chem chemical bonding

chem chemical bonding - Early History of Atomic Theories a)...

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Unformatted text preview: Early History of Atomic Theories a) Dalton Atomic theory Key Experimental work Law of definite composition: elements combine in a characteristic mass ratio Theoretical explanation Each atom has a particular combining capacity Atomic Theory Matter is composed of indestructible, indivisible atoms, which are identical for one element, but different from other elements Law of multiple proportions: Some atoms have more than one there may be more than one mass combining capacity ratio Law of conservation of mass: total mass remains Atoms are neither created nor destroyed constant in a chemical reaction In Dalton's atomic model, an atom is a solid sphere, similar to a billiard ball. b) Thomson Atomic Theory Matter is composed of atoms that contain electrons (negative particles) embedded in a positive material. The kind of element is characterized by the number of electrons in the atom. a) In Thomson's atomic model, the atom is a positive sphere with embedded electrons. b) This model can be compared to a raisin bun, in which the raisins represent the negative electrons and the bun represents the region of positive charge. c) Rutherford Atomic Model Key Experimental work Rutherford: a few positive alpha particles are deflected at large angles when fired at a gold foil Theoretical explanation Atomic Theory The positive charge in the atom An atom is composed of a very must be concentrated in a very tiny nucleus, which contains small volume of the atom positive charges and most of the mass of the atom. Very small negative electrons occupy most of the volume of the atom. A very strong nuclear force holds the positive charges within the nucleus. Most of the atom is empty space. Most materials are very stable and do not fly apart(break down) Rutherford: Most alpha particles pass straight through gold foil Chadwick's Theory Atoms are composed of protons, neutrons and electrons. Atoms of same elements have the same number of protons and electrons but may have a varying number of neutrons (isotopes of the elements) Origins of Quantum Mechanics 1)Plank's Quantum Hypothesis energies of the oscillationg atoms in the heated solid were multiples of a small quantity of energy(energy isnt constinous) Einstein suggested that light emitted by a hot solid is also quantizedit comes in bursts, not continuous stream of energy one little burst or packet of energy is known as a quantum of energy Example: The smallest quantity of money is pennies and any quantity of money can be expressed in terms of pennies eg $1.00 is 100 pennies or $1.00 can be made up from 2 quarters,3 dimes,3 nickels, and 5 pennies. Think of coins as representing the energy of the light quantathe penny is infrared, nickel is red, the dime is blue and the quarter is ultraviolet radiation.Heat without colour would be emitted as pennies only,red hot radiation would include nickels, white hot radiation would add dimes and blue hot would include more dimes and more quarters. as the temperature is increased, the proportion of each larger quantun becomes greater 2)Photoelectric Effect Each photon of light has a different energy, the energy is increasing from left to right. According to Einstein photons that fall on a metal surface have their energy transferred to the electrons of the metal. If the photon has more energy than the electrons need to leave the surface (different metals have different requirements) the electrons are ejected immediately. Bohr Atomic Theory 1)Problem with Rutherford model electrons could orbit the nucleus at any distance electrons circle round the nucleus, they are constantly changing their direction. According to classical electrodynamics (which deals with the motion of electrons), such electrons which either constantly change their direction or their velocity or both should continuously emit radiation. While doing so, they should lose energy, and thus spiral into the nucleus. This means every atom is unstable. Bohr's Model Electrons do not radiate energy as the orbit the nucleus. Each orbit corresponds to a state of constant energy (called a stationary state) Electrons can change their energy only by undergoing a transition from one stationary state to another. Quantum Numbers 1. Principal Quantum Number (n) (n) is the energy level of an electron n=1,2,3,4... 