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chemistry > C7 - Periodic table and energy > Flashcards

C7 - Periodic table and energy Flashcards

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1
Q

Who was responsible for the periodic table

A

Dmitri Mendeleev

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2
Q

How did Mendeleev organise the periodic table?

A

-In order of increasing atomic mass
-Grouped according to similar physical and chemical properties
-Left gaps for undiscovered elements
-Switched tellurium and iodine

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3
Q

Why didn’t Mendeleev accept noble gasses were elements?

A

They did not react

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4
Q

How is the periodic table arranged now

A

-Increasing atomic number
-7 horizontal periods
-18 vertical groups

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5
Q

What are groups?

A

Vertical column on periodic table
Elements with similar chemical properties and same number of outer-shell electrons

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6
Q

What are periods?

A

Horizontal rows in periodic table
Show trend in (physical) properties across a period

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7
Q

What’s periodicity?

A

Repeating trend in properties of element across each period of the periodic table

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8
Q

Periodic trend in electron configuration in period 2 and 3

A

-Period 2: 2s sub-shell fills with 2e-, 2p sub-shell fills with 6e-
-Period 3: 3s sub-shell fills with 2e-, 3p sub-shell fills with 6e-

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9
Q

What are blocks

A

Correspond to highest energy subshell
S block
P block
D block
F block

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10
Q

How are blocks arranged

A
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11
Q

What does ionisation energy measure?

A

How easily an atom loses electrons to form positive ions

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12
Q

What’s first ionisation energy?

A

Energy required to remove one electron from each atom in one mole of gaseous atoms of an element of form one mole of gaseous 1+ ions

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13
Q

First ionisation formula of Na

A

Na (g) —> Na+ (g) + e-

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14
Q

What factors affect ionisation energy?

A

Atomic radius
Nuclear charge
Electron shielding

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15
Q

How does atomic radius affect ionisation energy?

A

Greater distance between nucleus and out electrons
Less nuclear attraction

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16
Q

How does nuclear charge affect ionisation energy?

A

More protons
Greater attraction between nucleus and outer electrons

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17
Q

How does electron shielding affect ionisation energy?

A

Electrons are negatively charged
Inner shell electrons repel outer shell electrons
Reduce attraction between nucleus and outer electrons

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18
Q

What is shielding?

A

Repulsion between electrons in different inner shells
Reduces net attractive forces between nucleus and outer shell electrons

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19
Q

What is second ionisation energy?

A

Energy required to remove one electron form each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions

20
Q

Second ionisation energy formula of He

A

He+ (g) –> He2+ (g) + e-

21
Q

What can successive ionisation energies tell you?

A

-Large difference in ionisation energy suggests e- is in a different shell
-number of e- in outer shell

22
Q

What happens to first ionisation energy down a group?

A

Decreases

23
Q

Why does first ionisation energy decrease down a group?

A

-Atomic radius decreases
-More shells so shielding increases
-Nuclear attraction decreases
-First ionisation energy decreases

24
Q

What happens to first ionisation energy across a period?

A

Increases

25
Q

Why does first ionisation energy increase across a period?

A

-Nuclear charge increases
-Atomic radius decreases
-Nuclear attraction increases
-First ionisation energy increases

26
Q

Where does first ionisation energy fall in period 2?

A

-Beryllium to boron
-Nitrogen to oxygen

27
Q

Compare Be and B

A

-2p sub shell in in B has higher energy that 2s sub shell in Be
-Therefore, e- in B is easier to remove
-First ionisation of B is less

28
Q

Compare N and O

A

-Highest energy e- are in 2p sub shell in N and O
-Paired e- in 2p orbitals repel, making it easier to remove, in O
-Therefore, first ionisation of O is less than N

29
Q

At room temperature what state are most metals?

A

Solid

30
Q

What metal is not solid at room temperature?

A

Mercury - liquid

31
Q

What’s a metallic bond?

A

Electrostatic attraction between positive metal ions and delocalised electrons

32
Q

What is the structure and bonding of metals?

A

Giant metallic structure consisting of cations and delocalised electrons held together by metalling bonds

33
Q

Properties of metals

A

-Strong metallic bonds
-High electrical conductivity
-High melting and boiling points

34
Q

Explain the conductivity of metals

A

When voltage applied, delocalised electrons can move and carry charge

35
Q

Explain the high melting and boiling points of metals

A

To overcome strong electrostatic force of attraction between cations and delocalised electrons

36
Q

Are metals soluble

A

Insoluble

37
Q

What is a giant metallic lattice?

A

3D structure of positive ions and delocalised electrons, bonded together by metallic bonds

38
Q

What is a covalent bond?

A

Strong electrostatic attraction between a shared pair of electrons and nuclei of bonded atoms

39
Q

What is a giant covalent lattice?

A

3D structure of atoms bonded together by strong covalent bonds

40
Q

Properties of giant covalent structures

A

-High melting and boiling points
-Insoluble

41
Q

Explain the solubility of giant covalent structures

A

Covalent bonds are too strong to be broken by interactions with solvents

42
Q

Explain the melting and boiling point of giant covalent structures

A

Strong covalent bonds require high energy to overcome

43
Q

Explain the electrical conductivity of giant covalent structures

A

Don’t conduct electricity - all 4 outer shell electrons used in covalent bonding

44
Q

Which giant covalent structure do conduct electricity

A

Graphene
Graphite

45
Q

Periodic trend in melting point across period 2 and 3

A

Increases across groups 1-14
Sharp decrease from 14-15
Comparatively low from groups 15-18

46
Q

What does the sharp decrease in boiling points between group 14 and 15 indicate?

A

Change from giant structures to simple molecular structure

chemistry (21 decks)

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