# Topic 7 Chemical Energetics

The following are high-level thinking skills (KBAT) questions designed based on the learning outcomes provided:

What happens in exothermic and endothermic reactions?

An exothermic process releases heat, causing the temperature of the immediate surroundings to rise. An endothermic process absorbs heat and cools the surroundings.

Why does only catalyst affect the activation energy?

The only effect of the catalyst is to lower the activation energy of the reaction. A catalyst provides an alternative route for the reaction with a lower activation energy.

illustrate the reverse reaction within a reversible reaction calculation of endothermic and exothermic?

modal tolak untung, = total enthalpy used, = hess’ law = dH1 + dH2 + dH3 + …

What is enthalpy in simple terms?

Enthalpy is a concept used in science and engineering when heat and work need to be calculated. When a substance changes at constant pressure, enthalpy tells how much heat and work was added or removed from the substance. Enthalpy is similar to energy, but not the same.

Sketch the energy diagram of endothermic and exothermic reaction?

Given that as below, dH = -226 kJ mol-1, activation energy, Ea = + 134 kJ mol-1.

Sketch the energy level diagram and give the dH value of the reverse reaction?

What is the effect of temperature to the amount of reactants and products between the endothermic and exothermic reaction?

If the reaction is exothermic, an increase in temperature will cause the reverse reaction to occur, decreasing the amounts of the products and increasing the amounts of reactants. Lowering the temperature will produce the opposite response.

What happens to an exothermic reaction if the temperature is decreased?

Since exothermic reactions release heat, this means that heat is given off to the surroundings. If the temperature of an exothermic reaction is increased, the reaction will shift LEFT. If the temperature of an exothermic reaction is decreased, the reaction will shift RIGHT.

What is the effect of pressure to the amount of reactants and products between the endothermic and exothermic reaction? how does pressure affect an endothermic reaction?

If the pressure of a system is INCREASED, the reaction will shift toward the side with fewer moles of GAS. If the pressure of a system is DECREASED, the reaction will shift toward the side with more moles of GAS. If the pressure of the system with the above reaction increases, the reaction will shift RIGHT.

What happens when pressure increases in a reaction?

When you increase the pressure, the molecules have less space in which they can move. That greater density of molecules increases the number of collisions. When you decrease the pressure, molecules don’t hit each other as often and the rate of reaction decreases. Pressure is also related to concentration and volume.

is Melting endothermic or exothermic?

Melting is an endothermic reaction in which the total amount of heat in the substance, also known as the enthalpy, increases.

is Melting ice endothermic or exothermic?

Melting ice is endothermic, you can see this by putting a thermometer in a glass of warm water, adding an ice cube, and watching the temperature go down as the ice melts. The melting process needs heat to proceed and takes it from the warm water.

is frying an egg endothermic or exothermic?

Frying an egg is a chemical reaction. It is an example of an endothermic reaction, or one that takes in heat in order to make the reaction occur.

Define the lattice energy?

a measure of the energy contained in the crystal lattice of a compound, equal to the energy that would be released if the component ions were brought together from infinity. The attraction of the two ions releases energy and the process is exothermic. Lattice energy can be a very complicated process but is often simplified by using Coulomb’s law. LE = lattice energy. k = Q1 and Q2 = numerical ion charges.

Which ionic compound has lower lattice energy?

This model emphasizes two main factors that contribute to the lattice energy of an ionic solid: the charge on the ions, and the radius, or size, of the ions. The effect of those factors is: as the charge of the ions increases, the lattice energy increases. as the size of the ions increases, the lattice energy decreases.

is Lattice Energy negative or positive?

Lattice energy and lattice enthalpy The formation of a crystal lattice is exothermic, the value of ΔHlattice is negative because it corresponds (sepadan) to the coalescing (penggabungan) of infinitely separated gaseous ions (ion gas tanpa had) in vacuum to form the ionic lattice.

Which has the largest lattice energy?

MgO has the highest lattice energy.

What is difference lattice energy and enthalpy formation?

Lattice energy can be measured as the amount of energy required to break a bond between solid ions into a gas. Lattice enthalpy is the enthalpy change from when the solid structure is formed or broken.

