MGU Science Series — High School Chemistry
Part II: Atoms, Elements, and the Periodic Table
Energy Levels · Orbitals · Electron Configurations · Valence Electrons · Lewis Dot Symbols
Chapter 7 of 28
MGU Science Series  ·  High School Chemistry  ·  Library Textbook Edition

Chapter 7:
Electrons and
Atomic Structure

Energy levels, orbitals, electron configurations, valence electrons, Lewis dot symbols, and the octet rule — the electron model that makes chemical bonding and periodic trends understandable

■ Part II: Atoms, Elements, and the Periodic Table ■

■ 9 Sections  •  Chapter 7 of 28 ■

Follows Chapter 6: Atomic Structure and Isotopes  ·  Prepares students for Chapter 8: The Periodic Table

Chapter 7 — Front Matter

Purpose and Overview

You have already learned that atoms contain protons, neutrons, and electrons. Protons and neutrons are packed tightly in the nucleus at the center of the atom. Electrons move in the space surrounding that nucleus. You have also learned that the number of protons defines which element an atom is, and that adding or removing electrons creates charged ions.

But here is a question worth sitting with: if protons determine identity, why do chemists care so much about electrons? The answer is this: protons determine what an atom is, but electrons determine what an atom does.

When two atoms bond to form a compound, electrons are involved. When a metal conducts electricity, electrons are moving. When a substance dissolves in water, electrons are being shared or attracted. When neon glows in a sign, electrons are releasing energy as light. Electrons are the active, social particles of chemistry — they reach out from atoms, interact with other atoms, and make chemical reactions possible.

This chapter builds your understanding of electrons carefully. We will move from energy level ideas to the modern electron cloud model, introduce electron configurations, explain valence electrons, and teach you to draw Lewis dot symbols. By the end of this chapter, you will have the tools needed to understand chemical bonding, periodic trends, and the behavior of real chemical substances.

Opening Question

Sodium metal reacts explosively with water. Neon — the element right after sodium in terms of electron count — barely reacts with anything at all. Both atoms have nuclei. Both have electrons. What explains the difference? By the end of this chapter, you will be able to answer that question precisely, using the language of electron configuration and valence electrons.

9Sections
10Objectives
18Key Terms
8+Worked Examples

Learning Objectives

  1. 1Explain why electrons, rather than protons or neutrons, determine the chemical behavior of an atom.
  2. 2Describe quantized energy levels in atoms, identify them by principal quantum number, and calculate the maximum electrons each level holds.
  3. 3Explain the electron cloud model and contrast it with the older orbital-path model.
  4. 4Identify sublevels (s, p, d, f) and orbitals, state their electron capacities, and describe the shapes of s and p orbitals.
  5. 5Apply the Aufbau principle, Pauli exclusion principle, and Hund's rule to write complete electron configurations for elements 1–20.
  6. 6Identify the valence electrons of any main-group element from its electron configuration and from its periodic table group number.
  7. 7Draw Lewis dot symbols for main-group elements and interpret unpaired electrons and lone pairs.
  8. 8State the octet rule and explain how atoms reach stable electron arrangements by gaining, losing, or sharing electrons.
  9. 9Connect electron arrangements to the structure of the periodic table, chemical bonding, reactivity, and energy.
  10. 10Recognize and correct common misconceptions about electron behavior, orbital paths, and the octet rule.
Chapter 7 — Contents

Sections at a Glance

Sections 7.1 – 7.3
Electrons, Energy, and the Cloud Model
Why Electrons Matter · Energy Levels · The Electron Cloud Model
Section 7.1

Why Electrons Matter

Recall the basic structure of an atom. At the center sits the nucleus, which contains positively charged protons and neutral neutrons packed tightly together. Around the nucleus, in a region of space enormously larger than the nucleus itself, move the electrons — each carrying a negative charge of −1.

The nucleus is dense and small. If an atom were the size of a football stadium, the nucleus would be approximately the size of a marble at the center. Electrons would move through all that surrounding space like tiny, fast-moving specks. Because electrons occupy the outer regions of atoms, they are the first thing one atom encounters when it approaches another. When two atoms get close enough to interact chemically, it is their electrons that meet first. This is why electrons control bonding.

What Electrons Determine

Bonding — Whether an atom will form chemical bonds, what kinds, and how many.

Ion formation — Whether an atom tends to gain, lose, or share electrons in reactions.

Reactivity — How readily an element reacts with other substances.

Conductivity — Whether electrons can move freely through a material.

Energy — The energy stored in chemical bonds and released in reactions.

Key Idea

Protons determine what an atom is. Electrons determine what an atom does. A complete understanding of chemistry at the atomic level requires understanding electron arrangements.

