Understanding Chemical Bonding: Ionic, Covalent, and Metallic, Lecture notes of Reasoning

Download Understanding Chemical Bonding: Ionic, Covalent, and Metallic and more Lecture notes Reasoning in PDF only on Docsity!

Chapter 3.

Structures and Properties of

Substances

Chemical Bonding

The orbitals in the Periodic Table

The elements of the periodic table can be classified according to the type of orbital that is being filled.

• Elements that appear in the^ s^ block and the^ p^ block are called either the main group elements or the representative elements.

These elements are representative of a wide range of physical and chemical properties. Among

them there are highly reactive, moderately reactive and non reactive elements. While most are

solids at room temperature, roughly one quarter of them are gases, one is a liquid.

• Elements that appear in the d block are called the^ transition elements.
• the f block elements are called the^ inner transition elements

highly reactive elements, moderately reactive elements, and unreactive elements. While most main group elements are solids at room tempera- ture, roughly one quarter of them are gases, and one is a liquid. Elements that appear in the d block are called the transition elements. They mark the transition from the p orbital filling order to the d orbital filling order. By the same reasoning, the f block elements are called the inner transition elements , because they mark a transition from the d orbital filling order to the f orbital filling order.

3 s

2 s

3

4

5

6

7

2

1

1 (IA) 2 (IIA)

13 (IIIA)

14 (IVA)

15 (VA)

16 (VIA)

17 (VIIA)

18 (VIIIA)

3 (IIIB)

3 (IIIB)

4 (IVB)

5 (VB)

6 (VIB)

7 (VIIB)

8 9 10 (VIIIB)

11 (IB)

12 (IIB)

4 f

5 f

4 s

1 s

7 s

6 s

5 s

3 d

4 d

6 d

5 d

2 p

3 p

4 p

6 p

5 p

s block (main group elements)

f block (inner transition elements)

d block (transition elements)

p block (main elements)

The long form of the

periodic table, with the four energy

sublevel blocks identified.

Figure 3.

Chemical Bonding

Of the about 120 elements that occur in nature or that have been produced synthetically, only the noble

gases exist naturally as single, uncombined atoms.

In nature, systems of lower energy tend to be favored over systems of higher

energy. In other words, lower-energy systems tend to have greater stability than

higher-energy systems.

Bonded atoms, therefore, tend to have lower energy than single, uncombined

atoms

Defintion

Chemical bonds are electrostatic forces that hold atoms together in

compounds and involve the interaction of valence electrons.

Using Lewis Structures to Represent Atoms

To draw the Lewis structure of an atom: 1.replace its nucleus and inner electrons with its atomic symbol 2.add dots around the atomic symbol to symbolize the atom’s valence electrons (many chemists place the dots starting at the top and continue adding dots clockwise, at the right, then bottom, then left. then begin again at the top) Drawing a Lewis structure for a molecule lets you see exactly how many electrons are involved in each bond, and helps you to keep track of the number of valence electrons l bonding involves the interaction of valence electrons—t that occupy the outermost principal energy level of an at e used Lewis structures in previous studies to indicate th lectrons of atoms. Recall that to draw the Lewis structu lace its nucleus and inner electrons with its atomic sym dots around the atomic symbol to symbolize the atom’

. Many chemists place the dots starting at the top and c

ots clockwise, at the right, then bottom, then left. After e first four dots, you begin again at the top, as shown be is chapter, you will use Lewis structures often to represe s and the simplest formula unit of an ionic solid. Drawi ucture for a molecule lets you see exactly how many ele

Na

• •

S^ •

• •

P

• •
• •

Cl

• •

   - • - • 

Ar

Si

• • • •

Al

Mg

Ionic Bonding

Ionic bonding occurs between atoms of elements that have large differences in electronegativity

usually a metal (low electronegativity) and a non-metal (high electronegativity).

The units of ionic compounds such as sodium chloride and magnesium fluoride cannot be

separated easily by direct heating of the crystal salts.

