Atomic Structure and Interatomic Bonding

Transcription

1 Atomic Structure and Interatomic Bonding Kağan Yücetürk Ömer Bektaş Naci Bayhan Ulaş Erdoğan Burkay Alıcı

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3 Intorduction to History of Atom Our understanding of the physical world has grown at an incredible rate in the last 200 years. The key to the advances made in chemistry has been our growing knowledge about atoms. In this section we will look at some of the early historical discoveries that helped us build up a useful 'working model' of the atom.

4 Aristotle and Democritos

5 John Dalton The unit Evidence for Particles showed how the ancient Greeks had ideas about particles and atoms. But it wasn't until the start of the nineteenth century that a theory of atoms became linked to strong experimental evidence. It was then that an English scientist called John Dalton put forward his ideas about atoms. From his experiments and observations, he suggested that atoms were like tiny, hard balls. Each chemical element had its own atoms that differed from others in mass. Dalton believed that atoms were the fundamental building blocks of nature and could not be split. In chemical reactions, the atoms would rearrange themselves and combine with other atoms in new ways. In many ways, Dalton's ideas are still useful today. For example, they help us to understand elements, compounds, and molecules.

6 Dalton made a list of substances that he believed were elements. Decide whether the following substances from his list are elements or compounds. SODA Compound or Element? OXYGEN Compounda or Element? CARBON Compound or Element? GOLD Compound or Element?

7 J.J Thompson At the end of the nineteenth century, a scientist called J.J. Thomson discovered the electron. This is a tiny negatively charged particle that is much, much smaller than any atom. When he discovered the electron, Thomson was experimenting by applying high voltages to gases at low pressure. He noticed an interesting effect. Thomson did experiments on the beams of particles in his tube. They were attracted to a positive charge, so Thomson correctly concluded that they must be negatively charged themselves.

8 4. Other experiments showed that it would take about 2000 electrons to weigh the same as the lightest atom, hydrogen. He called the tiny, negatively charged particles electrons. But where had these tiny particles come from? Since they were so small. Thomson suggested that they could only have come from inside atoms. So Dalton's idea of the indestructible atom had to be revised.

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10 Ernest Rutherford The next development came about 10 years later. Two of Ernest Rutherford's students, Hans Geiger and Ernest Marsden, were doing an experiment at Manchester University with radiation. They were using the dense, positively charged particles (called alpha particles) as 'bullets' to fire at a very thin piece of gold foil. They expected the particles to barge their way straight through the gold atoms unimpeded by the diffuse positive charge spread throughout the atom that Thomson's model described.

11 In 1911, Ernest Rutherford interpreted these results and suggested a new model for the atom. He said that Thomson's model could not be right. The positive charge must be concentrated in a tiny volume at the centre of the atom, otherwise the heavy alpha particles fired at the foil could never be repelled back towards their source. On this model, the electrons orbited around the dense nucleus (centre of the atom).

12 Niels Bohr The next important development came in 1914 when Danish physicist Niels Bohr revised the model again. It had been known for some time that the light given out when atoms were heated always had specific amounts of energy, but no one had been able to explain this. Bohr suggested that the electrons must be orbiting the nucleus in certain fixed energy levels (or shells). The energy must be given out when 'excited' electrons fall from a high energy level to a low one.

13 o How did Bohr change Rutherford's model of the atom? A-) He said that the electrons were concentrated in the centre of the atom. B-) He said that electrons could only occupy specific energy levels as they orbit the nucleus. C-) He said that the electrons could orbit the nucleus in a completely random way.

14 Summary John Dalton introduced a new form of the ancient Greek idea of atoms at the beginning of the nineteenth century. In 1897, J.J. Thomson discovered the electron and suggested the 'plum pudding' model of the atom. In 1911, Rutherford suggested that electrons orbit the atomic nucleus like planets round the Sun. In 1914, Bohr modified Rutherford's model by introducing the idea of energy levels.

