Bonding Flashcards
(131 cards)
1
Q
What is Hund’s rule?
A
Electrons enter orbitals with parallel spins first (until all orbitals contain one electron)
2
Q
What does the last electron placed in an electron configuration represent?
A
The quantum number.
3
Q
What is the difference between abbreviated and unabbreviated electron configurations?
A
Unabbreviated shows the entire configuration for the element, while abbreviated shows the last noble gas (as a placeholder) and the rest of the configuration only.
4
Q
Abbreviated electron configuration for cobalt? (Just an example, no need to memorize)
A
[Ar] 4s^2 3d^7
Argon is the last noble gas before cobalt, so it acts as a placeholder.
5
Q
Where are valence electrons found and what is their significance?
A
They are found in the outermost energy level. They determine how reactive an element is.
6
Q
What is the octet rule?
A
8 total electrons in the valence energy level represents a stable electron configuration/arrangement.
7
Q
Noble gases are also called…
A
inert gases.
8
Q
Why are elements in the noble gases (group 18) stable?
A
They have eight electrons in the outermost energy level. They follow the octet rule.
9
Q
The noble gas electron configuration always ends in…Does this configuration represent a stable or unstable element?
A
Configuration ends in __s^2__p^6.
s2p6 represents a stable element/configuration.
10
Q
How can octet structure be achieved by an element that is not a noble gas (already stable)?
A
Valence electrons can be added or removed in the valence energy level to achieve octet structure (8 valence electrons, or s2p6 configuration).
11
Q
Elements in group 1 have how many valence electrons?
A
1.
12
Q
What is the significance if an element has four valence electrons?
A
The atom can either gain four electrons or lose four electrons to obtain octet structure.
13
Q
What does oxidation state show?
A
Charge. Ex: Li+ is the common oxidation state for lithium.
14
Q
What does the Lewis Dot diagram show?
A
The element’s valence electrons.
15
Q
Lanthanides and actinides are also known as…
A
Inner transition elements/metals.
16
Q
The most active metals are located…
A
In group 1 (1A) (Alkali metals)
17
Q
The most active nonmetals are located…
A
In group 17 (7A) (Halogens)
18
Q
What is the most active halogen?
A
Fluorine.
19
Q
Is an anion larger or smaller than its parent atom?
A
Larger.
20
Q
Is a cation larger or smaller than its parent atom?
A
Smaller.
21
Q
An element with both metallic and nonmetallic properties is called a…
A
metalloid or semi-metal.
22
Q
Elements in a group have the same number of…
A
Valence electrons.
23
Q
Elements in a period have the same number of…
A
Electron shells/energy levels.
24
Q
The majority of the elements on the periodic table are…(metals or nonmetals)
A
Metals.
25
As you go down a group, elements generally become (more or less) metallic.
More metallic.
26
What does a colored ion indicate?
A transition metal.
27
Group 1 elements are called...
How many valence electrons do they have?
What is their most common oxidation state?
Alkali metals. They have 1 valence electron. Their most common oxidation state is 1+.
28
Group 2 elements are called...
How many valence electrons do they have?
What is their most common oxidation state?
Alkaline earth metals. They have 2 valence electrons. Their most common oxidation state is 2+.
29
Groups 3-12 elements are called...
Transition metals.
30
Group 13 (3A) valence electrons? Oxidation state?
3 valence electrons. Oxidation state: 3+.
31
Group 14 (4A) valence electrons? Oxidation state?
4 valence electrons. Oxidation state: 4+/-.
32
Group 15 (5A) valence electrons? Oxidation state?
5 valence electrons. Oxidation state: 3-.
33
Group 16 (6A) valence electrons? Oxidation state?
6 valence electrons. Oxidation state: 2-.
34
Group 17 (7A) valence electrons? Oxidation state?
7 valence electrons. Oxidation state: 1-.
35
Group 18 (8A) valence electrons? Oxidation state?
8 valence electrons: Oxidation state: 0, because they are stable and non-reactive.
36
On the periodic table, where is the s-sublevel found?
Groups 1 and 2: alkali metals and alkaline earth metals. Also helium.
37
On the periodic table, where is the p-sublevel found?
Groups 13-18, excluding helium.
38
On the periodic table, where is the d-sublevel found?
Groups 3-12: transition metals.
39
On the periodic table, where is the f-sublevel found?
Lanthanide and actinide series.
40
After 4s on the periodic table, next in that row is:
3d. (NOT 4d!)
41
After 2s on the periodic table, next comes:
2p.
42
The f-sublevel contains only two rows of elements: The lanthanides and actinides. These two rows represent:
4f and 5f.
43
In each row of the "d" block there are only 10 elements. Why?
This is because the d sublevel can hold at maximum 10 electrons. The number of elements in each row of a certain sublevel corresponds to the maximum number of electrons that sublevel can hold.
44
What is the effect of distance between protons and electrons on force of attraction?
