Brainscape's Knowledge GenomeTM


Top Flashcards (Certified)
All Flashcards

chemistry > Module 3 - Inorganic and Physical Chemistry > Flashcards

Module 3 - Inorganic and Physical Chemistry Flashcards

1
Q

How are elements arranged in the periodic table?

A

By increasing atomic (proton) number, each consecutive element has 1 extra proton.

Vertical columns are groups (same number of outer shell electrons) and horizontal rows are periods (number of the highest energy electron shell).

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

How do groups and periods relate to properties of elements?

A

Groups - elements have similar properties; Periods - show repeating trends in physical and chemical properties.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

What is periodicity?

A

A repeating trend in properties of elements.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

What properties do we look at in periodicity?

A
  • Electron configuration
  • Ionisation energy
  • Structure
  • Melting points
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Describe the trend in electron configuration across a period.

A

Elements have different number of valence electrons, can have different types of orbitals in outer shell, outer electrons in the same quantum shell.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Describe the trend in electron configuration down a group.

A

Elements have the same number of valence electrons, same type of orbital in the outer shell, outer electrons in different quantum shells.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

How is the periodic table divided into blocks?

A

4 blocks corresponding to the highest energy sub-shell of the elements.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

State and explain the trend in first ionisation energy down a group.

A

Decreases down a group; although nuclear charge increases, the effect is outweighed by increased atomic radius and increased shielding.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

State and explain the trend in first ionisation energy across a period.

A

Increases across a period; for the equivalent level of shielding, nuclear charge increases and atomic radius decreases.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Describe the first ionisation energy trend in detail across period 2.

A

Rise from Li to Be, fall from Be to B followed by rise to C and N, fall from O followed by rise to F.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Explain why there is a fall between Be and B across period 2 despite ionisation energy ‘increasing’.

A

In boron, the 2p sub-shell has higher energy than the 2s sub-shell in beryllium, making it easier to remove a 2p electron.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Explain why there is a fall between N and O across period 2 despite ionisation energy ‘increasing’.

A

In oxygen, one of the 2p orbitals contains a pair of electrons, which leads to slight repulsion, making it easier to remove an electron compared to nitrogen.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Describe the melting point trend in detail across periods 2 and 3.

A

Increases from Group 1 to Group 4, sharp decrease in melting point between Group 4 and Group 5, remains comparatively low from Group 5 to Group 0.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Explain why group 4 elements occupy the peaks in melting point across periods 2 and 3.

A

They form giant covalent lattices which require a lot of energy to break due to strong covalent bonds.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Explain why elements with diatomic elements/single atoms occupy the troughs on the melting point graph.

A

They form simple molecular structures with weak London forces, requiring less energy to break and melt the substances.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

Explain why melting point increases from Group 1 to Group 2 to Group 3.

A

They are metals forming giant metallic lattices; metallic bonding increases in strength with more delocalised electrons.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

How do you determine if a reaction is REDOX?

A

If oxidation number of an element increases - oxidation; if it decreases - reduction.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

What is a reducing agent?

A

Something that donates electrons to enable reduction to occur.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

What is an oxidising agent?

A

Something that accepts electrons, enabling oxidation to occur.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

Explain the most common way group 2 elements react.

A

They react by redox, getting oxidised and losing 2 electrons to form 2+ ions with noble gas configuration.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

What are the 3 main reactions of group 2 metals?

A
  • With oxygen
  • With water
  • With dilute acids
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

How do group 2 metals react with oxygen?

A

They react with oxygen to form a metal oxide (MO) made up of M2+ and O2-.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

How do group 2 metals react with water?

A

They react to form an alkaline hydroxide (M(OH)2) and hydrogen gas; reactivity increases down the group.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

How do group 2 metals react with dilute acids?

