Chemistry 2 - Foundations in Chemistry Flashcards
(159 cards)
1
Q
What was stated in Dalton’s atomic theory?
A
Atoms are tiny particle made of elements.
They cannot be divided.
All the atoms in an element are the same.
Atoms of one element are different to those of other element.
2
Q
What did Thompson discover about electrons?
A
They have a negative charge.
They can be deflected by magnetic and electric fields.
They have very small mass.
3
Q
Explain the plum pudding model.
A
Atoms are made up of negative electrons moving around in a sea of positive charge.
4
Q
What were Rutherford’s proposal after the gold leaf experiment?
A
Most of the mass and positive charge of the atom are in the nucleus.
Electrons orbit the nucleus.
Most of the atom’s volume is the space between the nucleus and the electrons.
Overall positive and negative charges must balance.
5
Q
Explain the current model of the atom.
A
Protons and neutrons are found in the nucleus.
Electrons orbit in shells.
Nucleus is tiny compared to the total volume of the atom.
Most of an atom’s mass is in the nucleus.
Most of the atom is empty space between the nucleus and the electrons.
6
Q
What is the relative charge of a proton?
A
1+
7
Q
What is the relative charge of an electron?
A
1-
8
Q
Which particle has the same mass as a proton?
A
Neutrons.
9
Q
Which two particles make up the majority of an atom’s mass?
A
Protons and neutrons.
10
Q
Which letter is used to represent the atomic number of an atom?
A
Z
11
Q
What does the atomic number tell you about an element?
A
Atomic number = number of protons in an atom.
12
Q
Which letter represents the mass number?
A
A
13
Q
How is mass number calculated?
A
Mass number=number of protons+number of neutrons
14
Q
Define an isotope.
A
Atoms of the same element with different number of neutrons.
15
Q
Why do different isotopes of the same element react in the same way?
A
Neutrons have no impact on the chemical reactivity.
Reaction involve electrons, isotopes have the same number of electrons in the same arrangement.
16
Q
What are ions?
A
Charged particles that are formed when atoms lose or gain electrons.
17
Q
What is the charge of the ion when electrons are gained?
A
Negative.
18
Q
What is the unit used to measure atomic masses called?
A
Unified atomic mass unit, u
19
Q
Define relative atomic mass.
A
The weighted mean mass of an atom of an element compared with one twelfth of the mass of an atom of carbon-12.
20
Q
What is the unit of relative atomic mass?
A
It has no units.
21
Q
Define relative isotopic mass.
A
The mass of an atom of an isotope compared with one twelfth of the mass of an atom of carbon-12.
22
Q
The relative isotopic mass is the same as which number?
A
The mass number.
23
Q
What two assumptions are made when calculating the mass number?
A
The contribution of the electrons is neglected.
The mass of both protons and neutrons is taken as 1.0 u.
24
Q
Describe how to calculate the relative molecular mass and relative formula mass?
A
Both can be calculated by adding the relative atomic masses of each of the atoms making up the molecule or the formula.
25
What are the uses of mass spectrometry?
Identifying unknown compounds.
Finding relative abundance of each isotope of an element.
Determining structural information.
26
How does a mass spectrometer work?
The sample is made into positive ions.
They pass through the apparatus and are separated according to mass to charge ratio.
A computer analyses the data and produces a mass spectrum.
27
How is the group number related to the number of electrons?
Group number = number of electrons in the outer shell
28
What are groups in the periodic table?
Vertical columns in the table, elements in the same group have the same number of electrons in their outer shell, and as such similar chemical properties.
29
Do metals usually gain or lose electrons in reactions?
Lose electrons.
30
Which elements tend to not form ions, why?
Beryllium, boron, carbon and silicon. It requires a lot of energy to transfer outer shell elctrons.
31
What are molecular ions?
Covalently bonded atoms that lose or gain electrons.
32
What is the charge of an ammonium ion?
+1 → NH₄⁺
33
What is the charge of a hydroxide ion?
-1 → OH⁻
34
What is the charge of a nitrate ion?
-1 → NO₃⁻
35
What is the charge of a carbonate ion?
-1 → CO₃²⁻
36
What is the charge of a sulfate ion?