Practice 7.1Concept Check
  1. In your own words, explain why electrons — not protons or neutrons — are most responsible for how atoms interact chemically.
  2. Name three observable chemical phenomena (from everyday life) that involve electron behavior. For each, explain briefly what electrons are doing.
  3. If you removed all the electrons from a sodium atom, could it form chemical bonds with other atoms? Explain your reasoning.
Section 7.2

Energy Levels

Early models imagined electrons as tiny balls orbiting the nucleus like planets around the sun. This picture was useful but incorrect in important ways. Experiments showed that electrons in atoms can only have certain specific amounts of energy — not just any energy at all. This is the concept of quantized energy. An electron in an atom cannot have any energy it chooses. It must occupy one of a set of allowed energy states called energy levels.

Think of energy levels like the rungs of a ladder. You can stand on the first rung, the second rung, the third rung — but not between rungs. Electrons work the same way. They can have the energy of level 1, level 2, level 3, and so on, but not the energies in between.

Energy Level (n)Distance from NucleusRelative EnergyMax Electrons (2n²)
n = 1ClosestLowest2
n = 2FartherHigher8
n = 3Farther stillHigher still18
n = 4Farthest (common)Highest (common)32
Key Terms — Energy Levels

Principal quantum number (n) — the integer that labels an energy level; n = 1, 2, 3 …

Ground state — the lowest-energy arrangement of electrons in an atom; electrons occupy the lowest available levels.

Excited state — an arrangement in which one or more electrons have absorbed energy and moved to a higher level.

2n² rule — the maximum number of electrons in energy level n equals 2n².

Electrons, Light, and Flame Colors

When electrons absorb energy and jump to higher levels, then fall back down, they release that energy as light. Different elements have unique sets of energy levels, so they release light at unique wavelengths. This is why copper produces green flames, sodium produces yellow-orange ones, and neon signs glow red-orange. Each color is a signature of a specific set of electron energy jumps.

Worked Example 7.2.1
Distributing electrons among energy levels
  • ProblemSodium has 11 electrons. How are those electrons distributed among the first three energy levels?
  • Step 1Fill n = 1 first. It holds 2 electrons. Electrons placed: 2    Remaining: 11 − 2 = 9
  • Step 2Fill n = 2 next. It holds 8 electrons. Electrons placed: 8    Remaining: 9 − 8 = 1
  • Step 3The last electron goes into n = 3. 1 electron in n = 3
Answer: Sodium distributes as 2 | 8 | 1 across energy levels n = 1, 2, 3.
Practice 7.2Energy Level Practice
  1. Calculate the maximum number of electrons for energy levels n = 1 through n = 4 using the formula 2n². Show your work for each.
  2. Distribute the electrons of aluminum (13 electrons) among energy levels. Write the result in the form used in Worked Example 7.2.1.
  3. Distribute the electrons of chlorine (17 electrons) among energy levels. Write the distribution.
  4. An electron in n = 3 absorbs energy and jumps to n = 4. Is the electron now in a ground state or excited state? What will it do next, and what observable event might result?
  5. Two atoms both have electrons in energy level n = 3. Does this mean they are the same element? Explain.
Section 7.3

The Electron Cloud Model

Early models, especially Bohr's model from the early 1900s, pictured electrons moving in circular orbits around the nucleus like planets orbiting the sun. Bohr's model was a significant improvement and successfully explained some observations about hydrogen — but it had serious limitations for atoms with more than one electron.

More careful experiments and mathematical work revealed something surprising: electrons do not travel in neat, defined paths. In fact, it is not possible to know exactly where an electron is and exactly how fast it is moving at the same time. This is a fundamental feature of nature at the atomic scale, not a limitation of our instruments.

Because electrons do not have defined paths, chemists describe them in terms of probability. Instead of asking "where is this electron right now?", chemists ask "where is this electron most likely to be found?" The answer defines a region of space around the nucleus called the electron cloud.

Visualizing the Electron Cloud

Imagine shaking a can of spray paint over a sheet of paper many times, starting every shake from the same point. The paint would build up in some areas more than others. If you photographed the result, the dense regions would show where the paint landed most often. The electron cloud works similarly — where the cloud is dense, the probability of finding the electron is high. Where the cloud is thin or absent, the probability is low.

FeatureOld Orbit ModelModern Electron Cloud Model
Electron pathFixed circular orbitNo defined path; probability region
Electron locationExactly known at any momentDescribed as probability of being found
Atomic edgeSharp, defined boundaryCloud gradually thins; no sharp edge
Works for all atoms?Only hydrogen (approximately)All elements
Practice 7.3Electron Cloud Model
  1. Explain in your own words what the electron cloud represents. Is it a physical object? What does its density tell you?
  2. Why can't scientists simply watch an electron move through its orbit and report its exact position? What does this tell us about the nature of electrons?
  3. A student says "atoms have sharp, defined edges where the electrons stop." Is this accurate according to the electron cloud model? Explain.
  4. In what way is the electron cloud model more accurate than the Bohr model for atoms like carbon, oxygen, or iron?
Sections 7.4 – 7.5
Orbitals, Sublevels, and Electron Configurations
Orbitals as Probability Regions · Sublevels s, p, d, f · Writing Electron Configurations
Section 7.4

Orbitals as Regions of Probability

The energy levels described in Section 7.2 are divided into smaller units called sublevels (sometimes called subshells). Within each sublevel, electrons occupy regions of space called orbitals. An orbital is a specific three-dimensional region of space within an atom where there is a high probability of finding an electron. Each orbital can hold a maximum of 2 electrons.