15 Modern Atomic Theory Bohr's model of the atom is important because it introduced the concept of the quantum in explaining atomic properties. However, Bohr's model ultimately needed revision becuase it failed to explain the natue of atoms more complicated than hydrogen - the modern atomic theory. Louis de Brogllie introduces the wave/particle duality of matter (1921) Traditional (classical) physics had assumed that particles were particles and waves were waves and thats that. However, de Broglie suggested that particles could sometimes behave as waves and waves could sometimes bahave as particles. He suggested a simple equation that would relate the two: Particles have momentum (p), waves have wavelengths (l) and the two are related by the equation. λ=h/p h=planck's constant = 6.634x10-34 Js p=(mass)x(velocity)

16 Werner Heisenberg elucidated the Uncertainty Principle (1923) Classical physics had always assumed that precise location and velocity of objects was always possible. Heisenberg, however discovered that this was not necessarily the case at the atomic level. In particular, he stated that the act of observation interfered with the location and velocity of small particles such as electrons. This is the case because observation requires light and light has momentum.

17 Although the Schrodinger equation is too difficult to solve for any but the simplest atoms/molecules, we can nevertheless extract some essential conclusion from it: I)Energies are quantized:atoms and molecules cannot have any energy but only certain energies. This means that energies are "quantized". Il ) The orbitals, associated with each energy, determine where the electrons are located.each orbital is determined by a quantum number call the angular momentum quantum number "l". This quantum number can take on the values l=0 (s-orbital), l=1 (p-orbital), l=2 (d-orbital), l=3 (f-orbital) etc.

18 Historical Development of Atomic Theory

19 IMPORTANT CONCEPTS

20 The atomic mass unit (amu): Is often used to express atomic weight. 1 amu is defined as 1/12 of the atomic mass of the most common isotope of carbon atom that has 6 protons (Z=6) and six neutrons (N=6). Atomic number: Atomic number is number of protons. Atomic weight: Weighted average of the atomic masses of the atoms naturally occurring isotopes.

21 Bonding energy: Bond energy (E) is defined as the amount of energy required to break apart a mole of molecules into its component atoms. It is a measure of the strength of a chemical bond. Bond energy is also known as bond enthalpy (H) or simply as bond strength. Covalent bond: Electrons are shared between the molecules, to saturate the valency. Example - H2 Dipole (electric): A dipole is a separation of opposite electrical charges. Electronegative: Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons.

22 Electropositive: Ability of an atom to withdraw a shared pair of electrons in its combined state. Metals are more electropositive. Hydrogen bond: Hydrogen bonding, a special type of secondary bonding, is found to exist between some molecules that have hydrogen as one of the constituents. These bonding mechanisms are now discussed briefly. Ionic bond: Strong Coulomb interaction among negative atoms (have an extra electron each) and positive atoms (lost an electron). Example - Na+Cl-

23 Isotope: Isotopes are atoms that have the same number of protons and electrons but different numbers of neutrons and therefore have different physical properties. Metallic bond: the atoms are ionized, loosing some electrons from the valence band. Those electrons form a electron sea, which binds the charged nuclei in place. Mole: A mole is the amount of matter that has a mass in grams equal to the atomic mass in amu of the atoms.

24 Periodic table: A table of the chemical elements arranged in order of atomic number, usually in rows, so that elements with similar atomic structure (and hence similar chemical properties) appear in vertical columns. Polar molecule: A polar molecule is a molecule containing polar bonds where the sum of all the bond's dipole moments is not zero. Polar bonds form when there is a different between the electronegativity values of the atoms participating in a bond. Polar molecules also form when the spatial arrangement of chemical bonds leads to more positive charge on one side of the molecule than the other. Primary bonding: e- are transferred or shared strong ( KJ/mol or 1-10 ev/atom).

25 Quantum number: A quantum number is a value that is used when describing the energy levels available to atoms and molecules. Secondary bonding: no e- transferred or shared Interaction of atomic/molecular dipoles weak (< 100 KJ/mol or < 1 ev/atom). Secondary bonding: No e- transferred or shared Interaction of atomic/molecular dipoles weak (< 100 KJ/mol or < 1 ev/atom). Valence electron: Electrons that occupy the outermost filled Shell. They are responsible for bonding. van der Waals bond: Van der Waals, or physical bonds are weak in comparison to the primary or chemical ones; bonding energies are typically on the order of only 10 kj/mol ev/atom).

