CHEM 1120 Unit 3: Periodicity and Periodic Trends

Comprehensive practice flashcards covering the periodic properties of elements, effective nuclear charge, atomic and ionic radii, ionization energies, electron affinities, and chemical group trends.

Between $K^+$ and $Br^-$, which ion has the larger radius, and why?

$Br^-$ has a larger radius than $K^+$. $Br^-$ is isoelectronic with $Kr$ and has $4$ full shells, whereas $K^+$ is isoelectronic with $Ar$ and has $3$ full shells; having more electron shells results in a larger size.

Why is the first ionization energy of nitrogen ($N$) higher than that of carbon ($C$)?

$N$ has a greater effective nuclear charge ($Z_{eff}$) than $C$, which exerts a stronger nuclear attraction on the valence electrons, resulting in a higher first ionization energy.

Why is the first ionization energy of nitrogen ($N$) higher than that of oxygen ($O$)?

$N$ has a half-filled $2p$ subshell ($2p^3$) with one electron in each orbital, making it extra stable. In contrast, $O$ has $4$ electrons in the $2p$ subshell ($2p^4$) with one doubly occupied orbital; the repulsion between the two paired electrons makes them less stable, so removing one electron relieves repulsion and yields a stable half-filled subshell.

Which is greater: the first ionization energy of neon ($Ne$) or the second ionization energy of sodium ($Na$), and why?

The second ionization energy of $Na$ is greater. Both processes involve removing an electron from the $2p$ subshell of a full second shell, but $Na$ has more protons ($11$ protons vs. $10$ protons in $Ne$), creating a stronger nuclear attraction toward the electrons.

What organizational concept for the elements was proposed by British chemist John Newlands in 1865?

The 'Law of Octaves', which proposed that when elements are arranged by increasing atomic weight, similar chemical properties recur every eighth element, analogous to musical octaves.

Which historical representation of the periodic system is depicted in the image below, and who proposed it?

The Spiral Periodic Table, proposed in 1866 by Gustavus Detlef Heinrichs.

What periodic relationship did Dmitri Mendeleev formulate in 1869?

When the chemical elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically.

What placeholder name did Dmitri Mendeleev give in 1871 to the undiscovered element that was later identified as germanium in 1886?

Eka-silicon

What were Mendeleev's predicted values versus the observed values for the atomic mass and density of eka-silicon / germanium?

Predicted for eka-silicon: atomic mass $72$ and density $5.5\,g\,cm^{-3}$. Observed for germanium: atomic mass $72.6$ and density $5.47\,g\,cm^{-3}$.

What were the predicted formula and density for eka-silicon oxide, and what was observed for germanium oxide?

Predicted for eka-silicon oxide: $EsO_2$ with a density of $4.7\,g\,cm^{-3}$. Observed for germanium oxide: $GeO_2$ with a density of $4.703\,g\,cm^{-3}$.

What were the predicted and observed boiling points and densities for the chloride of eka-silicon / germanium?

Predicted for $EsCl_4$: boiling point below $100\,^\circ\text{C}$ and density of $1.9\,g\,cm^{-3}$. Observed for $GeCl_4$: boiling point of $86\,^\circ\text{C}$ and density of $1.887\,g\,cm^{-3}$.

How does effective nuclear charge ($Z_{eff}$) change across a period from left to right when the principal quantum number $n$ is fixed?

$Z_{eff}$ increases across a period from left to right while $n$ stays the same.

How does the principal quantum number $n$ change moving down a group, and what is its consequence?

$n$ increases down a group, which increases the total number of electron shells.

State Coulomb's Law equation as presented in the lecture notes.

$$F = k \frac{q_1 q_2}{r^2}$$

According to Coulomb's Law, what happens to atomic orbitals as $Z_{eff}$ increases while $n$ is fixed?

Atomic orbitals shrink (become smaller) and their stability increases.

According to Coulomb's Law, what happens to atomic orbitals as the principal quantum number $n$ increases?

