3.1.1 - Periodicity Flashcards
(112 cards)
1
Q
Hydrogen (H)
A
1s1
2
Q
Helium (He)
A
1s2
3
Q
Lithium (Li)
A
1s2, 2s1
4
Q
Beryllium (Be)
A
1s2, 2s2
5
Q
Boron (B)
A
1s2, 2s2, 2p1
6
Q
Carbon (C)
A
1s2, 2s2, 2p2
7
Q
Nitrogen (N)
A
1s2, 2s2, 2p3
8
Q
Oxygen (O)
A
1s2, 2s2, 2p4
9
Q
Fluorine (F)
A
1s2, 2s2, 2p5
10
Q
Neon (Ne)
A
1s2, 2s2, 2p6
11
Q
Sodium (Na)
A
1s2, 2s2, 2p6, 3s1
12
Q
Magnesium (Mg)
A
1s2, 2s2, 2p6, 3s2
13
Q
Aluminium (Al)
A
1s2, 2s2, 2p6, 3s2, 3p1
14
Q
Silicon (Si)
A
1s2, 2s2, 2p6, 3s2, 3p2
15
Q
Phosphorus (P)
A
1s2, 2s2, 2p6, 3s2, 3p3
16
Q
Sulfur (S)
A
1s2, 2s2, 2p6, 3s2, 3p4
17
Q
Chlorine (Cl)
A
1s2, 2s2, 2p6, 3s2, 3p5
18
Q
Argon (Ar)
A
1s2, 2s2, 2p6, 3s2, 3p6
19
Q
Potassium (K)
A
1s2, 2s2, 2p6, 3s2, 3p6, 4s1
20
Q
Calcium (Ca)
A
1s2, 2s2, 2p6, 3s2, 3p6, 4s2
21
Q
Scandium (Sc)
A
1s2, 2s2, 2p6, 3s2, 3p6, 3d1, 4s2
22
Q
Titanium (Ti)
A
1s2, 2s2, 2p6, 3s2, 3p6, 3d2, 4s2
23
Q
Vanadium (V)
A
1s2, 2s2, 2p6, 3s2, 3p6, 3d3, 4s2
24
Q
Chromium (Cr)
A
1s2, 2s2, 2p6, 3s2, 3p6, 3d5, 4s2
25
Manganese (Mn)
1s2, 2s2, 2p6, 3s2, 3p6, 3d5, 4s2
26
Iron (Fe)
1s2, 2s2, 2p6, 3s2, 3p6, 3d6, 4s2
27
Cobalt (Co)
1s2, 2s2, 2p6, 3s2, 3p6, 3d7, 4s2
28
Nickel (Ni)
1s2, 2s2, 2p6, 3s2, 3p6, 3d8, 4s2
29
Copper (Cu)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2
30
Zinc (Zn)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2
31
Gallium (Ga)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p1
32
Germanium (Ge)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p2
33
Arsenic (As)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p3
34
Selenium (Se)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p4
35
Bromine (Br)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p5
36
Krypton (Kr)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6
37
Rubidium (Rb)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 5s1
38
Strontium (Sr)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 5s2
39
Yttrium (Y)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 4d1, 5s2
40
Zirconium (Zr)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 4d2, 5s2
41
Niobium (Nb)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 4d4, 5s2
42
Molybdenum (Mo)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 4d5, 5s2
43
Technetium (Tc)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 4d5, 5s2
44
Ruthenium (Ru)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 4d7, 5s2
45
Rhodium (Rh)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 4d8, 5s2
46
Palladium (Pd)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 4d10, 5s2
47
Silver (Ag)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 4d10, 5s2
48
Cadmium (Cd)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 4d10, 5s2
49
Indium (In)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 4d10, 5s2, 5p1
50
Tin (Sn)
1s2, 2s2, 2p6, 3s2, 3p6, 3p10, 4s2, 4p6, 4d10, 5s2, 5p2
51
Describe what Dimitri’s periodic table was like in 1863
- 56 known elements
- arranged in order of atomic mass
- lined up elements of similar properties
- predicted properties of missing elements from group trends
52
Describe how the periodic table is arranged today
- 118 elements
- arranged in 7 rows (periods)
- 18 vertical groups
- element positions are linked to their physical and chemical properties
53
across the period are _________ trend in properties of elements called ___________ (ie from metals to non-metals)
- repeating
- periodicity
54
What is group 1 called?
