Chapter 9: ELECTRONS IN ATOMS AND THE PERIODIC TABLE

Chapter 9: ELECTRONS IN ATOMS AND THE PERIODIC TABLE

Problems: 1-3, 13-15, 19, 23-25, 31-32, 43, 45-46, 49c, 50a, 50b, 57c, 58 (b,c,d), 61-62, 69, 71-74, 77-88, 91-94

9.5 LIGHT: Electromagnetic Radiation
Light is a form of electromagnetic radiation, a type of energy that travels through space at a
constant speed, known as the speed of light (symbol c): 2.998×108 m/s (~186,000 mi./hour)
– While light may appear instantaneous to us, it’s really a wave traveling at this finite speed.

The term electromagnetic comes from the theory
proposed by Scottish scientist James Clerk Maxwell
that radiant energy consists of waves with an
oscillating electric field and an oscillating
magnetic field, which are perpendicular to one
another.

9.3 Electromagnetic Spectrum: continuum of radiant energy (see Fig. 9.4 on p. 280)
– The substances below are about the size of the wavelength indicated in the EM spectrum.

– e.g., an atom is about 10-10-10-9 m in size while a CD is about 10-3 m (or 1 mm) thick.

visible region: the portion of the EM spectrum that we can perceive as color

For example, a "red-hot" or "white-hot" iron bar freshly removed from a high-temperature
source has forms of energy in different parts of the EM spectrum
– red or white glow falls within the visible region, heat falls within the infrared region

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Thus, these electromagnetic waves have both a wavelength and a frequency:
wavelength (λ=Greek “lambda”): distance between successive peaks
frequency (ν=Greek “nu”): number of waves passing a given point in 1 s

How is energy related to wavelength and frequency?
– As the wavelength ↑, the frequency ↓, and the energy ↓
– As the wavelength ↓, the frequency ↑, and the energy ↑

Example: Which is higher in energy, red light at 700 nm or blue light at 400 nm?

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Classical Descriptions of Matter

John Dalton (1803)
– Atoms are hard, indivisible, billiard-like particles.
– Atoms have distinct masses (what distinguishes on type of atom from another).
– All atoms of an element are the same.

JJ Thomson (1890s)
– discovered charge-to-mass ratio of electrons

→ atoms are divisible because the electrons are one part of atom

Ernest Rutherford (1910)
– shot positively charged alpha particles at a thin foil of gold

→ discovery of the atomic nucleus

James Maxwell (1873)
– visible light consists of electromagnetic waves

Transition between Classical and Quantum Theory

Max Planck (1900); Blackbody Radiation
– heated solids to red or white heat
– noted matter did not emit energy in continuous bursts, but in whole-number multiples of

certain well-defined quantities

→ matter absorbs/emits energy in bundles = "quanta"

(single bundle of energy= "quantum")

Albert Einstein (1905); Photoelectric Effect

– Photoelectric Effect: Light shining on a clean metal → emission of electrons only when

the light has a minimum threshold frequency, ν0

– For ν < ν0 → no electrons are emitted
– For ν > ν0 → electrons are emitted, more e– emitted with greater intensity of light,

– Einstein applied Planck's quantum theory to light

→ light exists as a stream of "particles" called photons

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9.4 The Bohr Model: Atoms with Orbits

A Danish physicist named Niels Bohr used the results from the hydrogen emission spectrum
to develop a quantum model for the hydrogen atom.

Bohr Postulates: Bohr Model of the Atom

1. Energy-level Postulate
– Electrons move in discrete (quantized), circular orbits around the nucleus
– "tennis ball and stairs" analogy for electrons and energy levels
– a ball can bounce up to or drop from one stair to another, but it can never sit
halfway between two levels
– Each orbit has a specific energy associated with it, indicated as the principal energy
level or quantum number, n=1, 2, 3,...
– ground state or ground level (n = 1): lowest energy state for atom
– when the electron is in the lowest energy level in a hydrogen atom
– excited state: when the electron is in a higher energy level (n = 2,3,4,...)

2. Transitions Between Energy Levels
– When an atom absorbs energy
→ the electron can jump from a lower energy level to a higher energy level.
– When an electron drops from a higher energy level to a lower energy level
→ the atom releases energy, sometimes in the form of visible light.

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Emission Spectra: continuous or line spectra of radiation emitted by substances
– a heated solid (e.g. the filament in an incandescent light bulb) emits light that spreads out

to give a continuous spectrum = spectrum of all wavelengths of light, like a rainbow

Hydrogen Line Spectrum
– In contrast, when a sample of hydrogen is electrified, the resulting hydrogen emission

spectrum contains only a few discrete lines:

These discrete lines correspond to specific wavelengths → specific energies
→ The hydrogen atoms’ electrons can only emit certain energies

→ The energy of the electrons in the atom must also be quantized.
→ Planck’s postulate that energy is quantized also applies to the electrons in an atom.

