Topic 2: Microscopic World I Flashcards
(110 cards)
1
Q
List the first 20 elements in the Periodic Table.
A
Hydrogen (H), Helium (He), Lithium (Li), Beryllium (Be), Boron (B), Carbon (C), Nitrogen (N), Oxygen (O), Fluorine (F), Neon (Ne), Sodium (Na), Magnesium (Mg), Aluminium (Al), Silicon (Si), Phosphorus (P), Sulphur (S), Chlorine (Cl), Argon (Ar), Potassium (K), Calcium (Ca)
2
Q
What are the names of the following elements:
Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Br, Ag, I, Ba, Pt, Au, Hg, Pb?
A
Titanium, Vanadium, Chromium, Manganese, Iron, Cobalt, Nickel, Copper, Zinc, Bromine, Silver, Iodine, Barium, Platinum, Gold, Mercury, Lead
3
Q
Definition of an atom.
A
An atom is the fundamental particle of an element, it cannot be split into simpler particles by chemical methods.
4
Q
What are the three kinds of subatomic particles in an atom?
A
An atom consists of three kinds of subatomic particles: proton(p⁺), neutron(n) and electron(e⁻).
5
Q
What are the properties of subatomic particles?
^inlcude: Symbol, Relative charge, Relative mass, Position in the atom
A
Proton (p⁺), +1, 1, inside the nucleus
Neutron (n), 0, 1, inside the nucleus
Electron (e⁻), -1, ~0, around the nucleus
6
Q
Where does nucleus locate in an atom?
State its exception.
^include: What does a nucleus consist of?
A
Nucleus, which consists of proton(s) and neutron(s), is at the centre of an atom, except hydrogen-1(1H) which does not have neutron.
7
Q
Describe the movement of e⁻ in an atom.
A
e⁻ moves rapidly in different electron shells surrounding the nucleus.
8
Q
Why is most volume of an atom empty?
A
Most volume of an atom is empty
∵ Mass of an atom is concentrated at the nucleus, which is located at the center of the atom.
9
Q
Why is an atom electrically neutral?
A
An atom is electrically neutral
∵ It has the same no. of p⁺ and e⁻.
10
Q
What does the atomic number of an element equal to?
A
No. of p⁺ = No. of e⁻ in an atom
11
Q
What does the mass number of an element equal to?
A
No. of p⁺ + No. of n
12
Q
How to define whether atoms have the same atomic number or not?
A
Atoms of the same element have the same atomic number while atoms of different elements have different atomic numbers.
13
Q
How are mass number, atomic number and symbol of an element represented?
A
Mass number: upper left-handed corner
Atomic number: lower left-handed corner
Symbol: right-hand side
14
Q
Definition of isotopes.
A
Isotopes are atoms of the same element which have the same number of protons but different number of neutrons (same atomic number, different mass number).
Isotopes of elements occur in nature at different abundances(蘊藏量).
15
Q
Definition of relative abundance (or natural abundance / % abundance).
A
Relative abundance / natural abundance / % abundance is the distribution of isotopes in nature of an element.
16
Q
Why do isotopes have different physical properties, but same chemical properties?
A
Isotopes have different physical properties (e.g. melting point, boiling point) ∵ they have different masses.
Isotopes have same chemical properties ∵ they are of the same element having the same no. of e⁻ (same electronic arrangement).
17
Q
How do we measure the masses of atoms? Why?
A
We use a relative scale to measure the masses of atoms ∵ the mass of an atom is very small which causes inconvenience (e.g. the mass of a C atom is only 2 x 10^-23 g).
- The mass of a 12C atom is used as the standard for comparisons of the masses of different atoms.
18
Q
Definition of relative isotopic mass (R.I.M.).
A
Relative isotopic mass of an isotope of an element is the mass of the atom on the 12C = 12.00 scale.
Relative isotopic mass = Mass of an atom of the element / [(1/12) x Mass of an atom of carbon-12 = 1.67 x 10^-24]
- The mass of one carbon-12 (12C) atom is taken as exactly 12 units. That is, the mass of a 12,6C atom is defined to have a relative mass of 12.00.
19
Q
What are the similaries and differences between mass number and R.I.M.?
^include: Similarities and Origins
A
The mass number and R.I.M. are almost the same.
Both the values and the origins of the mass number and the relative isotopic mass are different.
- The relative isotopic mass of an atom is determined by experiment using advanced instruments.
(一定是小數) - The mass number of an atom is obtained from counting the no. of p⁺ and n present in the atom.
(一定是整數)
20
Q
Definition of relative atomic mass (R.A.M.).
A
Relative atomic mass of an element is the average isotopic mass of all the natural isotopes (or simply the average mass of an atom of the element) on the 12C = 12.00 scale.
