Topic 3 Chemical Bonding

By means of Lewis structure, shows the chemical bonding in (a) SO3 (b) CuCl42-

Classify the following compounds as ionic compound or covalent compound and state the type of chemical bond present in each of these compounds. (a) phospane, PH3 (b) lithium aluminium hydride, LiAlH4 (c) Ammonium chloride.

Show the formation of coordinate bonding in the following reactions and identify the donor and acceptor atoms.
(a) NH3 + AlCl3 –> NH3•AlCl3
(b) Fe3+ + 6H2O –> [Fe(H2O)6]3+
(c) AlCl3 + Cl- –> AlCl4-
(d) O + O2 –> O3

What is the dative bond/ coordinate bond?

What is ligand?

What is hybridisation of atomic orbitals?

What is the types of hybrid atomic orbitals?

What is sigma bond?

What is pi bond?

For each types of hybrid atomic orbitals, give the structure, bond angle, sketch the formation of the structure and number of sigma and pi bonds?

Draw the Lewis structures of (i) oxygen molecule and (ii) nitrogen molecule?

Explain the covalent bond in oxygen molecule, using the concept of overlapping atomic orbitals?

There are three electron groups in the Lewis structure for the oxygen molecule, O2. Hence, the oxygen atom uses sp2 hybrid orbitals for bond formation. The double bond in the O2 molecule consists of a sigma bond and pi bond. The sigma bond is formed by the head-on overlapping of sp2 hybrid orbitals, while the a bond is formed by the sideways overlapping of the p orbitals.

Draw the Lewis structures of (i) BeH, and (ii) BF3?

Explain why

(i) BeH, has a linear shape and

BeH2, has linear shape because Be uses sp hybrid orbitals for bond formation. The electronic configuration of beryllium is Be: 1s2 2s2. The 1s orbitals of hydrogen atoms then overlap with the sp hybrid orbitals of beryllium atoms to form two Be-H bonds.

(ii) BF, molecule has a trigonal planar shape.

BF3 has a trigonal planar shape because boron atom uses sp2 hybrid orbitals for bond formation. The electronic configuration of boron is B: 1s2 2s2 2p1. The 2pz orbitals of fluorine atoms overlap with the three sp2 hybrid orbitals of boron atom to form three B-F bonds.

In terms of hybrid orbitals, describe the chemical bonds in tetrachloromethane and explain the shape of the molecule?

The carbon atom in CCl4 uses sp3 hybrid orbitals to overlap with the p orbitals of chlorine atoms to form C-Cl single bonds. Hence, the CCl4 molecule has a tetrahedral shape.

What is VSEPR Theory?

Valence Shell Electron-Pair Repulsion (VSEPR) Theory There are eleven molecular shapes. The shapes of covalent molecules can be explained by using (a) the concept of overlapping and hybridisation of orbitals (b) the repulsion of the valence shell electron-pair.

If the Lewis structure (or dot-and-cross structure) can be written for a molecule or a polyatomic ion, the shape of this molecule or ion can be predicted using the electron-pair repulsion theory. The valence shell electron-pair repulsion theory (VSEPR theory) states that (a) the electron-pair around the central atom repel each other (b) the electron-pairs (bonding pairs and lone pairs) arrange themselves to be as far apart as possible to minimise the force of repulsion, and (c) the force of repulsion decreases in the order: lone pair-lone pair repulsion (strongest repulsion) > lone pair-bond pair repulsion (medium repulsion) > bond pair-bond pair repulsion (weakest repulsion).

It is important to distinguish between the geometry (orientation of electrons and the shape of molecules. For molecules and ions that do not contain lone pairs of electrons, the shape of the molecules or ions is the same as the geometry of electron-pairs. For molecules and ions that contain lone pairs of electrons, the shape of the molecules or ions is different from the geometry of electron-pairs.

Use electron-pair repulsion theory to draw the Lewis structures and predict their shapes of the following molecules and ions?

BeCl2, BF3, CO32-, SO4 2-, SO3 2-, NO3-, ClO4-, ClO2-, COCl2, I3-, F2O, AsH3, BH4-, H3O+, H2O, CH4, CH3+, CH3-, ICl2-, ICl4-, NH2-, NH3, NH4+, NH3 vs PH3, NH3 vs NF3, CO2, SO2, HCN, POCl3, SF6, POCl3, SOCl2, H3O+, ClF2+, ICl2-, PH4+, PCl6-, ClF3, SeCl3.