The ions that make up the ionic solid are arranged in a specific array of repeating units.

In solid sodium chloride, for example, the ions are arranged in a rigid lattice

structure. In such systems, the cations and anions are arranged so that

the system has the minimum possible energy

Lattice structure of sodium chloride

non-metals. For example, magnesium in Group 2 (IIA) and fluorine in Group 17 (VIIA) combine to form the ionic compound magnesium fluoride, MgF 2. Figure 4.1 shows a repeating unit in the crystal model of magnesium fluoride. The process that results in the formation of ions can be illustrated with an orbital diagram or with Lewis structures, as shown in Figures 4.2 and 4.3. Use them as a guide for the Practice Problems below. Through bonding, the atoms of each element obtain a valence electron configuration like that of the nearest noble gas. In this case, the nearest noble gas for both ions is neon. This observation reflects the octet rule.

F

• •

• •

• • Mg

2 + −

1 s 2 s 2 p

F

1 s 2 s 2 p

F

1 s 2 s 2 p 3 s

Mg

1 s 2 s 2 p

F−

1 s 2 s 2 p

F−

1 s 2 s 2 p

Mg^2 +

F

• •

• •

• •

F

• •

• •

• •

Mg + •^ F

• •
• •
• •

−

1. Write electron configurations for the following:

Practice Problems

MgF 2

Figure 4.

for MgF 2

Figure 4.

Ionic Bonding

E.g. magnesium in Group 2 and fluorine in Group 17 combine to form the ionic compound magnesium fluoride (MgF 2 ). The figure shows a repeating unit in the crystal model of magnesium fluoride.

The process that results in the formation of ions can be illustrated with Lewis structures

Because of the large differences in electronegativity, the atoms in an ionic compound usually

come from the s block metals and the p block non-metals.

Mg F

Ionic Bonding

Practice problem

• Write electron configurations for the following elements:

a. Li+

b.Ca2+

c.Br−

d.O^2 −

• Draw Lewis structures for these chemical species
• Draw orbital diagrams (box) and Lewis structures to show how the following pairs of

elements can combine. In each case, write the chemical formula for the product.

e.Li and S

f.Ca and Cl

g.K and Cl

h.Na and N

Properties of Ionic Solids

In general, ionic solids have the following properties:

• crystalline with smooth, shiny surfaces
• hard but brittle
• non-conductors of electricity and heat
• high melting points
• many ionic solids are also soluble in water (MgF^2 is an exception)

The amount of energy given off when an ionic crystal forms from the gaseous ions of its elements is

called the lattice energy (e.g. The lattice energy of MgF 2 is 2957 kJ/mol). The same amount of energy

must be added to break the ionic crystal back into its gaseous ions.

Characteristics of Covalent Bonding

Generally, electron-sharing enables each atom in a covalent bond to acquire a

noble gas configuration.

The period 2 non-metals from carbon to fluorine must fill their 2 s and 2 p orbitals to acquire a noble gas configuration like that of Ne ( octet rule ).

E.g. In the formation of the diatomic fluorine molecule, F 2 , the bonding (shared) pair of electrons gives each fluorine atom a complete valence level. lone pairs, are not involved in bonding.

cquire a noble gas configuration. For a hydrogen molecule, each atom

cquires a filled valence level like that of helium by treating the shared

air of electrons as if it is part of its own composition. As you can see in

igure 4.5, a single shared pair of electrons—a bonding pair — fills the

alence level of both hydrogen atoms at the same time.

The period 2 non-metals from carbon to fluorine must fill their 2 s and

heir three 2 p orbitals to acquire a noble gas configuration like that of

eon. Covalent bonding that involves these elements obeys the octet rule.

n the formation of the diatomic fluorine molecule, F 2 , for example, the

onding (shared) pair of electrons gives each fluorine atom a complete

alence level.

ach fluorine atom also has three unshared pairs of electrons. These

airs of electrons, called lone pairs , are not involved in bonding.