Atomic orbitals expand (become larger) and their stability decreases.

How is the nonbonding atomic radius (van der Waals radius) of an atom defined?

One-half the distance between adjacent nuclei in the atomic solid.

How is the bonding atomic radius (covalent radius) of an atom defined?

One-half the distance between two bonded atoms of the same element.

If the distance between two bonded bromine nuclei in $Br_2$ is $228\,pm$, what is the covalent radius of bromine?

$$\frac{228\,pm}{2} = 114\,pm$$

What is the general periodic trend for atomic size across a period and down a group?

Atomic size decreases from left to right across a period and increases from top to bottom down a group.

Why do the atomic sizes of transition elements stay roughly constant across each row?

Electrons are added to inner $d$ subshells, which effectively shield the outer valence electrons from the increasing nuclear charge, keeping effective nuclear charge and size relatively constant.

How does the size of a cation compare to its parent atom, and why?

Cations are smaller than their parent atoms because electron loss reduces electron-electron repulsion and often vacates the outermost shell, drawing remaining electrons closer to the nucleus.

How does the size of an anion compare to its parent atom, and why?

Anions are larger than their parent atoms because adding electrons increases electron-electron repulsion within the same shell without an increase in nuclear charge.

How does ionic radius change as positive charge increases for different cations of the same element, such as $Fe^{2+}$ versus $Fe^{3+}$?

Ionic radii decrease as the positive charge increases; therefore, $Fe^{2+} > Fe^{3+}$.

How do ionic radii trend down a group in the periodic table?

Ionic radii increase down a group due to the addition of principal energy levels (shells).

Rank the ions $Ca^{2+}$, $Sr^{2+}$, and $Mg^{2+}$ in order of decreasing size, and justify the ranking.

$Sr^{2+} > Ca^{2+} > Mg^{2+}$, because all three are Group 2 cations and ionic radius increases down the group as the number of electron shells increases.

Rank the isoelectronic ions $K^+$, $S^{2-}$, $Cl^-$, and $Ca^{2+}$ in order of decreasing size.

$S^{2-} > Cl^- > K^+ > Ca^{2+}$, because for an isoelectronic series ($18$ electrons), radius decreases as nuclear charge ($Z$) increases from $16$ protons ($S$) to $20$ protons ($Ca$).

What are the ionic radii in picometers for the isoelectronic ions $O^{2-}$, $F^-$, $Na^+$, and $Mg^{2+}$?

$O^{2-} = 140\,pm$, $F^- = 133\,pm$, $Na^+ = 102\,pm$, and $Mg^{2+} = 72\,pm$.

Between $Au^+$ and $Au^{3+}$, which ion is larger, and why?

$Au^+ > Au^{3+}$, because ionic radius decreases as the positive charge increases for cations of a given element.

What is the definition of ionization energy ($IE$)?

The minimum amount of energy needed to remove an electron from a gaseous ground-state atom.

Write the thermochemical equation and value for the first ionization energy ($IE_1$) of sodium ($Na$).

$Na(g) \rightarrow Na^+(g) + e^-$ with $IE_1 = 496\,kJ\,mol^{-1}$.

Write the thermochemical equation and value for the second ionization energy ($IE_2$) of sodium ($Na$).

$Na^+(g) \rightarrow Na^{2+}(g) + e^-$ with $IE_2 = 4560\,kJ\,mol^{-1}$.

What is the general periodic trend for first ionization energy ($IE_1$) across a period and down a group?

$IE_1$ increases from left to right across each period and decreases from top to bottom down each group.

Why is the first ionization energy of aluminum ($Al$, $577\,kJ\,mol^{-1}$) lower than that of magnesium ($Mg$, $738\,kJ\,mol^{-1}$)?

For $Al$, the removed electron comes from a higher-energy $3p$ orbital ($[Ne]3s^2 3p^1$) that is shielded by the $3s$ electrons, making it easier to remove than a $3s$ electron from $Mg$ ($[Ne]3s^2$).

