CHM012
Chemistry for Engineers
Topic 1:
Atomic Structure
and Periodicity
This powerpoint presentation is prepared by the faculty of the Chemistry Department, CSM, MSU-­IIT. | Page
Electromagnetic Radiation
Electromagnetic radiation (EMR) – is one of the ways that energy travel through space. The visible light is one type of EMR. Figure 1. Classification of electromagnetic spectrum.
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Electromagnetic Radiation
Characteristics of Electromagnetic radiation (EMR)
Relationship of wavelength and frequency:
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Example 1
The brilliant red colors seen in fireworks are due to the emission of light with wavelengths around 650 nm when
strontium salts such as Sr(NO3)2 and SrCO3 are heated. Calculate the frequency of red light of wavelength 6.50 x 10 2
nm.
Solution:
Practice Exercise 1:
The laser in an audio CD player uses light with a wavelength of 7.80x102 nm. Calculate the frequency of this light. (Ans. 3.84x1014 s-­‐1)
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Nature of Matter
✔ Max Planck found that energy can be gained or lost only in whole-­‐
number multiples of hv.
✔ Energy was found to be quantized, wherein a system can transfer
energy in whole quanta or “packets”. Thus energy has a particle-­‐like
properties.
where h is called Planck’s constant, 6.626x10-­‐34 Js
✔ Einstein suggested that electromagnetic radiation can be viewed as a
stream of “particles” called Photons. Where the energy of a photon is:
❖ Reading Assignment 1: Photoelectric Effect, (pp 57)
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Example 2
The blue color in fireworks is often achieved by heating copper (I) chloride (CuCl) to about 1200 oC. Then the compound
emits blue light having a wavelength of 450 nm. What is the increment of energy (the quantum) that is emitted at
4.50x102 nm by CuCl?
Solution:
Practice Exercise 2:
Microwave radiation has a wavelength on the order of 1.0 cm. Calculate the frequency and the energy of a single photon of this radiation. (Ans. 3x108 s-­‐1, 1.99x10-­‐23 J)
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Atomic Spectrum of Hydrogen
✔ Continuous spectrum (results when white light is passed through a prism) – contains all the wavelengths of visible
light
✔ Line spectrum – each line corresponds to a discrete wavelength:
Significance
✔ Only certain energies are allowed for the electron in the
hydrogen atom.
✔ Energy of the electron in the hydrogen atom is quantized.
❖ Reading Assignment 2: Atomic Spectrum of Hydrogen, (pp 61)
Figure 2. Emission and absorption spectrum of hydrogen Topic 1: Atomic Structure and Periodicity
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Assignment
Assignment 3:
1. Neutron diffraction is used in determining the structures of molecules.
a. Calculate the de Broglie wavelength of a neutron moving at 1.00% of the speed of light.
b. Calculate the velocity of a neutron with a wavelength of 75 pm (1 pm = 10-­‐12 m).
2. Why is it significant that the color emitted from the hydrogen emission spectrum is not white?
3. How does the emission spectrum support the idea of quantized energy levels?
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The Bohr Model
Assumptions
✔ Electrons in an atom can only occupy certain orbits (corresponding to certain energies).
✔ Electrons in permitted orbits have specific “allowed” energies; these energies will not be radiated from the atom.
✔ Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to
another; the energy defined by: E = hv.
Figure 3. Electronic Transitions in the Bohr Model for the Hydrogen Atom
a) An Energy-­‐Level Diagram for Electronic Transitions
b) An Orbit-­‐Transition Diagram, Which Accounts for the Experimental Spectrum
c) The resulting line spectrum on a photographic plate is shown
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The Bohr Model
✔ The energy absorbed or emitted from a single electron transition from one energy level to another:
ΔE = change in energy of the atom (energy of the emitted photon)
nfinal = integer; final distance from the nucleus
ninitial = integer; initial distance from the nucleus
Limitations of Bohr’s model
✔ It only works for hydrogen!
✔ Classical physics would result in an electron falling into the positively charged nucleus. Bohr simply assumed it would not!
✔ Circular motion is not wave-­‐like in nature.
Important ideas of Bohr’s model
✔ Electrons exist only in certain discrete energy levels.
✔ Energy is involved in the transition of an electron from one level to another.
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Example 3
Calculate the energy required to excite the hydrogen electron from level n=1 to level 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.
