Unit 4 The Periodic Table Properties Trends and

Unit 4 The Periodic Table Properties, Trends and Connections

Properties of Light • Sometimes light behaves like waves, and other times like particles. • Visible light is a kind of electromagnetic radiation, which is a form of energy that exhibits wavelike behavior as it travels through space.

The Electromagnetic Spectrum • Together, all the forms of electromagnetic radiation form the electromagnetic spectrum.

Visible Light • Visible Light is the narrow band of electromagnetic radiation that we can see. Visible Spectrum • It consists of a range of waves Color Wavelength Red 700 ‑ 650 nm with various Orange 649 ‑ 580 nm Yellow 579 ‑ 575 nm wavelengths. Green Blue Indigo Violet 574 ‑ 490 nm 489 ‑ 455 nm 454 ‑ 425 nm 424 ‑ 400 nm

The Speed of Light • The constant, c, equals the speed of light, and it is a fundamental constant of the universe. • All waves in the electromagnetic spectrum travel at the speed of light, c = 3. 00 x 108 m/s.

Properties of Light (continued) • Wavelength (λ) is the distance between corresponding points on adjacent waves. • Frequency (ν) is defined as the number of waves that pass a given point in a specific time, usually one second.

Wavelength vs. Frequency • Wavelength (λ) is inversely proportional to frequency (ν). In other words, when λ increases, ν decreases, and vice versa.

Wavelength vs. Frequency (continued) • The relationship between wavelength and frequency is described by the equation: c = λν • Where c is a constant (always the same number) equal to 3 x 108 m/s. • λ is the wavelength (in m). Problem-solving hint: 1 nm = 10 -9 m. • ν is the frequency (in s− 1 or Hz). Visual Concept

The Speed of Light Sample Problem A photon of light has a frequency of 4. 4 x 1014 Hz. Calculate its wavelength. Does it fall within the visible spectrum? If so, what color is it? Solution: Visible Spectrum Color Wavelength Use the equation: c = λν Red 620 ‑ 750 nm 590 ‑ 620 nm 3. 00 x 108 m/s = λ (4. 4 x 1014 Hz) Orange Yellow 570 ‑ 590 nm Green 495 ‑ 570 nm 3. 00 x 108 m/s 455 ‑ 495 nm -7 m Blue λ= = 6. 8 x 10 Indigo 425 ‑ 455 nm 14 4. 4 x 10 Hz Violet 400 ‑ 425 nm λ = 680 x 10 -9 m, or λ = 680 nm Yes, it is red light

The Photoelectric Effect • By the early 1900 s, scientists observed interactions of light and matter that couldn’t be explained by wave theory. • The photoelectric effect refers to the emission of electrons from a metal when light shines on the metal. Visual Concept

Energy and Frequency • In 1905, Einstein proposed an explanation: light energy must come in “packets”. • The energy in a light packet depends on its frequency according to the equation: E = hν • Where h is a constant (called Planck’s constant) equal to 6. 626 x 10 -34 J●s. • E is energy (J), and ν is frequency (s− 1 or Hz). Visual Concept

Energy and Frequency Sample Problem A friend tells you that he measured the energy of his microwaves at home and got a value of 6. 4 x 1010 J. What is the frequency of the microwaves? Does this value seem correct? Explain. Solution: Use the equation: E = hν E 6. 4 x 1010 J = v= 6. 626 x 10 -34 J●s h v = 9. 7 x 1043 Hz No, it is nowhere near the microwave range

The Particle Description of Light • A quantum of energy is the minimum quantity of energy that can be lost or gained by an atom. • A photon is a particle of electromagnetic radiation having zero mass and carrying a quantum of energy. • The energy of a particular photon is directly proportional to the frequency of the radiation. Visual Concept

Energy States • Ground state – The lowest energy state of an atom. • Excited state – an atom has a higher potential energy than it has in its ground state. • When an excited atom returns to its ground state, it gives off energy in the form of electromagnetic radiation.

Hydrogen’s Line Emission Spectrum

The Bohr Model • In 1913, Danish physicist Niels Bohr proposed a hydrogen-atom model that linked the atom’s electron to its line-emission spectrum.

The Bohr Model (continued) • According to the Bohr model, the electron can circle the nucleus only in allowed paths, or orbits. • The energy of the electron is higher when it is in orbits that are farther from the nucleus. Visual Concept

Electrons as Waves • In 1924, French scientist Louis de Broglie suggested that electrons act like waves confined to the space around an atomic nucleus. • It followed that the electron waves could exist only at specific frequencies – corresponding to the quantized energies of Bohr’s orbits.

