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Chemistry - Electron Configurations

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    Study Guide Chemistry – Electron Configurations 1. Quiz: Valence Electrons 1. Question Valence electrons are located in Answer Choices • energy levels closest to the nucleus • all energy levels • energy levels with the highest energy Correct Answer energy levels with the highest energy Why This Is Correct Valence electrons are the electrons found in the outermost energy level of an atom, which is also the highest - energy energy level . These electrons are the most important because they: • Participate in chemical bonding • Determine an element’s chemical reactivity • Explain why elements in the same group have similar properties 2. Question When all the electrons in an atom are located in the lowest energy levels and sublevels, the electrons are

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    Study Guide Answer Choices • in an excited state • in the ground state • in a state of chaos Correct Answer in the ground state Why This Is Correct The ground state is the lowest - energy, most stable arrangement of electrons in an atom. When electrons occupy the lowest possible energy levels and sublevels, the atom has minimum energy. An excited state occurs only when one or more electrons absorb energy and move to higher energy levels. “State of chaos” is not a scientifically recognized term in atomic theory. 3. Question Which of the following elements would be expected to form an ion with a +1 charge ? Answer Choices • Li: 1s² 2s¹ • F: 1s² 2s² 2p⁵ • Ne: 1s² 2s² 2p⁶ Correct Answer Li: 1s² 2s¹ Why This Is Correct Lithium (Li) has one valence electron in its outermost energy level (2s¹). To achieve a stable electron configuration, lithium tends to lose this one electron , forming a +1 ion (Li ⁺ ) .

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    Study Guide • Lithium (Li) → loses 1 electron → +1 charge • Fluorine (F) has 7 valence electrons and tends to gain 1 electron, forming a – 1 ion • Neon (Ne) already has a full outer shell and does not usually form ions Therefore, lithium is the element expected to form an ion with a +1 charge . 4. Question When forming a positive ion, electrons are lost from which level? Answer Choices • the level closest to the nucleus • the level farthest from the nucleus • any level within the electron cloud Correct Answer the level farthest from the nucleus Why This Is Correct When an atom forms a positive ion (cation) , it loses electrons . Electrons are removed from the outermost energy level , which is the level farthest from the nucleus . These outer electrons are: • held less tightly by the nucleus • higher in energy • easier to remove than inner (core) electrons Removing inner electrons would require much more energy and usually does not happen in normal chemical reactions. Example: Sodium (Na) loses its single outer electron to form Na ⁺ , but its inner electrons remain unchanged.

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    Study Guide 5. Question Which represents an excited state electronic configuration of electrons? Answer Choices • 1s² 2s² 2p¹ • 1s² 2s² 2p⁶ 3s¹ • 1s² 2s² 3s² Correct Answer 1s² 2s² 3s² Why This Is Correct An excited state occurs when one or more electrons absorb energy and move to a higher energy level instead of filling the lowest possible levels first. • In 1s² 2s² 3s² , electrons occupy the 3s orbital even though the 2p orbital is not filled . • This violates the Aufbau principle for ground states, indicating the atom is in an excited state . Why the other options are NOT excited states: • 1s² 2s² 2p¹ → Ground state (electrons fill lower energy orbitals first) • 1s² 2s² 2p⁶ 3s¹ → Ground state (lowest - energy correct filling order) 6. Question Isoelectronic ions are ions that have Answer Choices • different numbers of electrons • two extra electrons • the same number of electrons

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    Study Guide Correct Answer the same number of electrons Why This Is Correct Isoelectronic ions are defined as ions (or atoms) that contain the same total number of electrons , even though they may be different elements or have different charges. For example: • Na ⁺ has 10 electrons • Ne has 10 electrons • F ⁻ has 10 electrons Even though their numbers of protons and charges differ, they are isoelectronic because their electron counts are identical . 7. Question Which electron configuration represents the fluoride ion, F ⁻ ¹ ? Answer Choices • 1s² 2s² 2p⁴ • 1s² 2s² 2p⁵ • 1s² 2s² 2p⁶ Correct Answer 1s² 2s² 2p⁶ Why This Is Correct A neutral fluorine atom has 9 electrons with the configuration 1s² 2s² 2p⁵ . When fluorine forms the fluoride ion (F ⁻ ) , it gains one electron , giving it a total of 10 electrons . That extra electron fills the 2p sublevel, resulting in the stable (noble gas – like) configuration: 1s² 2s² 2p⁶ This configuration matches the correct answer.

