Exercises: Periodicity of properties

All elements in column 1 of the periodic table, the alkali metals, end their electron configuration in $s^1$.

1. a) What do the electron configurations of all elements in the same group have in common?
2. b) What consequence does this similarity have on their chemical properties? Give a specific example from Group 1.

In the laboratory, sodium is stored immersed in mineral oil because it reacts violently with water and even with moisture in the air.

1. a) Describe what is observed when a small piece of sodium is dropped into a trough of water.
2. b) What gas is given off in the reaction? Is the resulting solution acidic or alkaline?

The helium in balloons and the neon in illuminated signs are noble gases.

1. a) What is special about the electron configurations of noble gases?
2. b) Why do they hardly form any chemical compounds? What advantage does helium have over flammable hydrogen for filling balloons?

The chlorine used to disinfect swimming pools and the iodine in pharmacies are halogens (Group 17).

1. a) How many valence electrons do halogens have? What ion do they form when they gain an electron?
2. b) Explain, through the exchange of electrons between both atoms, why sodium and chlorine react to form sodium chloride, $\text{NaCl}$.

The third period of the periodic table, from left to right, is: Na, Mg, Al, Si, P, S, Cl and Ar.

1. a) State which of these elements are metals, which are non-metals and which one is a metalloid.
2. b) How does metallic character change when moving from left to right across a period?

A copper wire conducts electricity; a piece of sulfur does not.

1. a) State three typical properties of metals that non-metals do not share.
2. b) In which regions of the periodic table (left, center, right) are metals and non-metals located?

From the group to which each element belongs, state which ion it forms and justify your answer using the number of electrons it gains or loses to complete its outer shell:

1. a) magnesium (Group 2)
2. b) nitrogen (Group 15)
3. c) potassium (Group 1)

Sodium forms $\text{Na}^{+}$ ions and oxygen forms $\text{O}^{2-}$ ions.

1. a) How many $\text{Na}^{+}$ ions are needed to balance the charge of one $\text{O}^{2-}$ ion?
2. b) Write the formula of sodium oxide. What will be the formula of potassium oxide, which is in the same group as sodium?

When small pieces of lithium, sodium and potassium are dropped into three troughs of water, the reaction becomes progressively more violent. The diagram shows the shell structure of the three atoms:

1. a) Order Li, Na, K and Rb from lowest to highest reactivity with water and balance the equation for the reaction of sodium: $\text{Na} + \text{H}_2\text{O} \rightarrow \text{NaOH} + \text{H}_2$.
2. b) Explain the trend: the further down the alkali metal is in the group, is it easier or harder to remove its valence electron? Relate this to its reactivity.

You can do it on paper and send a photo with your answer.

The reactivity of the halogens varies in the opposite way to that of the alkali metals: fluorine is the most reactive and iodine the least.

1. a) When chlorine is bubbled into a potassium bromide solution, the solution turns brown because bromine is produced: $\text{Cl}_2 + \text{KBr} \rightarrow \text{Br}_2 + \text{KCl}$. Balance the equation and explain what it means that chlorine "displaces" bromine from its compounds.
2. b) Predict whether iodine would displace chlorine from a solution of $\text{NaCl}$. Justify your answer using the group trend.

Some measured atomic radii are: Li: $152\ \text{pm}$, Na: $186\ \text{pm}$ and K: $227\ \text{pm}$ in Group 1; and Na: $186\ \text{pm}$, Mg: $160\ \text{pm}$ and Al: $143\ \text{pm}$ in Period 3. The diagram shows the six atoms to scale:

1. a) Describe, using these data, the trend in atomic radius going down a group and moving across a period.
2. b) Explain it: what is added going down a group that makes the atom larger, and what increases moving across a period that attracts the electron cloud more strongly?

You can do it on paper and send a photo with your answer.

The first ionisation energy (energy required to remove the outermost electron, in $\text{kJ/mol}$) is: Na: $496$, Mg: $738$, Al: $578$ and Cl: $1251$ in Period 3; and Li: $520$, Na: $496$ and K: $419$ in Group 1. The graph plots all these values:

1. a) State the two trends in ionisation energy: moving across a period and going down a group.
2. b) Relate the low ionisation energy of alkali metals to their ease of forming $\text{X}^{+}$ ions and to their high reactivity.

Predict the formula of these ionic compounds from the usual charges of the ions, justifying the exchange of electrons:

1. a) aluminum ($\text{Al}^{3+}$) with sulphur ($\text{S}^{2-}$)
2. b) calcium ($\text{Ca}^{2+}$) with nitrogen ($\text{N}^{3-}$)

An element is a colourless, odourless and completely inert gas; industry uses it as a protective atmosphere in arc welding and to fill the space between the glass panes of double-glazed windows.

1. a) Which group of the periodic table does it belong to?
2. b) Its atom has $18$ electrons: which element is it? Write its electron configuration and explain, based on it, why it is inert.

Electronegativity measures the tendency of an atom to attract the electrons of a bond towards itself. Some values: F: $4.0$, O: $3.5$, Cl: $3.2$, S: $2.6$, Na: $0.9$ and Cs: $0.7$.

1. a) Use these data to verify that electronegativity increases across a period and decreases down a group.
2. b) Which element in the periodic table is the most electronegative, and in which corner is it located? Where, on the other hand, are the least electronegative elements?

Order the elements Na, Mg, S and Si from most to least metallic in character.

