Unit 3. Periodic Table

Question 1

What is the nuclear charge

Answer 1

The number of protons in the nucleus of an atom. Basically the atomic number

Question 2

Name of vertical columns and horizontal row

Answer 2

Group and period

Question 3

What does the period number tell you?

Answer 3

The number of occupied main energy levels in an atom

Question 4

What does the group number tell you?

Answer 4

The number of valence electrons in an atom of that element

Question 5

Where are metals, non-metals and metalloids located on the periodic table?

Answer 5

Metals located mainly on the left and middle in the s-, f- and d-block.
Non-metal elements are found on the right side in the p-block.
Metalloids (semi-metals) are located in a diagonal staircase that forms a boundary between metals and non metals.

Question 6

Name the name of the groups in the periodic table and where they are found, as well as their characteristic properties

Answer 6

Alkali Metals- Group 1- most reactive, react strongly with water
Alkaline earth metals- Group 2- reactive but less reactive then group 1
Halogens- Group 17- most reactive group of non-metals
Noble gases- Group 18- very unreactive monatomic gases
Transition metals- D-Block region (apart from zinc)- relatively stable with only moderate reactivity useful for construction

Question 7

What are the lanthanides and actinides

Answer 7

They are metallic elements that make up the f-block. Their two rows are separated from the rest of the periodic table. Atomic numbers from 89-103, they are radioactive. Lanthanides (57-71) are knows as rare earth metals

Question 8

What are the s- p- d- and f- block directly correlated to?

Answer 8

The electron configuration with the sub levels: s,p,d,f. The block in which an element is tell us which sub-level is in the process of being filled.

Question 9

How do you find the v.e of group 13 to 18?

Answer 9

Substract ten from group number.

Question 10

What is atomic radii

Answer 10

The size of an atom

Question 11

Trend in atomic radii?

Answer 11

Period
Starting with sodium (11) it decreases across the period until it reaches chlorine (17). Bc we gain an electron and a proton as we go across a period, electron shielding remains more or less constant, as the electrons fill the same main energy level. Therefore there is an increased attraction between the nucleus and outer electrons which pulls them closer.
Group
Going down a group, each element’s v.e occupies another extra main energy level which increases the atomic radii.

Question 12

What is electron shielding?

Answer 12

Electron shielding occurs when outer electrons are shielded from the attraction of the nucleus by inner electrons knows as shielding electrons.

Question 13

Which are more likely to become anions and cations, metals or non metals?

Answer 13

Metallic elements tend to loose electrons and form cations.
Non metallic elements tend to gain electrons and form anions

Question 14

What are isoelectronic species?

Answer 14

Isoelectronic species have the same electron configuration. Example: Na+, Mg2+. They have different numbers of proton but same of electrons.

Question 15

What is the ionic radius

Answer 15

The distance between the nucleus of an ion and the outermost energy shell of the ion

Question 16

When does the ionic radius increase?

Answer 16

When the number of proton decrease
More positively charged cations have the smallest ionic radius, and more negatively charged anions have the largest radius.

Question 17

Explain ionic radius trend in period 3

Answer 17

• first 4 cations have lost their outer energy levels of elections and now occupy one less energy level. Since the nucleus is pulling on fewer electrons, the attraction increases between the nucleus and the v.e, which reduces the ionic radius.
• anions are bigger than their parent atom as they gain electrons which reduces the attraction by the nucleus. The extra electrons also increase repulsion between electrons. This all makes the ionic radii increase.
• across a period the ionic radii decreases as the nuclear charge increases (protons higher)

Question 18

What is effective nuclear charge

Answer 18

The attraction felt by v.e which are further away from the nucleus and are shielded by other electrons.

Question 19

How do you calculate the effective nuclear charge?

Answer 19

Take sodium (11). The 11th electron in the last 3s sub-level is shielded by the other 10 electrons, which are closer to the nucleus, from the 11 protons. Therefore do 11 protons minus 10 electrons. That v.e effective nuclear charge is 1+.

Question 20

Nuclear charge trend

Answer 20

Period
- across increases as proton number increases but electron number in inner energy levels remains the same. It increases by one.
Group
- remains constant , as the attraction from the increasing numbers of protons in the nucleus is offset by the increase in occupied energy levels down the group.

Question 21

What is first Ionisation energy?

Answer 21

Ionisation Energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.

(A measure of the attraction between the positively charged nucleus and the negatively charged valence elections.)

Question 22

General trend for first Ionisation energy

Answer 22

Increase across period bc of increase in nuclear charge and decrease in atomic radii, which results in an increased attraction between the nucleus and the v.e

Decrease down a group is bc of the increase in atomic radius which causes weaker attraction between the nucleus and v.e of an atom.

Question 23

Exceptions to first ionization energy

Answer 23

Decrease in ionization energy from beryllium to boron, magnesium to aluminum, nitrogen to oxygen and phosphorus to sulfur.

