Electromagnetic Radiation ATOMIC STRUCTURE • Radiant energy that exhibits and wavelength-like behavior and travels through space at the speed Periodicity of light in a vacuum.

Figure 7.5: The electromagnetic Waves spectrum. • Waves have 3 primary characteristics: • 1. Wavelength: distance between two peaks in a wave. Symbol is 8 • 2. Frequency: number of waves per second that pass a given point in space. Symbol is < • 3. Speed: speed of light is 2.9979´108 m/s. Symbol is c 8< = c

Figure 7.4: Relation between Figure 7.3: Water Wave (ripple). wavelength and frequency.

1 Wavelength and frequency can be interconverted. The Wave Nature of Light • What is the wavelength of yellow light with a frequency of 5.09 x 1014 s-1? (Note: s-1, commonly •n = c/l referred to as Hertz (Hz) is defined as “cycles or waves per second”.)

• n = frequency (s-1) – If c = nl, then rearranging, we obtain l = c/n

• l = wavelength (m) (3.00 ´ 108 m /s) -7 l = 5.09 ´ 1014 s -1 = 5.89´10 m or 589 nm • c = speed of light (m s-1)

The Wave Nature of Light Quantum Effects and Photons • What is the frequency of violet light with a • By the early part of twentieth century, the wave wavelength of 408 nm? (see Figure 7.5) theory of light seemed to be well entrenched. • In 1905, Albert Einstein proposed that light had – If c = nl, then rearranging, we obtain n = c/l. both wave and particle properties as observed in the photoelectric effect. (see Figure 7.6) 3.00 ´ 108 m /s 14 -1 n = = 7.35´10 s • Einstein based this idea on the work of a German 408 ´ 10-9 m physicist, Max Planck.

B Quantum Effects and Photons

1 Planck’s Quantization of Energy

In order to explain the observation E = n h? that the ?’s of light emitted from a hot solid was dependent on the where n is the quantum number, h is temperature of the solid, Planck observed that the atoms in the solid Planck’s constant, and ? is the frequency of vibrated with a definite ?, depending vibration. Thus, the energy of the atoms is on the solid, and that the atom could only have certain energies of emitted in “packets” or quantized. vibration, E.

2 Planck’s Constant Quantum Effects and Photons Transfer of energy is quantized, and can only occur in discrete units, called quanta. • Photoelectric Effect hc – The photoelectric effect is the ejection of DE = hn = electrons from the surface of a metal when light l shines on it. (see Figure 7.6) • DE = change in energy, in J – Electrons are ejected only if the light exceeds a certain “threshold” frequency. -34 • h = Planck’s constant, 6.626 ´ 10 J s – Violet light, for example, will cause potassium • n = frequency, in s -1 to eject electrons, but no amount of red light • l = wavelength, in m (which has a lower frequency) has any effect.

Quantum Effects and Photons

• Photoelectric Effect Figure 7.6: – Einstein’s assumption that an electron is ejected The when struck by a single photon implies that it photoelectric behaves like a particle. effect. – When the photon hits the metal, its energy, hn is taken up by the electron. – The photon ceases to exist as a particle; it is said to be “absorbed.”

Figure 7.4: Electromagnetic Quantum Effects and Photons radiation exhibits wave properties • Photoelectric Effect and particulate properties. – The “wave” and “particle” pictures of light should be regarded as complementary views of the same physical entity. – This is called the wave -particle duality of light. – The equation E = hn displays this duality; E is the energy of the “particle” photon, and n is the frequency of the associated “wave .”

3 Energy and Mass Energy and Mass • Energy has mass hc E = • E = mc2 photon l • E = energy h • m = mass m = photon lc • c = speed of light (Hence the dual nature of light.)

“They could but make the best of it, and went around with Figure 7.9: Dispersion of white woebegone faces sadly complaining that on Mondays, Wednesdays, and Fridays they must look on light as a wave; on light by a prism. Tuesdays, Thursdays and Saturdays as a particle. On Sundays they simply prayed.” - Banesh Hoffman, The Strange Story of the Quantum, Dover Publications, Inc. New York 1959.

Figure 7.6: (a) A continuous spectrum Atomic Spectrum of Hydrogen containing all wavelengths of visible light. • Continuous spectrum: Contains all the (b) The wavelengths of light. hydrogen line spectrum • Line (discrete) spectrum: Contains only contains only a some of the wavelengths of light. few discrete wavelengths.

