, ionization energies measure the tendency of a neutral atom to resist the loss of electrons. Electron affinities are more difficult to measure than ionization energies.

An atom of Oxygen in the gas phase, for example, gives off energy when it gains an electron to form an ion of Oxygen.

O + e^–→ O^–– ∆H = Affinity = 141 kJ/mol

To use electron affinities properly, it is essential to keep track of sign. When an electron is added to a neutral atom, energy is released. This affinity is known as the first electron affinity and these energies are negative. By convention, the negative sign shows a release of energy. However, more energy is required to add an electron to a negative ion which overwhelms any the release of energy from the electron attachment process. This affinity is known as the second electron affinity and these energies are positive.

Electron affinity can be either positive or negative value. The greater the negative value, the more stable the anion is. Although affinity varies greatly across the periodic table, some patterns emerge. Generally, the elements on the right side of the periodic table will have large negative electron affinity. The electron affinities will become less negative as you go from the top to the bottom of the periodic table. However, nitrogen, oxygen, and fluorine do not follow this trend. Moreover, nonmetals have more positive affinity than metals. Atoms whose anions are more stable than neutral atoms have a greater affinity. Chlorine most strongly attracts extra electrons, while neon most weakly attracts an extra electron.

Affinities of Non metals vs. Affinities of Metals

• Metals:Metals like to lose valence electrons to form cations to have a fully stable shell. The electron affinity of metals is lower than that of nonmetals. Mercury most weakly attracts an extra electron.
• Nonmetals:Generally, nonmetals have more positive electron affinity than metals. Nonmetals like to gain electrons to form anions to have a fully stable electron shell. Chlorine most strongly attracts extra electrons. The electron affinities of the noble gases have not been conclusively measured, so they may or may not have slightly negative values.

Electronegativity

Electronegativity, symbol χ, is a chemical property that describes the tendency of an atom to attract electrons towards this atom. For this purposes, adimensionlessquantity thePauling scale, symbol χ, is the most commonly used.

The electronegativity of Oxygen is: χ = 3.44

In general, an atom’s electronegativity is affected by both its atomic number and the distance at which its valence electrons reside from the charged nucleus. The higher the associated electronegativity number, the more an element or compound attracts electrons towards it. The most electronegative atom, fluorine, is assigned a value of 4.0, and values range down to cesium and francium which are the least electronegative at 0.7. Electronegativity is related with ionization energy and electron affinity. Electrons with low ionization energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted by the positive nucleus on the negative electrons. Therefore the electronegativity is greatest at the top-right of the periodic table and decreases toward the bottom-left.

First Ionization Energy of Oxygen

First Ionization Energyof Oxygen is13.6181 eV.

Ionization energy, also calledionization potential, is the energy necessary toremove an electronfrom the neutral atom.

X + energy → X⁺+ e⁻

where X is any atom or molecule capable of being ionized, X⁺is that atom or molecule with an electron removed (positive ion), and e⁻is the removed electron.

A Oxygen atom, for example, requires the following ionization energy to remove the outermost electron.

O + IE → O⁺+ e⁻IE = 13.6181 eV

The ionization energy associated with removal of the first electron is most commonly used. Thenth ionization energy refers to the amount of energy required to remove an electron from the species with a charge of (n-1).

1st ionization energy

X → X⁺+ e⁻

2nd ionization energy

X⁺→ X²⁺+ e⁻

3rd ionization energy

X²⁺→ X³⁺+ e⁻

Ionization Energy for different Elements

There is an ionization energy for each successive electron removed. The electrons that circle the nucleus move in fairly well-defined orbits. Some of these electrons are more tightly bound in the atom than others. For example, only 7.38 eV is required to remove the outermost electron from a lead atom, while 88,000 eV is required to remove the innermost electron. Helps to understand reactivity of elements (especially metals, which lose electrons).

In general, the ionization energy increases moving up a group and moving left to right across a period. Ionization energy is is related with electronegativity and electron affinity. Electrons with low ionization energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted by the positive nucleus on the negative electrons. Therefore the electronegativity is greatest at the top-right of the periodic table and decreases toward the bottom-left. Moreover:

• Ionization energyis lowest for the alkali metals which have a single electron outside a closed shell.
• Ionization energyincreases across a row on the periodic maximum for the noble gases which have closed shells.

