How Many Core Electrons Are in A Chlorine Atom Chemistry MCQs

word image 2851 Introductory

Chemistry

Fifth Edition

Nivaldo J. Tro

Chapter 9

Electrons in

Atoms and the Periodic

Table

Dr. Sylvia Esjornson

Southwestern Oklahoma State University

Weatherford, OK

Blimps, Balloons, and Models of the Atom

  • word image 2852 On May 6, 1937, while landing in New Jersey on its first transatlantic crossing, the Hindenburg burst into flames, destroying the airship and killing 36 of the 97 passengers.
  • Apparently, as the Hindenburg was landing, leaking hydrogen gas ignited, resulting in an explosion that destroyed the airship.
  • The skin of the Hindenburg, which was constructed of a flammable material, may have also been partially to blame for its demise.

The Hindenburg was filled with hydrogen, a reactive and flammable gas. Question: What makes hydrogen reactive?

Blimps, Balloons, and Models of the Atom

  • word image 2853 Modern blimps are filled with helium, an inert gas.
  • The nucleus of the helium atom has two protons, so the neutral helium atom has two electrons—a highly stable configuration.
  • In this chapter, we learn about models that explain the inertness of helium and the reactivity of other elements.

Why is helium inert?

Niels Bohr and Erwin Schrödinger

Niels Bohr (left) and Erwin Schrödinger (right), along with Albert Einstein, played a role in the development of quantum mechanics, yet they were bewildered by their own theory of wave-particle duality for the electron.

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Light: Electromagnetic Radiation

  • The interaction of light with atoms helped to shape scientists’ models of the atom.
  • Light is a form of electromagnetic radiation.
  • Light is a type of energy that travels through space at a constant speed of 3.0 × 108 m/s (186,000 mi/s).
  • Light has properties of both waves and particles.

Light: Electromagnetic Radiation word image 2855

When a water surface is disturbed, waves are created that radiate outward from the site. The wave carries energy as it moves through the water.

Light: Electromagnetic Radiation

Wavelength: The wavelength of light, λ (lambda, pronounced “lam-duh”), is defined as the distance between adjacent wave crests.

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Light: Color of Light

  • White light, as produced by the sun or by a lightbulb, contains a spectrum of wavelengths and therefore a spectrum of color.
  • We see these colors—red, orange, yellow, green, blue, indigo, and violet—in a rainbow or when white light is passed through a prism.
  • Red light, with a wavelength of 750 nm (nanometers), has the longest wavelength of visible light.
  • Violet light, with a wavelength of 400 nm, has the shortest wavelength of visible light (1 nm = 1 × 10–9 m).
  • The presence of color in white light is responsible for the colors we see in our everyday vision.

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Light: Color of Light

word image 2858Components of white light

R O Y G B I V

• Light is separated into its constituent colors—red, orange, yellow, green, blue, indigo, and violet— when it is passed through a prism.

Light: Color in Objects

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• A red shirt appears red because it reflects red light; the shirt absorbs all of the other colors of light except the red light. Our eyes see only the reflected light, making the shirt appear red.

Light: Electromagnetic Radiation

  • Frequency: The frequency of light, ν (nu, pronounced “noo”), is defined as the number of cycles or crests that pass through a stationary point in one second.
  • Wavelength and frequency are inversely related—the shorter the wavelength, the higher the frequency.

Electromagnetic Radiation (Photons— Particles of Light)

  • Light can be viewed as a stream of particles.
  • A particle of light is called a photon.
  • We can think of a photon as a single packet of light energy.
  • The amount of energy carried in the packet depends on the wavelength of the light—the shorter the wavelength, the greater the energy.
  • Light waves carry more energy if their crests are closer together (higher frequency and shorter wavelength).
  • Violet light (shorter wavelength) carries more energy per photon than red light (longer wavelength).

The Electromagnetic Spectrum

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The Electromagnetic Spectrum

  • Visible light ranges from violet (shorter wavelength, higher energy) to red (longer wavelength, lower energy).
  • Photons of visible light do not damage biological molecules.
  • Photons of visible light do cause molecules in our eyes to rearrange, which sends a signal to our brains that results in vision.

EXAMPLE 9.1 Wavelength, Energy, and Frequency

• Arrange the three types of electromagnetic radiation—visible light, X-rays, and microwaves—in order of increasing:

  1. Wavelength
  2. Frequency
  3. Energy per photon

word image 2861Each Element Has Its Own Atomic Emission Spectrum of Light

  • Neon atoms inside a glass tube absorb electrical energy and then reemit the energy as red light.

Each Element Has Its Own Atomic Emission Spectrum of Light

  • word image 2862 Light emitted from a mercury lamp

(left) appears blue, and light emitted from a hydrogen lamp (right) appears pink.