4 When a strontium salt is dissolved in methanol (with a little water) and ignited, it gives a brilliant red flame. The red color is produced by Figure emission of light when electrons, excited by the energy of the burning 7.2: methanol, fall back to their ground states. Emission (line) spectra of some elements

If we pass light emitted from a heated gas, we obtain a line spectrum which contains only a few lines (or colors) or specific The Bohr Model wavelengths characteristic of the gas. In 1885, J.J. Balmer showed that for the H atom, the observed visible ?’s were given • Ground State: The lowest by the equation: possible energy state for an atom ?’ = 1/? = 1.097 x 107/m [(1/22)-(1/n2)] Where n is an integer > 2 (n = 1). The constant, 1.097 x 107, is called the Rydberg constant (R) and was determined experimentally.

2 Bohr’s postulates - developed to account Figure 7.8: for a) the stability of the H atom (e does Electronic not fall into the nucleus) and b) the transitions in observed line spectrum the Bohr a Energy Level Postulate - an electron model for the can have only specific energy values hydrogen in an atom, called energy levels. atom.

2 For the H atom: E = -(RH/n ) where RH is a constant with energy units and n is the principal quantum number

5 Figure 7.10: Energy -level b Transitions between energy levels - an diagram for the electron electron in an atom can change energy in the hydrogen atom. only by going from one energy level to another energy level. When it changes R = 2B2:e4/h 2 H energy, it undergoes a transition and Ef where : is the reduced - Ei = h? mass of system (e & p), e is the charge on an Bohr demonstrated that : electron, and h is 7 Planck’s constant, RH /hc = R = 1.097 x 10 /m 6.6256 x 10-27 g-cm2/s. If from Balmer’s formula. According to If we recall that 8< = c Bohr’s theory, light is emitted from a H

and E = h<, RH/hc = R atom when its electrons drops from a higher to lower energy level. See Fig 7.11

The Bohr Theory of the Hydrogen Atom Figure 7.11: • Bohr’s Postulates Transitions of – Bohr derived the following formula for the energy levels the electron in of the electron in the hydrogen atom.

the hydrogen Rh E = - 2 n = 1, 2, 3 ..... ¥ (for H atom) atom. n

– RH is a constant (expressed in energy units) with a value of 2.18 x 10-18 J.

The Bohr Theory of the The Bohr Theory of the Hydrogen Atom Hydrogen Atom • Bohr’s Postulates • Bohr’s Postulates – When an electron undergoes a transition from a higher – If we make a substitution into the previous equation that energy level to a lower one, the energy is emitted as a states the energy of the emitted photon, hn, equals E - E , photon. i f

Rh Rh Energy of emitted photon = hn = Ei - Ef hn = Ei - Ef = (- 2 ) - (- 2 ) ni nf – From Postulate 1, Rearranging, we obtain R R h E = - h 1 1 Ei = - 2 f 2 hn = R ( - ) n h 2 2 ni f nf ni

6 Figure 7.7: A change between A Problem to Consider two discrete energy levels emits a • Calculate the energy of a photon of light emitted photon of light. from a hydrogen atom when an electron falls from Recall that level n = 3 to level n = 1. E = h< 1 1 E = hn = Rh( 2 - 2) and that nf ni < = c/8 -18 1 1 E = (2.18´10 J)(1 2 - 3 2 ) E = 1.94 ´10-18J

Energy Changes in the Hydrogen The Bohr Theory of the Atom Hydrogen Atom • Bohr’s Postulates – Bohr’s theory explains not only the emission of •DE = E - E light, but also the absorption of light. final state initial state – When an electron falls from n = 3 to n = 2 energy level, a photon of red light (wavelength, hc 685 nm) is emitted. l = – When red light of this same wavelength shines DE on a hydrogen atom in the n = 2 level, the energy is gained by the electron that undergoes a transition to n = 3.