For example, sodium requires only 496 kJ/mol or 5.14 eV/atom to ionize it. On the other hand neon, the noble gas, immediately preceding it in the periodic table, requires 2081 kJ/mol or 21.56 eV/atom.

Oxygen – Melting Point and Boiling Point

Melting pointof Oxygen is-218.4°C.

Boiling pointof Oxygen is-183°C.

Note that, these points are associated with thestandard atmospheric pressure.

Boiling Point

In general, boiling is a phase change of a substance from the liquid to the gas phase. The boiling point of a substance is the temperature at which this phase change (boiling or vaporization) occurs. The temperature at which vaporization (boiling) starts to occur for a given pressure is also known as the saturation temperature and at this conditions a mixture of vapor and liquid can exist together. The liquid can be said to be saturated with thermal energy. Any addition of thermal energy results in a phase transition. At the boiling point the two phases of a substance, liquid and vapor, have identical free energies and therefore are equally likely to exist. Below the boiling point, the liquid is the more stable state of the two, whereas above the gaseous form is preferred. The pressure at which vaporization (boiling) starts to occur for a given temperature is called the saturation pressure. When considered as the temperature of the reverse change from vapor to liquid, it is referred to as the condensation point.

As can be seen, the boiling point of a liquid varies depending upon the surrounding environmental pressure. A liquid in a partial vacuum has a lower boiling point than when that liquid is at atmospheric pressure. A liquid at high pressure has a higher boiling point than when that liquid is at atmospheric pressure. For example, water boils at 100°C (212°F) at sea level, but at 93.4°C (200.1°F) at 1900 metres (6,233 ft) altitude. On the other hand, water boils at 350°C (662°F) at 16.5 MPa (typical pressure of PWRs).

In the periodic table of elements, the element with the lowest boiling point is helium. Both the boiling points of rhenium and tungsten exceed 5000 K at standard pressure. Since it is difficult to measure extreme temperatures precisely without bias, both have been cited in the literature as having the higher boiling point.

Melting Point

In general, melting is a phase change of a substance from the solid to the liquid phase. The melting point of a substance is the temperature at which this phase change occurs. The melting point also defines a condition in which the solid and liquid can exist in equilibrium. Adding a heat will convert the solid into a liquid with no temperature change. At the melting point the two phases of a substance, liquid and vapor, have identical free energies and therefore are equally likely to exist. Below the melting point, the solid is the more stable state of the two, whereas above the liquid form is preferred. The melting point of a substance depends on pressure and is usually specified at standard pressure. When considered as the temperature of the reverse change from liquid to solid, it is referred to as the freezing point or crystallization point.

See also: Melting Point Depression

The first theory explaining mechanism of melting in the bulk was proposed by Lindemann, who used vibration of atoms in the crystal to explain the melting transition. Solids are similar to liquids in that both are condensed states, with particles that are far closer together than those of a gas. The atoms in a solid are tightly bound to each other, either in a regular geometric lattice (crystalline solids, which include metals and ordinary ice) or irregularly (an amorphous solid such as common window glass), and are typically low in energy. The motion of individual atoms, ions, or molecules in a solid is restricted to vibrational motion about a fixed point. As a solid is heated, its particles vibrate more rapidly as the solid absorbs kinetic energy. At some point the amplitude of vibration becomes so large that the atoms start to invade the space of their nearest neighbors and disturb them and the melting process initiates. The melting point is the temperature at which the disruptive vibrations of the particles of the solid overcome the attractive forces operating within the solid.

As with boiling points, the melting point of a solid is dependent on the strength of those attractive forces. For example, sodium chloride (NaCl) is an ionic compound that consists of a multitude of strong ionic bonds. Sodium chloride melts at 801°C. On the other hand, ice (solid H₂O) is a molecular compound whose molecules are held together by hydrogen bonds, which is effectively a strong example of an interaction between two permanent dipoles. Though hydrogen bonds are the strongest of the intermolecular forces, the strength of hydrogen bonds is much less than that of ionic bonds. The melting point of ice is 0 °C.

Covalent bonds often result in the formation of small collections of better-connected atoms called molecules, which in solids and liquids are bound to other molecules by forces that are often much weaker than the covalent bonds that hold the molecules internally together. Such weak intermolecular bonds give organic molecular substances, such as waxes and oils, their soft bulk character, and their low melting points (in liquids, molecules must cease most structured or oriented contact with each other).

