Ap chem unit 7

Atomic Structure and Periodicity,[object Object],AP Chem: Unit 7,[object Object]

Electromagnetic Radiation,[object Object]

Electromagnetic Radiation,[object Object],One of the ways that energy travels through space is by electromagnetic radiation. ,[object Object],light from the sun,[object Object],X-rays,[object Object],microwaves,[object Object]

Wave Characteristics,[object Object],Wavelength (λ) – is the distance between two consecutive peaks or troughs in a wave.,[object Object],Frequency (ν) – is the number of waves (cycles) per second that pass a given point in space. ,[object Object],units – hertz or waves/sec (s-1),[object Object],Speed (c) – all types of electromagnetic radiation travel at the speed of light.,[object Object],2.9979 x 108m/s,[object Object],c = λν,[object Object]

The Nature of Matter,[object Object]

Wave and Particle Duality,[object Object],Planck found that matter could only absorb or emit energy in whole number multiples of the quantity hν. ,[object Object],h is Planck’s constant = 6.626 x10-34Js,[object Object],ΔE = hν,[object Object],Transfer of energy is not continuous but is quantized and can occur only in discrete amounts called quantum. Thus energy has particle properties as well as wave properties.,[object Object]

Wave and Particle Duality,[object Object],Einstein proposed that electromagnetic radiation was also quantized and could be viewed as a stream of “particles” called photons. ,[object Object],Ephoton = hv = hc/λ,[object Object]

The Photoelectric Effect,[object Object],The photoelectric effect refers to the phenomenon in which electrons are emitted from the surface of a metal when light strikes it. ,[object Object],No electrons are emitted by a metal below a specific threshold frequency (vo),[object Object],For light with frequency lower than the threshold frequency, no electrons are emitted regardless of intensity of the light.,[object Object]

The Photoelectric Effect,[object Object],For light with frequency greater than the threshold frequency, the number of electrons emitted increases with the intensity of the light.,[object Object],For light with frequency greater than the threshold frequency, the kinetic energy of the emitted electrons increases directly with frequency of the light.,[object Object]

The Photoelectric Effect,[object Object],These observations can be explained by assuming that electromagnetic radiation is quantized (consists of photons), and that the threshold frequency represents the minimum energy required to remove the electron from the metal’s surface.,[object Object],Minimum energy required to remove an electron = Eo = hvo,[object Object],KEelectron = ½ mv2 = hv – hvo,[object Object]

Planck and Einstein Conclusions,[object Object],Energy is quantized. It can occur only in discrete units called quanta.,[object Object],Electromagnetic radiation, which was previously thought to exhibit only wave properties, seems to show certain characteristics of particulate matter as well. This phenomenon is sometimes referred to as the dual nature of light.,[object Object]

Wave Particle Duality,[object Object],The main significance of the equation E = mc2 is that energy has mass. ,[object Object],m = E/c2,[object Object]

Louis de Broglie (1892-1987),[object Object],Since light which previously was thought to be purely wavelike, was found to have certain characteristics of particulate matter. But is the opposite also true? Does matter have that is normally assumed to be particulate exhibit wave properties?,[object Object]

Louis de Broglie (1892-1987),[object Object],de Broglie’s equation allows us to calculate the wavelength for a particle:,[object Object]

de Broglie’s Proof,[object Object]

Louis de Broglie (1892-1987),[object Object],Conclusion: Energy is really a form of matter, and all matter shows the same types of properties. All matter exhibits both particulate and wave properties. ,[object Object]

The Atomic Spectrum of Hydrogen,[object Object]

Spectrum,[object Object],A continuous spectrum results when white light passes through a prism and all wavelengths (colors) are shown.,[object Object],An emission spectrum produces only a few lines of color that is limited to discrete wavelengths produced by an atom. This is called a line spectrum and is specific to each atom.,[object Object]

Hydrogen Line Spectrum,[object Object],The significance of the line spectrum is that it indicates that only certain energies are allowed for the electron in the hydrogen atom. In other words the energy of the electron in the hydrogen atom is quantized ,[object Object]

Hydrogen Line Spectrum,[object Object]

The Bohr Model,[object Object]

Niels Bohr,[object Object],Bohr developed a quantum model for the hydrogen atom that allowed for only specific energy levels around the atom that corresponded with specific radii.,[object Object]

Niels Bohr (1885-1962),[object Object],The most important equation to come from Bohr’s model is the expression for the energy levels available to the electron in the hydrogen atom.,[object Object],Z is the nuclear charge, n is the energy level. ,[object Object]

Niels Bohr (1885-1962),[object Object],The most important equation to come from Bohr’s model is the expression for the energy levels available to the electron in the hydrogen atom.,[object Object],the negative sign calculates a lower energy closer to the atom, not the radiation of negative energy.,[object Object]