SublevelNumber of OrbitalsMax ElectronsShape
s12Sphere
p36Dumbbell (3 orientations)
d510Complex multi-lobe shapes
f714Complex multi-lobe shapes

Which Sublevels Belong to Which Energy Level

Energy Level (n)Sublevels PresentTotal Electrons (max)
n = 11s2
n = 22s, 2p8
n = 33s, 3p, 3d18
n = 44s, 4p, 4d, 4f32

s orbitals are spherical. An electron in an s orbital is equally likely to be found in any direction around the nucleus, so the probability cloud forms a sphere. Every energy level has exactly one s orbital.

p orbitals have a dumbbell shape — two lobes extending in opposite directions from the nucleus. Because there are three p orbitals in every p sublevel, they point in three perpendicular directions. Together, the three p orbitals cover all directions in three-dimensional space.

For this course, you do not need to memorize the detailed shapes of d and f orbitals. Understanding s and p orbitals well will prepare you for everything in high school chemistry, including molecular geometry in Chapter 12.

Common Confusion — Orbital vs. Energy Level

An orbital is a single region of space that holds at most 2 electrons. An energy level is a collection of sublevels. An sublevel is a collection of one or more orbitals. The three terms are not interchangeable. n = 2 is an energy level; 2p is a sublevel; one of the three 2p regions is a single orbital.

Practice 7.4Orbitals and Sublevels
  1. How many orbitals are in the 3p sublevel? How many electrons can it hold total?
  2. What is the maximum number of electrons the entire n = 3 energy level can hold? List the sublevels that contribute to this total and how many electrons each holds.
  3. Compare the shape of an s orbital to a p orbital. How do these shapes affect where electrons are likely to be found relative to the nucleus?
  4. A student says "the 2s and 3s orbitals have the same shape, so they must be the same energy." Identify the error in this reasoning.
  5. Why does the p sublevel have 3 orbitals rather than 1 or 5?
Section 7.5

Electron Configurations

An electron configuration is a written description of how the electrons in an atom are distributed among energy levels and sublevels. Electron configurations follow a set of rules. Once you know those rules, you can write the electron configuration for any element.

Three Rules for Filling Orbitals

Aufbau Principle: Electrons occupy the lowest available energy sublevel before filling higher ones. Build the configuration from the bottom up.

Pauli Exclusion Principle: Each orbital holds a maximum of 2 electrons. No two electrons in the same atom can be in the same orbital with the same spin.

Hund's Rule: When electrons fill orbitals in the same sublevel, they spread out one per orbital before any orbital gets a second electron. Electrons prefer separate orbitals to minimize repulsion.

The Filling Order

Sublevels fill in this order from lowest to highest energy:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p …
Notice that 4s fills before 3d. The periodic table is organized to reflect this order — use it as a guide when you need to remember the sequence.

Worked Example 7.5.1
Writing the electron configuration for hydrogen (1 electron)
  • ProcessOne electron goes into the lowest available sublevel, which is 1s.
Hydrogen: 1s¹
Worked Example 7.5.2
Writing the electron configuration for carbon (6 electrons)
  • Step 11s holds 2 electrons: place 2 there. 1s² (2 placed, 4 remaining)
  • Step 22s holds 2 electrons: place 2 there. 2s² (2 placed, 2 remaining)
  • Step 32p can hold 6 but only 2 remain: place 2. 2p² (2 placed, 0 remaining)
  • NoteHund's rule: the 2 electrons in 2p occupy two different 2p orbitals rather than sharing one.
Carbon: 1s² 2s² 2p²
Worked Example 7.5.3
Writing the electron configuration for phosphorus (15 electrons)
  • Fill 1s² → 2 placed, 13 remaining 2s² → 2 placed, 11 remaining 2p⁶ → 6 placed, 5 remaining 3s² → 2 placed, 3 remaining 3p³ → 3 placed, 0 remaining
Phosphorus: 1s² 2s² 2p⁶ 3s² 3p³
Worked Example 7.5.4
Writing the electron configuration for calcium (20 electrons)
  • Fill 1s² → 18 remaining 2s² → 16 remaining 2p⁶ → 10 remaining 3s² → 8 remaining 3p⁶ → 2 remaining 4s² → 0 remaining
  • Note4s fills before 3d. Calcium has no d electrons.
Calcium: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
Student Warning — The 4s/3d Order

The 4s sublevel fills before the 3d sublevel because 4s is slightly lower in energy than 3d for neutral atoms. Do not assume that a higher n value always means later filling. When in doubt, use the periodic table as your guide — elements in Period 4 of the s block (potassium, calcium) fill 4s before entering the d block.