The covalent bond that holds molecules of hydrogen, fluorine, and

ydrogen fluoride together is a single bond. It involves a single bonding

• • • •
• • F
• • F
• •
• •
• •
• • F
• • F
• •
• •

bonding pair lone pairs

or

Characteristics of Covalent Bonding

Some molecules are bonded together with two shared pairs of electrons. These are called double bonds. CO 2 is an example of a covalent molecule that consists of double bonds

Chapter 4 Stru

trons. Some molecules are bonded together with tw

ctrons. These are called double bonds. Carbon dioxi

a covalent molecule that consists of double bonds.

that are bonded with three shared pairs of electrons

s. Nitrogen, N 2 , another diatomic molecule, is a trip

• • N

N •^ or N

N^ •

C

• •

O

C •

• (^) or
• •
• •

O

• •
  • •

O

• •
• •

O

• •

Molecules that are bonded with three shared pairs of electrons have triple bonds. Nitrogen, N 2 , another diatomic molecule, is a triple-bonded molecule

Chapter 4 St

trons. These are called double bonds. Carbon dio

a covalent molecule that consists of double bonds

hat are bonded with three shared pairs of electron

. Nitrogen, N 2 , another diatomic molecule, is a tri - - - - - • N

N •^ or N

N^ •

C

• •

O

C •

• (^) or
• •
• •

O

• •
  • •

O

• •
• •

O

• •

Properties of Ionic Solids

In contrast to ionic solids, covalent compounds typically have the following

properties:

• exist as a soft solid, a liquid, or a gas at room temperature
• have low melting points and boiling points
• are poor conductors of electricity, even in solution
• may not be soluble in water

Diamond (C)^ Quartz (SiO 2 )

Predicting Covalent or Ionic Bonding

We can use the electronegativity difference between the bonding atoms to

predict the type of bond.

imply that covalent bonds are weaker than ionic bonds? Give evidence to justify your answer.

• −

mostly ionic (∆ EN > 1.7)

polar covalent (∆ EN 0.4 – 1.7)

mostly covalent (∆ EN < 0.4)

∆ EN

The relationship between bonding character and

Figure 4.

E.g.

• Two atoms with identical electronegativities, such as chlorine

(∆EN=3.16−3.16=0) share their electrons equally. They are

bonded covalently.

• In sodium chloride, chlorine (EN=3.16) attracts an electron

much more strongly than sodium (EN=0.93). Therefore, sodium’s

valence electron has a very high probability of being found near

chlorine. A high electronegativity difference is characteristic of

ionic compounds.

The relationship between bonding character and electronegativity difference

For atoms that have ∆EN between 0.4 and 1.7, the bond is polar

covalent. A polar covalent bond has an unequally shared pair of

electrons between two atoms.

This unequal sharing results in a bond that has partially positive

and partially negative poles.

Metallic Bonding

Based on electronegativity differences metals do not form ionic bonds with other metals. Similarly, metals do not have a sufficient number of valence electrons to form covalent bonds with one another. Metals do, however, share electrons. Unlike the electron sharing in covalent compounds, however, electron sharing in metals occurs throughout the entire structure of the metal.

M A

Co

Na

millions of atoms e−^ sea

Metals are composed of a densely packed core of metallic cations, within a delocalized region of shared, mobile valence electrons ( free-electron model ).

The force of attraction between the positively charged cations and the pool of valence electrons that moves among them constitutes a metallic bond.

Properties of the Metallic Bonding

The free-electron model explains many properties of metals:

• Conductivity : Metals are good conductors of electricity and heat because electrons can move freely throughout the metallic structure. This freedom of movement is not possible in solid ionic compounds, because the valence electrons are held within the individual ionic bonds in the lattice.
• Melting and Boiling Points: The melting and boiling points of Group 1 metals are generally lower than the melting and boiling points of Group 2 metals. Because the greater number of valence electrons and the larger positive charge of Group 2 atoms result in stronger metallic bonding forces