Solution:
Practice Exercise 3:
Calculate the energy required to remove the electron from a hydrogen atom in its ground state. (Ans. 2.178x10-­‐18 J)
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Assignment
Assignment 4:
1. What color of light is emitted when an excited electron in the hydrogen atom falls from:
a.
n = 5 to n = 2
b. n = 4 to n = 2
c.
n = 3 to n = 2 Which transition results in the longest wavelength of light?
❖ Reading Assignment 5: Quantum mechanical model of an atom, (pp 69)
2.
(a). Why does the Bohr model of hydrogen atom violate the uncertainty principle?
(b). In what way is the description of the electron using the wave function consistent with de Broglie’s hypothesis?
(c). What is meant by the term probability density?
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Quantum Numbers
❖ Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies.
❖ Each orbital describes a spatial distribution of electron density.
❖ An orbital is described by a set of three quantum numbers.
Principal Quantum Number (n)
✔ The principal quantum number, n, describes the energy level on which the orbital resides.
✔ The values of n are integers ≥ 1.
✔ These correspond to the values in the Bohr model.
Angular Momentum Quantum Number (l)
✔ This quantum number defines the shape of the orbital.
✔ Allowed values of l are integers ranging from 0 to n − 1.
✔ We use letter designations to communicate the different values of l and, therefore, the shapes and types of orbitals.
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Quantum Numbers
Magnetic Quantum Number (ml)
✔ The magnetic quantum number describes the three-­‐dimensional orientation of the orbital.
✔ Allowed values of ml are integers ranging from −l to l: −l ≤ ml ≤ l
✔ Therefore, on any given energy level, there can be up to 1 s-­‐orbital, 3 p-­‐orbitals, 5 d-­‐orbitals, 7 f-­‐orbitals,
and so forth.
❖ Orbitals with the same value of n form an electron shell.
❖ Different orbital types within a shell are subshells.
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Example 4
1. For principal quantum level n = 5, determine the number of allowed subshells (different values of l), and give the
designation of each.
Solution:
For n = 5, the allowed values of l run from 0 to 4 (n – 1 = 5 – 1). Thus the subshells and their designations are:
l=0
l=1
l=2
l=3
l=4
5s
5p
5d
5f
5g
2. For l = 2, determine the magnetic quantum numbers (ml) and the number of orbitals.
Solution:
magnetic quantum numbers = –2, – 1, 0, 1, 2
number of orbitals = 5
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s Orbital
✔The value of l for s orbitals is 0.
✔They are spherical in shape.
✔The radius of the sphere increases with the value of n.
✔For an ns orbital, the number of peaks is n.
✔For an ns orbital, the number of nodes (where there is zero probability of finding an electron) is n – 1.
✔As n increases, the electron density is more spread out and there is a greater probability of finding an electron further from the nucleus.
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p Orbital
✔The value of l for p orbitals is 1.
✔They have two lobes with a node between them.
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d Orbital
✔ The value of l for a d orbital is 2.
✔ Four of the five d orbitals have four lobes; the other resembles a p orbital with a doughnut around the center.
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f Orbital
✔ Very complicated shapes.
✔ Seven equivalent orbitals in a sublevel, l = 3
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Orbital Energy – Hydrogen Atom
✔ For a one-­‐electron hydrogen atom, orbitals on the same energy level have the same energy.
✔ Chemists call them degenerate orbitals.
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Orbital Energy – Many-­electron Atom
✔ As the number of electrons increases, so does the
repulsion between them.
✔ Therefore, in atoms with more than one electron, not
all orbitals on the same energy level are degenerate.
✔ Orbital sets in the same sublevel are still degenerate.
✔ Energy levels start to overlap in energy (e.g., 4s is
lower in energy than 3d.)
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Quantum Numbers
Spin Quantum Number (ms)
✔ In the 1920s, it was discovered that two electrons in the same orbital do not have exactly the same energy.
✔ The “spin” of an electron describes its magnetic field, which affects its energy.
✔ This led to the spin quantum number, ms.
✔ The spin quantum number has only two allowed values, +½ and –½.
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Pauli Exclusion Principle
✔No two electrons in the same atom can have exactly the same energy.
✔Therefore, no two electrons in the same atom can have identical sets of quantum numbers.
✔This means that every electron in an atom must differ by at least one of the four quantum number values: n, l, ml, and ms.
❖ Reading Assignment 6: Polyelectronic Atoms, (pp 77) and History of Periodic Table, (pp 79)
❖ Watch the video on the History of Periodic Table https://www.youtube.com/watch?v=83RSwczyyRY&feature=youtu.be&fbclid=IwAR1cC7G4LyIboEPzEYt_E5yyKwNbJCm8Tz231O
Zj4JzAiLmc_jbjIDvsVCE
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Electron Configuration
✔The way electrons are distributed in an atom is called its electron configuration.