Practice 7.5Electron Configurations

Part A: Write the complete electron configuration for each element. Show all sublevels.

  1. Helium (2 electrons)
  2. Boron (5 electrons)
  3. Oxygen (8 electrons)
  4. Neon (10 electrons)
  5. Magnesium (12 electrons)
  6. Chlorine (17 electrons)
  7. Argon (18 electrons)
  8. Potassium (19 electrons)
  9. Silicon (14 electrons)
  10. Sulfur (16 electrons)

Part B: Analysis questions.

  1. The configuration for an element is: 1s² 2s² 2p⁶ 3s². How many total electrons? What element is this? What is the highest energy level occupied?
  2. The configuration for an element is: 1s² 2s² 2p³. What element is this? How many unpaired electrons does it have? Explain how Hund's rule determines your answer.
  3. Write the configuration for an element in Period 3 with 6 valence electrons. Identify the element.
  4. True or false: 3d fills before 4s. Explain your answer fully.
Sections 7.6 – 7.8
Valence Electrons, Lewis Dot Symbols, and Stability
Valence Electrons · Lewis Dot Symbols · The Octet Rule · Stable Electron Arrangements
Section 7.6

Valence Electrons

You now know how to write a complete electron configuration for an atom. But not all electrons participate equally in chemical reactions. The electrons most involved in bonding and chemical behavior are the ones in the outermost occupied energy level — the level with the highest principal quantum number (n) that actually contains electrons. These outermost electrons are called valence electrons.

For most main-group elements (the s block and p block), valence electrons are the electrons in the outermost s and p sublevels. Electrons that are not valence electrons are called core electrons. They are in lower energy levels, tightly held, and do not participate in bonding under ordinary chemical conditions.

How to Find Valence Electrons

1. Write the electron configuration of the atom.

2. Identify the highest n value that appears in the configuration.

3. Count all electrons in sublevels with that n value.

Worked Example 7.6.1
Valence electrons of sodium (11 electrons)
  • Config1s² 2s² 2p⁶ 3s¹
  • FindHighest n = 3. Only 3s is present. It contains 1 electron.
Sodium has 1 valence electron.
Worked Example 7.6.2
Valence electrons of sulfur (16 electrons)
  • Config1s² 2s² 2p⁶ 3s² 3p⁴
  • FindHighest n = 3. Sublevels: 3s² and 3p⁴. Total: 2 + 4 = 6 valence electrons.
Sulfur has 6 valence electrons.

Valence Electrons and the Periodic Table

One of the most powerful features of the periodic table is that elements in the same group have the same number of valence electrons. This is not a coincidence — it is the reason those elements behave similarly in chemical reactions.

GroupValence ElectronsExample Elements
Group 1 (Alkali metals)1Li, Na, K, Cs
Group 2 (Alkaline earth metals)2Be, Mg, Ca, Ba
Group 133B, Al, Ga
Group 144C, Si, Ge
Group 155N, P, As
Group 166O, S, Se
Group 17 (Halogens)7F, Cl, Br, I
Group 18 (Noble gases)8 (He = 2)He, Ne, Ar, Kr
Practice 7.6Valence Electron Practice

Part A: Count the valence electrons for each element.

  1. Lithium (Li)
  2. Carbon (C)
  3. Fluorine (F)
  4. Argon (Ar)
  5. Aluminum (Al)
  6. Sulfur (S)
  7. Barium (Ba)
  8. Iodine (I)
  9. Selenium (Se)
  10. Cesium (Cs)

Part B: Reasoning questions.

  1. An element has 3 valence electrons. What group does it belong to? Name two elements in that group.
  2. An element has 6 valence electrons. How many electrons would it need to gain to complete its outer shell? What ion charge would it most likely form?
  3. Without writing a full electron configuration, predict the number of valence electrons for the element in Period 4, Group 15.
  4. Elements A and B are in the same group. Element A has the configuration 1s² 2s² 2p⁶ 3s¹. What is the valence electron count of element B? What can you predict about B's chemical behavior compared to A?
Section 7.7

Lewis Dot Symbols

Chemists use a simple visual shorthand to represent the valence electrons of an atom called a Lewis dot symbol. It was developed by American chemist Gilbert N. Lewis in the early twentieth century and remains one of the most widely used tools in chemistry.

A Lewis dot symbol consists of the chemical symbol of the element, surrounded by dots — each dot representing one valence electron.