Example,[object Object],What is the change in energy if an electron in level 6 (excited state) returns to level 1 (ground state) in a hydrogen atom?,[object Object],ni=6; nf=1; Z=1 (hydrogen nucleus contains a single proton),[object Object]

Example,[object Object],What is the change in energy if an electron in level 6 (excited state) returns to level 1 (ground state) in a hydrogen atom?,[object Object]

Example,[object Object],ΔE=Ef – Ei= E1 – E6=-2.117 x 10-18J,[object Object],The negative sign for the change in energy indicates that the atom has lost energy and is now more stable. This loss of energy produces a photon.,[object Object]

Example,[object Object],What is the corresponding wavelength for the energy produced from the electron jump?,[object Object],E = -2.117 x 10-18J,[object Object],9.383x10-8 m,[object Object]

Bohr Model Conclusions,[object Object],The model correctly fits the quantized energy levels of the hydrogen atom and postulates only certain allowed circular orbits for the electrons.,[object Object],As the electron becomes more tightly bound, its energy becomes more negative relative to the zero-energy reference state. As the electron is brought closer to the nucleus, energy is released from the system. ,[object Object]

Bohr Model Conclusions,[object Object]

Bohr Model Conclusions,[object Object],The energy levels calculated by Bohr closely agreed with the values obtained from the hydrogen emission spectrum but does not apply well to other atoms. The Bohr’s model is fundamentally incorrect but is very important historically because it paved the way for our current theory of atomic structure.,[object Object]

The Quantum Mechanical Model of the Atom,[object Object]

Quantum Mechanics,[object Object],Quantum Mechanics or Wave Mechanics were developed by three physicists: Heisenberg, de Broglie, and Schrodinger.,[object Object],Emphasis was given to the wave properties of the electron.,[object Object],The electron bound to the nucleus behaves similar to a standing wave.,[object Object]

Quantum Mechanics,[object Object],Like a standing wave, electrons can travel in patterns that allow for a common node. In other words, wave patterns around the nucleus must be in whole number wave patterns. But their exact movement is not known.,[object Object]

Heisenberg Uncertainty Principle,[object Object],There is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time. This limitation is small for large particles but substantial for electrons. ,[object Object]

Probability Distribution,[object Object],A probability distribution is used to indicate the probability of finding an electron in a specific position.,[object Object],Electron density map,[object Object],Radial probability distribution ,[object Object]

Probability Distribution,[object Object],For the hydrogen 1s orbital, the maximum radial probability occurs at a distance of 5.29x10-2nm or .529Å from the nucleus. This is the exact radius of the innermost orbit calculated in the Bohr Model.,[object Object],The definition most often used by chemists to describe the size of the hydrogen 1s orbital is the radius of the sphere that encloses 90% of the total electron probability,[object Object]

Quantum Numbers,[object Object]

Quantum Numbers,[object Object],Each orbital is characterized by a series of numbers called quantum numbers, which describe various properties of an orbital:,[object Object],Principal quantum number (n)- has integral values : 1,2,3,4. It describes the size and energy of the orbital. Energy Level,[object Object]

Quantum Numbers,[object Object],Angular momentum quantum number (l) – has integral values from 0 to n-1. This is related to shape of the atomic orbitals. Sublevel,[object Object],[object Object]

Quantum Numbers,[object Object],Electron spin quantum number (ms)- can only have one of two values, +½, -½. Electrons can spin in one of two opposite directions. ,[object Object]

Quantum Numbers,[object Object],In a given atom no two electrons can have the same set of four quantum numbers (n, l, ml , ms). This is called the Pauli exclusion principle; an orbital can only hold two electrons, and they must have opposite spins. ,[object Object]

Orbital Shapes and Energies,[object Object]

S orbitals,[object Object],The s orbitals have a characteristic spherical shape and contain areas of high probability separated by areas of zero probability. These areas are called nodal surfaces, or nodes.,[object Object]

S orbitals,[object Object],The number of nodes increases as n increases. The number of nodes = n - 1. ,[object Object]

P Orbitals,[object Object],P orbitals each have two lobes separated by a node at the nucleus. The p orbitals are labeled according to the axis of the xyz coordinate system along which the lobes lie. ,[object Object]

P Orbitals,[object Object],Cross section of electron probability of a p orbital,[object Object]

D Orbitals,[object Object],The five d orbitals first occur in energy level 3. They have two fundamental shapes. Four of the orbitals (dxz, dyz, dxy, and dx2-y2) have four lobes centered in the plane indicated in the orbital label. dx2-y2 lie along the x and y axes and dxylie between the axes. The fifth orbital dz2 has a unique shape with two lobes along the z axis and a belt centered in the xy plane. ,[object Object]

