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iB Chemistry > Topic 3 : Periodicity > Flashcards

Topic 3 : Periodicity Flashcards

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1
Q

Periodic Table

A
  1. The periodic table is a list of all known elements arranged in order of increasing atomic number, from 1 to 118.
  2. In addition to this, the elements are arranged in such a way that atoms with the same number of shells are placed together, and atoms with similar electronic configurations in the outer shell are also placed together.
  3. The elements are arranged in rows and columns.
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2
Q

Period

A

A row of elements thus arranged is called a period. The period number, n, is the outer energy level that is occupied by electrons.

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3
Q

Group

A
  1. A column of elements thus arranged is called a group
  2. The elements in each group share a similar outer-shell electronic configuration
  3. The outer electrons are known as the valence electrons.
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4
Q

Periodic Table - 4 Blocks

A
  1. All elements belong to one of four main blocks: the s-block, the p-block, the d-block and the f-block
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5
Q

Periodic Trends

A
  1. The physical and chemical properties of elements in the periodic table show clear patterns related to the position of each element in the table
  2. Elements in the same group show similar properties, and properties change gradually as you go across a period
  3. As atomic number increases, the properties of the elements show trends which repeat themselves in each period of the periodic table
  4. These trends are known as periodic trends and the study of these trends in known as periodicity
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6
Q

Atomic Radius

A
  1. The atomic radius of an element is a measure of the size of an atom
  2. It is the distance between the nucleus of an atom and the outermost electron shell
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7
Q

Atomic Radius trends

A

Atomic radii show predictable patterns across the periodic table

  1. They generally decrease across each period
  2. They generally increase down each group
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8
Q

Electron Shell Theory (atomic radius)

A
  1. Explains the trend of atomic radius (decreasing across each period and increasing down each group)
  2. Atomic radii decrease as you move across a period as the atomic number and effective nuclear charge increases. There are more protons in the nucleus
  3. Since there are more protons, the pull of the protons/force of attraction is larger so it pulls the electrons in making the atomic radius smaller.
  4. Atomic radii increase moving down a group due to electron shielding. There is an increased number of shells going down the group.
  5. The electrons in the inner shells repel the electrons in the outermost shells thus shielding them from the positive nuclear charge. As a result the pull of the nucleus is smaller which results in larger atomic radius
  6. Atomic radius sharply increases between the noble gas at the end of each period and the alkali metal at the beginning of the next period
  7. This is because the alkali metals at the beginning of the next period have one extra principal quantum shell
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9
Q

Ionic Radius

A

The ionic radius of an element is a measure of the size of an ion

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10
Q

Ionic Radius trends

A
  1. The trend down a group is the same as atomic radius – it increases down a group
  2. The trend across a period is not so straightforward as it depends on whether it is positive or negative ions being considered
  3. General rule of thump is that negative ions (anions) are a lot bigger than positive ions (cations)
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11
Q

Ionisation energy

A
  1. Ionisation is a process in which atoms lose or gain electrons and become ions
  2. The first ionisation (I1) energy of an element is the energy required to remove one electron from a gaseous atom
  3. The second ionisation (I2) energy involves the removal of a second electron
  4. As more electrons are removed, the positive charge of the nucleus relative to the electrons increases, making it more difficult for more electrons to be removed from the atom
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12
Q

Effective nuclear charge

A

the charge that the outermost electron experiences due to the positive nucleus

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13
Q

ionisation energy down a group

A
  1. ionisation energy goes down a group due to electron shielding
  2. The distance between the nucleus and outer electron increases
  3. The shielding by inner shell electrons increases
  4. The effective nuclear charge is decreasing as shielding increases hence the protons have a smaller force of attraction on the electrons and it is easier to lose one
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14
Q

ionisation energy across a period

A
  1. General trend across a period → ionisation energy increases
  2. This is because across a period the nuclear charge increases (pull of the protons - this is because there are more protons in the nucleus across a period)
  3. The distance between the nucleus and outer electron remains reasonably constant
  4. The shielding by inner shell electrons remains the same
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15
Q

4 factors that affect ionisation energy

A
  1. Size of the nuclear charge: the nuclear charge increases with increasing atomic number, which means that there are greater attractive forces between the nucleus and outer electrons, so more energy is required to overcome these attractive forces when removing an electron
  2. Distance of outer electrons from the nucleus: electrons in shells that are further away from the nucleus are less attracted to the nucleus so the further the outer electron shell is from the nucleus, the lower the ionisation energy
  3. Shielding effect of inner electrons: the shielding effect is when the electrons in full inner shells repel electrons in outer shells preventing them to feel the full nuclear charge so the greater the shielding of outer electrons by inner electron shells, the lower the ionisation energy
  4. Spin-pair repulsion: paired electrons in the same atomic orbital in a sub-shell repel each other more than electrons in different atomic orbitals; this makes it easier to remove an electron (which is why the first ionisation energy is always the lowest)
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16
Q

