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Chemistry > U3.1: Periodicity > Flashcards

U3.1: Periodicity Flashcards

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1
Q

Properties of metals

A
  • act as conductor of heat and electricity
  • forms cations by losing electrons
  • typically lustrous
  • Most of the elements are metals
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2
Q

Properties of non-metals

A
  • Doesn’t act as a conductor of anything
  • Forms anions by gaining electrons
  • Typically dull in solid state but most are non-metals
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3
Q

Property of metalloids

A
  • Semi metal
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4
Q

Properties of groups

A
  • Elements in same group have similar chemical properties and show a trend in physical properties
  • This is because they have the same number of valence electrons
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5
Q

Properties of periods

A
  • Elements in the same period have the same number of electron shells
  • This is because the number of electron shells affect the size of the element
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6
Q

List the metalloids

A

B, Si, Ge, As, Sb, Te, Po

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7
Q

Define effective nuclear charge

A

Net positive charge from the nucleus experienced by outer valence shell electrons once the screening effect of the core electrons has been considered

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8
Q

Define Periodicity

A

Refers to repeating trends or patterns of physical & chemical properties in elements

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9
Q

Define atomic radius

A

Distance from an atom’s nucleus to the outermost orbital of 1 of its electrons

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10
Q

Define isoelectric

A

Molecules that have the same electronic configuration

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11
Q

What is the trend in the atomic radius

A

Across a period:
- decreases as electrons are added to the same main energy level (n = 3)
- effective nuclear charge increases
- attraction between nucleus and valence electrons increases

Down a group:
- increases as number of electron shells increases

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12
Q

What is the trend in the ionic radius

A

POSITIVE IONS:
Across a period:
- decreases as the number of valence electrons decreases
- attraction between nucleus and valence electrons increases
- greater the number of electrons removed, the smaller the ionic radius

NEGATIVE IONS:
Across a period:
- increases as number of valence electrons increases
- attraction betwee the nucleus and valence electrons decreases
- greater the number of electrons gained, the larger the ionic radius

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13
Q

Define Ionisation energy

A

The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions

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14
Q

Define first ionisation energy

A

Energy required to remove one mole of the most loosely held electron from one mole of neutral gaseous atoms to form one mole of gaseous 1+ ions

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15
Q

What is the trend in ionisation energy

A

Across a period:
- increases as the number of protons increases in the nucleus and there are more electrons occupying the outer valence shell
- there’s an increase in effective nuclear charge
- there’s a greater number of electrons in the outer valence shell hence, there’s a stronger attraction to the nucleus
- More energy is required to remove an electron

Down a group:
- decreases as the atomic radius increases and electrons occupy higher energy levels
- Electrons in higher energy levels have a weaker attraction to the nucleus-
- Less energy is required to remove an electron

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16
Q

Define Electron Affinity

A

The amount of energy released when a neutral atom gains an electron to form an anion

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17
Q

What is the trend in electron affinity

A

Across a period:
- increases as there is an increasing effective nuclear charge
- stronger attraction between added electrons and nucleus

Down a group:
- decreases as elements further down have a larger atomic radius, with outer valence shells further away from the nucleus
- weaker attraction between the added electrons and nucleus

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18
Q

What are the excepts to the trend of electron affinity

A
  • F and Cl
  • Since they’re so small, it’s hard to add more electrons to its much higher electron density
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19
Q

How high is the electron affinity in metals and non metals

A
  • Metals have a low EA
  • Non-metals have a high EA
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20
Q

Define electronegativity

A

the tendency of an atom to attract electrons in a chemical bond

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21
Q

What is the trend in electronegativity

A

Across a period:
- increases as effective nuclear charge increases
- stronger attraction between nucleus of atom and shared pair of electrons

Down a group:
- decreases as atomic radius increases from additional energy levels
- Weaker attraction between nucleus of an atom and shared pair of electrons

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22
Q

How high is the electronegativity in metals and non-metals, why?