- Definition
= Average mass of an atom of the element / [(1/12) x Mass of an atom of carbon-12 = 1.67 x 10^-24] - For calculation
= (a%)Ma + (b%)Mb + (c%)Mc + …
a%: % abundance of isotope A
Ma: relative isotopic mass or mass no. of isotope A
21
Q
What are the uses of electron shells in atoms?
A
Atoms have electron shells for the accomodation of e⁻ which move around the nucleus.
22
Q
Definition of electronic arrangement (or electronic configuration).
A
Electronic arrangement is the distribution of e⁻ in different electron shells.
- Each electron shell is given a number 1, 2, 3, 4 and so on, starting from the one closest to the nucleus.
23
Q
What are the rules for writing electron arrangement?
A
- Each shell has the maximum no. of e⁻: 2n^2
3rd electron shell can hold up to 18 e⁻ (but usually hold 8 e⁻). - e⁻ are filled from the innermost electron shell to the outermost electron shell (filled from 1st electron shell).
24
Q
What are the electronic arrangements of helium (2), nitrogen (7), aluminium (13), sulphur (16), potassium (19), calcium (20)?
() atomic number, special cases
A
Helium: 2; Nitrogen: 2,5; Aluminium: 2,8,3; Sulphur: 2,8,6; Potassium: 2,8,8,1; Calcium: 2,8,8,2
25
How are elements arranged in the Periodic Table?
In the Periodic Table, elements are arranged in order of **increasing atomic number (no. of protons)**, not relative atomic mass.
26
Definition of period and group in the Periodic Table.
Period is the **horizontal row** in the Periodic Table
Group is the **vertical column** in the Periodic Table
27
What are the similarities and differences between elements in the same period?
Elements in the same period have **same no. of occupied electron shells**. However, they have **different no. of e⁻ in the outermost shell**, thus they have **different chemical properties**.
28
What are the similarities and differences between elements in the same group?
| ^except helium
Elements in the same group (except helium) have **same no. of e⁻ in the outermost shell**, thus they have **similar chemical properties** (e.g. noble gases do not have reactions; alkali metals can react with water vigorously). However, they have **different no. of occupied electron shells**.
29
What do the chemical properties of an element depend on?
The chemical properties (not physical properties) of an element depend on **the no. of e⁻ in the outermost shell** (not the no. of occupied e⁻ shells in its atoms).
30
How are metals, transitional metals, non-metals and metalloids distributed in the Periodic Table?
Across a period, from left to right, elements change from metals to metalloids and then to non-metals (**metallic character decreases from left to right)**.
* Metals: on the lower left-handed corner
* Transitional metals: in the middle
* Non-metals: on the upper right-handed corner
* Metalloids (or semi-metals): on the diagonal between metals and non-metals
31
What are the family names of elements in different groups?
Group I: Alkali metals (not alkaline metals)
Group II: Alkaline earth metals
Group VII: Halogens
Group 0: Noble gases
^Hydrogen does not belong to any group.
32
How do properties of elements change down a group?
1. Atomic size
* The size of atoms ↑ when going down a group ∵
no. of **occupied e⁻ shells ↑** down a group
2. Melting point (m.p.) and boiling point (b.p.)
* For Group I and II elements (metals), the m.p. and b.p. ↓ down the group ∵ **Atomic size ↑ down a group**, the attraction between the cations and the delocalized e⁻ in the metal ↓, this leads to **weaker metallic bond**.
* For Group VII and 0 elements (non-metals), the m.p. and b.p. ↑ down a group ∵ Molecular size ↑ down a group, this leads to **increased strength of van der Waals' force**.
3. Ability to attract e⁻
* The ability of an atom to attract incoming e⁻ ↓ down a group ∵ The **no. of occupied e⁻ shells ↑ down a group**, this leads to **weaker attraction between the incoming e⁻ and the nucleus** (as a result of longer distance).
33
What do the period and group no. of an element equal to?
The period no. of an element equals its **no. of occupied electron shells** (e.g. Cl is period 3 element, it has 3 occupied e⁻ shells).
The group no. of an element equals its **no. of outermost shell electons**.
34
How do properties of elements change across a period?
1. Atomic size
* The size of atoms ↓ when going across the period from left to right ∵ The **no. of p⁺ (or nuclear charge) ↑** across a period, this causes **increased attraction between e⁻ and the nucleus**.
2. Melting point (m.p.) and boiling point (b.p.)
* Increases from Group I to Group III ∵
They are metals and **metallic bond strength ↑** with the no. of outermost e⁻ involved in the metallic bond.
* Maximum at group IV ∵
C and Si have **giant convalent structures**, much energy is needed to break the **covalent bonds**.