A

They react with dilute acids to form a salt and hydrogen gas.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Explain the trend in reactivity down group 2.
Less energy required to oxidise, making the metal more reactive; ionisation energy decreases down the group due to increased atomic radius and shielding.
26
How do group 2 oxides react with water?
They react with water, releasing hydroxide ions and forming an alkaline solution of the metal hydroxide.
27
How does solubility of hydroxides change down group 2?
Solubility increases down the group, resulting in more OH- ions and a more alkaline solution.
28
What's the estimated pH of magnesium hydroxide solution?
Around 10.
29
What is the estimated pH of a barium hydroxide solution?
Around 13.
30
How do you show that solubility increases down group 2 for hydroxides?
Add spatula of each G2 oxide to water, shake, and measure pH of each solution.
31
How are group 2 oxides, hydroxides and carbonates classified?
They are bases - H+ acceptors.
32
Give an example of how G2 compounds are used in agriculture.
Calcium hydroxide is added to fields as lime to increase pH of acidic soils.
33
Give an example of how G2 compounds are used in medicine.
Used as antacids to treat acid indigestion; many use magnesium and calcium carbonates.
34
What are the characteristic physical properties of G2 metals?
* Fairly high melting points (compared to G1) * Light metals with low densities * Colorless (white) compounds
35
Describe characteristic physical properties of each halogen.
Varies among halogens; generally they exist as diatomic molecules.
36
State and explain the trend in boiling point down group 7.
Increases down the group due to stronger London forces from more electrons.
37
Explain the most common way group 7 elements react.
They get reduced, gaining 1 electron to form 1- halide ions with noble gas configuration.
38
What are halogen displacement reactions?
Redox reaction where a more reactive halogen displaces a less reactive halogen from a compound.
39
Explain how to use halogen-halide displacement reactions to show that reactivity of halogens decreases down the group.
Add solution of each halogen to aqueous solutions of the other halogens; if more reactive, it displaces and changes color.
40
What are the colors of the halogens in water vs in cyclohexane?
Varies based on halogen; non-polar halogens dissolve more readily in cyclohexane.
41
Why don't we test for fluorine and astatine?
Fluorine reacts with almost any substance; astatine is rare, radioactive, and has never been seen.
42
State and explain the trend in reactivity down group 7.
Reactivity decreases down the group due to increased atomic radius and shielding, reducing nuclear attraction.
43
What is disproportionation?
Redox reaction where the same element is both oxidised and reduced.
44
What is the reaction of chlorine with water?
Chlorine reacts with water to form chloric(I) acid and hydrochloric acid.
45
How does adding chlorine to water kill bacteria?
Disproportionation forms chloric(I) acid, which kills bacteria.
46
What is the reaction of chlorine with cold, dilute aqueous sodium hydroxide?
Forms sodium chlorate(I), sodium chloride, and water.
47
What are the benefits and risks of chlorine use?
* Benefits: Disinfects water * Risks: Toxicity and potential formation of chlorinated hydrocarbons.
48
What is enthalpy?
Measure of heat energy in a chemical system.
49
Explain why chemical reactions generally involve changes in enthalpy.
Reactants and products have different bonds and structures.
50
Give the formula for enthalpy change.
H∆ = H(products) - H(reactants).
51
State the law of conservation of energy.
Energy cannot be created or destroyed.
52
Explain what is meant by 'system' and 'surroundings' and 'universe'.
* System - reactants and products * Surroundings - apparatus, lab * Universe - everything including system and surroundings.
53
Define exothermic.
Energy transferred from system to surroundings.
54
Define endothermic.
Energy transfer from surroundings to the system.
55
Explain what happens when ∆H is -ve.
Reactants have more energy than products, releasing heat energy to surroundings.
56
Explain what happens when ∆H is +ve.
Reactants have less energy than products, gaining heat energy from surroundings.
57
Define activation energy.
Minimum energy required to start a reaction by breaking bonds.
58
Explain why most chemical reactions have an activation energy.
Necessary to break bonds in reactants to form new ones.
59
What are the standard conditions?
* Pressure = 100 kPa * Temperature = 25ºC = 298K * Concentration = 1 mol dm-3
60
What is an endothermic enthalpy profile diagram?
A diagram that shows the enthalpy changes during an endothermic reaction, including the activation energy (Ea).
61
Why do we have standard enthalpy changes & standard conditions?
To ensure consistency in measurements, as enthalpy change can vary slightly depending on conditions.
62
What are the standard conditions?
Pressure = 100kPa, Temperature = 25ºC (298K), Concentration = 1mol dm-3.
63
What is a standard state?
The physical state of a substance under standard conditions.
64
What is ∆rHᶿ?
Standard enthalpy change of reaction - enthalpy change accompanying a reaction in the molar quantities shown in a chemical equation under standard conditions.
65
What is ∆fHᶿ?
Standard enthalpy change of formation - enthalpy change when 1 mol of a compound is formed from its elements under standard conditions.
66
What is the ∆fHᶿ of elements?
0.
67
What is ∆neutHᶿ?
Standard enthalpy change of neutralisation - energy change accompanying the reaction of an acid with a base to form one mole of H2O(l) under standard conditions.
68
What is ∆cHᶿ?
Standard enthalpy change of combustion - enthalpy change when one mole of a substance reacts completely with oxygen under standard conditions.
69
How do we measure enthalpy changes?
