Module 3: Periodic Table and Energy AS

Question 1

What is Enthalpy, H?

Answer 1

The heat content that is stored in a chemical system

Question 2

Who arranged elements into seven Octaves?

Answer 2

John Newlands

Question 3

What is the definition for the activation energy?

Answer 3

The minimum energy required to start a chemical reaction

Question 4

What were the disavantages to Mendeleev’s periodic table?

Answer 4

Isotopes had yet to be discovered, and there was no place for them in Mendeleev’s table. It also didn’t include any noble gases which hadn’t been discovered yet.

Question 5

What were the advantages to Mendeleev’s periodic table?

Answer 5

He left space for unknown elements. His work also indicated that some accepted atomic weights were incorrect , while his table provided for variance from atomic weight order

Question 6

What is the standard enthalpy change of combustion?

Answer 6

The enthalpy change that takes place when one mole of a substance reacts completely with oxygen under standard conditions.

eg: CH₄ + 2O₂ –> CO₂ + 2H₂O

Question 7

What is the standard enthalpy change of formation of a compound?

Answer 7

The enthalpy change that takes place when one mole of a compound is formed from its constituent elements in their standard states.

NB: Forming 1 mole of an element in its standard state has a value of zero by definition

Question 8

What is the definition of Periodicity?

Answer 8

The trend in properties that’s repeated across each period.

Eg, All group 7 elements are very reactive because, having 7 outer shell electrons, they only need one more electron for a full outer shell.

Question 9

What is the unit for an enthalpy change?

Answer 9

kJmol^-1

Question 10

Which orbital would fill first: 3d or 4s?

Answer 10

4s

There are various ways to remember this; eg use your periodic table and notice that after 3p⁶ (Ar), the next element is an s block element (K is [Ar]4s¹).

If you use this method, note the pattern that the first shell contains only 1 type of orbital (s, so 1s); the second shell contains 2 types of orbital (s and p, so 2s and 2p); the third shell contains 3 types of orbital (s, p and d, so 3s, 3p and 3d). Notice that shell number is not the same as row number in the periodic table.

Question 11

What is ionisation?

Answer 11

When atoms lose or gain electrons

Question 12

What is ionisation energy?

Answer 12

The energy needed to form positive ions

Question 13

What is the first ionisation energy (1st I.E)?

Answer 13

• How easily an atom loses and electron to form a 1+ ion
• The energy required to remove one electron from each atom in one mole of the gaseous element to form one mole of gaseous 1+ ions

Question 14

Sketch an energy level diagram for the orbitals from 1s to 4p

Answer 14

Question 15

Define ionisation energy

Answer 15

The energy required to remove 1 mole of electrons from 1 mole of gaseous ions.

Note: Always endothermic

1st IE: X_(g) –> X⁺_(g) + e^-

2nd IE: X⁺_(g) –> X²⁺_(g) + e^-

Question 16

Describe the shapes of s, p and d orbitals

Answer 16

Question 17

Cr and Cu do not fill the 3d orbitals as you might expect. What is their electron configuration, and why?

Answer 17

Cr: [Ar]3d⁵4s¹ (Half-filled shells more stable)

Cu: [Ar]3d¹⁰4s¹ (filled 3d and 1 electron in 4s more stable)

Question 18

How does the atomic radius affect the nuclear attraction experienced by the outer electrons?

Answer 18

The larger the atomic radius, the smaller the nuclear attraction experienced by the outer electrons

Question 19

Describe how we fill electron orbitals.

Answer 19

This is the “Aufbau” (or Building up) principle. We start from the lowest energy orbital. Each sub-orbital can hold 2 electrons. We keep putting electrons into the orbitals from the bottom upwards.

s orbitals have only 1 sub orbital, so hold 2 electrons in total (and this explains why the s block is 2 elements wide)

p orbitals have 3 sub orbitals, so hold 6 electrons in total (and this explains why the p bolck is 6 elements wide)

d orbitals have 5 sub orbitals, so hold 10 electrons in total (and this explains why the d bolck is 10 elements wide)

Question 20

How does nuclear charge affect the attractive force on the outer electrons?

Answer 20

The higher the nuclear charge, the larger the attractive force on the outer electrons

This explains, within a period, the atoms of group 8 elements being smaller than the atoms of group 1 elements

Question 21

What is the sign of DELTA H for an exothermic reaction?

Answer 21

Negative,

Because the enthalpy of the products is smaller than the enthalpy of the reactants.

Question 22

What is electron shielding/screening?

Answer 22

When the inner shells of electrons block charge from the nucleus being experienced by the outer-shell electrons

Question 23

Do more inner shells in an atom create a larger or smaller shielding effect?

Answer 23

Larger

(More shells of electrons blocking nuclear charge)

Question 24

How does the shielding effect change the nuclear attraction experienced by outer electrons?

Answer 24

The larger the shielding effect, the smaller the nuclear attraction experienced by outer electrons

Question 25

Is the value for DELTA H in an endothermic reaction positive or negative.

Answer 25

Positive, The enthaply of the products is greater than the enthalpy of the reactants.

Question 26

What are successive ionisation energies?

Answer 26

They are a measure of the amount of energy requires to remove each electron in turn (For example, the second ionisation energy is how easily a 1+ ion loses an electron to form a 2+ ion)

Question 27

What factors effect nuclear attraction experienced by an electron?

Answer 27

Atomic radius Nuclear charge Electron shielding

Question 28

Why is each successive ionisation energy higher than the one before?