Rules for Drawing Lewis Dot Symbols

1. Write the element's chemical symbol.

2. Count the number of valence electrons.

3. Place dots one at a time on the four sides of the symbol (top, bottom, left, right). Each side is one position.

4. Once each side has one dot, begin pairing — add a second dot to each position that needs it.

5. Continue until all valence electrons are placed. Follow Hund's rule: spread out before pairing.

ElementSymbolValence e⁻UnpairedLone Pairs
HydrogenH110
HeliumHe201
LithiumLi110
BerylliumBe220
BoronB330
CarbonC440
NitrogenN531
OxygenO622
FluorineF713
NeonNe804
SodiumNa110
MagnesiumMg220
AluminumAl330
SiliconSi440
PhosphorusP531
SulfurS622
ChlorineCl713
ArgonAr804
What Lewis Dot Symbols Tell Us

Unpaired electrons (single dots) are available to form bonds. An atom with 3 unpaired electrons can form 3 bonds; an atom with 1 unpaired electron can form 1 bond. Lone pairs (paired dots) are not typically used for bonding in the same way, though they influence molecular geometry. The drive to pair unpaired electrons with unpaired electrons from other atoms is the basis of covalent bonding.

Practice 7.7Lewis Dot Practice

Part A: Draw the Lewis dot symbol for each element. State the number of unpaired electrons and lone pairs.

  1. Hydrogen (H)
  2. Beryllium (Be)
  3. Nitrogen (N)
  4. Chlorine (Cl)
  5. Potassium (K)
  6. Phosphorus (P)
  7. Oxygen (O)
  8. Neon (Ne)
  9. Bromine (Br)
  10. Silicon (Si)

Part B: Interpretation questions.

  1. An element has a Lewis dot symbol with 2 unpaired electrons and 2 lone pairs. How many valence electrons does it have total? What group does it belong to? Name one element with this symbol.
  2. An element has a Lewis dot symbol with 1 unpaired electron and 3 lone pairs. How many bonds can it typically form? Name one element with this symbol.
  3. An element has a Lewis dot symbol with 4 unpaired electrons and 0 lone pairs. How many bonds would it need to form to reach an octet? Name one element with this symbol.
  4. Why do the Lewis dot symbols of sodium and potassium look identical, even though these elements differ greatly in size and mass?
  5. Two elements both have 1 unpaired electron in their Lewis dot symbols. Are they necessarily the same element? Explain.
Section 7.8

Stable Electron Arrangements

Atoms do not bond randomly. They bond because bonding leads to a more stable, lower-energy arrangement. The most stable electron arrangement for most atoms involves having a full outer energy level. For most atoms, a full outer level means 8 valence electrons. This stability goal is called the octet rule: atoms tend to gain, lose, or share electrons until they have 8 valence electrons in their outermost level.

Three Ways Atoms Reach Stability

Gaining electrons (forming anions) — Nonmetals on the right side often gain electrons. Chlorine has 7 valence electrons; gaining 1 gives it 8, forming Cl⁻ with the argon configuration.

Losing electrons (forming cations) — Metals on the left side often lose electrons. Sodium has 1 valence electron in n = 3; losing it exposes n = 2 with 8 electrons, forming Na⁺ with the neon configuration.

Sharing electrons (covalent bonding) — Nonmetals share electrons with each other. Each atom counts shared electrons toward its own total, helping both atoms reach 8.

SituationApplies ToExampleResult
Duet ruleH, He (n = 1 holds only 2)H + H → H₂Each H counts 2 electrons
Octet ruleMost main-group elementsCl + Na → NaClBoth reach 8 valence electrons
Already stableNoble gases (full outer shell)Ne, Ar, KrDo not react under normal conditions
Student Warning — The Octet Rule Has Exceptions

The octet rule is a useful guideline, not an absolute law. Hydrogen and helium work with a duet. Boron commonly has only 6 electrons around it in compounds. Phosphorus and sulfur can have more than 8 valence electrons in some compounds. Transition metals often do not follow the rule at all. Use the octet rule as a helpful first prediction, not as an unbreakable rule.

Practice 7.8Stability and the Octet Rule
  1. Magnesium has 2 valence electrons. How many electrons would it need to lose to reach a stable (noble gas) configuration? What ion would it form? Which noble gas configuration does it match?
  2. Oxygen has 6 valence electrons. How many electrons would it need to gain to reach an octet? What ion would it form?
  3. Carbon has 4 valence electrons. Would carbon more likely gain or lose electrons to reach an octet? Or does it take a different route? Explain.
  4. Why are noble gases chemically unreactive while the halogens (Group 17) are among the most reactive elements?
  5. Using Lewis dot symbols, explain why two hydrogen atoms bond together to form H₂. What stability goal does each atom achieve?
Section 7.9
Electrons and Chemical Behavior
Connecting Electron Arrangements to the Periodic Table, Bonding, Reactivity, and Energy
Section 7.9

Electrons and Chemical Behavior

You have now studied energy levels, the electron cloud model, orbitals, electron configurations, valence electrons, Lewis dot symbols, and the octet rule. In this final section, we step back and connect all of these ideas to the larger story of chemistry.

