S2.1 The ionic model

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Practice exam-style IB Chemistry questions for The ionic model, aligned with the syllabus and grouped by topic.

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Question 1

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SL • Paper 1A

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A neutral atom has 13 protons and the electron arrangement 2,8,3. What are the numbers of protons and electrons in the ion it commonly forms?

A.

13 protons and 16 electrons

B.

10 protons and 13 electrons

C.

13 protons and 10 electrons

D.

16 protons and 10 electrons

Question 2

SL • Paper 1A

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SL • Paper 1A

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An atom has the electron arrangement 2,8,7. What charge is predicted for its most common ion?

A.

$7+$

B.

$2-$

C.

$1+$

D.

$1-$

Question 3

SL • Paper 1A

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SL • Paper 1A

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What is an ionic bond?

A.

The attraction between a shared electron pair and two nuclei

B.

The repulsion between ions of equal charge in a lattice

C.

The electrostatic attraction between oppositely charged ions

D.

The attraction between positive ions and delocalized electrons

Question 4

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SL • Paper 1A

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The formula of aluminium sulfate is based on $\text{Al}^{3+}$ and $\text{SO}_4^{2-}$ ions. What is the correct formula?

A.

$\text{Al}_2\text{SO}_4$

B.

$\text{AlSO}_4$

C.

$\text{Al}_2(\text{SO}_4)_3$

D.

$\text{Al}(\text{SO}_4)_3$

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Question 5

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SL • Paper 1A

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What is the formula of barium hydrogencarbonate?

A.

$\text{BaHCO}_3$

B.

$\text{Ba}_2(\text{HCO}_3)_2$

C.

$\text{Ba}(\text{HCO}_3)_2$

D.

$\text{Ba}_2\text{HCO}_3$

Question 6

SL • Paper 1A

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SL • Paper 1A

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Solid potassium bromide does not conduct electricity, but molten potassium bromide does. What is the best explanation?

A.

The ions become neutral atoms when the solid melts.

B.

Ions are fixed in the solid but mobile in the molten state.

C.

Electrons are delocalized only in the molten state.

D.

Covalent bonds form between ions in the molten state.

Question 7

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HL • Paper 1A

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An iron atom has the electron configuration $[\text{Ar}]4s^23d^6$ . What is the electron configuration of $\text{Fe}^{3+}$ ?

A.

$[\text{Ar}]4s^23d^3$

B.

$[\text{Ar}]4s^13d^4$

C.

$[\text{Ar}]3d^3$

D.

$[\text{Ar}]3d^5$

Question 8

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HL • Paper 1A

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For two acids $\text{HA}$ and $\text{HB}$ , $K_a(\text{HA})$ is larger than $K_a(\text{HB})$ . What does this imply about their conjugate bases?

A.

$\text{A}^-$ is more stable than $\text{B}^-$ because dissociation of $\text{HA}$ is more favourable.

B.

$\text{A}^-$ and $\text{B}^-$ must have the same stability because both are anions.

C.

$\text{A}^-$ is less stable than $\text{B}^-$ because dissociation of $\text{HA}$ is more favourable.

D.

$\text{B}^-$ is more stable than $\text{A}^-$ because $\text{HB}$ has the smaller $K_a$ .

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Question 9

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SL • Paper 2

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A chlorine atom has the electron configuration $1s^{2}2s^{2}2p^{6}3s^{2}3p^{5}$ .

Deduce the charge of the ion most commonly formed by chlorine.

[1]

Explain how the ion in (a) is formed, referring to its electron configuration.

[2]

Question 10

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HL • Paper 1A

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The successive ionization energies of a transition element increase gradually for the first two electrons removed, followed by a much larger increase for the third electron. What common ion charge is best supported by this pattern?

A.

$3+$

B.

$1+$

C.

$2+$

D.

$4+$

Question 11

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HL • Paper 1A

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What best explains why first-row transition elements often form cations with different charges?

A.

Their $4s$ and $3d$ subshells are close in energy.

B.

Their proton number changes during ionic bonding.

C.

Their outer shell must always reach a noble gas configuration.

D.