The lattice formation enthalpy is the enthalpy change when 1 mole of solid crystal is formed from its scattered gaseous ions. Lattice formation enthalpies are always negative.

Arrange the lattice energy of group 2 carbonate?

BeCO3 < MgCO3 < CaCO3 < SrCO3 < BaCO3.

What is the effect of heat on the Group 2 carbonates?

All the carbonates in this group undergo thermal decomposition (splitting up a compound by heating it) to the metal oxide and carbon dioxide gas.

MCO3 (s, white solids)→MO (s, hite solids) + CO2 (g).

Down the group (Be, Mg, Ca, Sr, Ba), the carbonates require more heating to decompose. The carbonates become more thermally stable down the group.

Why MgO has the highest lattice energy than CaO?

In both CaO and MgO, the charges are +2 and -2, so the choice is between these two oxides. The smallest ions will be able to get closest to each other. They will have the smallest distance between centres and will have the largest lattice energies. The smallest ions are at the top of the Periodic Table. Mg2+ is smaller than Ca2+ , so MgO has the largest lattice energy.

Why does MgO have higher lattice energy than NaCl?

Magnesium ions are smaller than sodium ions, and oxide ions are smaller than chloride ions. That means that the distance between the positive and negative ions is quite a lot less in MgO than in NaCl, and so the forces of attraction will be greater in MgO.

Which has more lattice enthalpy MgO or BaO?

BaO is more soluble in water than MgO because its atomic size gives the molecule a more iconic nature and less lattice energy. The MgO molecule has a cation with a smaller radius, indicating that there is strong attraction between the cation and anion in the lattice.

Define the lattice dissociation energy?

Explain the trend solubility of sulfate of group 2 when going down the group?

Solubility decreases as you go down the group. The lattice dissociation enthalpy and hydration enthalpy both decrease as you go down the group. The hydration enthalpy decreases more than the lattice dissociation enthalpy. Therefore the enthalpy of solution becomes more endothermic (or less exothermic).

Boardworks AS Chemistry. Trends in Group 2. Solubilities of group 2 sulfates. The solubility of the group 2 sulfates decreases down the group. Magnesium and calcium sulfate are considered to be soluble, whereas strontium and barium sulfate are considered to be insoluble. Group 2 hydroxide. Solubility. Note that this decrease in solubility down the group is the opposite of the trend for the solubility of the group 2 hydroxides. MgSO4. soluble. CaSO4. slightly soluble. Teacher notes. Beryllium sulfate is soluble, following the trend shown here. SrSO4. insoluble. BaSO4. insoluble.

The relationship between enthalpy of solution and solubility. The more endothermic (or less exothermic) the enthalpy of solution is, the less soluble the compound. So sulphates and carbonates become less soluble as you go down the Group but hydroxides become more soluble.

Why do Group 2 sulphates become less soluble? the hydration enthalpy decreases faster than the lattice enthalpy in the Group 2 sulphates, the solubility of Group 2 sulphates decreases when down the group.

Which sulphates are insoluble in water? Calcium sulphate is only very slightly soluble in water. Strontium and barium sulphates are virtually insoluble in water.

Why is MgSO4 more soluble than BaSO4? Having more hydration energy MgSO4 is more soluble than BaSO4. MgSO4 is soluble in water but BaSO4 is not becuase the size of Ba and SO4 ions is very large which leads to higher lattice enthalpy than that of hydration enthalpy.

Are Group 1 carbonates soluble? By contrast, the least soluble Group 1 carbonate is lithium carbonate. … Solubility of the carbonates increases as you go down Group 1. The hydroxides. The least soluble hydroxide in Group 1 is lithium hydroxide – but it is still possible to make a solution with a concentration of 12.8 g per 100 g of water at 20°C.

What are the solubility rules for ionic compounds?

Solubility Rules for Ionic Compounds in Water.

Rule 1: Compounds of NH4+ and group 1A (group 1, alkali metal) metal ions are soluble.

Rule 2: Compounds of NO3–, ClO4–, ClO3– and C2H3O2– (acetate ion, CH3CO2-) are soluble.

Rule 3: Compounds of Cl –, Br – and I – are soluble except those of Ag+, Cu+, Tl+, Hg22+ and Pb2+.

