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Single Slit Diffraction - Four Wavelengths



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Single Slit Diffraction - Four Wavelengths


This figure shows single slit diffraction, but the slit is the length of 4 wavelengths.
To examine this effect better, lets consider a single monochromatic wavelength. This will produce a wavefront that is all in the same phase. Downstream from the slit, the light at any given point is made up of contributions from each of these point sources. The resulting phase differences are caused by the different in path lengths that the contributing portions of the rays traveled from the slit.

The variation in wave intensity can be mathematically modeled. From the center of the slit, the diffracting waves propagate radially. The angle of the minimum intensity (θmin) can be related to wavelength (λ) and the slit's width (d) such that:

dsin⁡θmin=λ.

The intensity (I) of waves at any angle can also be calculated as a relation to slit width, wavelength and intensity of the original waves before passing through the slit:

I(θ)=I0(sin⁡(πx)πx)2,

where x is equal to:

dλsin⁡θ.
6.Atomic physics

6.1 Charged particles

In physics, a charged particle is a particle with an electric charge. It may be an ion, such as a molecule or atom with a surplus or deficit of electrons relative to protons. It can be the electrons and protons themselves, as well as other elementary particles, like positrons. It may also be an atomic nucleus devoid of electrons, such as an alpha particle, ahelium nucleus. Neutrons have no charge, so they are not charged particles unless they are part of a positively charged nucleus. Plasmas are a collection of charged particles, atomic nuclei and separated electrons, but can also be a gas containing a significant proportion of charged particles. Plasma is called the fourth state of matter because its properties are quite different from solidsliquids and gases.

1. Elastic scattering[edit]


It is the process of changing in direction of travelling particle due to the correlation with atom. Conservation of Energy is valid and momentum is preserved.

Rutherford scattering equation

o Coulomb’s Force: electrical repulsive force is acting on α-particle and nucleus.

o Elastic collision: sum of momentum is conserved before and after.

o Rutherford scattering equation

2. Inelastic collision[edit]


Inelastic collision takes most of the part in energy loss process of charged particle inside of the matter.

Stopping power

o Atom’s ionization caused by α-particle is called inelastic collision.

o α-particle loses momentum corresponding to ionization.

o Stopping power: Energy loss due to per unit length particle in matter. It is a power that interrupts the progress of heavy particle in matter.

Range equation R = Range, E = energy of heavy particle, S = stopping power

o Linear Energy Transfer (LET): absolute value of stopping power.

o Specific Ionization: the number of ion pairs produced per unit track length.

o Bragg curve: the graph of specific energy loss along the track of a charged particle. As it loses energy, stopping power value approximately increases along the 1/E then stops. Like this, the peak before its maximum range of stopping power called Bragg peak.

o Range: distance that heavy charged particle progressed until energy is completely lost by repeating ionizing and scattering with atom. In case of the particle that has certain energy, gets certain range. Range is defined as a relation of strength of α-ray and distance.


6.2 Quantum physic

Quantum mechanics (QM; also known as quantum physics or quantum theory), including quantum field theory, is a fundamental branch of physics concerned with processes involving, for example, atoms and photons. Systems such as these which obey quantum mechanics can be in a quantum superposition of different states, unlike in classical physics.

Quantum mechanics gradually arose from Max Planck's solution in 1900 to the black-body radiation problem (reported 1859) and Albert Einstein's 1905 paper which offered a quantum-based theory to explain the photoelectric effect (reported 1887).Early quantum theory was profoundly reconceived in the mid-1920s.

The reconceived theory is formulated in various specially developed mathematical formalisms. In one of them, a mathematical function, the wave function, provides information about the probability amplitude of position, momentum, and other physical properties of a particle.

Important applications of quantum theory include superconducting magnetslight-emitting diodes and the laser, the transistorand semiconductors such as the microprocessormedical and research imaging such as magnetic resonance imaging andelectron microscopy, and explanations for many biological and physical phenomena.

Scientific inquiry into the wave nature of light began in the 17th and 18th centuries, when scientists such as Robert HookeChristiaan Huygens and Leonhard Euler proposed a wave theory of light based on experimental observations.[2] In 1803, Thomas Young, an English polymath, performed the famous double-slit experiment that he later described in a paper titled On the nature of light and colours. This experiment played a major role in the general acceptance of the wave theory of light.

In 1838, Michael Faraday discovered cathode rays. These studies were followed by the 1859 statement of the black-body radiationproblem by Gustav Kirchhoff, the 1877 suggestion by Ludwig Boltzmann that the energy states of a physical system can be discrete, and the 1900 quantum hypothesis of Max Planck.[3] Planck's hypothesis that energy is radiated and absorbed in discrete "quanta" (or energy elements) precisely matched the observed patterns of black-body radiation.