* Low for Group V to Vll ∵
These elements usually have **simple molecular structures**, **weak van der Waals' forces** exist between molecules, only a small amount of energy is needed to break these forces.
* Very low for noble gases ∵
They are all **monoatomic** and thus have very small molecules sizes, **very weak van der Waals' forces** exist between molecules, only a very small amount of energy is needed to break these forces.
3. Ability to attract e⁻
* The ability of an atom to attract incoming e⁻ ↑ across a period ∵ The **no. of p⁺ (or nuclear charge) ↑** across a period, this causes **increased attraction between the incoming e⁻ and the nucleus**.
35
What are the properties of Group I metals / alkali metals?
* They are **soft** and can be cut by a knife ∵
They have weak metallic bonds.
* They are **shiny** when freshly cut, and **tarnish (失去光澤) easily** in air (∵ Reaction with oxygen gas).
* They have **low densities**: Li, Na and K float on water ∵
They have lower densities than water.
* They have **low m.p. and b.p**. compared to other metals (∵ They have weaker metallic bonds).
* They **react with water vigorously** to give hydrogen gas and alkaline solutions.
* The **reactivity** of the metals **↑ down the group** ∵
They react by losing outermost e⁻ and the strength of attraction between nucleus and the outermost e⁻ decreases down the group ∵ **Atomic size ↑ down the group**.
36
What are the electronic arrangement and no. of outermost shell e⁻ of Group I metals?
Electronic arrangement:
* Lithium (Li) - 2,1; Sodium (Na) - 2,8,1; Potassium (K) - 2,8,8,1; Rubidium (Rb) - 2,8,18,8,1; Caesium (Cs) - 2,8,18,18,8,1
No. of outermost shell e⁻ of Group I metals: 1
37
Describe the atomic size and reactivity of Group I metals.
Atomic size: ↑ down the group
Reactivity: ↑ down the group
38
What are observations when Group I metals react with water?
* Lithium (Li): The metal floats; Hydrogen gas evolves.
* Sodium (Na), Potassium (K): The metal floats; Hydrogen gas evolves; It moves about quickly on the water surface.
* Rubidium (Rb): The reaction is more vigorous than potassium.
* Caesium (Cs): The reaction is more vigorous than rubidium.
39
What are the properties of Group II metals / alkaline earth metals?
* They are **soft** but harder than Group I metals ∵
They have stronger metallic bonds than Group I metals.
* They are **shiny** when freshly cut, and **tarnish (失去光澤) easily** in air (∵ Reaction with oxygen gas).
* They **react with water** to give hydrogen gas and alkaline solutions.
* The reactivity of the metals ↑ down the group ∵ They react by losing outermost e⁻ and the strength of attraction between nucleus and the outermost e⁻ decreases down the group ∵ **Atomic size ↑ down the group**.
* They are **less reactive than Group I metals** ∵
They need to lose 2 outermost e⁻ during reactions while Group I metals need to lose 1 outermost e⁻ only.
40
What are the electronic arrangement and no. of outermost shell e⁻ of Group II metals?
Electronic arrangement:
* Beryllium (Be) - 2,2; Magnesium (Mg) - 2,8,2; Calcium (Ca) - 2,8,8,2; Strontium (Sr) - 2,8,18,8,2; Barium (Ba) - 2,8,18,18,8,2
No. of outermost shell e⁻ of Group II metals: 2
41
Describe the atomic size and reactivity of Group II metals.
Atomic size: ↑ down the group
Reactivity: ↑ down the group
42
What are the observations when Group II metals react with water?
* Beryllium (Be): /
* Magnesium (Mg): The metal sinks. Hydrogen gas evolves. (Heating is needed for Mg.)
* Calcium (Ca): The metal sinks. Hydrogen gas evolves.
* Strontium (Sr): The reaction is more vigorous than calcium.
* Barium (Ba): The reaction is more vigorous than strontium.
43
What are the properties of Group VII metals / halogens?
* They are **toxic**.
* They are **coloured**.
* Their **m.p. and b.p. ↑** down the group ∵
Molecular size ↑ down a group, this leads to increased strength of van der Waals' force.
* The **reactivity of the elements ↓** down the group ∵
The atomic size ↑ down the group, thus the ability of an atom to attract incoming e⁻ ↓ down a group.
44
What are the electronic arrangement and no. of outermost shell e⁻ of Group VII metals?
Electronic arrangement:
* Fluorine (F) - 2,7; Chlorine (Cl) - 2,8,7; Bromine (Br) - 2,8,18,7; Iodine (I) - 2,8,18,18,7; Astatine (At) - 2,8,18,32,18,7
No. of outermost shell e⁻ of Group VII metals: 7
45
Describe the atomic size and reactivity of Group VII metals.