Rule 4: Compounds of SO42– are soluble except those of Ca2+, Sr2+, Ba2+ and Pb2+.

Rule 5: Most other ionic compounds are insoluble.

What ions are always soluble?

1) Salts of ammonium and alkali metals (column 1A excluding hydrogen) are always soluble.

2) All chlorides, bromides, and iodides are soluble except when combined with Ag, Hg2+, and Pb which are insoluble.

3) Chlorates, acetates, and nitrates (CANs) are soluble.

Explain statements below regarding to solubility of an ionic compounds in water?

Statement-1 : Solubility of ionic compounds in water depends on both the lattice energy and the hydration energy.

Statement-2 : Ionic compounds dissolve in water when their hydration energy exceeds the lattice energy.

Answer: Statement-1 is true , statement-2 is true , statement-2 is a correct explanation for statement-1.

What is the relationship between enthalpy of solution and solubility?

The assumption is made that the more endothermic (or less exothermic) the enthalpy of solution is, the less soluble the compound. the more endothermic- less soluble. So sulphates and carbonates become less soluble as you go down the Group; hydroxides become more soluble. More +ve dH, less stable.

Can enthalpy of solution predict solubility?

In fact, most compounds that are soluble in water have positive enthalpies of solution. Therefore, from the equation ΔG = ΔH – TΔS we should predict that the solubility of every compound should increase with increasing temperature. That prediction turns out to be correct for nearly every solvent and solute.

Why are hydroxides more soluble down the group?

Solubility of the Hydroxides. Group 2 metal hydroxides become more soluble in water as you go down the group. The decrease in the lattice energy of the hydroxide salt and by the increase in the coordination number of the metal ion as you go down the column.

What is the enthalpy of a solution?

The enthalpy change of solution refers to the amount of heat that is released or absorbed during the dissolving process (at constant pressure). This enthalpy of solution (ΔHsolution) can either be positive (endothermic) or negative (exothermic).

Determine the enthalpy change of solution’s concept of sulubility by using the diagram’s below?

What is the difference between enthalpy of solution and enthalpy of hydration?

The enthalpy change of solution is the enthalpy change when 1 mole of an ionic substance dissolves in water to give a solution of infinite dilution. The hydration enthalpy is the enthalpy change when 1 mole of gaseous ions dissolve in sufficient water to give an infinitely dilute solution.

Which element has the highest hydration enthalpy?

lithium ion The attractions are stronger the smaller the ion. For example, hydration enthalpies fall as you go down a group in the Periodic Table. The small lithium ion has by far the highest hydration enthalpy in Group1, and the small fluoride ion has by far the highest hydration enthalpy in Group 7.

Explain the trend hydration energy of sulfate of group 2 when going down the group?

Explain the trend lattice dissociation energy of sulfate of group 2 when going down the group?

How to calculate the enthalpy change of reaction when given the standard enthalpy of formation? NH3 (-46 kJ mol-1), HF (-269 kJ mol-1) and NF3 (-114 kJ mol-1).

What is the formula of finding standard enthalpy change of formation?

standard enthalpy change of formation is equal to the sum of the standard enthalpies of formation of the products minus the sum of the standard enthalpies of formation of the reactants.

What is the enthalpy change of fluorine gas, F2?

zero

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7 Chemical Energetic

7.1 Enthalpy changes of reaction, dH

1. Arrange in table form, explain the similarity and differences of exothermic and endothermic?

2. Define enthalpy change of reaction, H, and state the standard conditions?

3. Define enthalpy change of formation, combustion, hydration, solution, neutralisation, atomisation, bond energy, ionisation energy and electron affinity;

4. Give a solution of a questions as an example on how we would calculate the heat energy change from experimental measurements using the relationship: heat change, q mcT or q = mc? Find your own question on any reference book offline or online.

5. Give a solution of a questions as an example on how we would calculate enthalpy changes from experimental results? Find your own question on any reference book offline or online.

7.2 Hess’ law

1. Give a solution of a questions as an example on how we would find enthalpy changes that cannot be determined directly by using Hess’ law?
e.g. an enthalpy change of formation from enthalpy changes of combustion.