What are the ground-state electron configuration and first ionization energy ($IE_1$) of silicon ($Si$)?

Electron configuration: $[Ne]3s^2 3p^2$; $IE_1 = 786\,kJ\,mol^{-1}$.

Why is the energy required to remove core electrons substantially higher than that required to remove valence electrons?

Core electrons are closer to the nucleus (lower principal quantum number $n$) and experience significantly less electron shielding, resulting in a much stronger electrostatic attraction to the nucleus.

What are the first three ionization energies ($IE_1$, $IE_2$, $IE_3$) for magnesium ($Mg$)?

$IE_1 = 738\,kJ\,mol^{-1}$, $IE_2 = 1450\,kJ\,mol^{-1}$, and $IE_3 = 7730\,kJ\,mol^{-1}$.

Why does aluminum ($Al$) show a dramatic increase in ionization energy between $IE_3$ ($2750\,kJ\,mol^{-1}$) and $IE_4$ ($11600\,kJ\,mol^{-1}$)?

The first three electrons removed are valence electrons from the $n=3$ shell; the fourth electron must be removed from the stable noble gas core ($[Ne]$, $n=2$ shell).

A third-period element exhibits successive ionization energies of $736$, $1450$, $7740$, $10200$, and $13600\,kJ\,mol^{-1}$. Identify the element.

Magnesium ($Mg$), because the primary jump occurs after the second ionization energy ($IE_2 = 1450\,kJ\,mol^{-1}$ to $IE_3 = 7740\,kJ\,mol^{-1}$), indicating two valence electrons.

How is electron affinity ($EA$) defined?

The energy change that occurs when an electron is added to a gaseous atom.

Write the chemical equation and first electron affinity ($EA_1$) value for chlorine.

$Cl(g) + e^- \rightarrow Cl^-(g)$ with $EA_1 = -349\,kJ\,mol^{-1}$.

Why do certain elements in Group 2 ($ns^2$), Group 15 ($ns^2 np^3$), and Group 18 ($ns^2 np^6$) have positive electron affinity values?

The added electron must occupy a new subshell or pair up in an already half-filled orbital, creating instability that makes the resulting anion less stable than the neutral atom plus a free electron.

What is almost always the sign of second electron affinity ($EA_2$), and why?

Positive, because adding an electron to an existing anion encounters strong electrostatic repulsion.

What is the general periodic trend for first electron affinity ($EA_1$) across a period?

$EA_1$ generally increases (becomes more negative/exothermic) from left to right across a period.

What are the periodic trends in metallic and nonmetallic character?

Metallic character decreases from left to right across a period and increases down a group; nonmetallic character increases from left to right across a period and decreases down a group.

Why are Group 18 noble gases exceptionally unreactive?

They have completely filled valence subshells, resulting in very high ionization energies and positive electron affinities.

How do first ionization energy ($IE_1$) and melting point compare between lithium ($Li$) and cesium ($Cs$)?

For $Li$: $IE_1 = 520\,kJ\,mol^{-1}$ and melting point is $181\,^\circ\text{C}$. For $Cs$: $IE_1 = 376\,kJ\,mol^{-1}$ and melting point is $29\,^\circ\text{C}$.

In the image below showing sealed ampoules of halogens, identify the halogen present in each ampoule from left to right.

From left to right: iodine ($I_2$, dark solid with purple vapor), chlorine ($Cl_2$, pale greenish-yellow gas), and bromine ($Br_2$, reddish-brown liquid and vapor).

According to Table 8.3, what are the atomic radius, electron affinity ($EA$), and boiling point of fluorine ($F$)?

Atomic radius: $72\,pm$; $EA$: $328\,kJ\,mol^{-1}$; boiling point: $-188\,^\circ\text{C}$.

What cryogenic application and minimum temperature limit are cited in the lecture for liquid helium?

Liquid helium is used as a cryogenic coolant to cool substances to temperatures as low as $1.2\,K$.