Exercises: Configuration, periodic table and properties

The periodic table is like a bookcase: each row is a period and each column is a group, as shown in the figure.

Explain what the electron configurations of elements in the same group have in common and what information the period number gives about the configuration of an element.

The aluminum in drink cans has the electron configuration $1s^2\ 2s^2\ 2p^6\ 3s^2\ 3p^1$.

1. a) Which period of the periodic table is it in? (Hint: look at the highest energy level containing electrons.)
2. b) How many valence electrons does it have?
3. c) Aluminum is in group $13$: what relationship do you observe between this group number and the valence electrons?

Lithium ($Z = 3$) and sodium ($Z = 11$) are alkali metals so reactive that in the lab they are stored submerged in oil so that they do not touch the air.

1. a) Write the electron configuration of each.
2. b) Point out what their valence shells have in common and which group of the periodic table they are in.

The fluorine ($Z = 9$) in some toothpastes and the chlorine ($Z = 17$) in swimming pools belong to group $17$, the halogens.

1. a) Write the electron configurations of the two elements.
2. b) Explain, based on these configurations, why they are in the same group.

The helium ($Z = 2$) that cools the magnets in hospital MRI machines and the argon ($Z = 18$) that fills incandescent light bulbs belong to group $18$, the noble gases.

1. a) Write the electron configurations of the two gases.
2. b) What is special about their outermost energy levels?
3. c) Relate this feature to the fact that noble gases hardly form compounds.

An element is in period $2$ and has $4$ valence electrons.

1. a) Write its electron configuration.
2. b) Which element is it?
3. c) Which group of the periodic table is it located in?

In the periodic table, metals are located on the left and nonmetals on the right, as shown in the figure.

For each of these elements, state whether it is a metal or a nonmetal and whether it is easier for it to lose or gain electrons:

1. a) sodium, from group $1$;
2. b) chlorine, from group $17$.

The periodic table orders the elements in order of increasing atomic number $Z$. Explain why this order matches the order in which orbitals are filled, taking into account that each element has one more proton (and, if neutral, one more electron) than the previous one.

Match heads contain sulfur, whose electron configuration is $1s^2\ 2s^2\ 2p^6\ 3s^2\ 3p^4$.

1. a) Deduce which period and group of the periodic table sulfur is in.
2. b) How many electrons does it need to complete its outermost energy level?
3. c) Predict which ion it will form and state its charge.

An element is located in group $2$ and period $3$ of the periodic table.

1. a) Write its electron configuration.
2. b) Which element is it? (Hint: it burns with a very intense white light and is used in flares and fireworks.)
3. c) Predict the ion it forms easily and write its symbol with the corresponding charge.

Aluminum (group $13$) and chlorine (group $17$) form the compound $\text{AlCl}_3$, present in some antiperspirants.

1. a) Predict the ion formed by each element, reasoning from their valence electrons.
2. b) Using the charges of these ions, explain why the formula is $\text{AlCl}_3$ and not $\text{AlCl}$.

Sodium and chlorine are in period $3$, but the sodium atom is significantly larger than the chlorine atom, as shown in the figure.

1. a) Write the electron configurations of the two elements and check that their valence electrons occupy the same energy level.
2. b) Explain why the chlorine atom is smaller, comparing the number of protons attracting that electron shell in each case.

Going down the group of alkali metals (lithium, sodium, potassium), atoms get progressively larger, as shown in the figure.

1. a) Write the electron configurations of the three elements ($Z$: Li = 3, Na = 11, K = 19).
2. b) Explain, looking at the number of occupied energy levels, why the atomic radius increases going down the group.

In a classroom experiment, small pieces of lithium, sodium, and potassium are dropped into water: lithium bubbles slowly, sodium moves quickly across the surface, and potassium burns with a violet flame.

1. a) Place the three metals in order from least to most reactive.
2. b) Explain this order taking into account how far the valence electron is from the nucleus in each atom and how easy it is to remove it.