2. Give a solution of a questions as an example on how we would construct energy level diagrams relating the enthalpy to reaction path and activation energy?

3. Give a solution of a questions as an example on how we would calculate enthalpy changes from energy cycles?

7.3 Born-Haber cycle

1. Define lattice energy for simple ionic crystals in terms of the change from gaseous ions to solid lattice?

2. Explain qualitatively the effects of ionic charge and ionic radius on the numerical magnitude of lattice energy values?

3. Construct Born-Haber cycle for the formation of simple ionic crystals?

7.4 The solubility of solids in liquids

1. Give a solution of a questions as an example on how we would construct energy cycles for the formation of aqueous solutions of ionic compounds?

2. Explain qualitatively the influence on solubility of the relationship between enthalpy change of solution, lattice energy of solid and enthalpy change of hydration or other solvent-solute interaction?

## 7 Chemical Energetics

### 7.1 Enthalpy changes of reaction, dH

Candidates should be able to:

(a) explain that most chemical reactions are accompanied by enthalpy changes (exothermic or endothermic);

(b) define enthalpy change of reaction, H, and state the standard conditions;

(c) define enthalpy change of formation, combustion, hydration, solution, neutralisation, atomisation, bond energy, ionisation energy and electron affinity;

(d) calculate the heat energy change from experimental measurements using the relationship: heat change, q mcT or q = mc ;

(e) calculate enthalpy changes from experimental results.

### 7.2 Hess’ law

Candidates should be able to:

(a) state Hess’ law, and its use to find enthalpy changes that cannot be determined directly, e.g. an enthalpy change of formation from enthalpy changes of combustion;

(b) construct energy level diagrams relating the enthalpy to reaction path and activation energy;

(c) calculate enthalpy changes from energy cycles.

### 7.3 Born-Haber cycle

Candidates should be able to:

(a) define lattice energy for simple ionic crystals in terms of the change from gaseous ions to solid lattice;

(b) explain qualitatively the effects of ionic charge and ionic radius on the numerical magnitude of lattice energy values;

(c) construct Born-Haber cycle for the formation of simple ionic crystals.

### 7.4 The solubility of solids in liquids

Candidates should be able to:

(a) construct energy cycles for the formation of aqueous solutions of ionic compounds;

(b) explain qualitatively the influence on solubility of the relationship between enthalpy change of solution, lattice energy of solid and enthalpy change of hydration or other solvent-solute interaction.

### Definition

Hess’ Law of Constant Heat Summation (or just Hess’s Law) states that regardless of the multiple stages or steps of a reaction, the total enthalpy change for the reaction is the sum of all changes.

Standard Enthalpy Change of Reaction, ΔHr: is the enthalpy change that occurs in a system when matter is transformed by a given chemical reaction, when all reactants and products are in their standard states.

Standard Enthalpy Change of Formation, ΔHf: the change in enthalpy when one mole of a substance is formed from its pure elements under standard condition (1 atm of pressure and 298.15 K).

Standard Enthalpy of Combustion, ΔHc: of a compound is the enthalpy change which occurs when one mole of the compound is burned completely in oxygen under standard conditions, and with everything in its standard state.

Standard Enthalpy Change of Neutralisation, ΔHneut: is the enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce 1 mole of water. Notice that enthalpy change of neutralisation is always measured per mole of water formed.

Standard Enthalpy Change of Atomisation, ΔHatm: the enthalpy change when 1 mol of an element in its standard state is atomised to produce 1 mol of gaseous atoms.

Standard Enthalpy Change of Hydration, ΔHhyd: the enthalpy change when 1 mole of gaseous ions dissolve in sufficient water to give an infinitely dilute solution. Hydration enthalpies are always negative.

Standard Enthalpy Change of Solution, ΔHsoln: the enthalpy change when 1 mole of an ionic substance dissolves in water to give a solution of infinite dilution.

Standard Bond Enthalpy: also known as bond energy is defined as the amount of energy required to break one mole of the stated bond.

Standard Ionisation Enthalpy: of elements is the amount of energy that an isolated gaseous atom requires to lose an electron in its ground state.

Heat of vaporization: the energy required to vaporize one mole of a liquid at a pressure of one atmosphere.

