Periodic Table Trends: Ionization Energy

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Questions and Answers

Explain how the number of protons in the nucleus of an atom affects its first ionisation energy and why?

As the number of protons increases, the positive charge of the nucleus also increases. This results in a stronger attraction for electrons, thus more energy is needed to remove an electron.

Describe the relationship between atomic radius and first ionisation energy. Why does this relationship exist?

As atomic radius increases first ionization energy decreases. This is because valence electrons are further from the nucleus and, thus, easier to remove.

How does electronegativity differ between elements that readily lose electrons versus those that readily gain electrons?

Elements that readily lose electrons have low electronegativity, while those that readily gain electrons have high electronegativity.

Explain how the number of electron shells affects the first ionisation energy of an element.

As the number of electron shells increases, the first ionisation energy decreases, because valence electrons are further from the nucleus and are shielded by inner electrons, which reduces the effective nuclear charge. Signup and view all the answers

What two factors determine the electronegativity of an element?

The number of electrons needed to fill its valence shell and its atomic radius. Signup and view all the answers

Describe how first ionisation energy changes as you move down a group in the Periodic Table.

The first ionisation energy decreases moving down a group. Signup and view all the answers

Does sodium (Na) or chlorine (Cl) have a higher electronegativity? Briefly explain why.

Chlorine (Cl) has a higher electronegativity because it needs only one electron to complete its valence shell, whereas sodium (Na) tends to lose an electron. Signup and view all the answers

Describe how first ionisation energy changes as you move across a period in the Periodic Table.

First ionisation energy increases from left to right across a period. Signup and view all the answers

Explain why electronegativity generally increases across a period in the periodic table.

Electronegativity increases across a period because the nuclear charge increases, leading to a stronger attraction for electrons. The shielding effect remains relatively constant, allowing the effective nuclear charge to increase. Signup and view all the answers

Describe the trend in electronegativity as you move down a group in the periodic table and provide a reason for this trend.

Electronegativity decreases down a group. This is because the number of electron shells increases, increasing the distance between the valence electrons and the nucleus, as well as increasing the shielding effect, reducing the effective nuclear charge. Signup and view all the answers

How does the effective nuclear charge influence the trend in atomic radius across a period?

As the effective nuclear charge increases across a period, the valence electrons are pulled closer to the nucleus, resulting in a decrease in atomic radius. Signup and view all the answers

Explain why metallic character decreases across a period.

Metallic character decreases across a period because elements become less likely to lose electrons due to increasing electronegativity and effective nuclear charge. It becomes more favorable to gain electrons. Signup and view all the answers

Describe the trend in atomic radius as you move down a group in the periodic table, and explain the underlying reason.

Atomic radius increases down a group. This is primarily due to the addition of new electron shells, which places the valence electrons farther from the nucleus. Signup and view all the answers

In terms of electron behavior, what makes an element exhibit strong metallic character?

An element with strong metallic character readily loses electrons to form positive ions (cations). They have low electronegativity and ionization energy. Signup and view all the answers

How does increased shielding affect the electronegativity of an atom?

Increased shielding reduces the effective nuclear charge experienced by the outer electrons, making it harder for the atom to attract additional electrons, and thus decreasing electronegativity. Signup and view all the answers

Explain the relationship between atomic radius and ionization energy. How does the trend in atomic radius influence the trend in ionization energy across a period?

Smaller atomic radius generally leads to higher ionization energy because the valence electrons are closer to the nucleus and experience a stronger attraction. As atomic radius decreases across a period, ionization energy increases. Signup and view all the answers

How does the atomic radius generally change as you move down a group (vertical column) in the periodic table? Briefly explain the reason for this trend.

Atomic radius increases down a group. This is because each subsequent element has an additional electron shell, increasing the distance of the outermost electrons from the nucleus. Signup and view all the answers

Define the term 'first ionisation energy'.

The first ionisation energy is the energy required to remove one electron from an atom in the gaseous phase. Signup and view all the answers

Comparing Sodium, Potassium and Lithium, which element would you expect to have the LOWEST first ionization energy? Explain your reasoning.

Potassium (K) would have the lowest first ionization energy, because it has the largest atomic radius, so its valence electron is furthest from the nucleus. Signup and view all the answers

How does the electronegativity trend change as you move across a period (horizontal row) from left to right on the periodic table? Briefly explain the underlying cause.

Electronegativity generally increases across a period. The increased nuclear charge attracts the valence electrons more strongly. Signup and view all the answers

Consider the elements Oxygen and Fluorine. Which element is more electronegative, and how does its position on the periodic table relate to its electronegativity?

Fluorine is more electronegative. As the element is further to the right in the same period it has more protons and therefore a greater effective nuclear charge. Signup and view all the answers

Explain how the metallic character of elements changes as you move down Group 1 (alkali metals). Relate your answer to ionization energy.

Metallic character increases down Group 1. Ionization energy decreases down the group. Signup and view all the answers

Predict which of the following elements has the most metallic character: Sodium, Aluminum, or Silicon. Explain your reasoning based on periodic trends.

Sodium (Na) has the most metallic character. Metallic character decreases from left to right across the period, so sodium should have the most metallic character of the three. Signup and view all the answers

How does 'shielding' by inner electrons affect the effective nuclear charge experienced by valence electrons, and how does this influence ionization energy?

Inner electrons shield valence electrons from the full positive charge of the nucleus, thus decreasing the effective nuclear charge experienced by the valence electrons. As shielding increases and effective nuclear charge decreases, the valence electrons are bound less tightly, resulting in a lower ionization energy. Signup and view all the answers

Flashcards

First Ionization Energy

The energy required to remove one electron from a neutral atom in its gaseous phase.

Small First Ionization Energy

Elements with a smaller first ionization energy tend to lose electrons to achieve stability.

Large First Ionization Energy

Elements with a larger first ionization energy tend to gain electrons to achieve stability.

Ionization Energy Across a Period

First ionization energy increases from left to right across a period due to increasing nuclear charge.