In 1896, Wilhelm Wien empirically determined a distribution law of black-body radiation,[4] known as Wien's law in his honor. Ludwig Boltzmann independently arrived at this result by considerations of Maxwell's equations. However, it was valid only at high frequencies and underestimated the radiance at low frequencies. Later, Planck corrected this model using Boltzmann's statistical interpretation of thermodynamics and proposed what is now called Planck's law, which led to the development of quantum mechanics.

Following Max Planck's solution in 1900 to the black-body radiation problem (reported 1859), Albert Einstein offered a quantum-based theory to explain the photoelectric effect(1905, reported 1887). Around 1900-1910, the atomic theory and the corpuscular theory of light[5] first came to be widely accepted as scientific fact; these latter theories can be viewed as quantum theories of matter and electromagnetic radiation, respectively.

Among the first to study quantum phenomena in nature were Arthur ComptonC. V. Raman, and Pieter Zeeman, each of whom has a quantum effect named after him. Robert Andrews Millikan studied the photoelectric effect experimentally, and Albert Einstein developed a theory for it. At the same time, Niels Bohr developed his theory of the atomic structure, which was later confirmed by the experiments of Henry Moseley. In 1913, Peter Debye extended Niels Bohr's theory of atomic structure, introducing elliptical orbits, a concept also introduced by Arnold Sommerfeld.[6] This phase is known as old quantum theory.

According to Planck, each energy element (E) is proportional to its frequency (ν):

{\displaystyle E=h\nu \ }





Max Planck is considered the father of the quantum theory.

where h is Planck's constant.

Planck cautiously insisted that this was simply an aspect of the processes of absorption and emission of radiation and had nothing to do with the physical reality of the radiation itself.[7] In fact, he considered his quantum hypothesis a mathematical trick to get the right answer rather than a sizable discovery.[8] However, in 1905 Albert Einstein interpreted Planck's quantum hypothesis realistically and used it to explain the photoelectric effect, in which shining light on certain materials can eject electrons from the material. He won the 1921 Nobel Prize in Physics for this work.

Einstein further developed this idea to show that an electromagnetic wave such as light could also be described as a particle (later called the photon), with a discrete quantum of energy that was dependent on its frequency.[9]



The 1927 Solvay Conference in Brussels.

The foundations of quantum mechanics were established during the first half of the 20th century by Max PlanckNiels BohrWerner HeisenbergLouis de BroglieArthur ComptonAlbert Einstein,Erwin SchrödingerMax BornJohn von NeumannPaul DiracEnrico FermiWolfgang PauliMax von LaueFreeman DysonDavid HilbertWilhelm WienSatyendra Nath BoseArnold Somerfield, and others. The Copenhagen interpretation of Niels Bohr became widely accepted.

In the mid-1920s, developments in quantum mechanics led to its becoming the standard formulation for atomic physics. In the summer of 1925, Bohr and Heisenberg published results that closed the old quantum theory. Out of deference to their particle-like behavior in certain processes and measurements, light quanta came to be called photons (1926). From Einstein's simple postulation was born a flurry of debating, theorizing, and testing. Thus, the entire field of quantum physics emerged, leading to its wider acceptance at the Fifth Solvay Conference in 1927.[citation needed]

It was found that subatomic particles and electromagnetic waves are neither simply particle nor wave but have certain properties of each. This originated the concept of wave–particle duality.[citation needed]

By 1930, quantum mechanics had been further unified and formalized by the work of David HilbertPaul Dirac and John von Neumann[10] with greater emphasis on measurement, the statistical nature of our knowledge of reality, and philosophical speculation about the 'observer'. It has since permeated many disciplines including quantum chemistryquantum electronicsquantum optics, and quantum information science. Its speculative modern developments include string theory and quantum gravity theories. It also provides a useful framework for many features of the modern periodic table of elements, and describes the behaviors of atoms during chemical bonding and the flow ofelectrons in computer semiconductors, and therefore plays a crucial role in many modern technologies.

While quantum mechanics was constructed to describe the world of the very small, it is also needed to explain some macroscopic phenomena such as superconductors, andsuperfluids.

The word quantum derives from the Latin, meaning "how great" or "how much".[13] In quantum mechanics, it refers to a discrete unit assigned to certain physical quantities such as the energy of an atom at rest (see Figure 1). The discovery that particles are discrete packets of energy with wave-like properties led to the branch of physics dealing with atomic and subatomic systems which is today called quantum mechanics. It underlies the mathematical framework of many fields of physics and chemistry, including condensed matter physicssolid-state physicsatomic physicsmolecular physicscomputational physicscomputational chemistryquantum chemistryparticle physicsnuclear chemistry, and nuclear physics. Some fundamental aspects of the theory are still actively studied.