Atomic size: ↑ down the group
Reactivity: ↓ down the group
46
What are the appearance of Group VII metals?
* Fluorine (F): Pale yellow gas
* Chlorine (Cl): Yellowish green gas
* Bromine (Br): Brown liquid
* Iodine (I): Purple solid
* Astatine (At): Black solid
47
What are the properties of Group 0 metals / noble gases?
* They are **colourless gases**.
* They are very stable / **unreactive** / inert (but the reactivity increases down the group 唔使識) ∵
They have **completely filled outmost shells**
(i.e. The stability of the noble gases can be explained by **duplet & octet rule**)
* The density of noble gas ↑ down the group.
Note: He and Ne have densities lower than air; Ar, Kr and Xe have densities higher than air.
48
What are the electronic arrangement and no. of outermost shell e⁻ of Group 0 metals?
Electronic arrangement:
* Helium (He) - 2; Neon (Ne) - 2,8; Argon (Ar) - 2,8,8; Krypton (Kr) - 2,8,18,8; Xenon (Xe) - 2,8,18,18,8
No. of outermost shell e⁻ of Group 0 metals:
* 8 (except He, which has 2 only)
49
What are the atomic size, reactivity and density of Group VII metals?
Atomic size: ↑ down the group
Reactivity: ↑ down the group
Density: ↑ down the group
50
Which noble gases can react with F₂?
Down the group, the element (e.g. Kr, Xe) can react with fluorine under special conditions, but He, Ne and Ar cannot.
51
Why are all noble gases unreactive?
| Include: Structures of noble gases
All noble gases are unreactive / stable / chemically inert ∵
They have **completely filled outermost shells**.
Noble gases are said to have **octet structures**, having 8 e⁻ in their outermost e⁻ shell. Helium is an exception that has a **duplet structure**, having only 2 e⁻ in its outermost e⁻ shell.
52
How do elements attain stable electronic structure?
Elements other than noble gases tend to **attain stable electronic structure of nearest noble gases** by:
1. losing outermost shell e⁻ to form positive ions / cations (for metals only)
2. gaining e⁻ to form negative ions / anions (for non-metals only)
3. sharing e⁻ with others (usually between non-metal atoms)
In most cases, the atoms in a stable chemical species (an element or a compound) have the **same electronic structures as the noble gases**.
53
Draw the electronic diagrams of the formation of metal ions.
| Na → Na+; O → O²⁻
Refer to notes p.33 (part 1).
54
What are the relationship between group number and charge of ion?
| ^Exceptions
Group I, II and III metals have ions with charges equal to their group numbers.
Group V, VI and VII elements have ions with charges equal to (group number - 8).
Exceptions:
* **IV: C and Si do not form ions** ∵
Too many e⁻ (four) have to be removed or ganied to attain a stable noble gas structure.
* **0: They do not form ions** ∵
They have completely filled outermost shells (they are already stable).
55
Definition of a simple ion.
A simple ion is an ion **formed from a single atom** (by losing or gaining e⁻).
56
Definition of polyatomic ions.
Polyatomic ions are ions formed from more than one atom.
e.g. **+1**: NH₄⁺ // **-1**: OH⁻, NO₂⁻, NO₃⁻, HCO₃⁻, HSO₄⁻, H₂PO₄⁻, MnO₄⁻ // **-2**: SO₃²⁻, SO₄²⁻, HPO₄²⁻, CO₃²⁻, CrO₄²⁻, Cr₂O₇²⁻ // **-3**: PO₄³⁻
57
List the common cations and anions.
Refer to notes p.36 (part 1).
58
What are halides?
**F⁻, Cl⁻, Br⁻ and I⁻ (anions formed from halogens)** are called halides.
59
What are metal-containing polyatomic anions?
**MnO₄⁻, CrO₄²⁻ and Cr₂O₇²⁻** are metal-containing polyatomic anions.
60
What are the colours of ions?
Most ions are **colourless**, e.g. Group I, II and III metal ions, O²⁻, F⁻, Cl⁻, Br⁻, I⁻, OH⁻, NO₃⁻, SO₄²⁻, PO₄³⁻.
Some ions are **coloured**, especially for those of transition metals / containing transition metals.
(Name of ion, Symbol, Colour in aqueous solution)
* Iron(II) ion, Fe²⁺, Green
* Iron(III) ion, Fe³⁺, Yellow (or brown)
* Cobalt(II) ion, Co²⁺, Blue / pink
* Nickel(II) ion, Ni²⁺, Green
* Copper(II) ion, Cu²⁺, Blue / greenish blue
* Chromium(III) ion, Cr³⁺, Green
* Chromate ion, CrO₄²⁻, Yellow
* Dichromate ion, Cr₂O₇²⁻, Orange
* Manganese(II) ion, Mn²⁺, Colourless
* Permanganate ion, MnO₄⁻, Purple
61
How are substances classified in accordance with bondings and structures?