Beryllium has the electron configuration 1s2 2s2

Boron has the electron configuration 1s2 2s2 2p1

Electrons in p orbitals are of higher energy and further from the nucleus than electrons in s orbitals, therefore they require less energy to remove. The same explanation can be applied for the drop in ionisation energy from magnesium to aluminium

Nitrogen has the electron configuration 1s2 2s2 2p3

Oxygen has the electron configuration 1s2 2s2 2p4

In oxygen, the electron is removed from a doubly occupied p-orbital. An electron in a doubly occupied orbital is repelled by the other electron and requires less energy to remove than an electron in a half-filled orbital. The same explanation applies for the decrease in ionisation energy from phosphorus to sulfur, except that the elect

Question 24

What is electronegativity?

Answer 24

Electronegativity is the attraction of an atom for a bonding pair of electrons.

Question 25

Trend of electronegativity

Answer 25

Increases across a period because of the increase in nuclear charge and the decrease in atomic radius. Down a group electronegativity decreases bc of increase in atomic radius. The increase in nuclear charge down a group is counteracted by the increase in shielding caused by extra energy levels.

Question 26

What is first electron affinity

Answer 26

The first electron affinity is the energy released when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous 1- ions.

Question 27

Are electron affinities exothermic or endothermic?

Answer 27

First Ionisation energies are exothermic bc energy is released when an electron is added to a neutral atom. Second electron affinities are endothermic bc the repulsion that occurs when a second electron is added to an already negative ion.

Question 28

Which elements have the highest electron affinities?

Answer 28

Group 17, they are the most exothermic as they are relatively small atoms that can accommodate an electron in their unfilled outer energy level.

Question 29

Trend for electron affinity

Answer 29

Period Across a period electron affinity increases, as we progress towards the filling of the outer valence shell of the atom. Group 17 releases more energy when aiming an electron as it achieves a filled valence shell and is more stable. Group Down the group the electron affinity decreases, as the additional electron gained is entering a new energy level further from the nucleus. This added electron has a weaker attraction to the nucleus and therefore releases less energy when added.

Question 30

How are metals characterized?

Answer 30

- having fewer v.e - larger atomic radii - lower electronegativity and lower Ionisation energies - tendency to loose v.e easily

Question 31

How are non metallics characterized

Answer 31

- having more v.e - smaller atomic radii - higher electronegativity and Ionisation energy - tend to gain electrons.

Question 32

Trend of metallic character

Answer 32

Decreases across period Increases down a group

Question 33

Trend in melting point in period 3

Answer 33

Across a period the structure and bonding gradually changes from metallic to giant covalent to molecular covalent. Sodium, mg and aluminum are metallic elements. the strength of the metallic bonding increases from sodium to aluminum, consequently so does the melting point. The melting point reaches a peak with silicone (right after Aluminum), due to its giant covalent structure. The elements to the right of silicone are non metallic, and although the covalent bonds within each molecule are strong, there are only relatively weak intermolecular forces between the molecules, so they have lower melting points.

Question 34

What does the melting point rely on?

Answer 34

Intermolecular forces and structure

Question 35

Trend of bonding and structure in period 3 oxides.

Answer 35

Period 3 oxides show a gradual decrease in metallic character across the period. Metal oxides have giant ionic structures, which are sold under standard conditions due to the strong electrostatic attractions between the ions. Non metals oxides are molecular compounds that exist as individual molecules. Bc of the weak intermolecular forces between their molecules, they are usually gas or liquids under standard conditions. They are giant covalent. Na2O is ionic bc of the difference in electronegativity But Cl2O is covalent bc both element are very electronegative an therefore the difference is less.

Question 36

Acid base trend in period 3 oxides

Answer 36

All oxides before aluminum are basic. Al is amphoteric. Anything after Al is acidic. Note: when term basic, amphoteric and acidic used, it means that they form basic, amphoteric and acidic solutions when reacted with water.

Question 37

Trends of alkali metals (group 1)

Answer 37

- they are very soft - first three elements have low densities that they can float on water - melting points and boiling point are quite low down the group due to the metallic bonding getting weaker as the ionic radii of the metals increase - they are stored in oil to prevent them from reacting with the air - they easily loose their v.e due to low first ionization energy - they undergo vigorous reactions with water to from metal hydroxide and hydrogen gas, which forms a solution with a high ph. - reactivity increases down group

Question 38

Trend in the halogens (group 17)

Answer 38

- very reactive group of non-metals - they have characteristic colors. Cl= dense pale green, smelly and poisonous. Br= red/brown liquid, smelly and poisonous. Iodine= grey solid with purpule vapor, smelly and poisonous - physical state changes down the group, with chlorine being a gas to bromine a liquid. This is related to the increasing mole mass which results in stronger intermolecular forces between molecules. - halogens are diatomic - melting/boiling points increase down the group, due to increase IMF - reactivity decreases down group - react with group 1 to form salts.

Question 39

What is a displacement reaction?

Answer 39

A reaction in which the more reactive halogen displaces ions of the less reactive halogen from solution. They are accompanied by color changes