Electron Affinity: the change in energy (in kJ/mole) of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion. In other words, the neutral atom’s likelihood of gaining an electron.

Lattice Energy: the energy required to break apart an ionic solid and convert its component atoms into gaseous ions.

Born-Haber Cycle: applies Hess’s law to calculate the lattice enthalpy by comparing the standard enthalpy change of formation of the ionic compound (from the elements) to the enthalpy required to make gaseous ions from the elements.

## 7 Chemical Energetics

### 7.1 Enthalpy changes of reaction, ΔH

Candidates should be able to:
(a) explain that most chemical reactions are accompanied by enthalpy changes (exothermic or endothermic);

Standard Enthalpy Change of Reaction, ΔHr: is the enthalpy change that occurs in a system when matter is transformed by a given chemical reaction, when all reactants and products are in their standard states.

(b) define enthalpy change of reaction, H, and state the standard conditions;

(c) define: enthalpy change of formation,

Standard Enthalpy Change of Formation, ΔHf: the change in enthalpy when one mole of a substance is formed from its pure elements under standard condition (1 atm of pressure and 298.15 K).

define: combustion,

Standard Enthalpy of Combustion, ΔHc: of a compound is the enthalpy change which occurs when one mole of the compound is burned completely in oxygen under standard conditions, and with everything in its standard state.

define: hydration,

Standard Enthalpy Change of Hydration, ΔHhyd: the enthalpy change when 1 mole of gaseous ions dissolve in sufficient water to give an infinitely dilute solution. Hydration enthalpies are always negative.

define: solution,

Standard Enthalpy Change of Solution, ΔHsoln: the enthalpy change when 1 mole of an ionic substance dissolves in water to give a solution of infinite dilution.

define: neutralisation,

Standard Enthalpy Change of Neutralisation, ΔHneut: is the enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce 1 mole of water. Notice that enthalpy change of neutralisation is always measured per mole of water formed.

define: atomisation,

Standard Enthalpy Change of Atomisation, ΔHatm: the enthalpy change when 1 mol of an element in its standard state is atomised to produce 1 mol of gaseous atoms.

define: bond energy,

Standard Bond Enthalpy: also known as bond energy is defined as the amount of energy required to break one mole of the stated bond.

define: ionisation energy

Standard Ionisation Enthalpy: of elements is the amount of energy that an isolated gaseous atom requires to lose an electron in its ground state.

define: electron affinity

Electron Affinity: the change in energy (in kJ/mole) of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion. In other words, the neutral atom’s likelihood of gaining an electron.

(d) calculate the heat energy change from experimental measurements using the relationship:
heat change, q mcT or q = mc ;

(e) calculate enthalpy changes from experimental results.

## 7.2 Hess’ law

Candidates should be able to:
(a) state Hess’ law, and its use to find enthalpy changes that cannot be determined directly, e.g. an enthalpy change of formation from enthalpy changes of combustion;

Hess’ Law of Constant Heat Summation (or just Hess’s Law) states that regardless of the multiple stages or steps of a reaction, the total enthalpy change for the reaction is the sum of all changes.

(b) construct energy level diagrams relating the enthalpy to reaction path and activation energy;

(c) calculate enthalpy changes from energy cycles.

### 7.3 Born-Haber cycle

Candidates should be able to:
(a) define lattice energy for simple ionic crystals in terms of the change from gaseous ions to solid lattice;

Lattice Energy: the energy required to break apart an ionic solid and convert its component atoms into gaseous ions.

(b) explain qualitatively the effects of ionic charge and ionic radius on the numerical magnitude of lattice energy values;

(c) construct Born-Haber cycle for the formation of simple ionic crystals.

Born-Haber Cycle: applies Hess’s law to calculate the lattice enthalpy by comparing the standard enthalpy change of formation of the ionic compound (from the elements) to the enthalpy required to make gaseous ions from the elements.

### 7.4 The solubility of solids in liquids

Candidates should be able to:

(a) construct energy cycles for the formation of aqueous solutions of ionic compounds;

(b) explain qualitatively the influence on solubility of the relationship between enthalpy change of solution, lattice energy of solid and enthalpy change of hydration or other solvent-solute interaction.