“A curious thing about the atom-model work of Bohr, prior to 1923 or 1924, was that if you look at the then-current papers Quantum Mechanics you get the impression that everybody in the world was terrif- ically excited about the Bohr model and believed in it hook, • Bohr’s theory established the concept of atomic line, and sinker, including the electron orbits as they are energy levels but did not thoroughly explain the used in the ads for the atomic age nowadays. Bohr, on the “wave-like” behavior of the electron. other hand, was constantly making remarks, speeches, and admonitions to the effect that this is temporary and we ought to • Current ideas about atomic structure depend on be looking for a way to do it right.” the principles of quantum mechanics , a theory - Edward U. Condon, to the Philosophical Society of Washington, that applies to subatomic particles such as December 2, 1960. electrons.

7 Quantum Mechanics Quantum Mechanics

• Bohr’s theory established the concept of atomic • The first clue in the development of quantum energy levels but did not thoroughly explain the theory came with the discovery of the de Broglie “wave-like” behavior of the electron. relation. – In 1923, Louis de Broglie reasoned that if light • Current ideas about atomic structure depend on exhibits particle aspects, perhaps particles of the principles of quantum mechanics , a theory matter show characteristics of waves. that applies to subatomic particles such as – He postulated that a particle with mass m and a electrons. velocity v has an associated wavelength. – The equation l = h/mv is called the de Broglie relation.

Quantum Mechanics Quantum Mechanics

• If matter has wave properties, why are they • If matter has wave properties, why are they not commonly observed? not commonly observed? For most objects, – The de Broglie relation shows that a baseball the wavelength is outside emr spectrum! (0.145 kg) moving at about 60 mph (27 m/s) – Electrons have wavelengths on the order of a few has a wavelength of about 1.7 x 10-34 m. picometers (1 pm = 10-12 m). X-ray region of emr -34 kg×m2 spectrum 6.63 ´ 10 s -34 l = (0.145 kg)(27 m /s) = 1.7 ´10 m – Under the proper circumstances, the wave – This value is so incredibly small that such character of electrons should be observable. waves cannot be detected.

The Wavelength of an Electron

What is the wavelength of an electron moving at 6 x 106 m/s? Quantum Mechanics • Heisenberg’s uncertainty principle is a relation that states that the product of the uncertainty in position Using deBroglie’s equation: (Dx) and the uncertainty in momentum (mDvx) of a particle is a minimum of h/4p. 8 = h/mv = (6.63 x 10-34J-s)/[(9.11x10-31kg)(6 x 106 m/s)] = 1.21 x 10-10 m or 0.121 nm h (Dx)(mDvx ) ³ Remember that 1 J = 1 kg-m2/s2 4p What is the energy associated with this electron? – When m is large (for example, a baseball) the uncertainties are small, but for electrons, high -34 8 -10 E = h< = hc/8 = (6.63 x 10 J-s)(3.0 x 10 m/s)/(1.21 x 10 m ) uncertainties make it impossible to predict an exact = 1.64 x 10-15 J orbit.

8 Heisenberg Uncertainty Principle The Meaning of the Heisenberg Uncertainty Principle h We know that the electron has a mass of 9.11 x 10-31kg and moves Dx × D(mv) ³ with an average speed of 5 x106 m/s in the H atom. The velocity is 4p the only uncertainty in the momentum so we can calculate the uncertainty in the electrons position in the H atom using Heisenberg’s Principle: • x = position ) B ) -34 B -31 6 • mv= momentum x = h/(4 m v) = (6.63x10 J-s)/[4 (9.11 x 10 kg )(5 x10 m/s)] = 1 x 10-9 m (or 1 nm) • h = Planck’s constant • The more accurately we know a particle’s The diameter of an H atom is about 0.2 nm, so the uncertainty in the position, the less accurately we can know its electron’s position is greater than the size of the H atom! We simply can’t say where the electron is “located” in the H atom. Now, try momentum. the calculation using a 70.0 kg man running at 10.0 m/s! You will find that the uncertainty in the man’s position is Very small.

Figure 7.9: The Quantum Mechanics standing waves • Based on the wave properties of the atom caused by the H$ y = Ey vibration of a guitar string • y = wave function fastened at both • H$ = mathematical operator ends. Each dot • E = total energy of the atom represents a node • A specific wave function is often called an (a point orbital. of zero displacement).