1. Take known amount of starting material. 2. Measure temperature in reaction mixture. 3. Calculate heat energy exchanged: (q /J=mc Delta T). 4. Calculate enthalpy change: (∆H = {q /kJ over n}). 5. Decide sign of ∆H based on whether exothermic or endothermic.
70
What are the common equipment required for measuring ∆Hc?
Calorimeter, thermometer, and reaction vessel.
71
What is the method for measuring ∆rH of two solutions?
Mix two known volumes of solutions and measure temperature change.
72
What are the reasons for inaccuracies when measuring ∆cH?
1. Heat loss to surroundings. 2. Incomplete combustion. 3. Evaporation of the liquid. 4. Non-standard conditions.
73
What assumptions are made when measuring ∆rH?
1. Density of solution is the same as water. 2. Specific heat capacity of solution is the same as water. 3. Heat capacity of equipment is negligible.
74
Define average bond enthalpy.
Energy required to break 1 mol of a specified type of bond in a gaseous molecule.
75
What is the sign of bond enthalpy?
+ve because energy is always required to break bonds - endothermic.
76
What are the limitations of average bond enthalpies?
1. Actual bond enthalpy varies depending on chemical environment. 2. Average is calculated from actual bond enthalpies in different environments.
77
Explain bond breaking and making in terms of energy and enthalpy changes.
Bond breaking requires energy (endothermic, +∆H), while bond making releases energy (exothermic, -∆H).
78
State the formula for calculating enthalpy changes from average bond enthalpies.
∆rH = ∑(bond enthalpies in reactants) - ∑(bond enthalpies in products).
79
What is Hess' law?
∆H for any chemical reaction is independent of the intermediate stages as long as the initial and final conditions are the same.
80
What is the rate of reaction?
How fast a reactant is being used up or how fast a product is being formed.
81
What is the formula and units for rate?
Rate = change in concentration / change in time (usually mol/dm³/s).
82
Explain the shape of the rate curve.
Fastest at the start with a steep gradient, slows down as reactants are used up, levels off to a plateau when reactants are completely used up.
83
What factors affect the rate of a chemical reaction?
* Concentration * Temperature * Catalyst use * Surface area of solid reactants
84
State collision theory.
Reacting particles must collide for a reaction to occur.
85
What conditions must be met for a collision to result in a chemical reaction?
* Particles must collide with the correct orientation * Particles must have sufficient energy to overcome the activation energy barrier.
86
How does concentration affect the rate of reaction?
Increasing concentration increases the rate due to more frequent successful collisions.
87
How does increasing pressure affect the rate?
Increased pressure increases rate as particles are closer together, leading to more frequent collisions.
88
How do you calculate rate from a graph?
By calculating the gradient of the curve.
89
How do you find the instantaneous rate?
The gradient of the tangent at a specific point in time.
90
How do you find the initial rate?
The gradient of the tangent drawn at point t=0.
91
Define catalyst.
A substance that increases the rate of a chemical reaction without being used up, providing an alternative route with lower activation energy.
92
Define homogeneous catalyst.
A catalyst in the same physical state as the reactants.
93
Define heterogeneous catalyst.
A catalyst in a different physical state from the reactants.
94
What are the benefits of catalysts?
* Reduced energy consumption * Environmental benefits * Economic importance in catalytic processes.
95
How do catalysts influence energy consumption and the environment?
They increase reaction rates, reducing the need for heat or high pressure, which may lower fossil fuel use and CO2 emissions.
96
How do catalysts influence atom economy and waste?
They may allow processes with higher atom economy, reducing waste.
97
What is the economic importance of catalytic processes?
They make products faster and use less energy, cutting costs and increasing profitability.
98
State features of a Boltzmann distribution.
* No molecules have 0 energy, curve starts at origin * Area under curve = number of molecules * No maximum energy for a molecule, asymptotic to the x-axis.
99
Explain the effect of temperature on the rate of reaction.
Higher temperatures increase the number of molecules with energy greater than or equal to activation energy, leading to more frequent successful collisions.
100
Explain why pH cannot be used as a measure of rate.
A very large change in H+ concentration is required to cause any change in pH due to the logarithmic scale.
101
When should we not monitor mass loss?
If there is a very small mass change resulting in large percentage error.
102
When should we not monitor volume of gas given off?
If the gas is water soluble.
103
Draw Boltzmann distribution to show effect of catalysts on rate of rxn.
104
State general answer structure to state which method of monitoring rate should be used.
Answer structure: Reason why using method What apparatus What property is being measured How youd expect it to change eg: “Because a precipitate is formed, use a colorimeter to measure the decrease in transmittance of light.”
105
State ways how rate of change can be monitored.
Conductivity probe if ions are produced - measure increase in conductivity Gas syringe/cylinder/balance if gas produced Colorimeter - color change or precipitate forms
106
State usual assumptions made when measuring ∆rH
Density of solution is the same as water SHC of solution is the same as SHC of water Heat capacity of equipment is negligible
107
Give equipment required for measuring ∆rH of two solutions / solid and solution

chemistry (5 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2025 Bold Learning Solutions. Terms and Conditions