Electrons Explain the Periodic Table

Elements are arranged by atomic number — and as atomic number increases, electrons are added one by one to energy levels and sublevels. When a new energy level begins filling, a new period of the table starts. When the same sublevel fills across a row, elements in that row share the same general electron arrangement pattern. When a column of elements all have the same number of valence electrons, they share similar chemical behavior. The periodic table is essentially a map of electron configurations.

Electrons Explain Bonding

Atoms bond to achieve full outer energy levels. Ionic bonds form when one atom transfers electrons to another, creating oppositely charged ions that attract. Covalent bonds form when atoms share electrons, both counting the shared pair toward their valence totals. The type of bond, the number of bonds, and the geometry of molecules all trace back to electron arrangements.

Electrons Explain Reactivity

Some atoms react violently, some react slowly, and some barely react at all. Alkali metals (Group 1) have 1 valence electron — easily lost, producing a very stable configuration. They react with almost anything. Noble gases have full outer shells and nothing to gain or lose, so they do not react. Halogens (Group 17) have 7 valence electrons — one short of a full shell. They aggressively attract electrons from other atoms. Reactivity follows directly from the electron count.

Electrons Explain Energy

When electrons move between energy levels, energy is absorbed or released. When bonds form, energy is released. When bonds break, energy is absorbed. The energy stored in fuels — gasoline, natural gas, hydrogen — is ultimately stored in the electron arrangements of molecules. Burning releases that energy because the products have lower-energy electron arrangements than the reactants.

The Invisible Made Useful

Electrons cannot be seen directly — they are smaller than any traditional microscope can image. Yet chemists have built an extraordinarily detailed and reliable model of electron behavior, one that accurately predicts the colors of light emitted by atoms, the angles of bonds in molecules, the conductivity of metals, the strength of acids, the solubility of compounds, and countless other properties. This is what makes the electron model one of the great achievements of human understanding.

Practice 7.9Connecting Concepts
  1. Explain in your own words why elements in the same group have similar chemical behavior. Use the term "valence electrons" in your answer.
  2. Potassium reacts vigorously with water. Argon, the element just before potassium in atomic number, does not react with water at all. Explain this difference using electron configurations and the octet rule.
  3. When wood burns, energy is released. Using what you know about electrons and bonds, describe in general terms why energy is released during combustion.
  4. The element bromine (Br, Group 17) reacts similarly to chlorine (Cl, Group 17). What electron feature do they share? Why does this explain their similar reactivity?
Chapter 7 — Back Matter

Chapter Summary

Key Terms Review

Aufbau principle
The rule that electrons fill the lowest available energy sublevel before occupying higher ones. Build configurations from the bottom up.
core electrons
Electrons in inner energy levels that do not participate in bonding; also called inner electrons. They shield valence electrons from the nucleus.
electron cloud
A probability-based map of where an electron is likely to be found around a nucleus. Dense regions indicate high probability; thin regions indicate low probability.
electron configuration
A notation describing how electrons are distributed among the energy levels and sublevels of an atom. Example: 1s² 2s² 2p⁶.
energy level
One of the allowed, quantized energy states an electron can occupy in an atom. Labeled by the principal quantum number n = 1, 2, 3, …
excited state
An electron arrangement in which one or more electrons have absorbed energy and moved to a higher energy level than the ground state.
ground state
The lowest-energy electron arrangement of an atom; electrons occupy the lowest available energy levels.
Hund's rule
Electrons fill orbitals in the same sublevel one at a time before any orbital gets a second electron. Electrons spread out before pairing.
Lewis dot symbol
A notation showing the chemical symbol surrounded by dots representing valence electrons. Each dot is one valence electron.
lone pair
A pair of valence electrons that is not involved in bonding; also called a nonbonding pair. Lone pairs influence molecular geometry.
octet rule
The tendency of main-group atoms to gain, lose, or share electrons until they have 8 valence electrons. A useful guideline with known exceptions.
orbital
A three-dimensional region of space within an atom where there is a high probability of finding an electron. Each orbital holds at most 2 electrons.
Pauli exclusion principle
No two electrons in the same atom can occupy the same orbital with the same spin. Each orbital holds at most 2 electrons.
principal quantum number (n)
The integer that labels an energy level (n = 1, 2, 3, …). Higher n means higher energy and greater average distance from the nucleus.
quantized energy
Energy that can only take specific, allowed values rather than any value in a continuous range. Electrons in atoms have quantized energies.
sublevel
A subdivision of an energy level, labeled s, p, d, or f. Each sublevel contains a specific number of orbitals and holds a specific maximum number of electrons.
unpaired electron
A valence electron that occupies an orbital by itself; available to form a bond with another atom. Counted as single dots in Lewis dot symbols.
valence electrons
Electrons in the outermost occupied energy level of an atom. These electrons participate in bonding and determine chemical behavior.

