3.2 PHYSICAL CHEMISTRY Flashcards
(62 cards)
1
Q
Define chemical energy
A
- The energy held within the bonds between atoms
2
Q
Define enthalpy
A
- A measure of the heat/thermal energy content of a substance
3
Q
Define enthalpy change ∆H
A
- The heat exchange with the surroundings during a chemical reaction
4
Q
Define system and surroundings
A
System = the atoms and bonds in the chemical reaction
Surroundings = everything around
5
Q
Describe the relationship between energy and system/surroundings
A
- Heat/energy loss in a chemical system = heat/energy gain to the surroundings
- Heat/energy gain in a chemical system = heat/energy loss to the surroundings
6
Q
Draw a labled exothermic diagram for enthalpy change
A
7
Q
Draw a labled endothermic diagram for enthalpy change
A
8
Q
Describe the term exothermic reaction in three ways
A
1) Heat energy is given out to surroundings from chemical reaction
2) More energy is released when product bonds are formed than the energy absorbed when breaking the reactant bonds
3) The enthalpy of reactants higher than enthalpy of products (thus, negative value)
9
Q
Is bond breaking an exothermic or endothermic process?
A
- Endothermic
10
Q
Is bond forming an exothermic or endothermic process?
A
- Exothermic
11
Q
State and lable the two equations needed to work out ∆H from heat energy
A
12
Q
Define activation energy Eₐ
A
- The minimum amount of energy required for a reaction to take place
13
Q
Define calorimetry
A
- Working out the enthalpy change of a reaction
14
Q
State what is measured during calorimetry to determine enthalpy change
A
- The temperature change of a chemical system (the reaction)
15
Q
Conclude the type of reaction if a temperature rise is detected during calorimetry
A
- The reaction is exothermic
- Heat energy has been released to the surroundings
16
Q
Conclude the type of reaction if a temperature drop is detected during calorimetry
A
- The reaction is endothermic
- Heat energy has been taken in from the surroundings
17
Q
State the equation used after calorimetry
A
18
Q
State why a polystyrine cup is used in a beaker during calorimetry
A
- Polystyrine acts as an insulator to reduce heat loss by evapourisation
19
Q
Draw and lable a graph for calorimetry
A
20
Q
Describe the term copper calorimetry for combustion enthalpy
A
- When the heat energy from the combustion of a fuel is used to increase the temperature of a known mass of water
21
Q
State the three reasons why experimental values differ from data values for combustion enthalpy
A
1) Non-standard conditions (298K 100kPa)
2) Incomplete combustion
3) Potential heat loss to surroundings
22
Q
Define average bond enthalpy
A
- The energy needed to break 1 mole of gaseous bonds in GASEOUS MOLECULES to form gaseous atoms
23
Q
State why bond enthalpy
values are always positive
A
- Because bond breaking is an endothermic process
24
Q
Describe what a smaller bond enthalpy indicates about that bond
A
Smaller bond enthalpy = weaker bond = less energy to break = more reactive
25
State the **equation** for **∆Hᶠ** in a **chemical reaction**
26
State the **equation** for **∆Hᶜ** in a **chemical reaction**
27
Define **standard enthalpy change of formation ∆Hᶠ**
- The enthalpy change when **1 mole** of a **compound** is **formed** from its **constituent elements** under standard conditions
28
Define **standard enthalpy change of neutralisation ∆Hₙₑᵤₜ**
- The enthalpy change when **1 mole** of **water** is **formed** from a **neutralisation reaction** under standard conditions
29
Define **standard enthalpy change of combustion ∆Hᶜ**
- The enthalpy change when **1 mole** of a **substance** is **completely combusted** under standard conditions
30
Draw an **enthaply change** of **reaction** from **enthalpy change** of **combustion**
31
Draw an **enthaply change** of **reaction** from **enthalpy change** of **formation**
32
State why **initial rate** (**T=0**) has the **highest rate** of **reaction**
- Because **concentrations** of **reactants** are **highest** at the start of a reaction
- **Gradient** of **tangent** to the curve is also **steepest** here
33
State and describe the **five** factors that affect the **rate** of **reaction**
1) **Temperature** (**↑ temp** = **↑ rate**)
2) **Pressure** (for **gases**) (**↑ pressure** = **↑ rate**)
3) **Concentration** (**↑ conc** = **↑ rate**)
4) **Surface area** (**↑ SA** = **↑ rate**)
