LO 1.5 The student is able to explain the distribution of electrons in an atom or ion based upon data. (Sec 7.12) LO 1.6 The student is able to analyze.

LO 1.5 The student is able to explain the distribution of electrons in an atom or ion based upon data. (Sec 7.12) LO 1.6 The student is able to analyze data relating to electron energies for patterns and relationship. (Sec 7.12) LO 1.7 The student is able to describe the electron structure of the atom, using PES (photoelectron spectroscopy) data, ionization energy data, and/or Coulomb’s Law to construct explanations of how the energies of electrons within shells in atoms vary. (Sec 7.12) LO 1.9 The student is able to predict and/or justify trends in atomic properties based on location on the periodic table and/or the shell model. (Sec ) LO 1.10 Students can justify with evidence the arrangement of the periodic table and can apply periodic properties to chemical reactivity. (Sec ) LO 1.12 The student is able to explain why a given set of data suggests, or does not suggest, the need to refine the atomic model from a classical shell model with the quantum mechanical model. (Sec )

LO 1.13 Given information about a particular model of the atom, the student is able to determine if the model is consistent with specific evidence. (Sec 7.11) LO 1.15 The student can justify the selection of a particular type of spectroscopy to measure properties associated with vibrational or electronic motions of molecules. (Sec 7.1)

AP Learning Objectives, Margin Notes and References
LO 1.15 The student can justify the selection of a particular type of spectroscopy to measure properties associated with vibrational or electronic motions of molecules. Additional AP References LO 1.15 (see APEC #1, “Energy Levels and Electron Transitions”) LO 1.15 (see Appendix 7.4, “Molecular Spectroscopy: An Introduction”)

Different Colored Fireworks
Copyright © Cengage Learning. All rights reserved

Questions to Consider Why do we get colors?
Why do different chemicals give us different colors? Copyright © Cengage Learning. All rights reserved

Electromagnetic Radiation
One of the ways that energy travels through space. Three characteristics: Wavelength Frequency Speed Example- microwaves, X-rays, radiant heart Copyright © Cengage Learning. All rights reserved

Characteristics Wavelength ( ) – distance between two consecutive peaks or troughs in a wave. Frequency ( ) – number of waves (cycles) per second that pass a given point in space Speed (c) – speed of light (2.9979×108 m/s) Inverse relationship between λ and ν. Hz = sec-1 = cycles/sec Copyright © Cengage Learning. All rights reserved

The Nature of Waves

Classification of Electromagnetic Radiation
Visible 400nm-700nm Copyright © Cengage Learning. All rights reserved

Energy can be gained or lost only in whole number multiples of .
A system can transfer energy only in whole quanta (or “packets”). Energy seems to have particulate properties too. Max Plank- working with black body radiation realized this and proposed quantized energy. Copyright © Cengage Learning. All rights reserved

Electromagnetic radiation is a stream of “particles” called photons.
Energy is quantized. Electromagnetic radiation is a stream of “particles” called photons. Planck’s constant = h = × Js Einstein analyzed photoelectric effect and theorized that all electromagnetic radiation itself was quantized Copyright © Cengage Learning. All rights reserved

The Photoelectric effect
To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE Copyright © Cengage Learning. All rights reserved

Refers to a phenomenon in which electrons are emitted from the surface of a metal when light strikes it. Characterized by 1. Studies in which the frequency of light is varied show that no electrons are emitted by a given metal below a specific threshold frequency regardless of the intensity (number of photons available) of the light. 2. For light with frequency higher than the threshold frequency, the number of electrons emitted increases with the intensity of the light. Copyright © Cengage Learning. All rights reserved

3. For light with frequency greater than the threshold frequency, the kinetic energy of the emitted electrons increases linearly with the frequency of light. Extra energy above the threshold translates as KE for the electron. KE elec=1/2mv2 = hv –hv0 Copyright © Cengage Learning. All rights reserved

Energy has mass Energy is quantized. It can only occur in discrete units called quanta. Dual nature of light: Electromagnetic radiation (and all matter) exhibits wave properties and particulate properties. E = mc2 Energy has mass. Mass of a photon = h/λc (relative mass, nit rest mass) Copyright © Cengage Learning. All rights reserved