Structures →
1. **Metallic bond → Griant metallic structure**
All metals, e.g. Li, Na, K, Mg, Ca, Al, Zn, Fe ...
2. **Ionic bond → Giant ionic structure**
All ionic compounds, e.g. NaCl, CaF₂, Al₂O₃, CuSO₄, FeCl₂, NH₄Cl, (NH₄)₂SO₄
3. **Covalent bond → Giant covalent structure, Simple molecular structure, Macromolecular structure**
(a non-metal element / a covalent compound)
**GCS**: Diamond, Graphite, Silicon, Silicon dioxide (quartz, SiO₂), Boron;
**SMS**: Noble gases (no bonds) and most covalent substances except those with giant covalent structures), e.g. He, Ar, O₂, I₂, P₄, S₈, CO₂, CH₄, C₂H₄)
Special case:
* Some substances **contain both covalent bonds and ionic bonds**, e.g. NH₄Cl, in which the bonding between N and H is covalent bond, while that between NH₄⁺ and Cl⁻ is ionic bond.
62
How are e⁻ tranferred within atoms when a metal react with a non-metal?
When a metal reacts with a non-metal,
* atoms of the **metal donate e⁻** to atoms of the non-metal to form **cations**
* atoms of the **non-metal accept e⁻** to form **anions**
→ transfer of e⁻ from metal atoms to non-metal atoms.
63
How are cations and anions formed?
Both cations and anions attain **the nearest noble gas electronic structures**.
* Metal atoms (Group I to III) usually lose 1 to 3 e⁻ to form cations.
* Non-metal atoms (Group V to VII metals) usually gain 1 to 3 e⁻ to form anions.
64
Definition of ionic bond.
| ^Where can it be found? // Exceptions // Bonding
The cation and anion are held together by strong electrostatic force / attraction called ionic bond. It usually formed between **a metal and a non-metal**
Exceptions:
* **Ammonium compounds** are ionic compounds, which do not contain metals, e.g. NH₄Cl (ammonium chloride), (NH₄)₂SO₄ (ammonium sulphate).
The compound formed is called an **ionic compound**.
* Ionic bond is **non-directional** in nature.
i.e. Cations and anions attract each other **in all directions** in both solid and molten (liquid) states.
65
Draw the electronic diagrams of the formation of ionic compounds.
| NaCl, MgF₂, K₂O
Refer to notes p.39 (part 1).
66
What are steps to predict the chemical formula of ionic compounds?
Refer to notes p.42 (part 1).
1. Write down the symbols of ions in the compound.
2. Write down the charges of each ion on the top of each symbol.
3. Cross-multiply the numbers.
4. Combine the symbols and symplify the ratio if necessary.
67
What are the precautions in writing a chemical formula?
* Total positive charge must be equal to the total negative charge ∵ An ionic compound is electrically neutral.
* In all the cases, the symbol of the cation is written first, followed by the symbol of the anion.
Refer to notes p.42 (part 1).
68
What are the rules in naming ionic compounds?
* Name the cation first, followed by the anion.
* If the metal ion exist in more than one form (usually transition metals), the charge is indicated in a bracket after the name of the metal.
* Simple anion end with 'ide'. Polyatomic anion end with 'ite' or 'ate' (except hydroxide).
Refer to notes p.43 (part 1).
69
Describe + Draw the structure of ionic compounds (Giant ionic structure).
| Sodium chloride crystal - NaCl (s); Caesium chloride crystal - CsCl (s)
Structure: Refer to notes p.7 (part 2)
All ionic compounds have giant ionic structures.
In a giant ionic structure,
* The **oppositely charged ions** in an ionic compound are held together regularly by strong electrostatic forces (ionic bonds) to form **a giant lattice**.
70
What are the properties of ionic compounds?
Refer to notes p.10 (part 2).
* Ionic compounds are usually in **crystalline (結晶) form**, having repeating arrangement of ions.
* Ionic compounds have **high melting points and boiling points** ∵ The ions are held by strong ionic bonds. Much energy is needed to break these ionic bonds.
* In solid state, ionic compounds do not conduct electricity ∵ The ions are not mobile / free to move.
In **molten state (liquid state) or aqueous state (dissolved in water)**, the ions become mobile, so they can conduct electricity.
* Ionic compound is **hard** ∵ Ions are held together closely by strong ionic bonds and relative motion of the ions is restricted.