Answer 28

* As one electron is removed, there is less repulsion between the remaining electrons, meaning each shell is drawn slightly closer to the nucleus * The positive nuclear charge outweighs the negative charge every time an electron is removed * As the distance between the electrons and the nucleus **decreases**, the nuclear attraction **increases** * More energy is needed to remove the electrons

Question 29

How do you find out the value of *K*_c?

Answer 29

Products ÷ Reactants eg: A + 2B ⇔ 2C *K*_c = [C]² / [A][B]²

Question 30

What is Calorimetry?

Answer 30

the quantitative study of energy in a chemical reaction This is typically achieved by conducting an experiment in such a way that a substance with known mass and known specific heat capacity is warmed or cooled by the reaction. The change in thermal energy can be found using Q = mc(DELTA)T

Question 31

What is the trend in ionisation energies across a period?

Answer 31

Increase, due to: Increasing number of protons creating higher attraction Electrons being added to the same shell, drawing it inwards Same number of inner shells so shielding will hardly change

Question 32

what is meant by the term 'specific heat capacity'?

Answer 32

the energy required to raise the temperature of 1g of a substance by 1K

Question 33

what is the biggest experimental error and how can you reduce it?

Answer 33

**_Heat loss_** - insulating the equipment with draught shields

Question 34

what is a **_Bomb Calorimeter_**?

Answer 34

a sophisticated piece of equipment which minimises heat loss as much as possible, it uses pure oxygen to ensure complete combustion is acheived.

Question 35

Does the first ionisation energy of an atom increase or decrease across a period?

Answer 35

Increases

Question 36

State the trend in solubility of group 2 metal hydroxides

Answer 36

Solubility increases down the group: Be(OH)2 is insoluble Mg(OH)2 is slightly soluble Ba(OH)2 is the most soluble

Question 37

State the trend in alkalinity of Group 2 hydroxides

Answer 37

Alkalinity increases down the group: Be(OH)2 is insoluble, so zero concentration of OH- ions Mg(OH)2 is slightly soluble, very low concentration of OH- ions Ba(OH)2 is the most soluble, so highest concentration of OH- ions

Question 38

State the physical states (at STP) and colours of the halogens

Answer 38

Flourine: gas, (very pale green/yellow) Chlorine: gas, pale green Bromine: liquid, orange Iodine: solid, dark purple NOTE: Bromine liquid will give off brown fumes; Iodine solid sublimes, forming a purple vapour.

Question 39

State the colours of chlorine, bromine and iodine: * In water * In cyclohexane

Answer 39

Chlorine: Pale green in water; pale green in cyclohexane Bromine: Orange in water; orange in cyclohexane Iodine: Brown/orange in water; violet in cyclohexane

Question 40

Give an ionic equation for the reaction of silver ions with any halide ion (X^- = halide)

Answer 40

Ag⁺_(aq) + X^-_(aq) --\> AgX_(s)

Question 41

Give the colours of the silver halides

Answer 41

Silver chloride: White Silver bromide: Cream Silver iodide: yellow

Question 42

How is ammonia used to identify silver halides?

Answer 42

Silver chloride: Dissolves in **DILUTE** ammonia Silver bromide: Dissolves in **CONCENTRATED** ammonia only Silver iodide: **INSOLUBLE** in ammonia

Question 43

State the test for Ammnium ions

Answer 43

Ammonium ions react with hydroxide (from sodium hydroxide solution) to produce ammonia and water: NH₄⁺_(aq) + OH^-_(aq) --\> NH_3(g) + H₂O_(l) Add a small amount of ammonium containing salt to sodium hydroxide solution in a test tube and warm gently. Balance a piece of moist, pink litmus on the mouth of the test tube. If ammonium ions were present, the ammonia formed turns the pink litmus blue.

Question 44

To test for carbonate ions...

Answer 44

Add a sample of the carbonate to a soluion of dilute acid. The effervescing suggests CO₂, which can be confirmed by bubbling the gas through limewater. If the limewater goes cloudy, the gas evolved was CO₂, formed by the reaction below: CO₃^2-_(s) + 2H⁺_(aq) --\> H₂O_(l) + CO_2(g)

Question 45

To test for sulfate ions...

Answer 45

Barium sulfate is insoluble. If you have a solution which you suspect includes sulfate, combing a few drops of this with a few drops of a solution containing barium ions will produce a white precipitate of barium sulfate: Ba²⁺_(aq) + SO₄^2-_(aq) --\> BaSO_4(s)

Question 46

Give 2 uses of group 2 metal compounds

Answer 46

The 2 uses you need to know arise from the basicity of group 2 compounds (Remember, Group 2 is also called the "Alkaline Earth" metals) * Calcium hydroxide is added to fields by farmers to increase the pH of acidic soils * Magnesium hydroxide and calcium carbonate are used to neutralise excess stomach acid (Magnesium hydroxide is often used as a suspension, "milk of magnesia" and calcium carbonate is often in tablets as both compounds are very poorly soluble in water)

Question 47

Give uses of chlorine

Answer 47

Chlorine is used to kill bacteria in drinking water Risks from using chlorine are that * The gas is toxic (so used at low concentrations) * Chlorine can react with hydrocarbons formed by rotting vegetation. The compounds made might be carcinogenic. However, the risk from this to human health is lower than the risk of disease contracted through use of untreated drinking water.

Question 48

Define average bond enthalpy

Answer 48

The energy to break one mole of a specified type of bond in a gaseous molecule: * Energy is always required to break bonds * Bond enthalpies are always endothermic * Bond enthalpies always have positive values * The actual bond enthalpy can vary depending on the chemical environment of the bond. The average is calculated using actual values but in different environments.

Question 49

Define Hess' Law

Answer 49