The Wave Nature of Matter
Louis de Broglie suggested that if light can have material properties, matter should exhibit wave properties. He demonstrated that the relationship between mass and wavelength was  = h mv Copyright © Cengage Learning. All rights reserved

The Wave Nature of Matter
All matter exhibits both particulate and wave properties. Large objects exhibit predominately particulate properties, very small matter (photons) exhibit more wave like properties, matter with intermediate mass (electrons) clearly exhibit both properties. Copyright © Cengage Learning. All rights reserved

Line spectrum – each line corresponds to a discrete wavelength:
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: Hydrogen emission spectrum Copyright © Cengage Learning. All rights reserved

Refraction of White Light

The Line Spectrum of Hydrogen

Significance Only certain energies are allowed for the electron in the hydrogen atom. Energy of the electron in the hydrogen atom is quantized. Energy levels are quantized. If any energy level was allowed, then you would have a continuous spectrum ΔE =h ν = hc/λ Copyright © Cengage Learning. All rights reserved

CONCEPT CHECK! Why is it significant that the color emitted from the hydrogen emission spectrum is not white? How does the emission spectrum support the idea of quantized energy levels? If the levels were not quantized, we’d probably see white light. This is because all possible value of energy could be released, meaning all possible colors would be emitted. All the colors combined make white light. Since only certain colors are observed, this means that only certain energy levels are allowed. An electron can exist at one level or another, and there are regions of zero probability in between. Copyright © Cengage Learning. All rights reserved

LO 1.12 The student is able to explain why a given set of data suggests, or does not suggest, the need to refine the atomic model from a classical shell model with the quantum mechanical model.

Ground state – lowest possible energy state (n = 1)
Electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits. Bohr’s model gave hydrogen atom energy levels consistent with the hydrogen emission spectrum. Ground state – lowest possible energy state (n = 1) Copyright © Cengage Learning. All rights reserved

Electronic Transitions in the Bohr Model for the Hydrogen Atom b) An Orbit-Transition Diagram, Which Accounts for the Experimental Spectrum Lines in the visible spectrum correspond to transitions from higher levels to n=2 level. Copyright © Cengage Learning. All rights reserved

For 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 Copyright © Cengage Learning. All rights reserved

The model correctly fits the quantized energy levels of the hydrogen atom and postulates only certain allowed circular orbits for the electron. As the electron becomes more tightly bound, its energy becomes more negative relative to the zero-energy reference state (free electron). As the electron is brought closer to the nucleus, energy is released from the system. Ground state is the lowest possible state vs. excited state Copyright © Cengage Learning. All rights reserved

Bohr’s model is incorrect. This model only works for hydrogen.
Electrons do not move around the nucleus in circular orbits. Copyright © Cengage Learning. All rights reserved

Which transition results in the longest wavelength of light?
EXERCISE! What color of light is emitted when an excited electron in the hydrogen atom falls from: n = 5 to n = 2 n = 4 to n = 2 n = 3 to n = 2 Which transition results in the longest wavelength of light? blue, λ = 434 nm green, λ = 486 nm orange/red, λ = 657 nm For each transition, use ΔE = hc / λ = (–2.178×10–18)[(1/nf) – (1/ni)]. Solve for λ in each case. a) blue (λ = 434 nm) b) green (λ = 486 nm) c) orange/red (λ = 657 nm) The longest wavelength of light is from transition n = 3 to n = 2 (letter c). Copyright © Cengage Learning. All rights reserved

LO 1.12 The student is able to explain why a given set of data suggests, or does not suggest, the need to refine the atomic model from a classical shell model with the quantum mechanical model.

Quantum Mechanics Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated. It is known as quantum mechanics. Figure 7.11 | The hydrogen electron visualized as a standing wave around the nucleus. The circumference of a particular circular orbit would have to correspond to a whole number of wavelengths, as shown in (a) and (b), or else destructive interference occurs, as shown in (c). This is consistent with the fact that only certain electron energies are allowed; the atom is quantized. (Although this idea encouraged scientists to use a wave theory, it does not mean that the electron really travels in circular orbits.)

