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# Atomic Structure Periodicity - PowerPoint PPT Presentation

Atomic Structure & Periodicity. What is the nature of the atom? How can atomic structure account for the periodic properties observed? What is quantum mechanics?. Electromagnetic Radiation. Electromagnetic Radiation a way that energy travels through space exhibits wavelike behavior

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## PowerPoint Slideshow about 'Atomic Structure Periodicity' - wilmet

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Presentation Transcript

• What is the nature of the atom?

• How can atomic structure account for the periodic properties observed?

• What is quantum mechanics?

• a way that energy travels through space

• exhibits wavelike behavior

• travels at the speed of light

• Three primary characteristics of waves

• wavelength

• frequency

• speed

• Wavelength

• l = lamda

• distance between two consecutive peaks

• units could be meters, nm, Angstroms, cm, etc.

• Frequency

• n = nu

• number of waves or cycles that pass a certain point per second

• units may be cycles/second or Hertz (Hz) or just sec-1

• Speed

• all types of EM (electromagnetic) radiation travel at the speed of light

• Speed = c = 3.00 x 108 m/sec

• There is an inverse relationship between wavelength and frequency.

• l a 1/n

• in other words

• ln = c

• The shorter the wavelength, the higher the frequency.

• Low frequency, long wavelength

• An important means of energy transfer

• Electromagnetic Spectrum:

• Calculate the frequency of red light of wavelength 6.50 x 102 nm.

• (Answer: n = 4.61 x 1014 Hz)

• Classical Physics…pre-1900’s

• could predict the motion of planets

• could explain the dispersion of light by a prism

• the assumption was that physicists at that time knew all that there was to know about physics

• Pre-1900’s

• Matter and Energy are distinct

• Matter is particulate

• has mass

• position in space could be specified

• Energy is wavelike

• no mass

• delocalized

• 1900

• Max Planck (1858 - 1947)

• Studied the radiation profiles emitted by solid bodies heated to incandescence (I.e., heated so hot that the objects gave off light)

• These profiles could not be explained with classical physics

• Classical Physics - matter could absorb or emit any quantity of energy

• Planck observed that energy could be gained or lost only in whole number multiples of a quantity, hn

• h = Planck’s constant h = 6.626 x 10-34 J.sec

• n = the frequency of the energy absorbed or emitted

• Max Planck’s equation for the change in energy for a system:

• DE = nhn ( where n = 1, 2, 3, …)

• From Planck’s work, we learn that energy is quantized

• energy occurs in discrete packets called quanta; one packet of energy is called a quantum

• THUS, Energy seems to have particulate properties!

• The blue color in fireworks is often achieved by heating copper (I) chloride to about 1200oC. Then the compound emits blue light with a wavelength of 450 nm. What is the quantum that is emitted at 4.50 x 102 nm by CuCl?

• (Answer: DE = 4.41 x 10-19 J)

• Albert Einstein

• proposed that electromagnetic radiation is quantized

• EMR can be viewed as a stream of “particles” known as photons

• The energy of the photon then is:

• Ephoton = hn = hc

l

• Albert Einstein (cont.)

• E = mc2

• energy has mass

• m = E

c2

• use this equation to calculate the mass associated with a given quantity of energy

• Mass of a photon:

• Ephoton = hc and m = E

l c2

• so m = hc/l

c2

• m = h ( so the mass of a photon depends on its wavelength)

cl

• Arthur Compton’s work (1922) with X rays and electrons showed that photons have the mass calculated

• Conclusions:

• Energy is quantized

• Energy can occur is discrete units called quanta

• EMR can show characteristics of particulate matter (photons) as well as wavelike characteristics

• This phenomenon is known as the dual nature of light

• So light can be particulate as well as wavelike

• Can matter then be wavelike as well as particulate?

• Louis de Broglie (1892 - 1987)

• for EMR: m = h/lc

• for a particle: m = h/lv

• v = velocity …because matter does not travel at the speed of light

• so rearrange to solve for the wavelength of a particle:

• l = h (de Broglie’s equation)

mv

• Compare the wavelength for an electron ( a particle with m = 9.11 x 10-31 kg) traveling at a speed of 1.0 x 107m/s with that for a ball ( a particle with m = 0.10 kg) traveling at 35 m/s

• (Answer: le = 7.27 x 10-11 m; lb = 1.9 x 10-34 m)

• How can de Broglie’s equation be tested?