-2 → SO₄²⁻
37
What is an empirical formula?
The simplest whole number ratio of atoms of each element present in a compound.
38
Describe how to calculate empirical formula?
Divide the amount of each element by its molar mass.
Divide the answers by the smallest value obtained.
If there is a decimal, divide by a suitable number to make it into a whole number.
39
When an acid is added to water what ion is released into the solution?
Hydrogen ions, H⁺
40
Define acid.
A proton donor.
41
Describe the dissociation of a strong acid.
Full dissociation.
42
Define base.
A proton acceptor.
43
Which base is often used to treat acid indigestion?
Magnesium hydroxide.
44
What are alkalis?
Bases that can dissolve in water to form aqueous hydroxide ions.
45
What are amphoteric substance?
Substances that can act as acids and bases.
46
What is formed when acid reacts with carbonate?
Salt, carbon dioxide and water.
47
What is a salt?
A compound that is formed when the H⁺ ions of acids are replaced by a metal ion or another positive ion.
48
What is formed when acids react with metal oxides?
Salt and water.
49
What is formed when acid reacts with alkali?
Salt and water.
50
What is formed when acids react with metal?
Salt and hydrogen.
51
Why are the products same when acid reacts with alkali or metal oxides?
Both alkali and metal oxides are types of bases.
52
How are ammonium salts formed?
When acids react with aqueous ammonia.
53
What are hydrated crystals?
A crystalline structure containing water.
54
What does anhydrous crystals mean?
A crystalline form that contains no water.
55
What does a dot formula indicate?
The amount of water present in a crystalline stucture.
56
Write down the methods to carry out a titrations.
Using a pipette, measure the volume of a solution.
Add the solution into a conical flask and add an indicator into it,
Add the other solution into a burette and record the initial volume.
Slowly add the solution in the burette and record the volume.
Slowly add the solution in the burette into the conical flask.
Swirl the mixture continuously until the end point is reached.
Repeat until concordant results are obtained.
57
What is an oxidation number?
The number of electrons an atom uses to bond with any other atoms.
58
What is the oxidation of an uncombined element, such as C, H₂ or O₂?
0
59
What is the oxidation number of combined oxygen, such as in H₂O?
Almost always -2 but there are some cases where it isn't.
60
What is the oxidation number of oxygen in peroxides?
-1
61
What is the oxidation number of combined hydrogen, such as in NH₃ and H₂S?
+1
62
What is the oxidation number of combined hydrogen in metal hydrides, such as in LiH?
-1
63
What is the oxidation number of a simple ion?
The charge on the ion.
64
What is the oxidation number of combined fluorine, such as in NaF and CaF₂?
-1
65
When an element has more than one stable oxidation number, how is it indicated?
It is written as a roman numeral.
66
What is the oxidation number of Fe in iron(III) chloride?
+3
67
What are oxyanions?
Negative ions that have an element along with oxygen.
68
What is the oxidation number of S in SO₄²⁻?
+6
Because combined oxygen has an oxidation number of -2,
4×(-2)=-8
The charge on the compound is -2, so the sum of oxidation numbers must equal -2.
So; -2-(-8)=+6
69
Define oxidation in terms of electron transfer and oxidation number.
Oxidation is;
loss of electrons.
an increase in oxidation number.
70
Define reduction in terms of electron transfer and oxidation number.
Reduction is;
gain of electrons.
a decrease in oxidation number.
71
What is a redox reaction?
A reaction in which bot oxidation and reduction takes place.
72
What is the oxidation number of a metal?
0, because it is an uncombined element.
73
What does the principal quantum number indicate?
The shell occupied by the electrons.
74
What is a shell?
A group of orbitals with the same principal quantum number.
75
How many electrons can the first shell hold?
2
76
How many electrons can the second shell hold?
8
77
How many electrons can the third shell hold?
18
78
How may electrons can the fourth shell hold?
32
79
What is an orbital?
A region around the nucleus that can hold up to two electrons with opposite spins.
80
How many electrons can an orbital hold?
2
81
What are the 4 types of orbitals?
s-orbitals
p-orbitals
d-orbitals
f-orbitals
82
What is the shape of an s-orbital?
Spherical
83
What is the shape of a p-orbital?