As distance between two oppositely charged particles (protons and electrons) increases, force of attraction decreases.
45
What is coulombic attraction?
Attraction between two oppositely charged particles.
46
As you move down a group, how does distance between valence electrons and nucleus change?
As you move down a group, distance between the valence electrons and the nucleus increases (because you are adding energy levels).
47
As you move down a group, how does force of attraction between valence electrons and nucleus change?
As you move down a group, force of attraction decreases because distance between electrons and protons increases.
48
How does the number of protons in the nucleus affect attractive force between electrons and protons?
As number of protons increases, attractive force increases.
49
For each proton added, how much does force increase?
Each proton increases force of attraction by 2.30 x 10^ -8 N if distance is .10 nm
50
Does force of attraction change when electrons are added?
Slightly, because electrons repel each other, but not by much. Each electron still gets approximately the full force of the nucleus. (Ex: If there are 3 protons and 3 electrons, each electron receives 6.90 x 10^-8.)
51
What is atomic radius?
A way to express the size of an atom
52
How does one find atomic radius?
Atomic radius= 1/2 the distance between the nuclei of 2 identical atoms that are bonded together in a diatomic molecule.
53
Atomico radius is measured in...
picometers (pm), which is 1 x 10^-12 m.
54
How does atomic radius change across a period? Why?
From left to right, atomic radius decreases across a period because protons are added to the nucleus but number of energy levels stays the same, so attractive force increases.
55
How does atomic radius change down a group? Why?
From top to bottom, atomic radius increases down a group because energy levels are added.
56
Define ionization energy.
The energy required to remove the first valence electron from a specific atom.
57
Ionization energy is measured in...
kj/mol
58
How does the ionization energy of valence electrons compare to that of inner electrons?
IE of valence is lower than IE of inner electrons.
59
What is electron shielding?
When inner electrons shield the valence electrons from the attractive force of the nucleus.
60
How does IE change across a period? Why?
Left to right, IE increases because more protons are added to nucleus but number of energy level stays the same, so force of attraction is stronger, making it harder to remove electrons.
61
How does IE change down agroup? Why?
Top to bottom, IE decreases because the increase in energy levels (increase in distance between electrons and protons) causes attractive force to decrease; easier to remove electrons.
62
What is electronegativity?
A measure of the ability of an atom to attract the electrons when the atom is part of a compound.
63
What is electron affinity?
The energy released when electrons are gained.
64
How is electronegativity measured?
By a scale: Pauling electronegativity scale.
65
What is the most electronegative element?
Fluorine: 3.98
66
Nonmetals usually have high/low electronegativity because...
High electronegativity because they have many valence electrons and typically gain electrons to become anions.
67
Metals usually have high/low electronegativity because...
Low electronegativity because they have few valence electrons and typically lose them to become cations.
68
How does electronegativity change down a group?
Top to bottom, electronegativity decreases because increase in atomic radius leads to weaker pull on electrons; easier to lose them.
69
How does electronegativity change across a period?
Left to right, electronegativity increases because of increase in number of protons and nuclear charge: Harder to lose electrons.
70
Which group has lowest electronegativity? Highest?
Alkali metals, halogens.
71
Do noble gases have electronegativity?
No, because they do not typically form compounds.
72
Do transition metals vary in electronegativity?
Not by much.
73
What is a chemical bond?
The force of attraction between two or more ions or atoms.
74
How are bonds created?
Bonds are created when atoms share or transfer valence electrons.
75
Why do atoms or ions form bonds?
To reach a stable octet structure/arrangement of valence electrons.
76
What is a covalent bond?
A bond that forms when nonmetal atoms share electrons to form molecules.
77
What is an ionic bond?
A bond that forms when one atom transfers electrons to another atom, resulting in the formation of ions. It occurs between a metal and a nonmetal atom and forms a crystal lattice instead of molecules.
78
What is a metallic bond?
A metal cation becomes attracted to its own valence electrons and the electrons of other ions of the same metal. It forms a lattice structure. This only occurs with metals.
79
What is the difference between polar and nonpolar covalent bonds?
In a polar bond, electrons are shared unequally between the ions because one of them has a higher electronegativity. (This results in one side of the molecule having a charge.) In a nonpolar bond, there is no charge because electrons are shared equally due to equal electronegativities.
80
Regarding bonding, nonmetals want to (gain/lose) electrons to form (anions/cations).
Nonmetals (groups 5-7) want to gain electrons to form anions.
81
Regarding bonding, metals want to (gain/lose) electrons to form (anions/cations).
Metals (groups 1-2 and transition) want to lose electrons to form cations.
82
If there is no difference in electronegativity, electrons are shared...
Equally.
83
Define electronegativity.
A measure of the attraction (pull) a particular atom has for electrons in a chemical bond. Measured on the Pauling Scale from 0-4.
84
If electonegativity is high, the atom...
wants to gain electrons.