5
4p
✔The most stable organization is the lowest possible energy, called the ground state.
✔Each component consists of ❖a number denoting the energy level;
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Electron Configuration
✔The way electrons are distributed in an atom is called its electron configuration.
5
4p
✔The most stable organization is the lowest possible energy, called the ground state.
✔Each component consists of ❖a number denoting the energy level;
❖a letter denoting the type of orbital;
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Electron Configuration
✔The way electrons are distributed in an atom is called its electron configuration.
5
4p
✔The most stable organization is the lowest possible energy, called the ground state.
✔Each component consists of ❖a number denoting the energy level;
❖a letter denoting the type of orbital;
❖a superscript denoting the number of electrons in those orbitals.
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Orbital Diagrams
✔Each box in the diagram represents one orbital.
✔Half-­‐arrows represent the electrons.
✔The direction of the arrow represents the relative spin of the electron.
Example 5
The orbital diagram of oxygen: Total no. of electrons: 8
Electron configuration, O: 1s22s22p4
Aufbau Principle (“building up”)
üFill up electrons in lowest energy orbitals
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Hund's Rule
“For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.”
✔ This means that, for a set of orbitals in the same sublevel, there must be one electron in each orbital before pairing and the electrons have the same spin, as much as possible.
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Condensed Electron Configuration
✔Elements in the same group of the periodic table have the same number of electrons in the
outer most shell. These are the valence electrons.
✔The filled inner shell electrons are called core electrons. These include completely filled d or f
sublevels.
✔We write a shortened version of an electron configuration using brackets around a noble gas
symbol and listing only valence electrons.
Li à 1s2 2s1
Li à [He] 2s1
He à 1s2
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Filling of Orbitals in the Periodic Table
✔We fill orbitals in increasing order of energy.
✔Different blocks on the periodic table correspond to different types of orbitals: s = blue, p = pink (s and p are representative elements); d = orange (transition elements); f = tan (lanthanides and actinides, or inner transition elements)
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Electron Configuration Anomalies
✔Some irregularities occur when there are enough electrons to half-­‐fill s and d orbitals on a given row.
❖For instance, the electron configuration for chromium is
[Ar] 4s1 3d5
rather than the expected
[Ar] 4s2 3d4.
❖This occurs because the 4s and 3d orbitals are very close in energy.
❖These anomalies occur in f-­‐block atoms with f and d orbitals, as well.
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Example 6
Give the electron configurations for sulfur (S) and cadmium (Cd)
Solution:
Sulfur is element 16 and resides in Period 3, where the 3p orbitals are being filled. Since, sulfur is the
fourth among the “3p elements,” it must have four 3p electrons. Its configuration is:
S: 1s22s22p63s23p4 or [Ne]3s23p4
Cadmium is element 48 and is located in Period 5 at the end of the 4d transition metals. It is the tenth
element in the series and thus has 10 electrons in the 4d orbitals, in addition to the 2 electrons in the 5s
orbital. The configuration is:
Cd: 1s22s22p63s23p64s23d104p65s24d10 or [Kr]5s24d10
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Periodicity
✔Periodicity is the repetitive pattern of a property for elements based on atomic number.
✔The following properties are discussed in this chapter:
• Sizes of atoms and ions
• Ionization energy
• Electron affinity
✔First, we will discuss a fundamental property that leads to may of the trends, effective nuclear charge.
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Effective Nuclear Charge
✔Many properties depend on attractions between valence electrons and the nucleus.
✔Electrons are both attracted to the nucleus and repelled by other electrons.
✔The forces an electron experiences depend on both factors.
✔The effective nuclear charge, Zeff, is found this way:
Zeff = Z − S
where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons.
✔Effective nuclear charge is a periodic property:
❖ It increases across a period.
❖ It decreases down a group.
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Atomic Radius
✔The bonding atomic radius is half the internuclear distance when atoms are bonded.
✔The bonding atomic radius tends to
— decrease from left to right across a period (Zeff ↑).
— increase from top to bottom of a group (n ↑).
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Ionic Radius
✔Determined by interatomic distances in ionic compounds
✔Ionic size depends on
• the nuclear charge.
• the number of electrons.
• the orbitals in which electrons reside.
✔Cations are smaller than their parent atoms:
• The outermost electron is removed and repulsions between electrons are reduced.