What is polar covalent bonds?

What is non polar covalent bonds?

How is the polarisation happens, where and when?

What is dipole?

What is dipole moment?

What is/ are the factors influencing the formation of ionic compounds? What makes an ionic compound stable?

element with low ionisation energy combines with an element with higher electron affinity. ionic bonding is the resukt of strong electrostatic force of attraction between positive and negative ions. hence the strength of an ionic bond is (i) proportional to the charge on the ions and (ii) inversely proportional to the distance between the ions. The higher the charge on tge ions and/ or the smaller the ionic radius, the stronger is the attractive force between the ions, hence the stronger the ionic bond.

What is the definition of ionisation energy?

minimum amount of energy required to remove the most loosely bound electron, the valence electron, of an isolated neutral gaseous atom or molecule. it is a measure of the difficulty in removing an electron from an atom or ion or the tendency of an atom or ion to surrender an electron.

What is the definition of electron affinity?

the amount of energy released when an electron is attached to a neutral atom or molecule in the gaseous state to form a negative ion.

Draw Lewis structure of the following ionic compounds? KF, BaO, Na2O, MgBr2.

The element X has one electron and the element Y has six electrons in their outermost shell respectively. (a) use the lewis diagram to show the transfer of electrons when X combines with Y. (b) what is the formula of the compound formed between X and Y?

 

What makes a covalent bond strong?

covalent bond is measured by bond dissociation energy. bond dissociation energy (or bond energy) is the energy required to break one mole of a covalent bond in a gaseous molecule. the stronger the covalent bond, the greater the bond energy.

Make comparisons of ionic and covalent compounds’ properties for physical states, bond strength, melting/ boiling point, electrical conductivity, solubility?

 

Classify the following compounds as ionic compound or covalent compound and state the type of chemical bond present in each of these compounds. (a) beryllium chloride (b) barium chloride (c) iodine trichloride (d) calcium oxide.

X, Y and Z represent the elements with proton numbers 9, 16, and 19 resepctively.
(a) Write the electronic configurations of X, Y and Z in terms of s and p orbitals. (b) Predict the type of bonding you would expect to form between (i) X and Y (ii) X and Z (iii) Y and Z (c) Write the dot-and-cross diagrams for the compounds formed.

What is coordinate / dative bond?

Draw Lewis structure of ammonium ion hydronium ion, Al2Cl6 and [ Fe(CN)6]3- ion, use arrow to indicate the dative bond?

 

3.1 Ionic bonding

What are the 3 types of chemical bonds?

1. Metallic bond.

2. Ionic bond.

3. Covalent bond.

What are the 4 types of chemical bonds?

There are four different types of chemical bonds:

1. polar covalent,

2. nonpolar covalent,

3. ionic, and

4. hydrogen bonds.

Why is chemical bonding important in our life?

Three types of chemical bonds are important in human physiology, because they hold together substances that are used by the body for critical aspects of homeostasis, signaling, and energy production, to name just a few important processes. These are ionic bonds, covalent bonds, and hydrogen bonds.

How to determine ionic bond’s strength?

We measure the strength of a ionic bond by the energy required to break it. Strength of lonic Bonding Is Measured Experimentally by “Lattice Energy” Lattice Energy is the energy change associated with the formation of one mole of a crystalline ionic solid from its gaseous ions.

For example: Na+ (g) + CI- (g) → NaCl (s); dH = lattice energy for Nacl (dHlattice)

The strength of the ionic bond is directly dependent upon the quantity of the charges and inversely dependent on the distance between the charged particles. A cation with a 2+ charge will make a stronger ionic bond than a cation with a 1+ charge.

Is lattice energy positive or negative?

Negative. Energy is released (exothermic) when things that attract get to come closer together.

F2O, v-shaped or bent shape.

AsH3, trigonal pyrimidal.

NH4+, tetrahedral.

BH4-, tetrahedral.

H3O+, trigonal pyramidal.

CH3+, trigonal planar.

CH3-, trigonal pyramidal.

What is the lattice energy of following ionic compounds?

NaBr : −747 kJ / mol

LiH -906

LiF -1009

LiCl -829

LiBr -789

LiI -734

RbCl -687

CuCl -979

CuBr -976

CuI -958

MgF2 -2908

MgBr2 -2406

MgCl2 : -2477 kJ / mol

MgF2 : −2922

BeCl2 :  kJ / mol

Among LiCl, RbCl, BeCl2 ​ and MgCl2 ​ which two are the highest and the lowest ionic compounds?