* Ionic compounds are **brittle** ∵ Ionic bonds exist between cations and anions (1M). When a force is applied, ions of the same charge may come together and cause repulsion, the solid thus breaks apart (1M).
* Ionic compounds usually have **high density** ∵ Ions are held together closely by strong ionic bond.
* Ionic compounds may contain **water of crystallisation**, e.g. CuSO₄·5H₂O(s), CoCl₂·6H₂O(s).
→ The solubility of CuSO₄ in water decreases with decreasing temp. Thus, CuSO₄·5H₂O(s) crystals are formed upon cooling.
* In general, ionic compounds are **soluble in water** (refer to solubility table), but **insoluble in non-aqueous solvents**, such as hexane, tetrachloromethane (CCl₄) and benzene.
71
List the general solubility rule of ionic compounds in water.
Refer to notes p.11 (part 2).
72
How do atoms of non-metals combine?
When atoms of non-metals combine, they tend to **share their outermost shell e⁻** to form octet structures (or duplet structures). There is **no e⁻ transfer involved** (which occurs in ionic bond formation).
73
How are covalent compounds formed?
| ^Exception
Atoms in covalent compounds attain **noble gas electronic structures (octet or duplet)** in most cases.
Exception:
* BF₃ (B in BF₃ has only 6 outermost e⁻, Topic 7)
74
Definition of covalent bond.
| ^Where can it be found?// Exceptions // Bonding
The atoms of non-metals are held by strong electrostatic forces between **shared e⁻ and the positive nuclei of bonded atoms**. The force / attraction is called covalent bond.
It is formed between **non-metal atoms**.
The compound formed is called a **covalent compound**.
75
What are the types of substances formed by covalent bond?
| ^Element, compound or mixture
The substance formed can be an **element** if 2 (or more) atoms of the same element combine.
* e.g. H₂, O₂, N₂, Cl₂, Br₂, P₄ and S₈
The substance formed can also be a **compound** if atoms of different elements combine.
* e.g. H₂O, NH₃, CH₄ and CO₂
76
Definition of a bond pair.
A bond pair is **a shared electron pair**.
* One shared electron pair between two atoms forms a **single bond**.
* Two shared electron pairs between two atoms forms a **double bond**.
* Three shared electron pairs between two atoms forms a **triple bond**.
77
Definition of a lone pair.
A lone pair is **an unshared electron pair** in the **outermost shell** of an atom.
78
Draw the electronic diagrams of covalent compounds.
| Cl₂, NH₃
Refer to notes p.54 (part 1).
79
What are steps to predict the molecular formulae of covalent compounds?
Refer to notes p.55 (part 1).
1. Write down the electronic arrangment of the atoms involved.
2. Write down the no. of e⁻ each atom needs to attain a noble gas electronic arrangement.
3. Cross multiply the numbers.
4. Combine the symbols and simplify the ratio if necessary.
80
How do noble gases exist? Why?
Noble gases are **monoatomic** (exist as individual atoms) ∵ They have **completely filled outermost shells**. An individual noble gas atom can be called a monoatomic molecule.
81
What are the rules in naming covalent compounds?
* The more electropositive element (further left of the periodic table) is listed before the more electronegative element (further right of the periodic table).
Note: **Hydrogen** is squeezed between nitrogen and oxygen.
* The second element is given an **-ide** ending.
* Prefixes are used to denote how many atoms of each element are present in the compound, i.e. **1: mono-; 2: di-; 3: tri-; 4: tetra-; 5: penta-; 6: hexa-**.
Note: The mono- prefix is usually not used for the first element in the formula; The "o" and "a" endings of these prefixes are dropped when they are attached to "oxide".
82
What are the common names of
H₂O, NH₃ and CH₄?
(Formula, Common name, Systematic name)
H₂O: Water, Dihydrogen monoxide
NH₃: Ammonia, Trihydrogen nitride
CH₄: Methane, Tetrahydrogen carbide
83
Definition of dative covalent bond.
Dative covalent bond (dative bond / coordinate bond) is a special case of **covalent bond**, in which **both e⁻ in the single bond** are donated from the **same atom**.
84
NH₃ reacts with BF₃ to give H₃NBF₃. Describe the bond formation between BF₃ and NH₃.
* In BF₃, there are **three bonding electron pairs** / there is **a vacant site** in the outermost shell of the B atom. (1M)
('BF₃ is an electron-deficient species in which B has only 6 outermost e⁻'.)
* By accepting **the lone pair of e⁻** from the N atom of NH₃ / **forming dative bond** with N, the B atom attains the stable electronic configuration of neon (a noble gas). (1M)
85
Describe the formation of dative covalent bond using H₃O⁺ as example.
