Grade: Grade 10 Subject: Science Unit: Chemical Bonding SAT: ProblemSolving+DataAnalysis ACT: Science

Ionic and Covalent Bonds

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Chemical Bonds

A chemical bond is a lasting attraction between atoms that enables the formation of chemical compounds. Bonds form because atoms seek to achieve a more stable electron configuration, typically a full outer shell (octet rule for most atoms).

The Two Main Types of Chemical Bonds

Ionic Bonds

Ionic Bond

An ionic bond is formed when one or more electrons are transferred from one atom to another, creating oppositely charged ions that are attracted to each other. This typically occurs between metals (which lose electrons) and nonmetals (which gain electrons).

Formation of Ionic Bonds:

  1. A metal atom loses one or more valence electrons to become a positively charged cation.
  2. A nonmetal atom gains those electrons to become a negatively charged anion.
  3. The opposite charges attract, holding the ions together in an ionic compound.

Example: Sodium Chloride (NaCl)

Sodium (Na) has 1 valence electron. Chlorine (Cl) has 7 valence electrons.
Na loses its valence electron: Na --> Na+ + e-
Cl gains that electron: Cl + e- --> Cl-
The Na+ and Cl- ions attract each other, forming NaCl.

Covalent Bonds

Covalent Bond

A covalent bond is formed when two atoms share one or more pairs of electrons. This typically occurs between two nonmetal atoms that have similar electronegativity values.

Types of Covalent Bonds:

Type Electrons Shared Representation Example
Single Bond 2 (1 pair) Single line (-) H-H in H2
Double Bond 4 (2 pairs) Double line (=) O=O in O2
Triple Bond 6 (3 pairs) Triple line (triple bond) N triple bond N in N2

Polar vs. Nonpolar Covalent Bonds

Polar Covalent Bond

A polar covalent bond occurs when electrons are shared unequally between atoms due to a difference in electronegativity. The more electronegative atom attracts the shared electrons more strongly, creating partial charges.

Nonpolar Covalent Bond

A nonpolar covalent bond occurs when electrons are shared equally between atoms with the same or very similar electronegativity values.

Electronegativity and Bond Type

Electronegativity Difference Bond Type Electron Behavior
0.0 - 0.4 Nonpolar Covalent Electrons shared equally
0.5 - 1.7 Polar Covalent Electrons shared unequally
1.8 or greater Ionic Electrons transferred

Comparing Ionic and Covalent Compounds

Property Ionic Compounds Covalent Compounds
Elements Involved Metal + Nonmetal Nonmetal + Nonmetal
State at Room Temp Usually solid Gas, liquid, or solid
Melting/Boiling Point High Generally low
Electrical Conductivity Conducts when molten or dissolved Generally does not conduct
Solubility in Water Often soluble Varies; many are insoluble
Hardness Hard but brittle Soft or waxy (molecular)

Lewis Dot Structures

Lewis dot structures show how valence electrons are arranged in molecules. Dots represent electrons, and lines represent shared pairs (bonds).

Steps to Draw Lewis Structures

  1. Count the total number of valence electrons from all atoms.
  2. Draw single bonds between atoms (2 electrons each).
  3. Distribute remaining electrons as lone pairs to complete octets.
  4. If octets aren't satisfied, form double or triple bonds.

Polyatomic Ions

Polyatomic ions are groups of covalently bonded atoms that carry an overall charge. They act as a single unit in ionic compounds.

Name Formula Charge
Hydroxide OH- 1-
Nitrate NO3- 1-
Sulfate SO4 2- 2-
Ammonium NH4+ 1+
Carbonate CO3 2- 2-
Phosphate PO4 3- 3-

SAT/ACT Connection

Science sections may present data about compound properties and ask you to identify the bond type. Remember: high melting points and electrical conductivity when dissolved typically indicate ionic compounds, while low melting points and poor conductivity suggest covalent (molecular) compounds.