Quantum mechanics is essential to understanding the behavior of systems at atomic length scales and smaller. If the physical nature of an atom were solely described byclassical mechanics, electrons would not orbit the nucleus, since orbiting electrons emit radiation (due to circular motion) and would eventually collide with the nucleus due to this loss of energy. This framework was unable to explain the stability of atoms. Instead, electrons remain in an uncertain, non-deterministic, smearedprobabilistic wave–particleorbital about the nucleus, defying the traditional assumptions of classical mechanics and electromagnetism.

Quantum mechanics was initially developed to provide a better explanation and description of the atom, especially the differences in the spectra of light emitted by differentisotopes of the same chemical element, as well as subatomic particles. In short, the quantum-mechanical atomic model has succeeded spectacularly in the realm where classical mechanics and electromagnetism falter.



6.3 Nuclear physic

Nuclear physics is the field of physics that studies atomic nuclei and their constituents and interactions. The most commonly known application of nuclear physics is nuclear power generation, but the research has led to applications in many fields, including nuclear medicine and magnetic resonance imagingnuclear weaponsion implantation in materials engineering, and radiocarbon dating ingeology and archaeology.

The field of particle physics evolved out of nuclear physics and is typically taught in close association with nuclear physics.

The history of nuclear physics as a discipline distinct from atomic physics starts with the discovery of radioactivity by Henri Becquerel in 1896,[1] while investigating phosphorescence in uranium salts.[2] The discovery of the electron by J. J. Thomson[3] a year later was an indication that the atom had internal structure. At the beginning of the 20th century the accepted model of the atom was J. J. Thomson's "plum pudding" model in which the atom was a positively charged ball with smaller negatively charged electrons embedded inside it.

In the years that followed, radioactivity was extensively investigated, notably by the husband and wife team of Pierre Curie and Marie Curie and by Ernest Rutherford and his collaborators. By the turn of the century physicists had also discovered three types of radiation emanating from atoms, which they named alphabeta, and gamma radiation. Experiments by Otto Hahn in 1911 and by James Chadwick in 1914 discovered that the beta decay spectrum was continuous rather than discrete. That is, electrons were ejected from the atom with a continuous range of energies, rather than the discrete amounts of energy that were observed in gamma and alpha decays. This was a problem for nuclear physics at the time, because it seemed to indicate that energy was not conserved in these decays.

The 1903 Nobel Prize in Physics was awarded jointly to Becquerel for his discovery and to Pierre Curie and Marie Curie for their subsequent research into radioactivity. Rutherford was awarded the Nobel Prize in Chemistry in 1908 for his "investigations into the disintegration of the elements and the chemistry of radioactive substances".

In 1905 Albert Einstein formulated the idea of mass–energy equivalence. While the work on radioactivity by Becquerel and Marie Curie predates this, an explanation of the source of the energy of radioactivity would have to wait for the discovery that the nucleus itself was composed of smaller constituents, the nucleons.


7. Astronomy

Cosmology

Cosmology (from the Greek κόσμος, kosmos "world" and -λογία, -logia "study of"), is the study of the origin, evolution, and eventual fate of the universePhysical cosmology is the scholarly and scientific study of the origin, evolution, large-scale structures and dynamics, and ultimate fate of the universe, as well as the scientific laws that govern these realities.[1]

The term cosmology was first used in English in 1656 in Thomas Blount's Glossographia,[2] and in 1731 taken up in Latin byGerman philosopher Christian Wolff, in Cosmologia Generalis.[3]



Religious or mythological cosmology is a body of beliefs based on mythologicalreligious, and esoteric literature and traditions ofcreation and eschatology.

Physical cosmology is studied by scientists, such as astronomers and physicists, as well as philosophers, such asmetaphysiciansphilosophers of physics, and philosophers of space and time. Because of this shared scope with philosophy,theories in physical cosmology may include both scientific and non-scientific propositions, and may depend upon assumptions that can not be tested. Cosmology differs from astronomy in that the former is concerned with the Universe as a whole while the latter deals with individual celestial objects. Modern physical cosmology is dominated by the Big Bang theory, which attempts to bring together observational astronomy and particle physics;[4] more specifically, a standard parameterization of the Big Bang with dark matter and dark energy, known as the Lambda-CDM model.