Ionisation energy between the last element in one period and the first element in the next period

A
  1. There is a rapid decrease in ionisation energy between the last element in one period and the first element in the next period caused by:
  2. The increased distance between the nucleus and the outer electrons
  3. The increased shielding by inner electrons
  4. These two factors outweigh the increased nuclear charge
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17
Q

Ionisation energy between Beryllium and Boron

A
  1. There is a slight decrease in 1st ionisation energy between beryllium and boron as the fifth electron in boron is in the 2p sub-shell which is further away from the nucleus than the 2s sub-shell of beryllium
  2. Beryllium has a first ionisation energy of 900 kJ mol-1 as its electron configuration is 1s2 2s2
  3. Boron has a first ionisation energy of 801 kJ mol-1 as its electron configuration is 1s2 2s2 2p1
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18
Q

Ionisation energy between nitrogen and oxygen

A
  1. There is a slight decrease in 1st ionisation energy between nitrogen and oxygen due to spin-pair repulsion in the 2p sub-shell of oxygen
  2. Nitrogen has a first ionisation energy of 1402 kJ mol-1 as its electron configuration is 1s2 2s2 2p3
  3. Oxygen has a first ionisation energy of 1314 kJ mol-1 as its electron configuration is 1s2 2s2 2p4
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19
Q

successive ionisation energies of an element

A
  1. The successive ionisation energies of an element increase as removing an electron from a positive ion is more difficult than from a neutral atom
  2. As more electrons are removed the attractive forces increase due to decreasing shielding and an increase in the proton to electron ratio
  3. The increase in ionisation energy, however, is not constant and is dependent on the atom’s electronic configuration
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20
Q

successive ionisation energies of an element (first to third)

A
  1. The first electron removed has a low ionisation energy as it is easily removed from the atom due to the spin-pair repulsion of the electrons in the 4s orbital
  2. The second electron is a little more difficult to remove than the first electron as you are removing an electron from a positively charged ion
  3. The third electron is much more difficult to remove than the second one corresponding to the fact that the third electron is in a principal quantum shell which is closer to the nucleus (3p)
  4. The big jumps on the graph (of successive ionisation energies) show the change of shell and the small jumps are the change of sub-shell
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21
Q

Electron Affinity

A
  1. When atoms gain electrons they become negative ions or anions
  2. Electron affinity (EA) can be thought of as the opposite process of ionisation energy
  3. First electron affinity is the amount of energy released when one mole of electrons is gained by one mole of atoms of an element in the gaseous state to form one mole of gaseous ions
  4. The first electron affinity is always exothermic. However, the second electron affinity can be an endothermic process
  5. Electron affinities are measured under standard conditions which are 298 K and 100 kPa
  6. The units of EA are kilojoules per mole (kJ mol-1)
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22
Q

Trends in electron affinity

A
  1. Electron affinities generally decrease down a group
  2. As the atoms become larger the attraction for an additional electron is less, since the effective nuclear charge is reduced due to increased shielding. It is harder for the nucleus to attract an electron
  3. Electron affinity become less exothermic going down the group. An exception to this is fluorine whose electron affinity is smaller than expected
  4. This is because fluorine is such a small atom and an additional electron in the 2p sub-shell experiences considerable repulsion with the other valence electrons
  5. Noble gases do not form negative ions so they do not show up in charts
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23
Q

Electronegativity

A
  1. Electronegativity is the ability of an atom to attract a pair of electrons towards itself in a covalent bond
  2. It attracts elements that are already in a covalent bond
  3. This phenomenon arises from the positive nucleus’s ability to attract the negatively charged electrons, in the outer shells, towards itself
24
Q

Pauling Scale

A
  1. The Pauling scale is used to assign a value of electronegativity for each atom
  2. Fluorine is the most electronegative atom on the Periodic Table, with a value of 4.0 on the Pauling Scale. It is best at attracting electron density towards itself when covalently bonded to another atom
25
Q

Factors that affect electronegativity

A
  1. Nuclear charge
  2. Atomic Radius
  3. Electron Shielding
26
Q

Nuclear charge affect on electronegativity

A
  1. Attraction exists between the positively charged protons in the nucleus and negatively charged electrons found in the energy levels of an atom
  2. An increase in the number of protons leads to an increase in nuclear attraction for the electrons in the outer shells
  3. Therefore, an increased nuclear charge results in an increased electronegativity
27
Q