Exothermic lattice energy resulting in greater stability in the ionic compound. Lattice energies affect the solubilities of ionic substances in water. In general, the higher the lattice energy, the less soluble a compound is in water. LE, Solubility in water.

RbCl -687, 0.91 g/mL

LiCl -829, 0.84 g/mL

BeCl2 covalent molecule, 0.15 g/mL

MgCl2 : -2477, 0.054 g/mL

Rb has the largest ionic radius with small ion charge which is +1 and thus RbCl has the smallest lattice energy.

MgCl2 has the largest ion charge which is +2 thus it is the highest ionic strength. BeCl2 ​has a covalent bonding with simple covalent molecule (not ionic compound) because of very small and too much energy is required to remove two electrons from its valence 2s subshell.

Which has more lattice energy NaCl or MgCl2?

MgCl2 = 2477; NaCl = 769 kJ mol-1. Mg has higher ion charge +2 than Na +1. Higher charge makes ionic bond stronger. Higher lattice energy implies better stability meaning stronger bonds.

FeF2 -2912

CuF2 -3042

CoF2 2962

NaH -811

NaF -904

NaCl -769

NaBr -736

NaI -688

CsI :  kJ / mol

CsCl −657

CsBr −632

AgF -969

AgCl -916

AgBr -900

AgI -895

CaF2 -2611

CaI2 -492

CaO -3464

CaS -3093

CrF2 -2879

NiF2 -3046

KH -714

KF -801

KCl -698

KBr -672

KI -632

KI3 :  kJ / mol

AuCl -1042

AuI -1050

TlCl -748

BaF2 -2368

TiF2 -2749

ZnF2 -2971

ZnS -3619

CdS -3402

HgS -3573

Al2O3 -15,326

AlF3 : +5215 kJ / mol

AlCl3 : covalent compound

BF3 :  kJ / mol

NH4Br :  kJ / mol

MgO : -3795.0 kJ / mol

CaO : -3414 kJ / mol

BaO : -3054 kJ / mol

What do you think about determine the magnitude of this energy (lattice energy)?

Force of attraction between the ions. When larger negative lattice energy => stronger force of ionic bond.

Determine the strength of ionic bonds by using formula?

The strength of an ionic bond is determined by the charges of the ions and the distance between them. The larger the charges and the smaller the ions the stronger the bonds will be Bond strength then is proportional to (Q1 x Q 2) / (r^2). Where Q1 and Q2 represent ion charges and r is the sum of the ionic radii. (do not need to know formula).

How Coulomb’s law determine lattice energy of ionic bond strength?

Is AlCl3 classified as ionic or covalent?

Alcl3 is a covalent compound and not ionic compound. Because the Al atom has higher polarizing power which attracts the cl electron easily and causes electron sharing.

Why is AlCl3 in covalent and AlF3 in ionic?

In between AlCl3 and AlF3, cation (Al) is the same both the compound. Hence AlCl3 becomes covalent. On the other hand, due to the the smaller size of fluorine, aluminium does not polarised fluoride ion to such a great extent. So AlF3 becomes ionic.

What is AlF3 lattice energy?

Lattice Energy for AlF3 (s): +5215 kJ/mol.

What is MgO lattice energy?

Lattice Energy for MgO. −3795 kJ/mol.

Which has lowest lattice energy LiF, NaCl, KBr, CsI?

cesium iodide Lattice energy decreases as the size of ions increases. Compared to other compounds, the size of cation and anion is highest in cesium iodide. Hence the lattice energy is lowest for cesium iodide.

Which has highest lattice energy in alkali metal halide?

The lattice energies for the alkali metal halides is therefore largest for LiF and smallest for CsI, as shown in the table below. The ionic bond should also become stronger as the charge on the ions becomes larger.

Which of the following has the highest lattice energy NaCl, KI, MgO, BaO, CaO?

the answer: MgO

NaCl: 787.3 kJ/mol

KI: 632.0 kJ/mol

MgO: −3795 kJ/mol

BaO:  kJ/mol

CaO: −3414.0 kJ/mol

Which has lowest lattice energy MgO, Na2O, NaF, MgCl2, CaO ?

MgO

3.2 Covalent bonding

What is lewis structure?