We do not know the detailed pathway of an electron.
Heisenberg uncertainty principle: There is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time. Δx = uncertainty in a particle’s position Δ(mν) = uncertainty in a particle’s momentum h = Planck’s constant Copyright © Cengage Learning. All rights reserved

Physical Meaning of a Wave Function (Ψ)
The wave equation is designated with a lower case Greek psi (). The square of the wave equation, 2, gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time. (Orbital) Ψ is a function of the electron position in 3D space.ψ2 gives rise to atomic orbitals (tells only the probability of location, not how it got there.) Copyright © Cengage Learning. All rights reserved

Probability Distribution for the 1s Wave Function (ψ2 orbital)
The maximum point in the curve is a result of 2 opposing effects. An increase probability of finding the electron closer to the nucleus. The volume of the sphere increases as you move away from the nucleus so the probability of finding the electron decreases, but summing more positions. Copyright © Cengage Learning. All rights reserved

Relative Orbital Size Difficult to define precisely.
Orbital is a wave function. Picture an orbital as a three-dimensional electron density map. Hydrogen 1s orbital: Radius of the sphere that encloses 90% of the total electron probability. (probability never becomes zero as head out of the nucleus) Copyright © Cengage Learning. All rights reserved

Each orbital describes a spatial distribution of electron density.
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. Quantum numbers describe orbital properties

Principal Quantum Number, n
The principal quantum number, n, describes the energy level on which the orbital resides. The values of n are integers ≥ 0. As n increases, the orbital becomes larger, and the electron is further away from the nucleus, and has higher energy.

Azimuthal 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.

Azimuthal Quantum Number, l
Value of l 1 2 3 Type of orbital s p d f

Magnetic Quantum Number, ml
Describes the three-dimensional orientation of the orbital. Values 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, etc.

Quantum Numbers Orbitals with the same value of n form a shell.
Different orbital types within a shell are subshells.

A shell with principal quantum number n will consist of exactly n subshells. (Ex. n = 2 consists of 2 subshells 2s and 2p. For a given subshell l, there are 2l + 1 orbitals. (Ex. Each p (l = 1) subshell consists of three orbitals. The total number of orbitals in a shell is n2.

EXERCISE! For principal quantum level n = 3, determine the number of allowed subshells (different values of l), and give the designation of each. # of allowed subshells = 3 l = 0, 3s l = 1, 3p l = 2, 3d The allowed values of l run from 0 to 2, so the number of allowed subshells is 3. Thus the subshells and their designations are: l = 0, 3s l = 1, 3p l = 2, 3d Copyright © Cengage Learning. All rights reserved

EXERCISE! For l = 2, determine the magnetic quantum numbers (ml) and the number of orbitals. magnetic quantum numbers = –2, – 1, 0, 1, 2 number of orbitals = 5 The magnetic quantum numbers are -2, -1, 0, 1, 2. The number of orbitals is 5. Copyright © Cengage Learning. All rights reserved

1s Orbital To play movie you must be in Slide Show Mode
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Three Representations of the Hydrogen 1s, 2s, and 3s Orbitals
Contour drawings show where 90% probability Copyright © Cengage Learning. All rights reserved

2px Orbital To play movie you must be in Slide Show Mode

2py Orbital To play movie you must be in Slide Show Mode

2pz Orbital To play movie you must be in Slide Show Mode

The Boundary Surface Representations of All Three 2p Orbitals
+ and – are math signs not charges of the orbitals. They indicate phases. Copyright © Cengage Learning. All rights reserved

3dx2-y2 Orbital To play movie you must be in Slide Show Mode

3dxy Orbital To play movie you must be in Slide Show Mode

3dxz Orbital To play movie you must be in Slide Show Mode

3dyz Orbital To play movie you must be in Slide Show Mode

3dz2 To play movie you must be in Slide Show Mode

The Boundary Surfaces of All of the 3d Orbitals
Dxy and dx2-y2 are both in the xy palne. Xy is b/n axis and x2-y2 is on the axes. Copyright © Cengage Learning. All rights reserved

Representation of the 4f Orbitals in Terms of Their Boundary Surfaces
Very complex and will not discuss any participation in bonding in this course. Copyright © Cengage Learning. All rights reserved