• The wavelength of an electron is the same length as the distance between the atoms in a typical crystal

• A crystal diffracts electrons just as it diffracts EMR (in the form of X-rays)

• Therefore, electrons do have an associated wavelength

• EMR was found to possess particulate properties

• Particles, like electrons, were found to have an associated wavelength

• Matter and Energy are not distinct!

• Energy is really a form of matter!

• Large pieces of matter - predominately particulate

• Very small pieces of matter (e.g. photons) are predominately wavelike (but can exhibit particulate properties)

• Intermediate pieces of matter (e.g. electrons) are wavelike as well as particulate

• Add energy (in the form of a spark) to H2(g)

• H2 molecules absorb energy, some H-H bonds are broken, and the H atoms get excited

• this excess energy is released in the form of light

• an emission spectrum (the pattern of light emitted) is always the same for a particular element (like a fingerprint for an element)

• passing H’s emission spectrum through a prism results in a few characteristic lines…hence the term line spectrum

• H’s line spectrum shows that hydrogen’s electrons are quantized

• I.e., only certain energies are allowed for the electron

• if not, then the spectrum would be continuous, like the rainbow observed when light passes through a prism

• fits in with Max Planck’s postulates

• DE = hn = hc

l

• Niels Bohr (1885 - 1962)

• 1913 developed a quantum model for the hydrogen atom

• the electron in a hydrogen atom moves around the nucleus in certain allowed orbits

• used classical physics and made some new assumptions to calculate these orbits

• Bohr’s model had to account for the line spectrum of hydrogen

• the hydrogen atom has energy levels

• E = -2.178 x 10-18 J (Z2/n2)

• Z = the nuclear charge (in Hydrogen, Z = 1)

• n = (1, 2, 3, …) the energy level ...the larger the value for n, the larger the orbit radius…

• negative sign indicates that an electron bound to a nucleus has lower energy than an electron infinitely far away from the nucleus

• n = 1 …the ground state for hydrogen…the electron is closest to the nucleus

• When energy is added to the H atom, the electron jumps up to a higher energy level

• When giving off energy, the electron falls back to the ground state, or the lowest energy state

• Calculate the energy required to excite the hydrogen electron from n = 1 to n = 2. Also, calculate the wavelength of light that must be absorbed by a hydrogen atom in its ground state to reach this excited state.

• (Answer: DE = 1.633 x 10-18J; l = 1.216 x 10-7m)

• The energy levels calculated agree with the line spectrum for hydrogen

• Cannot be applied to atoms other than hydrogen

• Still a good model because it introduced the quantization of energy in atoms

• So the Bohr model did not work for other atoms

• Needed a new approach

• Three physicists to the rescue:

• Heisenberg (1901 - 1976)

• de Broglie (1892 - 1987)

• Schrodinger (1887 - 1961)

• Quantum Mechanics = Wave Mechanics

• Emphasized the wave nature of the electron

• The electron behaves like a “standing wave”

• similar to the stationary waves of string instruments like the guitar and violin

• existance of nodes

• only whole number of half wavelengths are allowed

• Schrodinger’s equation

• a mathematical treatment of the electron

• too complicated for this course!

• Electron’s position in space is described by a wave function

• a specific wave function is an orbital

• An orbital

• not a Bohr model orbit

• not a circular path

• actually, when we describe an orbital, we do not know exactly how the electron is moving

• Heisenberg Uncertainty Principle

• we cannot know accurately both the position and the momentum of a particle (such as the electron) at any given time

• I.e., the more accurately we know the position, the less accurately we know the momentum…and vice versa

• So what does the wave function tell us?