Atomic radius affect on electronegativity

A
  1. The atomic radius is the distance between the nucleus and electrons in the outermost shell
  2. Electrons closer to the nucleus are more strongly attracted towards its positive nucleus
  3. Those electrons further away from the nucleus are less strongly attracted towards the nucleus
  4. Therefore, an increased atomic radius results in a decreased electronegativity
28
Q

Electron shielding affect on electronegativity

A
  1. Filled energy levels can shield the effect of the nuclear charge causing the outer electrons to be less attracted to the nucleus
  2. The addition of shells, and sub-shells in an atom will cause electrons to have less of an attraction.
  3. Therefore, an increased number of shells and sub-shells will result in decreased electronegativity
29
Q

electronegativity down a group

A
  1. There is a decrease in electronegativity going down the group
  2. The nuclear charge increases as more protons are being added to the nucleus
  3. However, each element has an extra filled electron shell, which increases shielding
  4. The addition of the extra shells increases the distance between the nucleus and the outer electrons resulting in larger atomic radii
  5. Overall, there is decrease in attraction between the nucleus and outer bonding electrons
  6. Effective nuclear charge has decreased down the group
30
Q

electronegativity across a period

A
  1. Electronegativity increases across a period
  2. The nuclear charge increases with the addition of protons to the nucleus
  3. Shielding remains the same across the period as no new shells are being added to the atoms
  4. The nucleus has an increasingly strong attraction for the bonding pair of electrons of atoms across the period
  5. This results in smaller atomic radii
31
Q

Properties of metals

A
  1. one-three outer shell electrons
  2. Metallic bonding due to loss of outer shell electrons
  3. good conductors of electricity
  4. Basic oxides (a few are amphoteric)
  5. Many react with acids
  6. Malleable, can be bent and shaped
  7. High melting and boiling point
32
Q

Properties of non-metals

A
  1. four-seven electrons in the outer shell
  2. Covalent bonding by sharing of outer shell electrons
  3. Poor conductors of electricity
  4. Acidic oxides (some are neutral)
  5. Does not react with acids
  6. Flaky and brittle
  7. Low melting and boiling point
33
Q

Reasons for properties of metals

A
  1. The low ionisation energies and low electronegativities of metals can account for the ability of their valence electrons to move away from the nucleus
  2. This is known as ‘delocalisation‘ of the electrons
  3. These properties increase from left to right as you transition from metal to metalloid to non-metal
34
Q

Reasons for properties of non-metals

A
  1. The high electronegativity and electron affinity of non-metals can be related to their tendency to share electrons and form covalent bonds, either with themselves or other non-metal elements
  2. The similarities in electronegativities of the diagonal band of metalloids which divides the metals from the non-metals explains the behaviour of metalloids
35
Q

Oxides across a period

A
  1. The acid-base character of the oxides provides evidence of chemical trends in the periodic table
  2. The broad trend is that oxides change from basic through amphoteric to acidic across a period
36
Q

Period 3 elements (oxides)

A
  1. In order : Basic, Basic, Amphoteric, Acidic, Acidic, Acidic
  2. The acidic and basic nature of the Period 3 elements can be explained by looking at their structure, bonding and the Period 3 elements’ electronegativity
  3. The difference in electronegativity between oxygen and Na, Mg and Al is the largest
  4. Electrons will therefore be transferred to oxygen when forming oxides giving the oxide an ionic bond
  5. The oxides of Si, P and S will share the electrons with the oxygen to form covalently bonded oxides
  6. The oxides of Na and Mg which show purely ionic bonding produce alkaline solutions with water as their oxide ions (O2-) become hydroxide ions (OH–)
  7. The oxides of P and S which show purely covalent bonding produce acidic solutions with water because when these oxides react with water, they form an acid which donates H+ ions to water
  8. The metallic oxides form hydroxides when they react with water
  9. The non-metallic oxides form oxoacids when they react with water
37
Q

sulfur trioxide with water

A
  1. SO3(g) + H2O(l) → H2SO4(aq)

2. pH change (1) - strongly acidic

38
Q

Sulfurous acid with water

A

H2SO4(aq) + H2O(l) → H3O+ (aq) + HSO4–(aq)

39
Q

sodium oxide with water

A
  1. Na2O(s) + H2O(l) → 2NaOH(aq)

2. pH change (14) - strongly alkaline

40
Q

magnesium oxide with water

A
  1. MgO(s) + H2O(l) → Mg(OH)2(aq)