A
  • metals have low EN as they lose electrons easily
  • non-metals have high EN as they gain electrons to complete their outer valence shell
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23
Q

What is one exception to conducting electricity whilst being a non-metal

A

Graphite

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24
Q

Define metallic character

A

Tendency of an element to lose electrons to form cations

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25
Q

Define non-metallic character

A

Tendency of an element to gain electrons to form anions

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26
Q

Trend in metallic character?

A

From bottom left to top right of periodic table, MC increases

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27
Q

Trend in non-metallic character?

A

From bottom left to top right of periodic table,
NMC increases

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28
Q

All Chemical trends of Group 1 (Alkali Metals) Down the group

A
  • Atomic/ionic radius increases: Increasing e- shells
  • EN and1st IE decreases: Increased distance between V. e- and nucleus hence, easier to remove
  • MP decreases: Atomic radius increases hence, metallic bonds become weaker
  • Reactivity increases: V. e’s easier to lose due to greater shielding of inner e-
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29
Q

All physical properties of Group 1 (Alkali Metals) Down the group

A
  • Density increases
  • Softness increases
  • MP increases
  • Good conductors
  • Lustrous
30
Q

What types of reactions can Group 1 (Alkali Metals) undergo

A
  • Metals reacting with O2 (tarnishing not rusting)
  • Metals reacting with water to form OH- (a base) & H2
  • Metals reacting with halides
31
Q

Qualitative data of Lithium’s reaction with water

A
  • Solid floats
  • Solid moves slowly across surface
  • Gentle fizzing
  • Solution turns pink (given that U. Indicator’s alrd been put in)
32
Q

Qualitative data of Sodium’s reaction with water

A
  • Solid melts
  • Solid creates a molten ball and moves around quickly on surface
  • Fizzing
  • Solution turns pink (given that U. Indicator’s alrd been put in)
33
Q

Qualitative data of Potassium’s reaction with water

A
  • Solid melts
  • Solid creates a molten ball which quickly ignites to produce a lilac flame
  • Vigorous fizzing
  • Solution turns pink (given that U. Indicator’s alrd been put in)
34
Q

All Chemical trends of Group 17 (Halogens) Down the group

A
  • Atomic/ionic radius increases: Increasing e- shells
  • 1st IE decreases: V. e-‘s further away from nucleus hence, easier to remove
  • EN decreases: Increased distance between V. e- and nucleus hence, easier to remove
  • MP increases : VDW Forces become greater with more e-
  • Reactivity decreases : V. shell gets further away from nucleus hence, attraction between V. e- & nucleus gets weaker therefore e- are less easily gained
35
Q

All physical Properties of Group 17 (Halogens) Down the group

A
  • Volatility decreases
  • Solubility in H2O decreases
  • MP increases
  • Dullness of colour increases
36
Q

What are the colours observed for the first 4 Halogens (F, Cl, Br, I) in elemental state (ES) and aqueous state (AS)

A

F: Pale yellow gas (ES) & Yellow (AS)
Cl: Yellow-green gas (ES) & Pale yellow-green (AS)
Br: Reddish-brown liquid (ES) & Orange (AS)
I: Grey solid with purple vapours when it sublimes (ES) & Brown (AS)

37
Q

What do displacement reactions of a halogen and a halide solution depend on and give an example and explanation.

A

Electron Affinity of the element (which then affects reactivity)

For example: Cl2 has a higher EA than Br2 hence, Cl2 displaces Br in KBr forming KCl and Br2

38
Q

Define amphoteric

A

Can act as both an acid and a base

39
Q

What is the acid-base character, pH and structure of period 3 oxides Na2O, MgO, Al2O3, SiO2, P4O10, SO2 / SO3, Cl2O

A

Acid Base Character:
- Basic: Na2O, MgO
- Amphoteric: Al2O3
- Acidic: Si2O, P4O10, SO2/SO3, Cl2O

pH:
- High (>7): Na2O, MgO
- Amphoteric: (high in powder [>7], low in water [<7])
- Low (<7): Si2O, P4O10, SO2/SO3, Cl2O