Figure 7.5: (a) Diffraction occurs Figure 7.10: when electromagnetic radiation is scattered The hydrogen from a regular array of objects, such as the ions electron in a crystal of sodium chloride. The large spot in the center is from the visualized as main incident beam of X rays. (b) Bright spots in a standing the diffraction pattern result from constructive wave around interference of waves. The waves are in phase; that is, their peaks the nucleus. match. (c) Dark areas result from destructive interference of waves. The waves are out of phase; the peaks of one wave coincide with the troughs of another wave

9 Quantum Mechanics Heisenberg’s Uncertainty Principle: (? x)(? px) = h/4p or • Although we cannot precisely define an electron’s (? x)(? vx) = h/4pm, orbit, we can obtain the probability of finding an -34 electron at a given point around the nucleus. h = 6.63 x 10 Js For subatomic particles which – Erwin Schrodinger defined this probability in a mathematical expression called a wave function, have very small masses, the un- denoted y (psi). certainties in position and speed – The probability of finding a particle in a region of can be quite large. We can only space is defined by y2 . (see Figures 7.18 and 7.19) talk about the probability, ? 2, of where we might find them.

Probability Distribution Figure 7.11: (a) The probability 4 square of the wave function distribution for the hydrogen 1s 4 probability of finding an electron at a given orbital in three position dimensional space. (b) The probability of 4 Radial probability distribution is the finding the electron probability distribution in each spherical at points along a line shell. drawn from the nucleus outward in any direction for the hydrogen 1s orbital.

Figure 7.12: (a) Cross section of the hydrogen 1s orbital probability distribution divided into successive thin spherical shells. (b) The radial probability distribution. Figure 7.13: Two representations of the hydrogen 1s, 2s, and 3s orbitals.

10 Figure 7.25: Cutaway diagrams showing the spherical shape of S orbitals . Quantum Numbers (QN) • 1. Principal QN (n = 1, 2, 3, . . .) - related to size and energy of the orbital. • 2. Angular Momentum QN (l = 0 to n - 1) - relates to shape of the orbital.

• 3. Magnetic QN (m l = l to -l) - relates to orientation of the orbital in space relative to other orbitals.

1 1 • 4. Electron Spin QN (m s = + /2, - /2) - relates to the spin states of the electrons.

Figure 7.14: Representation of the 2p Figure 7.15: A cross section of the orbitals . (a) The electron probability distributed for a 2p orbital. (b) The boundary electron probability distribution for a surface representations of all three 2p orbitals . 3p orbital.

11 Figure 7.17: Representation of Figure 7.16: Representation of the 4f orbitals in terms of their the 3d orbitals. boundary surface.

A Summary of the H Atom Figure 7.18: Orbital energy In the quantum (wave) mechanical model, the electron is a “standing levels for the hydrogen atom. wave” and can be described by a series of wave functions (orbitals) That describe the possible energies and spatial distributions available to the electron. From the Heisenberg Uncertainty Principle, the model cannot predict the exact electron motions – the Q2 represents the probability distri- bution of the electrons in that orbital. We depict the orbital as an electron density map. The size of an orbital is arbitrarily defined as the surface that contains 90% of the total electron density. The H atom has many types of orbitals – in the ground state, the single electron resides in the 1s orbital. The electron can be excited to higher energy orbitals if the appropriate amount of energy is put into the atom.

Figure 8.3: A representation of electron spin.

Electron Spin

• In Chapter 7, we saw that electron pairs residing in the same orbital are required to have opposing spins. – This causes electrons to behave like tiny bar magnets. (see Figure 8.3) – A beam of hydrogen atoms is split in two by a magnetic field due to these magnetic properties of the electrons. (see Figure 8.2)

Return to slide 2

12 I Electronic Structure of Atoms Figure 8.2: The Stern-Gerlach A Electron Spin and the Pauli Exclusion Principle experiment.

1 The Stern-Gerlach Experiment demonstrated interaction between the magnetic field of the spinning electron (in a stream of Ag atoms) and an external magnetic field. When ms = + ½ same poles are aligned; when ms = - ½ opposite poles are aligned

Figure 7.20: A comparison of the Pauli Exclusion Principle radial probability distributions of the 2s and 2p orbitals. • In a given atom, no two electrons can have the same set of four quantum numbers

(n, l, ml, ms). • Therefore, an orbital can hold only two electrons, and they must have opposite spins.

Figure 7.21: (a) The radial probability Figure 7.22: The orders of the energies of the distribution for an electron in a 3s orbital. (b) orbitals in the first three levels of The radial probability distribution for the 3s, polyelectronic atoms. 3p, and 3d orbitals .