The Heisenberg Uncertainty Principle • In 1927, German physicist Werner Heisenberg realized that an attempt to locate an electron with a photon knocks the electron off its course. • The Heisenberg uncertainty principle states that it is impossible to determine simultaneously both the position and velocity of an electron or any other very small particle. Visual Concept

The Schrödinger Wave Equation • In 1926, Austrian physicist Erwin Schrödinger developed an equation that treated electrons in atoms as waves. • Together with Heisenberg and others, Schrödinger laid the foundation for modern quantum theory. Visual Concept

“If you think you understand quantum mechanics, you don’t understand quantum mechanics. ” - Richard Feynman “Shut up and calculate!” - David Mermin

Quantum Theory • Quantum theory describes mathematically the wave properties of electrons and other very small particles. • There are four different types of quantum numbers used: 1. 2. 3. 4. Principal quantum # (n) – energy level. Angular momentum quantum # (l) - sublevel. Magnetic quantum # (m) - orbital. Spin quantum # (s).

Principal Quantum Number • The Principal Quantum Number (n) indicates the main energy level occupied by an electron. • As n increases, the electron’s energy and its distance from the nucleus increases.

Angular Momentum Quantum Number • The Angular Momentum Quantum Number (l) (also called the sublevel) indicates the shape of the orbital. • The number of sublevels allowed for each energy level is equal s orbital to n. p orbital d orbital sphere dumbbell cloverleaf f orbital complex

Magnetic Quantum Number • The Magnetic Quantum Number (m) indicates the orientation of an orbital around the nucleus. Sublevel Orbitals s 1 p 3 d 5 f 7

Spin Quantum Number • The Spin Quantum Number (s) indicates the fundamental spin state of an electron in an orbital. • There are only two possible values for s, +½ and –½. • A single orbital can hold a maximum of two electrons, but the electrons must have opposite spin states.

Quantum Numbers Overview

Electron Configuration Rules • According to the Aufbau Principle, an electron occupies the lowest-energy orbital that can receive it. • The order of increasing energy is shown on the vertical axis. Each box represents an orbital. (diagram on pg. 325)

Electron Configuration Rules (continued) • According to the Pauli exclusion principle, no two electrons in the same atom can have the same set of four quantum numbers.

Electron Configuration Rules (continued) • According to Hund’s Rule, orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin state. Wrong Correct

Orbital Notation • An orbital containing one electron is represented as: • An orbital containing two electrons is represented as: • The lines are labeled with the principal quantum number and sublevel letter. For example, the orbital notation for helium is written as follows: He 1 s

Orbital Notation Sample Problem 1 Write the orbital notation for Carbon. Solution: Carbon is atomic number 6, so it has 6 electrons. The first two electrons go in the 1 s orbital. The next two electrons go in the 2 s orbital. The final two electrons go in the 2 p orbitals. Carbon 1 s 2 s 2 p

Electron Configuration Notation • Electron-configuration notation eliminates the lines and arrows of orbital notation. • Instead, the number of electrons in a sublevel is shown by a superscript. • Example: Carbon Orbital Notation 1 s 2 s Electron Configuration 1 s 22 p 2 2 p

Mendeleev and Periodicity • The first periodic table of the elements was published in 1869 by Russian chemist Dmitri Mendeleev. • Mendeleev left empty spaces in his table and predicted elements that would fill 3 of the spaces. • By 1886, all 3 of these elements had been discovered. Visual Concept

Mosley and the Periodic Law • In 1911, the English scientist Henry Moseley discovered that the elements fit into patterns better when they were arranged according to atomic number, rather than atomic weight. • The Periodic Law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers.

The Periodic Table • Elements in the periodic table arranged into vertical columns, called groups or families, that share similar chemical properties. • Elements are also organized horizontally in rows, or periods.

Blocks of the Periodic Table s p d f

Electron Configuration Notation Sample Problem 1 a. Write electron configuration for Selenium (Se). b. How many unpaired electrons are in an atom of Selenium? Solution: a. 1 s 2 2 p 6 3 s 2 3 p 6 4 s 2 3 d 10 4 p 4 b. Only consider the 4 p 4 electrons, since all electrons will be paired in filled orbitals. 2 electrons are unpaired 4 p

Noble Gas Notation • The Group 18 elements (He, Ne, Ar, Kr, Xe, and Rn) are called the noble gases. • Noble gas notation is an abbreviated electron configuration. • Use square brackets around the noble gas at the end of the prior period to replace part of the configuration. • Example: Calcium Electron Configuration 1 s 22 p 63 s 23 p 64 s 2 Noble Gas Notation [Ar]4 s 2

Noble Gas Notation Sample Problem 1 a. Write the noble gas notation for Gold (Au). b. How many inner-shell electrons does this atom have? Solution: a. [Xe] 6 s 2 4 f 14 5 d 9 b. The outer shell is the one with the highest #. There are 2 e- in energy level 6 (6 s 2). All the rest are inner-shell electrons. 79 total e- - 2 outer-shell e- = 77 inner-shell e-

Electron-Dot Notation • Electron-dot notation is indicated by dots placed around the element’s symbol. Only the valence electrons are shown. Inner-shell electrons are not shown.