2. Secondary Quantum Number (l) At each energy level (n) there is a subenergy level (l) when n=0, l = 0 n=1, l = 0 or 1 0 1 2 n=2, l = 0 or 1 or 2 p d l can also be referred to with a letter s 3 f Quantum Numbers (Cont'd) 3. Magnetic Quantum Number ( ) = l to + l when l = 0 (s), = 0 1 orbital at the (s) level l = 1 (p), = 1, 0, +1 3 orbitals at the (p) level l = 2 (d), = 2, 1, 0, +1, +2 5 orbitals at the (d) level 4. Spin Quantum Number ( ) Every electron has either a + (clockwise) or ( counter clockwise) spin The 1st electron has a positive spin The two electrons in a sub shell have opposite spins Atomic Structure and the Periodic Table Creating Energy Level Diagrams Three Main Principles Must Be Followed Pauli exclusion Principle According to Pauli's principle an orbital can accommodate a maximum of 2 electrons & the must have opposite spins because no two electrons in an atom can have the same 4 quantum numbers Aufbau Principle In the ground state of an atom, an electron enters the lowest energy level first, and subsequent electrons are placed in the order of increasing energy levels Hund's rule Electron pairing cannot take place until one electron occupies each of the several orbitals at the same energy level Exceptions Gold Copper Iron Silver Atomic Structure and the Periodic Table (Cont'd) Creating Energy Level Diagrams for Cations and Anions Anions To create diagrams for anions, add electrons according to the aufbau principle Cations To create diagrams for cations, remove electrons from the orbital with the highest principal quantum number first Electron Configuration Procedure Determine the position of the Ex. Scandium: Atomic #21 element and the number of 1s 2s 2p 3s 3p 4s 3d electrons in the atom or ion Wave Mechanics and Orbitals Quantum theory was able to provide improvements in the understanding of electron energy states but many questions such as , "What is the electron doing in the atom and where does it spend its time?", were still a problem. Louis de Broglie proposed the idea of wave mechanics (or quantum mechanics). He stated that a particle could act as both a particle, and a wave. A difficulty with wave mechanics was that a particle as a wave was hard to understand and contrary to our experience. A way to solve this was to continue to picture the electron as a particle travelling in waves so that its location can only be specified as a statistical probability Heisenberg's uncertainty principle: It is impossible to simultaneously know exact speed and position of a particle. Wave Mechanics and Orbitals (Cont'd) Quantum mechanics does not actually provide a description of how an electron gets from one point to another. However the wave equations can be used to provide a 3D probability of an electron in an orbital based on its quantum numbers. Electron probability density : A mathematical or graphical representation of the chance of finding an electron in a given space. The electrons in s orbitals are shown in a spherical shape, p orbitals are shown as dumbbells and d orbitals are shown with multiple petals. Applications of Quantum Mechanics Quantum mechanics has had enormous success in explaining many of the features of our world. It has been used in physics and chemistry and new discoveries are being made because of it. Its applications can be used in many different aspects of science as well as our everyday lives. Some examples are MRIs, the electron microscope, the microchip, and the laser. Lasers are used everywhere in society. They are used in bar code scanners, printers, and in medical offices. The invention of MRIs has been very beneficial for those working in medicine because MRIs are very powerful and can be used for things that an Xray would be unable to accomplish. For example, an MRI can differentiate between similar types of tissues and that would be very beneficial for doctors who diagnose cancer or are dealing with sensitive tissues around the brain or spinal cord. Chapter 3 Review According to Einstein photons that fall on a metal surface have their energy transferred to the electrons of the metal. If the photon has more energy than the electrons need to leave the surface (different metals have different requirements) the electrons are ejected immediately. Electrons do not radiate energy as the orbit the nucleus. Each orbit corresponds to a state of constant energy (called a stationary state) Electrons can change their energy only by undergoing a transition from one stationary state to another. Quantum Numbers : (n) is the energy level, (l) is the energy subshell, ( ) is the orientation of the subshell, and ( ) is the spin of the electron To draw diagrams follow the three main principles Pauliexclusion: Each orbital has a max. of 2 electrons