Theoretical astrophysicist David N. Spergel has described cosmology as a "historical science" because "when we look out in space, we look back in time" due to the finite nature of the speed of light.[5]

Physics and astrophysics have played a central role in shaping the understanding of the universe through scientific observation and experiment. Physical cosmology was shaped through both mathematics and observation in an analysis of the whole universe. The universe is generally understood to have begun with the Big Bang, followed almost instantaneously by cosmic inflation; an expansion of space from which the universe is thought to have emerged 13.799 ± 0.021 billion years ago.[6] Cosmogony studies the origin of the Universe, and cosmography maps the features of the Universe.

In Diderot's Encyclopédie, cosmology is broken down into uranology (the science of the heavens), aerology (the science of the air), geology (the science of the continents), and hydrology (the science of waters).[7]

Metaphysical cosmology has also been described as the placing of man in the universe in relationship to all other entities. This is exemplified by Marcus Aurelius's observation that a man's place in that relationship: "He who does not know what the world is does not know where he is, and he who does not know for what purpose the world exists, does not know who he is, nor what the world is."[8]

7.2 Astrophysics

Astrophysics is the branch of astronomy that employs the principles of physics and chemistry "to ascertain the nature of theheavenly bodies, rather than their positions or motions in space."[1][2] Among the objects studied are the Sun, other starsgalaxies,extrasolar planets, the interstellar medium and the cosmic microwave background.[3][4] Their emissions are examined across all parts of the electromagnetic spectrum, and the properties examined include luminositydensitytemperature, and chemical composition. Because astrophysics is a very broad subject, astrophysicists typically apply many disciplines of physics, including mechanics,electromagnetismstatistical mechanicsthermodynamicsquantum mechanicsrelativitynuclear and particle physics, and atomic and molecular physics.

In practice, modern astronomical research often involves a substantial amount of work in the realms of theoretical and observational physics. Some areas of study for astrophysicists include their attempts to determine: the properties of dark matterdark energy, andblack holes; whether or not time travel is possible, wormholes can form, or the multiverse exists; and the origin and ultimate fate of the universe.[3] Topics also studied by theoretical astrophysicists include: Solar System formation and evolutionstellar dynamics andevolutiongalaxy formation and evolutionmagnetohydrodynamicslarge-scale structure of matter in the universe; origin of cosmic raysgeneral relativity and physical cosmology, including string cosmology and astroparticle physics.

Astrophysics can be studied at the bachelorsmasters, and Ph.D. levels in physics or astronomy departments at many universities.

Although astronomy is as ancient as recorded history itself, it was long separated from the study of terrestrial physics. In the Aristotelianworldview, bodies in the sky appeared to be unchanging spheres whose only motion was uniform motion in a circle, while the earthly world was the realm which underwent growth and decay and in which natural motion was in a straight line and ended when the moving object reached its goal. Consequently, it was held that the celestial region was made of a fundamentally different kind of matter from that found in the terrestrial sphere; either Fire as maintained by Plato, or Aether as maintained by Aristotle.[5][6] During the 17th century, natural philosophers such as Galileo,[7] Descartes,[8] and Newton[9] began to maintain that the celestial and terrestrial regions were made of similar kinds of material and were subject to the same natural laws.[10] Their challenge was that the tools had not yet been invented with which to prove these assertions.[11]

For much of the nineteenth century, astronomical research was focused on the routine work of measuring the positions and computing the motions of astronomical objects.[12][13] A new astronomy, soon to be called astrophysics, began to emerge when William Hyde Wollaston andJoseph von Fraunhofer independently discovered that, when decomposing the light from the Sun, a multitude of dark lines (regions where there was less or no light) were observed in the spectrum.[14] By 1860 the physicist, Gustav Kirchhoff, and the chemist, Robert Bunsen, had demonstrated that the dark lines in the solar spectrum corresponded to bright lines in the spectra of known gases, specific lines corresponding to unique chemical elements.[15] Kirchhoff deduced that the dark lines in the solar spectrum are caused by absorption bychemical elements in the Solar atmosphere.[16] In this way it was proved that the chemical elements found in the Sun and stars were also found on Earth.