Their ions usually gain electrons into the $4p$ subshell.

Question 12

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HL • Paper 1A

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A Lewis representation of sulfate with two $\text{S}=\text{O}$ bonds and two $\text{S}-\text{O}$ bonds is often preferred to a representation with four single $\text{S}-\text{O}$ bonds. What is the formal-charge basis for this preference?

A.

It removes the need for any oxygen lone pairs.

B.

It makes the total charge on sulfate equal to zero.

C.

It places all negative formal charge on sulfur.

D.

It reduces the number and magnitude of non-zero formal charges.

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Question 13

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SL • Paper 2

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Ionic compounds are represented by formulas that show the simplest whole-number ratio of ions.

Deduce the formula of aluminium nitrate.

[2]

Deduce the name of $Fe_2O_3$ .

[2]

Question 14

SL • Paper 2

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SL • Paper 2

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Solid sodium chloride is commonly represented by the formula $NaCl$ .

Outline why $NaCl$ does not represent a separate molecule in the solid state.

[2]

State what is meant by ionic bonding being non-directional.

[1]

Question 15

SL • Paper 2

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SL • Paper 2

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A student tests the electrical conductivity of potassium bromide in two physical states.

Predict the relative electrical conductivity of solid and molten potassium bromide.

[1]

Explain the prediction in (a) using the ionic model.

[2]

Question 16

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SL • Paper 1B

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A spreadsheet was used to compare atoms and their most common monatomic ions for four main-group elements.

Species	Protons / count	Electrons / count
Sodium atom	11	11
Sodium ion	11	10
Magnesium atom	12	12
Magnesium ion	12	10
Oxygen atom	8	8
Oxide ion	8	10
Chlorine atom	17	17
Chloride ion	17	18

Identify the two species in the table that are anions.

[1]

Deduce the charge of the magnesium ion using the electron and proton data.

[1]

Explain why the formation of the chloride ion is described as reduction.

[2]

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Question 17

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A student prepared a table of selected ions to practise writing formulas of ionic compounds.

Ion	Formula	Charge
aluminium ion	$\text{Al}^{3+}$	$3+$
nitrate ion	$\text{NO}_3^-$	$1-$
barium ion	$\text{Ba}^{2+}$	$2+$
phosphate ion	$\text{PO}_4^{3-}$	$3-$
ammonium ion	$\text{NH}_4^+$	$1+$
carbonate ion	$\text{CO}_3^{2-}$	$2-$

Deduce the formula of aluminium nitrate.

[1]

Deduce the formula of barium phosphate.

[1]

Explain why brackets are needed in the formula of ammonium carbonate.

[2]

Question 18

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HL • Paper 1A

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In a Lewis representation of $\text{SO}_4^{2-}$ with four single $\text{S}-\text{O}$ bonds, what formal charge is assigned to the sulfur atom?

A.

$4+$

B.

$0$

C.

$2+$

D.

$2-$

Question 19

SL • Paper 2

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SL • Paper 2

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Lattice enthalpy can be used as a measure of the strength of bonding in ionic compounds.

State whether $NaF$ or $KF$ has the larger lattice enthalpy of dissociation.

[1]

Explain the answer to (a).

[2]

State why $MgO$ has a much larger lattice enthalpy than $NaF$ .

[1]

Question 20

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SL • Paper 2

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When aqueous silver nitrate is mixed with aqueous sodium chloride, a white precipitate forms.

Explain how water molecules can separate ions from an ionic lattice when a salt dissolves.

[2]

Write the ionic equation for the formation of the precipitate.

[1]

State what the formation of the precipitate indicates about silver chloride.

[1]

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Question 21

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HL • Paper 2

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Iron is a first-row transition element. The electron configuration of an iron atom may be written as $[Ar]4s^{2}3d^{6}$ .

Write the electron configuration of $Fe^{2+}$ .

[1]

Write the electron configuration of $Fe^{3+}$ .

[1]

Explain why transition elements such as iron can form cations with different charges.