26 Periodic Table of the Elements

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28 chlorine nitrogen gold helium silver oxygen mercury hydrogen sodium niobium neodymium carbon

29 Elements Science has come along way since Aristotle s theory of Air, Water, Fire, and Earth. Scientists have identified 90 naturally occurring elements, and created about 28 others.

30 Elements The elements, alone or in combinations, make up our bodies, our world, our sun, and in fact, the entire universe.

31 The most abundant element in the earth s crust is oxygen.

32 Periodic Table The periodic table organizes the elements in a particular way. A great deal of information about an element can be gathered from its position in the period table. For example, you can predict with reasonably good accuracy the physical and chemical properties of the element. You can also predict what other elements a particular element will react with chemically. Understanding the organization and plan of the periodic table will help you obtain basic information about each of the 118 known elements.

33 Key to the Periodic Table Elements are organized on the table according to their atomic number, usually found near the top of the square. The atomic number refers to how many protons an atom of that element has. For instance, hydrogen has 1 proton, so it s atomic number is 1. The atomic number is unique to that element. No two elements have the same atomic number.

34 What s in a square? Different periodic tables can include various bits of information, but usually: atomic number symbol atomic mass number of valence electrons state of matter at room temperature.

35 Atomic Number This refers to how many protons an atom of that element has. No two elements, have the same number of protons. Bohr Model of Hydrogen Atom Wave Model

36 Atomic Mass Atomic Mass refers to the weight of the atom. It is derived at by adding the number of protons with the number of neutrons. This is a helium atom. Its atomic Hmass is 4 (protons plus neutrons). What is its atomic number?

37 Atomic Mass and Isotopes While most atoms have the same number of protons and neutrons, some don t. Some atoms have more or less neutrons than protons. These are called isotopes. An atomic mass number with a decimal is the total of the number of protons plus the average number of neutrons.

38 Atomic Mass Unit (AMU) The unit of measurement for an atom is an AMU. It stands for atomic mass unit. One AMU is equal to the mass of one proton.

39 Atomic Mass Unit (AMU) There are 6 X or 600,000,000,000,000, 000,000,000 amus in one gram. (Remember that electrons are 2000 times smaller than one amu).

40 Symbols C Carbon Cu Copper All elements have their own unique symbol. It can consist of a single capital letter, or a capital letter and one or two lower case letters.

41 Common Elements and Symbols

42 Valence Electrons The number of valence electrons an atom has may also appear in a square. Valence electrons are the electrons in the outer energy level of an atom. These are the electrons that are transferred or shared when atoms bond together.

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44 Properties of Metals Metals are good conductors of heat and electricity. Metals are shiny. Metals are ductile (can be stretched into thin wires). Metals are malleable (can be pounded into thin sheets). A chemical property of metal is its reaction with water which results in corrosion.

45 Properties of Non-Metals Sulfur Non-metals are poor conductors of heat and electricity. Non-metals are not ductile or malleable. Solid non-metals are brittle and break easily. They are dull. Many non-metals are gases.

46 Properties of Metalloids Silicon Metalloids (metal-like) have properties of both metals and non-metals. They are solids that can be shiny or dull. They conduct heat and electricity better than nonmetals but not as well as metals. They are ductile and malleable.

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49 Families Periods Columns of elements are called groups or families. Elements in each family have similar but not identical properties. For example, lithium (Li), sodium (Na), potassium (K), and other members of family IA are all soft, white, shiny metals. All elements in a family have the same number of valence electrons. Each horizontal row of elements is called a period. The elements in a period are not alike in properties. In fact, the properties change greatly across even given row. The first element in a period is always an extremely active solid. The last element in a period, is always an inactive gas.

50 IONIC BOND COVALENT BOND

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57 METALLIC BONDING In many metals (e.g. Al, Cu, Ag, Pd, etc.), the atoms share their valence electrons but not only or necessarily with neighboring atoms but rather with all the atoms in the matter. This is facilitated by the electrons forming a sea or cloud that engulfs the rest of the atom (i.e. the positively charged nucleus with its remaining non-valence electrons).

