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IB Chemistry > Unit 2: Periodic Table Trends > Flashcards

Unit 2: Periodic Table Trends Flashcards

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1
Q

What are the 6 main trends in the periodic table covered in this unit?

A

Atomic radii, ionization energy, ionic radii, electron affinity, electronegativity, reactivity

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2
Q

What is the nuclear charge?

A

Positive charge from the nucleus/# of protons

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3
Q

What are shielding electrons?

A

The electrons located between the nucleus and the valence electrons

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4
Q

What is the effective nuclear charge?

A

The attractive force of the nucleus that is felt by valence electrons

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5
Q

How is Zeff. calculated?

A

Nuclear charge - shielding electrons

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6
Q

What does a higher Zeff mean?

A

A stronger pull on electrons

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7
Q

Why do valence electrons feel less of an attractive force from the nucleus?

A

Because they are at higher energy levels which means they are further from the nucleus resulting in less of an electrostatic force and because the shielding electrons repel the valence electrons

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8
Q

What is the general trend for the atomic radii in the periodic table?

A

Increases down a group, decreases across a period

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9
Q

Why does the atomic radii generally increase down a group?

A

Due to higher energy levels further from nucleus and there is an increase in shielding electrons repelling valence electrons

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10
Q

Why does atomic radii decrease across a period?

A

Due to an increase in nuclear charge while shielding and energy level remains the same (increase in Zeff across period)

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11
Q

What is the exception to atomic radii decreasing across a period?

A

There can be a slight jump up when moving into p block due to slightly higher energy and different shape

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12
Q

What are the two ways of measuring atomic radius?

A

Bonding diameter/covalent radius and non bonded radius/Vander waals radius

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13
Q

What do you compare when comparing atomic radii?

A

Energy level, shielding electrons, and effective nuclear charge

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14
Q

Why does ionization energy decrease going down a group?

A

because the electrons are in higher energy levels and further from the nucleus and the radius is bigger which means the electrons are less attracted to the nucleus

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15
Q

Why does ionization energy increase across a period?

A

because the atomic radius is decreasing due to same energy level and same shielding but greater effective nuclear charge which pulls electrons in tighter

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16
Q

What are the exceptions to the trend in ionization energy?

A

There is a decrease from group 2 to group 13 due to change to p block which is slightly higher energy and easier to ionize, and there is a decrease from group 15 to group 16 because group 15 has a half filled p orbital while group 16 has one orbital with paired electrons which require pairing energy therefore is easier to ionize

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17
Q

What does a parent atom become after losing electrons?

A

a smaller cation with a larger effective nuclear charge

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18
Q

What does a parent atom become when it gains electrons?

A

a larger anion with a smaller effective nuclear charge

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19
Q

How does the trend in ionic radii differ from the trend in atomic radii?

A

at group 14, there is an increase in size of radii due to the ability of the atoms to become anions rather than cations

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20
Q

What is electron affinity?

A

the change in energy when one mole of electrons is added to one mole of a neutral atom in the gas state

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21
Q

What does a larger number mean for electron affinity?

A

the more likely the atom is to gain electrons/higher affinity

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22
Q

What does a negative electron affinity mean?

A

the process is exothermic/energy is being released

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23
Q

Why do noble gases have no values listed for electron affinity?

A

because they do not gain electrons

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24
Q

Why are there second electron affinity values positive for oxygen and sulfur?