[2]

Question 22

HL • Paper 2

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HL • Paper 2

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Nitric acid has a much larger acid dissociation constant, $K_a$ , than carbonic acid for its first dissociation. The conjugate bases formed are nitrate, $NO_3^-$ , and hydrogencarbonate, $HCO_3^-$ .

State which conjugate base is more stable.

[1]

Explain the relationship between conjugate base stability and $K_a$ of the conjugate acid.

[2]

Question 23

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HL • Paper 2

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Hydrogensulfate, $HSO_4^-$ , and dihydrogenphosphate, $H_2PO_4^-$ , can each act as acids to form polyatomic anions. The $K_a$ value for $HSO_4^-$ is larger than that for $H_2PO_4^-$ .

Identify the conjugate bases formed when $HSO_4^-$ and $H_2PO_4^-$ each donate one proton.

[2]

Discuss what the larger $K_a$ value for $HSO_4^-$ suggests about the relative stability of the two conjugate bases.

[2]

Question 24

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SL • Paper 1B

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Electronegativity values were used to estimate the ionic character of several binary compounds. A value of Δχ greater than about 1.8 was taken as a guideline for mainly ionic bonding.

Compound	Element 1	$\chi_1$	Element 2	$\chi_2$
AlCl3	Al	1.61	Cl	3.16
MgCl2	Mg	1.31	Cl	3.16
MgO	Mg	1.31	O	3.44
SiO2	Si	1.90	O	3.44

Calculate Δχ for magnesium oxide from the values in the table.

[1]

Identify the compound in the table with the greatest ionic character.

[1]

Explain why electronegativity difference is a model rather than a complete classification of bonding.

[3]

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Question 25

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SL • Paper 1B

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The electrical conductivity of an ionic compound was measured as a solid, as a molten liquid and as an aqueous solution.

Describe the change in conductivity when sodium chloride is melted.

[1]

Explain the difference in conductivity between solid sodium chloride and molten sodium chloride.

[2]

Suggest why pure water has a much lower conductivity than aqueous sodium chloride.

[1]

Question 26

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SL • Paper 1B

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A student mixed pairs of aqueous ionic solutions and recorded whether a precipitate formed.

cation / anion	$\text{Cl}^-(\text{aq})$	$\text{NO}_3^-(\text{aq})$	$\text{CO}_3^{2-}(\text{aq})$	$\text{SO}_4^{2-}(\text{aq})$
$\text{Ag}^+(\text{aq})$	precipitate	no precipitate	precipitate	precipitate
$\text{Ca}^{2+}(\text{aq})$	no precipitate	no precipitate	precipitate	precipitate
$\text{Na}^+(\text{aq})$	no precipitate	no precipitate	no precipitate	no precipitate
$\text{K}^+(\text{aq})$	no precipitate	no precipitate	no precipitate	no precipitate

Identify one pair of ions from the grid that forms a precipitate.

[1]

Deduce the formula and name of the precipitate formed when $Ca^{2+}(aq)$ is mixed with $CO_3^{2-}(aq)$ .

[2]

Explain why the formation of a precipitate is evidence for an ionic compound with low solubility in water.

[2]

Question 27

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HL • Paper 2

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The table shows successive ionization energies for a first-row transition element.

Ionization number	Ionization energy / kJ mol $^{-1}$
1st	650
2nd	1410
3rd	2830
4th	4510
5th	6300
6th	12360

State what a much larger jump in successive ionization energy indicates.

[1]

Explain how successive ionization energy data support variable oxidation states in transition elements.

[2]

Question 28

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HL • Paper 2

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Two possible Lewis representations of the sulfate ion are shown.

Determine the formal charge on sulfur and on each oxygen atom in Structure A.

[2]

Identify the preferred representation and justify the choice using formal charge.

[2]

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Question 29

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The melting points of three period 3 chlorides are compared.

Describe the pattern shown by the melting point data.

[1]

Explain why the melting point of aluminium chloride is not predicted well by a purely ionic model.

[3]

Question 30

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SL • Paper 1B

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Lattice enthalpies and cation radii were compared for the group 1 chlorides.

Describe the relationship shown by the graph.

[1]

Explain the trend using the ionic model.