with opposite spins Aufbau: Electrons enter the lowest energy level first, and subsequent electrons are placed in the order of increasing energy levels Hund's Rule: Electron pairing cannot take place until one electron occupies each of the several orbitals at the same energy level Wave Mechanics and Orbitals - A particle can act as both a particle, and a wave. - Heisenberg's uncertainty principle: It is impossible to simultaneously know exact speed and position of a particle. - Electron probability density: A mathematical or graphical representation of the chance of finding an electron in a given space. - The electrons in s orbitals are shown in a spherical shape, p orbitals are shown as dumbbells and d orbitals are shown with a multiple petal shape. Applications of Quantum Mechanics - Quantum mechanics has had enormous success in explaining many of the features of our world. - Its applications can be used in many different aspects of science as well as our everyday lives. Some examples are MRIs, the electron microscope, the microchip, and the laser. Lewis Theory of Bonding Timeline of theories: 1852: 1858: (we 1874: passes 1904: Edward Frankland stated that each element has a fixed valance that determines its bonding capacity. Friedrich extended the idea and illustrated a bond as a dash between bonding atoms now call it a structural diagram). Jacobus van't Hoff and Joseph Le Bel extented these structures to 3D. The had also revised the theory in order to explain why certain substances can change light as it through. Richard Abegg, a German chemist, suggested that the stability of noble gases was because of their full outermost shell. In other words, this tell us that the atoms would either lose or gain electrons to become stable by achieving a full outermost shell. As they lose and gain electrons, they become ions and these ions are held together by electrostatic charge, resulting in ionic bonding. 1916: Gilbert Lewis, an American chemist, used many known chemical formulas, concept of valance, the octet rule & the electronshell model to explain chemical bonding. His work gave a clear understanding of chemical bonding, especially covalent bonding. Key ideas of Lewis Theory: Atoms and ions are stable when they have noble gas like electron structure Electrons are stable when they are paired Atoms form bonds to achieve stable octet Ionic bond: exchange of electrons between metal and nonmetal atoms Covalent bonds: sharing of electrons between nonmetal atoms Rules for drawing Lewis Structure: Example: SO3 STEP 1: Total number of valence electrons subtracted] = 6 + (3)(6) = 24 [for each negative charge, 1 extra electron is added and for a positive charge, an electron is STEP 2: To achieve stable octet STEP 3: Difference between 1 and 2 STEP 4: Number of covalent bonds = 8 + (3)(8) = 32 24 = 8 / 2 = 32 = 8 = 4 Exceptions to the Lewis structure is often, especially beyond the second period. For instance, the atoms of hydrogen through boron do not achieve an octet of electrons when they form molecules. Also, nitrogen has 1 paired electron and 3 unpaired in its Lewis symbol. This tells us that nitrogen has a bonding capacity of 3. Extending the Lewis Theory of Bonding Polyatomic ions could not be easily explained by Lewis Theory Nevil Sidgwik showed that the Lewis structures can work if it is not required that each atom should contribute an electron to the covalent bond Coordinate Covalent bond: both the electrons in a covalent bond come from the same atom. He also acknowledged that an octet is desirable, but not necessary in all molecules and polyatomic ions. The Nature of the Chemical Bond While an atom has only 1 nucleus to consider, a molecule has more than 1 nuclei and so the application of quantum mechanics get harder. To simplify our work, we can start with individual atoms and their orbitals, and then build the molecule using these orbitals. Valence bond theory: a covalent bond is formed when 2 orbitals overlap and produce a new combined orbital containing 2 electrons of opposite spin. This is an exothermic reaction: energy is lost as covalent bonds are formed Orbital overlap for H2 Orbital overlap for H2O. Although, here it looks like the angle is 90, it is actually 105. Other factors for this result will be discussed later on. Hybrid Orbitals: Hybridization: a combination of orbitals that create a new set of orbitals that take part in covalent bonding Hybrid orbital: an atomic orbital formed by a combination of at least 2 different orbitals These orbitals are spontaneously formed when bonding occurs and do not exist in an isolated