Concept Check

Answer each question in a complete sentence or two. These test understanding, not calculation.

  1. What is the difference between an energy level and a sublevel? Give an example of each.
  2. Why do electrons in atoms have quantized energy rather than any energy value they choose?
  3. Explain in your own words what an electron cloud represents. Is it a physical object? What does its density tell you?
  4. Why do p sublevels hold more electrons than s sublevels?
  5. What does it mean for an electron to be in an "excited state"? How does it return to the ground state, and what happens when it does?
  6. Why do elements in the same group of the periodic table have similar chemical behavior?
  7. A student says: "Valence electrons are the most important electrons because they are the biggest." Identify the error and correct it.
  8. Why is neon chemically unreactive while fluorine (the element just before it in atomic number) is extremely reactive?
  9. Explain why sodium tends to lose 1 electron in chemical reactions rather than 2 or 3.
  10. In what way is the octet rule a consequence of electron configuration, rather than an arbitrary rule?

Common Misconceptions

Misconception 1 — Electrons orbit the nucleus like planets
"Electrons travel in fixed circular orbits around the nucleus." Correction: Electrons do not follow defined paths. They occupy regions of probability called orbitals. The location of an electron at any moment cannot be precisely determined. The orbit image is a useful introductory picture but does not reflect how electrons actually behave.
Misconception 2 — Higher energy levels are always farther from the nucleus
"An electron in n = 3 is always farther from the nucleus than an electron in n = 2." Correction: Higher n values correspond to greater average distances from the nucleus, but electrons in any level can be found at various distances. The statement is approximately true and useful as a guide, but not precisely exact.
Misconception 3 — Electron configurations never change
"The electron configuration of an element is the same whether it is a neutral atom or part of a compound." Correction: Electron configurations describe neutral atoms. When atoms bond, they gain, lose, or share electrons — so their electron arrangement changes. The neutral atom configuration is a starting point, not a fixed description for all situations.
Misconception 4 — The octet rule is absolute
"All atoms always end up with exactly 8 valence electrons after bonding." Correction: The octet rule is a useful guideline with real exceptions. Hydrogen and helium use a duet (2 electrons). Boron commonly forms compounds with 6 electrons. Phosphorus and sulfur can have more than 8. Transition metals often do not follow the rule at all.
Misconception 5 — Lone pairs are wasted electrons
"Lone pairs are leftover electrons that didn't bond — they don't do anything." Correction: Lone pairs are valence electrons not involved in bonding between atoms, but they are not wasted. They influence molecular geometry, contribute to polarity, and can participate in certain bonding types. They are a full part of the electron arrangement.
Misconception 6 — Valence electrons are weakly held because they are "outer"
"Valence electrons are the outermost electrons, so they must be very loosely held by all atoms." Correction: How tightly valence electrons are held depends on nuclear charge and shielding by core electrons. Fluorine, for example, holds its valence electrons very tightly despite them being outer electrons. This variation is the basis of periodic trends you will study in Chapter 8.

Skill Practice

Part AEnergy Level Calculations
  1. Calculate the maximum number of electrons in each energy level from n = 1 to n = 4. Show your work using 2n².
  2. An atom has 20 electrons. Distribute them among energy levels, showing how many go in each level (n = 1, 2, 3, 4).
  3. Which energy level first has a d sublevel? How many electrons does that d sublevel hold?
Part BElectron Configurations

Write the complete electron configuration for each element.

  1. Hydrogen (H)
  2. Helium (He)
  3. Lithium (Li)
  4. Carbon (C)
  5. Nitrogen (N)
  6. Fluorine (F)
  7. Neon (Ne)
  8. Sodium (Na)
  9. Aluminum (Al)
  10. Silicon (Si)
  11. Phosphorus (P)
  12. Chlorine (Cl)
  13. Argon (Ar)
  14. Potassium (K)
  15. Calcium (Ca)
Part CValence Electrons and Lewis Dot Symbols
  1. For each element in Part B, identify the number of valence electrons and draw the Lewis dot symbol.
  2. Which elements in Part B have the same Lewis dot symbol pattern? Group them and explain why they share the same pattern.
  3. Using Lewis dot symbols, show why a chlorine atom (Cl) and a sodium atom (Na) might transfer an electron to reach stable arrangements. Show the resulting ions.