Place the fluorine ($Z = 9$), chlorine ($Z = 17$) and bromine ($Z = 35$) atoms in order of increasing atomic radius. Justify your answer by reasoning about how many occupied electron shells each atom has (fluorine fills up to shell $2$, chlorine up to $3$ and bromine up to $4$).

Two unknown elements have atomic numbers $Z = 3$ and $Z = 19$.

1. a) Write their electron configurations.
2. b) Deduce whether they are in the same group and state which group it is.
3. c) Predict one chemical property they share.

An element X has the electron configuration $1s^2\ 2s^2\ 2p^6\ 3s^2\ 3p^6\ 4s^1$.

1. a) Identify it using the periodic table and deduce its period and group.
2. b) Is it a metal or a nonmetal? Which ion will it easily form?
3. c) Predict the formula of the compound it forms with oxygen, which gives $\text{O}^{2-}$ ions, and justify your answer.

The radius of the sodium atom is about $0.190\ \text{nm}$, whereas that of its $\text{Na}^+$ ion is only $0.098\ \text{nm}$. In contrast, the chlorine atom measures $0.099\ \text{nm}$ and its $\text{Cl}^-$ ion measures $0.181\ \text{nm}$, as shown in the diagram.

1. a) State how the size of each element changes when it becomes an ion.
2. b) Explain both changes by comparing the number of protons and electrons before and after the ions are formed (you may use electron configurations to help you).

The $\text{N}^{3-}$, $\text{O}^{2-}$, $\text{F}^{-}$, $\text{Na}^{+}$ and $\text{Mg}^{2+}$ ions all have the same electron configuration, as shown in the diagram.

1. a) Write this configuration.
2. b) However, their radii are different: $0.065$; $0.098$; $0.133$; $0.140$ and $0.171\ \text{nm}$ (not in order). Assign each radius to its corresponding ion and explain the criterion used, bearing in mind that they all have $10$ electrons but a different number of protons in the nucleus.

In the alkali metals group, reactivity increases going down the group: potassium reacts much more violently with water than lithium does. In the halogens group, the exact opposite occurs: fluorine is the most reactive, and going down the group, reactivity decreases, as shown in the diagram.

Explain both trends using a single concept: what each family needs to do with electrons (lose them or gain them) to achieve a noble gas configuration, and how far their valence shell is from the nucleus.

Wall plaster contains an element X whose $\text{X}^{2+}$ ion is isoelectronic with argon ($Z = 18$).

1. a) Calculate the atomic number of X and identify it using the periodic table.
2. b) Deduce its group and period.
3. c) Predict whether it will be a reactive metal and explain what would happen to a small piece of X if it were dropped into water (compare it with sodium).

Neon never forms chemical compounds, but sodium and chlorine readily combine to form common salt, $\text{NaCl}$.

1. a) Write the electron configurations of the three elements ($Z$: Ne = 10, Na = 11, Cl = 17).
2. b) Explain why neon "does not need" to react.
3. c) Explain why sodium and chlorine "need each other" to obtain a noble gas configuration.

Lithium, sodium and potassium are very similar to each other, and the same is true for fluorine, chlorine and bromine.

1. a) Write the electron configurations of Li ($Z = 3$), Na ($Z = 11$) and K ($Z = 19$), and highlight what is repeated in their valence shells.
2. b) Based on this repetition, explain why the chemical properties of elements recur periodically across the periodic table.

Five elements from periods $2$, $3$ and $4$ have approximate atomic radii of $0.064$; $0.099$; $0.155$; $0.190$ and $0.235\ \text{nm}$. They are fluorine, chlorine, lithium, sodium and potassium, but it is not known which is which.

1. a) Assign each radius to its element, reasoning using the two periodic trends (how the radius changes across a period and down a group).
2. b) Arrange the five elements in order of increasing atomic number and explain how the two trends combine across the periodic table.