💡 Examples

Example 1: Identifying Bond Type by Elements

Problem: Classify each compound as ionic or covalent: (a) MgCl2, (b) CO2, (c) KBr, (d) H2O

Solution:

(a) MgCl2: Mg is a metal (Group 2), Cl is a nonmetal (Group 17). Metal + nonmetal = Ionic

(b) CO2: C is a nonmetal (Group 14), O is a nonmetal (Group 16). Nonmetal + nonmetal = Covalent

(c) KBr: K is a metal (Group 1), Br is a nonmetal (Group 17). Metal + nonmetal = Ionic

(d) H2O: H is a nonmetal, O is a nonmetal. Nonmetal + nonmetal = Covalent

Example 2: Using Electronegativity Difference

Problem: Determine the bond type between hydrogen (EN = 2.1) and chlorine (EN = 3.0).

Solution:

Step 1: Calculate electronegativity difference: |3.0 - 2.1| = 0.9

Step 2: Compare to the scale:

  • 0.0 - 0.4: Nonpolar covalent
  • 0.5 - 1.7: Polar covalent
  • 1.8+: Ionic

Step 3: 0.9 falls in the 0.5 - 1.7 range.

Answer: The H-Cl bond is polar covalent. Chlorine is more electronegative, so it attracts the shared electrons more strongly.

Example 3: Drawing a Lewis Structure

Problem: Draw the Lewis structure for water (H2O).

Solution:

Step 1: Count valence electrons:

  • O has 6 valence electrons
  • Each H has 1 valence electron
  • Total: 6 + 1 + 1 = 8 valence electrons

Step 2: Draw the skeleton structure with O in the center, bonded to both H atoms.

Step 3: Each O-H bond uses 2 electrons (4 total for 2 bonds).

Step 4: Remaining electrons: 8 - 4 = 4 electrons (2 lone pairs on O).

Answer: O is bonded to two H atoms with single bonds, and O has two lone pairs of electrons. Each H atom shares its electron with O.

Example 4: Explaining Ionic Compound Properties

Problem: Explain why NaCl conducts electricity when dissolved in water but not as a solid.

Solution:

Step 1: In solid NaCl, Na+ and Cl- ions are locked in a rigid crystal lattice structure.

Step 2: Although the ions are charged, they cannot move freely, so they cannot carry electric current.

Step 3: When NaCl dissolves in water, the crystal lattice breaks apart.

Step 4: The Na+ and Cl- ions become free to move throughout the solution.

Answer: Dissolved NaCl conducts electricity because the dissociated ions can move freely and carry electric charge. In the solid state, ions are fixed in position and cannot move to conduct electricity.

Example 5: Writing Formulas with Polyatomic Ions

Problem: Write the formula for calcium phosphate.

Solution:

Step 1: Identify the ions:

  • Calcium ion: Ca2+ (Group 2 metal)
  • Phosphate ion: PO4 3-

Step 2: Balance the charges to get a neutral compound:

  • Ca2+ has a +2 charge
  • PO4 3- has a -3 charge
  • Need total positive charges to equal total negative charges

Step 3: Find the least common multiple: LCM of 2 and 3 is 6.

  • Need 3 Ca2+ ions: 3 x (+2) = +6
  • Need 2 PO4 3- ions: 2 x (-3) = -6

Answer: Ca3(PO4)2 - The parentheses around PO4 indicate 2 phosphate units.