Oxygen – Thermal Conductivity

Thermal conductivity of Oxygen is 0.02674 W/(m·K).

The heat transfer characteristics of a solid material are measured by a property called the thermal conductivity, k (or λ), measured in W/m.K. It is a measure of a substance’s ability to transfer heat through a material by conduction. Note that Fourier’s law applies for all matter, regardless of its state (solid, liquid, or gas), therefore, it is also defined for liquids and gases.

The thermal conductivity of most liquids and solids varies with temperature. For vapors, it also depends upon pressure. In general:

Most materials are very nearly hom*ogeneous, therefore we can usually write k = k (T). Similar definitions are associated with thermal conductivities in the y- and z-directions (ky, kz), but for an isotropic material the thermal conductivity is independent of the direction of transfer, kx = ky = kz = k.

Thermal Conductivity of Metals

Thermal Conductivity of Nonmetals

For nonmetallic solids,kis determined primarily byk_ph, which increases as the frequency of interactions between the atoms and the lattice decreases. In fact, lattice thermal conduction is the dominant thermal conduction mechanism in nonmetals, if not the only one. In solids, atoms vibrate about their equilibrium positions (crystal lattice). The vibrations of atoms are not independent of each other, but are rather strongly coupled with neighboring atoms. The regularity of the lattice arrangement has an important effect onk_ph, with crystalline (well-ordered) materials likequartz having a higher thermal conductivity than amorphous materials like glass. At sufficiently high temperatures k_ph∝ 1/T.

Thequantaof the crystal vibrational field are referred to as ‘‘phonons.’’ A phonon is a collective excitation in a periodic, elastic arrangement of atoms or molecules in condensed matter, like solids and some liquids. Phonons play a major role in many of the physical properties of condensed matter, like thermal conductivity and electrical conductivity. In fact, for crystalline, nonmetallic solids such as diamond, k_phcan be quite large, exceeding values of k associated with good conductors, such as aluminum. In particular, diamond has the highest hardness and thermal conductivity (k = 1000 W/m.K) of any bulk material.

Thermal Conductivity of Liquids and Gases

In physics, afluidis a substance that continually deforms (flows) under an applied shear stress.Fluidsare a subset of the phases of matter and includeliquids,gases, plasmas and, to some extent, plastic solids. Because the intermolecular spacing is much larger and the motion of the molecules is more random for the fluid state than for the solid state,thermal energy transportis less effective. Thethermal conductivityof gases and liquids is therefore generally smaller than that of solids. In liquids, the thermal conduction is caused by atomic or molecular diffusion. In gases, the thermal conduction is caused by diffusion of molecules from higher energy level to the lower level.

Thermal Conductivity of Gases

The effect of temperature, pressure, and chemical species on the thermal conductivityof a gas may be explained in terms of thekinetic theory of gases. Air and other gases are generally good insulators, in the absence of convection. Therefore, many insulating materials (e.g.polystyrene) function simply by having a large number ofgas-filled pocketswhichprevent large-scale convection. Alternation of gas pocket and solid material causes that the heat must be transferred through many interfaces causing rapid decrease in heat transfer coefficient.

Thethermal conductivity of gasesis directly proportional to the density of the gas, the mean molecular speed, and especially to themean free pathof molecule. The mean free path also depends on the diameter of the molecule, with larger molecules more likely to experience collisions than small molecules, which is the average distance traveled by an energy carrier (a molecule) before experiencing a collision. Light gases, such ashydrogenandheliumtypically havehigh thermal conductivity. Dense gases such as xenon and dichlorodifluoromethane have low thermal conductivity.

In general, the thermal conductivity of gases increases with increasing temperature.

Thermal Conductivity of Liquids

As was written, in liquids, the thermal conduction is caused by atomic or molecular diffusion, but physical mechanisms for explaining the thermal conductivity of liquids are not well understood. Liquids tend to have better thermal conductivity than gases, and the ability to flow makes a liquid suitable for removing excess heat from mechanical components. The heat can be removed by channeling the liquid through a heat exchanger. The coolants used in nuclear reactors include water or liquid metals, such as sodium or lead.