1. a) Justify the order based on their positions in the third period.
2. b) Explain what macroscopic property, such as electrical conductivity, would change when going through the series in that order.

The oxides of the Period 3 elements are: $\text{Na}_2\text{O}$, $\text{MgO}$, $\text{Al}_2\text{O}_3$, $\text{SiO}_2$, $\text{SO}_3$ and $\text{Cl}_2\text{O}_7$. When they react with water, the first two give alkaline solutions and the last two give acidic solutions; $\text{Al}_2\text{O}_3$ and $\text{SiO}_2$ are practically insoluble, although $\text{SiO}_2$ has acidic character (it reacts with bases).

1. a) Relate the acidic or basic character of the oxide to the position of the element in the period and to its metallic character.
2. b) Aluminum oxide is amphoteric: it reacts as a base with acids and as an acid with bases. Does this ambiguity fit with its intermediate position in the period? Justify your answer.
3. c) Predict the character (acidic, basic or amphoteric) of potassium oxide, $\text{K}_2\text{O}$, and selenium dioxide, $\text{SeO}_2$, from the positions of K and Se in the periodic table.

Measured radii: Na: $186\ \text{pm}$ and $\text{Na}^{+}$: $95\ \text{pm}$; Cl: $99\ \text{pm}$ and $\text{Cl}^{-}$: $181\ \text{pm}$. The diagram shows atoms and ions to scale:

1. a) Which ion is smaller than its neutral atom and which is larger? Numerically compare each pair.
2. b) Explain both changes by counting electrons and shells: what does sodium lose upon ionisation and why does chlorine increase so much in size when gaining an electron (think about electron-electron repulsion)?
3. c) The $\text{Na}^{+}$ ion and the $\text{F}^{-}$ ion are isoelectronic (they have $10$ electrons): predict which of the two will be larger and justify your answer by comparing their nuclear charges ($Z = 11$ for Na and $Z = 9$ for F).

You can do it on paper and send a photo with your answer.

The first three successive ionisation energies of a Period 3 element are $E_{i,1} = 738\ \text{kJ/mol}$, $E_{i,2} = 1451\ \text{kJ/mol}$ and $E_{i,3} = 7733\ \text{kJ/mol}$. The graph plots the three energies to scale:

1. a) Calculate how many times larger $E_{i,2}$ is than $E_{i,1}$ and how many times larger $E_{i,3}$ is than $E_{i,2}$.
2. b) Between which ionisations is the "jump"? How many valence electrons does that atom have? Justify your answer.
3. c) Knowing that its configuration ends in $3s^2$, identify the element and explain why removing the third electron costs so much energy: where does that electron come from?

Selenium ($Z = 34$) is an essential trace element that appears in the composition of some dietary supplements.

1. a) Write its electron configuration and deduce its period and group.
2. b) Predict: its most common ion, the formula of its compound with sodium, and the formula of its hydride (the compound with hydrogen analogous to water).
3. c) Compare its reactivity with that of sulphur, which is directly above it in the same group, and justify it using the trend for non-metals going down a group.

The photoelectric cells in some industrial safety barriers contain caesium: visible light is enough to dislodge electrons from its surface.

1. a) Relate this property to the first ionisation energy of caesium ($376\ \text{kJ/mol}$, the lowest among alkali metals) and to its position in Group 1.
2. b) Sodium has $E_{i,1} = 496\ \text{kJ/mol}$: using the group trend, explain why caesium is more suitable than sodium for this application.
3. c) Sodium and potassium are stored under oil: use group trends to explain what would happen if they were stored in the open air and why lithium can be handled with slightly less urgency.

The reaction of sodium with water is $2\ \text{Na} + 2\ \text{H}_2\text{O} \rightarrow 2\ \text{NaOH} + \text{H}_2$.

1. a) In a laboratory demonstration, $0.2\ \text{mol}$ of sodium is used: how many moles of $\text{H}_2$ are given off? What mass of hydrogen is this ($M(\text{H}) = 1\ \text{g/mol}$)?
2. b) What mass of $\text{NaOH}$ ($M = 40\ \text{g/mol}$) is formed in that same demonstration?
3. c) The reaction releases a lot of heat and the hydrogen can burn: explain, using the reactivity trend of Group 1, why carrying out the same demonstration with potassium is even more dangerous.

If the first ionisation energy is plotted against atomic number $Z$ for the first $20$ elements, a sawtooth graph is obtained: peaks at He, Ne and Ar, troughs at Li, Na and K, and a general rise within each period. The diagram shows this representation:

1. a) Explain, using their electron configurations, why the peaks correspond to the noble gases.
2. b) Explain why the troughs correspond to the alkali metals.
3. c) Without looking up values, predict at which element around $Z = 36$ the next peak and the next trough will appear: justify both predictions.

The compound $\text{NaCl}_2$ does not exist, whereas $\text{MgCl}_2$ does; nor does $\text{MgCl}$ exist.

1. a) Using the electron configurations of sodium and magnesium, explain what would have to happen for the $\text{Na}^{2+}$ ion to exist and why this does not occur (think about energy required for the second ionisation).
2. b) Explain why the $\text{Mg}^{+}$ ion does not appear in ionic compounds, whereas $\text{Mg}^{2+}$ does.
3. c) Summarise by generalising: how does the periodic table "decide" the formulae of ionic compounds? Write the rule for a Group $g$ metal that forms $\text{X}^{g+}$ ions combined with a halogen that forms $\text{Y}^{-}$ ions.