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    Study Guide 8. Question Which electron configuration represents an atom with its electrons in the ground state ? Answer Choices • 1 s² 2 s² 2 p⁵ 3 s¹ • 1 s² 2 s² 2 p⁵ 3 s² • 1 s² 2 s² 2 p⁶ 3 s¹ Correct Answer 1 s² 2 s² 2 p⁶ 3 s¹ Why This Is Correct An atom is in its ground state when all electrons occupy the lowest possible energy levels before moving to higher ones. • Electrons must completely fill lower - energy orbitals (like 2p ) before entering higher - energy orbitals (like 3s ). • In 1 s² 2 s² 2 p⁶ 3 s¹ , the 2p sublevel is fully filled, and only then does an electron occupy the 3s sublevel. • This configuration follows the Aufbau principle , so it represents the ground state. Why the Others Are Incorrect • 1 s² 2 s² 2 p⁵ 3 s¹ → An electron is promoted to 3s before 2p is full → excited state • 1 s² 2 s² 2 p⁵ 3 s² → Still skips filling 2p completely → excited state Therefore, the correct ground - state configuration is 1 s² 2 s² 2 p⁶ 3 s¹ . 2. The Periodic Table The periodic table is more than just a list of elements. Its layout shows us how electrons are arranged in atoms. As we move across the table, electrons are added step by step into different energy levels called orbitals or subshells . Understanding this pattern helps us explain why elements behave the way they do.

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    Study Guide 2.1 Filling of the s Subshell The first two columns on the left side of the periodic table are called the alkali metals (Group 1) and alkaline earth metals (Group 2). As we move down these columns, electrons are added to the s subshell . • Group 1 elements have 1 electron in the s subshell • Group 2 elements have 2 electrons in the s subshell Because these s electrons are loosely held, these elements easily lose them. That is why: • Group 1 elements usually form +1 ions • Group 2 elements usually form +2 ions Helium is an exception. Even though it has two s electrons, it is very stable and does not normally react. Figure 1.Filling of the s subshells. 2 .2 Filling of the 2p Subshell After the s subshell is filled, electrons begin to fill the p subshell . The elements from boron (B) to neon (Ne) show how electrons fill the 2p subshell . • The p subshell can hold up to 6 electrons • Each step across the table adds one more electron to the p subshell These elements gradually become more stable as the p subshell fills.

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    Study Guide Figure 2. Filling of the 2 p subshell. The same filling pattern repeats in the next energy level. Elements from aluminum (Al) to argon (Ar) show the filling of the 3p subshell . • Again, the p subshell fills from 1 to 6 electrons • When the 3p subshell is full, the atom becomes very stable Figure 3. Filling of the 3p Subshell 2.3 Transition Metals and the 3d Subshell The large block of metals in the middle of the periodic table are called the transition metals . They extend from scandium (21) to mercury (80) . In these elements: • Electrons are added to the d subshell • The 3d subshell can hold up to 10 electrons As you move across this block, the number of electrons in the 3d subshell increases gradually. Figure 4 . Filling of the 3d subshell

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    Study Guide 2.4 Anomalous Electron Configurations Some elements do not follow the usual filling pattern. Chromium (Cr) and Copper (Cu) are important examples. • One electron moves from the 4s subshell to the 3d subshell • This happens because half - filled or fully filled d subshells are more stable Because of this behavior: • Transition metals often have multiple oxidation states • For example, vanadium can form +2, +3, +4, or +5 ions This reminds us that electron configurations have exceptions , not just rules. 2.5 Lanthanides and Actinides The two rows shown at the bottom of the periodic table are: • Lanthanides • Actinides They are placed separately to keep the table compact and easy to read. • Lanthanides: elements 57 – 71 • Actinides: elements 89 – 103 These elements involve the filling of f subshells , which can hold 14 electrons . Figure 5. The correct placement of the lanthanides and actinides in the periodic table.

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    Study Guide 2.6 Special Features of Actinides • Actinides have large, unstable nuclei • Many undergo spontaneous radioactive decay • Elements with atomic numbers 93 and higher are synthetically produced 2.7 Why Electron Configuration Matters The outermost electrons of an atom are the most important. They determine: • Chemical behavior • Atomic size • Charge • Ability to form bonds When you understand how electrons fill subshells, the periodic table starts to make sense: • Elements in the same column behave similarly • This is because they have the same type of outer electron configuration Summary • The periodic table shows how electrons fill s, p, d, and f subshells • Electron arrangement explains chemical properties • Patterns exist, but exceptions are important • Knowing electron configuration helps you understand chemistry better 3. Quiz: The Periodic Table 1. Question Another name for the lanthanide series is

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