The thermal conductivity of nonmetallic liquids generally decreases with increasing temperature.

Coefficient of Thermal Expansion of Oxygen

Linear thermal expansion coefficient of Oxygen is — µm/(m·K)

Thermal expansion is generally the tendency of matter to change its dimensions in response to a change in temperature. It is usually expressed as a fractional change in length or volume per unit temperature change. Thermal expansion is common for solids, liquids and for gases. Unlike gases or liquids, solid materials tend to keep their shape when undergoing thermal expansion. A linear expansion coefficient is usually employed in describing the expansion of a solid, while a volume expansion coefficient is more useful for a liquid or a gas.

The linear thermal expansion coefficient is defined as:

where L is a particular length measurement and dL/dT is the rate of change of that linear dimension per unit change in temperature.

The volumetric thermal expansion coefficient is the most basic thermal expansion coefficient, and the most relevant for fluids. In general, substances expand or contract when their temperature changes, with expansion or contraction occurring in all directions.

The volumetric thermal expansion coefficient is defined as:

where L is the volume of the material and dV/dT is the rate of change of that volume per unit change in temperature.

In a solid or liquid, there is a dynamic balance between the cohesive forces holding the atoms or molecules together and the conditions created by temperature. Therefore higher temperatures imply greater distance between atoms. Different materials have different bonding forces and therefore different expansion coefficients. If a crystalline solid is isometric (has the same structural configuration throughout), the expansion will be uniform in all dimensions of the crystal. For these materials, the area and volumetric thermal expansion coefficient are, respectively, approximately twice and three times larger than the linear thermal expansion coefficient (α_V = 3α_L). If it is not isometric, there may be different expansion coefficients for different crystallographic directions, and the crystal will change shape as the temperature changes.

Oxygen –Specific Heat, Latent Heat of Fusion, Latent Heat of Vaporization

Specific heatof Oxygen is0.92 J/g K.

Latent Heat of Fusionof Oxygen is(O2) 0.444 kJ/mol.

Latent Heat ofVaporization of Oxygen is(O2) 6.82 kJ/mol.

Specific Heat

Specific heat, or specific heat capacity,is a property related tointernal energythat is very important in thermodynamics. Theintensive propertiesc_vandc_pare defined for pure, simple compressible substances as partial derivatives of theinternal energyu(T, v)andenthalpyh(T, p), respectively:

where the subscriptsvandpdenote the variables held fixed during differentiation. The propertiesc_vandc_pare referred to asspecific heats(orheat capacities) because under certain special conditions they relate the temperature change of a system to the amount of energy added by heat transfer. Their SI units areJ/kg KorJ/mol K.

Different substancesare affected todifferent magnitudesby theaddition of heat. When a given amount of heat is added to different substances, their temperatures increase by different amounts.

Heat capacity is anextensive propertyof matter, meaning it is proportional to the size of the system.Heat capacity Chas the unit of energy per degree or energy per kelvin. When expressing the same phenomenon as anintensive property, theheat capacityis divided by the amount of substance, mass, or volume, thus the quantity is independent of the size or extent of the sample.

Latent Heat of Vaporization

In general, when a materialchanges phasefrom solid to liquid, or from liquid to gas a certain amount of energy is involved in this change of phase. In case of liquid to gas phase change, this amount of energy is known as theenthalpy of vaporization, (symbol ∆H_vap; unit: J) also known as the(latent) heat of vaporizationor heat of evaporation. As an example, see the figure, which descibes phase transitionsof water.

Latent heat is the amount of heat added to or removed from a substance to produce a change in phase. This energy breaks down the intermolecular attractive forces, and also must provide the energy necessary to expand the gas (thepΔV work). When latent heat is added, no temperature change occurs. The enthalpy of vaporization is a function of the pressure at which that transformation takes place.

The temperature at which vaporization (boiling) starts to occur for a given pressure is also known as the saturation temperature and at this conditions a mixture of vapor and liquid can exist together. The liquid can be said to be saturated with thermal energy. Any addition of thermal energy results in a phase transition. At the boiling point the two phases of a substance, liquid and vapor, have identical free energies and therefore are equally likely to exist. Below the boiling point, the liquid is the more stable state of the two, whereas above the gaseous form is preferred.