Lewis structures, also called electron-dot structures or electron-dot diagrams, are diagrams that show the bonding between atoms of a molecule, and the lone pairs of electrons that may exist in the molecule. A Lewis structure can be drawn for any covalently-bonded molecule, as well as coordination compounds (a central metal atom surrounded by nonmetal atoms or groups of atoms, called ligands).

Lewis Diagrams for ionic bond

Lewis Diagrams for Covalent Bond

Lewis Diagrams for Multiple Bonds

Draw the Lewis structure of covalent molecules octet rule as exemplified by

NH3, CCl4, H2O, CO2, N2O4, N2H2, Cl2, O2, N2, C2H2, CS2,

NH4+, H3O+, CO3²-, SO4²-, HCO3-, C2O4²-, PO4³-, ClO2-, CrO4²-, CN-, NO3-,

and

exception to the octet rule as exemplified by BeCl2, BF3, NO, NO2, PCl5, SF6.

What is octet rule?

The octet rule refers to the tendency of atoms to prefer to have eight electrons in the valence shell. When atoms have fewer than eight electrons, they tend to react and form more stable compounds. When discussing the octet rule, we do not consider d or f electrons. Only the s and p electrons are involved in the octet rule, making it useful for the main group elements (elements not in the transition metal or inner-transition metal blocks); an octet in these atoms corresponds to an electron configurations ending with s2p6 .

What is bond angle in chemistry?

A bond angle is the angle between two bonds originating from the same atom in a covalent species. Geometrically, a bond angle is an angle between two converging lines.

What is principle of valence shell electron pair repulsion theory?

The valence-shell electron-pair repulsion (VSEPR) theory states that electron pairs repel each other whether or not they are in bond pairs or in lone pairs. Thus, electron pairs will spread themselves as far from each other as possible to minimize repulsion.

The valence electron pairs surrounding an atom tend to repel each other and will, therefore, adopt an arrangement that minimizes this repulsion. This in turn decreases the molecule’s energy and increases its stability, which determines the molecular geometry.

What are the main features of Vsepr theory?

1) The VSEPR model is used to predict the geometry of molecules.

2) The electron pairs around an atom are assumed to arrange themselves to reduce electron repulsion.

3) The molecular geometry is determined by the position of the bonding electron pairs .

Valence shell electron-pair repulsion (VSEPR) theory is a covalent bond theory that uses the repulsive forces between single electrons and pairs of electrons about the central atom to predict their relative positions around the atomic nuclei.

According to VSEPR theory, electron pairs around a central atom repel each other. Because of the repulsion between electrons, the shape with minimum energy is the one in which the distance between electron pairs is maximized. Therefore, the shape with the lowest energy is likely to be the actual form of the molecule. Electron pairs can either be inside a bond between atoms or exist as a lone pair. There are generally between two and six electron pairs (either in the bonds of bonding atoms or as lone pairs) around any central atom, and the electron pairs can be arranged in either of five ways, each with a different geometry. Electron-pair geometry is the shape description for all electron pairs (bonding and nonbonding) about a central atom. Electron-pair geometry maximizes the distance between every pair of electrons around a central atom. For example, if there are three pairs of electrons, increasing the angle between two electron pairs in a trigonal planar arrangement would push one of the pairs closer to the third pair, which would increase the potential energy of the third pair. When the distance between all of the electron pairs is maximized, the potential energy of the system is at a minimum. Systems always favor a position of minimum energy, so arrangements that maximize the distance between atoms and/or lone pairs are favored. Generally, large angles represent a greater degree of separation and are more favorable. Angles of 90° are least favorable.

VSEPR theory predicts only one possible shape, linear, for molecules with two bonding electron pairs. It predicts two possible shapes for molecules with three electron pairs: bent and trigonal planar. The theory predicts three possible shapes for molecules with four electron pairs: tetrahedral, trigonal pyramidal, and bent. VSEPR theory predicts four possible shapes for molecules with five electron pairs: trigonal bipyramidal, sawhorse, T-shape, and linear. For molecules with six electron pairs, there are three possible shapes: octahedral, square pyramidal, and square planar.

3.3 Metallic bonding

3.4 Intermolecular forces: van der Waals forces and hydrogen bonding

What is van der Waals force of attraction?

Van der Waals forces include attraction and repulsions between atoms, molecules, and surfaces, as well as other intermolecular forces. They differ from covalent and ionic bonding in that they are caused by correlations in the fluctuating polarizations of nearby particles (a consequence of quantum dynamics).