[2]

Predict, with a reason, whether rubidium chloride or lithium chloride has the higher melting point.

[2]

Question 31

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HL • Paper 1B

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Successive ionization energies were measured for a first-row transition element, X. The element forms common ions $X^{2+}$ and $X^{3+}$ .

Describe the pattern in the first three ionization energies shown on the graph.

[1]

Explain why the gradual increase supports the formation of more than one cation by this transition element.

[2]

Explain why the $4s$ electrons are removed before the $3d$ electrons when a first-row transition metal forms a cation.

[2]

Question 32

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HL • Paper 1B

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The electron configurations of a transition element atom and two of its ions were obtained from spectroscopic data.

Species	Proton number	Electron configuration
Neutral atom	26	$[\text{Ar}]3d^{6}4s^{2}$
Ion 1	26	$[\text{Ar}]3d^{6}$
Ion 2	26	$[\text{Ar}]3d^{5}$

Identify the element from the proton number and electron configuration.

[1]

Deduce the charges of the two ions shown in the table.

[2]

Explain why this element is classified as a transition element.

[2]

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Question 33

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Polyatomic anions can be considered as conjugate bases of acids. Data for several acids and their conjugate bases are shown.

Acid	Conjugate base	pKa
Nitric acid	Nitrate ion	-1.4
Phosphoric acid	Dihydrogen phosphate ion	2.15
Ethanoic acid	Ethanoate ion	4.76
Carbonic acid	Hydrogencarbonate ion	6.35

Identify the conjugate base of nitric acid from the table.

[1]

Using the data, state which conjugate base is most stable.

[1]

Explain the relationship between conjugate base stability and the value of $K_a$ for the conjugate acid.

[3]

Question 34

SL • Paper 2

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SL • Paper 2

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Magnesium reacts with nitrogen to form solid magnesium nitride. The reaction can be interpreted using electron transfer and the ionic model.

I.

Deduce the ion formed by a magnesium atom and give its electron configuration.

[2]

II.

Deduce the ion formed by a nitrogen atom and give its electron configuration.

[2]

I.

Deduce the empirical formula of magnesium nitride and write the balanced equation for its formation from the elements.

[2]

II.

Explain why the formation of magnesium nitride from its elements is a redox reaction.

[2]

Question 35

SL • Paper 2

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SL • Paper 2

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Polyatomic ions are present in many fertilizer salts. Ammonium phosphate and calcium nitrate both contain ions that must be treated as charged units when formulas are written.

I.

Deduce the formula of ammonium phosphate from $NH_4^+$ and $PO_4^{3-}$ , explaining the charge balance.

[2]

II.

Deduce the formula of calcium nitrate from $Ca^{2+}$ and $NO_3^-$ , explaining why brackets are used.

[2]

Compare the bonding within and between the ions in ammonium phosphate.

[2]

Explain why solid ammonium phosphate does not conduct electricity but an aqueous solution of ammonium phosphate does.

[2]

Question 36

SL • Paper 2

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SL • Paper 2

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A student compares sodium chloride and silver chloride using solubility and conductivity tests. The observations are shown in Table 2.

Observation	Sodium chloride	Silver chloride
Appearance of solid	white crystalline solid	white precipitate
Conductivity of solid / mS cm $^{-1}$	0.0	0.0
After water is added	dissolves completely; clear colourless solution forms	only a small amount dissolves; cloudy suspension remains
Conductivity after water is added / mS cm $^{-1}$	12.0	0.2

I.

Deduce one observation that supports the conclusion that sodium chloride is not molecular.

[1]

II.

Explain why neither solid sodium chloride nor solid silver chloride conducts electricity.

[2]

Explain why an aqueous sodium chloride solution conducts electricity but a suspension containing solid silver chloride has much lower conductivity.

[2]

Evaluate the statement: "All ionic compounds dissolve in water."

[2]

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Question 37

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SL • Paper 2

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The ions $O^{2-}$ , $F^-$ , $Na^+$ and $Mg^{2+}$ are isoelectronic with neon.

I.

State the electron configuration common to these four ions.

[1]

II.

Explain why isoelectronic species are not the same chemical species.