Dumb-bell shaped
84
How many orbitals are found in an s-subshell?
1
85
How many electrons can be held in an s-subshell?
2
86
How many orbitals are found a p-subshell?
3
87
How many electrons can be held in a p-subshell?
6
88
How many orbitals are found in a d-subshell?
5
89
How many electrons can be held in a d-subshell?
10
90
How many orbitals are found in an f-subshell?
7
91
How many electrons can be held an f-subshell?
14
92
When using "electrons in box" representation, what shape is used to represent the electrons?
Arrows, electrons are represented as vertical arrows in a box, where their direction is their spin. each box is an orbital, and boxes are grouped to their subshells, which are spaced further apart from other subshell groups.
93
What letter is used to represent shell number?
𝑛
94
From which shell onwards are s-orbitals present?
𝑛=1
95
From which shell onwards are p-orbitals present?
𝑛=2
96
From which shell onwards are d-orbitals present?
𝑛=3
97
From which shell onwards are f-orbitals present?
𝑛=4
98
What are the rules by which electrons are arranged in a shell?
Electrons are added one at a time.
The lowest available energy level (subshell) is filled first.
Each energy level must be filled before the next one can fill.
Each orbital is filled singly before pairing.
4s is filled before 3d.
99
Why does the 4s orbital fill before the 3d orbital?
The 4s orbital has a lower energy than 3d before it is filled.
100
What is the electron configuration of krypton?
1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶
101
Which electrons are lost when an atom becomes a positive ion?
Electrons in the highest energy levels.
102
What are the three main types of chemical bonds?
Ionic, covalent and metallic.
103
Define ionic bonding.
The electrostatic attraction between positive and negative ions.
104
Give an example of an ionically bonded substance.
NaCl (sodium chloride- salt)
| or any other suitable answer
105
Define covalent bonding.
The electrostatic attraction between a shared pair of electrons and their nuclei.
106
Define metallic bonding.
Electrostatic attraction between the positive metal ions and the sea of delocalised electrons.
107
Electrons in which shell are represented in a dot and cross diagram?
The outer shell.
108
Why do giant ionic lattices conduct electricity when liquid but not when solid?
In a solid state, the ions are in fixed positions and thus cannot move. When they are in a liquid state, the ions are mobile and thus can freely carry the charge. This is also why solutions of ionic substances are conductive, since the charges dissociate and can freely carry charge.
109
Do giant ionic lattices have high or low melting and boiling points? Explain why.
They have high melting and boiling point, because a large amount of energy is required to overcome the electrostatic bonds.
110
In what type of solvents do ionic lattices dissolve?
Polar solvents, such as water.
111
Why are ionic compounds soluble in water?
Water has a polar bond. Hydrogen atoms have a δ⁺ charge and oxygen atoms have a δ⁻ charge. These charges are able to attract charged ions.
112
What is it called when atoms are bonded by a single pair of shared electrons?
A single bond.
113
How many covalent bonds does carbon form?
4
114
How many covalent bonds does oxygen form?
2
115
What is a lone pair?
Electrons in the outer shell that are not involved in bonding.
116
What is formed when atoms share two pairs of electrons?
A double bond.
117
What is formed when atoms share three pairs of electrons?
A triple bond.
118
What is average bond enthalpy?
A measure of the average energy needed to break the bond.
119
What is a dative covalent bond?
A bond where both of the shared electrons are supplied by one atom.
120
How are oxonium ions formed?
Formed when acid is added to water, H₃O⁺.
121
What does expansion of the octet mean?
When a bonded atom has more than 8 electrons in the outer shell.
122
What are the types of covalent structure?
Simple molecular and giant covalent lattice.
123
Describe the bonding in simple molecular structures.
Atoms within the same molecule are held by strong covalent bonds and different molecules are held by weak intermolecular forces.
124
Do simple molecular structures have high or low melting and boiling points? Explain why.
They have low melting and boiling points, this is because a small amount of energy is enough to overcome the intermolecular forces.
125
Can simple molecular structures conduct electricity?
No, they are non conductors.
126
Why don't simple molecular structures conduct electricity?
They have no free charged particles to move around.
127
Simple molecular structures dissolve in what type of solvent?
Non-polar solvents.