Structure:
- Ionic Lattice: Na2O, MgO
- Covalent Network Lattice: SiO2
- Small Cage-Like Structure: P4O10
- Covalent molecule: SO2/SO3, Cl2O

40
Q

Structure of Period 3 elements

A

Metallic:
- Na, Mg, Al

Covalent network lattice (Metalloid):
- Silicon

Covalent Molecule:
- P, S, Cl

41
Q

Basic oxides, of group 1 and 2 metals, reactions with water (Na2O and MgO)

A
  • Na2O(s) + H2O(l) -> 2NaOH(aq)
  • MgO(s) + H2O(l) -> Mg(OH)2(aq)
42
Q

Acidic oxides, of group 1 and 2 metals, reactions with water (CO2, SO2/SO3)

A
  • CO2(g) + H2O(l) -> H2CO3(aq)
  • SO2(g) + H2O(l) -> H2SO3(aq)
  • SO3(g) + H2O (l) -> H2SO4(aq)
43
Q

Implications of CO2, SO2/SO3 and NO2?

A

Ocean acidification: CO2
Acid Rain: SO2/SO3 & NO2

44
Q

2 major global concerns regarding acid rain and ocean acidification?

A

Acid Rain:
- Damages infrastructure
- Impacts crop yield

Ocean Acidification:
- Kills marine life in small bodies of water like lakes due to accumulation of acid hence, body of water becomes more acidic which affects marine food chains

45
Q

Define Oxidation state

A
  • The degree of oxidation an atom in a compound experiences
46
Q

What is oxidation

A
  • Loss(v) of H2 or electrons
  • Gain(v) of O2
  • Increase (v) in oxidation state (becomes more (v) positive

(v) = vice versa for reduction

47
Q

Oxidation species are often on which side of the half equation

A

Left side

48
Q

Main rules for oxidation states?

A
  • Ox. charge is +1 not 1+ (that’s ionic charge)
  • Element or compound by itself has overall ox. state 0
  • Groups 1/2/3 always are +1/+2/+3
  • Monoatomic ions ox. state is their ion charge
  • D-block transition metals have variable ox. states
  • Halogens have -1ox. state (+1 with O2)
  • Hydrogen has +1 ox. state (-1 with metals)
  • O2 has -2 ox. state (-1 in peroxide)
  • Fluorine has -1 ox. state
49
Q

Which oxidation state is commonly shown by all period 4 transition elements, with the exception of scandium?

A
  • Oxidation state shown by all period 4 transition elements is +2
  • Scandium has +3 Ox. state
50
Q

Rule linking roman numbers and ox. states? Give example of PbO2.

A
  • Ox. state of first element is often included in the name of the species
  • PbO2 is Lead(IV) oxide
51
Q

What is the reason for discontinuity in IE in groups 2 and 13 between B and Be

A

B: 1s2,2s2,2p1
Be: 1s2, 2s2

  • Be has higher IE than B due to the paired 2s2 electrons being closer to the nucleus
  • B’s inner electrons shield the 2p1 e- from the nucleus
    Therefore, this makes the outer valence electron in B easier to remove
52
Q

What is the reason for discontinuity in IE in groups 15 and 16 between N and O

A
  • N is more stable as its shells are half-filled
    Therefore greater IE is required to remove outer valence electron
  • O is less stable as the paired electrons in 2p orbital occupy the same region of space
    Therefore there’s increased repulsion between the electrons and less IE is required to remove an outer valence electron
53
Q

Why will the 2nd IE always be greater than the 1st

A
  • As the electrons are closer to the nucleus, more energy is required to overcome the stronger forces of attraction to remove an electron from a cation
54
Q

Define Successive Ionisation:

A

The process of removing successive electrons from an atom or positive ion

55
Q

Why does the IE decrease down the group

A
  • Atomic radius increases
  • Further distance between outer valence electrons and nucleus
  • Greater shielding by inner electrons
  • Less energy is required to remove the outer valence electrons
56
Q

Why does the IE increase across a period

A
  • Atomic radius decreases
  • Decreased distance between outer valence electrons and nucleus
  • Increasing nuclear charge
  • Each element across period 3 has the same number of shells
  • More energy is required to remove the outer valence electrons
57
Q

What are the 6 characteristics of transition metals

A
  • Varying Ox. States
  • High melting points
  • Magnetic Properties
  • Catalytic Properties
  • Formation of coloured compounds
  • Formation of complex ions
58
Q

Define the term ‘Ligand’

A

A species that donates a lone pair of e- to form a coordination bond with a transition metal ion.

59
Q

State the requirement of a ligand

A
  • Must have 1 lone pair of e- on at least 1 atom in a molecule
60
Q

How to find out the coordination number?

A

Coordination number is the total number of ligands

61
Q

What are the 3 types of ligands and explain what they mean

A
  • (Unidentate) Monodentate: having only 1 lone pair
  • Bidentate: having 2 lone pairs
  • (Multidentate) Polydentate: Ligands having more than 1 coordination bond
62
Q

What are transition metals

A
  • Elements whose atoms have incomplete d orbitals
63
Q

Why can each element form a number of ions with different oxidation states

A
  • e- in 4s subshell have less energy than 3d subshell
  • Ions are formed as the e- are lost from the 4s and 3d subshells
  • e- are removed from the 4s shell before the 3d subshell thus resulting in a number of ions with different ox. states
64
Q

How can transition elements form ions with an ox. state of +2?

A

Elements lose 2 e- from the 4s subshell thus, a +2 ox. state of transition elements is shown when s-electrons are removed

65
Q

What does the close proximity of 4s and 3d sublevels contribute to

A
  • The ability of e- from both sublevels to act as valence electrons in transition elements as they’re very close to each other.
  • Thus, this means that e- can be lost easily without much energy
66
Q

Equation of Na2O and P4O10 with water. Differentiate them.

A
  • Na2O(s) + H2O (l) –> 2NaOH(s). The solid formed is a base
  • P4O10(s) + 6H2O(l) –> 4H3PO4(aq). The solution formed is an acid.
67
Q

Another name for ligands and why are they known as ligands

A
  • Lewis base is another name for a ligand
  • They donate a pair of electrons
68
Q

State the substances that are involved in forming acid rain. Provide the chemical equations too

A
  • SO2: S + O2 –> SO2
  • SO3: 2SO2 + O2 –> 2SO3
  • NO and NO2: N2 + O2 –> 2NO (at very high temperatures)
69
Q

State the problems and reductions of the substances involved in making acid rain

A

SO2: Acid Rain - reduce by removing Sulphur from fuel before burning
SO3: Acid Rain - reduce by removing Sulphur from fuel before burning
NO & NO2: Acid Rain and Respiratory Problems - use a catalytic converter for engines

70
Q

Outline the conditions causing ocean acidification and what it affects as well as the pH of the water for each condition

A
  • More CO2: Lowers the amount of Carbonate (CO3-) ions thus, more acidic conditions and leads to less corals
  • Less CO2: Rises amount of Carbonate (CO3-) ions thus, less acidic conditions and leads to more corals
71
Q

Define what a chelate is

A

A ligand that can form 2 or more coordinate bonds with a transition metal

72
Q

Explain why the (transition metal complex) is coloured

A
  • It has a partially filled d-orbital
  • electrons can get promoted from a lower energy level to a higher energy level within the d-orbitals where the colour’s observed from
  • thus, the energy is absorbed from the VL region of the spectrum resulting in complimentary colours are observed
  • It’s the ligand field splitting that causes greater differences in energy levels where the different energy levels corresponds to the wavelength of light absorbed

Chemistry (18 decks)

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