13 “ Things on a very small scale behave like nothing that you have any direct experience about. They do not behave like waves, Figure 8.8: they do not behave like particles, they do not behave like clouds, Orbital energies or billiard balls, or weights on springs, or like anything that you for the have ever seen.” scandium atom - Richard P. Feynman (Z=21). The Feynman Lectures on Physics

Aufbau Principle Aufbau Principle • The Aufbau principle is a scheme used to reproduce the ground state electron configurations • As protons are added one by one to of atoms by following the “building up” order. the nucleus to build up the elements, – Listed below is the order in which all the possible sub- shells fill with electrons. electrons are similarly added to these 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f hydrogen-like orbitals.

– You need not memorize this order. As you will see, it can be easily obtained.

Electron Configuration Electron Configuration

• An “electron configuration” of an atom is a • An orbital diagram is used to show how the particular distribution of electrons among orbitals of a sub shell are occupied by electrons. available sub shells. – The notation for a configuration lists the sub- – Each orbital is represented by a circle. shell symbols sequentially with a superscript – Each group of orbitals is labeled by its sub shell indicating the number of electrons occupying notation. that sub shell. – For example, lithium (atomic number 3) has 1s 2s 2p two electrons in the “1s” sub shell and one – Electrons are represented by arrows: electron in the “2s” sub shell 1s2 2s1. up for ms = +1/2 and down for m s = -1/2

14 The Pauli Exclusion Principle The Pauli Exclusion Principle • The maximum number of electrons and • The Pauli exclusion principle, which their orbital diagrams are: summarizes experimental observations, states that no two electrons can have the same four Maximum Number of Number of quantum numbers. Sub shell Orbitals Electrons – In other words, an orbital can hold at most two s (l = 0) 1 2 electrons, and then only if the electrons have opposite spins. p (l = 1) 3 6 d (l =2) 5 10 f (l =3) 7 14

Hund’s Rule Orbital Diagrams • Consider carbon (Z = 6) with the ground • The lowest energy configuration state configuration 1s 22s 22p2. for an atom is the one having the – Three possible arrangements are given in the following orbital diagrams. maximum number of unpaired 1s 2s 2p electrons allowed by the Pauli Diagram 1: principle in a particular set of Diagram 2: degenerate orbitals. Diagram 3: – Each state has a different energy and different magnetic characteristics.

Orbital Diagrams Magnetic Properties • Hund’s rule states that the lowest energy arrangement • Although an electron behaves like a tiny magnet, two (the “ground state”) of electrons in a sub-shell is electrons that are opposite in spin cancel each other. obtained by putting electrons into separate orbitals of Only atoms with unpaired electrons exhibit magnetic the sub shell with the same spin before pairing susceptibility. electrons. – A paramagnetic substance is one that is – Looking at carbon again, we see that the ground weakly attracted by a magnetic field, usually state configuration corresponds to diagram 1 the result of unpaired electrons. when following Hund’s rule. – A diamagnetic substance is not attracted by a magnetic field generally because it has only paired electrons. 1s 2s 2p

15 Figure 7.29: A diagram that summarizes the order in which the orbitals fill in Aufbau Principle polyelectronic atoms. • With boron (Z=5), the electrons begin filling the 2p subshell. Z=5 Boron 1s22s 22p1 or [He]2s 22p1 Z=6 Carbon 1s22s 22p2 or [He]2s 22p2 Z=7 Nitrogen 1s22s 22p3 or [He]2s 22p3 Z=8 Oxygen 1s22s 22p4 or [He]2s 22p4 Z=9 Fluorine 1s22s 22p5 or [He]2s 22p5 Z=10 Neon 1s22s 22p6 or [He]2s 62p6

Configurations and the Aufbau Principle Periodic Table • With sodium (Z = 11), the 3s sub shell • Electrons that reside in the outermost shell of an begins to fill. atom - or in other words, those electrons outside Z=11 Sodium 1s22s 22p63s1 or [Ne ]3s1 the “noble gas core”- are called valence electrons. Z=12 Magnesium 1s22s 22p23s2 or [Ne ]3s2 – These electrons are primarily involved in chemical reactions. – Then the 3p sub shell begins to fill. – Elements within a given group have the same Z=13 Aluminum 1s22s 22p63s23p1 or [Ne]3s23p1 “valence shell configuration.” • – This accounts for the similarity of the chemical • properties among groups of elements. Z=18 Argon 1s22s 22p63s23p6 or [Ne]3s23p6

Figure 8.12: A periodic table illustrating the building-up order. Valence Electrons The electrons in the outermost principle quantum level of an atom. Atom Valence Electrons Ca 2 N 5 Br 7

Inner electrons are called core electrons.