2. pH change (10) - weakly alkaline

41
Q

Phosphorous pentoxide with water

A
  1. P4O10(s) + 6H2O(l) → 4H3PO4(aq)

2. pH change (2) - strongly acidic

42
Q

Nitrogen Dioxide with water

A
  1. 2NO2(aq) + H2O(l) → HNO3(aq) + HNO2(aq)

2. pH change (1) - strongly acidic

43
Q

sulfur dioxide with water

A
  1. SO2(g) + H2O(l) → H2SO3(aq)

2. 2. pH change (1) - strongly acidic

44
Q

Explaining period 3 oxides

A
  1. Metal and non-metal elements generally form ionic compounds so the elements Na to Al have giant ionic structures
  2. The oxides become more ionic as you go down the group as the electronegativity decreases
  3. The oxides become less ionic as you go across a period as the electronegativity increases
  4. The oxides of non-metals such as S, N and P form molecular covalent compounds
45
Q

Group 1 Metals

A
  1. The group 1 metals are called the alkali metals because they form alkaline solutions with high pH values when reacted with water
  2. Group 1 metals are lithium, sodium, potassium, rubidium, caesium and francium
  3. They all end in the electron configuration ns1
46
Q

Physical properties of the group 1 metals

A
  1. Are soft and easy to cut, getting softer and denser as you move down the group
  2. Have shiny silvery surfaces when freshly cut
  3. Conduct heat and electricity
  4. They all have low melting points and low densities and the melting point decreases going down the group as the atomic radius increases and the metallic bonding gets weaker
47
Q

Chemical properties of the group 1 metals

A
  1. They react readily with oxygen and water vapour in air so they are usually kept under oil to stop them from reacting
  2. Group 1 metals will react similarly with water, reacting vigorously to produce an alkaline metal hydroxide solution and hydrogen gas
48
Q

Reactions of lithium and water

A
  1. 2LI(s) + 2H2O (l) → 2LiOH(aq) + H2(g)
  2. Reaction is slower than with sodium, it bubbles of H2 gas
  3. lithium doesn’t melt due to its high melting point
49
Q

Reactions of sodium and water

A
  1. 2Na(s) + 2H2O (l) → 2NaOH(aq) + H2(g)
  2. Bubbles of H2 gas
  3. melts into a shiny ball that bounces and dashes around the surface
  4. NaOH formed which produces a highly alkaline solution
50
Q

Reactions of potassium and water

A
  1. 2K(s) + 2H2O (l) → 2KOH(aq) + H2(g)
  2. Reacts more violently than sodium, bubbles of H2 gas
  3. melts into a shiny ball that bounces and dashes around the surface
  4. Hot enough to burn H2 which forms lilac flame
51
Q

Alkali metals with halogens

A
  1. All the alkali metals react vigorously with halogens
  2. The reaction results in an alkali metal halide salt
  3. The reaction becomes increasingly vigorous going down group 1 because the atoms of each element get larger going down the group
  4. This means that the ns1 electron gets further away from the nucleus and is shielded by more electron shells.
  5. The further an electron is from the positive nucleus, the easier it can be lost in reactions
52
Q

The halogens

A
  1. These are the non-metals that are poisonous and include fluorine, chlorine, bromine, iodine and astatine
  2. Halogens are diatomic, meaning they form molecules of two atoms
  3. All halogens have seven electrons in their outer shell
  4. They form halide ions by gaining one more electron to complete their outer shells
53
Q

Trends in melting and boiling point of halogens

A

The density and melting and boiling points of the halogens increase as you go down the group

54
Q

trend in reactivity in halogens

A
  1. Reactivity of halogens decreases as you go down the group
  2. The halogens electron configurations all end in ns2np5
  3. Each outer shell contains seven electrons and when they react, they will need to gain one outer electron to get a full outer shell of electrons
  4. Going down the group, the electron affinity decreases and the atomic radius increases
  5. As you go down the group, the number of shells of electrons increases so shielding also increases
  6. This means that the outer electrons are further from the nucleus so there are weaker electrostatic forces of attraction that attract the extra electron needed
  7. The electron is attracted less readily, so the lower down the element is in the group the less reactive it is
55
Q

Reaction of the halogens with halide ions in displacement reactions

A
  1. A halogen displacement occurs when a more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide
  2. The reactivity of halogens increases as you move up the group
  3. Out of the 3 halogens, chlorine, bromine and iodine, chlorine is the most reactive and iodine is the least reactive
56
Q

Amphoteric oxides

A
  1. Some metallic oxides are amphoteric oxides.
  2. They react with both acids and alkalis to produce salts.
  3. One example is aluminium oxide (Al2O3).

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