85
If electronegativity is low, the atom...
wants to lose electrons.
86
How does electronegativity differentiate between ionic and covalent bonds?
If difference in electronegativity between two atoms is greater than 1.7, the bond is ionic. If EN difference < 1.7, the bond is covalent.
87
What is the relationship between electronegativity and ionic character?
As electronegativity increases, ionic character increases.
88
What are the three rules for writing ionic compound formulas?
1. Positive and negative charges must balance
2. Positive (cation- metal) written first
3. Negative (anion- nonmetatl) written last
89
Ionic formulas are...This means that they are made up of two single elements.
Binary.
90
How do you deal with the charges and subscripts when writing an ionic formula?
The metal's charge becomes the nonmetal's subscript number; the nonmetal's charge becomes the metal's subscript number.
91
Aluminum's most common oxidation state is 3+. Oxygen's most common oxidation state is 2-. Write the formula using these two elements and name the compound.
Al2O3. (2 and 3 are subscript.)
| Aluminum oxide.
92
How do you indicate the number of atoms of each element in an ionic compound?
Use subscripts.
93
When writing an ionic formula, the anion's ending changes to...
-ide. Ex: oxygen becomes oxide.
94
Binary means...
The compound is made of 2 single elements.
95
When writing an ionic formula, how do you name the cation/metal?
Its name stays the same.
96
What is the overall charge on any ionic compound?
Zero/neutral charge.
97
Consider the formula: Mg2S2. What is incorrect?
It can be reduced because there are two of each type of atom. It should say, MgS.
98
What is the difference between a molecular and empirical formula?
A molecular formula tells you exactly how many atoms of each element are in a compound, while empirical formulas reduce the subscripts in a molecular formula by the greatest common factor to simplify it. Ex: C2H4 becomes CH2.
99
A compound formed by a covalent bond is called...
A molecule.
100
A compound formed by an ionic bond is called...
An ion.
101
What is the largest class of molecules?
Organic molecules (which mean they contain carbon).
102
What are ternary ionic compounds?
Compounds that are made of three or more elements ionically bonded.
103
Can you use the cross cross method for compounds that contain polyatomic ions?
Yes.
104
What are polyatomic ions?
Ions made of more than one atom.
105
How do you separate polyatomic ions when part of a compound?
Use parentheses.
106
Some transition metals have more than one oxidation state. How do you show this when writing an ionic formula name?
Put the charge as a Roman numeral in parenthesis after the element name. Ex: Fe 2+ is named iron (II)
107
(SO4)2- is named
Sulfate
108
Covalent formula writing is also called
Molecular formula writing
109
Covalent formulas are also called
Molecular formulas
110
Covalent compounds form...and are made of....
Form molecules and are made of two nonmetals combining to form binary molecular compounds
111
What is a binary compound?
A compound made of two atoms
112
What kind of bond holds molecular compounds together?
Covalent bonds (sharing electrons)
113
Do ionic charges matter in molecular compound naming?
No.
114
How do you indicate the number of atoms of each element in a molecular compound?
Use prefixes
115
Why is it necessary to use prefixes when naming molecular compounds?
Because when two nonmetals combine, they can do it in multiple ways.
116
Name this compound: NO
Nitrogen monoxide
117
Covalent formulas are made of
Molecules
118
Name this compound: N2O5
Dinitrogen pentoxide
119
```
Prefix for 1
2
3
4
5...
```
1: mono
2: di
3: tri
4: tetra
5: penta
6: hexa
7: hepta
8: octa
9: nona
10: deca
120
What is solubility?
The ability to dissolve
121
Water is known as...
The universal solvent
122
In this solution, which is the solvent and which is the solute? Water+KoolAid
Water=solvet, KoolAid=solute
123
Why are substances like copper chloride water-soluble?
They are ionic instead of nonpolar covalent, and water is polar covalent. (Polar covalent also dissolves in water)
124
Why are substances like copper chloride insoluble in oil?
Because oil is nonpolar covalent, which cannot dissolve ionic compounds.
125
Why are substances like soap soluble in both water and oil?
Because one end of a soap molecule is nonpolar and the other is polar, which caters to both oil (nonpolar) and water (polar).
126
What is required for electrons to be exchanged between atoms?
The difference in electronegativity must be greater than 1.7.
127
Nonpolar covalent is also called
Pure covalent
128
Which two elements have the lowest electronegativity?
Cesium and francium
129
Are chargers used to determine formula names for a covalent formula writing?
No.
130
NO Which is the correct name: nitrogen monoxide or mononitrogen monoxide?
Nitrogen monoxide. Don't use the mono prefix for the first element named.
131
N2O Which is the correct name: Dinitrogen oxide or dinitrogen monoxide?
Dinitrogen monoxide. Use "mono" and all other prefixes for the second element named; use all prefixes except for "mono" for the first element in a covalent bond.