Emission Spectra of the Elements Are

Not Continuous

A white-light spectrum is continuous, with some radiation emitted at every wavelength. The emission spectrum of an individual element includes only certain specific wavelengths.

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Light Emitted by Hydrogen Contains Distinct Wavelengths That Are Specific to Hydrogen

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Niels Bohr Developed a Simple Model to Explain These Results

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word image 2866The Energy Is Quantized

The energy of each Bohr orbit, specified by a quantum number n = 1, 2, 3 is fixed, or quantized.

Bohr orbits are like steps of a ladder, each at a specific distance from the nucleus and each at a specific energy. It is impossible for an electron to exist between orbits in the Bohr model.

Excitation and Emission

word image 2867 • When a hydrogen atom absorbs energy, an electron is excited to a higher-energy orbit. The electron then relaxes back to a lower-energy orbit, emitting a photon of light.

Hydrogen Emission Lines

word image 2868

The Bohr Model: Atoms with Orbits

  • The great success of the Bohr model of the atom was that it predicted the lines of the hydrogen emission spectrum.
  • However, it failed to predict the emission spectra of other elements that contained more than one electron.
  • For this and other reasons, the Bohr model was replaced with a more sophisticated model called the quantum-mechanical or wave-mechanical model.

The Quantum-Mechanical Model:

Atoms with Orbitals

  • The quantum-mechanical model of the atom replaced the Bohr model in the early twentieth century. In the quantum-mechanical model,

Bohr orbits are replaced with quantum-mechanical orbitals.

  • Orbitals are different from orbits in that they represent probability maps that show a statistical distribution of where the electron is likely to be found.

The Quantum-Mechanical Model:

Atoms with Orbitals

  • Quantum mechanics revolutionized physics and chemistry because, in the quantum-mechanical model, electrons do not behave like particles flying through space.
  • We cannot, in general, describe their exact paths.
  • An orbital is a probability map that shows where the electron is likely to be found when the atom is probed; it does not represent the exact path that an electron takes as it travels through space.

Baseball Paths and Electron Probability Maps

word image 2869 • To describe the behavior of a “pitched” electron, you would have to construct a probability map of where it would cross home plate.

Principal Quantum Numbers for Orbitals

  • In the quantum-mechanical model, a number and a letter specify an orbital (or orbitals).
  • The lowest-energy orbital in the quantum-mechanical model is called the 1s orbital.
  • It is specified by the number 1 and the letter s.
  • The number is called the principal quantum number (n) and specifies the principal shell of the orbital.

Ground States and Excited States

  • The single electron of an undisturbed hydrogen atom at room temperature is in the 1s orbital.
  • This is called the ground state, or lowest energy state, of the hydrogen atom.
  • The absorption of energy by a hydrogen atom can cause the electron to jump (or make a transition) from the 1s orbital to a higher-energy orbital. When the electron is in a higher-energy orbital, the hydrogen atom is said to be in an excited state.
  • All the atoms of each element have one ground state and many excited states. word image 2870

word image 2871Energy Increases with Principal Quantum Number

  • The higher the principal quantum number, the higher the energy of the orbital.
  • The possible principal quantum numbers are n = 1, 2, 3 with energy increasing as n increases.
  • Since the 1s orbital has the lowest possible principal quantum number, it is in the lowest-energy shell and has the lowest possible energy.

Shapes of Quantum-Mechanical Orbitals

  • The letter indicates the subshell of the orbital and specifies its shape.
  • The possible letters are s, p, d, and f, each with a different shape.
  • Orbitals within the s subshell have a spherical shape.

word image 2872

Representations of Orbitals

  • Orbitals are sometimes represented by dots, where the dot density is proportional to the probability of finding the electron.
  • The dot density for the 1s orbital is greatest near the nucleus and decreases farther away from the nucleus.
  • The electron is more likely to be found close to the nucleus than far away from it. word image 2873

Representations of Orbitals

Dot density and shape representations of the 1s orbital: The dot density is proportional to the probability of finding the electron. The greater dot density near the middle represents a higher probability of finding the electron near the nucleus.

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The Number of Subshells in a Given Principal Shell Is Equal to the Value of n

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The 2s Orbital Is Similar to the 1s Orbital, but Larger in Size

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The 2p Orbitals: This figure Shows Both the Dot Representation (Left) and Shape Representation (Right) for Each p Orbital

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Orbitals When n = 3

  • The next principal shell, n = 3, contains three subshells specified by s, p, and d.
  • The s and p subshells contain the 3s and 3p orbitals, similar in shape to the 2s and 2p orbitals, but slightly larger and higher in energy.
  • The d subshell contains five d orbitals.

word image 2880

The 3d Orbitals: This Figure Shows Both the Dot Representation (Left) and Shape Representation (Right) for Each d Orbital word image 2881

word image 2882 Introductory

Chemistry

Fifth Edition

Nivaldo J. Tro

Chapter 10

Chemical Bonding

Dr. Sylvia Esjornson

Southwestern Oklahoma State University

Weatherford, OK

Electronegativity and Polarity: Why Oil and Water Don’t Mix

  • word image 2883 If you combine oil and water in a container, they separate into distinct regions. Why? Something about water molecules causes them to bunch together into one region, expelling the oil molecules into a separate region.