Among those who extended the study of solar and stellar spectra was Norman Lockyer, who in 1868 detected bright, as well as dark, lines in solar spectra. Working with the chemist, Edward Frankland, to investigate the spectra of elements at various temperatures and pressures, he could not associate a yellow line in the solar spectrum with any known elements. He thus claimed the line represented a new element, which was called helium, after the GreekHelios, the Sun personified.[17][18]

In 1885, Edward C. Pickering undertook an ambitious program of stellar spectral classification at Harvard College Observatory, in which a team of woman computers, notablyWilliamina FlemingAntonia Maury, and Annie Jump Cannon, classified the spectra recorded on photographic plates. By 1890, a catalog of over 10,000 stars had been prepared that grouped them into thirteen spectral types. Following Pickering's vision, by 1924 Cannon expanded the catalog to nine volumes and over a quarter of a million stars, developing the Harvard Classification Scheme which was accepted for world-wide use in 1922.[19]

In 1895, George Ellery Hale and James E. Keeler, along with a group of ten associate editors from Europe and the United States,[20] established The Astrophysical Journal: An International Review of Spectroscopy and Astronomical Physics.[21] It was intended that the journal would fill the gap between journals in astronomy and physics, providing a venue for publication of articles on astronomical applications of the spectroscope; on laboratory research closely allied to astronomical physics, including wavelength determinations of metallic and gaseous spectra and experiments on radiation and absorption; on theories of the Sun, Moon, planets, comets, meteors, and nebulae; and on instrumentation for telescopes and laboratories.[20]

In 1925 Cecilia Helena Payne (later Cecilia Payne-Gaposchkin) wrote an influential doctoral dissertation at Radcliffe College, in which she applied ionization theory to stellar atmospheres to relate the spectral classes to the temperature of stars.[22] Most significantly, she discovered that hydrogen and helium were the principal components of stars. This discovery was so unexpected that her dissertation readers convinced her to modify the conclusion before publication. However, later research confirmed her discovery.[23]

By the end of the 20th century, further study of stellar and experimental spectra advanced, particularly as a result of the advent of quantum physics.[24]



5. Тexts on chemistry in english for high school
listening

Atoms, Molecules and Ions
All matter, whether living or nonliving, is made of the same tiny building blocks, called atoms. An atom is the smallest basic unit of matter. All atoms have the same basic structure, composed of three smaller particles.

• Protons: A proton is a positively charged particle in an atom’s nucleus. The nucleus is the dense center of an atom

. • Neutrons: A neutron has no electrical charge, has about the same mass as a proton, and is also found in an atom’s nucleus.

• Electrons: An electron is a negatively charged particle found outside the nucleus. Electrons are much smaller than either protons or neutrons. Different types of atoms are called elements, which cannot be broken down by ordinary chemical means. Which element an atom is depends on the number of protons in the atom’s nucleus. For example, all hydrogen atoms have one proton, and all oxygen atoms have 16 protons. Only about 25 different elements are found in organisms. Atoms of different elements can link, or bond, together to form compounds. Atoms form bonds in two ways.

• Ionic bonds: An ion is an atom that has gained or lost one or more electrons. Some atoms form positive ions, which happens when an atom loses electrons. Other atoms form negative ions, which happens when an atom gains electrons. An ionic bond forms through the electrical force between oppositely charged ions. [1]

• Covalent bonds: A covalent bond forms when atoms share one or more pairs of electrons. A molecule is two or more atoms that are held together by covalent bonds.

2.1 Laws of Chemical Combination—The basic laws of chemical combination are the law of conservation of mass, the law of constant composition, and the law of multiple proportions. Each played an important role in Dalton’s development of the atomic theory.

2.2 John Dalton and the Atomic Theory of Matter—Dalton developed his atomic theory to account for the basic laws of chemical combination. The theory centered around the existence of indivisible small particles of matter called atoms and addressed the unique nature of chemical elements, the formation of chemical compounds from atoms of different elements, and the atomic nature of chemical reactions.

2.3 The Divisible Atom—Of the fundamental particles found in atoms, the three of most concern to chemists are protons, neutrons, and electrons. Protons and neutrons make up the nucleus, and their combined number is the mass number, A, of the atom. The number of protons is the atomic number, Z. Electrons are found outside the nucleus, and their number is also equal to the atomic number. The negative charge on an electron is equal in magnitude to the positive charge on a proton. All atoms of an element have the same atomic number, but they may have different numbers of neutrons and hence different mass numbers. Atoms containing the same number of protons (atomic number) but different numbers of neutrons (mass number) are isotopes of an element. Chemical symbols for isotopes are commonly written in the form  with Abeing the mass number and Z the atomic number of the element E.

2.4 Atomic Masses—The atomic mass of an element is a weighted average value calculated from the masses and relative abundances of its naturally occurring isotopes. Theatomic mass unit represents the standard unit of measure of atomic masses; it is exactly  of the mass of a carbon-12 atom.

2.5 The Periodic Table: Elements Organized—The periodic table is an arrangement of the elements by atomic number into rows and columns. This arrangement places elements having similar properties in the same vertical groups (families). This arrangement also allows for the classification of elements as metal, nonmetal, or metalloid.