atom. Example: In case of Carbon [1s2, 2s2, 2p2] a 2s electron jumps to the empty p orbital, when activation energy is provided, to form sp3 hybrid orbital. Now, there are 4 unpaired electrons resulting in the valency of 4 for carbon. Forms of Hybridization Double Bonds: Sigma bond: Pi Bond: orbitals overlap endtoend orbitals overlap sidebyside Example: C2H4 The new idea is that from the hybridization is said to be partial. After an electron jumps to the 2p shell, instead of counting all four valence electrons, we take the first three to form sp2 and the last one is just a normal 2pz shell with one electron. The 2 sp2 orbitals are used to form sigma bonds with the four hydrogens and carbon. The 2pz orbitals form the pi bond, which is a region of electron density above and below the sigma bond, joining the 2 carbon atoms. Also, this additional pair of electrons provide greater attraction toward the nuclei making a double bond shorter. Triple Bonds: Example: C2H2 Similar to the partial hybridization in double bonds, after the 2s2 electron jumps to the 2pz shell, sp hybrid is formed leaving 2py and 2pz as normal shells with one electron each. The two sp orbitals form the linear sigma bonds with hydrogen and the other p shells form two pi bonds above and below the sigma bond. VSEPR Theory A simpler theory than the valence bond theory created by Australian Ronald Nyholm and Englishman Ron Gillespie in 1957. VSEPR: valenceshellelectronpairrepulsion theory; based on the electrical repulsion of bonded and unbonded electron pairs in a molecule or polyatomic ion. Using the VSEPR Theory: First, we have to draw the Lewis structure and then consider the arrangement. The main idea is that all pairs of electrons repel each other and try to get as far away as possible. Repulsion: Lone pairLone pair > Lone pairBond Pair > Bond pairBond pair Molecular Shapes These bond pairs and lone pair can be used to create a general formula in order to predict the shape. In any general formula (ex: AX2E2), X represents the bond pairs and E represents the lone pairs. Polar Molecules A covalent bond is formed when a pair of electrons is shared between two atoms There are two types of covalent bonds: Polar covalent bonds Unequal sharing of electrons Form polar molecules Have partially different charges at opposite ends of the molecule More polar = more ionic Nonpolar covalent bonds Equal share of electrons Form nonpolar molecules Have partially similar charges at the opposite ends of a molecule More nonpolar = more covalent What is electronegativity? The measure of the ability of an atom to attract electrons towards it. The greater the electronegativity of an atom, the greater is its ability to attract electrons. Electronegativity increases as you go and as you go of the periodic table. This is because the nuclear pull increases towards the top and towards the right hand side of the periodic table. What are the two factors necessary in determining the polarity of a molecule? Symmetry between molecules Shape of the molecule Intermolecular Forces Intermolecular forces are the forces of attraction that exist between molecules They are also known as Van der Waals forces and they give a molecule it's characteristic properties. Intermolecular forces < covalent bonds There are three types of intermolecular forces: 1) DipoleDipole force: Attractive force between dipoles (polar molecules) The strength of the force depends on the polarity of the molecules 2) London force: Attraction of the nucleus of one atom towards the electrons of another atom The strength of this force depends on the number of electrons that were attracted 3) Hydrogen bonding: The attraction of hydrogen atoms towards more electronegative atoms with lone paris of electrons Energy change is an indication that hydrogen bonds are being formed The theory of hydrogen bonding is important to understand because they support many of out life processes. Like proteins and DNA. Liquids: The physical properties of liquids depend on the strength of the intermolecular bonds that hold each molecule together. Some of these properties are: Surface Tension: The property of a liquid that allows it to behave like an elastic. Molecules near the surface undergo a net attraction towards the molecules that are available and the molecules that are present inside the liquid form bonds equally with all the other molecules. Because of this property, surface tension, a liquid can support light objects (such as insects). Capillary action: Occurs due to the intermolecular forces. This is the property of a liquid which allows it to