Review Questions

Short Answer
  1. What is the maximum number of electrons that can fit in energy level n = 3? Show how you calculated this.
  2. Write the electron configuration for chlorine and identify its valence electrons.
  3. How many orbitals are in the 3p sublevel? How many electrons can it hold?
  4. What is the difference between a lone pair and an unpaired electron in a Lewis dot symbol? Which type is more available for bonding?
  5. An atom has the configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹. Identify the element. How many valence electrons does it have? Draw its Lewis dot symbol.
  6. Explain what the Aufbau principle states and give one example of how it influences where electrons are placed in a configuration.
  7. How does writing a Lewis dot symbol help predict how many bonds an atom can form?
Extended Response
  1. Compare the electron configurations of sodium (Na) and potassium (K). How are they similar? How are they different? What does the similarity predict about their chemical behavior?
  2. A student writes the electron configuration for nitrogen as 1s² 2s² 2p⁵ instead of 1s² 2s² 2p³. What is the error? What principle was violated? Correct the configuration and explain what the correct version says about nitrogen's chemistry.
  3. Describe what happens to a strontium atom (Sr, 38 electrons) in terms of electron arrangement when it forms the Sr²⁺ ion. Write the electron configuration of the ion and compare it to the configuration of a nearby noble gas.
  4. Arsenic (As) is in Group 15, Period 4. Without writing a detailed configuration table: (a) How many valence electrons does arsenic have? (b) Draw its Lewis dot symbol. (c) How many bonds would you predict arsenic can form? (d) What noble gas configuration is arsenic "trying to reach"?

Chapter 7 Checkpoint

Chapter 7 Checkpoint
Complete without assistance — assesses all major skills and concepts from Chapter 7
Section 1 — Vocabulary and Concepts
  1. Explain the difference between an energy level and a sublevel. Give an example of each.
  2. What does it mean for energy to be quantized? How does this apply to electrons in atoms?
  3. What does the electron cloud model tell us about where an electron is? Is the electron cloud a physical object?
  4. State the octet rule. Name one element that is an exception and explain why.
Section 2 — Electron Configuration
  1. Write the complete electron configuration for magnesium (Mg, 12 electrons).
  2. Write the complete electron configuration for sulfur (S, 16 electrons).
  3. Write the complete electron configuration for calcium (Ca, 20 electrons).
  4. An element has the configuration 1s² 2s² 2p⁶ 3s² 3p³. Identify the element and state how many valence electrons it has.
Section 3 — Valence Electrons and Lewis Dots
  1. How many valence electrons does each of the following have? (a) Nitrogen, (b) Oxygen, (c) Chlorine, (d) Argon
  2. Draw Lewis dot symbols for nitrogen (N), oxygen (O), and fluorine (F). State the number of unpaired electrons in each.
  3. Explain why nitrogen can form 3 bonds while oxygen typically forms 2, based on their Lewis dot symbols.
Section 4 — Stability and Bonding
  1. Calcium has 2 valence electrons. What ion does it form, and why? Which noble gas configuration does it match?
  2. Fluorine has 7 valence electrons. What ion does it form, and why? Which noble gas configuration does it match?
  3. Why do the elements lithium, sodium, and potassium all have similar chemical properties despite being very different in atomic mass and size?
Section 5 — Applying Electron Concepts
  1. A student writes the configuration for argon as 1s² 2s² 2p⁶ 3s² 3p⁵. What is wrong? Write the correct configuration. How many valence electrons does argon actually have?
  2. An unknown element has 5 valence electrons and is in Period 2. Identify the element, write its electron configuration, and draw its Lewis dot symbol.
  3. Explain why sodium (Na) reacts readily with water while argon (Ar) does not, using the concept of valence electrons and electron stability.
Looking Ahead
Chapter 8: The Periodic Table

Chapter 7 revealed how electrons are arranged in atoms — in energy levels, sublevels, and orbitals — and how valence electrons determine chemical behavior. In Chapter 8, you will see how all of these electron arrangements are organized into the structure of the periodic table. You will learn how atomic number and electron configuration determine every element's position on the table, how periods and groups emerge directly from the pattern of electron filling, and how the table is divided into metals, nonmetals, and metalloids. Chapter 8 will also introduce periodic trends — predictable changes in atomic size, ionization energy, and electron affinity — that connect everything you have learned about electron arrangement to the larger picture of how elements behave in the real world.

MGU Library Connection
  • Chapter 6: Atomic Structure and Isotopes — protons, neutrons, electrons, atomic number, and ions (prerequisite for this chapter)
  • Chapter 8: The Periodic Table — periodic trends, groups, periods, metals, nonmetals, metalloids (follows from this chapter)
  • Chapter 9: Ionic Bonding and Ionic Compounds — valence electrons and electron transfer (uses concepts from Sections 7.6–7.8)
  • Chapter 10: Covalent Bonding and Molecular Compounds — electron sharing and Lewis structures (uses Lewis dot concepts from Section 7.7)
  • Chapter 12: Molecular Geometry and Intermolecular Forces — orbital shapes and lone pairs directly applied
  • Appendix C: Periodic Table Reference — blocks, groups, periods, and electron configuration patterns