Energies of Orbitals For a one-electron hydrogen atom, orbitals on the same energy level have the same energy. That is, they are degenerate. Ground state is 1s. Any higher is considered excited state. Copyright © Cengage Learning. All rights reserved

Electron Spin Electron spin quantum number (ms) – can be +½ or -½.
Pauli exclusion principle - in a given atom no two electrons can have the same set of four quantum numbers. An orbital can hold only two electrons, and they must have opposite spins. Spectral data indicated that e- has a magnetic moment with 2 possible orientations when the atom is placed in a magnetic field. Copyright © Cengage Learning. All rights reserved

Spin Quantum Number, ms This led to a fourth quantum number, the spin quantum number, ms. The spin quantum number has only 2 allowed values: +1/2 and −1/2. Copyright © Cengage Learning. All rights reserved

Atoms with more than one electron. Electron correlation problem:
Since the electron pathways are unknown, the electron repulsions cannot be calculated exactly, so we make approximations. Treat each electron as if it were moving in a field of charge that is the net result of the nuclear attraction and the average repulsion of all the other electrons. Copyright © Cengage Learning. All rights reserved

The net effect is that the electron is not bound nearly as tightly to the nucleus as it would be if the other electrons were not present. The valence electrons are screened or shielded from the nuclear charge by the repulsions of the other electrons. Copyright © Cengage Learning. All rights reserved

Penetration Effect A 2s electron penetrates to the nucleus more than one in the 2p orbital. This causes an electron in a 2s orbital to be attracted to the nucleus more strongly than an electron in a 2p orbital. Thus, the 2s orbital is lower in energy than the 2p orbitals in a polyelectronic atom. Orbital in polyelectron atom are the same shape, but different energies than the H orbitals b/c of this Copyright © Cengage Learning. All rights reserved

Orbital Energies To play movie you must be in Slide Show Mode

A Comparison of the Radial Probability Distributions of the 2s and 2p Orbitals
2s orbital e- penetrates the shielding and gets closer to the nuclear charge, therefore it has lower energy of the orbital. Copyright © Cengage Learning. All rights reserved

The Radial Probability Distribution of the 3s Orbital

A Comparison of the Radial Probability Distributions of the 3s, 3p, and 3d Orbitals
When electrons are placed in a particular quantum level, they “prefer” the orbitals in the order s, p, d, and then f. Copyright © Cengage Learning. All rights reserved

Energies of Orbitals As the number of electrons increases, though, so does the repulsion between them. Therefore, in many-electron atoms, orbitals on the same energy level are no longer degenerate. Copyright © Cengage Learning. All rights reserved

LO 1.9 The student is able to predict and/or justify trends in atomic properties based on location on the periodic table and/or the shell model. LO 1.10 Students can justify with evidence the arrangement of the periodic table and can apply periodic properties to chemical reactivity. LO 1.13 Given information about a particular model of the atom, the student is able to determine if the model is consistent with specific evidence.

Aufbau Principle As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to hydrogen-like orbitals. An oxygen atom has an electron arrangement of two electrons in the 1s subshell, two electrons in the 2s subshell, and four electrons in the 2p subshell. Oxygen: 1s22s22p4 Aufbau- German for building up Electron configuration vs. Orbital diagram Copyright © Cengage Learning. All rights reserved

Hund’s Rule The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate (same energy) orbitals. Copyright © Cengage Learning. All rights reserved

Orbital Diagram A notation that shows how many electrons an atom has in each of its occupied electron orbitals. Oxygen: 1s22s22p4 Oxygen: 1s 2s p Paramagnetic – attracted to a magnet (parallel spin) Diamagnetic – repels a magnet (all paired electrons) Copyright © Cengage Learning. All rights reserved

Valence Electrons The electrons in the outermost principal quantum level of an atom. 1s22s22p6 (valence electrons = 8) The elements in the same group on the periodic table have the same valence electron configuration. Valence vs. core e- Copyright © Cengage Learning. All rights reserved