– Each element has a unique line spectrum.

→ Emission spectra can be used to identify unknown elements in chemical analysis.
→ The element’s line spectrum is often called its "atomic fingerprint".

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Other examples of emission line spectra for mercury and neon to compare with hydrogen.

9.6 THE QUANTUM-MECHANICAL ORBITALS
Limitations of the Bohr Model → Quantum Mechanical Model
– Unfortunately, the Bohr Model failed for all other elements that had more than one proton

and more than one electron. (The multiple electron-nuclear attractions, electron-electron
repulsions, and nuclear repulsions make other atoms much more complicated than
hydrogen.)

Quantum Mechanical Model
In 1920s, a new discipline, quantum mechanics, was developed to describe the motion of
submicroscopic particles confined to tiny regions of space.
– Quantum mechanics makes no attempt to specify the position of a submicroscopic particle

at a given instant or how the particle got there
– It only gives the probability of finding submicroscopic particles (e.g. food court analogy)
→ Instead we “take a snapshot” of the atom at different times and “see” where the electrons

are likely to be found (See Fig. 9.16 on p. 296).

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ORBITALS AND THEIR SHAPES

Erwin Schrödinger (1926)
– developed a model that predicts the probability of finding the electron near a given point

– probability density for an electron is called the "electron cloud"
→ “shape” of atomic orbitals

Energy Levels and Sublevels
– For all other elements (with more than 1 proton and more than 1 electron), principal

energy levels (numbered 1, 2, 3,…) are further divided into energy sublevels (s, p, d, f).

– Principal Energy Level (n=1, 2, 3,…):
– Indicates the size and energy of the orbital occupied by the electron
– As n increases, the orbital becomes larger, so the electron spends more time further
away from the nucleus.
→ The further the electron is from the nucleus, the higher its energy.

principal energy level (or
shell), n: n=1,2,3,...

energy sublevels: s, p, d, and f
(or subshells)

These sublevels consist of
orbitals with specific shapes
corresponding to the probability
of finding the electron in a given
region in space.
→ An electron within a given energy sublevel doesn't orbit around the nucleus.

→ Instead, it has a high probability of being found within a given volume corresponding to
the orbital and its energy.

s orbitals: spherical
– size of the orbitals increase with principal quantum number, n

→ 1s < 2s < 3s, etc.

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p orbitals: dumbbell-shaped
– 3 types: px, py, pz (where x, y, and z indicates axis on which orbital aligns)
– The figures below shows the boundary surface representations of the p orbitals.

Note: There are also d and f orbitals, but we will not be studying these orbitals in this class.

ELECTRON CONFIGURATIONS:
– Shorthand descriptions of the arrangement of electrons within an atom

REMEMBER the following!
– s orbitals can hold 2 electrons
– a set of p orbitals can hold 6 electrons
– a set of d orbitals can hold 10 electrons
– a set of f orbitals can hold 14 electrons

Writing Electron Configurations

1. Electrons are distributed in orbitals of increasing energy, with the lowest energy orbitals
filled first. (Consider the parking garage analogy.)

2. Once an orbital has the maximum number of electrons it can hold, it is considered “filled.”
Remaining electrons must then be placed into the next higher energy orbital, and so on.

Orbitals in order of increasing energy: (See p. 299, Fig. 9.24)

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 5d < 6p

Ex. 1 He → _____ e− electron configuration for He: _____________________

Ex. 2 C → _____ e− electron config for C: ___________________________________

Ex. 3 S → _____ e− electron config for S: ___________________________________

Ex. 4 Cl → _____ e− electron config for Cl: ___________________________________

Ex. 5 K → _____ e− electron config for K: ___________________________________

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9.7 ELECTRON CONFIGURATIONS AND THE PERIODIC TABLE

Blocks of Elements

The shape of the Periodic Table actually corresponds to the order of energy sublevels.
– Consider the figure below to see how electrons for each element are distributed into energy

sublevels.

Electron configurations of atoms with many electrons can become cumbersome.
→ Core notation using Noble gas configurations:

– Elements in the last column of the Periodic Table are called “noble gases.”

– Since noble gases are at the end of each row in the Periodic Table, all of their electrons
are in filled orbitals.

[He] = 1s2
[Ne] = 1s2 2s2 2p6
[Ar] = 1s2 2s2 2p6 3s2 3p6

– Such electrons are called “core electrons” since they are more stable (less reactive)
when they belong to completely filled orbitals.

→ Noble gas electron configurations can be used to abbreviate the “core electrons” of all
elements

→ Electron configurations using Noble gas abbreviations are called “core notation”

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Electron Configurations using Core Notation:
a. Electron configuration for S using full notation: 1s2 2s2 2p6 3s2 3p4

Electron configuration for S using core notation: [Ne] 3s2 3p4
b. Electron configuration for Al using full notation: ________________________________

Electron configuration for Al using core notation: ________________________________

c. Electron configuration for Ca using full notation: ________________________________

Electron configuration for Ca using core notation: ________________________________
Note: Be able to write electron configurations for elements #1-20.