• Nothing we can visualize

• The square of the wave function gives us the probability of finding an electron in a particular location in space

• electron density = electron probability = atomic orbital

• Atomic Orbital

• a volume that encloses 90% of the total electron probability…

• a volume where the electron can be found 90% of the time

• Quantum numbers

• Schrodinger’s equation has many solutions

• I.e., there are many orbitals described by the wave functions

• each orbital can be described by a set of quantum numbers

• Principal quantum number

• n

• n = 1, 2, 3, …

• related to the size and energy of the orbital

• as n increases, the orbital becomes larger, so the electron is farther from the nucleus, and has higher energy

• Angular momentum quantum number

• l

• l = 0 …n - 1

• related to the shape of the atomic orbital

• l = 0 (s orbital)

• l = 1 ( p orbital)

• l = 2 (d orbital)

• l =3 ( f orbital)

• Magnetic Quantum Number

• ml

• ml = -l…0…l

• related to the orbital’s orientation in space

• Electron Spin Quantum Number

• ms

• ms= + 1/2

• an electron can spin in one of two opposite directions

• Orbital Shape and Energies

• Nodal Surfaces = nodes

• areas of zero probability of finding an electron

• as n increases, the number of nodal surfaces increases

• Degenerate orbitals (in hydrogen)

• have the same value of n

• have the same energy

• Add energy to an atom

• electron becomes excited

• electron is transferred to a higher energy orbital

• Electrons spin in two directions

• a spinning charge produces a magnetic field

• Pauli Exclusion Principle

• Wolfgang Pauli (1900 - 1958)

• in an atom, no two electrons can have the same set of quantum numbers…I.e., e-’s in the same orbital have opposite spin

• When there are many electrons in an atom, what is happening to these electrons?

• Electrons

• have kinetic energy as they move around

• feel attractive forces from the nucleus

• are repelled by other electrons

• Because of the Heisenberg Uncertainty Principle, the repulsions between e-’s cannot be calculated exactly

• Approximate forces on electrons

• Approximate which has greater effect, attraction from nucleus or repulsion from other electrons

• Example: Na atom

• 11 protons, 11 electrons

• look at the outermost 3s electron

• it is attracted to the 11 protons in the nucleus

• however, it doesn’t feel all 11 protons because it is shielded by (repulsions of) the inner electrons

• In a polyelectronic atom

• shape of orbitals are the same as for hydrogen

• size and energies will be different than for hydrogen’s orbital

• due to interplay between nuclear attraction and electron repulsion

• In polyelectronic atoms

• orbitals on the same energy level are not degenerate

• they do not have the same energy

• Es < Ep < Ed < Ef

• energy of an orbital depends on the electron’s ability to penetrate and spend time close to the nucleus

• Periodic Table

• constructed to represent the patterns observed in the chemical properties of elements

• Johann Dobereiner (1780 - 1849)

• found several groups of three elements with similar properties

• e.g., Cl, Br, and I

• John Newlands, 1846

• suggested elements be arranged in octaves

• some properties seem to repeat every eighth element

• Julius Lothar Meyer (1830 - 1895)

• Dmitri Mendeleev (1834 - 1907)

• both men independently developed what is the modern periodic table

• Mendeleev correctly predicted the properties and existence of Sc, Ga, and Ge (ekasilicon)

• Modern periodic table arranges elements by atomic number rather than atomic mass

• Aufbau (building up) Principle

• electrons are added one by one to hydrogen-like atomic orbitals

• Orbital Filling Diagrams show this principle

• Hund’s rule

• lowest energy arrangement occurs when electrons occupy separate orbitals with parallel spins

• I.e., w/in a sublevel, there is no doubling up w/in an orbital until each orbital on the sublevel has one electron

• Valence electrons

• electrons in the outermost principle quantum level of an atom

• electrons in the highest s and p orbitals of an atom

• most important electrons to chemists

• electrons involved in bonding

• Inner electrons = core electrons

• Elements in the same group have the same outer electron configuration

• have the same number of valence electrons

• perhaps this is the reason for the similarities in chemical properties?