2.6 Molecules and Molecular Compounds—A chemical formula, the generic term for the various notations used to represent compounds, indicates the relative numbers of atoms of each type in a compound. An empirical formula expresses the simplest atom ratio, and a molecular formula reflects the actual composition of a molecule.Structural formulas describe the arrangement of atoms within molecules. Molecular models are also used to represent the structure and shape of molecules. For example, for acetic acid:



A molecular compound consists of molecules. In a binary molecular compound, the molecules are made up of atoms of two elements. In naming these compounds, the numbers of atoms in the molecules are denoted by prefixes; the names also feature -ide endings.



Examples: NI3 = nitrogen triiodide S2F4 = disulfur tetrafluoride

2.7 Ions and Ionic Compounds—Ions are formed by the loss or gain of electrons by single atoms or groups of atoms. Positive ions are cations, and negative ions areanions. An ionic compound is made up of cations and anions held together by electrostatic attractions. Chemical formulas of ionic compounds are based on an electrically neutral combination of cations and anions called a formula unit, such as NaCl.

The names of some monatomic cations include roman numerals to designate the charge on the ion. The names of monatomic anions are those of the nonmetallic elements, modified to an -ide ending. Polyatomic ions contain more than one atom. For polyatomic anions, the prefixes hypo- and per- and the endings -ite and -ate are commonly used. A hydrate is an ionic compound that includes a fixed number of water molecules associated with the formula unit.

Examples:













2.8 Acids, Bases, and Salts—According to the Arrhenius theory, an acid produces H+ in water and a base produces OH-. A neutralization reaction between an acid and a base forms water and an ionic compound called a salt. Binary acids have hydrogen and a nonmetal as their constituent elements. Their names feature the prefix hydro- and the ending -ic attached to the stem of the name of the nonmetal. Ternary oxoacids have oxygen as an additional constituent element, and their names use prefixes (hypo- andper-) and endings (-ous and -ic) to indicate the number of O atoms per molecule.

Examples:



2.9 Organic Compounds—Organic compounds are based on the element carbon. Hydrocarbons contain only hydrogen and carbon. Alkanes have carbon atoms joined together by single bonds into chains or rings, with hydrogen atoms attached to the carbon atoms. Alkanes with four or more carbon atoms can exist as isomers, which are molecules that have the same molecular formula but different structures and properties.



Molecules and Ions

Although atoms are the smallest unique unit of a particular element, in nature only the noble gases can be found as isolated atoms. Most matter is in the form of ions, or compounds.

Molecules and chemical formulas

A molecule is comprised of two or more chemically bonded atoms. The atoms may be of the same type of element, or they may be different.

Many elements are found in nature in molecular form - two or more atoms (of the same type of element) are bonded together. Oxygen, for example, is most commonly found in its molecular form "O2" (two oxygen atoms chemically bonded together).

Oxygen can also exist in another molecular form where three atoms are chemically bonded. O3 is also known as ozone. Although O2 and O3 are both compounds of oxygen, they are quite different in their chemical and physical properties. There are seven elements which commonly occur as diatomic molecules. These include H, N, O, F, Cl, Br, I.

An example of a commonly occurring compound that is composed of two different types of atoms is pure water, or "H2O". The chemical formula for water illustrates the method of describing such compounds in atomic terms: there are two atoms of hydrogen and one atom of oxygen (the "1" subscript is omitted) in the compound known as "water". There is another compound of Hydrogen and Oxygen with the chemical formula H2O2 , also known as hydrogen peroxide. Again, although both compounds are composed of the same types of atoms, they are chemically quite different: hydrogen peroxide is quite reactive and has been used as a rocket fuel (it powered Evil Kenievel part way over the Snake River canyon).

Most molecular compounds (i.e. involving chemical bonds) contain only non-metallic elements.

Molecular, Empirical, and Structural Formulas

Empirical vs. Molecular formulas


  • Molecular formulas refer to the actual number of the different atoms which comprise a single molecule of a compound.

  • Empirical formulas refer to the smallest whole number ratios of atoms in a particular compound.

Compound

Molecular Formula

Empirical Formula

Water

H2O

H2O

Hydrogen Peroxide

H2O2

HO

Ethylene

C2H4

CH2

Ethane

C2H6

CH3

Molecular formulas provide more information, however, sometimes a substance is actually a collection of molecules with different sizes but the same empirical formula. For example, carbon is commonly found as a collection of three dimensional structures (carbon chemically bonded to carbon). In this form, it is most easily represented simply by the empirical formula "C" (the elemental name).

