Periodic Trends (College Board AP® Chemistry): Flashcards
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Define periodic table.
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Define periodic table.
Periodic table is a tabular arrangement of chemical elements organized by increasing atomic number, from 1 to 118, into horizontal rows (periods) and vertical columns (groups).
True or False?
Elements in the same group of the periodic table share similar chemical properties.
True.
Groups are vertical columns whose members have the same number of valence electrons, which determines their chemical behavior.
What do the periods (horizontal rows) of the periodic table represent?
Periods represent the principal energy shells of an element. Elements in the same period have their valence electrons in the same outermost shell.
The periodic table is divided into four block types: ..........-block, p-block, d-block, and f-block, named after the subshells occupied by valence electrons.
The periodic table is divided into four block types: s-block, p-block, d-block, and f-block, named after the subshells occupied by valence electrons.
True or False?
d-block elements have at least one d-electron and one s-electron, but no p or f electrons in their outer shell.
True.
d-block elements (transition metals) fill d-subshells. Their outer shell contains d and s electrons, but no p or f electrons.
Germanium (Ge) is in Period 4 and Group 14 of the p-block. What is its full electron configuration?
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p2.
Period 4 places valence electrons in the fourth shell; Group 14 gives four valence electrons; the p-block position places the last two electrons in 4p.
p-block elements have at least one ..........-electron in their valence shell.
p-block elements have at least one p-electron in their valence shell.
Define ionization energy.
Ionization energy is the energy required to remove one mole of electrons in the gaseous state from one mole of neutral gaseous atoms, forming one mole of positively charged gaseous ions.
True or False?
Atomic radius and ionization energy follow the same periodic trend across a period.
False.
Atomic radius and ionization energy follow opposite trends. As you move left to right across a period, ionization energy increases while atomic radius decreases, both driven by increasing effective nuclear charge.
Why does ionization energy decrease as you move down a group in the periodic table?
Moving down a group adds inner electron shells, which increases the shielding effect and the distance between the nucleus and valence electrons. This weakens the Coulombic attraction experienced by valence electrons, making it easier to remove an outer electron.
As you move across a period from left to right, atomic radius generally .......... because the increasing number of protons raises the effective nuclear charge and strengthens Coulombic attraction.
As you move across a period from left to right, atomic radius generally decreases because the increasing number of protons raises the effective nuclear charge and strengthens Coulombic attraction.
Define electronegativity.
Electronegativity is a measure of the ability of an atom to attract a pair of electrons when it forms a covalent bond.
True or False?
Across a period, the effective nuclear charge experienced by valence electrons increases from left to right.
True.
As you move left to right across a period, protons are added to the nucleus while electrons fill the same shell. Shielding barely changes, so valence electrons experience a progressively higher effective nuclear charge.
Why does the ionic radius increase as the negative charge on an ion increases?
When electrons are added to form a more negative ion, the nuclear charge stays the same while electron–electron repulsion increases among the valence electrons. This repulsion pushes the electron cloud outward, resulting in a larger ionic radius.
Across Period 2 from lithium to neon, first ionization energy generally .......... because the nuclear charge increases while the number of inner shielding electrons stays approximately constant.
Across Period 2 from lithium to neon, first ionization energy generally increases because the nuclear charge increases while the number of inner shielding electrons stays approximately constant.
How can the position of an element in the periodic table be used to predict its atomic radius?
An element's period indicates the number of inner shells (shielding), and its group position indicates effective nuclear charge relative to others in the same period. Elements further right in a period have a smaller atomic radius (higher effective nuclear charge); elements further down a group have a larger atomic radius (more shielding and greater nucleus–valence electron distance).
True or False?
Strontium (Sr) has a smaller atomic radius than bromine (Br) because strontium is in Period 5 and has greater shielding.
False.
Strontium is in Period 5, which means more inner shells and greater shielding, giving it a weaker Coulombic attraction to its valence electrons. This makes strontium larger than bromine, not smaller.
Of potassium (K), calcium (Ca), bromine (Br), and strontium (Sr), .......... has the smallest atomic radius because it is furthest right in its period and has the greatest effective nuclear charge relative to the others in the same period.
Of potassium (K), calcium (Ca), bromine (Br), and strontium (Sr), bromine has the smallest atomic radius because it is furthest right in its period and has the greatest effective nuclear charge relative to the others in the same period.
Chlorine (Cl) is directly below fluorine (F) in Group 17. Predict whether chlorine's ionization energy is greater or smaller than fluorine's, and explain why.
Chlorine has a smaller ionization energy than fluorine. Chlorine has more inner electron shells, increasing the shielding effect and the distance between the nucleus and valence electrons. This weakens the Coulombic attraction, so less energy is needed to remove an outer electron compared to fluorine.
True or False?
The periodic table can be used to predict chemical properties of elements even when experimental data is unavailable.
True.
Periodic trends in ionization energy, atomic radius, electronegativity and electron affinity are predictable from an element's position in the table, allowing properties to be estimated without direct measurement.
Define valence electrons.
Valence electrons are the electrons present in the outermost shell of an atom. They dictate an atom's reactivity and ability to form bonds.
True or False?
Elements in the same group of the periodic table have the same number of valence electrons.
True.
Group membership is defined by the number of valence electrons. This shared count causes elements in the same group to display similar chemical behavior.
What is the difference between a cation and an anion?
A cation is a positively charged ion formed when an atom loses electrons; an anion is a negatively charged ion formed when an atom gains electrons. Metals typically form cations; nonmetals typically form anions.
Atoms form ions to achieve the electron configuration of their closest .........., which has a stable, full outer shell.
Atoms form ions to achieve the electron configuration of their closest noble gas, which has a stable, full outer shell.
True or False?
An ionic bond forms between two nonmetals.
False.
An ionic bond forms between a metal and a nonmetal. Two nonmetals share electrons and form a covalent bond instead.
Using the periodic table, predict the typical ion charge for elements in Group 6, and explain why.
Group 6 elements typically form 2− ions. They have six valence electrons and gain two more to complete their outer shell and achieve the electron configuration of the nearest noble gas. The charge can be calculated as: group number − 8 = 6 − 8 = −2.
Elements in Group 2 tend to lose .......... electrons from their valence shell, forming ions with a 2+ charge.
Elements in Group 2 tend to lose two electrons from their valence shell, forming ions with a 2+ charge.
Why do Group 4 elements typically not form ionic compounds?
Group 4 elements would need to lose or gain four electrons to reach a noble gas configuration, making ionic bonding energetically unfavorable. These elements bond covalently instead.
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