Electronegativity and Polarity: Why Oil and

Water Don’t Mix

word image 2884

  • The two bonds between O and H each consist of an electron pair—two electrons shared between the oxygen atom and the hydrogen atom.
  • The oxygen and hydrogen atoms each donate one electron to this electron pair; however, they don’t share them equally.
  • The oxygen atom takes more than its fair share of the electron pair.

Electronegativity

  • The ability of an element to attract electrons within a covalent bond is called electronegativity.
  • Oxygen is more electronegative than hydrogen, which means that, on average, the shared electrons are more likely to be found near the oxygen atom than near the hydrogen atom.
  • Consider this representation of one of the two OH bonds:

word image 2885

  • The oxygen atom (getting the larger share) has a partial negative charge, symbolized by δ− (delta minus).
  • The hydrogen atom (getting the smaller share) has a partial positive charge, symbolized by δ+ (delta plus).
  • The result of this uneven electron sharing is a dipole moment, a separation of charge within the bond.

Polar Covalent Bonds

  • Covalent bonds that have a dipole moment are called polar covalent bonds.
  • The magnitude of the dipole moment, and the polarity of the bond, depend on the electronegativity difference between the two elements in the bond and the length of the bond.
  • For a fixed bond length, the greater the electronegativity difference, the greater the dipole moment and the more polar the bond.

Electronegativity

  • The value of electronegativity is assigned using a relative scale on which fluorine, the most electronegative element, has an electronegativity of 4.0.
  • Linus Pauling introduced the electronegativity scale used here. He arbitrarily set the electronegativity of fluorine at 4.0 and computed all other values relative to fluorine.

word image 2886

Identical Electronegativities

  • If two elements with identical electronegativities form a covalent bond, they share the electrons equally, and there is no dipole moment.
  • In Cl2, the two Cl atoms share the electrons evenly. This is a pure covalent bond. The bond has no dipole moment; the molecule is nonpolar.

word image 2887

Large Electronegativity Difference

  • If there is a large electronegativity difference between the two elements in a bond, such as what normally occurs between a metal and a nonmetal, the electron is completely transferred and the bond is ionic.
  • In NaCl, Na completely transfers an electron to Cl. This is an ionic bond.

word image 2888

Intermediate Electronegativity Difference

  • If there is an intermediate electronegativity difference between the two elements, such as between two different nonmetals, then the bond is polar covalent.
  • In HF, the electrons are shared, but the shared electrons are more likely to be found on F than on H. The bond is polar covalent.

word image 2889

word image 2890

The degree of bond polarity is a continuous function. The guidelines given here are approximate.

word image 2891

Polar Bonds and Polar Molecules

  • Does the presence of one or more polar bonds in a molecule always result in a polar molecule? The answer is no.
  • A polar molecule is one with polar bonds that add together— they do not cancel each other—to form a net dipole moment.
  • If a diatomic molecule contains a polar bond, then the molecule is polar.
  • For molecules with more than two atoms, it is more difficult to tell polar molecules from nonpolar ones because two or more polar bonds may cancel one another.

Polar Bonds and Polar Molecules

Consider carbon dioxide:

word image 2892

  • Each bond is polar because the difference in electronegativity between oxygen and carbon is 1.0.
  • CO2 has a linear geometry, the dipole moment of one bond completely cancels the dipole moment of the other, and the molecule is nonpolar.

word image 2893

Vector Notation for Dipole Moments

  • We can represent polar bonds with arrows (vectors) that point in the direction of the negative pole and have a plus sign at the positive pole (as just shown for carbon dioxide).
  • If the arrows (vectors) point in exactly opposing directions as in carbon dioxide, the dipole moments cancel.
  • In the vector representation of a dipole moment, the vector points in the direction of the atom with the partial negative charge.

Polar Bonds and Polar Molecules

word image 2894 Consider water (H2O):

  • Each bond is polar because the difference in electronegativity between oxygen and hydrogen is 1.4.
  • Water has two dipole moments that do not cancel, and the molecule is polar.

word image 2895 word image 2896 word image 2897

Attraction

word image 2898 word image 2899 word image 2900

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Conceptual Checkpoint 10.7

Which bond would you expect to be more polar: the bond in

HCl

or the bond in

HBr

?

