Periodic trends
60 flashcards covering Periodic trends for the MCAT Chem / Phys / Psych / Soc section.
Periodic trends refer to the predictable patterns in the properties of elements as you move across periods or down groups in the periodic table. For example, atomic radius generally decreases from left to right across a period due to increasing nuclear charge, while it increases down a group because of added electron shells. These trends also include changes in ionization energy, electronegativity, and electron affinity, which help explain why elements behave differently and are crucial for understanding chemical reactivity and bonding.
On the MCAT, periodic trends appear in the Chemical and Physical Foundations section, often through multiple-choice questions that ask you to compare element properties or predict trends. Common traps include confusing how factors like shielding and effective nuclear charge affect these patterns, so watch for questions that test exceptions or misapplications. Focus on mastering the underlying reasons for trends and practicing with periodic table data to apply them quickly. A solid grasp here can boost your score on related passages and standalone items.
Remember to memorize the key trends and their exceptions for quick recall.
Terms (60)
- 01
Periodic trends
Periodic trends are patterns in the properties of elements as they are arranged in the periodic table, such as changes in atomic size, ionization energy, and electronegativity across periods and down groups.
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Atomic radius
Atomic radius is the distance from the nucleus to the outermost electron shell of an atom, and it generally decreases across a period due to increasing nuclear charge and increases down a group as additional electron shells are added.
- 03
Atomic radius across a period
Across a period from left to right, atomic radius decreases because the increasing number of protons pulls the electrons closer to the nucleus, overcoming electron-electron repulsion.
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Atomic radius down a group
Down a group, atomic radius increases as each successive element has an additional electron shell, which is farther from the nucleus despite increased nuclear charge.
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Ionic radius
Ionic radius is the size of an ion, where cations are smaller than their parent atoms due to loss of electrons and reduced electron-electron repulsion, while anions are larger due to gained electrons and increased repulsion.
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Ionic radius of cations
For cations, ionic radius is smaller than the atomic radius because removing electrons reduces the number of electron shells or decreases electron repulsion, allowing the nucleus to pull remaining electrons closer.
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Ionic radius of anions
For anions, ionic radius is larger than the atomic radius because adding electrons increases electron-electron repulsion, expanding the electron cloud.
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Ionization energy
Ionization energy is the energy required to remove an electron from a gaseous atom or ion, and it generally increases across a period and decreases down a group.
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Ionization energy across a period
Across a period, ionization energy increases due to decreasing atomic radius and increasing effective nuclear charge, making it harder to remove an electron.
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Ionization energy down a group
Down a group, ionization energy decreases because the outermost electrons are farther from the nucleus and experience more shielding, making them easier to remove.
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Exceptions in ionization energy
Exceptions occur in ionization energy trends, such as between group 2 and 13 elements, where elements like boron have lower ionization energy than beryllium due to electron configuration stability.
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Electron affinity
Electron affinity is the energy change when an electron is added to a neutral atom in the gaseous state, and it generally becomes more negative across a period and less negative down a group.
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Electron affinity across a period
Across a period, electron affinity typically becomes more negative as atomic size decreases and effective nuclear charge increases, making it easier to attract an additional electron.
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Electron affinity down a group
Down a group, electron affinity becomes less negative because the added electron is farther from the nucleus, reducing the attraction and making the process less favorable.
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Electronegativity
Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond, and it increases across a period and decreases down a group.
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Electronegativity across a period
Across a period, electronegativity increases due to decreasing atomic radius and increasing effective nuclear charge, enhancing the atom's pull on shared electrons.
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Electronegativity down a group
Down a group, electronegativity decreases as atomic radius increases and the outermost electrons are farther from the nucleus, weakening the atom's ability to attract electrons.
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Pauling scale
The Pauling scale is a numerical measure of electronegativity based on bond dissociation energies, assigning values from about 0.7 for cesium to 4.0 for fluorine to compare elements' electron-attracting abilities.
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Effective nuclear charge
Effective nuclear charge is the net positive charge experienced by an outer electron, calculated as the atomic number minus the shielding effect of inner electrons, and it increases across a period.
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Shielding effect
The shielding effect is the reduction in the nuclear attraction for an outer electron due to the repulsion from inner electrons, which remains relatively constant down a group and increases slightly across a period.
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Metallic character
Metallic character refers to the tendency of an element to exhibit metallic properties like conductivity, and it increases down a group and decreases across a period.
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Metallic character across a period
Across a period, metallic character decreases as elements transition from metals to nonmetals, with increasing ionization energy and electronegativity making elements less likely to lose electrons.
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Metallic character down a group
Down a group, metallic character increases because larger atomic radii and lower ionization energies make it easier for atoms to lose electrons and form positive ions.
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Non-metallic character
Non-metallic character is the tendency to gain electrons and form negative ions, and it increases across a period and decreases down a group.
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Reactivity of alkali metals
Alkali metals in group 1 increase in reactivity down the group due to decreasing ionization energy, making it easier for them to lose their single valence electron in reactions.