F Orbitals,[object Object],The f orbitals first occur in level 4 and have shapes more complex than those of the d orbitals. These orbitals are not involved in the bonding in any of the compounds that we will consider. ,[object Object]

Orbital Energies ,[object Object],For the hydrogen atom, the energy of a particular orbital is determined by its value of n. Thus all orbitals with the same value of n have the same energy – they are said to be degenerate.,[object Object]

Polyelectronic Atoms,[object Object]

Polyelectronic Atoms,[object Object],Polyelectronic atoms are atoms with more than one electron. To look at these atoms, three energy contributions must be considered:,[object Object],Kinetic energy of the electrons as they move around the nucleus.,[object Object],The potential energy of attraction between the nucleus and the electrons.,[object Object],The potential energy of repulsion between the two electrons.,[object Object]

Polyelectronic Atoms,[object Object],Since electron pathways are unknown, dealing with the repulsions between electrons cannot be calculated exactly.,[object Object],This is called the electron correlation problem.,[object Object]

Polyelectronic Atoms,[object Object],The electron correlation problem occurs with all polyelectronic atoms. To deal with this, we assume each electron is moving in a field of charge that is the net result of the nuclear attraction and the average repulsions of all the other electrons.,[object Object],In other words,…..,[object Object]

Polyelectronic Atoms,[object Object],A valence electron is attracted to the highly charged nucleus and still repelled by the other ‘inner’ electrons. The net effect is that the electron is not bound nearly as tightly to the nucleus as it would be if it were alone. ,[object Object], This is a screened or shielded affect. ,[object Object]

Polyelectronic Atoms,[object Object],Because of this shielded affect. orbitals within a principal energy level do not have the same energy (degenerate). Sublevels vary in energy within a principal quantum level.,[object Object],s<p<d<f,[object Object]

Polyelectronic Atoms,[object Object],Hydrogen vs. Polyelectronic,[object Object]

History of the Periodic Table,[object Object]

Early Greeks,[object Object],Earth ,[object Object],Air,[object Object],Fire ,[object Object],Water,[object Object]

Dobereiner (1780-1849),[object Object],Johann Dobereiner was the first chemist to recognize patterns and found several groups of three elements that have similar properties.,[object Object],chlorine, bromine and iodine,[object Object],called triads.,[object Object]

Newlands,[object Object],John Newlands suggested that elements should be arranged in octaves, based on the idea that certain properties seemed to repeat for every eighth element in a way similar to the musical scale.,[object Object]

Meyer and Mendeleev,[object Object],The present form of the periodic table was conceived independently by two chemists: Meyer and Mendeleev. Usually Mendeleev is given most of the credit, because it was he who emphasized how useful the table could be in predicting the existence and properties of still unknown elements. ,[object Object]

Meyer and Mendeleev,[object Object],In 1872 when Mendeleev first published his table, the elements gallium, scandium, and germanium were unknown. Mendeleev correctly predicted the existence and properties of these elements from the gaps in his periodic table. Mendeleev also corrected the atomic masses of several elements. ,[object Object]

Mendeleev’s Table,[object Object]

The Aufbau Principle and the Periodic Table,[object Object]

Three rules for Orbital Configuration,[object Object],Aufbau principle – As protons are added, so are electrons, and fill in orbitals in order of energy levels.,[object Object],Pauli Exclusion – Two electrons with opposite spins can occupy an orbital.,[object Object],Hund’s rule – The lowest energy configuration for an atom is the one with one unpaired electrons in each degenerate orbital. (Electrons don’t like roommates),[object Object]

Valence Electrons,[object Object],Valence electrons are the electrons in the outermost principal quantum level of an atom. These are the most important electrons because they are involved in bonding. ,[object Object],The inner electrons are known as core electrons.,[object Object]

Valence Electrons,[object Object],The elements in the same group have the same valence electron configuration. Elements with the same valence electron configuration show similar chemical behavior. ,[object Object]

Transition Metals,[object Object],Transition metals have electron configurations that fill in the order of 4s before 3d. Copper and Chromium have a configuration that is observed different than what is expected.,[object Object],Expected: Cr: 1s22s22p63s23p64s23d4,[object Object],Observed: 1s22s22p63s23p64s13d5,[object Object],Expected: Cu:1s22s22p63s23p64s23d9,[object Object],Observed: 1s22s22p63s23p64s13d10,[object Object]

Transition Metals,[object Object]

Additional Orbital Rules,[object Object],The (n+1)s orbital always fills before the nd orbitals. The s orbitals fill prior to the d orbitals due to the vicinity of the nucleus.,[object Object],After lanthanum, which has the configuration of [Xe] 6s25d1, a group of 14 elements called the lanthanide series, or the lanthanides occurs. This seris of elements corresponds to the filling of the seven 4f orbitals. ,[object Object]