A

because the electron to electron repulsion must be overcome by adding energy

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25
What is the general trend for electron affinity?
Numbers become more negative left to right
26
Why do halogens have the larges electron affinity?
because their sublevel is almost full
27
What are the exceptions to the general trend for electron affinity?
group 2 has a more positive electron affinity than expected due to having a full s orbital, therefore they are less attracted to electrons, and group 15 also has a more positive electron affinity than expected due to having a half filled p sublevel which does not require pairing energy
28
What is electronegativity?
the relative attraction that an atom has for a shared pair of electrons in a bond
29
What scale is used for the values of electronegativity?
the pauling scale
30
What is the general trend for electronegativity?
it increases from the bottom left to the top right
31
What is a dipole?
a partial charge assigned to a molecule according to the electronegativity/unequal sharing of electrons
32
How do you assign dipoles?
delta+ = partially positive = lower electronegativity, delta- = partially negative = higher electronegativity
33
What does calculating the difference in electronegativity tell us?
The type of bond
34
What are the different types of bonds?
non-polar covalent, polar covalent, ionic
35
What are the boundaries for the electronegativity difference value for determining the type of bond?
less than 0.5 = non-polar covalent, 0.5-1.8 = polar covalent, more than 1.8 = ionic
36
Why do atoms with a larger radii not attract bonding electrons as easily?
because bonding electrons are in higher energy levels and further from nucleus and they experience more sheilding from inner electrons
37
Why do atoms with a smaller radii have a stronger attraction to their bonding electrons?
because the energy level and shielding is the same across a period and the effective nuclear charge increases across a period
38
What are some characteristics of alkali metals (group 1)?
they are too reactive to be found in nature, soft and shiny, stored in oil to prevent reaction with air and water, lowest ionization energy
39
How do elements from group 1 react with water?
forms hydrogen gas and a basic aqueous solution through single displacement reaction
40
What are some characteristics of noble gases (group 18)?
they have full sublevels, unreactive, colourless, monatomic/exist as single atoms, high ionization energy
41
What are some characteristics of halogens (group 17)?
they usually exist as diatomic molecules, their reactivity decreases down the group
42
Why does the reactivity of halogens decrease down the group?
because the radii increase and the attraction for outer electrons decreases
43
What do halogens and group 1 metals react together to form?
halides
44
What do the most vigorous reactions occur between?
elements with the largest electronegativity difference
45
What happens in a single displacement reaction when a halide reacts with a halogen on its own that is more reactive than the halogen in the ionic compound?
the more reactive halogen displaces the ions of the less reactive halogen from its compound
46
What is required for a single displacement reaction to occur between a halide and another halogen?
the halogen by itself must be more reactive/higher in group 17 than the one in the ionic compound
47
What do halogens form with silver?
insoluble salts
48
What does it mean for an element to be reduced?
it gained electrons
49
What does it mean for an element to be oxidized?
it lost electrons
50
What kind of structure do ionic compounds of the period 3 metals make with oxygen?
giant ionic structures with an endlessly repeating lattice
51
What are some characteristics of the ionic compounds of period 3 metals with oxygen?
they have a high melting and boiling point due to stron attractions between particles, theyconduct electricity when dissolved in a solution becuase the ions are free to move around
52
What structure does the oxide of silicon have?
the giant structure with an endlessly repeating lattice
53
What happens to the ionic properties across period 3 as we bond with oxygen?
they are decreasing from left to right
54
What molecules would have a change in electronegativity of 0?
diatomic molecules
55
What characterizes a polar covalent bond?
there is a positive and negative side of the bond due to unequal sharing of electrons
56
Why happens with oxides down a group?
they become more ionic due to the larger difference in electronegativity
57
What are oxidation states?
the numbers assigned to show the number of electrons transferred when forming a bond
58
What oxide is basic?
metal oxide
59
What oxide is acidic?
non-metal oxide
60
What oxide is amphoteric (acts as acid or a base)?
metalloid oxide
61
What will the product be when the reactants is a metal oxide and water?
metalOH (aq) (base)
62
What will non-metal period 3 elements react with water to form?
an acid
63
What is the product(s) of sulfur trioxide reacting with water?
sulfuric acid
64
What is the formula for sulfuric acid?
H2SO4
65
What is the formula for hydrochloric acid?
HCl
66
How would a metalloid such as aluminum act as an acid/base?
if it lowers pH = acid, if it makes pH higher = base
67
Why is rain naturally slightly acidic?
because CO2 is dissolved in water which produces carbonic acid
68
What is the formula for carbonic acid?
H2CO3
69