VALENCE ELECTRONS

core electrons: innermost electrons belonging to filled electron shells

valence electrons: Electrons in the outermost shell
– Since atoms want filled electron shells to be most stable, they’ll combine with other atoms

with unfilled shells (gaining or losing e–s) to get stability.

→ Valence electrons lead to chemical bonds and reactions between atoms.
→ An element’s chemical properties are determined by its number of valence electrons.

The electron configurations using core notation represent the core electrons with the Noble
gas, and the remaining electrons are the valence electrons.

valence electrons

For example, consider the electron configuration for Ca: [Ar] 4s2
– In Ca, the first 18 e–s are the core electrons, and the 2 e–s in 4s are valence electrons.

For Main Group (A) elements, Group # → # of valence electrons page 10 of 13

– Elements in Group IA: Each has 1 valence electron
– Elements in Group IIA: Each has 2 valence electrons
– Elements in Group IIIA: Each has 3 valence electrons
– Elements in Group IVA: Each has 4 valence electrons
– Elements in Group VA: Each has 5 valence electrons
– Elements in Group VIA: Each has 6 valence electrons
– Elements in Group VIIA: Each has 7 valence electrons
– Elements in Group VIIIA: Each has 8 valence electrons

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Electron-Dot (or Lewis) Symbols
– Show the atom of an element with

1. Element symbol representing the nucleus and core electrons
2. Dots representing the valence e–

Rules for writing Electron Dot Symbol
1. Write down the element symbol
2. Determine the number of valence electrons using the group number
3. Assume the atom has four sides, and distribute electrons with one electron per side
before pairing electrons.

Write the Lewis symbol for each of the following:

boron: phosphorus: oxygen: fluorine:

9.8 The Explanatory Power of the Quantum-Mechanical Model
4.7 Ions: Losing and Gaining Electrons
Although we do not delve into the quantitative aspects of the quantum-mechanical model in
this course, calculations show that atoms with the number of valence electrons as the
noble gases (2 valence electrons for helium and 8 valence electrons for all the other noble
gases) are very low in energy and are therefore stable.
Thus, elements tends to gain or lose electrons, so they are isoelectronic with (have the same
number of electrons as) a Noble gas to become more stable.
Ex. 1: Indicate the number of protons and electrons for the following:

Na Na+

S S2–

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Ex. 2: Give the formula for the ion formed by each of the following elements:

calcium: _______ chlorine: _______ magnesium: _______ barium: _______

nitrogen: _______ oxygen: _______ phosphorus: _______ potassium: _______

isoelectronic: has the same number of electrons
Thus, Na+ is isoelectronic with _______, and S2– is isoelectronic with _______.

Ex. 1: Circle all of the following ions that are isoelectronic with argon:

K+ Sr2+ Al3+ P3− Ti4+ Ca2+ O2− Mg2+

9.9 PERIODIC TRENDS: Atomic Size, Ionization Energy, and Metallic Character
Atomic Radius (or Size): distance from the nucleus to the outermost electrons

Periodic Trend for Atomic Radius

– Increases down a group: More p+, n, and e– → bigger radius

– Decreases from left to right along a period:
– Effective nuclear charge: # of protons – # of outermost electrons
– Number of p+ and e– increases, but electrons going into same orbitals.

– The higher the effective nuclear charge → smaller radius because nucleus pulling e– in

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Example: Compare an Al atom with a Cl atom below:

Trend from top to bottom → like a snowman

Trend from left to right → like a snowman

that fell to the right

METALLIC CHARACTER: Tendency to behave like a metal rather than a nonmetal

Periodic Trend for Metallic Character:
– Decreases from left to right along a period:

Metals concentrated on left-hand side of P.T., nonmetals on right-hand side
– Increases down a group: Looking at groups IVA and VA, go from nonmetals (C & N) to

semimetals (Si & As) to metals (Sn & Bi)

→ Same snowman trends as for atomic radius!

IONIZATION ENERGY: Na(g) + ionization energy → Na+(g) + e−

– Energy required to remove an electron from a neutral atom to form an ion

Periodic Trend for Ionization Energy
– Decreases down a group:

Bigger the atom, the further away electrons are from protons in nucleus

→ electrons held less tightly and are more easily removed
– Increases from left to right along a period:

– Elements with fewer (1–3) valence electrons can more easily give up electrons to gain
noble gas configuration (stability)

– Elements with more (4–7) valence electrons can more easily gain electrons to gain
noble gas configuration (stability)

Trend from top to bottom → like an upside-
down snowman

Trend from left to right → like a upside-down
snowman that fell to the right

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