This comprehensive guide provides answers and detailed explanations for common periodic trends review problems. Understanding periodic trends is fundamental to mastering chemistry, so let's dive into the key concepts and solidify your understanding.
Understanding Periodic Trends
Before we tackle specific problems, let's briefly review the major periodic trends:
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Atomic Radius: This refers to the size of an atom. Atomic radius generally increases down a group (due to added electron shells) and decreases across a period (due to increasing nuclear charge pulling electrons closer).
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Ionization Energy: This is the energy required to remove an electron from an atom. Ionization energy generally increases across a period (due to increasing nuclear charge) and decreases down a group (due to increasing distance from the nucleus and shielding).
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Electron Affinity: This is the energy change when an electron is added to a neutral atom. Electron affinity generally increases across a period and decreases down a group, though the trend is less consistent than ionization energy.
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Electronegativity: This measures an atom's ability to attract electrons in a chemical bond. Electronegativity generally increases across a period (due to increasing nuclear charge) and decreases down a group (due to increasing atomic size and shielding).
Example Problems and Solutions
Let's work through some typical periodic trend problems. Remember to consult a periodic table!
Problem 1: Which element has a larger atomic radius: Sodium (Na) or Chlorine (Cl)?
Answer: Sodium (Na) has a larger atomic radius than Chlorine (Cl).
Explanation: Sodium and Chlorine are in the same period (Period 3). Atomic radius decreases across a period. Therefore, the element further to the left (Sodium) will have a larger atomic radius.
Problem 2: Which element has a higher ionization energy: Oxygen (O) or Sulfur (S)?
Answer: Oxygen (O) has a higher ionization energy than Sulfur (S).
Explanation: Oxygen and Sulfur are in the same group (Group 16). Ionization energy decreases down a group. The element higher up in the group (Oxygen) will have a higher ionization energy because its valence electrons are closer to the nucleus and experience a stronger attraction.
Problem 3: Arrange the following elements in order of increasing electronegativity: Lithium (Li), Fluorine (F), and Nitrogen (N).
Answer: Li < N < F
Explanation: Electronegativity increases across a period and decreases down a group. Fluorine is furthest to the right and highest up, making it the most electronegative. Lithium is furthest to the left and lowest down, making it the least electronegative. Nitrogen falls between the two.
Problem 4: Explain why the atomic radius of noble gases is generally larger than the atomic radius of halogens in the same period.
Answer: Noble gases have a completely filled valence shell, leading to slightly greater electron-electron repulsion which increases the atomic radius compared to halogens in the same period which have one less electron in their valence shell.
Problem 5: Why does ionization energy generally increase across a period?
Answer: As you move across a period, the number of protons in the nucleus increases, resulting in a stronger positive charge. This stronger nuclear charge attracts the electrons more tightly, requiring more energy to remove an electron (higher ionization energy).
Further Practice
To further solidify your understanding, try working through additional problems from your textbook or online resources. Focus on visualizing the periodic table and understanding the underlying reasons for each trend. Remember to always consult a periodic table for element locations and properties. Consistent practice is key to mastering periodic trends!