16 Figure 7.26: Electron Configurations and the configurations for potassium Periodic Table • The following slide illustrates how the periodic through krypton. table provides a sound way to remember the Aufbau sequence. – In many cases you need only the configuration of the outer elements. – You can determine this from their position on the periodic table. – The total number of valence electrons for an atom equals its group number.

Figure 7.25: The electron configurations Figure 7.27: The orbitals being filled for in the type of orbital occupied last for the elements in various parts of the periodic first 18 elements. table.

Figure 7.24: Mendeleev's early periodic table, published in 1872. Note the spaces left for missing elements with atomic masses 44, 68, 72, 100.

17 Figure 8.9: A periodic table.

Information Contained in the Configurations and the Periodic Table Periodic Table • 1. Each group member has the same valence • Note that elements within a given family electron configuration (these electrons have similar configurations. primarily determine an atom’s chemistry). • 2. The electron configuration of any – The Group IIA elements are sometimes called representative element. the alkaline earth metals. Beryllium 1s22s 2 • 3. Certain groups have special names (alkali metals, halogens, etc). Magnesium 1s22s22p63s 2 • 4. Metals and nonmetals are characterized by Calcium 1s22s22p63s23p64s 2 their chemical and physical properties.

Configurations and the Broad Periodic Table Periodic Table Classifications • Note that elements within a given family have similar configurations. • Representative Elements (main group): filling s and p orbitals (Na, Al, Ne, O) – For instance, look at the noble gases. • Transition Elements : filling d orbitals Helium 1s2 (Fe, Co, Ni) Neon 1s22s 22p6 • Lanthanide and Actinide Series (inner Argon 1s22s22p63s 23p6 transition elements): filling 4f and 5f Krypton 1s22s22p63s23p63d104s 24p6 orbitals (Eu, Am, Es)

18 Figure 7.36: Special names for groups in the periodic table.

Ionization Energy

• The quantity of energy required to remove an electron from the gaseous atom or ion.

112

Figure 7.31: The values of first Figure 7.32: Trends in ionization ionization energy for the energies (kJ/mol) for the elements in the first six periods. representative elements.

19 Periodic Trends • First ionization energy: • increases from left to right across a period; • decreases going down a group.

This means that metallic character increases as we go from top to bottom of a group

116

A vial containing potassium metal. The sealed Sodium metal vial contains an inert gas to protect the is so reactive potassium from reacting with oxygen. that it is stored under º kerosene to 4 K (s) + O 2 (g) protect it from 2 K 2O (s) the oxygen in the air.

4 Na (s) + O 2 (g) º 2 Na2O (s)

Potassium reacts violently with water.

2 K (s) + H 2O (l) º 2 KOH (aq) + H 2 (g)

20 Electron Affinity Periodic Properties • Electron Affinity • The electron affinity is the energy • The energy change associated change for the process of adding an with the addition of an electron to a electron to a neutral atom in the gaseous gaseous atom. state to form a negative ion. - - –For a chlorine atom, the first electron affinity • X(g) + e ® X (g) is illustrated by: Cl([Ne ]3s23p5)+ e- ® Cl- ([Ne ]3s23p6) Electron Affinity = -349 kJ/mol

Periodic Properties • Electron Affinity – The more negative the electron affinity, the more stable the negative ion that is formed. – Broadly speaking, the general trend goes from lower left to upper right as electron affinities become more negative. – Table 8.4 gives the electron affinities of the main-group elements.

Figure 7.33: The electron affinity values for atoms among the first 20 elements that form stable, isolated X- ions.

21 Figure 7.34: The radius of an atom (r) is defined as half the distance between the nuclei Periodic Trends in a molecule consisting of identical atoms. • Atomic Radii: • decrease going from left to right across a period; • increase going down a group.

Figure 7.28: The periodic table with atomic symbols, Figure 7.35: atomic numbers, and partial electron configurations. Atomic radii (in picometers ) for selected atoms.

22