128
Give examples of giant covalent structures.
Diamond, graphite and silicon dioxide (SiO₂)
129
List some properties of giant covalent structures?
High melting and boiling points.
Non-conductors of electricity, except graphite.
Insoluble in polar and non-polar solvents
130
How does graphite conduct electricity?
Delocalised electrons present between the layers are able to move freely, carrying the charge.
131
Why do giant covalent structures have high melting and boiling points?
There are many strong covalent bonds within the molecules, which need to be broken, which requires a lot of energy.
132
Draw and describe the structure of a diamond.
A 3d tetrahedral structure of C atoms, with each C atom bonded to four others.
133
What does the shape of a molecule depend on?
The number of electron pairs in the outer shell.
| The number of these electrons which are bonded and lone pairs.
134
What is the shape and bond angle in a shape with 2 bonded pairs and 0 lone pairs?
Linear
| 180°
135
What is the shape and bond angle in a shape with 3 bonding pairs and 0 lone pairs?
Trigonal planar
| 120°
136
What is the shape and bond angle in a shape with 4 bonded pairs and 0 lone pairs?
Tetrahedral
| 109.5°
137
What is the shape and bond angle in a shape with 5 bonded pairs and 0 lone pairs?
Trigonal bipyramid
| 90° and 120°
138
What is the shape and bond angle in a shape with 6 bonded pairs and 0 lone pairs?
Octahedral
| 90°
139
What is the shape and bond angle in a shape with 3 bonded pairs and 1 lone pair?
Pyramidal
| 107°
140
What is the shape and bond angle in a shape with 2 bonded pairs and 2 lone pairs?
Non-linear
| 104.5°
141
By how many degrees does each lone pair reduce the bond angle?
2.5°
142
Define electronegativity.
The ability of an atom to attract the pair of electrons (the electron density) in a covalent bond.
143
In which direction of the periodic table does electronegativity increase?
Top right, towards fluorine.
144
What does it mean when the bond is non-polar?
The electrons in the bond are evenly distributed.
145
What is the most electronegative element?
Fluorine.
146
How is a polar bond formed?
Bonding atoms have different electronegativities.
147
Why is H₂O polar, whereas CO₂ is non-polar?
CO₂ is a symmetrical molecule, so there is no overall dipole.
148
What is meant by intermolecular force?
The attractive force between neighbouring molecules.
149
What are the 2 types of intermolecular forces?
Hydrogen bonding and Van der Walls' forces.
150
What is the strongest type of intermolecular force?
Hydrogen bonding.
151
What are the 2 interactions that can be referred as Van der Waals' forces?
Permanent dipole-induced dipole interaction
| Permanent dipole-permanent dipole interaction
152
Describe permenant dipole-induced dipole interactions.
When a molecule with a permanent dipole is close to other non-polar molecules it causes the non-polar molecules to become slightly polar, leading to attraction.
153
Describe permanent dipole-permanent dipole interactions.
Some molecules with polar bonds have permanent dipoles, permanent dipole-permanent dipole interactions are the forces of attraction between those dipoles and those of neighbouring molecules.
154
Describe London forces.
London forces are caused by random movements of electrons.
This leads to instantaneous dipoles.
An instantaneous dipole induces a dipole in nearby molecules.
The induced dipoles attract one another.
155
Are London forces greater in smaller or larger molecules?
Larger due to the higher number of electrons.
156
Does boiling point increase or decrease down the noble gas group? Why?
Boiling point increases down the group, because the number of electrons increases and hence the strength of London forces also increases.
157
What conditions are needed for hydrogen bonding to occur?
O-H, N-H or an F-H bond and a lone of electrons on O, F or N.
Because O, N and F are highly electronegative, the H nucleus is left exposed.
The hydrogen bonds are the strong forces of attraction between the H nucleus and a lone pair of electrons on O, N or F.
158
Why is ice less dense than liquid water?
In ice, the water molecules are arranged in an orderly pattern. It has an open lattice with hydrogen.
In water, this lattice is collapsed and so the molecules are closer together.
159
Why does water have a higher melting and boiling point than expected?
Hydrogen bonds are stronger than other intermolecular forces, so extra strength is required to overcome these forces.