What is an external source that contributes to acid rain?
fossil fuels
70
When is sulfurous acid formed?
SO2 + H2O
71
What is the formula for sulfurous acid?
H2SO3
72
What is NO produced from?
internal combustion engines
73
What does NO react with O2 to produce?
NO2
74
What is the formula for nitric acid?
HNO3
75
What is an example of why CO2 from combustion of fossil fuels is harmful?
because it dissolves in water which inhibits shell growth in marine animals
76
What does an oxidation reaction infer?
electron(s) transferred
77
How can oxidation happen?
addition of oxygen, removal of hydrogen, loss of electron(s), increase of oxidation state
78
What is the oxidation state of an element?
0
79
What does an increase in oxidation state mean?
the atom has lost some control over electrons
80
What does a decrease in oxidation state mean?
the atom has gained control over electrons
81
What is the first rule of oxidation states?
Atoms in the free element have an oxidation state of zero
82
What is the second rule of oxidation states?
in simple ions, the oxidation state is the same as the charge on the ion
83
What is the third rule for oxidation states?
the oxidation states of all the atoms in a neutral compound must add up to zero
84
what is the fourth rule for oxidation states?
the oxidation states of all the atoms in a polyatomic ion must add up to the charge on the ion
85
What is the fifth rule for oxidation states?
the usual oxidation state for an element in a compound is the same as the charge on its most common ion
86
What is the sixth rule for oxidation states?
F has the oxidation state of -1 in all its compounds as it is the most electronegative element
87
What is the seventh rule for oxidation states?
O has the oxidation state of -2 except when in peroxides where it is -1. When bonded to F, it is positive as it is less electronegative than F
88
What is the eighth rule for oxidation states?
Cl has the oxidation state of -1 except when bonded to the more electronegative F and O
89
What is the ninth rule for oxidation states?
H has the oxidation state of +1 except when bonded to the group 1 and 2 metals when it forms ionic hydrides
90
What is the tenth rule for oxidation states?
the oxidation state of a transition metal in a complex ion can be worked out from the charge on the ligands
91
What are some common ligands?
H2O, NH3, Cl-, CN-, OH-
92
How do you name compounds using oxidation numbers?
put the oxidation number of the "main" element in brackets between the two words (normal first word, second word could be acid, oxide, ion, etc)
93
What are the key features of transition metals (9)?
variable oxidation states, incomplete d-sublevels, often form coloured compounds, expanded octets are possible, conduct heat and electricity, used as catalysts, magnetic properties which depend on oxidation number and coordinate bonds, hight melting points, zinc not technically transition metal
94
What does it mean to be a catalyst in chemistry?
increases rate of reaction without interfering
95
Why is zinc not technically a transition metal?
because it has a full d sublevel and doesn't have variable oxidation states
96
What is important to remember when ionizing a transition metal with orbital diagrams?
remove the 4s electrons before 3d
97
How does the ionization energy trend differ for transition metals compared to s and p block elements?
it does not increase as drastically
98
Why does the ionization energy not increase as drastically for transition metals than for s and p block elements?
it is due to pairing electrons, because electron enter an inner shell orbital instead of a valence orbital, and because inner orbitals have greater shielding
99
What are the exceptions to ionization energy trends for transition metals?
Cr and Cu
100
What does it mean if a transition metal has stable high oxidation states?
it has unstable low oxidation states
101
What does it mean if a transition metal has stable high and low oxidation states?
it has unstable middle oxidation states
102
What does it mean if a transition metal has stable low oxidation states?
it has unstable high oxidation states
103
Which transition metals are stable high/unstable low?
Sc, Ti, V (first 3)
104
Which transition metals are unstable middle?
Cr, Mn (4-5)
105
Which transtion metals are stable low/unstable high?
Fe, Co, Ni, Cu, Zn (last 5)
106
What do transition metals make?
complex ions
107
What is a coordination bond?
type of bond where both electrons come from lone pair in a ligand)
108
What determines the coordination number for complex ions?
the number of bonds there are
109
What are coloured complexes because of for transition metals?
the partially filled d sublevel
110
How does colour work in terms of what is absorbed vs emitted for transition metals?
colour seen = wavelength emitted, wavelength of complimentary colour = colour absorbed
111
When a d sublevel splits, which orientations have higher energy?
the two along the bond axis compared to the three that are between the axis
112
What is the change in energy the result of when a d sublevel splits?
electrons moving from a lower energy orientation to a higher energy orientation
113
What does charge density in complex ions have to do with?
the ligands
114
What does a greater charge density mean?
a larger split in d orbitals/more energy
115
Which equation could determine energy from d subevel split?
E=(hc)/wavelength

IB Chemistry (22 decks)

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