What is the effect in a liquid which has a strong van der Waals forces?

The liquid has high enthalpy of vapourisation.

3 Chemical Bonding

3.1 Ionic bonding

Candidates should be able to:

(a) describe ionic (electrovalent) bonding as exemplified by NaCl and MgCl2. other ionic compounds: NaCl, KF, BaO, Na2O, MgBr2, MgCl2, Al2O3, BaCl2, CaO

3.2 Covalent bonding

Candidates should be able to:

(a) draw the Lewis structure of covalent molecules (octet rule as exemplified by NH3, CCl4, H2O, CO2, N2O4 and exception to the octet rule as exemplified by BF3, NO, NO2, PCl5, SF6);
(b) draw the Lewis structure of ions as exemplified by SO42-, CO32-, NO3- and CN-;
(c) explain the concept of overlapping and hybridisation of the s and p orbitals as exemplified by BeCl2, BF3, CH4, N2, HCN (or CN-), NH3 and H2O molecules;
(d) predict and explain the shapes of and bond angles in molecules and ions using the principle of valence shell electron pair repulsion, e.g. linear, trigonal planar, tetrahedral, trigonal bipyramid, octahedral, V-shaped, T-shaped, seesaw and pyramidal;
(e) explain the existence of polar and non-polar bonds (including CCl, CN, CO, CMg) resulting in polar or/ and non-polar molecules;
(f) relate bond lengths and bond strengths with respect to single, double and triple bonds;
(g) explain the inertness of nitrogen molecule in terms of its strong triple bond and non-polarity;
(h) describe typical properties associated with ionic and covalent bonding in terms of bond strength, melting point and electrical conductivity;
(i) explain the existence of covalent character in ionic compounds such as Al2O3, Ali3 and LiI;
(j) explain the existence of coordinate (dative covalent) bonding as exemplified by H3O+, NH4+, Al2Cl6 and [Fe(CN)6]3 (or CN-).

3.3 Metallic bonding

Candidates should be able to:

(a) explain metallic bonding in terms of electron sea model.

3.4 Intermolecular forces: van der Waals forces and hydrogen bonding

Candidates should be able to:

(a) describe hydrogen bonding and van der Waals forces (permanent, temporary and induced dipole);
(b) deduce the effect of van der Waals forces between molecules on the physical properties of substances;
(c) deduce the effect of hydrogen bonding (intermolecular and intramolecular) on the physical properties of substances.

3 Chemical Bonding

Puppies demonstrating covalent bonding (A) The two puppies represent atoms, their bones represent one of their electrons. (B) Both puppies share both bones. Image by Byron Inouye

In a reaction, there is a change in chemical bonding.
Some of the bonds in the reactants are broken, and new bonds are made to form the products.
It costs energy to break bonds, but energy is released when new bonds are made.

covalent bond, electronegativity, polar covalent

covalent compound / molecule
H2
Cs2
Cs2
Br2
NH3
H2O
CCl4
Cl2
HF
SiH4
I2
SiO2
N2H2

3.1 Ionic bonding

Candidates should be able to:

(a) describe ionic (electrovalent) bonding as exemplified by NaCl and MgCl2. other ionic compounds: NaCl, KF, BaO, Na2O, MgBr2, MgCl2, Al2O3, BaCl2, CaO

ionic compounds/  molecules

KF
BaO
Na2O
MgBr2
Al2O3
BaCl2
CaO

3.2 Covalent bonding

Candidates should be able to:

(a) draw the Lewis structure of covalent molecules (octet rule as exemplified by NH3, CCl4, H2O, CO2, N2O4 and exception to the octet rule as exemplified by BF3, NO, NO2, PCl5, SF6);

NH3,

 

CCl4,

H2O,

CO2,

N2O4

exception

BF3,

 

Exception:

(b) draw the Lewis structure of ions as exemplified by SO42-, CO32-, NO3- and CN-;

incomplete: BeCl2, BeF2
expanded: SF6, SF4 , PCl5,
odd: NO , NO2
SO4 2-,
CO3 2-,
NO3-,
CN-;

incomplete: BeCl2

BeCl2. 1 Be × 2e- = 2e- 2 Cl × 7e- = 14e- 16e- Cl Be Cl. – 4e- 12e-

BeF2

 

expanded:

SF6,

 

SF4 ,

PCl5,

 

 

odd: NO , NO2

NO,

The valence electrons you have available are: 1 N + 1 O = 1×5 + 1×6 = 11.
With an odd number of electrons (11), we cannot give every atom an octet. We can write two possible structures. The formal charge on each atom is:

Top structure: N = 5 – 3 – ½(4) = 0; O = 6 – 4 – ½(4) = 0
Bottom structure: N = 5 – 4 – ½(4) = -1; O = 6 – 3 + ½(4) = +1

The “best” Lewis structure is one that has the fewest formal charges — the top structure.

