Atomic structure periodicity
Download
1 / 79

Atomic Structure & Periodicity - PowerPoint PPT Presentation


  • 112 Views
  • Uploaded on

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

loader
I am the owner, or an agent authorized to act on behalf of the owner, of the copyrighted work described.
capcha
Download Presentation

PowerPoint Slideshow about 'Atomic Structure & Periodicity' - wilmet


An Image/Link below is provided (as is) to download presentation

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.


- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript
Atomic structure periodicity
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

  • Electromagnetic Radiation

    • a way that energy travels through space

    • exhibits wavelike behavior

    • travels at the speed of light


Waves
Waves

  • Three primary characteristics of waves

    • wavelength

    • frequency

    • speed


Waves1
Waves

  • Wavelength

    • l = lamda

    • distance between two consecutive peaks

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


Waves2
Waves

  • 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


Waves3
Waves

  • Speed

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

    • Speed = c = 3.00 x 108 m/sec


Waves4
Waves

  • 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


Electromagnetic radiation1
Electromagnetic Radiation

  • An important means of energy transfer

  • Electromagnetic Spectrum:


Electromagnetic radiation2
Electromagnetic Radiation

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

  • (Answer: n = 4.61 x 1014 Hz)


The nature of matter
The Nature of Matter

  • 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


The nature of matter1
The Nature of Matter

  • 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


The nature of matter2
The Nature of Matter

  • 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


The nature of matter3
The Nature of Matter

  • 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


The nature of matter4
The Nature of Matter

  • 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 nature of matter5
The Nature of Matter

  • 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)


The nature of matter6
The Nature of Matter

  • 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


The nature of matter7
The Nature of Matter

  • 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


The nature of matter8
The Nature of Matter

  • 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


The nature of matter9
The Nature of Matter

  • 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


The nature of matter10
The Nature of Matter

  • So light can be particulate as well as wavelike

  • Can matter then be wavelike as well as particulate?


The nature of matter11
The Nature of Matter

  • 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


The nature of matter12
The Nature of Matter

  • 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)


The nature of matter13
The Nature of Matter

  • 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


The nature of matter14
The Nature of Matter

  • 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!


The nature of matter15
The Nature 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


The atomic spectrum of hydrogen
The Atomic Spectrum Of Hydrogen

  • 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


The atomic spectrum of hydrogen1
The Atomic Spectrum of Hydrogen

  • 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


The bohr model of the hydrogen atom
The Bohr Model of the Hydrogen Atom

  • 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


The bohr model of the hydrogen atom1
The Bohr Model of the Hydrogen Atom

  • 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


The bohr model of the hydrogen atom2
The Bohr Model of the Hydrogen Atom

  • 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


The bohr model of the hydrogen atom3
The Bohr Model of the Hydrogen Atom

  • 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 bohr model of the hydrogen atom4
The Bohr Model of the Hydrogen Atom

  • 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


The quantum mechanical model of the atom
The Quantum Mechanical Model of the Atom

  • 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)


The quantum mechanical model of the atom1
The Quantum Mechanical Model of the Atom

  • 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


The quantum mechanical model of the atom2
The Quantum Mechanical Model of the Atom

  • 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


The quantum mechanical model of the atom3
The Quantum Mechanical Model of the Atom

  • 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


The quantum mechanical model of the atom4
The Quantum Mechanical Model of the Atom

  • 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


The quantum mechanical model of the atom5
The Quantum Mechanical Model of the Atom

  • 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


The quantum mechanical model of the atom6
The Quantum Mechanical Model of the Atom

  • Atomic Orbital

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

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


The quantum mechanical model of the atom7
The Quantum Mechanical Model of the Atom

  • 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


The quantum mechanical model of the atom8
The Quantum Mechanical Model of the Atom

  • 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


The quantum mechanical model of the atom9
The Quantum Mechanical Model of the Atom

  • 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)


The quantum mechanical model of the atom10
The Quantum Mechanical Model of the Atom

  • Magnetic Quantum Number

    • ml

    • ml = -l…0…l

    • related to the orbital’s orientation in space


The quantum mechanical model of the atom11
The Quantum Mechanical Model of the Atom

  • Electron Spin Quantum Number

    • ms

    • ms= + 1/2

    • an electron can spin in one of two opposite directions


The quantum mechanical model of the atom12
The Quantum Mechanical Model of the Atom

  • Orbital Shape and Energies

    • Nodal Surfaces = nodes

      • areas of zero probability of finding an electron

      • as n increases, the number of nodal surfaces increases


The quantum mechanical model of the atom13
The Quantum Mechanical Model of the Atom

  • 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


Electron spin and the pauli principle
Electron Spin and the Pauli Principle

  • 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


Polyelectronic atoms
Polyelectronic Atoms

  • 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


Polyelectronic atoms1
Polyelectronic Atoms

  • 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


Polyelectronic atoms2
Polyelectronic Atoms

  • 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


Polyelectronic atoms3
Polyelectronic Atoms

  • 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


Polyelectronic atoms4
Polyelectronic Atoms

  • 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


History of the periodic table
History of the Periodic Table

  • 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

        • recognized “triads” of elements


History of the periodic table1
History of the Periodic Table

  • 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


The aufbau principle and the periodic table
The Aufbau Principle and the Periodic Table

  • 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


The aufbau principle and the periodic table1
The Aufbau Principle and the Periodic Table

  • 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


The aufbau principle and the periodic table2
The Aufbau Principle and the Periodic Table

  • 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?


The aufbau principle and the periodic table3
The Aufbau Principle and the Periodic Table

  • Some exceptions to the Aufbau principle

    • no good explanation at this time

    • Memorize:

      • Cr [Ar]4s13d5

      • Cu [Ar]4s13d10


The aufbau principle and the periodic table4
The Aufbau Principle and the Periodic Table

  • 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,


Periodic trends
Periodic Trends

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

    • ionization energy

    • electron affinity

    • atomic size

    • metallic character


Periodic trends1
Periodic Trends

  • 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


Periodic trends2
Periodic Trends

  • 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


Periodic trends3
Periodic Trends

  • 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


Periodic trends4
Periodic Trends

  • 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 trends5
Periodic Trends

  • 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


Periodic trends6
Periodic Trends

  • 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


Periodic trends7
Periodic Trends

  • 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


Periodic trends8
Periodic Trends

  • 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)


Periodic trends9
Periodic Trends

  • 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 trends10
Periodic Trends

  • 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 trends11
Periodic Trends

  • 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


Periodic trends12
Periodic Trends

  • Atomic Radius

    • 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 trends13
Periodic Trends

  • Periodic Trend for Atomic Radii

    • Going across a row

      • atomic radii decreases

      • 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 trends14
Periodic Trends

  • Periodic Trend for Atomic Radii

    • going down a column

      • atomic radii increases

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


Periodic trends15
Periodic Trends

  • 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


Periodic trends16
Periodic Trends

  • 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 properties of a group
The Properties of a Group

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


The properties of a group1
The Properties of a Group

  • 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


The properties of a group2
The Properties of a Group

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