✔Anions are larger than their parent atoms:
• Electrons are added and repulsions between electrons are increased.
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Ionic Radius – Isoelectronic Series
✔In an isoelectronic series, ions have the same number of electrons.
✔Ionic size decreases with an increasing nuclear charge.
✔An Isoelectronic Series (10 electrons)
• Note increasing nuclear charge with decreasing ionic radius as atomic number increases
O2–
1.26 Å
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F–
1.19 Å
Na+
1.16 Å
Mg2+
0.86 Å
Al3+
0.68 Å
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Electron Configuration of Ions
Cations: The electrons are lost from the highest energy level (n value).
Example:
Li is 1s2 2s1 Li+ is 1s2 (losing a 2s electron).
Fe is 1s22s22p63s23p63d64s2
Fe2+ is 1s22s22p63s23p63d6 (losing two 4s electrons).
Anions: The electron configurations are filled to ns2np6. Example:
F is 1s22s22p5
F– is 1s22s22p6 (gaining one electron in 2p).
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Example 7
Predict the trend in radius for the following ions:
Be2+, Mg2+, Ca2+, and Sr2+
Solution: • All of these ions are formed by removing two electrons from an atom of a Group 2A element
• In going from beryllium to strontium, we are going down the group, so the sizes increase:
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Ionization Energy, I
✔The ionization energy is the minimum energy required to remove an electron from the ground state of a gaseous atom or ion.
The first ionization energy is that energy required to remove the first electron.
The second ionization energy is that energy required to remove the second electron, etc.
✔Note: the higher the ionization energy, the more difficult it is to remove an electron!
✔ It requires more energy to remove each successive electron.
✔ When all valence electrons have been removed, it takes a great deal more energy to remove the next electron.
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Ionization Energy -­ Trends
1)
2)
3)
I1 generally increases across a period.
I1 generally decreases down a group.
The s-­‐ and p-­‐block elements show a larger range of values for I1. (The d-­‐block generally increases slowly across the period; the f-­‐block elements show only small variations.)
Factors that Influences Ionization Energy
✔ Smaller atoms have higher I values.
✔ I values depend on effective nuclear charge and average distance of the electron from the nucleus.
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Ionization Energy – Irregularities ✔The trend is not followed when the added valence electron in the next element
• enters a new sublevel (higher energy sublevel);
• is the first electron to pair in one orbital of the sublevel (electron repulsions lower energy).
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Example 8
Consider atoms with the following electron configurations: 1s22s22p6
1s22s22p63s1
1s22s22p63s2 Which atom has the largest first ionization energy, and which one has the smallest second ionization energy? Explain your choices.
Solution:
• The atom with the largest value of l1 is the one with the configuration 1s22s22p6 (this is the neon atom), because this element is found at the right end of Period 2
• Since the 2p electrons do not shield each other very effectively, l1 will be relatively large
• The other configurations given include 3s electrons, which are effectively shielded by the core electrons and are farther from the nucleus than the 2p electrons in neon
• Thus l1 for these atoms will be smaller than for neon
• The atom with the smallest value of l2 is the one with the configuration 1s22s22p63s2 (the magnesium atom)
• For magnesium, both l1 and l2 involve valence electrons
• For the atom with the configuration 1s22s22p63s1 (sodium), the second electron lost (corresponding to l2) is a core electron (from a 2p orbital)
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Electron Affinity
✔ Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom:
Cl + e− ⎯⎯→ Cl−
✔ It is typically exothermic, so, for most elements, it is negative!
✔ Not much change in a group.
✔ Across a period, it generally increases. Three notable exceptions include the following:
1) Group 2A: s sublevel is full
2) Group 5A: p sublevel is half-­‐full
3) Group 8A: p sublevel is full
Note: For Group 8A the electron affinity for many of these elements is positive (X– is unstable).
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Assignment
Assignment 7:
1. Explain why the graph of ionization energy versus atomic number (across a row) is not linear. Where are the exceptions?
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Assignment
2. Which of the following would require more energy to remove an electron? Why?
Sodium vs. Chlorine
Lithium vs. Cesium
3. Which element has the larger second ionization energy? Why?
Lithium vs. Beryllium
4. Which of the following should be the larger atom? Why?
Sodium vs. Chlorine
Lithium vs. Cesium
5. Which is larger? Why?
•
The hydrogen 1s orbital •
The lithium 1s orbital 6. Which is lower in energy? Why?
•
The hydrogen 1s orbital •
The lithium 1s orbital Topic 1: Atomic Structure and Periodicity
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