Additional Orbital Rules,[object Object]

Additional Orbital Rules,[object Object],After actinium, a group of 14 elements called the actinide series or actinides occurs. ,[object Object],The groups 1A, 2A, 3A…, the group numbers indicate the total number of valence electrons for the atoms in these groups.,[object Object]

Periodic Trends in Atomic Properties,[object Object]

Ionization Energy,[object Object],Ionization energy is the energy required to remove an electron from a gaseous atom or ion when the atom or ion is assumed to be in its ground state:,[object Object],X(g)X+(g) + e-,[object Object]

Ionization Energy,[object Object],It is always the highest-energy electron (the one bound least tightly) that is removed first. The first ionization energy (I1) is the energy required to remove that first electron. The second ionization energy (I2) is considerably larger. ,[object Object]

Ionization Energy,[object Object],The first electron is removed from a neutral atom, the second from a +1 cation. The increase in positive charge binds the electrons more firmly and the ionization energy increases. The trend continues for consecutive electrons removed.,[object Object],Core electrons are always held tighter than valence. ,[object Object]

Ionization Energy,[object Object],First ionization energy increases from left to right across a period.,[object Object],First ionization energy decreases in going down a group.,[object Object]

Ionization Energy,[object Object]

Electron Affinity,[object Object],Electron Affinity is the change in energy change associated with the addition of an electron to a gaseous atom:,[object Object],X(g) + e- X-(g),[object Object]

Electron Affinity,[object Object],If the addition of the electron is exothermic the corresponding value for electron affinity will carry a negative sign.,[object Object],The more negative the energy, the greater the quantity of energy released.,[object Object]

Electron Affinity,[object Object],Electron affinities generally become more negative from left to right across a period and becomes more positive down a group.,[object Object],As with Ionization energy. Some exceptions occur due to repulsions and electron configuration. ,[object Object]

Electron Affinity,[object Object]

Atomic Radius,[object Object],Atomic radii are measured by the distances between atoms in chemical compounds. ,[object Object],Covalent atomic radii are assumed to be half the distance between atoms in covalent bonds.,[object Object],For metallic atoms, the metallic radii are obtained from half the distance between metal atoms in a solid metal crystal,[object Object]

Atomic Radius,[object Object],Atomic radii decrease in going from left to right across a period because of increasing nuclear charge and decreasing shielding.,[object Object],Atomic radius increases down a group, because of the increases in the orbitals sizes associated with principal quantum numbers. ,[object Object]

The Properties of a Group: The Alkali Metals,[object Object]

Information and the Periodic Table,[object Object],It is the number and type of valence electrons that primarily determine an atom’s chemistry,[object Object],The organization of the period table allows the prediction of electron configuration without memorization.,[object Object]

Information and the Periodic Table,[object Object],Groups on the periodic table have specialized names: Alkali metals, Alkaline earth metals, Halogens, …etc.,[object Object],The most basic division of elements in the periodic table is into metals and non-metals. This division affects chemical properties.,[object Object],Metals tend to give up electrons and have low ionization energies. The opposite is true for non-metals. ,[object Object]

Information and the Periodic Table,[object Object]

Information and the Periodic Table,[object Object],Metalloids are elements along the division line and exhibit both metallic and nonmetallic properties under certain circumstances. These elements are sometimes called semimetals. ,[object Object]

The Alkali Metals,[object Object],Lithium, sodium, potassium, rubidium, cesium, and francium are the most chemically reactive of the metals. Hydrogen is found in group 1 but behaves as a nonmetal because its very small and the electron is bound tightly to the nucleus. ,[object Object]

The Alkali Metals,[object Object],Going down the group the first ionization energy decreases and the atomic radius increases. The overall density increases due to the increase of atomic mass relative to atomic size (therefore more mass per unit volume). ,[object Object]

The Alkali Metals,[object Object],There is a smooth decrease in melting point and boiling points in Group 1 that is not typical for other groups. ,[object Object],The most important chemical property of Group 1 is its ability to lose its valence electrons. Group 1 are very reactive. ,[object Object]

The Alkali Metals,[object Object],Hydration energy of an ion represents the change in energy that occurs when water molecules attach to the metal cation. ,[object Object],The hydration energy is greatest with Li+ because it has the most charge density (charge per unit volume). This means that polar water molecules are more strongly attracted to the small Li+ ions ,[object Object],The order of reducing abilities in an aqueous reaction is Li > K > Na,[object Object]

Ap chem unit 7

Ap chem unit 7

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Ap chem unit 7