Alkali metals
55
What is Group 2 called?
Alkali earth metals
56
What is Group 3-12 called?
Transition elements
57
What is Group 15 called?
Pnictogens
58
What is Group 16 called?
Chalogens
59
What is Group 17 called?
Halogens
60
What is Group 18 called?
Noble gases
61
What does ionisation energy measure?
This measures how easily an atom loses electrons to form positive ions
62
What is the first ionisation energy?
The energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions
63
Write the first ionisation equation for Sodium
Na (g) -> Na+ (g) + e-
64
Factors affecting ionisation energy
- atomic radius
- nuclear charge
- electron shielding
65
How many ionisation energies does an element have?
An element has as many ionisation energies as there are electrons. For example, helium has 2 electrons and 2 ionisation energies.
66
What is the second ionisation energy?
Is the energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+
67
In Fluorine, why is there a large increase between the 7th and 8th ionisation energies?
It suggests that the eighth electron must be removed from a different shell. The first shell (n=1, closer to the nucleus) contains 2 electrons. The second shell (n=2, the outer shell) contains 7 electrons.
68
Successive ionisation energies allow predictions to be made about:
- the number of electrons in the outer shell
- the group of the element in the periodic table
- the identity of an element
69
What is Aubaf Principle?
States that electrons are filled into atomic orbitals in the increasing order of orbital energy level.
70
How many electrons can each orbital hold?
Two electrons
71
What are metal/ non-metals near the divide with in between properties called?
Semi-metals or metalloids
72
Which group is the divide between metals and non-metals clearest?
Group 14 (4)
73
Which metal is liquid at room temperature?
Mercury
74
What is a positive ion called?
Cation
75
What elements do metallic bonding?
Between metals and metals
76
What is metallic bonding?
Metallic bonding is the strong electrostatic attraction between cations (positive ions) and delocalised electrons. The cations are fixed in position, maintaining the structure and shape of the metal. The delocalised elections are mobile and are able to move throughout the structure. Only the electrons move.
77
A metal structure, billions of metals atoms are held together by metallic bonding in a _____ ________ _______.
Giant metallic lattice
78
What properties do most metals have?
- Strong metallic bonds -> attraction between positive ions and delocalised electrons.
- High electrical conductivity -> in both solid and liquid state.
- High melting and boiling points.
79
What does melting point depend on in metallic bonds?
The strength of the metallic bonds that hold the atoms together in the giant metallic lattice
80
For most metals, high temperatures are necessary to provide the large amount of energy needed to overcome the strong _____________ __________ between the _______ and _________.
- Electrostatic attraction
- Cations
- Electrons
81
Are metals soluble?
No - they don’t dissolve.
82
Which type of bonding is between non-metals and non-metals?
Covalent bonding.
83
In a solid state, covalently bonded molecules form what type of structure that is held together by what?
- Simple molecular structure
- Weak intermolecular forces
84
Do molecules that form simple molecular lattices have high or low boiling/ melting points?
Low
85
Non-metals like _____, ______ and _______ have very different lattice structures - whereby the billions of atoms are held together by a network of what?
- Boron, carbon and silicon
- Form a network of strong covalent bonds to form a giant covalent structure
86
What structure does carbon in its diamond form, form?
Tetrahedral structure
87
How many bonds can a carbon atom have?
4 bonds
88
What does periodicity refer to in chemistry?
Periodicity refers to the repeating patterns in the properties of elements across different periods and groups in the periodic table.
89
What is the structure of the periodic table?