The Orbitals Being Filled for Elements in Various Parts of the Periodic Table
Ns fills before n-1d b/c of penetration to the vicinity of the nucleus so the s orbital is lower in energy than the d. 6s25d14f……. Copyright © Cengage Learning. All rights reserved

EXERCISE! Determine the expected electron configurations for each of the following. a) S 1s22s22p63s23p4 or [Ne]3s23p4 b) Ba [Xe]6s2 c) Eu [Xe]6s24f7 a) 16 electrons total; 1s22s22p63s23p4 or [Ne]3s23p4 b) 56 electrons total; [Xe]6s2 c) 63 electrons total; [Xe]6s24f7 Copyright © Cengage Learning. All rights reserved

Some Anomalies For instance, the electron configuration for copper is
[Ar] 4s1 3d5 rather than the expected [Ar] 4s2 3d4. Copyright © Cengage Learning. All rights reserved

Electron Configurations of Ions
When electrons are added to an atom to form an ion, they are added to the empty or partially filled orbital with the lowest value for n. F ([He]2s22p5) → F- ([He]2s22p6) When electrons are removed from an atom to form a cation, they are always removed from the occupied orbital with the largest value of n. Fe ([Ar]3d64s2) → Fe2+ ([Ar]3d6) Copyright © Cengage Learning. All rights reserved

Write the electron configuration for (a) Ca2+, (b) Co3+, and (c) S2–.
PRACTICE EXERCISE Write the electron configuration for (a) Ga3+, (b) Cr3+, and (c) Br–. Copyright © Cengage Learning. All rights reserved

LO 1.5 The student is able to explain the distribution of electrons in an atom or ion based upon data. LO 1.6 The student is able to analyze data relating to electron energies for patterns and relationship. LO 1.7 The student is able to describe the electron structure of the atom, using PES (photoelectron spectroscopy) data, ionization energy data, and/or Coulomb’s Law to construct explanations of how the energies of electrons within shells in atoms vary. LO 1.9 The student is able to predict and/or justify trends in atomic properties based on location on the periodic table and/or the shell model. LO 1.10 Students can justify with evidence the arrangement of the periodic table and can apply periodic properties to chemical reactivity.

Periodic Trends Ionization Energy Electron Affinity Atomic Radius
Use quantum mechanical model to explain trends

Effective Nuclear Charge
Properties of atoms depend not only on their electron configuration, but also on how strongly their outer electrons are attracted to the nucleus. The force of attraction between an electron and the nucleus depends on: 1. The magnitude of the net nuclear charge 2. The average distance between the outer electrons and the nucleus

In a many-electron atom, the nuclear charge that an electron experiences depends on both the attraction of the nucleus and the repulsion of the other electrons. Since there are so many electron-electron repulsions, we estimate the net attraction of each electron to the nucleus by treating each electron as though it were moving a net electric field created by the nucleus and the electron density of the other electrons.

The net electric field is viewed as a single positive charge located in the nucleus called the effective nuclear charge (Zeff). The effective nuclear charge acting on an electron is smaller that the actual nuclear charge because it takes into account the electron repulsions.

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 (core)electrons because the core electrons are most effective at screening the valence electrons from the nucleus. Electrons in the same valence shell only screen slightly.

The 3s electron has some probability of being inside the Ne core. As a consequence of this "penetration," the 3s electron will experience a greater attraction from the nucleus than expected. Thus the effective nuclear charge experienced by the 3s electron is actually 2.5+.

Because the 2s probability function has a peak closer to the nucleus, the 2s orbital is less effectively screened by the core orbitals and therefore the 2s orbital has a higher effective nuclear charge than the 2p orbital. This greater attraction to the nucleus leads to a lower energy for the 2s orbital than the 2p orbital.

Trends in Zeff Row – The effective nuclear charge increases across a row because the core electrons stays the same, but the actual nuclear charge increases. Column – The effective nuclear charge increases slightly because the larger electron cores are less able to screen the outer electrons from the nuclear charge. This increase is considered insignificant.

Ionization Energy Energy required to remove an electron from a gaseous atom or ion. X(g) → X+(g) + e– Mg → Mg+ + e– I1 = 735 kJ/mol (1st IE) Mg+ → Mg2+ + e– I2 = 1445 kJ/mol (2nd IE) Mg2+ → Mg3+ + e– I3 = 7730 kJ/mol *(3rd IE) *Core electrons are bound much more tightly than valence electrons.