NO2,

 

(c) explain the concept of overlapping and hybridisation of the s and p orbitals as exemplified by BeCl2, BF3, CH4, N2, HCN (or CN-), NH3 and H2O molecules;

BeCl2

BF3

CH4

N2

HCN (or CN-)

NH3

H2O

(d) predict and explain the shapes of and bond angles in molecules and ions using the principle of valence shell electron pair repulsion, e.g. linear, trigonal planar, tetrahedral, trigonal bipyramid, octahedral, V-shaped, T-shaped, seesaw and pyramidal;

e.g. linear, trigonal planar, tetrahedral, trigonal bipyramid, octahedral,
V-shaped, T-shaped, seesaw and pyramidal;
linear,
trigonal planar,
tetrahedral,
trigonal bipyramid,
octahedral,
V-shaped,
T-shaped,
seesaw,
pyramidal;

(e) explain the existence of polar and non-polar bonds (including CCl, CN, CO, CMg) resulting in polar or/ and non-polar molecules;

http://www.softschools.com/quizzes/chemistry/polarity_electronegativity/quiz941.html

CCl ,
CO,
CMg

(f) relate bond lengths and bond strengths with respect to single, double and triple bonds;

(g) explain the inertness of nitrogen molecule in terms of its strong triple bond and non-polarity;

(h) describe typical properties associated with ionic and covalent bonding in terms of bond strength, melting point and electrical conductivity;

(i) explain the existence of covalent character in ionic compounds such as Al2O3, Ali3 and LiI;

(j) explain the existence of coordinate (dative covalent) bonding as exemplified by H3O+, NH4+, Al2Cl6 and [Fe(CN)6]3 (or CN-).

 

H3O+,
NH4+,
Al2Cl6
[Fe(CN)6]3 (or CN-).

3.3 Metallic bonding

Candidates should be able to:

(a) explain metallic bonding in terms of electron sea model.

 

3.4 Intermolecular forces: van der Waals forces and hydrogen bonding

Candidates should be able to:

(a) describe hydrogen bonding and van der Waals forces (permanent, temporary and induced dipole);

two types of van der Waals forces:
1. permanent dipole forces (dipole-dipole), (polar molecules).
2. temporary dipole-induced dipole forces (London dispersion), (temporary polar molecules).
3. permanent induced dipole-induced dipole (London dispersion) (non-polar molecules)

1. permanent dipole forces (dipole-dipole), (polar molecules).
2. temporary dipole-induced dipole forces (London dispersion), (non-polar molecules).

3. permanent induced dipole-induced dipole (London dispersion) (non-polar molecules)

Summary of permanent, temporary and induced dipole

(b) deduce the effect of van der Waals forces between molecules on the physical properties of substances;

(c) deduce the effect of hydrogen bonding (intermolecular and intramolecular) on the physical properties of substances.
electronegativity
Other links for online notes
http://www.softschools.com/quizzes/chemistry/polarity_electronegativity/quiz941.html

ionic compound / molecule

KF
BaO
Na2O
MgBr2
Al2O3
BaCl2
CaO

covalent compound / molecule

H2
Br2
NH3
H2O
CCl4
Cl2
HF
SiH4
I2
SiO2
N2H2

multiple covalent bond

double bond: O2 , CO2 , N2O4 ,
triple bond: N2 , C2H2 , HCN (or CN-)

ions covalent bond

ClO2-
CN-
NO2-
CO3 2-
SO4 2-
NO3-
H3O+
OH-

dative / coordinate bond

NH4+
Al2Cl6
[Fe(CN)6]3- (or CN-)
CO

exception to the octet rule

1. incomplete octet

2. Expanded octet

3. Odd electron molecules

incomplete octet

BeCl2
BF3

Expanded octet

PCl5
SF6

Odd electron molecules

NO
NO2

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