Note: The answers to all Conceptual Checkpoints appear at the end of the chapter.

The Explanatory Power of the Quantum-Mechanical Model

Some observations:

  • Sodium tends to form Na+ ions, and fluorine tends to form Fions.
  • Some elements are metals, and others are nonmetals.
  • word image 2902 The noble gases are chemically inert, and the alkali metals are chemically reactive. word image 2903

The Explanatory Power of the Quantum-Mechanical Model

  • The chemical properties of elements are largely determined by the number of valence electrons they contain.
  • Their properties vary in a periodic fashion because the number of valence electrons is periodic.

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The Noble Gases

  • Calculations show that atoms with 8 valence electrons (or 2 for helium) are predicted to be particularly low in energy and therefore stable.

word image 2905 word image 2906

  • The noble gases are chemically stable, and thus relatively inert or nonreactive as accounted for by the quantum model.
  • Elements with electron configurations close to the noble gases are the most reactive because they can attain noble gas electron configurations by losing or gaining a small number of electrons.

word image 2907The Alkali Metals

  • word image 2908 Alkali metals (Group 1) are among the most reactive metals since their outer electron configuration (ns1) is 1 electron beyond a noble gas configuration.
  • If they can react to lose the electron, they attain a noble gas configuration.
  • This explains why the Group 1 metals tend to form 1+ cations.

The Alkaline Earth Metals

  • word image 2909 word image 2910 The alkaline earth metals (Group 2) all have electron configurations ns2 and are therefore 2 electrons beyond a noble gas configuration.
  • In their reactions, they tend to lose 2 electrons, forming 2+ ions and attaining a noble gas configuration.

The Halogens

  • The halogens (Group 7) all have ns2np5 electron configurations and are therefore 1 electron short of a noble gas configuration.

word image 2911 word image 2912

  • In their reactions, halogens tend to gain 1 electron, forming 1− ions and attaining a noble gas configuration.

Elements That Form Predictable Ions

word image 2913

Periodic Trends: Atomic Size Has Two Factors

  • #1: As you move to the right across a period in the periodic table, atomic size decreases.
  • The atomic size of an atom is determined by the distance between the outermost electrons and the nucleus.
  • The size of an orbital depends on the principal quantum number.
  • With each step across a period, the number of protons in the nucleus is increasing.
  • This increase in the number of protons results in a greater pull on the electrons from the nucleus, causing atomic size to decrease.

Periodic Trends: Atomic Size Has Two Factors

  • #2: As you move down a column in the periodic table, atomic size increases.
  • As you move down a column in the periodic table, the highest principal quantum number, n, increases.
  • Since the size of an orbital increases with increasing principal quantum number, the electrons that occupy the outermost orbitals are farther from the nucleus as you move down a column.

Periodic Properties: Atomic Size

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Periodic Properties:

Ionization Energy

• Ionization energy increases as you move to the right across a period and decreases as you move down a column in the periodic table.

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Periodic Properties: Ionization Energy word image 2916

Periodic Properties: Metallic Character

  • Metals tend to lose electrons in their chemical reactions, while nonmetals tend to gain electrons.
  • As you move across a period in the periodic table, ionization energy increases, which means that electrons are less likely to be lost in chemical reactions.
  • Metallic character decreases as you move to the right across a period and increases as you move down a column in the periodic table.

Periodic Properties: Metallic Character

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Chapter 9 in Review

  • Light, a form of electromagnetic radiation, exhibits both wavelike and particle-like behavior. Particles of light are called photons.
  • The Bohr model: The emission spectrum of hydrogen can be explained by the Bohr model for the hydrogen atom. Each orbit is specified by a quantum number (n), which also specifies the orbit’s energy.

Chapter 9 in Review

  • The quantum-mechanical model describes electron orbitals, which are electron probability maps that show the relative probability of finding an electron in various places surrounding the atomic nucleus.
  • An electron configuration indicates which orbitals are occupied for a particular atom. Orbitals are filled in order of increasing energy and obey the Pauli exclusion principle (each orbital can hold a maximum of two electrons with opposing spins) and Hund’s rule (electrons occupy orbitals of identical energy singly before pairing).

Chapter 9 in Review

  • The periodic table: Elements within the same column of the periodic table have similar outer electron configurations and the same number of valence electrons and therefore have similar chemical properties.
  • The periodic table is divisible into blocks (s block, p block, d block, and f block) in which particular sublevels are filled.
  • As you move across a period to the right in the periodic table, atomic size decreases, ionization energy increases, and metallic character decreases.
  • As you move down a column in the periodic table, atomic size increases, ionization energy decreases, and metallic character increases.
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