Ionization Energy In general, as we go across a period from left to right, the first ionization energy increases. Why? Electrons added in the same principal quantum level do not completely shield the increasing nuclear charge caused by the added protons. Electrons in the same principal quantum level are generally more strongly bound from left to right on the periodic table due to a stronger Zeff.

Ionization Energy In general, as we go down a group from top to bottom, the first ionization energy decreases. Why? The electrons being removed are, on average, farther from the nucleus. As n increases, size of orbital increases

The Values of First Ionization Energy for the Elements in the First Six Periods

However, there are two apparent discontinuities in this trend.

The first occurs between Groups 2 and 13.
Electron removed from p-orbital rather than s-orbital Electron farther from nucleus Small amount of repulsion by s electrons.

The second occurs between Groups 15 and 16.
Electron removed comes from doubly occupied orbital. Repulsion from the other electron in orbital helps in its removal.

Successive Ionization Energies (KJ per Mole) for the Elements in Period 3

CONCEPT CHECK! Explain why the graph of ionization energy versus atomic number (across a row) is not linear. electron repulsions Where are the exceptions? some include from Be to B and N to O The graph is not linear due to electron repulsions. Some exceptions include from Be to B and N to O.

Which atom would require more energy to remove an electron? Why? Na Cl
CONCEPT CHECK! Which atom would require more energy to remove an electron? Why? Na Cl Cl would require more energy to remove an electron because the electron is more tightly bound due to the increase in effective nuclear charge.

Which atom would require more energy to remove an electron? Why? Li Cs
CONCEPT CHECK! Which atom would require more energy to remove an electron? Why? Li Cs Li would require more energy to remove an electron because the outer electron is on average closer to the nucleus (so more tightly bound).

CONCEPT CHECK! Which has the larger second ionization energy? Why? Lithium or Beryllium Lithium has the larger second ionization energy because then a core electron is trying to be removed which will require a lot more energy than a valence electron.

CONCEPT CHECK! Three elements are indicated in the periodic table below. Based on their locations, predict the one with the largest second ionization energy. Red

Electron Affinity Energy change associated with the addition of an electron to a gaseous atom. X(g) + e– → X–(g) *I1 measures the ease with which an atom loses an electron, whereas electron affinity measures the ease with which an atom gains an electron. NOT in a molecule

Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row. Halogens have the most negative electron affinity because they are 1 electron shy of a filled p subshell. F has a particularly higher (less negative) electron affinity because of the smaller size of its 2p orbitals.

Electron Affinity The electron affinity for the noble gases are positive because any additional electron would have to be placed in an empty higher-energy subshell.

Electron Affinity There are again, however, two discontinuities in this trend.

Electron Affinity The first occurs between Groups 1 and 2.
Added electron must go in the previously empty, higher-energy p-orbital, not s-orbital. Electron is farther from nucleus and feels repulsion from s-electrons. Be and Mg do not add an electron exothermically and thus do not form a stable isolated X- ion.

Electron Affinity The second occurs between Groups 14 and 15.
Group 15 has half-filled p subshells. Extra electron must go into occupied orbital, creating repulsion. N does not add an electron exothermically and thus does not form a stable isolated X- ion.

Electron Affinity Electron affinities do not change greatly you move down a group. As the average distance between the added electron and the nucleus increases, the electron-nucleus attraction decreases. The orbital holding the added electron is larger, therefore there is less electron repulsion.

Covalent Atomic Radius
The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei. The bonding radius is always smaller than the nonbonding radius due to electron cloud overlap. Can not measure the specific size of an atom b/c no defined barrier

Atomic Radius In general as we go across a period from left to right, the atomic radius decreases. Effective nuclear charge increases, therefore the valence electrons are drawn closer to the nucleus, decreasing the size of the atom. In general atomic radius increases in going down a group. Orbital sizes increase in successive principal quantum levels.