[2]

Compare the ionic radii of this isoelectronic series, in terms of nuclear charge.

[2]

Use the ionic model to explain why magnesium oxide has a higher lattice enthalpy than sodium fluoride.

[2]

Question 38

HL • Paper 1B

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HL • Paper 1B

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Two possible Lewis representations of the sulfate ion were analysed using formal charge.

State the total charge on the sulfate ion in each representation.

[1]

Compare the distribution of formal charges in the two representations.

[2]

Evaluate which representation is preferred by the formal-charge model.

[2]

Question 39

HL • Paper 1B

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HL • Paper 1B

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Melting points and bonding indicators were compared for period 3 metal chlorides.

Chloride	Cation charge / $e$	Cation radius / pm	$\Delta$ EN (Pauling)	Melting point / $^\circ\text{C}$
$NaCl$	+1	102	2.2	801
$MgCl_2$	+2	72	1.8	714
$AlCl_3$	+3	54	1.5	192

Describe the trend in melting point from $NaCl$ to $AlCl_3$ shown by the data.

[1]

Explain why charge alone does not account for the melting point of $AlCl_3$ in the data.

[2]

Discuss how lattice enthalpy and the bonding continuum together explain the data.

[2]

Question 40

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HL • Paper 1B

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Successive ionization energy data for calcium and a first-row transition element were plotted on the same axes.

Identify which curve is calcium.

[1]

Explain the large jump for calcium after removal of two electrons.

[2]

Evaluate how the data support the idea that transition elements can form ions with different charges.

[2]

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Question 41

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SL • Paper 2

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Lattice enthalpy data for some ionic compounds are shown in Table 1.

Compound	Lattice enthalpy / kJ mol $^{-1}$
LiCl	853
NaCl	787
KCl	715
LiF	1036
MgO	3795

I.

Using the data, compare the lattice enthalpies of the group 1 chlorides and explain the trend.

[2]

II.

Explain why the group 2 oxide in the table has a much larger lattice enthalpy than the group 1 fluoride.

[2]

Predict and explain how the difference in lattice enthalpy between lithium chloride and potassium chloride affects their melting points and volatility.

[2]

Explain why molten lithium chloride conducts electricity whereas solid lithium chloride does not.

[2]

Question 42

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SL • Paper 2

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Melting point and electrical conductivity data for three period 3 chlorides are shown in Table 3.

Chloride	Melting point / $^\circ\text{C}$	Electrical conductivity when molten / $S\ m^{-1}$	Volatility / relative
NaCl	801	250	low
MgCl_2	714	200	low
AlCl_3	190	1	high

I.

Identify the chloride least well described by a simple ionic lattice model.

[1]

II.

Explain the evidence from the data for your answer to (a)(i).

[2]

Compare the expected lattice strength of sodium chloride and magnesium chloride using ionic charge and ionic radius.

[2]

Discuss the limitations of classifying bonding as purely ionic or purely covalent using these chlorides as examples.

[3]

Question 43

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HL • Paper 2

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Iron forms both $Fe^{2+}$ and $Fe^{3+}$ ions. A graph compares successive ionization energies for iron with those for sodium.

I.

Iron has the electron configuration $[Ar]4s^2 3d^6$ . Deduce the electron configurations of $Fe^{2+}$ and $Fe^{3+}$ .

[2]

II.

Explain why the $4s$ electrons are removed before the $3d$ electrons when iron forms cations.

[1]

Compare the ionization energy patterns to explain why iron can form ions with different charges whereas sodium mainly forms $Na^+$ .

[3]

Deduce the formula of iron(III) sulfate and explain the use of brackets in the formula.

[2]

Question 44

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HL • Paper 2

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Polyatomic anions are conjugate bases of acids. Table 4 gives acid dissociation constants for selected conjugate acids.

Conjugate acid	Conjugate base	$K_a$
$\text{HSO}_4^-$	$\text{SO}_4^{2-}$	$1.2 \times 10^{-2}$
$\text{HNO}_3$	$\text{NO}_3^-$	$2.0 \times 10^{1}$
$\text{H}_2\text{CO}_3$	$\text{HCO}_3^-$	$4.3 \times 10^{-7}$

I.