• Some exceptions to the Aufbau principle

• no good explanation at this time

• Memorize:

• Cr [Ar]4s13d5

• Cu [Ar]4s13d10

• Many exceptions in the Lanthanide and Actinide series…it’s not necessary to memorize these exceptions

• one of the reasons for the irregularity of filling orbitals is that the 5f and 6d orbitals and the 4f and 5d orbitals have very similar energies,

• Model of the atom can be used to account for important atomic properties

• ionization energy

• electron affinity

• atomic size

• metallic character

• Ionization energy

• the energy required to remove an electron from a gaseous atom or ion:

• X(g) ---> X+(g) + e-

• the electron is removed from an atom in the ground state

• the highest energy electron is removed first

• i.e., the e- farthest from the nucleus is removed first

• Al(g) --> Al+(g) + e- I.E.1 = 580 kJ/mol

• Al+(g) --> Al+2(g) + e- I.E.2 = 1815 kJ/mol

• Al+2(g) --> Al+3(g) + e- I.E.3 = 2740 kJ/mol

• Al+3(g) --> Al+4(g) + e- I.E.4 = 11,600 kJ/mol

• Al: 1s22s22p63s23p1

• 1st I.E. - remove the 3p electron

• 2nd I.E. - remove one of the 3s electrons

• requires more than 3 times the energy than removing the first electron

• 2nd e- feels a greater positive charge (e- is being removed from a positive ion), and is held more tightly

• Largest jump in I.E. occurs between I.E.3 and I.E.4

• I.E.4 is high due to the removal of a 2p electron which experiences a large positive charge

• a 2p electron is a core or inner electron which experiences a greater nuclear charge than the valence electrons which are shielded by the core electrons

• Periodic Trend for I.E.

• Going across a row

• I.E. Increases

• going across a row, e-’s are being added to the outer energy level

• there is a greater nuclear charge as well which is felt by each of the valence electrons

• Going down a column

• I.E. Decreases

• going down a row, electrons are being added to energy levels that are farther from the nucleus

• even though there are more protons, the electrons are getting farther away, and feel less of the nuclear charge, and so are easier to remove

• Exceptions to the general trend for I.E.

• Explain exceptions in terms of electron repulsions

• two electrons in one orbital do repel each other, removing one of the electrons would eliminate this repulsion, thus stabilizing the atom...so less energy is required to remove one of these “repulsive” electrons

• Electron Affinity

• energy change when an electron is added to a gaseous atom

• X(g) + e- ---> X-(g)

• If energy is added when an electron is added, the sign for E.A. is positive (endothermic)

• If energy is released when an electron is added , the sign for E.A. is negative (exothermic)

• The more negative the E.A., the more energy is released when an electron is added to an atom

• the negative ion is more “stable,” i.e., lower in energy, than the neutral atom

• Periodic Trend for E.A.

• Going across a row

• becomes more negative

• some exceptions to the general trend

• explain in terms of electron repulsion

• if an electron is added to a previously occupied orbital, repulsion, or destabilization, occurs, so less energy will be given off (E.A. Will be less negative)

• Periodic Trend for E.A.

• Going down a column

• e-’s are farther from the nucleus

• outer e-’s don’t feel the nuclear charge as strongly

• E.A. Becomes more positive going down a column

• differences in E.A. Is less going down a column than in going across a period

• The atomic radius cannot be measured exactly

• defined as half the distance between two nuclei of identical atoms

• otherwise, the atomic radii for an element are estimated from the element’s various covalent compounds

• Periodic Trend for Atomic Radii

• Going across a row

• increased # of protons means each outer electron feels a greater nuclear charge. This results in the electrons being pulled closer to the nucleus, and the atomic radius will decrease

• Periodic Trend for Atomic Radii

• going down a column

• e- are being added to higher energy levels that are farther from the nucleus

• Metallic Character

• Metals tend to form positive ions easily

• Metals have low I.E.

• High metallic character then would be found in elements at the lower left hand corner of the periodic table

• Nonmetals

• have high ionization energy

• tend to form negative ions

• tend to have large negative E.A.’s

• most reactive nonmetals are found in the upper right hand corner of the Periodic Table

• The periodic table is a tool for predicting properties of an element or of a group of elements

• The Groups of Representative elements exhibit similar chemical properties

• these chemical properties vary in a predictable way...

• Each group member have the same number of valence electrons

• the number of valence electrons determine an element’s chemistry

• An element’s electron configuration is very important as an aid to predicting an element’s properties