Structural formulas

Sometimes the molecular formulas are drawn out as structural formulas to give some idea of the actual chemical bonds which unite the atoms.



Structural formulas give an idea about the connections between atoms, but they don't necessarily give information about the actual geometry of such bonds.



Ions

The nucleus of an atom (containing protons and neutrons) remains unchanged after ordinary chemical reactions, but atoms can readily gain or lose electrons.

If electrons are lost or gained by a neutral atom, then the result is that a charged particle is formed - called an ion.

For example, Sodium (Na) has 11 protons and 11 electrons. However, it can easily lose 1 electron. The resulting cation has 11 protons and 10 electrons, for an overall net charge of 1+ (the units are electron charge). The ionic state of an atom or compound is represented by a superscript to the right of the chemical formula: Na+, Mg2+ (note the in the case of 1+, or 1-, the '1'is omitted). In contrast to the Na atom, the Chlorine atom (Cl) easily gains 1 electron to yield the chloride ion Cl- (i.e. 17 protons and 18 electrons).



In general, metal atoms tend to lose electrons, and nonmetal atoms tend to gain electrons.

Na+ and Cl- are simple ions, in contrast to polyatomic ions such as NO3- (nitrate ion) and SO42- (sulfate ion). These are compounds made up of chemically bonded atoms, but have a net positive or negative charge.

The chemical properties of an ion are greatly different from those of the atom from which it was derived.

Predicting ionic charges

Many atoms gain or lose electrons such that they end up with the same number of electrons as the noble gas closest to them in the periodic table.

The noble gasses are generally chemically non-reactive, they would appear to have a stable arrangement of electrons.



Other elements must gain or lose electrons, to end up with the same arrangement of electrons as the noble gases, in order to achieve the same kind of electron stability.
Example: Nitrogen

Nitrogen has an atomic number of 7; the neutral Nitrogen atom has 7 protons and 7 electrons. If Nitrogen gained three electrons it would have 10 electrons, like the Noble gas Neon (10 protons, 10 electrons). However, unlike Neon, the resulting Nitrogen ion would have a net charge of N3- (7 protons, 10 electrons).



The location of the elements on the Periodic table can help in predicting the expected charge of ionic forms of the elements.

This is mainly true for the elements on either side of the chart.


Ionic compounds

Ions form when one or more electrons transfer from one neutral atom to another. For example, when elemental sodium is allowed to react with elemental chlorine an electron transfers from a neutral sodium to a neutral chlorine. The result is a sodium ion (Na+) and a chlorine ion, chloride (Cl-):



The oppositely charged ions attract one another and bind together to form NaCl (sodium chloride) an ionic compound.



An ionic compound contains positively and negatively charged ions

It should be pointed out that the Na+ and Cl- ions are not chemically bonded together. Whereas atoms in molecular compounds, such as H2O, are chemically bonded.



Ionic compounds are generally combinations of metals and non-metals.

Molecular compounds are general combinations of non-metals only.

Pure ionic compounds typically have their atoms in an organized three dimensional arrangement (a crystal). Therefore, we cannot describe them using molecular formulas. We can describe them usingempirical formulas.



If we know the charges of the ions comprising an ionic compound, then we can determine the empirical formula. The key is knowing that ionic compounds are always electrically neutral overall.



Therefore, the concentration of ions in an ionic compound are such that the overall charge is neutral.

In the NaCl example, there will be one positively charged Na+ ion for each negatively charged Cl- ion.

What about the ionic compound involving Barium ion (Ba2+) and the Chlorine ion (Cl-)?

1 (Ba2+) + 2 (Cl-) = neutral charge

Resulting empirical formula: BaCl2[2]

Atomic structure

The Nucleus

The atomic nucleus is the central area of the atom. It is composed of two kinds of subatomic particles: protons and neutrons.




Diagram showing the atomic structure with the protons and neutrons held together to form the dense area of the nucleus.

Atoms are the building blocks of all matter. Everything you can see, feel and touch is all made of atoms. There are even things you cannot see, feel, hear or touch that are also made of atoms. Basically, everything is made up of atoms.

In 1909, Ernest Rutherford led Hans Geiger and Ernest Marsden through what is known as the Gold Foil Experiments. During the experiments they would shoot particles through extremely thin sheets of gold foil. In 1911, Rutherford came to the conclusion that the atom had a dense nucleus because most of the particles shot straight through, but some of the particles were deflected due to the dense nucleus of the gold atoms. This theory would eliminate the idea that the atom was structured more like plum pudding. The plum pudding model was the leading model of atomic structure until Rutherford's findings.





