Stability of atoms and ions

Isolated atoms are rarely found in nature: except for noble gases, elements occur forming compounds or metals.

1. a) Explain what it means for an atom to be "stable" and what atoms gain by bonding with one another.
2. b) What happens to the energy of the system when atoms form a bond? Does it increase or decrease?

The octet rule guides the chemistry of most elements.

1. a) State it in your own words.
2. b) Which electrons in the atom does it refer to?
3. c) What configuration do atoms reach when they satisfy it? Give an example of an element that already possesses it naturally.

Neon ($Z = 10$) is used in illuminated signs because it is a completely inert gas: it does not react with anything.

1. a) Write its electron configuration.
2. b) How many electrons does it have in its outer shell?
3. c) Explain, from this configuration, why noble gases do not form compounds under normal conditions.

When sodium combines with chlorine, each sodium atom loses an electron and becomes the $\text{Na}^+$ ion. The diagram represents the loss of the electron.

1. a) Define what a cation is and what an anion is.
2. b) State the sign of the charge for each type of ion.
3. c) Is the $\text{Na}^+$ ion a cation or an anion? Reason your answer.

Magnesium ($Z = 12$) reacts by losing electrons and forming the $\text{Mg}^{2+}$ ion.

1. a) How many electrons does each magnesium atom lose?
2. b) How many electrons does the formed ion have in total?
3. c) Write the process as a half-equation: $\text{Mg} \rightarrow \text{Mg}^{2+} + \ldots$ (complete the number of electrons released).

The chlorine ($Z = 17$) in swimming pool water acts as a disinfectant and, when reacting, each atom gains an electron forming the $\text{Cl}^{-}$ ion.

1. a) How many electrons does the chloride ion have?
2. b) Is it a cation or an anion?
3. c) What is its net electrical charge? Justify your answer by counting protons and electrons.

A particle contains $12$ protons and $10$ electrons.

1. a) Calculate its net charge.
2. b) Is it a neutral atom, a cation or an anion?
3. c) Identify the element and write the symbol of the ion.

Chlorine ($Z = 17$) and argon ($Z = 18$) are neighbours in the periodic table.

1. a) Write the electron configuration of the chlorine atom and that of the $\text{Cl}^{-}$ ion.
2. b) Compare the configuration of the chloride ion with that of the argon atom: what do you observe?
3. c) Using this result, explain why chlorine tends to gain an electron when it reacts.

The aluminium ($Z = 13$) in soft drink cans and the sulphur ($Z = 16$) in matches are very different elements, but both form stable ions.

1. a) Write the electron configuration of each atom.
2. b) Predict the most stable ion that each will form by applying the octet rule.
3. c) Write the configuration of each ion and name the noble gas whose configuration it matches.

To abbreviate configurations, noble gas notation is used: that of potassium ($Z = 19$) is written as $[\text{Ar}]\ 4s^1$.

1. a) Write its full electron configuration, without abbreviating.
2. b) How many valence electrons does potassium have?
3. c) Which ion will it form most easily? Write its configuration and explain why it is exceptionally stable.

These species seem very different: $\text{N}^{3-}$, $\text{O}^{2-}$, $\text{F}^{-}$, $\text{Na}^{+}$, $\text{Mg}^{2+}$, and $\text{Al}^{3+}$. Data: $Z(\text{N}) = 7$, $Z(\text{O}) = 8$, $Z(\text{F}) = 9$, $Z(\text{Na}) = 11$, $Z(\text{Mg}) = 12$, $Z(\text{Al}) = 13$.

1. a) Calculate the number of electrons in each species.
2. b) What do all of them have in common?
3. c) Explain what it means for them to be isoelectronic and state which noble gas they are isoelectronic with.

When converting into an ion, an atom changes in size: the radius of the sodium atom is $186\ \text{pm}$ and that of the $\text{Na}^{+}$ ion is $102\ \text{pm}$; that of the chlorine atom is $99\ \text{pm}$ and that of the $\text{Cl}^{-}$ ion is $181\ \text{pm}$. The four sizes are compared to scale in the figure.

1. a) Write the electron configurations of the two atoms and the two ions.
2. b) Explain why the cation is much smaller than the atom from which it originates.
3. c) Explain why the anion is much larger than its atom.

The species $\text{O}^{2-}$, $\text{F}^{-}$, $\text{Na}^{+}$, and $\text{Mg}^{2+}$ all have $10$ electrons, but their radii are $140\ \text{pm}$, $133\ \text{pm}$, $102\ \text{pm}$, and $72\ \text{pm}$, respectively. The figure compares the four radii to scale.

1. a) Explain why, for the same number of electrons, the size of the species depends on the number of protons in the nucleus.
2. b) Order the four species from largest to smallest radius and relate that order to their atomic number.
3. c) Without further data, predict which will have a larger radius between the $\text{N}^{3-}$ ion and the $\text{Al}^{3+}$ ion, and justify your prediction.

These are the energies required to remove successive electrons from a magnesium atom ($Z = 12$): $E_1 = 738\ \text{kJ/mol}$, $E_2 = 1451\ \text{kJ/mol}$, and $E_3 = 7733\ \text{kJ/mol}$. The figure represents the three energies.

1. a) Locate between which ionisations the large jump occurs and justify it using the electron configuration of magnesium.
2. b) Explain why magnesium forms the $\text{Mg}^{2+}$ ion and never $\text{Mg}^{3+}$ in its compounds.
3. c) Predict, reasoning from their configurations, between which ionisations the large jump will appear in sodium ($Z = 11$) and in aluminium ($Z = 13$).

Removing an electron from a neon atom costs $2081\ \text{kJ/mol}$, but removing it from the $\text{Na}^{+}$ ion, which has exactly the same electron configuration, costs $4562\ \text{kJ/mol}$. In contrast, it costs a neutral sodium atom only $496\ \text{kJ/mol}$ to lose its $3s$ electron. The figure compares the three ionisation energies.

1. a) Explain why it is harder to ionise $\text{Na}^{+}$ than $\text{Ne}$ if both have the same electrons.
2. b) Using the data, reason why sodium easily forms the $\text{Na}^{+}$ ion but never appears as $\text{Na}^{2+}$ in chemical compounds.
3. c) Connect your answers to the idea of noble gas configuration stability.