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Reactivity of halogens
Halogens in group 17 increase in reactivity up the group due to decreasing atomic radius and increasing electronegativity, making it easier for them to gain an electron.
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Melting point trends for metals
Melting points of metals generally increase across a period up to group 8 and then decrease, due to varying strengths of metallic bonds influenced by the number of valence electrons.
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Boiling point trends
Boiling points typically increase down a group for metals due to stronger metallic bonds from larger atoms, but vary across periods based on intermolecular forces.
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Density trends
Density of elements generally increases down a group and across periods up to the transition metals, due to increasing atomic mass and decreasing atomic volume.
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Lanthanide contraction
Lanthanide contraction is the steady decrease in atomic radii of the lanthanide elements due to poor shielding by 4f electrons, which affects the sizes of subsequent elements in the periodic table.
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Periodic law
The periodic law states that the properties of elements are a periodic function of their atomic numbers, explaining the repeating patterns observed in the periodic table.
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Main group elements
Main group elements are those in groups 1, 2, and 13-18, whose properties follow clear periodic trends due to their valence electron configurations.
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Transition metals
Transition metals in groups 3-12 exhibit less pronounced periodic trends compared to main group elements because of d-orbital filling, affecting their sizes and ionization energies.
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Acidic properties of oxides
Oxides of elements become more acidic across a period from left to right, as non-metals form acidic oxides while metals form basic ones.
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Basic properties of oxides
Oxides of metals, especially those lower in a group, are basic and react with acids, while non-metal oxides are acidic.
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Common trap: Atomic vs. ionic radius
A common mistake is confusing atomic radius with ionic radius; remember that cations are always smaller than their atoms, while anions are larger.
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Strategy for predicting trends
To predict periodic trends, always consider the direction across a period or down a group and the underlying factors like effective nuclear charge and electron shielding.
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Valence electrons and trends
The number of valence electrons influences periodic trends, as elements with fewer valence electrons are more metallic and reactive on the left side of the periodic table.
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Electron configuration impact
Electron configuration affects trends by determining stability; for example, half-filled or fully filled subshells can lead to anomalies in ionization energy.
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Group trends in reactivity
Reactivity generally increases down groups for metals and decreases for non-metals, based on ease of electron loss or gain.
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Period trends in electronegativity
Electronegativity peaks at the top right of the periodic table, excluding noble gases, due to small atomic sizes and high effective nuclear charges.
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Noble gases and trends
Noble gases have very low reactivity and electronegativity due to full valence shells, making them outliers in periodic trends.
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Example: Sodium vs. chlorine
Sodium has a larger atomic radius and lower electronegativity than chlorine, making sodium more metallic and reactive as a metal.
In the periodic table, sodium in period 3 group 1 has an atomic radius of 186 pm, while chlorine in period 3 group 17 has 99 pm.
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Ionization energy of noble gases
Noble gases have the highest ionization energies in their periods due to their stable electron configurations.
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Electronegativity of fluorine
Fluorine has the highest electronegativity at 4.0 on the Pauling scale, making it the most electron-attracting element.
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Atomic radius of hydrogen
Hydrogen's atomic radius is smaller than alkali metals but larger than halogens in the same period, due to its single electron.
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Trends in alkali metal reactivity
Alkali metals react more vigorously with water down the group, as seen with lithium being less reactive than potassium.
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Halogen displacement reactions
A more reactive halogen can displace a less reactive one from its compounds, following the trend of decreasing reactivity down the group.
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Effective nuclear charge formula
Effective nuclear charge is approximated as Z minus S, where Z is the atomic number and S is the shielding constant from inner electrons.
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Common trap: Shielding vs. size
Do not confuse shielding effect with atomic size; shielding affects how much the nucleus attracts outer electrons, while size is the actual distance.
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Period 2 trends overview
In period 2, atomic radius decreases from lithium to neon, while ionization energy and electronegativity increase, with exceptions like beryllium to boron.
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Group 17 electron affinity
Group 17 elements have highly negative electron affinities, increasing up the group, as they need one electron to complete their octet.
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Melting point of group 1 metals
Melting points of group 1 metals decrease down the group due to weaker metallic bonds from larger atomic sizes.
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Density of transition metals
Transition metals have high densities due to their compact atomic structures from d-orbital electrons.
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Example: Ionization energy of Be and B
Beryllium has a higher ionization energy than boron because beryllium's electron configuration is more stable with a full s-subshell.
Be (1s2 2s2) requires more energy to remove an electron than B (1s2 2s2 2p1).
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Electronegativity in bonding
Differences in electronegativity determine bond polarity, with greater differences leading to more ionic bonds.
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Trends in acid strength
Binary acids increase in strength down a group due to weaker H-X bonds from larger atomic sizes.
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Atomic radius and periodic position
Elements in the same group have similar chemical properties due to the same number of valence electrons, but their sizes increase down the group.
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Ionization energy anomalies
Anomalies in ionization energy, like nitrogen having higher energy than oxygen, occur due to half-filled subshell stability.
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Strategy for MCAT questions on trends
For MCAT questions, visualize the periodic table and recall that most trends follow effective nuclear charge and distance from the nucleus.