Atomic Numbers

The atomic nucleus is in the center of the atom. The number of protons and neutrons in the atom define what type of atom or element it is. An element is a bunch of atoms that all have the same type of atomic structure. For instance, hydrogen is an element. Every hydrogen atom is made up of 1 proton, 0 neutrons, and 1 electron.

The composition of the atomic nucleus gives us lots of information about the element it represents. The number of protons inside the nucleus gives us theatomic number. The protons have a positive (+) charge. In order for the atom to have a neutral charge, the electrons (-) need to balance it out with their negative charge. Therefore, in a neutral atomthere are just as many protons as electrons. So, if you know the atomic number and know the charge of the atom then the number of electrons is easy to find. For instance, hydrogen has 1 proton, 1+, so in order for the hydrogen atom to be neutral it must have 1- charge. Therefore, hydrogen has 1 electron.

Where do the neutrons fit in all of this? Well,neutrons are neutral. To keep it all straight I use the first letters: Neutrons are Neutral, and Protons arePositive. I then remember Electrons through the process of Elimination.

Although the neutrons do not give the atom any charge, they still hold their own weight in the importance of the atomic structure. The neutron is the largest of the subatomic particles. When put the neutrons and protons together we get the atomic mass. The electrons are so small that their mass only counts for .01%. The electrons are not inside of the nucleus; instead they are flying around like crazy on the outside of the nucleus.

Since the atomic number gives us the number of protons in an atom and the atomic mass gives us the number of protons and neutrons, we can find the number of neutrons by subtracting the atomic number from the atomic mass.

Atomic mass - atomic number = number of neutrons.

The atomic number of an atom gives each element its identity. You can find out which element it is by its atomic number and reverse the process to find out what the atomic number is if you know which element you are working with.

Let's run through all of the numbers with an element, oxygen.

Oxygen
Atomic Number: 8

Atomic Mass: 16[3]
The ability of atoms to lose or to gain electrons.

Next, let's review two atomic properties important to bonding that are related to the position of the element on the periodic table. They are the tendency or ability of atoms tolose electrons and the tendency or ability to gain electrons.

First, let's consider the ability to lose electrons. This is related to ionization energy, which you studied in a previous lesson. The ionization energy, of course, is the amount of energy that it takes to remove an electron from an atom. You have learned that the ionization energies are lowest for the elements down and on the left hand side of the periodic table and increase as you go up and all the way across to the right including the inert gases.

The ionization energy measures how hard it is to lose or remove an electron. High ionization energy means that it is hard to lose electrons. Low ionization energy means that it easy to lose electrons. The elements on the left side lose their electrons fairly easily and the elements on the right side of the periodic table do not lose their electrons very easily. Taking vertical position on the table into account, the elements that are lower on the table lose electrons more easily and the elements that are higher have a harder time losing electrons. Thus the overall trend is from most easily losing electrons on the lower left to least easily losing electrons on the upper right. Keep that trend in mind.



Ability to Lose Electrons

 

 

 

 

 

 

 

 

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The ability to gain electrons is also related to the position on the periodic table. You should recall that as you go from left to right on the periodic table, the attraction for electrons increases and the ability to gain electrons increases. This is true all the way across the periodic table except/em> for the inert gases. There is an abrupt drop in the ability to gain electrons when we get to the inert gases. This is because their energy level is full and any additional electrons will have to start a new energy level.

Ability to Gain Electrons

 

 

 

 

 

 

 

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1.4 Classification of Chemical Bond Types

Types of Chemical Bonds

A group of atoms bonded to one another form a molecule.

If the molecule has more than one type of element present it is a compound.


Different types of bonds hold molecules and compounds together.


These 2 types of bonds are 1. ionic and 2. covalent.

Atoms start off with the same number of positive protons and negative electrons. This way the opposite charges cancel each other out.

Charged atoms, or ions, can form when atoms lose or gain electrons- remember that atoms will gain or lose electrons in order to have a full outer shell.

If an atom starts off with 9 protons and 9 electrons, the positive and negative charges are balanced out. However, this atom only has 7 electrons in its outer shell, so it wants 1 more electron to have 8 and be happy. But when the atom gains an extra negative electron, it now has 10 negative electrons and 9 positive protons. Therefore its overall charge is -1. If an atom has one electron in its outer shell, it will usually give that electron away and use the next lower shell as a "full" outer shell. When it gives a negative electron away, it becomes a positively charged ion.

Positive and negative ions are attracted to one another and bond together in ionic bonds.

A salt is a dry solid composed of atoms connected by ionic bonds Ex- table salt.

A covalent bond results when two atoms share electrons, thereby completing their valence shells



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