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59 METALLIC BONDING Primary bond for metals Good electric and thermal conductivities (because of the free electrons) Bonding may be weak or strong Metallic bonding is non directional

60 SECONDARY BONDING Van Der Waals, or physical bonds and hydrogen bonds are weak in comparison to the primary or chemical ones. Secondary bonding forces arise from atomic or molecular dipoles.

61 An electric dipole exists whenever there is some separation of positive and negative portions of an atom or molecule. The bonding results from the coulombic attraction between the positive end of one dipole and the negative region of an adjacent one

62 Van Der Waals Geckos, harmless tropical lizards, are extremely fascinating and extraordinary animals. They have very sticky feet that cling to virtually any surface. This characteristic makes it possible for them to rapidly run up vertical walls and along the undersides of horizontal surfaces. In fact, a gecko can support its body mass with a single toe

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64 The secret to this remarkable ability is the presence of an extremely large number of microscopically small hairs on each of their toe pads. When these hairs come in contact with a surface, weak forces of attraction (i.e., Wan Der Waals forces) are established between hair molecules and molecules on the surface.

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66 The fact that these hairs are so small and so numerous explains why the gecko grips surfaces so tightly. To release its grip, the gecko simply curls up its toes, and peels the hairs away from the surface.

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68 Hydrogen Bonds A special type of permanent moleculeinduced dipole bonds is called Hydrogen Bonds. It exists in compounds like HF, H2O, and NH3 (see sketch of this bond).

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70 In these atoms the hydrogen proton acts as a strong positive pole that can form bonds with negative poles from other molecules. This is the strongest secondary bond and is responsible for relatively high melting temperature like for water.

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72 Terminology Periodic Table : A table of the chemical elements arranged in order of atomic number, usually in rows, so that elements with similar atomic structure (and hence similar chemical properties) appear in vertical columns. Element : Pure matter is the element that comes from the same type of atoms. Atomic Number : The number of protons in an atom defines the element and is called the atomic number. Mole Mass : The mass of a mole (atom, molecule, or ion) is called the molar mass. Ionic Bond: The chemical bond formed by electron exchange between metal and amethal atoms is ionic bond.

73 Hydrogen Bond: A hydrogen bond is a strong bond between a molecule and an electronegative atom, such as oxygen, nitrogen, or fluorine, which is the result of a partial plus charge of hydrogen attached to another, or electronegative, atom of the same molecular weight. Electron: Electron / Lead is a minor particle with a minor minus (-) charge. Electonegativity: Electronegativity means that each of the atoms that make up a bond; Expresses the binding power of binding electron İsotope: Atoms with the same atom number but different mass numbers are called isotopes

74 Van der Walls Bonds: Van der Waals bonds are bonds that link molecules or groups of atoms with weak electrostatic attraction. Cation: Cation is the name given to ions whose electrons are less than their protons and consequently the effect of shielding electrons decreases. Anion: Positive (+) charged ions are called cations, and negative (-) charged ions are called anions. Polar Bond: Polar bond is the bond type in which the net force of different kinds of atoms comes together and is different from zero. Apolar Bond: Apolar bond is a bondless bond.

75 Unstable Atom : Unstable atom is the atom that tends to receive or give electrons. Stable Atom : Stable atoms are atoms with 8 electrons at the last energy level Mole: A mole is the amount of matter that has a mass in grams equal to the atomic mass in amu of the atoms. Electropositive: Ability of an atom to withdraw a shared pair of electrons in its combined state. Metals are more electropositive Secondary bonding: No e- transferred or shared Interaction of atomic/molecular dipoles weak

76 Valence electron: Electrons that occupy the outermost filled Shell. They are responsible for bonding. Quantum number: A quantum number is a value that is used when describing the energy levels available to atoms and molecules. Secondary bonding: no e- transferred or shared Interaction of atomic/molecular dipoles weak The atomic mass unit (amu): is often used to express atomic weight. Bonding energy: Bond energy (E) is defined as the amount of energy required to break apart a mole of molecules into its component atoms.