Electron-Dot Notation Sample Problem a. Write the electron-dot notation for hydrogen. b. Write the electron-dot notation for nitrogen. Solution: a. Hydrogen is in group 1. It has one valence electron. • H a. Nitrogen is in group 15. It has 5 valence electrons. • • • N • •

Atomic Radii • Atomic radius – one-half the distance between the nuclei of identical atoms Group 1 that are bonded together. • Atomic radii tend to increase as you go down a group because electrons occupy successively higher energy levels farther away from the nucleus.

Atomic Radii (continued) • Atomic radii tend to decrease as you go across a period because as more electrons are added they are pulled closer to the more highly charged nucleus. Period 2

Atomic Radii Sample Problem Of the elements Mg, Cl, Na, and P, which has the largest atomic radius? Explain. Solution: Na has the largest radius. All of the elements are in the 3 rd period, and atomic radii decrease across a period.

Ionic Radii: Cations • Atoms with 1, 2, or 3 valence electrons tend to lose them to form positive ions, which are called cations. • Cations are always much smaller than their corresponding atoms.

Ionic Radii: Anions • Atoms with 5, 6, or 7 valence electrons tend to gain more in order to have an octet (8 electrons ) in their outer shell. • Gaining extra electrons forms negative ions, called anions. • Anions are always much larger than their corresponding atoms.

Ionization Energy • The energy required to remove an electron from a neutral atom of an element in the gaseous state is called the ionization energy (IE). • Ionization energy tends to increase across each period because a higher nuclear charge more strongly attracts electrons in the same energy level.

Ionization Energy (continued) • Ionization energy tends to decrease down each group because electrons farther from the nucleus are removed more easily.

Ionization Energy Sample Problem Consider two elements, A and B. A has an IE of 419 k. J/mol. B has an IE of 1000 k. J/mol. Which element is more likely to be in the s block? Which will be in the p block? Which is more likely to form a positive ion? Solution: Element A is most likely to be in the s-block since IE increases across the periods. Element B would most likely lie at the end of a period in the p block. Element A is more likely to form a positive ion since it has a much lower IE than B.

Electron Affinity and Electronegativity • Electron affinity is the energy change that occurs when an electron is acquired by a neutral atom. • Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound. • Electronegativity applies to atoms in a compound, while electron affinity is a property of isolated atoms.

Electron Affinity and Electronegativity (continued) • Electron affinity and electronegativity both tend to increase across periods, and decrease (or stay the same) down a group.

Electronegativity Sample Problem Of the elements Ga, Br, and Ca, which has the highest electronegativity? Explain. Solution: All of these elements are in the fourth period. Br has the highest atomic number and is farthest to the right in the period. Br would have the highest electronegativity since electronegativity increases across a period.

Group 1: Alkali Metals • Group 1 elements are called alkali metals. • Alkali metals have a silvery appearance and are soft enough to cut with a knife. • They are extremely reactive and are not found in nature as free elements. • They must be stored under oil or kerosene. Visual Concept

Group 2: Alkaline Earth Metals • Elements in group 2 are known as the alkaline earth metals. • Group 2 metals are harder, denser and stronger than alkali metals, and have higher melting points. • Less reactive than group 1, but still too reactive to be found in nature as free elements.

Group 17: Halogens • Elements in group 17 are known as the halogens. • Halogens are the most reactive nonmetals, reacting vigorously with metals to form salts • Most halogens exist in nature as diatomic molecules (i. e. F 2, Cl 2, Br 2 and I 2. ) Visual Concept

Group 18: Noble Gases • Elements in group 18 are known as noble gases. • They are completely nonreactive and don’t form compounds under normal conditions. • A new group was added to the periodic table in 1898 for the noble gases.

Chapter 5 – Section 2: Electron Configuration and the Periodic Table d-block: Transition Metals • Elements in the d-block are called transition metals. • They have typical metallic properties such as conduction of electricity and high luster. • Less reactive than group 1 and 2 elements. • Some (i. e. platinum & gold) are so unreactive they usually don’t form compounds.

f-block: Lanthanides & Actinides • Elements in the period 6 of the f-block are called lanthanides (or rare-earth). • Lanthanides are shiny metals similar in reactivity to alkaline earth metals. • Elements in period 7 of the f-block are called actinides. • Actinides are all radioactive, and many of them are known only as man-made elements.