Use the $K_a$ data to rank the conjugate bases $NO_3^-$ , $HCO_3^-$ and $SO_4^{2-}$ in order of increasing stability.

[2]

II.

Explain why a larger $K_a$ value indicates a more stable conjugate base.

[2]

Deduce the formulas of magnesium hydrogencarbonate and ammonium sulfate.

[2]

Compare the bonding in ammonium sulfate with the bonding within the sulfate ion.

[2]

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Question 45

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HL • Paper 2

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Cobalt forms more than one chloride, including $CoCl_2$ and $CoCl_3$ . Cobalt atoms have the electron configuration $[Ar]4s^2 3d^7$ . Successive ionization energy data for cobalt are shown.

I.

Deduce the oxidation state of cobalt in each chloride and give the systematic name of each compound.

[2]

II.

Deduce the electron configurations of $Co^{2+}$ and $Co^{3+}$ .

[2]

Use the successive ionization energy pattern to explain why cobalt can form both $Co^{2+}$ and $Co^{3+}$ ions.

[2]

Discuss why the charges of transition metal ions are less easily predicted from group number than the charges of group 1 and group 2 ions.

[2]

Question 46

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HL • Paper 2

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Barium ions form low-solubility salts with some polyatomic anions. Solutions of barium nitrate are mixed separately with sodium sulfate and sodium carbonate. Acid dissociation data for the conjugate acids of sulfate and carbonate ions are also provided.

Mixture	Anion present	Conjugate acid	K_a at $25^\circ\text{C}$
$\text{Ba(NO}_3)_2$ (aq) + $\text{Na}_2\text{SO}_4$ (aq)	$\text{SO}_4^{2-}$ (aq)	$\text{HSO}_4^-$ (aq)	$1.2 \times 10^{-2}$
$\text{Ba(NO}_3)_2$ (aq) + $\text{Na}_2\text{CO}_3$ (aq)	$\text{CO}_3^{2-}$ (aq)	$\text{HCO}_3^-$ (aq)	$4.7 \times 10^{-11}$

I.

Deduce the net ionic equations for the formation of the barium sulfate and barium carbonate precipitates.

[2]

II.

Use the $K_a$ data to compare the stability of $SO_4^{2-}$ and $CO_3^{2-}$ as conjugate bases.

[2]

Discuss why both precipitates can have low solubility even if the anions differ in conjugate-base stability.

[2]

Evaluate how the electrical conductivity of the mixtures changes after precipitation occurs.

[2]

Question 47

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HL • Paper 2

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The sulfate ion is present in many ionic compounds. Two possible Lewis representations of $SO_4^{2-}$ are shown, with no formal charges displayed.

I.

Calculate the formal charges on sulfur and on each oxygen atom in representation A.

[2]

II.

Calculate the formal charges on sulfur, on a double-bonded oxygen atom and on a single-bonded oxygen atom in representation B.

[2]

Evaluate which representation is preferred using formal charge arguments.

[2]

Deduce the formula of aluminium sulfate and identify the bonding between aluminium ions and sulfate ions and within the sulfate ion.

[2]

Question 48

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HL • Paper 2

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Sulfate salts contain the polyatomic ion $SO_4^{2-}$ . Table 5 gives cation radius and lattice enthalpy data for selected sulfate salts.

Salt	Cation radius / pm	Lattice enthalpy / kJ mol $^{-1}$
Sodium sulfate	102	2510
Magnesium sulfate	72	2890
Aluminium sulfate	53	13800

I.

Use charge balance to deduce the formulas of magnesium sulfate and aluminium sulfate.

[2]

II.

Using the data, compare the charge density of $Mg^{2+}$ and $Na^+$ and relate this to their lattice enthalpies.

[2]

Explain why cation charge and radius affect lattice enthalpy in sulfate salts.

[2]

Evaluate whether formal charge considerations in the sulfate ion alone are sufficient to predict the melting points of sulfate salts.

[2]

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S1.5 Ideal gases

S2.2 The covalent model

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