* The O atom in H₂O has **lone pairs of e⁻**. (1M)
* H⁺ does not have e⁻ in its outermost shell. (1M)
* **Dative covalent bond** formed between the O atom in H₂O and H⁺ by sharing electron pair. (1M)
(Also accept graphical answer)
86
Draw the electronic diagrams of covalent compounds with dative covalent bond.
| NH₄⁺ (ammonium ion), H₃O⁺ (hydronium ion), H₃NBF₃
Refer to notes p.66 (part 1).
87
Describe the structure of covalent substances
(Giant covalent structure).
The atoms (millions or more) are held together by strong **covalent bonds** to form **a giant network**.
Notes:
* Covalent substances having giant covalent structures are not regarded as molecular substances (they **do not contain molecules**).
* The term 'molecule' can only be used to describe substances with simple molecular structures.
88
Describe + Draw the structure of diamond (a form of carbon)
Structure: Refer to notes p.15 (part 2)
* Electronic arrangement of C = 2,4 (2 occupied electron shells and 4 outermost e⁻)
* Each C atom can form a maximum of four single bonds with other atoms, using **all of its outermost e⁻**.
* Each C atom in diamond forms **strong covalent bonds** with **four** other C atoms.
* Each C atom in diamond is **tetrahedrally** bonded to four other C atoms (all bond angles = 109.5°).
89
Describe + Draw the structure of graphite (a form of carbon).
Structure: Refer to notes p.15 (part 2)
* Graphite has a **layered structure**.
* C atoms **within the same layer** are held together by **strong covalent bonds**. Each C atom forms strong covalent bonds with **three** other C atoms **using three of its four outermost e⁻**.
* **Van der Waals' forces** (a type of intermolecular force / attraction) exist **between the layers**.
90
(?) Why does graphite have delocalized e⁻, but diamond has not?
Diamond form four bonds using all of its outermost shell e⁻, while graphite **only forms three bonds using three of its four outermost shell e⁻**. The remaining one electron acts as delocalized electron, and thus graphite can conduct electricity.
91
What are the applications of graphite?
1. **Making dry cells** (as electrodes 電極) ∵
Good electrical conductivity due to presence of delocalized e⁻ (not related to v.d.w. force).
2. **Making pencils** ∵
Graphite layers are only attracted by weak van der Waals' forces and graphite layers flake off easily onto the paper → good marking properties.
3. As a **solid lubricant** (潤滑劑) where 'wet' lubricant (e.g. oil) cannot be used ∵
Graphite layers are only attracted by weak van der Waals' forces so that the layers can slide past each other easily.
92
Describe + Draw the structure of silicon.
Structure: Refer to notes p.15 (part 2)
* The structure of silicon is the same as diamond.
* Each Si atom forms strong covalent bonds with **four** other Si atoms using all the four outermost e⁻.
93
Describe + Draw the structure of silicon dioxide (SiO₂, quartz).
Structure: Refer to notes p.16 (part 2)
* SiO₂ is a compound, not an element.
Note: Diamond, graphite and silicon are elements.
* Ratio of Si atom : O atom = 1:2
Each silicon atom is bonded to **four** oxygen atoms.
Each oxygen atom is bonded to **two** silicon atoms.
94
Draw the structure of silicon carbide.
Refer to notes p.16 (part 2).
95
What are the properties of substances with giant covalent structure?
* They have **very high m.p. and b.p.** ∵
The atoms are held togther by strong covalent bonds.
During melting / boiling, most / all of the strong covalent bonds are needed to be broken, which requires a lot of heat energy.
* They **do not conduct electricity** ∵ There are neither mobile ions nor delocalized e⁻.
Exception: Graphite can conduct electricity ∵ It has delocalized e⁻.
* They are very **hard** ∵ Atoms are held together by strong covalent bonds to form a giant network. Relative motion of atoms is restricted.
Exception: Graphite is soft ∵ It has a layered structure with weak van der Waals' forces between the layers. This allows the layers to slide past each other.
* They are **insoluble in all solvents** ∵ Breaking the strong covalent bonds requires a lot of energy.
96
Why diamond (m.p. = 3500°C) has a higher melting point than iron (m.p. = 1538°C)?
Diamond: break **most** of covalent bonds (strong)
Iron: break **some** of metallic bonds (strong)
Wrong answer:
It is because a covalent bond is stronger than a metallic bond.
97
Describe + Draw the structure of covalent substances (Simple molecular structure).
Structure: Refer to notes p.18,19 (part 2); Solid + Liquid state
Covalent substances having simple molecular structures are regarded as molecular substances (they are consist of molecules).
* They exist as **separate / discrete molecules**.
* In most cases, atoms within the same molecules are held together by **strong covalent bonds**. Separate molecules are held together by **intermolecular forces**.