rise or fall when it comes in contact with a solid surface. For example, collect some water in a tray and add red food colouring to it. Now take a white flower with a long stem and dip the tip of the stem in the coloured water. After a while, the flower would be slightly red in colour. This change in colour was the effect of capillary action. STRUCTURE OF WATER: When temperature is decreased, the arrangement of the molecules is compressed in size. This allows them to form a compact structure and have an increase in density. This is the normal process for many liquids. But water undergoes a different scenario. When temperature is decreased, the water molecules rearrange themselves to form a hexagonal shape with spaces in the middle. This structure allows ice to decrease in density and be lighter than water. Solids share many characteristics: having a definite shape and volume, they do not flow, are incompressible. But there are many other characteristics that distinguish them. IONIC CRYSTALS: The positive ions are attracted to the negative ions to form a strong ionic bond. This arrangement is repeated to form an ionic crystal. They are hard, have high melting and boiling points, conduct electricity in their liquid state. IONIC BONDS > INTERMOLECULAR BONDS METALLIC CRYSTALS: They have a crystal lattice. They have a continuous and compact structure. The nuclei are held together by the electrons. They are shiny (absorb and give off light), flexible, conduct electricity (as the electrons are free to move around). MOLECULAR CRYSTAL: They have a crystal lattice like ionic compounds. The only difference is that the atoms are neutrally charged. They have low melting points, are not hard and nonconductors of electricity. They are held together by intermolecular force Molecular < ionic and/or covalent bonds The Structure and Properties of Solids COVALENTNETWORK CRYSTALS: They are a network of covalent bonds formed by carbons. Each carbon is bonded to three others. Their interlocking structure is thought to be responsible for the strength of the crystals. They are hard, brittle, and insoluble, have high melting points and are nonconductors of electricity. The electrons are held within the atom or in the covalent bonds. So, they aren't free to move around. This is why they are nonconductors of electricity. They are the hardest materials on earth. Example, a Diamond. Carbon can also form other Covalent Networks. Such as: -Tetrahedral arrangements (diamonds) -Layers of sheets (Graphite) -Spherical molecules (buckyballs) -Long tubes (nanotubes) Semiconductors: Semiconductors are materials that allows current to flow under certain circumstances. They are prepared by adding small amounts of impurities to a pure conductor. This process is called doping and it changes the original conductor's electrical properties. Electrons in a semiconductor only require a small amount of energy to jump to higher levels. It becomes easier for them to join other atoms. Conducting electricity: Insulators > semiconductors > conductors Semiconductors as power supplies for many satellites, in solar cells, transistors and many other things. Chapter 4 Review Lewis Theory: atoms and ions are stable when they have a full octet and atoms form bonds to reach stable octet Valence bond theory: a covalent bond is formed when 2 orbitals overlap and produce a new combined orbital containing 2 electrons of opposite spin. Hybridization: when electrons are excited, they jump from one orbital to another forming new hybrid orbitals, which form the covalent bonds Sigma bond: orbitals overlap endto end Pi bond: orbitals overlap sidebyside; used for double bonds VSEPR: valenceshellelectronpairrepulsion theory; based on the electrical repulsion of bonded and unbonded electron pairs in a molecule or polyatomic ion. The property of a molecule of being structurally different at opposite ends of its long axis is called polarity. The measure of the ability of an atom to attract electrons towards is known as the atom's electronegativity. Intermolecular forces are the forces of attraction that exist between molecules. There are three types of forces: DipoleDipole, London Dispersion and Hydrogen Bonding. The physical properties of a liquid depends on the strength of the intermolecular forces. For example, surface tension and capillary action. Water has an exceptional structure that allows it to have unique properties. The different categories of solids are: Ionic Crystals, Metallic Crystals, Molecular Crystals, and Covalent Network Crystals. ...
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