✏️ Practice

1. Which compound is most likely to be ionic?

A) CH4

B) H2O

C) NaBr

D) CO2

2. What type of bond is formed when electrons are shared unequally between two atoms?

A) Ionic bond

B) Nonpolar covalent bond

C) Polar covalent bond

D) Metallic bond

3. If the electronegativity difference between two bonded atoms is 2.1, the bond is best classified as:

A) Nonpolar covalent

B) Polar covalent

C) Ionic

D) Metallic

4. Which property is typical of ionic compounds?

A) Low melting points

B) Poor conductivity when dissolved

C) High melting points

D) Gaseous at room temperature

5. In a molecule of ammonia (NH3), how many total valence electrons are present?

A) 5

B) 6

C) 7

D) 8

6. Which bond involves the transfer of electrons?

A) Single covalent

B) Double covalent

C) Ionic

D) Polar covalent

7. A triple bond consists of how many shared electrons?

A) 2

B) 4

C) 6

D) 8

8. The formula for magnesium oxide is MgO. What are the charges of the ions in this compound?

A) Mg+ and O-

B) Mg2+ and O2-

C) Mg3+ and O3-

D) Mg+ and O2-

9. Which pair of elements would form a nonpolar covalent bond?

A) Na and Cl

B) H and H

C) H and O

D) Ca and F

10. What is the correct formula for aluminum sulfate?

A) AlSO4

B) Al2SO4

C) Al2(SO4)3

D) Al3(SO4)2

Click to reveal answers
  1. C - NaBr is formed between sodium (metal) and bromine (nonmetal), making it ionic.
  2. C - Polar covalent bonds involve unequal sharing due to electronegativity differences.
  3. C - An electronegativity difference of 2.1 exceeds 1.8, indicating an ionic bond.
  4. C - Ionic compounds have high melting points due to strong electrostatic attractions between ions.
  5. D - Nitrogen has 5 valence electrons, each hydrogen has 1: 5 + 3(1) = 8 total.
  6. C - Ionic bonds form when electrons are transferred from a metal to a nonmetal.
  7. C - A triple bond shares 3 pairs of electrons, which equals 6 electrons.
  8. B - Mg loses 2 electrons to become Mg2+, O gains 2 electrons to become O2-.
  9. B - Two identical atoms (H-H) share electrons equally, forming a nonpolar covalent bond.
  10. C - Al3+ and SO4 2-: need 2(3+) and 3(2-) to balance charges, giving Al2(SO4)3.

✅ Check Your Understanding

Question 1: Why do metals typically form cations while nonmetals typically form anions in ionic compounds?

Reveal Answer

Metals have low ionization energies and few valence electrons, making it energetically favorable for them to lose electrons and achieve a stable noble gas configuration. This creates positively charged cations. Nonmetals have high electronegativity and almost-full valence shells, making it favorable for them to gain electrons to complete their octet. This creates negatively charged anions. The combination of metal cations and nonmetal anions creates ionic compounds.

Question 2: Explain why covalent compounds generally have lower melting points than ionic compounds.

Reveal Answer

Ionic compounds consist of ions held together by strong electrostatic attractions throughout a crystal lattice. Breaking these bonds requires substantial energy, resulting in high melting points. Covalent compounds are often discrete molecules held together by weak intermolecular forces (like van der Waals or hydrogen bonding) between molecules. While the covalent bonds within molecules are strong, the forces between molecules are weak and require less energy to overcome. Therefore, covalent compounds typically melt at lower temperatures.

Question 3: Carbon dioxide (CO2) and silicon dioxide (SiO2) both contain double bonds between the central atom and oxygen. However, CO2 is a gas at room temperature while SiO2 (quartz) is a very hard solid. Explain this difference.

Reveal Answer

CO2 forms discrete molecular units (O=C=O) held together by weak intermolecular forces, making it a gas at room temperature. SiO2, however, forms a covalent network solid where each silicon atom is covalently bonded to four oxygen atoms in a continuous three-dimensional network. There are no individual molecules - the entire crystal is one giant molecule. Breaking this extended network requires breaking many strong covalent bonds, giving SiO2 a very high melting point (about 1700C) and making it extremely hard.

Question 4: A compound has a melting point of 801C, conducts electricity when dissolved in water, and forms crystals. Based on these properties, what type of bonding does this compound have? What additional evidence would help confirm your answer?

Reveal Answer

The compound is most likely ionic based on the evidence: (1) High melting point (801C) indicates strong bonds throughout a crystal structure. (2) Conducts electricity when dissolved means free-moving ions are present in solution. (3) Crystal formation is typical of ionic compounds with their regular lattice arrangements. Additional confirming evidence could include: testing if it conducts electricity when molten (ionic compounds would), checking if it dissolves in polar solvents like water (many ionic compounds do), or determining the elements involved (a metal-nonmetal combination would confirm ionic bonding).

🚀 Next Steps

  • Review any concepts that felt challenging
  • Move on to the next lesson when ready
  • Return to practice problems periodically for review