Remarks:
* **Noble gases** are consist of **monoatomic molecules** (i.e. atom = molecule).
* Van der Waals' force is a type of intermolecular forces.
98
Describe the strength of intermolecular force compare to other bonds.
Intermolecular force is a **much weaker attractive force** compared to covalent bond (also ionic bond and metallic bond).
99
How does the strength of v.d.w. force / intermolecular force related to molecular size?
The strength of van der Waals' force / intermolecular force ↑ with molecular size.
Example (Refer to notes p.20):
* No. of occupied e⁻ shells: Ar (3) > Ne (2) > He (1)
* Molecular size: Ar > Ne > He
* Strength of v.d.w. force: Ar > Ne > He
100
What are the properties of substances with simple molecular structures?
101
How do metal atoms form cations
(positive ions)?
The outermost shell (valence shell) e⁻ of metal atoms leave the atoms readily to form cations (positive ions).
102
What are delocalized electrons (or mobile electrons)?
Delocalized electrons are electrons that are **not taken away** by other species and they can **move freely** around metal ions.
* The delocalized / mobile e⁻ are shared by **all the atoms** in the metal.
103
Definition of metallic bond.
The strong electrostatic attraction between **cation and the 'sea' of delocalized e⁻** is called metallic bond.
* Metallic bond is **non-directional**.
104
Describe + Draw the structure of metals
(Giant metallic structure).
Structure: Refer to notes p.2 (part 2).
All metals have giant metallic structures.
In a giant metallic structure,
* The atoms are **closely packed** together in a **regular pattern** (to form a giant lattice of cations).
* The cations (positive metal ions) are surrounded by **a 'sea' of delocalized e⁻**.
* The strong electrostatic attractions between the cations and the 'sea' of delocalized e⁻ are the **metallic bonds**.
105
Why are metals electrical / heat conductor?
The delocalized e⁻ can move under **the influence of an electric field** → a metal is an electrical conductor.
The delocalized e⁻ can move upon **heating** → a metal is a heat conductor.
106
What are the requirements for a substance to conduct electricity?
Refer to notes p.70 (part 1)
Either one of the followings:
**1. Presence of delocalised / mobile e⁻**
* All metals [(s) and (l) = molten]
* Graphite, a form of carbon (s)
**2. Presence of mobile ions**
* Ionic compounds [(l) and (aq), not (s)]
* Acids and alkalis (aq)
107
How does the strength of metallic bond change across the period?
For metals of the same period, **the strength of metallic bond ↑ from left to right** ∵
**No. of outermost e⁻ ↑** and **atomic size ↓** from left to right.
Example:
* No. of outermost shell e⁻: Na (1) < Mg (2) < Al (3)
* Atomic size: Na > Mg > Al
* Strength of metallic bond: Na < Mg < Al
* Melting point and boiling point: Na < Mg < Al
Note: Melting and boiling of a metal requires breaking the metallic bonds. Breaking stronger bonds required more energy.
108
How does the strength of metallic bond change across the group?
For metals of the same group, **the strength of metallic bond ↓ down the group** ∵
As the **atomic size ↑** down the group (as a result of ↑ed no. of occupied e⁻ shells), the **attractions between the protons in the nucleus of cations and the delocalized e⁻ ↓** down the group.
Example:
* Atomic size (or no. of occupied e⁻ shells): K (4) > Na (3)
* Strength of metallic bond: K < Na
* Melting point and boiling point: K (63°C, 759°C) < Na (98°C, 883°C)
Note: Melting and boiling of a metal requires breaking the metallic bonds. Breaking stronger bonds required more energy.
109
What are the properties of metals?
In general, metals (except mercury: liquid, weak bond strength) have **high melting points and boiling points** as the **metallic bonds are strong** and much energy is needed to break them.
* **Group I metals** have **relatively low melting points and boiling points** as a result of weaker metallic bonds (they have only 1 outermost e⁻) compared to other metals.
* Metals usually have **higher m.p. and b.p. than substances with simple molecular structures** ∵
Metallic bond is stronger than intermolecular force / van der Waals' force and thus more energy is needed to break the metallic bonds.
Most metals have **high densities** ∵ of the strong metallic bonds so that the metal atoms are usually packed closely in the solid.
* Most metals have densities higher than water (exception: **Li, Na and K**).
Most metals are **hard** ∵ of the strong metallic bonds. Relative motion of the atom is restricted.
Metals are **ductile and malleable** ∵ the layers of atoms can slide over each other.
* When they do this, the metals do not break as the **non-directional metallic bonds** still hold the atoms together.
* Refer to notes p.6 (part 2)
Metals are **good conductor of heat and electricity** ∵ the delocalized e⁻ can move freely.
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