Latent Heat ofFusion

In case ofsolid toliquid phase change, the change in enthalpy required to change its state is known as theenthalpy of fusion, (symbol ∆H_fus; unit: J) also known as the(latent)heat of fusion. Latent heat is the amount of heat added to or removed from a substance to produce a change in phase. This energy breaks down the intermolecular attractive forces, and also must provide the energy necessary to expand thesystem (thepΔV work).

The liquid phase has a higher internal energy than the solid phase. This means energy must be supplied to a solid in order to melt it and energy is released from a liquid when it freezes, because the molecules in the liquid experience weaker intermolecular forces and so have a higher potential energy (a kind of bond-dissociation energy for intermolecular forces).

The temperature at which the phase transition occurs is the melting point. The melting point also defines a condition in which the solid and liquid can exist in equilibrium. Adding a heat will convert the solid into a liquid with no temperature change. At the melting point the two phases of a substance, liquid and vapor, have identical free energies and therefore are equally likely to exist. Below the melting point, the solid is the more stable state of the two, whereas above the liquid form is preferred. The melting point of a substance depends on pressure and is usually specified at standard pressure. When considered as the temperature of the reverse change from liquid to solid, it is referred to as the freezing point or crystallization point.

Oxygen – Electrical Resistivity and Electrical Conductivity

Electrical resistivity of Oxygen is — nΩ·m.

Electrical resistivity and its converse, electrical conductivity, is a fundamental property of a material that quantifies how strongly it resists or conducts the flow of electric current. A low resistivity indicates a material that readily allows the flow of electric current. The symbol of resistivity is usually the Greek letter ρ (rho). The SI unit of electrical resistivity is the ohm-metre (Ω⋅m). Note that, electrical resistivity is not the same as electrical resistance. Electrical resistance is expressed in Ohms. While resistivity is a material property, resistance is the property of an object.

Conductors – Semiconductors – Resistors

Substances in which electricity can flow are called conductors. Conductors are made of high-conductivity materials such as metals, in particular copper and aluminium.

Insulators, on the other hand, are made of a wide variety of materials depending on factors such as the desired resistance.

Semiconductors are materials, inorganic or organic, which have the ability to control their conduction depending on chemical structure, temperature, illumination, and presence of dopants. The name semiconductor comes from the fact that these materials have an electrical conductivity between that of a metal, like copper, gold, etc. and an insulator, such as glass. They have an energy gap less than 4eV (about 1eV). In solid-state physics, this energy gap or band gap is an energy range between valence band and conduction band where electron states are forbidden. In contrast to conductors, electrons in a semiconductor must obtain energy (e.g. from ionizing radiation) to cross the band gap and to reach the conduction band.

To understand the difference betweenmetals,semiconductorsandelectrical insulators, we have to define the following terms from solid-state physics:

• Valence Band. In solid-state physics, thevalence bandandconduction bandare the bands closest to theFermi leveland thus determine the electrical conductivity of the solid. In electrical insulators and semiconductors, the valence band is the highest range of electron energies in which electrons are normally present at absolute zero temperature. For example, asiliconatom has fourteen electrons. In the ground state, they are arranged in the electron configuration[Ne]3s²3p². Of these,four are valence electrons, occupying the 3s orbital and two of the 3p orbitals. The distinction between the valence and conduction bands is meaningless in metals, because conduction occurs in one or more partially filled bands that take on the properties of both the valence and conduction bands.
• Conduction Band. In solid-state physics, the valence band and conduction band are the bands closest to the Fermi level and thus determine the electrical conductivity of the solid. In electrical insulators and semiconductors, theconduction bandis the lowest range ofvacant electronic states. On a graph of the electronic band structure of a material, the valence band is located below the Fermi level, while theconduction bandis located above it. In semiconductors,electronsmay reach the conduction band, when they areexcited, for example, byionizing radiation(i.e. they must obtain energy higher thanE_gap). For example, diamond is a wide-band gap semiconductor (E_gap= 5.47 eV) with high potential as an electronic device material in many devices. On the other side, germanium has a small band gap energy (E_gap= 0.67 eV), which requires to operate the detector at cryogenic temperatures. The distinction between the valence and conduction bands is meaningless in metals, because conduction occurs in one or more partially filled bands that take on the properties of both the valence and conduction bands.
• Band Gap. In solid-state physics, theenergy gapor theband gapis an energy range between valence band and conduction band where electron states are forbidden. In contrast to conductors, electrons in a semiconductor must obtain energy (e.g. from ionizing radiation) to cross the band gap and to reach the conduction band.Band gapsare naturally different for different materials. For example, diamond is a wide-band gap semiconductor (E_gap= 5.47 eV) with high potential as an electronic device material in many devices. On the other side, germanium has a small band gap energy (E_gap= 0.67 eV), which requires to operate the detector at cryogenic temperatures.
• Fermi Level. The term “Fermi level” comes fromFermi-Dirac statistics, which describes a distribution of particles over energy states in systems consisting of fermions (electrons) that obey thePauli exclusion principle. Since they cannot exist in identical energy states, Fermi level is the term used to describe the top of the collection ofelectron energy levelsatabsolute zero temperature. TheFermi levelis the surface ofFermi seaat absolute zero where no electrons will have enough energy to rise above the surface. In metals, the Fermi level lies in the hypothetical conduction band giving rise to free conduction electrons. In semiconductors the position of the Fermi level is within the band gap, approximately in the middle of the band gap.
• Electron-hole Pair. In the semiconductor,free charge carriersareelectronsandelectron holes(electron-hole pairs). Electrons and holes are created byexcitation of electronfrom valence band to the conduction band. An electron hole (often simply called a hole) is the lack of an electron at a position where one could exist in anatomor atomic lattice. It is one of the two types of charge carriers that are responsible for creating electric current in semiconducting materials. Since in a normal atom or crystal lattice the negative charge of the electrons is balanced by the positive charge of the atomic nuclei, the absence of an electron leaves a net positive charge at the hole’s location. Positively charged holes can move from atom to atom in semiconducting materials as electrons leave their positions. When an electron meets with a hole, they recombine and these free carriers effectively vanish. The recombination means an electron which has been excited from the valence band to the conduction band falls back to the empty state in the valence band, known as the holes.

Oxygen – Crystal Structure

A possible crystal structure of Oxygen is cubic structure.

In metals, and in many other solids, the atoms are arranged in regular arrays called crystals. A crystal lattice is a repeating pattern of mathematical points that extends throughout space. The forces of chemical bonding causes this repetition. It is this repeated pattern which control properties like strength, ductility, density, conductivity (property of conducting or transmitting heat, electricity, etc.), and shape. There are 14 general types of such patterns known as Bravais lattices.

The three most common basic crystal patterns are:

• Body-centered Cubic. In a body-centered cubic (BCC) arrangement of atoms, the unit cell consists of eight atoms at the corners of a cube and one atom at the body center of the cube. In a body-centered cubic arrangement, a unit cell contains (8 corner atoms × ⅛) + (1 center atom × 1) = 2 atoms. The packing is more efficient (68%) than simple cubic and the structure is a common one for alkali metals and early transition metals. Metals containing BCC structures include ferrite, chromium, vanadium, molybdenum, and tungsten. These metals possess high strength and low ductility.
• Face-centered Cubic.In a face-centered cubic (FCC) arrangement of atoms, the unit cell consists of eight atoms at the corners of a cube and one atom at the center of each of the faces of the cube. In a face-centered cubic arrangement, a unit cell contains (8 corner atoms × ⅛) + (6 face atoms × ½) = 4 atoms. This structure, along with its hexagonal relative (hcp), has the most efficient packing (74%). Metals containing FCC structures include austenite, aluminum, copper, lead, silver, gold, nickel, platinum, and thorium. These metals possess low strength and high ductility.
• Hexagonal Close-packed. In a hexagonal close-packed (HCP) arrangement of atoms, the unit cell consists of three layers of atoms. The top and bottom layers contain six atoms at the corners of a hexagon and one atom at the center of each hexagon. The middle layer contains three atoms nestled between the atoms of the top and bottom layers, hence, the name close-packed. Hexagonal close packed (hcp) is one of the two simple types of atomic packing with the highest density, the other being the face centered cubic (fcc). However, unlike the fcc, it is not a Bravais lattice as there are two nonequivalent sets of lattice points. Metals containing HCP structures include beryllium, magnesium, zinc, cadmium, cobalt, thallium, and zirconium. HCP metals are not as ductile as FCC metals.

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