Atomic Radii for Selected Atoms
Figure 7.35 | Atomic radii (in picometers) for selected atoms. Note that atomic radius decreases going across a period and increases going down a group. The values for the noble gases are estimated, because data from bonded atoms are lacking.

Which should be the larger atom? Why? Na Cl
CONCEPT CHECK! Which should be the larger atom? Why? Na Cl Na should be the larger atom because the electrons are not bound as tightly due to a smaller effective nuclear charge.

Which should be the larger atom? Why? Li Cs
CONCEPT CHECK! Which should be the larger atom? Why? Li Cs Cs should be the larger atom because of the increase in orbital sizes in successive principal quantum levels (to accommodate more electrons).

CONCEPT CHECK! Natural gas used in home heating and cooking is odorless. Because natural gas leaks pose the danger of explosion or suffocation, various smelly substances are added to the gas to allow detection of a leak. One such substance is methyl mercaptan, CH3SH, whose structure is shown below. Use Figure 7.35 to predict the lengths of the C—S, C—H, and S—H bonds in this molecule. Cs should be the larger atom because of the increase in orbital sizes in successive principal quantum levels (to accommodate more electrons).

Which is lower in energy?
CONCEPT CHECK! Which is larger? The hydrogen 1s orbital The lithium 1s orbital Which is lower in energy? The hydrogen 1s orbital is larger because the electrons are not as tightly bound as the lithium 1s orbital (lithium has a higher effective nuclear charge and will thus draw in the inner electrons more closely). The lithium 1s orbital is lower in energy because the electrons are closer to the nucleus.

Atomic Radius of a Metal
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Atomic Radius of a Nonmetal
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EXERCISE! Arrange the elements oxygen, fluorine, and sulfur according to increasing: Ionization energy S, O, F Atomic size F, O, S Ionization Energy: S, O, F (IE increases as you move up a column and to the right across a period.) Atomic Size: F, O, S (Atomic radius increases as you move to the left across a period and down a column.)

LO 1.9 The student is able to predict and/or justify trends in atomic properties based on location on the periodic table and/or the shell model. LO 1.10 Students can justify with evidence the arrangement of the periodic table and can apply periodic properties to chemical reactivity.

The Periodic Table – Final Thoughts
It is the number and type of valence electrons that primarily determine an atom’s chemistry. Electron configurations can be determined from the organization of the periodic table. Certain groups in the periodic table have special names. Copyright © Cengage Learning. All rights reserved

Special Names for Groups in the Periodic Table

Metals Versus Nonmetals
Nonmetals – large Ie and large Ea, tend to gain e- to form anions when reacting with a metal. Most reactive in the upper right corner. Metals – Low Ie, Low Ea, tend to lose e- and form cations. Most reactive metals in the lowers left corner due to very low Ie. Copyright © Cengage Learning. All rights reserved

The Alkali Metals Li, Na, K, Rb, Cs, and Fr
Most chemically reactive of the metals React with nonmetals to form ionic solids Going down group: Ionization energy decreases Atomic radius increases Density increases (mass increases faster than size) Melting and boiling points smoothly decrease H acts like a nonmetal b/c very small size so the e- is very tightly bound. Copyright © Cengage Learning. All rights reserved

The Alkali Metals Alkali metals easily reduce nonmetals due to their low ionization energy. (Cs>Rb>K>Na>Li) 2Na(s) + S(s) → Na2S(s) Reduce water also, but the trend in reducing ablilty is the reverse due to lithium’s highly exothermic hydration energy. (Li>K>Na) 2M(s) + 2H2O(l) → H2(g) + 2M+(aq) + 2OH-(aq) +energy Lithium is so small it has a large charge to mass ratio so highly exothermic hydration energy. Surprisingly Li does not react as vigorously with water as Na and K because of it’s high melting point. The others melt when reacting with water which increases their surface area and causes a more violent reaction. Copyright © Cengage Learning. All rights reserved

Produce bright colors when placed in flame.

LO 1.5 The student is able to explain the distribution of electrons in an atom or ion based upon data. (Sec 7.12) LO 1.6 The student is able to analyze.

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LO 1.5 The student is able to explain the distribution of electrons in an atom or ion based upon data. (Sec 7.12) LO 1.6 The student is able to analyze.

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