5) **Catalyst** (**+ cat** = **↑ rate**)
34
State in terms of **collision theory** the effect of these **factors (except catalysts)** on the **rate** of **reaction**
- They ensure that **particles** **collide** with **sufficient energy** to **overcome** **activation energy** **frequently**
- Thus, **ROR increases**
35
State the **equation** for **rate** of **reaction**
36
State the **units** for **rate** of **reaction**
**mol dm⁻³ s⁻¹**
37
State the **three** ways in which **changing quantitys** can be **measured** for **rate** of **reaction**
1) **Concentration** of a **reactant** or **product** (titration)
2) **Gas volumes** of **products** (gas syringe)
3) **Mass** of **products forming**/**decreasing mass of reactants** (balance)
38
State the **equation** for **rate** of **reaction** on a **graph**
39
Describe how to determine **rate** of **reaction** from a **graph**
1) Draw a **tangent** to the **curve** wherever required
2) Use the formula **gradient=∆y/∆x**
40
Define **catalyst**
- **Increase** the **rate** of a **reaction** without being **used up** by the overall reaction
- Allows the **reaction** to **proceed** via a **different route** by **lowering activation energy**
41
State and explain the **two** types of **catalysts**
1) **Homogenous catalysts** (where the **catalyst** and **reactants** are in the **same** **phase/state**)
2) **Heterogenous catalysts** (where the **catalyst** and **reactants** are in the **different** **phase/state**)
42
Describe the **two** benefits of using **catalysts**
1) **Reduces** the **energy demand** from **combustion** of **fossil fuels**
2) **Increases economic sustainability** because it **lowers temperatures**
43
Explain what the **boltzmann distribution** shows
- The **distribution** of **energies** of **particles** at a particular **temperature**
- The **proportion** of **particles** with **sufficient energy** to **overcome** **activation energy**
44
Draw and lable a generic **boltzmann distribution** and state the features
45
Explain the effect of **temperature** on **reaction rate** via a labled **boltzmann distribution**
46
Explain the effect of **catalysts** on **reaction rate** via a labled **boltzmann distribution**
47
State the **three** features required for an **equilibrium** to be **dynamic**
1) **Closed system**
2) **Concentrations** of **reactants** and **products** remaining **constant**
3) **Rate** of **forward** and **backward** reactions to be **equal**
48
State the **three** features that affect **position** of **equilibrium**
1) **Temperature**
2) **Gaseous pressure**
3) **Concentration**
49
Explain fully the effect of **changing concenteration** on **equilibrium position**
50
Explain fully the effect of **changing gaseous pressure** on **equilibrium position**
51
Explain fully the effect of **changing temperature** on **equilibrium position**
52
State why **catalysts** do not affect the **position** of **equilibrium**
- Because they **speed up** the **rate** of the **foward** and **backward** reaction by **equal amounts**
53
Describe **industrial conditions**
- They tend to be a **comprimise** between **rate** of **reaction** and **yeild** of **production** (**equilibrium position**)
- To ensure **profits** while still being **safe**
54
Draw a **concentration-time graph** identifying **dynamic equilibrium** of **reactants** and **products** when **equilibrium** lies to the **LEFT**
55
Draw a **concentration-time graph** identifying **dynamic equilibrium** of **reactants** and **products** when **equilibrium** lies to the **RIGHT**
56
Draw a **concentration-time graph** identifying **dynamic equilibrium** of **reactants** and **products** when **equilibrium** lies in the **MIDDLE**
57
Define **homogenous equilibrium**
- Where all **species** are in the same **phase/state**
58
State the **equilibrium constant Kc** expression
59
State what is meant is **equilibrium constant Kc** has a value of **1**
- **Equilibrium** lies in the **MIDDLE** of reactants and products
60
State what is meant is **equilibrium constant Kc** has a value **bigger** than **1**
- **Equilibrium** lies to the **RIGHT**
- (**Products/forward** is **favoured**)
61
State what is meant is **equilibrium constant Kc** has a **smaller** than **1**
- **Equilibrium** lies to the **LEFT**
- (**Reactants/backwards** is **favoured**)
62
State the **units** of **equilibrium constant Kc**
**mol dm⁻³**