The periodic table is arranged in periods (rows) and groups (columns). Elements in the same group have similar chemical properties, and properties vary across periods.
90
What is a period in the periodic table?
A period is a horizontal row in the periodic table. As you move across a period, the atomic number increases, and the properties of the elements change in a predictable pattern.
91
What is a group in the periodic table?
A group is a vertical column in the periodic table. Elements in the same group share similar chemical and physical properties, particularly in terms of their outer electron configuration.
92
How does atomic radius change across a period?
Atomic radius decreases across a period from left to right because the increasing nuclear charge pulls the electrons closer to the nucleus, reducing the size of the atom.
93
How does atomic radius change down a group?
Atomic radius increases as you go down a group because additional electron shells are added, which increases the distance between the outer electrons and the nucleus.
94
What is first ionization energy?
First ionization energy is the energy required to remove one mole of electrons from one mole of atoms in the gas phase to form positive ions (cations).
95
How does first ionization energy change across a period?
First ionization energy increases across a period because the atomic radius decreases and the nuclear charge increases, making it harder to remove an electron.
96
How does first ionization energy change down a group?
First ionization energy decreases down a group because the outer electrons are farther from the nucleus and experience more shielding from inner electrons, making them easier to remove.
97
What is electron shielding?
Electron shielding occurs when inner-shell electrons repel outer-shell electrons, reducing the effective nuclear charge felt by the outer electrons. This makes them easier to remove.
98
How does electronegativity change across a period?
Electronegativity increases across a period because the atomic radius decreases, and the nucleus has a stronger attraction for bonding electrons.
99
How does electronegativity change down a group?
Electronegativity decreases down a group because the atomic radius increases, and the outer electrons are further from the nucleus, reducing the pull on bonding electrons.
100
What is metallic character?
Metallic character refers to the tendency of an element to lose electrons and form positive ions. Metals are typically more metallic in character than non-metals.
101
How does metallic character change across a period?
Metallic character decreases across a period because elements become more non-metallic as they gain electrons and have higher electronegativity.
102
How does metallic character change down a group?
Metallic character increases down a group because the outer electrons are further from the nucleus, making it easier for elements to lose electrons and display metallic properties.
103
What is melting point and how does it vary across periods?
Melting point is the temperature at which a solid turns into a liquid. In a period, the melting point generally increases across the metals (due to stronger metallic bonds) and decreases for non-metals (due to weaker intermolecular forces).
104
How does melting point change down a group?
Melting point generally increases down a group for metals due to stronger metallic bonding with more electrons. For non-metals, the melting point often decreases because the forces between molecules are weaker.
105
What is the trend in conductivity across periods for metals and non-metals?
Metals generally have high electrical conductivity, which decreases as you move across the period to non-metals, which have low conductivity.
106
How does the atomic size relate to ionization energy?
The smaller the atomic size, the higher the ionization energy, because the electrons are closer to the nucleus and more tightly held.
107
How does ionization energy affect the reactivity of elements?
Elements with low ionization energies (e.g., alkali metals) are more reactive because it is easier for them to lose electrons, while elements with high ionization energies (e.g., noble gases) are less reactive.
108
What is the trend in reactivity for metals across a period?
The reactivity of metals decreases across a period because metals lose electrons less easily as their ionization energies increase.
109
What is the trend in reactivity for non-metals across a period?
The reactivity of non-metals increases across a period because they are more likely to gain electrons as the atomic size decreases and electronegativity increases.
110
What happens to the properties of elements as you move across a period?
As you move across a period, the elements transition from metallic to non-metallic character. The atomic size decreases, ionization energy increases, and electronegativity also increases.
111
What is periodicity in terms of atomic structure?
Periodicity refers to the recurring trends in the properties of elements due to the repeating pattern of electron configurations across periods and groups in the periodic table.
112
What are transition elements and how do their properties vary across a period?
Transition elements are elements found in the d-block of the periodic table. Their properties, such as ionization energy, color, and reactivity, vary in more complex ways than main-group elements, often showing a variety of oxidation states.
